Electrochemical Analysis of New Materials for Proton-Exchange Membrane Fuel Cells: Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage and “Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction By Leah Katherine Rubin Shen A dissertation submitted in partial satisfaction of the requirements for the degree of Doctor of Philosophy in Chemistry in the Graduate Division of the University of California, Berkeley Committee in charge: Professor John Arnold, Chair Professor Angelica M. Stacy Professor Bryan D. McCloskey Summer 2015
Transcript
Electrochemical Analysis of New Materials for Proton-Exchange Membrane Fuel Cells: Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage and
“Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction
By
Leah Katherine Rubin Shen
A dissertation submitted in partial satisfaction of the
requirements for the degree of
Doctor of Philosophy
in
Chemistry
in the
Graduate Division
of the
University of California, Berkeley
Committee in charge:
Professor John Arnold, Chair Professor Angelica M. Stacy
Professor Bryan D. McCloskey
Summer 2015
Electrochemical Analysis of New Materials for Proton-Exchange Membrane Fuel Cells: Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage and
“Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction
Electrochemical Analysis of New Materials for Proton-Exchange Membrane Fuel Cells: Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage and
“Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction
by
Leah Katherine Rubin Shen
Doctor of Philosophy in Chemistry
University of California, Berkeley
Professor John Arnold, Chair
Chapter 1: An Overview of Proton-Exchange Membrane Fuel Cell Research Needs
Proton-exchange membrane (PEM) fuel cells are a promising technology for sustainable energy, having the potential to replace vehicle combustion engines and enable grid-level storage of intermittent renewable energy sources. However, PEM fuel cells are currently fueled by hydrogen, which is difficult to store, transport, and distribute efficiently and safely. They also require precious metal catalysts for both hydrogen oxidation at the anode and oxygen reduction at the cathode. In particular, the platinum cathode catalyst reduces oxygen at a large overpotential, leading to lower overall cell voltage due to kinetic losses. Widespread adoption of PEM fuel cells is therefore limited by both the difficulties of hydrogen storage and distribution and the expense of the catalysts used.
To mitigate these difficulties, new materials for hydrogen storage and oxygen reduction catalysis are needed. This first part of this work explores the use of nitrogen heterocycles for “virtual” hydrogen storage. A virtual hydrogen storage system uses a liquid organic carrier that is electrochemically dehydrogenated to produce protons and electrons separately, allowing the compound to behave like hydrogen in the fuel cell environment. The second part of this work studies small molecule oxygen reduction catalysts using first-row transition metals. These catalysts could be an inexpensive alternative to platinum, but must be anchored within the fuel cell in order to work optimally. The catalysts studied herein were thus additionally designed to be covalently attached to an electrode surface.
Chapter 2: Electro-Dehydrogenation of Indoline: Direct Oxidation at the Electrode and Redox Catalysis
Indoline was chosen as a model for studying virtual hydrogen storage feasibility. The direct electrode reactivity of indoline favors a dimerization pathway. However, in the presence of excess base, either an electrochemical-chemical (EC) or chemical-electrochemical (CE) pathway is favored, depending on the strength of the base. The use of redox mediators in conjunction with base lowers the reaction voltage by 250 mV (ferrocene) or 750 mV (decamethylferrocene). Bulk electrolysis studies with ferrocene and imidazole demonstrate 25-35% indole formation, with no
2
dependence on catalyst concentration. However, bulk electrolysis in the presence of decamethylferrocene and 1,1,3,3-tetramethylguanidine produces up to 50% dehydrogenation to indole at 0.13 V vs. NHE, with more indole formation at higher catalyst concentrations.
Chapter 3: Electrochemical Analysis of a Fully Hydrogenated Carbazole as a Virtual Hydrogen Storage Compound
An electrochemical study of N-ethyldodecahydrocarbazole (NEC-H12), a compound that stores 5.7% hydrogen by weight, is presented. This first study of a fully hydrogenated nitrogen heterocycle indicates that the direct electrode oxidation of NEC-H12 involves a slow electron transfer step with a rate-competitive follow-up reaction. Addition of a strong base, 1,1,3,3-tetramethylguanidine, increases the oxidative current up to five-fold. Effective redox catalysis with ferrocene is also demonstrated, lowering the voltage for NEC-H12 oxidation by nearly 0.4 V. Finally, while bulk electrolysis studies of NEC-H12 mostly result in an intractable mixture of products, use of a platinum electrode in dichloromethane results in substantial dehydrogenation, with a pathway that likely goes through N-ethyl-1,2,3,4-tetrahydrocarbazole.
Chapter 4: Oxygen Reduction Catalysis by First-Row Transition Metal Corrole Complexes
Oxygen reduction catalysis is demonstrated with copper, cobalt, and iron corrole complexes. Of these, cobalt(III) and iron(III) complexes of 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole demonstrate good activity for oxygen reduction in solution-phase studies (acidic acetonitrile) and in surface studies using rotating ring-disk voltammetry (aqueous sulfuric acid). The cobalt(III) complex favors two-electron reduction to hydrogen peroxide while the iron(III) complex favors four-electron reduction to water, both of which are consistent with previous literature reports.
Chapter 5: “Clickable” Propargyl- and Azido-Modified Metallocorrole Complexes: Biscorrole Characterization and Covalent Attachment
Electrochemical characterization of “click”-functionalized metallocorrole complexes is presented. Voltammetric studies indicate that addition of “click” functionality to the periphery of the corrole ligand does not affect the redox processes at the metal center, validating the previous study of related complexes for oxygen reduction catalysis. Huisgen azide-alkyne cycloaddition of propargyl- and azide-functionalized metallocorroles results in copper-copper and copper-iron biscorrole complexes, which display electronic independence of the metal centers. Finally, preliminary studies of covalent attachment of copper(III) 5,15-bis(pentafluorophenyl)-10-(4-propargyloxyphenyl)corrole to an azide-functionalized electrode surface are reported.
Conclusions and Future Outlook:
This work describes studies to improve hydrogen storage technologies and oxygen reduction catalysis for PEM fuel cells. Future work recommended in the area of virtual hydrogen storage includes: further elucidation and optimization of electrochemical dehydrogenation pathways, development of catalysts capable of performing proton-electron dehydrogenation, electrochemical study of other nitrogen heterocycles and other hydrogen-carrying fuels such as alcohols, and development of new membrane technologies that can support new fuel types. In the area of “clickable” catalysts, future work includes optimization of surface attachment and comparison of catalysis for covalently-bound vs. physisorbed catalysts. More broadly, this
3
enables the development and study of other click-functionalized metallocorrole complexes with applications in medical imaging, drug delivery, or environmental sensing.
i
For my parents, who created a childhood that nurtured my intellect and thirst for learning For my sisters, Tessa and Maia, who keep me grounded and remind me what life can be like
And for Joy, who loves me and believes in me more than I ever thought possible.
ii
Table of Contents
List of Figures ................................................................................................................................ iv List of Schemes ............................................................................................................................ viii List of Tables ................................................................................................................................. ix List of Equations ..............................................................................................................................x Acknowledgements ........................................................................................................................ xi
Chapter 1: An Overview of Proton-Exchange Membrane Fuel Cell Research Needs 1.1 Why Proton-Exchange Membrane Fuel Cells?.......................................................................1 1.2 Research Needs for PEM Fuel Cells .......................................................................................2 1.3 Introduction to Cyclic Voltammetry .......................................................................................4 1.4 Conclusions and Overview .....................................................................................................6 1.5 References ...............................................................................................................................6
Part I: Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage
Chapter 2: Electro-Dehydrogenation of Indoline: Direct Oxidation at the Electrode and Redox Catalysis 2.1 Introduction ...........................................................................................................................11 2.2 Results and Discussion .........................................................................................................11 2.2.1 Direct Electrode Oxidation of Indoline.....................................................................11 2.2.2 Effect of Added Base on Indoline Oxidation ............................................................17 2.2.3 Effect of Redox Catalysis on Indoline Oxidation .....................................................20 2.2.4 Controlled-Potential Electrolysis and Indole Formation ..........................................31 2.3 Conclusions ...........................................................................................................................33 2.4 Experimental Methods and Materials ...................................................................................34 2.5 References .............................................................................................................................35
Chapter 3: Electrochemical Analysis of a Fully Hydrogenated Carbazole as a Virtual Hydrogen Storage Compound 3.1 Introduction ...........................................................................................................................37 3.2 Results and Discussion .........................................................................................................37 3.2.1 Synthesis of N-Ethyldodecahydrocarbazole, N-Ethyloctahydrocarbazole and N-
Ethyl-1,2,3,4-tetrahydrocarbazole ............................................................................37 3.2.2 Direct Electrode Oxidation of N-Ethyldodecahydrocarbazole .................................39 3.2.3 Effect of Bases on N-Ethyldodecahydrocarbazole Voltammetry .............................48 3.2.4 Redox Catalysis of N-Ethyldodecahydrocarbazole Oxidation .................................51
iii
3.2.5 Bulk Electrolysis and Product Analysis of N-Ethyldodecahydrocarbazole ..............54 3.3 Conclusions ...........................................................................................................................58 3.4 Experimental Methods and Materials ...................................................................................58 3.5 References .............................................................................................................................60
Part II: “Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction
Chapter 4: Oxygen Reduction Catalysis by First-Row Transition Metal Corrole Complexes 4.1 Introduction ...........................................................................................................................65 4.2 Results and Discussion .........................................................................................................65 4.2.1 Voltammetric Characterization of Copper, Iron and Cobalt Corrole Complexes ....65 4.2.2 Solution-Phase Oxygen Reduction Reactivity ..........................................................70 4.2.3 Analysis of Oxygen Reduction Reactivity by Rotating Ring-Disk Voltammetry ....73 4.3 Conclusions ...........................................................................................................................78 4.4 Experimental Methods and Materials ...................................................................................78 4.5 References .............................................................................................................................79
Chapter 5: “Clickable” Propargyl- and Azido-Modified Metallocorrole Complexes: Biscorrole Characterization and Covalent Attachment 5.1 Introduction ...........................................................................................................................82 5.2 Results and Discussion .........................................................................................................82 5.2.1 Preparation of “Clickable” Metallocorroles and Homo- and Hetero-bimetallic
Appendix: Computational Evaluation of Virtual Hydrogen Storage Compounds and Comparison with Direct Electrode Oxidation Potentials A.1 Effect of Anode Reaction Potential on Overall Cell Voltage ...............................................94 A.2 Computational Evaluation of Substrates ...............................................................................94 A.3 Experimental Methods and Materials ...................................................................................96 A.4 References .............................................................................................................................96
iv
List of Figures
Figure 1.1: Schmatic of a proton-exchange membrane fuel cell .................................................1
Figure 2.1: Oxidation of 10 mM indoline ..................................................................................12
Figure 2.2: 10 mM indoline at multiple scan rates ....................................................................12
Figure 2.3: 40 mM indoline, 5 successive scans ........................................................................13
Figure 2.4: 100 mM indoline, 5 successive scans ......................................................................14
Figure 2.5: Peak current as a function of scan rate at varying concentrations of indoline between 5 mM and 100 mM. ...................................................................................15
Figure 2.6: Comparison of 0.5 mM ferrocene to 0.5 mM indoline. ..........................................16
Figure 2.7: Anodic peak potential as a function of indoline concentration at 100 mV/s and as a function of scan rate at 0.5 mM .................................................................17
Figure 2.8: Onset potentials for 10 mM indoline alone and with 100 mM imidazole, triethylamine, diethylenetriamine, or 1,1,3,3-tetramethylguanidine. ......................18
Figure 2.9: Comparison of 10 mM indoline alone, with 100 mM imidazole, or with 100 mM 1,1,3,3-tetramethylguanidine ...........................................................................19
Figure 2.10: 0.5 mM indoline with amounts of 2,6-lutidine increasing from 0.25 mM to 10 mM (left). Anodic peak potential as a function of scan rate for 0.5, 2, and 10 mM 2,6-lutidine added to 0.5 mM indoline (right) .................................................19
Figure 2.11: 0.5 mM indoline with amounts of 1,1,3,3-tetramethylguanidine increasing from 0.25 mM to 10 mM (left). Anodic peak potential as a function of scan rate for 10 mM 1,1,3,3-tetramethylguanidine added to 0.5 mM indoline (right). ......................................................................................................................20
Figure 2.13: Reduction potentials for 1 mM TCNQ, 2 mM BQ, 1 mM NQ and 1 mM AQ. ......21
Figure 2.14: 1 mM TCNQ with 10 mM indoline.........................................................................23
Figure 2.15: 1 mM NQ and 10 mM indoline with 100 mM imidazole. 1 mM NQ and 100 mM imidazole alone ................................................................................................23
Figure 2.17: Oxidation potentials for eight different ferrocene derivatives, showing the influence of different substituents ...........................................................................25
Figure 2.18: 1 mM 1,1’-diaminoferrocene and 1 mM 1,1’-bis(dimethylamino)ferrocene with 10 mM indoline ...............................................................................................26
Figure 2.19: 1 mM 1,1’-dimethylferrocene with 10 mM indoline, 10 mM indoline alone .........27
Figure 2.20: 1 mM ferrocene with 10 mM indoline, 10 mM indoline alone. ..............................27
Figure 2.21: 1 mM ethynylferrocene with 10 mM indoline, 10 mM indoline alone. ..................28
v
Figure 2.22: 1 mM decamethylferrocene with 10 mM indoline and 100 mM 1,1,3,3-tetramethylguanidine, 10 mM indoline and 100 mM tetramethylguanidine alone ........................................................................................................................28
Figure 2.23: 1 mM 1,1’-diaminoferrocene with 10 mM indoline and 100 mM 1,1,3,3-tetramethylguanidine, 1 mM 1,1’-diaminoferrocene and 100 mM tetramethylguanidine alone .....................................................................................29
Figure 2.24: 1 mM ferrocene with 10 mM indoline and 100 mM imidazole, 10 mM indoline and 100 mM imidazole alone ....................................................................29
Figure 2.25: 1 mM 1,1’-dimethylferrocene with 10 mM indoline and 100 mM imidazole, 10 mM indoline and 100 mM imidazole alone .......................................................30
Figure 2.26: Charge transferred over time for electrolysis of 10 mM indoline with 1-5 mM ferrocene ..................................................................................................................31
Figure 2.27: Charge transferred over time for electrolysis of 10 mM indoline with 100 mM imidazole and 1-5 mM ferrocene .....................................................................32
Figure 3.1: Some isomers of NEC-H12. .....................................................................................38
Figure 3.2: 1 mM NEC-H12 .......................................................................................................39
Figure 3.3: 1 mM NEC-H12 at scan rates ranging from 25 mV/s to 100 V/s .............................40
Figure 3.4: Peak current as a function of scan rate at concentrations of NEC-H12 between 0.5 mM and 20 mM .................................................................................................42
Figure 3.5: 1 mM ferrocene vs. 1 mM NEC-H12 .......................................................................43
Figure 3.6: 1 mM NEC-H12 at rotation rates ranging from 100 to 2500 rpm ............................44
Figure 3.7: Koutecký-Levich plot of data from Figure 3.6 at 0.6, 0.8 and 1 V. ........................45
Figure 3.8: 10 mM NEC-H12 vs. 10 mM of an 8.5% mixture of NEC-H8 in NEC-H12 ............47
Figure 3.9: 1 mM NEC-H4 .........................................................................................................47
Figure 3.10: 10 mM NEC-H12 with no base, 100 mM imidazole, 100 mM 2-ethyl-4-methylimidazole, 100 mM triethylamine, 100 mM diethylenetriamine, or 100 mM 1,1,3,3-tetramethylguanidine ...........................................................................48
Figure 3.11: Oxidation potentials for 100 mM imidazole, 2-ethyl-4-methylimidazole, triethylamine, diethylenetriamine, or 1,1,3,3-tetramethylguanidine .......................49
Figure 3.12: 10 mM NEC-H12 with varying concentrations of 1,1,3,3-tetramethylguanidine ...............................................................................................50
Figure 3.13: 1 mM ferrocene with 10 mM NEC-H12 ...................................................................51
Figure 3.14: 1 mM ferrocene with 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine, 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine alone .....................................................................................52
Figure 3.15: 1 mM ferrocene with varying concentrations of 1,1,3,3-tetramethylguanidine, showing the catalysis of 1,1,3,3-tetramethylguanidine oxidation by ferrocene ......52
vi
Figure 3.16: 1 mM ferrocene with 100 mM 1,1,3,3-tetramethylguanidine, and with 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine .............................................53
Figure 3.17: 1 mM ferrocene and 10 mM NEC-H12 alone (pink) and with concentrations of 1,1,3,3-tetramethylguandine ranging from 5 to 1000 mM ..................................53
Figure 3.18: Characteristic 1H NMR shifts of the ethyl CH2 group for N-ethylcarbazoles in varying states of hydrogenation ..............................................................................55
Figure 3.19: Tentative products from controlled-potential electrolysis of NEC-H12 in the presence of ferrocene and controlled-current electrolysis of NEC-H12 by two and four electrons ....................................................................................................56
Figure 3.20: 1H NMR spectrum of 1 mM NEC-H12 after 102 hours of electrolysis at a platinum electrode in dichloromethane ...................................................................57
Figure 4.1: Copper, iron and cobalt corrole complexes studied in this work ............................66
Figure 4.2: 1 mM (C6F5)2(p-OMePh)corroleCu ........................................................................67
Figure 4.3: 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2 .............................................................68
Figure 4.4: 0.54 mM ((C6F5)2(p-OMePh)corroleFe)2(μ-O) .......................................................69
Figure 4.5: 1 mM (C6F5)2(p-OMePh)corroleCo ........................................................................69
Figure 4.6: 100 mM acetic acid in oxygen-saturated solution with 1 mM (C6F5)2(p-OMePh)corroleCu ...................................................................................................71
Figure 4.7: 100 mM acetic acid in oxygen-saturated solution with 0.54 mM ((C6F5)2(p-OMePh)corroleFe)2(μ-O) ........................................................................................71
Figure 4.8: 100 mM acetic acid in oxygen-saturated solution with 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2 ........................................................................................72
Figure 4.9: 100 mM acetic acid in oxygen-saturated solution with 1 mM (C6F5)2(p-OMePh)corroleCo ...................................................................................................72
Figure 4.10: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleCu adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode .................................................................................74
Figure 4.11: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleFe(Et2O)2 adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode ......................................................................75
Figure 4.12: Koutecký-Levich analysis of oxygen reduction by (C6F5)2(p-OMePh)corroleFe(Et2O)2 ........................................................................................75
Figure 4.13: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleCo adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode .................................................................................77
Figure 4.14: Comparison of oxygen reduction by blank carbon ink and copper(III), iron(III), and cobalt(III) complexes .........................................................................77
vii
Figure 5.1: 1 mM (C6F5)2(p-OMePh)corroleCu, (C6F5)2(p-O(CH2CCH)Ph)corroleCu, or (C6F5)2(p-O(CH2(C2HN3(CH2CH2OCH2CH2OH)))Ph)corroleCu ..........................85
Figure 5.2: 1 mM copper-copper biscorrole complex (left). 1 mM (C6F5)2(p-OMePh)corroleCu (right) ........................................................................................86
Figure 5.3: <1 mM copper-iron biscorrole complex (left). 1 mM (C6F5)2(p-OMePh)corroleCu (green, right) and 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2 (red, right) ......................................................................86
Figure 5.4: Azide-modified glassy carbon disk reacted with (C6F5)2(p-O(CH2CCH)Ph)corroleCu, initially and after exposure to an aqueous saturated EDTA solution for 10 minutes .................................................................88
Figure 5.5: Azide-modified glassy carbon disk reacted with (C6F5)2(p-O(CH2CCH)Ph)corroleCu, initially and after exposure to an aqueous saturated solution of disodium diethyldithiolcarbamate for 10 minutes .................88
Scheme 1.1: Anode, cathode, and overall reactions taking place in a PEM fuel cell ....................2
Scheme 1.2: Demonstration of the difference between a conventional hydrogen storage system and a “virtual” hydrogen storage system .......................................................3
Scheme 2.1: Proposed dimerization mechanism for indoline, showing the relationship of 0.5 electrons transferred for every molecule of indoline. ........................................15
Scheme 2.2: Comparison of EC vs. CE mechanisms. ..................................................................18
Scheme 2.3: Comparison of RRD, EC and CE mechanisms. ......................................................20
Scheme 2.4: General scheme showing interaction between a quinone (Q) and a hydrogen-carrying substrate (LH2) ..........................................................................................22
Scheme 2.5: General scheme showing the interaction of ferrocene with a substrate (LH2) ........24
Scheme 3.1: Synthesis of NEC-H12 from NEC. ...........................................................................38
Scheme 3.2: Synthesis of NEC-H4 from 1,2,3,4-tetrahydrocarbazole. ........................................39
Scheme 4.1: Pathways for oxygen reduction to hydrogen peroxide or water ..............................73
Scheme 5.1: Syntheses of (C6F5)2(p-OMePh)corroleCu (5-1), (C6F5)2(p-OMePh)corroleFe(Et2O)2 (5-2), (C6F5)2(p-O(CH2CCH)Ph)corroleCu (5-3), (C6F5)2(p-O(CH2CCH)Ph)corroleFe(Et2O)2 (5-4), and (C6F5)2(m-CH2N3)Ph)corroleCu (5-5) ......................................................................................83
Scheme 5.2: (C6F5)2(p-O(CH2CCH)Ph)corroleCu (5-3) reacting with a small test substrate to form (C6F5)2(p-O(CH2(C2HN3(CH2CH2OCH2CH2OH)))Ph)corroleCu (5-6). ........................................................................................................................84
Scheme 5.3: Formation of biscorrole complexes 5-7 and 5-8 using Husigen azide-alkyne cycloaddition ...........................................................................................................84
ix
List of Tables
Table 2.1: Values of n for different concentrations of indoline. ..............................................15
Table 2.2: Reduction potentials for the quinones seen in Figure 2.13. ....................................22
Table 2.3: Redox potentials for the ferrocenes seen in Figure 2.16 .........................................25
Table 2.4: Summary of preparative electrolysis data for 10 mM indoline with 1–5 mM ferrocene ..................................................................................................................32
Table 2.5: Summary of preparative electrolysis data for 10 mM indoline with 100 mM imidazole and 1−5 mM ferrocene ...........................................................................33
Table 2.6: Summary of preparative electrolysis data for 10 mM indoline with 100 mM 1,1,3,3-tetramethylguanidine and 1, 3 and 5 mM decamethylferrocene .................33
Table 3.1: Values for α and n for concentrations of NEC-H12 between 0.5 mM and 20 mM ..........................................................................................................................42
Table 3.2: Comparison of n calculated from step-potential data for 1 mM and 10 mM NEC-H12 on both glassy carbon and platinum electrodes .......................................46
Table 3.3: Summary of changes to the onset potential for NEC-H12 oxidation in the presence of base.......................................................................................................48
Table 3.4: Values for α, n, and Epa vs. log(v) for 10 mM NEC-H12 and concentrations of 1,1,3,3-tetramethylguanidine between 0.5 mM and 100 mM .................................50
Table 3.5: Characteristic 1H NMR shifts of the ethyl CH2 group and major mass spectrometry peaks for N-ethylcarbazoles in varying states of hydrogenation .......54
Table 3.6: Comparison of controlled-potential electrolyses performed with 10 mM NEC-H12 using a carbon felt electrode in acetonitrile .............................................55
Table 3.7: Comparison of controlled-potential electrolyses performed with 1 mM NEC-H12 ...........................................................................................................................56
Table 4.1: Comparison of redox processes seen in copper corrole complexes ........................67
Table 4.2: Comparison of redox processes seen in iron corrole complexes .............................68
Table 4.3: Comparison of redox processes seen in cobalt corrole complexes .........................70
Table 4.4: Number of electrons and % hydrogen peroxide calculated for the reduction of oxygen by (C6F5)2(p-OMePh)corroleFe(Et2O)2 at different potentials. ..................76
Table 5.1: Assignment of oxidation and reduction processes for the bimetallic biscorroles and comparison with related compounds..............................................87
Table A.1: Computed and experimental first oxidation potentials for proposed virtual hydrogen storage compounds ..................................................................................95
x
List of Equations
Equation 2.1: Peak current for an irreversible process ..............................................................13
Equation 2.2: Peak current for a reversible process ..................................................................14
Equation 2.3: Relationship of total charge to moles of analyte and number of electrons transferred per molecule .....................................................................................31
Equation 3.1: Relationship between peak current (Ep) and half peak current (E1/2) for a reversible process ................................................................................................40
Equation 3.2: Relationship between peak current (Ep) and half peak current (E1/2) for an irreversible process .............................................................................................40
Equation 3.3: See Equation 2.1 .................................................................................................41
Equation 3.4: See Equation 2.2 .................................................................................................43
Equation 3.5: Comparison of peak currents for equimolar solutions ........................................44
Equation 4.1: See Equation 3.7 .................................................................................................74
Equation 4.2: Equation for calculating number of electrons using ring and disk currents from rotating ring-disk voltammetry ..................................................................76
Equation 4.3: Equation for calculating % hydrogen peroxide formed using ring and disk currents from rotating ring-disk voltammetry.....................................................76
Equation A.1: Overall cell voltage as a function of anode and cathode voltages ......................94
Equation A.2: Relationship of standard redox potentials to Gibbs’ free energy ........................94
Equation A.3: Calculation of theoretical one-electron oxidation potentials from gas-phase ionization potentials and solvation energies .......................................................94
xi
Acknowledgments
It feels fitting that this should be the last thing I write for my dissertation, as I can think of no better way to cap this experience than by thanking everyone who had a hand in helping me bring it to fruition.
To John Arnold, thank you for your support in allowing me to design my own graduate experience and explore new areas of research within the group. I have appreciated your kindness as a mentor, and your willingness to back my less-conventional career goals.
To John Kerr, thank you also for adopting me into your research group at LBNL and for your insights into the strange and wonderful workings of electrochemistry.
To Elise, thank you for your insights into the virtual hydrogen storage project and for establishing a solid foundation for electrochemistry in the Arnold Group. It was my pleasure to study indoline alongside you and learn how to approach an electrochemical problem from you.
To Ashleigh, thank you for teaching me how to run my very first CV, for being an incredibly useful senior lab member, and for your witty commentary over the years.
To Heather, thank you for being an amazing friend and mentor and for first making me feel like I could have friends in my research group. Sharing a lab space with you and collaborating on the corrole oxygen reduction project was one of the highlights of my research experience. Thank you also for teaching me (almost) everything I know about Canada, enabling me to be a less embarrassed global citizen. I look forward to seeing how you are going to change the world, because I know you will.
To Brendon, you are the best undergrad I could have ever asked for. You pushed me to understand better and to do more, and our shared love of electrochemistry kept me sane. I am so proud of you and excited to hear about your research now and in the future.
To Alison, thank you for being my “lone wolf project” subgroup partner, for sharing the experience of doing data-heavy research in a sea of synthesis, and for commiserating about the general trials and tribulations of grad life.
To the other members of the Arnold Group I have had the pleasure to know: Pete, Sergio, Thomas, Dan, Ben, Bernard, Jessica, Nick, Trevor, Mary, Clément, Oanh, María, Katherine, Michael, Lauren, Alex, Angela, and others. Thank you all for keeping the lab running smoothly and for making it a pleasant and collegial atmosphere in which to work. I am lucky to have known you all, and wish you all the best in the future.
To other members of the Kerr Group: Pete, Sébastien, and Kyle. Thank you for your guidance and support on the virtual hydrogen storage project, and particular thanks to Sébastien for both synthesizing hydrogenated carbazoles and teaching me new French words.
To other researchers I have had the opportunity to collaborate with, thank you for your contributions to these projects. They are far more than any one person could do. Thanks to Grigorii, Lakshmi, Matt, Davide, Oana, and everyone else on the EFRC team; to Jerzy, Neil and José-María from UC Fees; and to Thomas and Matt for their azide-modification expertise.
xii
To David, thank you for being my toxicology guru and for tackling the question of comparative nitrogen heterocycle toxicity with me.
To Marty and Meg, thank you for your mentorship and your tireless contributions to the Berkeley Center for Green Chemistry and the Greener Solutions program. My involvement with BCGC has been the absolute highlight of my time in graduate school, and I am so grateful for the work you have both done that made that possible.
To my two amazing interdisciplinary BCGC teams: Karen, Heather, Kierston, Joe, Jen, and Sara. Thank you all so much for being not only amazing to work with, but for being such cool people that we all wanted to hang out afterwards! I am so glad to have met all of you, and to have formed such deep friendships with you.
To Claire, thank you for being my I-House bestie our first year, and for providing a needed counter-balance to the many strong personalities we spent mealtimes with.
To Abby, Latisha, Cory, Steph, Kate, Jen, Jordan, Ashley, Beatriz, Erin, Raul, Josh, David, Kurt, Peter and Mike. Thank you for being a wonderful group of friends to go through this crazy experience with. We’ve slogged through phys org and inorganic problem sets together, stayed up late writing NSF proposals together, studied for quals together, bounced research ideas off each other, sought expert advice from each other, borrowed chemicals and instruments from each other, talked about career paths with each other, and gotten lunch, coffee, dinner, or drinks together countless times. Abby, thank you especially for your thesis cheerleading these past couple of months. Latisha, thank you for surviving C2M with me and for making that experience way more fun than it would have been otherwise. It is not an exaggeration when I say that I absolutely would not have finished this degree without all of you. I am so proud of all of you and excited to see the amazing and diverse places our lives take us.
To my other wonderful non-chemist friends in the Bay Area: Autumn, John, Meredith, Colleen, Justin, Paul, and Jacob. Thank you for providing a safe space to escape from the grad school life, and for always thinking I know lots of things about chemistry. I have been so grateful to have you all nearby, even I don’t get to see you as often as I’d like (and thank you for loving and supporting me despite that!).
To my sister Tessa, thank you for being my first best friend, even if we didn’t fully realize it until much later. You have tolerated my decidedly literal self starting from a very young age, have supported me and my quirks throughout our lives, and maybe still think I’m somewhat cool, even though we are both grown up. Thank you for being a feminism and social justice sounding board, for sharing in the wild and wonderful experience of being our parents’ children, and for always feeding me delicious things when I show up at your place. My life is absolutely richer for having you in it.
To my sister Maia, thank you for counterbalancing my scientific, nerdy, quantitatively-inclined self with your creativity and artistry. You let me go off to college when you were only four, forgave me (mostly) for abandoning you when I finally came back to California, and have developed into an amazing young woman whom I now consider a peer and a close friend. Thank you for filling my life with beauty through your paintings and photographs, and for introducing me to new music. Thank you, also, for sharing in the experience of growing up in our family, and
xiii
for always being ready to provide witty commentary on our parents and their particular idiosyncrasies. I am so proud of you and cannot wait to see where your art and life take you.
To my parents, Barbara and Lew, who provided a safe, supportive, educational, and nurturing childhood. Thank you for having a house full of books, and for encouraging us to always look things up in the dictionary, atlas, or encyclopedia. Thank you for sending me to schools that supported me both as a scholar and as a human being. Thank you for encouraging a wide range of interests and supporting my involvement in music and theater. And thank you for loving me and supporting me in the decisions that are right for me, even when they take me halfway around the world. Dad, thank you for inspiring my interest in energy issues, and for encouraging me to study science because “you can always get less technical later”. Mom, thank you for always providing excellent life and financial advice, and for doing so impartially. I would not be the person I am today without you, and I am grateful to you both.
And finally, to my wonderful, amazing, beloved wife Joy. There is very little I can say to thank you adequately for all the support you have given me throughout this process. We met six months before I took my qualifying exam, and since then you have listened to my rants, read over my fellowship and job application essays, shared in the ups and downs of my day-to-day research life, and told me that I was not only surviving, but sometimes even succeeding at this grad school thing. I am indebted to you for your emotional and physical support, from tasks as tangible as cooking and doing laundry to the intangible yet invaluable contribution of seeing me for who I really am and loving me so, so much anyway. I look forward to new adventures together and many long years of post-graduate school happiness with you. Thank you, deeply, from the bottom of my heart and with all my love.
– Leah
1
Chapter 1 An Overview of Proton-Exchange Membrane Fuel Cell Research Needs
1.1 Why Proton-Exchange Membrane Fuel Cells?
One of the greatest challenges facing our world today is the gap between our projected energy consumption and the ability of our planet to sustain this level of consumption. Our energy needs are largely powered, at this point, by burning carbon-based fuels such as petroleum, natural gas, and coal.1 This leads to a dependence that is unsustainable, due to both limitations in supply and negative environmental impacts. Extraction of these fuels from the environment can be quite environmentally damaging, leading to water and air pollution, ecosystem disruption, and negative impacts on human health.2,3 Additionally, relatively non-toxic emissions of carbon dioxide are contributing to a growing concern regarding climate change and the natural disasters that may be associated.4 In short, there is a need for fundamental research that can lead to innovative changes in our energy landscape.
An energy-storage technology that shows promise is the proton-exchange membrane (PEM) fuel cell. While there are several types of fuel cells, PEM fuel cells are particularly promising for transportation, as they start up quickly and have a favorable power-to-weight ratio.5 Fuel cells use essentially the same architecture as batteries, with two electrodes separated by an electrically insulating membrane (Figure 1.1).5 A different chemical reaction occurs at each electrode, generating a driving force for electrons to flow from one electrode to the other. However, since electrons cannot pass through the membrane, they must flow through an external circuit connecting the two electrodes, generating an electrical current.
Figure 1.1: Schematic of a proton-exchange membrane fuel cell.5
PEM fuel cells use hydrogen gas as a fuel, which oxidizes at the anode. This reaction produces protons and electrons, with two protons and two electrons being produced for each molecule of hydrogen. The protons are allowed to pass through the polymeric proton-exchange membrane
2
while the electrons travel around the outside of the cell through a circuit. Meanwhile, at the cathode, oxygen is reduced by those same electrons, and the reduced oxygen combines with the protons that have passed through the membrane, forming water (Scheme 1.1). The overall cell voltage is given by subtracting the voltage of the anode reaction from the voltage of the cathode reaction.
Scheme 1.1: Anode, cathode, and overall reactions taking place in a PEM fuel cell. Ecell is given by Ecathode – Eanode. Voltages reported vs. the Normal Hydrogen Electrode (NHE).
Since the only byproduct of PEM fuel cell usage is water, PEM fuel cells represent a significant environmental improvement over engines that burn carbon-based fuels.
1.2 Research Needs for PEM Fuel Cells
It is important to realize that while fuel cells are a very useful means of converting stored chemical energy into electricity, they are not themselves a source of energy. Fuel cells still require an added fuel source in order to produce any electricity. PEM fuel cells, as described above, use hydrogen as their fuel source, and one of the major limitations in PEM fuel cell usage is the feasibility of hydrogen transport and distribution.6 This is particularly true for mobile applications, which require lightweight, safe, and efficient on-board storage and delivery of hydrogen to the fuel cell.
Storage of hydrogen gas is most efficient at either high pressures or low temperatures, but high pressure hydrogen poses a significant safety risk, and low temperature hydrogen, a significant loss of fuel due to evaporation.6 Additionally, both high pressure and low temperature storage require specialized storage tanks, which add significant weight and volume. As alternatives to high-pressure or cryogenic storage, various other storage options have been proposed, including adsorbents such as metal-organic frameworks and carbon nanotubes, and chemical storage systems such as metal and non-metal hydrides.7 Since our current fueling systems are based on liquid fuel, the ideal carrier will be liquid, lightweight, easily dehydrogenated and rehydrogenated, and inexpensive.
Although many types of hydrogen storage materials have been considered, liquid organic compounds could offer a solution that is reversible, efficient, and lightweight.8 Reversibility of any storage system is key, in order to ensure longevity of the storage system. This would allow for a core scaffold that can be recycled, with only hydrogen being required to regenerate it. Reversible hydrogenation and dehydrogenation has been experimentally and computationally proven with a number of storage materials, in particular nitrogen heterocycles.9–16 Most studies of hydrogenated liquid organics focus on optimizing the dehydrogenation process for producing hydrogen in situ, but an alternative is a “virtual hydrogen storage” system, where a liquid organic
3
carrier is electrochemically dehydrogenated to produce protons and electrons directly in the fuel cell (Scheme 1.2).17
A virtual hydrogen storage system prevents the formation of free hydrogen, addressing safety concerns due to its flammability. Furthermore, hydrogen storage compounds that can be used directly in a fuel cell allow for the possibility of actually increasing the fuel cell voltage. This is because the overall voltage is a function of the reactions at both the cathode and the anode. In a conventional PEM fuel cell, the hydrogen oxidation reaction determines the anode voltage. This is true even when the hydrogen is stored and then released in situ. However, in a virtual hydrogen storage system, the anode voltage depends on the oxidation of the storage molecule, as opposed to hydrogen. This reaction could take place at voltages more negative than the voltage for hydrogen oxidation, increasing overall fuel cell voltage and enabling us to get more energy from each gallon of fuel.
Scheme 1.2: Demonstration of the difference between a conventional hydrogen storage system (top, red) and a “virtual” hydrogen storage system (bottom, blue). In both cases the dehydrogenated scaffold, N-ethylcarbazole, can be recycled and rehydrogenated for future use.
An optimal virtual hydrogen storage system would use a liquid fuel with high energy density that oxidizes at a potential equal to or negative of the potential for hydrogen oxidation (0.00 V vs. NHE). This fuel would also be able to be reversibly hydrogenated and dehydrogenated, maintaining the underlying scaffold (Scheme 1.2). In order to realize this goal, electrochemical studies of possible storage molecules are required to determine the potentials necessary for oxidation and the reversibility.
Another significant factor limiting widespread adoption of PEM fuel cells is the use of platinum catalysts for both the cathodic and anodic reactions. Platinum catalysts are expensive, easily poisoned by contaminants such as carbon monoxide, and contribute to unnecessary voltage
4
losses due to the high overpotential for oxygen reduction.5,18 In order for PEM fuel cells to become a widely-viable technology, a robust, inexpensive catalyst must be developed.
Multiple approaches to reducing the use of platinum in oxygen reduction catalysis have been studied, including platinum nanoparticles, platinum alloys, doped carbon materials, metal oxides, and bio-inspired catalysts such as macrocyclic metal complexes.18 Examples of the latter include porphyrin, phthalocyanine, and corrole complexes, and examples of each have demonstrated oxygen reduction capability.19–22 While transition metal porphyrin chemistry is well-established, analogous corrole complexes have been reported more recently. Corrole ligands are trianionic, stabilizing higher metal oxidation states than their dianionic porphyrin counterparts, which can make them particularly reactive towards small molecules.23,24 Multiple metallocorroles have been prepared and characterized in recent years, using alkali metals, d-block transition metals, lanthanides, and actinides.24,25 These complexes show promise in a variety of applications, with iron and cobalt corrole complexes in particular having demonstrated utility in catalyzing oxygen reduction.21,23,26–32
While recent examples of oxygen reduction by metallocorroles demonstrate that these small molecule catalysts can be used in oxygen reduction, they cannot practically be used in the solution phase in a device due to problems with aggregation and catalyst deactivation. There is thus a need for incorporating a tethering arm into the corrole ligand that can be attached to a modified electrode surface,33 which would preclude catalyst aggregation. Tethering could also allow the corrole complexes to function as pseudo-bimetallic systems due to complex proximity. This could also mimic bimetallic systems that are more selective for four-electron reduction to water than similar monometallic complexes.21,34
1.3 Introduction to Cyclic Voltammetry
The study of electron transfer and electrochemical reactions at electrode surfaces has been well-elucidated over the years, with numerous books35–38 and reviews39–45 having been written on the subject. One of the most common techniques applied is cyclic voltammetry, due to its versatility in identifying oxidation and reduction processes for a compound of interest. Cyclic voltammetry can also elucidate information about the nature of the electron transfer(s) and any subsequent chemical reactions. Although multiple electrochemical techniques are used in this work, cyclic voltammetry is the most common and is thus explained below. Other techniques will be introduced as needed in the following chapters.
Cyclic voltammetry uses a three-electrode system, consisting of a working electrode, a counter electrode, and a reference electrode. The working electrode is typically composed of a substance that is robust, electrically conductive, and otherwise inert under the analytical conditions. Typical substances are glassy carbon, platinum, and gold, with glassy carbon being the primary choice for this work. Glassy carbon was chosen because it is widely used and provides a large voltage window for analysis. Additionally, during the evaluation of possible oxygen reduction catalysts, platinum would be competitive and would therefore interfere with the analysis. The working electrode must also have a well-defined and small surface area, such that the currents produced can be well-modeled. The counter electrode can be composed of similar substances to the working electrode, with the main concern being that, in contrast to the working electrode, it
5
must have a very large surface area. This work uses a high surface area platinum mesh counter electrode, which is a common choice.
Multiple types of reference electrodes exist, with varying applications. The key for any reference electrode is that it has a consistent chemical composition throughout the length of the experiment. The normal hydrogen electrode (NHE) is not always practical for actual experimental work, but is considered the standard to which other electrodes are referenced. For aqueous work, both the saturated calomel electrode (SCE) and the silver-silver chloride electrode (Ag/AgCl) are commonly used. SCE has a voltage of 0.241 V vs. NHE and Ag/AgCl (saturated KCl) has a voltage of 0.197 V vs. NHE (other concentrations of chloride salts will shift the potential of the Ag/AgCl electrode).35 For non-aqueous work, reference electrodes containing water and chloride are problematic, and a non-aqueous silver-silver ion (Ag/Ag+) reference electrode is preferred. This can be prepared with silver nitrate in acetonitrile. A molecule with well-defined electrochemical behavior, such as ferrocene, can also be added as an external reference. The Ag/Ag+ reference electrode (0.01 M AgNO3) has a potential of 0.548 V vs. NHE, and ferrocene has a potential of 0.630 vs. NHE.46
Figure 1.2: Sample cyclic voltammogram, showing an oxidation (positive current) followed by a reduction (negative current) after switching the direction of the potential sweep.
To obtain a cyclic voltammogram, the three electrodes are placed into a solution containing the analyte and attached to a potentiostat, which controls the potential at the working electrode. The counter electrode provides the other half of the circuit, so that there is a complete circuit between the working electrode and the counter electrode. There is also a circuit between the working electrode and the reference electrode, so that the voltage at the working electrode is always being precisely measured. The voltage is then swept in a linear fashion, first increasing and then decreasing (or vice versa) and the resulting current response is measured. For a molecule with a
-30
-20
-10
0
10
20
30
40
-0.4 -0.3 -0.2 -0.1 0 0.1 0.2 0.3Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
6
well-defined, reversible redox process, two peaks are seen: a peak in the anodic direction due to oxidation current, and a peak in the cathodic direction due to reduction current (Figure 1.2). These data are typically rendered with potential in volts on the x axis and current in amperes on the y axis, although current density (current divided by working electrode area) is also seen. Conventions for which axis direction corresponds to oxidation and which to reduction vary; this work uses positive voltages and currents to refer to oxidation. The shape of the cyclic voltammogram is a function of both electrode potential, which is the driving force for the oxidation or reduction reaction, and mass transport, which corresponds to the diffusion of the analyte to the electrode.
1.4 Conclusions and Overview
The following chapters summarize efforts undertaken to study both liquid organic hydrogen storage compounds and new “clickable” oxygen reduction catalysts using electrochemical techniques. Part I addresses fundamental limitations in our understanding of the electrochemical of partially and fully hydrogenated nitrogen heterocycles. Chapter 2 discusses the electrochemical behavior of indoline, a model hydrogen storage compound chosen for study due to its relatively simple dehydrogenation pathway. Chapter 3 expands further on this work with a study of N-ethyldodecahydrocarbazole and related compounds. This provides the first electrochemical study of a fully hydrogenated N-heterocycle. Part II switches focus to the cathode reaction, discussing the development of corrole complexes that catalyze oxygen reduction and can be covalently attached to a functionalized electrode surface. Chapter 4 characterizes the oxygen reduction capabilities of copper, iron and cobalt corrole complexes. Finally, Chapter 5 discusses the characterization of “click”-functionalized metallocorroles and some preliminary efforts to attach them to surfaces. The work overall contributes to the development of new technological solutions for hydrogen storage and platinum-free oxygen reduction catalysis, addressing limitations of conventional hydrogen-powered fuel cells that could enable us to move away from a society powered by fossil fuels.
1.5 References
1. United States Department of Energy. Department of Energy - Energy Sources. http://www.energy.gov/energysources/index.htm (accessed May 3, 2011).
2. Mascarelli, A. After the Oil. Nature 2010, 467, 22–24. 3. Hendryx, M.; Ahern, M. M. Relations between Health Indicators and Residential
Proximity to Coal Mining in West Virginia. Am. J. Public Health 2008, 98, 669–671. 4. Warren, R.; Price, J.; Fischlin, A.; de la Nava Santos, S.; Midgley, G. Increasing Impacts
of Climate Change upon Ecosystems with Increasing Global Mean Temperature Rise. Clim. Change 2011, 106, 141–177.
5. United States Department of Energy. FCT Fuel Cells: Basics. http://www1.eere.energy.gov/hydrogenandfuelcells/fuelcells/basics.html (accessed May 10, 2011).
6. Schlapbach, L.; Zuttel, A. Hydrogen-Storage Materials for Mobile Applications. Nature 2001, 414, 353–358.
7. Durbin, D. J.; Malardier-Jugroot, C. Review of Hydrogen Storage Techniques for on Board Vehicle Applications. Int. J. Hydrogen Energy 2013, 38, 14595–14617.
7
8. Makowski, P.; Thomas, A.; Kuhn, P.; Goettmann, F. Organic Materials for Hydrogen Storage Applications: From Physisorption on Organic Solids to Chemisorption in Organic Molecules. Energy Environ. Sci. 2009, 2, 480–490.
9. Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H. Hydrogen Storage by Reversible Hydrogenation of Pi-Conjugated Substrates. US Pat. 7,101,530 2006.
10. Moores, A.; Poyatos, M.; Luo, Y.; Crabtree, R. H. Catalysed Low Temperature H2 Release from Nitrogen Heterocycles. New J. Chem. 2006, 30, 1675–1678.
11. Clot, E.; Eisenstein, O.; Crabtree, R. H. Computational Structure–Activity Relationships in H2 Storage: How Placement of N Atoms Affects Release Temperatures in Organic Liquid Storage Materials. Chem. Commun. 2007, 2231–2233.
12. Sung, J. S.; Choo, K. Y.; Kim, T. H.; Tarasov, A. L.; Tkachenko, O. P.; Kustov, L. M. A New Hydrogen Storage System Based on Efficient Reversible Catalytic Hydrogenation/Dehydrogenation of Terphenyl. Int. J. Hydrogen Energy 2008, 33, 2721–2728.
13. Araujo, C. M.; Simone, D. L.; Konezny, S. J.; Shim, A.; Crabtree, R. H.; Soloveichik, G. L.; Batista, V. S. Fuel Selection for a Regenerative Organic Fuel Cell/Flow Battery: Thermodynamic Considerations. Energy Environ. Sci. 2012, 5, 9534–9542.
14. Sotoodeh, F.; Smith, K. J. An Overview of the Kinetics and Catalysis of Hydrogen Storage on Organic Liquids. Can. J. Chem. Eng. 2013, 9999, 1–14.
15. Yang, M.; Dong, Y.; Fei, S.; Pan, Q.; Ni, G.; Han, C.; Ke, H.; Fang, Q.; Cheng, H. Hydrogenation of N-Propylcarbazole over Supported Ruthenium as a New Prototype of Liquid Organic Hydrogen Carriers (LOHC). RSC Adv. 2013, 3, 24877–24881.
16. Dong, Y.; Yang, M.; Yang, Z.; Ke, H.; Cheng, H. Catalytic Hydrogenation and Dehydrogenation of N-Ethylindole as a New Heteroaromatic Liquid Organic Hydrogen Carrier. Int. J. Hydrogen Energy 2015, doi:10.1016/j.ijhydene.2015.05.196.
17. Crabtree, R. H. Hydrogen Storage in Liquid Organic Heterocycles. Energy Environ. Sci. 2008, 1, 134.
18. Morozan, A.; Jousselme, B.; Palacin, S. Low-Platinum and Platinum-Free Catalysts for the Oxygen Reduction Reaction at Fuel Cell Cathodes. Energy Environ. Sci. 2011, 4, 1238–1254.
19. Collman, J. P.; Marrocco, M.; Denisevich, P.; Koval, C.; Anson, F. C. Potent Catalysis of the Electroreduction of Oxygen to Water by Dicobalt Porphyrin Dimers Adsorbed on Graphite Electrodes. J. Electroanal. Chem. Interfacial Electrochem. 1979, 101, 117–122.
20. Zagal, J.; Páez, M.; Tanaka, A. A.; dos Santos Jr., J. R.; Linkous, C. A. Electrocatalytic Activity of Metal Phthalocyanines for Oxygen Reduction. J. Electroanal. Chem. 1992, 339, 13–30.
21. Kadish, K. M.; Frémond, L.; Ou, Z.; Shao, J.; Shi, C.; Anson, F. C.; Burdet, F.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Cobalt(III) Corroles as Electrocatalysts for the Reduction of Dioxygen: Reactivity of a Monocorrole, Biscorroles, and Porphyrin-Corrole Dyads. J. Am. Chem. Soc. 2005, 127, 5625–5631.
22. Wang, B. Recent Development of Non-Platinum Catalysts for Oxygen Reduction Reaction. J. Power Sources 2005, 152, 1–15.
23. Collman, J. P.; Kaplun, M.; Decréau, R. A. Metal Corroles as Electrocatalysts for Oxygen Reduction. Dalton Trans. 2006, 554–559.
24. Aviv-Harel, I.; Gross, Z. Aura of Corroles. Chem. Eur. J. 2009, 15, 8382–8394.
8
25. Buckley, H. L.; Arnold, J. Recent Developments in out-of-Plane Metallocorrole Chemistry across the Periodic Table. Dalton Trans. 2015, 44, 30–36.
26. Kadish, K. M.; Shao, J.; Ou, Z.; Frémond, L.; Zhan, R.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Electrochemistry, Spectroelectrochemistry, Chloride Binding, and O2 Catalytic Reactions of Free-Base Porphyrin-Cobalt Corrole Dyads. Inorg. Chem. 2005, 44, 6744–6754.
27. Kadish, K. M.; Frémond, L.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Cobalt(IV) Corroles as Catalysts for the Electroreduction of O2: Reactions of Heterobimetallic Dyads Containing a Face-to-Face Linked Fe(III) or Mn(III) Porphyrin. J. Inorg. Biochem. 2006, 100, 858–868.
28. Kadish, K. M.; Shen, J.; Frémond, L.; Chen, P.; El Ojaimi, M.; Chkounda, M.; Gros, C. P.; Barbe, J.-M.; Ohkubo, K.; Fukuzumi, S.; et al. Clarification of the Oxidation State of Cobalt Corroles in Heterogeneous and Homogeneous Catalytic Reduction of Dioxygen. Inorg. Chem. 2008, 47, 6726–6737.
29. Kadish, K. M.; Frémond, L.; Shen, J.; Chen, P.; Ohkubo, K.; Fukuzumi, S.; El Ojaimi, M.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Catalytic Activity of Biscobalt Porphyrin-Corrole Dyads toward the Reduction of Dioxygen. Inorg. Chem. 2009, 48, 2571–2582.
30. Dogutan, D. K.; Stoian, S. A.; McGuire, Jr., R.; Schwalbe, M.; Teets, T. S.; Nocera, D. G. Hangman Corroles: Efficient Synthesis and Oxygen Reaction Chemistry. J. Am. Chem. Soc. 2011, 133, 131–140.
31. Schechter, A.; Stanevsky, M.; Mahammed, A.; Gross, Z. Four-Electron Oxygen Reduction by Brominated Cobalt Corrole. Inorg. Chem. 2012, 51, 22–24.
32. Wang, Z.; Lei, H.; Cao, R.; Zhang, M. Cobalt Corrole on Carbon Nanotube as a Synergistic Catalyst for Oxygen Reduction Reaction in Acid Media. Electrochim. Acta 2015, 171, 81–88.
33. Devadoss, A.; Chidsey, C. E. D. Azide-Modified Graphitic Surfaces for Covalent Attachment of Alkyne-Terminated Molecules by “Click” Chemistry. J. Am. Chem. Soc. 2007, 129, 5370–5371.
34. Rosenthal, J.; Nocera, D. G. Role of Proton-Coupled Electron Transfer in O-O Bond Activation. Acc. Chem. Res. 2007, 40, 543–553.
35. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons: New York, NY, 2001.
36. Zanello, P. Inorganic Electrochemistry: Theory, Practice and Applications; Royal Society of Chemistry: Cambridge, UK, 2003.
37. Savéant, J.-M. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry; John Wiley & Sons: Hoboken, NJ, 2006.
38. Compton, R. G.; Banks, C. E. Understanding Voltammetry, 2nd ed.; Imperial College Press: London, UK, 2011.
39. Marcus, R. A. Chemical and Electrochemical Electron-Transfer Theory. Annu. Rev. Phys. Chem. 1964, 15, 155–196.
40. Costentin, C. Electrochemical Approach to the Mechanistic Study of Proton-Coupled Electron Transfer. Chem. Rev. 2008, 108, 2145–2179.
41. Evans, D. H. One-Electron and Two-Electron Transfers in Electrochemistry and Homogeneous Solution Reactions. Chem. Rev. 2008, 108, 2113–2144.
43. Savéant, J.-M. Electrochemical Approach to Proton-Coupled Electron Transfers: Recent Advances. Energy Environ. Sci. 2012, 7718–7731.
44. Henstridge, M. C.; Laborda, E.; Rees, N. V.; Compton, R. G. Marcus-Hush-Chidsey Theory of Electron Transfer Applied to Voltammetry: A Review. Electrochim. Acta 2012, 84, 12–20.
45. Weinberg, D. R.; Gagliardi, C. J.; Hull, J. F.; Fecenko Murphy, C.; Kent, C. A.; Westlake, B. C.; Paul, A.; Ess, D. H.; Granville McCafferty, D.; Meyer, T. J. Proton-Coupled Electron Transfer. Chem. Rev. 2012, 112, 4016–4093.
46. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
10
Part I:
Hydrogenated N-Heterocyclic Compounds for Virtual Hydrogen Storage
11
Chapter 2
Electro-Dehydrogenation of Indoline: Direct Oxidation at the Electrode and Redox Catalysis
2.1 Introduction
The successful development of a functioning virtual hydrogen storage system rests on a number of factors, one of which is the fuel selection itself. Nitrogen heterocycles have been proposed as good hydrogen storage materials due to their thermodynamically favorable dehydrogenation reactivity, which has been demonstrated both computationally and experimentally.1–4 Compounds where the nitrogen is present in a five-membered ring, such as pyrroles, indoles, and carbazoles, have been shown to be particularly thermodynamically accessible. Although fully hydrogenated compounds can store more protons and electrons by weight, partially hydrogenated compounds with only a few protons and electrons available are easier to study from a mechanistic perspective.
Indoline was chosen as a model fuel compound for study due to its partial hydrogenation, which was anticipated to provide an easily-elucidated mechanism. It has been studied as a hydrogen storage compound using thermal catalytic dehydrogenation,2,5 and has been shown to react by a two proton, two electron mechanism to form indole, its fully dehydrogenated counterpart.6 Previous electrochemical studies of indoline have studied its behavior alone and in the presence of bases;7,8 this work builds on that by adding redox catalysis. Both voltammetry and preparative electrolysis are used herein to characterize the electrochemical reactivity of indoline and its potential for use as a virtual hydrogen storage carrier.9
2.2 Results and Discussion
2.2.1 Direct Electrode Oxidation of Indoline
Indoline oxidizes at 0.34 V vs. Ag/Ag+ at a glassy carbon electrode in acetonitrile (Figure 2.1). Evaluation at multiple scan rates (Figure 2.2) shows an anodic shift in oxidation peak potential with increasing scan rate and the lack of a substantial return reduction peak, indicating an electrochemically irreversible oxidation or a chemical reaction following the electron transfer (or both).
The return peaks observed when the potential is scanned cathodically after oxidation are broad, misshapen and much shallower than the sharp oxidation peak (Figure 2.2). This suggests that either the cathodic process has a slow electron transfer or, more likely, there is slow formation of a reaction product following the initial oxidation. Assuming the latter, the reaction product is also electrochemically active and is reduced at the electrode between 0.1 and −0.2 V vs. Ag/Ag+. The presence of two reductions and two oxidations suggest perhaps a dimeric product, containing two moieties in electronic communication with each other. The more extended conjugation system of a dimer would also lead to a lower redox potential, as seen here.
12
Figure 2.1: Oxidation of 10 mM indoline. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.2: 10 mM indoline at multiple scan rates. 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode.
Upon repeated scans, the oxidation current diminishes. The changes over multiple scans are more dramatic at higher concentrations and scan rates, and further evidence for the formation of a new species following the initial electron transfer is seen. For example, successive scans of a 40 mM solution of indoline at 2000 mV/s produce two reduction peaks, at 0.03 and −0.19 V, and two subsequent additional oxidation peaks, at 0.10 and −0.12 V (Figure 2.3). Indole was examined separately and found to oxidize above 0.9 V, so these additional peaks are not due to indole formation. At even higher concentrations (100 mM indoline), the new peaks appear to coalesce into one redox couple, at −0.23 (red) and −0.11 (ox) V (Figure 2.4).
The peak current for an irreversible process is defined as10
Equation 2.1
where n is the total number of electrons involved in the process, α is the transfer coefficient, A is the area of the electrode (cm2), D is the diffusion coefficient of the analyte (cm2/s), C is the concentration of the analyte in bulk solution (mol/mL), and v is the scan rate (V/s). The transfer coefficient describes the transition state of an electron transfer reaction. If the transition state is equally similar to both the reduced and oxidized forms, α is equal to 0.5. Greater similarity to the starting form results in α being less than 0.5, and vice versa.
Figure 2.3: 40 mM indoline, five successive scans. 0.1 M NEt4BF4 in acetonitrile, 2000 mV/s, glassy carbon electrode.
-1000
0
1000
2000
3000
4000
5000
6000
-0.4 -0.2 0 0.2 0.4 0.6
Scan 1Scan 2Scan 3Scan 4Scan 5
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
14
Figure 2.4: 100 mM indoline, five successive scans. 0.1 M NEt4BF4 in acetonitrile, 2000 mV/s, glassy carbon electrode.
The relationship given in Equation 2.1 was used to further analyze voltammetric results for indoline at different concentrations and scan rates. The area of the electrode used was 0.143 cm2. This is the real surface area, as opposed to the geometric surface area, and was calculated from ferrocene voltammetry and a similar peak current definition for reversible processes (Equation 2.2).10 The diffusion coefficient for indoline was estimated using molecular weight as a proxy, giving a value of 2.7 x 10−5 cm2/s.11 The transfer coefficient was assumed to be 0.5.
Equation 2.2
Plots of ip vs. v1/2 for indoline result in linear fits, with slopes equal to (299000)nα1/2AD1/2C (Figure 2.5). The values of n calculated for each concentration are shown in Table 2.1. With the exception of the data for 5 mM, analysis at all concentrations indicates that n is 0.5, suggesting that a second molecule of indoline is involved in the chemical reaction following the initial oxidation, meaning the effective concentration of indoline oxidized is halved and thus only half the expected number of electrons are transferred. These data, combined with the new peaks seen in Figures 2.3 and 2.4, suggest a dimerization mechanism (Scheme 2.1). There is precedent for dimerization and oligomerization of indoline, as seen in the electrochemical synthesis of polyindoline.12
-2000
0
2000
4000
6000
8000
10000
12000
-0.4 -0.2 0 0.2 0.4 0.6 0.8
Scan 1Scan 2Scan 3Scan 4Scan 5
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
15
Figure 2.5: Peak current as a function of scan rate at concentrations of indoline between 5 mM and 100 mM. 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode.
[indoline], mM slope of ip vs. v1/2 n 5 1.22 x 10−3 1.55
10 0.94 x 10−3 0.60 20 1.93 x 10−3 0.61 40 3.17 x 10−3 0.50 100 7.45 x 10−3 0.47
Table 2.1: Values of n for different concentrations of indoline.
Scheme 2.1: Proposed dimerization mechanism for indoline, showing the relationship of 0.5 electrons transferred for every molecule of indoline.
0
0.002
0.004
0.006
0.008
0.01
0.012
0 0.5 1 1.5
5 mM10 mM20 mM40 mM100 mM
anod
ic p
eak
curr
ent (
A)
(scan rate, V/s)1/2
2 LH2 2 [LH2]●+
+ 2e−
1. Initial electron transfer
2 [LH2]●+
[H2L-LH2]2+
2. Radical-radical coupling
[H2L-LH2]2+
+ 2 LH2 HL-LH + 2 [LH3]+
3. Deprotonation 4 LH2 HL-LH + 2 [LH3]
+ + 2e
−
RRD LH2 =
16
Additional data13 further support the dimerization mechanism. A comparison of peak heights between indoline and ferrocene give a value for n of 0.6 (Figure 2.6). Analysis of the shift in peak potential (Ep) at different concentrations or scan rates is a well-studied means of analyzing a chemical reaction following an electron transfer. Plots of Ep vs. log(v) and Ep vs. log(concentration) give slopes of 22 and 24 mV per decade, respectively (Figure 2.7), which is consistent with a radical-radical dimerization mechanism (RRD).14 Finally, controlled-potential (0.34 V) electrolysis13 of indoline at a carbon felt electrode led to complete consumption of indoline, as seen by liquid chromatography and gas chromatography. Two polymeric products were detected by gel permeation chromatography (GPC), with approximate molecular weights of 600 and 3500. Given that the presumed dimer of indoline oxidizes at a potential more negative than the oxidation potential of indoline (Figures 2.3 and 2.4), it follows that at the electrolysis potential, the dimer will continue to react, forming higher weight polymeric products.
Figure 2.6: Comparison of 0.5 mM ferrocene (orange) to 0.5 mM indoline (blue). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode. Data courtesy of Elise Deunf.
-10
-5
0
5
10
15
20
-0.4 -0.3 -0.2 -0.1 0 0.1 0.2 0.3 0.4
0.5 mM ferrocene0.5 mM indoline
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
17
Figure 2.7: Anodic peak potential as a function of indoline concentration at 100 mV/s and as a function of scan rate at 0.5 mM. 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode. Data courtesy of Elise Deunf.
2.2.2 Effect of Added Base on Indoline Oxidation
In a fuel cell environment, a proton-exchange membrane helps drive the oxidation reaction by providing a path for protons from the anode to the cathode. In the absence of such a membrane, bases of different strength were added to indoline to evaluate the effects on the voltammetric reaction. Previous work has shown that the oxidation potential can be lowered in the presence of base7,8 and this was also seen here, with stronger bases resulting in greater shifts. The onset of oxidative current for indoline occurs around 0.15 V (Figure 2.8), but the addition of a base shifts the current onset potential to anywhere from −0.10 V (imidazole, pKa = 14.2 in acetonitrile15) to −0.50 V (1,1,3,3-tetramethylguanidine, pKa = 23.3 in acetonitrile16). The dependence on base strength is likely because deprotonated indoline has more electron density on the nitrogen, making it easier to oxidize than indoline17 and the equilibrium between indoline and deprotonated indoline will shift towards more deprotonation with stronger bases.
Comparison of indoline oxidation with imidazole or 1,1,3,3-tetramethylguanidine demonstrates further mechanistic differences in indoline oxidation in the presence of weak vs. strong bases. As can be seen in Figure 2.9, both imidazole (a weak base) and tetramethylguanidine (a strong base) increase the current at or around the peak potential of indoline (0.34 V vs. Ag/Ag+). However, tetramethylguanidine also produces a second oxidative peak at a more negative potential (−0.19 V vs. Ag/Ag+). This could be explained by the formation of a nearly- or fully-deprotonated form of indoline that can be oxidized at a lower potential, as also seen in Figure 2.8.18 In mechanistic terms, this could indicate that we are transitioning from an EC (electrochemical-chemical) mechanism to a CE (chemical-electrochemical) mechanism. In other words, an initial chemical
0.33
0.34
0.35
0.36
0.37
0.38
-1.5 -1 -0.5 0 0.5
-0.8 -0.6 -0.4 -0.2 0 0.2 0.4
Ep vs. log(v)Ep vs. log(C)
log(v, V/s)
Peak
Pot
entia
l (V
vs. A
g/A
g+ )
log(C, mM)
18
reaction, the deprotonation of indoline, is now the first step (C) and the electron transfer (E) is the second step (Scheme 2.2).
Scheme 2.2: Comparison of EC vs. CE mechanisms.
Figure 2.8: Onset potentials for 10 mM indoline alone (blue) and with 100 mM imidazole (light green), 100 mM triethylamine (purple), 100 mM diethylenetriamine (orange) or 100 mM 1,1,3,3-tetramethylguanidine (pink). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Additional analysis13 provides further insight into the mechanism of indoline oxidation in the presence of base. Figure 2.10 illustrates the shift from a dimerization mechanism to an EC mechanism in the presence of a weak base. Increasing amounts of base (left) lead to an increase in the peak current and a cathodic shift in the peak potential. Scan rate analysis (right) reveals a shift of approximately 30 mV per decade, consistent with an EC mechanism. Due to the change in the presence of base, the following chemical reaction is likely deprotonation. Figure 2.11 displays the same “pre-oxidation” seen above in Figure 2.9, corresponding to the oxidation of partially or fully deprotonated indoline. The scan rate analysis is inconclusive, as a slope of 52 mV per decade indicates that the reaction no longer taking place under “pure kinetic conditions”, where the rate of the following chemical reaction is fast enough to have an influence on the voltammetric response.14 However, as stated above, the presence of the pre-oxidation suggests that we have switched to a CE mechanism.
0
10
20
30
40
50
60
70
80
-0.6 -0.5 -0.4 -0.3 -0.2 -0.1 0 0.1 0.2
10 mM indoline+ 100 mM imidazole+ 100 mM triethylamine+ 100 mM diethylenetriamine+ 100 mM tetramethylguanidine
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
LH2 [LH2]●+
+ e− LH2 + B [LH]
− + BH
+
[LH2]●+
+ B [LH]● + BH
+ [LH]
− [LH]
● + e
−
EC CE
19
Figure 2.9: Comparison of 10 mM indoline alone (blue) and with 100 mM imidazole (light green) or 100 mM 1,1,3,3-tetramethylguanidine (pink). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.10: 0.5 mM indoline with amounts of 2,6-lutidine increasing from 0.25 mM to 10 mM (left). Anodic peak potential vs. log(scan rate) for 0.5, 2, and 10 mM 2,6-lutidine added to 0.5 mM indoline (right). 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode. Data courtesy of Elise Deunf.
-200
0
200
400
600
800
1000
1200
1400
-0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8 1
10 mM indoline+ 100 mM imidazole+ 100 mM tetramethylguanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-10
0
10
20
30
40
-0.2 -0.1 0 0.1 0.2 0.3 0.4 0.5
0.5 mM indoline+ 0.25 mM lutidine+ 0.5 mM lutidine+ 1 mM lutidine+ 2 mM lutidine+ 5 mM lutidine+ 10 mM lutidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
0.28
0.3
0.32
0.34
0.36
0.38
-1.2 -1 -0.8 -0.6 -0.4 -0.2 0 0.2 0.4
1 eq. lutidine4 eq. lutidine20 eq. lutidine
Slopes: 32, 29, 35 mV per decade
log(v, V/s)
Peak
Pot
entia
l (V
vs. A
g/A
g+ )
20
Figure 2.11: 0.5 mM indoline with amounts of 1,1,3,3-tetramethylguanidine increasing from 0.25 mM to 10 mM (left). Anodic peak potential vs. log(scan rate) for 10 mM 1,1,3,3-tetramethylguanidine added to 0.5 mM indoline (right). 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode. Data courtesy of Elise Deunf.
Scheme 2.3: Comparison of RRD, EC and CE mechanisms.
The mechanism of indoline oxidation is therefore affected by the presence and strength of added base, allowing access to both EC and CE mechanisms.
2.2.3 Effect of Redox Catalysis on Indoline Oxidation
Redox catalysis has been shown to enhance electrochemical reactivity of organic substrates by lowering overpotentials.19,20 Both ferrocene- and quinone-based compounds have been shown to act as homogenous catalysts in electrochemical reactions,19–21 and were studied as possible redox catalysts for indoline oxidation. These catalysts operate via an outer-sphere electron transfer mechanism: the given catalyst is oxidized (or reduced) at the electrode and then transfers an electron from (or to) the substrate in solution.
Quinones are well-known two-electron, two-proton transfer agents, which could potentially facilitate the two-electron oxidation of indoline. Dichlorodicyanoquinone (DDQ), for example, was recently demonstrated as a two-electron, two-proton dehydrogenation catalyst for benzylaniline.6,22 The mechanism of dehydrogenation depended on the catalyst-substrate ratio; stoichiometric levels of DDQ led to hydride transfer and hydroquinone formation, but catalytic amounts resulted in formation of the DDQ radical anion.
0
10
20
30
40
50
-0.6 -0.4 -0.2 0 0.2 0.4 0.6
0.5 mM indoline+ 0.25 mM tetramethylguanidine+ 0.5 mM tetramethylguanidine+ 1 mM tetramethylguanidine+ 2 mM tetramethylguanidine+ 5 mM tetramethylguanidine+ 10 mM tetramethylguanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
0.32
0.33
0.34
0.35
0.36
0.37
0.38
0.39
0.4
-1.2 -1 -0.8 -0.6 -0.4 -0.2 0 0.2 0.4
20 eq. tetramethylguanidine
log(v, V/s)
Peak
Pot
entia
l (V
vs. A
g/A
g+ )
Slope: 52 mV per decade
2 LH2 2 [LH2]●+
+ 2e−
LH2 [LH2]●+
+ e− LH2 + B [LH]
− + BH
+
2 [LH2]●+
[H2L-LH2]2+
[LH2]●+
+ B [LH]● + BH
+ [LH]
− [LH]
● + e
−
Direct electrode oxidation In presence of weak base In presence of strong base
EC CE RRD
21
Four different quinones were screened as possible catalysts for indoline oxidation (Figure 2.12): tetracyanoquinodimethane (TCNQ), 1,4-benzoquinone (BQ), 1,4-naphthoquinone (NQ), and anthraquinone (AQ). Although the electrode oxidation potential of indoline is quite positive, at 0.34 V vs. Ag/Ag+, the thermodynamic oxidation potential is likely to be lower. Because of this, even quinones with quite negative reduction potentials (Figure 2.13, Table 2.2) were screened.
Figure 2.12: Quinone-based catalysts.
The double reduction peaks seen in Figure 2.13 are characteristic of quinone reduction in aprotic solvents, as the quinone is reduced in two distinct one-electron steps, from quinone to the radical anion (Q−) and then to the diradical (Q2−).23 In protic solvents, this reduction is usually accompanied by two protonations, forming hydroquinone (although not in the case of TCNQ, as it has no oxygen atoms to protonate). In either case, the reaction is readily reversible.
Figure 2.13: Reduction potentials for 1 mM TCNQ (green, top), 2 mM BQ (blue, 2nd from top), 1 mM NQ (purple, 3rd from top) and 1 mM AQ (pink, bottom). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Table 2.2: Reduction potentials for the quinones seen in Figure 2.13. Values given in V vs. Ag/Ag+.
In the case of quinone interaction with a substrate, we would anticipate a two-electron oxidation of the substrate, coupled with a two-electron reduction of the quinone. The anticipated voltammetric response assuming catalytic substrate oxidation would be to see a decrease or even elimination of quinone reduction current, and an increase in the associated oxidation current. This would imply that the quinone pre-reduces by oxidizing the substrate, leading to less or no reduction current. It then will re-oxidize at the electrode, producing an anodic current (Scheme 2.4).
Scheme 2.4: General scheme showing interaction between a quinone (Q) and a hydrogen-carrying substrate (LH2). Pre-reaction produces a doubly-reduced quinone (Q2−) and a doubly-oxidized substrate. The reduced quinone then re-oxidizes at the electrode and the substrate loses two protons to form the dehydrogenated product (L).
If the process is catalytic, as soon as the quinone is oxidized at the electrode it will immediately oxidize another molecule of substrate and be available for oxidation again. Assuming that the substrate oxidation and re-oxidation of quinone are both suitably fast, an increase in oxidation current over the response seen from the quinone by itself will be seen. This oxidation current may correspond to oxidation of either the quinone diradical or of hydroquinone, depending on the timescale and extent of quinone protonation. In the case of indoline, we ultimately anticipate the loss of two protons, but the proton loss and the formation of hydroquinone may not occur on the voltammetry timescale. In the case of hydroquinone formation, an oxidation peak will still be seen, but it will be shifted anodically, as hydroquinone is more difficult to oxidize than the quinone diradical. Although TCNQ has a similar voltammetry profile to the other quinones, it differs in that it does not form hydroquinone, even in protic media.
Addition of indoline to TCNQ caused a dramatic decrease in reduction current (Figure 2.14). However, the corresponding decrease in oxidation current and shifts in potential do not indicate a reversible redox interaction with indoline. Rather, these suggest that some other process is preventing the quinone from re-oxidizing at the electrode, perhaps a non-reversible reaction between the two. This result led to TCNQ not being pursued further.
No changes were seen in the voltammograms of BQ, NQ, or AQ in the presence of indoline, likely because their redox potentials are simply too negative to interact with indoline. Since it was demonstrated in the previous section that adding a base to indoline will lower its oxidation
23
potential (Figure 2.8), imidazole was added to the quinone-indoline solutions to see if that would facilitate catalysis.
Figure 2.14: 1 mM TCNQ, without (green) and with (light green) 10 mM indoline. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.15: 1 mM NQ and 10 mM indoline, without (light blue) and with (dark blue) 100 mM imidazole. 1 mM NQ and 100 mM imidazole alone (pink, dashed). 0.1 M NEt4BF4 in acetonitrile, 200 mV/s, glassy carbon electrode.
-60
-40
-20
0
20
40
-1 -0.8 -0.6 -0.4 -0.2 0 0.2 0.4
1 mM TCNQ alone1 mM TCNQ, 10 mM indoline
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
-100
-50
0
50
100
-2 -1.6 -1.2 -0.8 -0.4
1 mM NQ, 10 mM indoline1 mM NQ, 10 mM indoline, 100 mM imidazole1 mM NQ, 100 mM imidazole
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
24
For all three quinones, there were changes in the voltammograms upon addition of imidazole, notably, that the current increased for the first reduction and the second reduction peak completely disappeared (Figure 2.15). However, further investigation revealed that these changes were due entirely to interaction between the quinones and imidazole, and that indoline was not involved. This is likely due to the imidazole generating a protic environment, leading to concerted formation of hydroquinone. This is, however, not useful for our application, and BQ, NQ, and AQ were also not pursued further.
Ferrocene and its derivatives are also quite well-known as stable redox shuttles. Unlike the quinones, the anticipated voltammetric response for ferrocenes will be different (Scheme 2.5). Ferrocene must first oxidize to ferrocenium at the electrode before it can oxidize a substrate. We would then expect to see an increase in current at the ferrocene oxidation potential in the presence of a suitable substrate, as the substrate reduces the ferrocenium back to ferrocene and the ferrocene is able to be oxidized again. This allows for more current than would be seen in the absence of an oxidizable substrate.
Scheme 2.5: General scheme showing the interaction of ferrocene with a substrate (LH2). Ferrocene will initially oxidize at the electrode , then will oxidize LH2, returning to its initial state.
Figure 2.16: Ferrocene-based catalysts.
Eight different ferrocene derivatives (Figure 2.16) were evaluated to determine their redox potentials, in order to choose the most appropriate potential match for indoline. The presence of
electron-donating or -withdrawing substituents on the cyclopentadienyl rings of ferrocene derivatives can have a significant impact on the redox potential of the FeII/III transition,24–26 and as a result we see a range of redox potentials that spans over a volt (Figure 2.17, Table 2.3).
Figure 2.17: Oxidation potentials for eight different ferrocene derivatives, showing the influence of different substituents. L−R: 1,1’-bis(dimethylamino)ferrocene (red), 1,1’-diaminoferrocene (orange), decamethylferrocene (yellow), 1,1’-dimethylferrocene (light green), (hydroxymethyl)ferrocene (green), ferrocene (light blue), ethynylferrocene (blue), and 1,1’-dibromoferrocene (purple). 1 mM ferrocene, 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Table 2.3: Redox potentials for the ferrocenyl compounds seen in Figure 2.16. Values given in V vs. Ag/Ag+.
As mentioned before, the electrode oxidation potential for indoline may be more positive than the thermodynamic redox potential. This is because there may be a significant overpotential associated with direct oxidation at the electrode. Since the exact thermodynamic redox potential for indoline is not known, all ferrocene derivatives with more negative potentials than the indoline electrode oxidation could be redox catalysts. For use in a fuel cell application, the aim is to reduce overpotential as much as possible, because this will increase the overall cell voltage.
As can be seen in Figure 2.17, 1,1’-dibromoferrocene oxidizes at too positive a potential to catalyze indoline oxidation, since the indoline will oxidize at the bare electrode before 1,1’-dibromoferrocene does. (Hydroxymethyl)ferrocene has a potential almost identical to ferrocene, thus it was assumed that it would behave similarly to ferrocene. The advantage of (hydroxymethyl)ferrocene over ferrocene is that it is water-soluble, which makes it a good potential mediator in the aqueous environment of a fuel cell.
When 1,1’-diaminoferrocene and 1,1’-bis(dimethylamino)ferrocene were examined with indoline, no catalysis was observed (Figure 2.18). Decamethylferrocene, while not shown here, also does not catalyze indoline oxidation.13 This suggests that these catalysts oxidize at potentials more negative than the oxidation potential of indoline, and that the redox potential for indoline is somewhere between−0.42 V (redox potential of decamethylferrocene) and 0.34 V (indoline oxidation potential at bare electrode) vs. Ag/Ag+.
Figure 2.18: 1 mM 1,1’-diaminoferrocene, without (orange) and with (gold, dashed) 10 mM indoline (left). 1 mM 1,1’-bis(dimethylamino)ferrocene, without (red) and with (pink, dashed) 10 mM indoline. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
1,1’-dimethylferrocene, ferrocene, and ethynylferrocene were also all examined for their ability to catalyze indoline oxidation. As can be seen in Figures 2.19–2.21, ferrocene is the best option for indoline catalysis. 1,1’-dimethylferrocene oxidizes at too negative a potential to produce more than a very mild catalytic effect (Figure 2.19). While adding indoline to ethynylferrocene produces significant current at the catalyst peak potential (Figure 2.21), the overlay of indoline oxidation (gold, dashed) shows that a large part of this current is likely due to direct oxidation of indoline at the electrode.
Ferrocenyl catalysts were also examined with indoline in the presence of base, to see if an excess of base would facilitate deprotonation and increase catalytic current at the ferrocene peak potential. In particular, there was interest in using a strong base to pre-react with indoline, facilitating catalysis by some of the ferrocenes with more negative oxidation potentials.
-40
-20
0
20
40
60
80
-0.8 -0.7 -0.6 -0.5 -0.4
1 mM Fc(NMe2)21 mM Fc(NMe2)2, 10 mM indoline
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-30
-20
-10
0
10
20
30
40
50
-1 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4 -0.3 -0.2
1 mM Fc(NH2)21 mM Fc(NH2)2, 10 mM indoline
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
27
Figure 2.19: 1 mM 1,1’-dimethylferrocene, without (dark green) and with (light green) 10 mM indoline. 10 mM indoline alone (gold, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.20: 1 mM ferrocene, without (dark blue) and with (light blue) 10 mM indoline. 10 mM indoline alone (gold, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-40
-20
0
20
40
60
80
-0.2 -0.1 0 0.1 0.2
1 mM FcMe21 mM FcMe2, 10 mM indoline10 mM indoline
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-40
-20
0
20
40
60
80
-0.2 -0.1 0 0.1 0.2
1 mM Fc1 mM Fc, 10 mM indoline10 mM indoline
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
28
Figure 2.21: 1 mM ethynylferrocene, without (dark purple) and with (light purple) 10 mM indoline. 10 mM indoline alone (gold, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.22: 1 mM decamethylferrocene, without (gold) and with (blue) 10 mM indoline and 100 mM 1,1,3,3-tetramethylguanidine (Me4guanidine). 10 mM indoline and 100 mM tetramethylguanidine alone (red, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode. Data courtesy of Elise Deunf.
-100
0
100
200
300
400
500
600
-0.1 0 0.1 0.2 0.3 0.4 0.5 0.6
1 mM ethynylFc1 mM ethynylFc, 10 mM indoline10 mM indoline
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-20
0
20
40
60
80
-0.7 -0.6 -0.5 -0.4 -0.3
1 mM FcMe101 mM FcMe10, 10 mM indoline, 100 mM Me4guanidine10 mM indoline, 100 mM Me4guanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
29
Figure 2.23: 1 mM 1,1’-diaminoferrocene, without (orange) and with (green) 10 mM indoline and 100 mM 1,1,3,3-tetramethylguanidine (Me4guanidine). 1 mM 1,1’-diaminoferrocene and 100 mM tetramethylguanidine alone (blue, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 2.24: 1 mM ferrocene, without (blue) and with (green) 10 mM indoline and 100 mM imidazole. 10 mM indoline and 100 mM imidazole alone (brown, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-50
0
50
100
150
200
250
-0.8 -0.6 -0.4 -0.2 0
1 mM Fc(NH2)21 mM Fc(NH2)2, 10 mM indoline, 100 mM Me4guanidine1 mM Fc(NH2)2, 100 mM Me4guanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-200
0
200
400
600
800
1000
1200
1400
-0.3 -0.2 -0.1 0 0.1 0.2 0.3
1 mM Fc1 mM Fc, 10 mM indoline, 100 mM imidazole10 mM indoline, 100 mM imidazole
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
30
Figure 2.25: 1 mM 1,1’-dimethylferrocene, without (green) and with (light blue) 10 mM indoline and 100 mM imidazole. 10 mM indoline and 100 mM imidazole alone (brown, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Decamethylferrocene was shown to be an effective catalyst in the presence of 1,1,3,3-tetramethylguanidine (Figure 2.22).13 1,1,3,3-tetramethylguanidine was also added to 1,1’-diaminoferrocene, but was found to interact with the catalyst even in the absence of indoline (Figure 2.23). Weaker bases were not effective at facilitating catalysis by 1,1’-diaminoferrocene. Additionally, even 1,1,3,3-tetramethylguanidine, the strongest base used, could not facilitate indoline oxidation by 1,1’-bis(dimethylamino)ferrocene.
For ferrocenes with more positive oxidation potentials, weaker bases such as imidazole can be used to assist catalysis. The addition of imidazole produces a dramatic increase in current at the catalyst peak potential in the presence of indoline for both ferrocene (Figure 2.24) and 1,1’-dimethylferrocene (Figure 2.25). The increase in current density is not attributable to indoline and imidazole alone or to catalyst and imidazole alone; all three components must be present.
The ultimate purpose of the redox mediator screening was to find a suitable catalyst-base combination to be used to study controlled-potential electrolysis of indoline. Two systems were selected for further study: indoline with ferrocene and imidazole, and indoline with decamethylferrocene and 1,1,3,3-tetramethylguanidine. This was to allow for study of both EC and CE mechanisms, depending on base strength. Additionally, the voltage attainable with of decamethylferrocene (0.13 V vs. NHE) is comparable to the hydrogen oxidation potential, which could enable us to get close to the same voltage as a conventional PEM fuel cell.
-200
0
200
400
600
800
1000
-0.3 -0.2 -0.1 0 0.1 0.2 0.3
1 mM FcMe21 mM FcMe2, 10 mM indoline, 100 mM imidazole10 mM indoline, 100 mM imidazole
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
31
2.2.4 Controlled-Potential Electrolysis and Indole Formation
In order to further study the mechanisms of indoline reactivity and to assess if indoline can be electrochemically dehydrogenated to indole, controlled-potential electrolyses were run. Electrolyses were carried out for 20 hours on 100 mL of a 10 mM indoline solution containing varying concentrations of ferrocene and imidazole. The total charge passed can be defined as
Equation 2.3
where n is the number of electrons per analyte molecule, F is Faraday’s constant (96485 C/mol) and N is the moles of analyte in solution. Equation 2.3 was used to calculate the average number of electrons per molecule of indoline in solution. The percentages of indoline consumed and indole formed were calculated by dividing the amount of each in the final solution by the amount (of the same substance) in the initial solution, using a high-performance liquid chromatography (HPLC) method calibrated for quantitative detection of both indoline and indole.
Table 2.4 summarizes the results for 10 mM indoline with ferrocene, no imidazole present. Electrolysis potentials were set at the peak potential of ferrocene, to minimize direct indoline oxidation at the electrode. The charge passed over time can be seen in Figure 2.26.
Figure 2.26: Charge transferred over time for electrolysis of 10 mM indoline with 1–5 mM ferrocene. 0.1 M NEt4BF4 in acetonitrile, carbon felt electrode.
0
50
100
150
200
0 5 10 15 20
1 mM ferrocene2 mM ferrocene3 mM ferrocene4 mM ferrocene5 mM ferrocene
Table 2.4: Summary of preparative electrolysis data for 10 mM indoline with 1–5 mM ferrocene. 0.1 M NEt4BF4 in acetonitrile, carbon felt electrode. (a) Data courtesy of Elise Deunf; electrolysis was performed on only 50 mL solution. (b) HPLC data for starting solution unavailable; calculations performed assuming a starting concentration of 10 mM indoline.
Similar data can be found in Figure 2.27 and Table 2.5, which summarize the electrolysis data for 10 mM indoline with 100 imidazole and 1–5 mM ferrocene.
Figure 2.27: Charge transferred over time for electrolysis of 10 mM indoline with 100 mM imidazole and 1–5 mM ferrocene. 0.1 M NEt4BF4 in acetonitrile, carbon felt electrode.
There is a dependence of charge on catalyst concentration up to 2 mM, but not beyond. In both cases, large jumps in total charge are seen between 1 mM and 2 mM ferrocene, but the charge for 2–5 mM ferrocene stays fairly constant. This indicates that up to 2 mM ferrocene, the rate of initial electron transfer is relevant and enhanced, but beyond that other factors (such as mass transport) may have a greater influence.
0
100
200
300
400
0 5 10 15 20
1 mM ferrocene2 mM ferrocene3 mM ferrocene4 mM ferrocene5 mM ferrocene
Table 2.5: Summary of preparative electrolysis data for 10 mM indoline with 100 mM imidazole and 1–5 mM ferrocene. 0.1 M NEt4BF4 in acetonitrile, carbon felt electrode.
Another conclusion drawn is that while ferrocene is necessary to form some indole, between 25 and 35% indole formation is seen in all cases, with no clear dependence on the amount of ferrocene. Additionally, imidazole does not actually improve the amount of indole formed, and may even reduce it, although HPLC quantitation is not accurate enough to state that definitively.
In contrast, the results of preparative electrolyses performed using decamethylferrocene and 1,1,3,3-tetramethylguanidine show a clear dependence on catalyst concentration, with up to 50% indole formation possible when 5 mM catalyst are used (Table 2.6).13 This suggests that the rate of electron transfer is at least partially rate-determining for the overall process under this mechanism.
[Decamethylferrocene] (mM) % Indoline Consumed % Indole Formed
none 70% 12% 1 87% 16% 3 85% 42% 5 100%(a) 51%
Table 2.6: Summary of preparative electrolysis data for 10 mM indoline with 100 mM 1,1,3,3-tetramethylguanidine and 1, 3 and 5 mM decamethylferrocene. 0.1 M NEt4BF4 in acetonitrile, carbon felt electrode. Data courtesy of Elise Deunf. (a) Due to poor solubility of the catalyst in acetonitrile, experiment was run at 40 °C.
2.3 Conclusions
Indoline has been shown to react by a dimerization mechanism in the absence of any other components. Addition of base converts this mechanism to either an EC mechanism (weak bases) or a CE mechanism (stronger bases), with a favoring of deprotonation as the following chemical step. Quinone- and ferrocene-based redox catalysts have been evaluated, and multiple base-catalyst combinations for catalyzing indoline oxidation have been determined. The use of a weak base, imidazole, and ferrocene as redox mediator does not produce more than 35% or so of indole, with no clear dependence on catalyst concentration. However, the use of a strong base and decamethylferrocene enables up to 50% conversion to indole at a low applied potential (−0.4 V vs. Ag/Ag+), validating its possible use in a PEM fuel cell as part of a virtual hydrogen storage system.
34
2.4 Experimental Methods and Materials
Ferrocene was obtained commercially and purified by sublimation before use. Tetraethylammonium tetrafluoroborate (NEt4BF4) was obtained commercially and recrystallized from ethanol. All other reagents used were obtained commercially and used as received. Solutions for electrochemical analysis were made using HPLC-grade acetonitrile and 0.1 M NEt4BF4 as the supporting electrolyte.
Diaminoferrocene (Fc(NH2)2) was synthesized by a previous group member according to a published procedure.26
Synthesis of 1,1’-bis(dimethylamino)ferrocene: Fc(NMe2)2 was prepared according to a modified literature procedure27 under air-free conditions. Diaminoferrocene (0.432 g, 2 mmol) was placed in a Schlenk flask. Acetic acid (10 mL) was added and the solution stirred under nitrogen, turning from light yellow to greenish brown. Paraformaldehyde (1.21 g, 40 mmol) was added, turning the solution opaque and yellow-green. Sodium cyanoborohydride (1.257 g, 20 mmol) was added and the solution became foamy and bright yellow. The mixture was stirred for 16 hours under nitrogen. Concentrated sodium hydroxide was added, and the solution turned red, then orange as the reaction flask became warm. The product was extracted into hexanes, giving a dark red solution. This solution was washed with water and brine, dried over sodium sulfate, filtered through a frit, and concentrated. It was then filtered through alumina with more hexanes, reconcentrated, and cooled to −40 °C, forming flat orange crystals (0.1063 g, 19.5% yield). 1H NMR (400 MHz, CDCl3): δ 4.10 (s, 4H), 3.85 (s, 4H), 2.58 (s, 12H). UV-Vis (nm): 300. ESI-MS (+) Calcd: 272.0970 for C14H20N2Fe1 Observed: 272.0964.
Nuclear magnetic resonance (NMR) spectra were obtained at ambient temperature in CDCl3 using Bruker AV-300, AVB-400, or AVQ-400 spectrometers at the University of California, Berkeley, and referenced to the residual solvent peak. UV-visible spectra were determined with a Varian Cary 50 UV-Vis spectrophotometer. Mass spectral data were obtained at the University of California, Berkeley Microanalytical Facility
Cyclic voltammetry experiments were performed using a Gamry 600 Reference potentiostat. Voltammetry solutions were deoxygenated before use and kept under a blanket of inert gas (either nitrogen or argon) during analysis. A three-electrode cell was used, with a glassy carbon working electrode (3 mm diameter), a platinum mesh counter electrode, and a non-aqueous silver ion reference electrode. The reference electrode was prepared by immersing a silver wire in a 0.01 M solution of silver nitrate in supporting electrolyte, and was kept separate from the bulk solution by a Vycor frit. Potentials are reported vs. Ag/Ag+ (0.01 AgNO3), which is 0.548 V vs. NHE and −0.087 V vs. ferrocene (Fc).28
Controlled-potential electrolyses were performed using a PAR 276 potentiostat. The working electrode was composed of carbon felt (99%, Alfa Aesar), which was rinsed with dilute hydrochloric acid and isopropyl alcohol prior to use. The counter electrode was platinum mesh, separated from the bulk solution by a porous frit. The reference electrode was prepared with a silver wire in a 0.01 M solution of AgNO3 in acetonitrile, contained in a glass tube with a ceramic junction. Potentials for electrolysis were chosen to be just negative of the oxidation peak
35
of the catalyst or substrate (in the absence of catalyst) after an initial cyclic voltammogram of the solution using a platinum disk working electrode.
Indoline consumption and indole formation from the bulk electrolyses were quantified by reverse-phase liquid chromatography (LC), using a C18 column and water-acetonitrile gradient elution on an Agilent 1120 LC instrument using a UV-Vis detector at 254 nm. LC solvents were filtered through polyvinylidene membranes and degassed prior to use. o-Terphenyl was added to the LC samples as a standard for calibration and quantitation.
2.5 References
1. Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H. Hydrogen Storage by Reversible Hydrogenation of Pi-Conjugated Substrates. US Pat. 7,101,530 2006.
2. Moores, A.; Poyatos, M.; Luo, Y.; Crabtree, R. H. Catalysed Low Temperature H2 Release from Nitrogen Heterocycles. New J. Chem. 2006, 30, 1675–1678.
3. Clot, E.; Eisenstein, O.; Crabtree, R. H. Computational Structure–Activity Relationships in H2 Storage: How Placement of N Atoms Affects Release Temperatures in Organic Liquid Storage Materials. Chem. Commun. 2007, 2231–2233.
4. Araujo, C. M.; Simone, D. L.; Konezny, S. J.; Shim, A.; Crabtree, R. H.; Soloveichik, G. L.; Batista, V. S. Fuel Selection for a Regenerative Organic Fuel Cell/Flow Battery: Thermodynamic Considerations. Energy Environ. Sci. 2012, 5, 9534–9542.
5. Dean, D.; Davis, B.; Jessop, P. G. The Effect of Temperature, Catalyst and Sterics on the Rate of N-Heterocycle Dehydrogenation for Hydrogen Storage. New J. Chem. 2011, 35, 417–422.
6. Luca, O. R.; Wang, T.; Konezny, S. J.; Batista, V. S.; Crabtree, R. H. DDQ as an Electrocatalyst for Amine Dehydrogenation, a Model System for Virtual Hydrogen Storage. New J. Chem. 2011, 35, 998–999.
7. Rainka, M. P.; Peters, A.; Soloveichik, G. Base Effects on Electrochemical Oxidation of Indoline. Int. J. Hydrogen Energy 2013, 38, 3773–3777.
8. Peters, A. J.; Rainka, M. P.; Krishnan, L.; Laramie, S.; Dodd, M.; Reimer, J. A. Electrochemical Characterization of Hydrogen-Bonding Complexation between Indoline and Nitrogen Containing Bases. J. Electroanal. Chem. 2013, 691, 57–65.
9. Portions of the data presented in this chapter were obtained by a colleague, Dr. Elise Deunf. These data are presented here in support of the larger argument, and are noted and cited as needed in the text.
10. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons: New York, NY, 2001.
11. Valencia, D. P.; González, F. J. Estimation of Diffusion Coefficients by Using a Linear Correlation between the Diffusion Coefficient and Molecular Weight. J. Electroanal. Chem. 2012, 681, 121–126.
12. Özden, M.; Ekinci, E.; Karagözler, A. E. Synthesis and Optimization of Permselective Polymer (Polyindoline) Film. J. Solid State Electrochem. 1998, 2, 427–431.
13. Deunf, E.; Rubin, L. K.; Arnold, J.; Kerr, J. B. Electro-Dehydrogenation of a Liquid Hydrogen Carrier Fuel: Direct Oxidation at the Electrode and Redox Catalysis. Unpublished Work.
36
14. Savéant, J.-M. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry; John Wiley & Sons: Hoboken, NJ, 2006.
15. Streitwieser, A.; Kim, Y.-J. Ion Pair Basicity of Some Amines in THF: Implications for Ion Pair Acidity Scales. J. Am. Chem. Soc. 2000, 122, 11783–11786.
16. Gierczyk, B.; Wojciechowski, G.; Brzezinski, B.; Grech, E.; Schroeder, G. Study of the Decarboxylation Mechanism of Fluorobenzoic Acids by Strong N-Bases. J. Phys. Org. Chem. 2001, 14, 691–696.
17. Costentin, C. Electrochemical Approach to the Mechanistic Study of Proton-Coupled Electron Transfer. Chem. Rev. 2008, 108, 2145–2179.
18. Costentin, C.; Hajj, V.; Robert, M.; Savéant, J.-M.; Tard, C. Effect of Base Pairing on the Electrochemical Oxidation of Guanine. J. Am. Chem. Soc. 2010, 132, 10142–10147.
19. Andrieux, C. P.; Merz, A.; Savéant, J.-M.; Tomahogh, R. Preceding Chemical Reaction Mechanisms in Homogeneous Electron Transfer Reactions. Mediated Electrochemical Reduction of Highly Reactive Benzylic Halides. J. Am. Chem. Soc. 1984, 106, 1957–1962.
21. Houmam, A. Electron Transfer Initiated Reactions: Bond Formation and Bond Dissociation. Chem. Rev. 2008, 108, 2180–2237.
22. Driscoll, P. F.; Deunf, E.; Rubin, L.; Arnold, J.; Kerr, J. B. Electrochemical Redox Catalysis for Electrochemical Dehydrogenation of Liquid Hydrogen Carrier Fuels for Energy Storage and Conversion. J. Electrochem. Soc. 2013, 160, G3152–G3158.
23. Quan, M.; Sanchez, D.; Wasylkiw, M. F.; Smith, D. K. Voltammetry of Quinones in Unbuffered Aqueous Solution: Reassessing the Roles of Proton Transfer and Hydrogen Bonding in the Aqueous Electrochemistry of Quinones. J. Am. Chem. Soc. 2007, 129, 12847–12856.
24. Silva, M. E. N. P. R. A.; Pombeiro, A. J. L.; Fraústo da Silva, J. J. R.; Herrmann, R.; Deus, N.; Castilho, T. J.; Silva, M. F. C. G. Redox Potential and Substituent Effects at Ferrocene Derivatives. Estimates of Hammett σp and Taft Polar σ* Substituent Constants. J. Organomet. Chem. 1991, 421, 75–90.
25. Batterjee, S. M.; Marzouk, M. I.; Aazab, M. E.; El-Hashash, M. A. The Electrochemistry of Some Ferrocene Derivatives: Redox Potential and Substituent Effects. Appl. Organomet. Chem. 2003, 17, 291–297.
26. Shafir, A.; Power, M. P.; Whitener, G. D.; Arnold, J. Synthesis, Structure, and Properties of 1,1′-Diamino- and 1,1′-Diazidoferrocene. Organometallics 2000, 19, 3978–3982.
27. Metallinos, C.; Zaifman, J.; Dudding, T.; Van Belle, L.; Taban, K. Asymmetric Lithiation of Boron Trifluoride-Activated Aminoferrocenes: An Experimental and Computational Investigation. Adv. Synth. Catal. 2010, 352, 1967–1982.
28. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
37
Chapter 3 Electrochemical Analysis of a Fully Hydrogenated Carbazole as a Virtual
Hydrogen Storage Compound
3.1 Introduction
One of the factors influencing the fuel selection for a hydrogen storage system is the gravimetric storage density of the fuel, that is, how many protons and electrons are stored in a fuel molecule relative to the overall molecular weight. N-ethyldodecahydrocarbazole (NEC-H12), a fully hydrogenated compound that stores 5.7% hydrogen by weight, was first proposed as a hydrogen storage material that could be reversibly hydrogenated and dehydrogenated over a rhodium catalyst at a relatively modest temperature (125 °C).1 Since then, multiple studies have examined NEC-H12 hydrogenation and dehydrogenation mechanisms, delved into the influences of temperature, pressure, and catalyst material, and developed new catalysts.2–26 However, little attention has been paid to the electrochemical behavior of NEC-H12, with the one exception being a study that examined the open circuit voltage and currents produced by NEC-H12 in a fuel cell.27 That study concluded that while NEC-H12 was capable of producing very high open circuit voltages, the dehydrogenation mechanisms were complex and would require substantial optimization.
The work described in this chapter is an attempt to further understand how NEC-H12 might operate in a virtual hydrogen storage system, wherein the dehydrogenation process proceeds as stepwise removal of protons and electrons. While electrochemical dehydrogenation for fuel cells has been studied extensively for alcohols,28–32 particularly methanol, and for partially hydrogenated heterocycles such as indoline,33–35 this is the first example of an electrochemical study of a fully hydrogenated N-heterocycle.
The direct electrode reactivity of NEC-H12 was studied in the presence of both bases and redox catalysts, similar to the work with indoline outlined in Chapter 2. While outer-sphere redox mediation does little to guide the dehydrogenation reaction towards the formation of the desired N-ethylcarbazole product, it does lower the overall oxidation potential, which is desirable for maximizing overall fuel cell voltage. Increases in current are seen with the addition of a strong base, although the voltage shifts possible with indoline are not seen here, likely due to the lack of an acidic proton. Also demonstrated herein is the formation of dehydrogenated products using a platinum electrode.
3.2 Results and Discussion
3.2.1 Synthesis of N-Ethyldodecahydrocarbazole, N-Ethyloctahydrocarbazole and N-Ethyl-1,2,3,4-tetrahydrocarbazole
NEC-H12 is readily synthesized from N-ethylcarbazole (NEC) at high pressure (1000 psi hydrogen) and moderate temperature (150 °C) (Scheme 3.1). Although the bulk of the reaction occurs rapidly,1 allowing the reaction to proceed for 20 hours afforded nearly quantitative conversion to the fully hydrogenated product. The other components of the reaction mixture were identified by gas chromatography-mass spectrometry (GC-MS) as partially hydrogenated
38
intermediates: N-ethyloctahydrocarbazole (NEC-H8) and a very small amount of N-ethyltetrahydrocarbazole (NEC-H4). Both of these are known stable intermediates in the dehydrogenation of NEC-H12.1,2,4,5,10
Scheme 3.1: Synthesis of NEC-H12 from NEC.
NEC-H12 was isolated as a light yellow oil (80% yield), which turns dark orange over time under ambient conditions. It was stored at 0 °C in a foil-wrapped flask under a blanket of nitrogen in an attempt to prevent discoloration, although even darkened NEC-H12 appears to still be relatively pure by NMR. “Relatively pure”, in this case, means that there is no indication that NEC-H12 has decomposed into NEC-H8, NEC-H4, or NEC. It is also worth noting that NEC-H12 is a mixture of isomers, depending on the orientation of the hydrogens on the tertiary carbons in the middle ring (Figure 3.1),4,5 making it somewhat difficult to resolve individual hydrogen signals by NMR spectroscopy.
Figure 3.1: Some isomers of NEC-H12.
Although N-ethyloctahydrocarbazole (NEC-H8) was not isolated as the pure compound, a mixture of NEC-H12 and NEC-H8 was obtained as a byproduct from the synthesis of NEC-H12, with 8.5% of the mixture measured as NEC-H8 by integration of peaks in the 1H NMR spectrum.
N-ethyl-1,2,3,4-tetrahydrocarbazole (NEC-H4) can be prepared by base-catalyzed ethylation of the nitrogen of 1,2,3,4-tetrahydrocarbazole (Scheme 3.2). It, too, was isolated as a light yellow oil (83% yield), which was also sensitive to storage conditions. Although significant decomposition can be observed when it is stored under ambient conditions, it can be readily re-purified by column chromatography, and it is stable when stored in the dark under nitrogen.
39
Scheme 3.2: Synthesis of NEC-H4 from 1,2,3,4-tetrahydrocarbazole.
3.2.2 Direct Electrode Oxidation of N-Ethyldodecahydrocarbazole
Cyclic voltammetry studies with NEC-H12 indicate an oxidation potential of 0.48 V vs. Ag/Ag+ (Figure 3.2). The peak current shifts anodically with increasing scan rate, and no return peak is seen, even at scan rates up to 100 V/s (Figure 3.3). This indicates an irreversible process, potentially with a follow-up chemical reaction. Irreversibility in an electrochemical sense refers to the speed of the electrode reaction. In cyclic voltammetry, the data reflect a competition between mass transport to the electrode and the electrode reaction. If the electrode reaction is sufficiently fast, molecules will react immediately as they are brought to the electrode. This is the hallmark of a reversible process. Conversely, if the electrode reaction is slower than the rate of mass transport, the process is considered irreversible, leading to dependence of the peak potential on scan rate, decreased peak current, and a large separation between the forward and return peaks or no return peak at all. However, it is important to note that a reversible (fast) electron transfer followed by a sufficiently fast chemical reaction will also not display a return peak and will also display a shift of peak potential with scan rate.
Figure 3.2: 1 mM NEC-H12. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-5
0
5
10
15
20
25
30
-0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8
1 mM NEC-H12
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
40
Figure 3.3: 1 mM NEC-H12 at scan rates ranging from 25 mV/s to 100 V/s. 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode.
Another way of distinguishing between a reversible and an irreversible process is by looking at the shape of the oxidation or reduction peak. Peak shape is characterized by the potential difference between the peak current (Ep) and half the peak current (E1/2), that is, the potential at the peak current and the potential at which that current is at half its peak value. For a reversible process, this difference is defined as36
Equation 3.1
which becomes 57 mV at room temperature (298 K). Irreversible processes usually have a somewhat broader peak, with being greater than 57 mV, and are defined by36
Equation 3.2
which becomes 47.7/α mV at 298 K. α is a parameter known as the transfer coefficient that describes the transition state of an electron transfer reaction. If the transition state is equally similar to both the reduced and oxidized forms, α is 0.5. Greater similarity to the starting form results in α being less than 0.5, and vice versa. For NEC-H12, the breadth of the peak (Figure 3.2) as given by is 112.1 mV, indicating an irreversible process where α is 0.43.
The electrode reactivity of NEC-H12 was characterized further by cyclic voltammetry at a number of concentrations and scan rates. The peak current for an irreversible voltammetric process is defined as37
Equation 3.3
where n is the total number of electrons involved in the process, α is the transfer coefficient (defined above), A is the area of the electrode (cm2), D is the diffusion coefficient of the compound (cm2/s), C is the concentration of the compound in bulk solution (mol/mL), and v is the scan rate (V/s).
It is worth noting here that this type of analysis is only effective for irreversible processes where the overall number of electrons is given by n and the first step is a rate-determining irreversible heterogeneous one-electron transfer.37 More complex combinations of electron transfer steps and chemical steps cannot modeled very well by the equations that have been developed to describe simpler processes.
An appropriate range of scan rates was selected by examination of ferrocene at a wide range of scan rates; from 10 mV/s to 100 V/s. Ferrocene displays a reversible one-electron process. The ideal peak-to-peak separation for a reversible process is about 60 mV/s, with actual experimental values for ferrocene in acetonitrile (0.1 M NEt4BF4) being closer to 70 mV/s. The data for scan rates between 25 and 1000 mV/s conform to this peak-to-peak separation and are not overly corrupted by solution resistance. Above these scan rates, the electron transfer starts to appear only quasi-reversible, suggesting that the currents are high enough to contribute to significant iR drop38 (peak-to-peak separations range from 85 to 185 mV). Although the data at 10 mV/s appear reasonable with ferrocene, data at this scan rate often appear anomalous due to self-stirring of the solution, so it is also not included in analyses.
The true surface area of the electrode used was also calculated from ferrocene data. Although the geometric surface area is 0.071 cm2
, the true surface area was found to be 0.143 cm2, likely due to microscopic defects on the surface. For the calculations that follow, 0.143 cm2 was the value used for A.
Concentrations of NEC-H12 used ranged from 0.5 to 20 mM. Scan rates were varied between 25 and 1000 mV/s to give linear plots of ip vs. v1/2 (Figure 3.4); the slopes were then used to find n (Equation 3.3). The diffusion coefficient was estimated as 2.2 x 10−5 cm2/s, which is the average of the diffusion coefficients for two molecules similar in size to NEC-H12: anthracene (D = 2.23 x 10−5 cm2/s) and anthraquinone (D = 2.17 x 10−5 cm2/s).39 α was calculated for each run using
(Equation 3.2), and the values over all scan rates for a given concentration were averaged. The slopes and values for α and n found for each concentration are given in Table 3.1.
Although there is some variation in the values found for n, the data generally cluster around 0.5. This can be explained by considering the possible chemical reactions that may follow the electron transfer. A follow-up reaction that involves dimerization of the oxidized species with an unoxidized molecule, for example, would have an overall effective concentration that is half of
42
the actual concentration in solution, due to half the molecules being involved in follow-up reactions only (and not in the initial electron transfer). Similarly, a follow-up reaction that involves a molecule of NEC-H12 deprotonating another, oxidized, molecule would lead to a similar value for n. This is similar to what was seen in the voltammetric study of indoline (Chapter 2).
Figure 3.4: Peak current as a function of scan rate at concentrations of NEC-H12 between 0.5 mM and 20 mM. 0.1 M NEt4BF4 in acetonitrile, glassy carbon electrode.
[NEC-H12], mM slope of ip vs. v1/2 α n 0.5 0.40 x 10−4 0.39 0.64 1 0.67 x 10−4 0.47 0.49 2 1.38 x 10−4 0.48 0.50 5 3.16 x 10−4 0.38 0.51
10 5.96 x 10−4 0.36 0.50 20 10.7 x 10−4 0.36 0.44
Table 3.1: Values for α and n for concentrations of NEC-H12 between 0.5 mM and 20 mM.
Quantitative parameters have been established for probing the nature of chemical reactions following an electron transfer. One method involves examining the shift in peak potential (Ep) as a function of scan rate. The value of the shift per decade (factor of 10 change in scan rate) indicates a subset of chemical reactions that may follow after the electron transfer.
A plot of Ep vs. log(v) for 1 mM NEC-H12 indicates a shift of approximately 30 mV per decade, which is characteristic of EC (electrochemical-chemical), RSD-ECE(radical-substrate dimerization, electrochemical-chemical-electrochemical), or RSD-DISP1 (radical-substrate dimerization, disproportionation) mechanisms.40 An EC mechanism would involve an electron
0
2E-4
4E-4
6E-4
8E-4
1E-3
0.2 0.4 0.6 0.8 1
0.5 mM1 mM2 mM5 mM10 mM20 mM
(scan rate, V/s)1/2
anod
ic p
eak
curr
ent (
A)
43
transfer followed by a first-order homogeneous chemical reaction, while the two RSD mechanisms involve a second-order reaction between an oxidized (or reduced) molecule and an unoxidized (or unreduced) molecule following the electron transfer. An analysis of concentration dependence should distinguish between first- and second-order reactions, as the peak potential will shift only for the second-order reactions. However, no clear dependence of the peak potential on NEC-H12 concentration was seen. Nevertheless, the peak shape for an EC mechanism is predicted to be fairly sharp, and as a broad peak is seen here, an RSD mechanism is likely. This is also supported by the measurement of 0.5 for the number of electrons transferred. However, there is no way to distinguish between RSD-ECE and RSD-DISP1 using these data. It is worth noting here that an RSD mechanism could also point to a deprotonation reaction where a second molecule of NEC-H12 deprotonates the first, oxidized molecule, as the nitrogen of NEC-H12 is reasonably basic (pKa = 11.71 for protonated NEC-H12
41).
Another way of determining the value of n is to compare the peak currents for equimolar amounts of the molecule of interest (NEC-H12) and an electrochemically well-behaved compound such as ferrocene. The peak current for a reversible process is defined as37
Equation 3.4
which is similar to the definition of peak current for an irreversible process (Equation 3.3).
Figure 3.5: 1 mM ferrocene (orange, dashed) vs. 1 mM NEC-H12 (aqua). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-40
-20
0
20
40
60
-0.2 0 0.2 0.4 0.6
1 mM ferrocene1 mM NEC-H12
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
44
When equimolar solutions are examined with the same electrode, the comparison simplifies to
Equation 3.5
such that any errors in electrode area measurement are eliminated. Solutions of 1 mM NEC-H12 and 1 mM ferrocene were examined (Figure 3.5), and the resulting peak current data (ip) were plotted vs. v1/2. The slopes were used to calculate n for NEC-H12, given that Dferrocene is 2.24 x 10−5 cm2/s.42 Using these values, nα1/2 is 0.40. If α is 0.5, n is 0.63, which is consistent with the half-electron mechanism proposed earlier.
Rotating-disk voltammetry was also used to probe the behavior of NEC-H12. Figure 3.6 shows 1 mM NEC-H12 examined at a range of rotation rates, from 100 to 2500 rpm.
Figure 3.6: 1 mM NEC-H12 at multiple rotation rates ranging from 100 to 2500 rpm. 0.1 M NEt4BF4 in acetonitrile, 10 mV, glassy carbon rotating electrode.
The limiting current reached at a rotating electrode can be described by the Levich equation (Equation 3.6). However, Levich analysis is more appropriate for reversible processes.
Koutecký-Levich analysis uses the following equation (Equation 3.7), which accounts for the possibility of slow electron transfer at the electrode.
Equation 3.7
where i is the current (A), n the number of electrons transferred, F is Faraday’s constant (96485 C/mol), A is the electrode area (cm2), D is the diffusion coefficient of the analyte (cm2/s), ω is the rotation rate (rpm), ν is the kinematic viscosity (cm2/s), and C is the analyte concentration (mol/mL). The parameter iK is called the exchange current, and refers to the current exchanged at both electrodes at equilibrium, when net current for the system is zero. Koutecký-Levich analysis was applied to the rotating-disk data seen in Figure 3.6 (Figure 3.7).
Figure 3.7: Koutecký-Levich plot of data from Figure 3.6 at 0.6, 0.8 and 1 V.
The slopes in Figure 3.7 were used to calculate n. For our system, A is 0.126 cm2, D is 2.2 x 10−5 cm2/s, ν is 4.536 x 10−3 cm2/s,43 and C is 1 x 10−6 mol/mL, meaning n is 0.78.
Yet another method of determining n comes from a technique called step-potential voltammetry. The starting potential is chosen to be in a region where the compound of interest does not reduce or oxidize. The potential is then “stepped” to a region where reduction or oxidation will occur. The resulting current, assuming linear diffusion control (an appropriate assumption for a planar electrode), can be described using the Cottrell Equation (Equation 3.8):
Equation 3.8
0
5
10
15
20
25
30
0 0.02 0.04 0.06 0.08 0.1 0.12
0.6 V0.8 V1 V
1/(i,
mA
)
1/( , rpm)1/2
46
As before (Equation 3.6), n equals the number of electrons transferred, F is Faraday’s constant, A is the surface area of the electrode, C is the concentration of the bulk solution, D is the diffusion coefficient, and t is the time of the experiment (s). A plot of i vs. t1/2 should be linear, and relevant values such as n and D can be extracted from the slope.
Three different solutions were examined: 1 mM NEC-H12, 10 mM NEC-H12, and 1 mM ferrocene. Ferrocene was used to eliminate errors in electrode area measurement, as was done previously in Figure 3.5. Plots of i vs. t1/2 were generated for all t greater than 0.25 s. This was done to eliminate deviations from linearity seen in the initial data points due to a small delay in reaching the set potential. Slopes of these plots were used to calculate n, with comparison to the ferrocene data in order to eliminate A as a variable. Values for D used were as before: 2.2 x 10−5 cm2/s for NEC-H12 and 2.24 x 10−5 cm2/s for ferrocene.
[NEC-H12], mM Electrode Material n 1 glassy carbon 0.84 1 platinum 0.85 10 glassy carbon 0.64 10 platinum 0.62
Table 3.2: Comparison of n calculated from step-potential data for 1 mM and 10 mM NEC-H12 on both glassy carbon and platinum electrodes.
As Table 3.2 shows, the values for n for 1 mM NEC-H12 are slightly higher than were measured with any other technique. However, the values for 10 mM NEC-H12 are close to 0.5, as was measured before. If n is 0.5, this would imply that the following chemical step is rate-determining, whereas if n is 1, the initial electron transfer is rate-determining. A value in between may indicate that the rates of both are similar. Here, the full collection of data for NEC-H12 suggests that n is between 0.5 and 1, that the electron transfer and follow-up chemical reaction have similar rates, and that the reaction following initial electron transfer is second-order (either a dimerization or a deprotonation).
NEC-H8 and NEC-H4 were also examined by cyclic voltammetry. The NEC-H8/NEC-H12 mixture shows a slight cathodic shift in oxidation potential, implying that NEC-H8 is easier to oxidize than NEC-H12 (Figure 3.8). However, NEC-H4 displays two oxidation peaks, at 0.55 and 0.71 V vs. Ag/Ag+, meaning it is more difficult to oxidize than NEC-H12 (Figure 3.2). NEC-H4 also displays a reduction shoulder around 0.66 V, implying that the second oxidation process is reversible (∆Ep is 47 mV). This is consistent with a previous study of N-methyl-1,2,3,4-tetrahydrocarbazole, where the two oxidation peaks seen were attributed to formation of the radical cation and then reversible oxidation to the dication, ultimately leading to a dimeric product.44
47
Figure 3.8: 10 mM NEC-H12 (aqua) and 10 mM of an 8.5% mixture of NEC-H8 in NEC-H12 (blue, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 3.9: 1 mM NEC-H4. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-50
0
50
100
150
200
250
300
-0.4 -0.2 0 0.2 0.4 0.6 0.8 1
10 mM NEC-H1210 mM NEC-H8/NEC-H12 mixture
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
-10
0
10
20
30
40
50
-0.4 -0.2 0 0.2 0.4 0.6 0.8 1
1 mM NEC-H4
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
48
3.2.3 Effect of Bases on N-Ethyldodecahydrocarbazole Voltammetry
NEC-H12 oxidation was examined in the presence of both weak and strong bases. Since it was determined that the chemical reaction following the initial oxidation of NEC-H12 involves another molecule of NEC-H12, it is possible that the chemical reaction following the electron transfer could be a deprotonation. Thus, oxidizing NEC-H12 in the presence of base could both promote deprotonation as the chemical step. This could also allow for more electron transfer, resulting in a larger value for n.
An initial screening of bases with protonated pKa values ranging from 14.2 to 23.3 (in acetonitrile) showed that unlike indoline, the oxidation potential for NEC-H12 does not shift with increasing base strength (Figure 3.10, Table 3.3), likely due to the fact that it does not possess an acidic proton.
Figure 3.10: 10 mM NEC-H12, with no base (green), 100 mM imidazole (aqua), 100 mM 2-ethyl-4-methylimidazole (blue), 100 mM triethylamine (purple), 100 mM diethylenetriamine (pink), or 100 mM 1,1,3,3-tetramethylguanidine (red). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Base Onset Potential (V)(a) pKa in Water(b) pKa in
Table 3.3: Summary of changes to the onset potential for NEC-H12 oxidation in the presence of base. (a) Defined as the point at which current density reaches approximately 10 μA. (b) pKa for the protonated, conjugate acid form.
-500
0
500
1000
1500
2000
2500
3000
-0.2 0 0.2 0.4 0.6 0.8
10 mM NEC-H12+ 100 mM imidazole+ 100 mM 2-ethyl-4-methylimidazole+ 100 mM triethylamine+ 100 mM diethylenetriamine+ 100 mM 1,1,3,3-tetramethylguanidine
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
49
At least part of the additional current seen when bases are added to NEC-H12 is due to the oxidation of the bases themselves at the electrode (Figure 3.11). Because NEC-H12 oxidizes at a more positive potential than indoline does, more interference is seen from base oxidation in the case of NEC-H12. In particular, the weaker bases (imidazole and 2-ethyl-4-methylimidazole) produce essentially no additional current. The larger currents seen in Figure 3.10 for the weaker bases (aqua and blue lines) are almost entirely due to the addition of currents from NEC-H12 and the bases at the electrode. However, analysis of the additional current produced in addition to the added electrode currents of NEC-H12 and base for the stronger bases does show a dependence on pKa: diethylenetriamine, triethylamine and 1,1,3,3-tetramethylguanidine produce additional currents of 50, 300, and 800 μA, respectively, beyond what would be anticipated for NEC-H12 and the bases independently. However, although the currents produced with strong bases are more than expected from simple simultaneous oxidation of fuel and base, the shift in oxidation onset potential is fairly small, meaning that use of redox catalysts with more negative oxidation potentials is not feasible with NEC-H12.
Figure 3.11: Oxidation potentials for 100 mM imidazole (orange), 100 mM 2-ethyl-4-methylimidazole (red), 100 mM triethylamine (green), 100 mM diethylenetriamine (blue), or 100 mM 1,1,3,3-tetramethylguanidine (purple). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
10 mM NEC-H12 was examined increasing amounts of the strongest base, 1,1,3,3-tetramethylguanidine, with concentrations ranging from 5 mM to 1 M. The data were then analyzed for electron count and potential shift per decade.
When increasing amounts of base are used, an increase in current at the NEC-H12 oxidation peak potential is seen for up to 10 equivalents of base (Figure 3.12).47 The value of n increases as well, reaching a maximum of two electrons. This supports the theory that the follow-up chemical reaction here is a deprotonation, as increased amounts of base would be anticipated to increase the rate of deprotonation, allowing for additional electron transfers on the voltammetry timescale
-500
0
500
1000
1500
2000
2500
-0.5 0 0.5 1
100 mM imidazole100 mM 2-ethyl-4-methylimidazole100 mM triethylamine100 mM diethylenetriamine100 mM 1,1,3,3-tetramethylguanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
50
beyond the half electron measured for NEC-H12 oxidation alone. If n is 2, this suggests an ECE mechanism, where the initial electron transfer and following chemical reaction are then succeeded by a second electron transfer. This might also explain the limitation in peak current, as a point is reached where the second electron transfer becomes rate-limiting.
Figure 3.12: 10 mM NEC-H12 with no base (orange), 5 mM (pink), 20 mM (purple), 50 mM (dark blue), 200 mM (light blue), or 500 mM (aqua) 1,1,3,3-tetramethylguanidine. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Table 3.4: Values for α, n, and Epa vs. log(v) for 10 mM NEC-H12 and concentrations of 1,1,3,3-tetramethylguanidine between 0.5 mM and 100 mM.
Table 3.4 shows calculated transfer coefficients (α) and electron counts (n) for all concentrations of base examined. Although the peak current does not increase substantially above the value for 100 mM 1,1,3,3-tetramethylguanidine, the transfer coefficient continues to decrease, indicating
0
200
400
600
800
1000
1200
-0.2 0 0.2 0.4 0.6
10 mM NEC-H12+ 5 mM tetramethylguanidine+ 20 mM tetramethylguanidine+ 50 mM tetramethylguanidine+ 200 mM tetramethylguanidine+ 500 mM tetramethylguanidine
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
51
that the further addition of base causes the peak to broaden. This is why the value for n continues to increase slightly even though the peak current does not. Unfortunately, there was no clear trend seen in the shift of the peak potential with scan rate, meaning that detailed analysis of the reactions following the initial electron transfer is not possible here. We might anticipate that as the chemical reaction becomes faster, the initial electron transfer becomes rate-limiting, leading to a larger shift in peak potential, but these data are inconclusive. Regardless, there is clear evidence of base catalysis of NEC-H12 oxidation, suggesting that the current obtained from NEC-H12 can be enhanced by the addition of a strong base.
3.2.4 Redox Catalysis of N-Ethyldodecahydrocarbazole Oxidation
NEC-H12 redox catalysis is unexplored in the literature, and no outer-sphere electrochemical catalysis has been reported. However, redox catalysis proved to be quite successful with indoline (Chapter 2), and was thus of interest for NEC-H12 oxidation as well. Ferrocene was examined as a possible redox catalyst for NEC-H12, as it oxidizes in a relevant potential range, and was found to be competent at catalyzing the oxidation of NEC-H12, as can be seen from the increase in current at the ferrocene peak potential in Figure 3.13.
Figure 3.13: 1 mM ferrocene, without (orange) and with (blue) 10 mM NEC-H12. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
The addition of 1,1,3,3-tetramethylguanidine increases the catalytic current even further (Figure 3.14). However, this current is not all entirely due to the combination of fuel, catalyst, and base, as ferrocene can also catalyze the oxidation of 1,1,3,3-tetramethylguanidine (Figure 3.15). Nevertheless, about two-thirds of the current increase seen in Figure 3.14 over the current from ferrocene alone is due to the presence of NEC-H12, as can be seen by comparing the current for ferrocene and 1,1,3,3-tetramethylguanidine (blue) with the current from all three (green, dashed) in Figure 3.16.
-20
0
20
40
60
80
100
120
-0.6 -0.4 -0.2 0 0.2 0.4
1 mM ferrocene1 mM ferrocene, 10 mM NEC-H12
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
52
Figure 3.14: 1 mM ferrocene, without (orange) and with (purple) 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine. 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine alone (gold, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 3.15: 1 mM ferrocene with no base (orange) and 1 mM (light aqua), 10 mM (aqua) and 100 mM (blue) 1,1,3,3-tetramethylguanidine, showing the catalysis of 1,1,3,3-tetramethylguanidine oxidation by ferrocene. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-50
0
50
100
150
200
250
300
350
-0.6 -0.4 -0.2 0 0.2 0.4
1 mM ferrocene+ 10 mM NEC-H12, 100 mM tetramethylguanidine10 mM NEC-H12, 100 mM tetramethylguanidine alone
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
-20
0
20
40
60
80
100
-0.4 -0.3 -0.2 -0.1 0 0.1 0.2 0.3
1 mM ferrocene+ 1 mM tetramethylguanidine+ 10 mM tetramethylguanidine+ 100 mM tetramethylguanidine
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
53
Figure 3.16: 1 mM ferrocene, alone (orange), with 100 mM 1,1,3,3-tetramethylguanidine (blue) and with 10 mM NEC-H12 and 100 mM 1,1,3,3-tetramethylguanidine (green, dashed). 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 3.17: 1 mM ferrocene and 10 mM NEC-H12 alone (pink) and with concentrations of 1,1,3,3-tetramethylguandine ranging from 5 to 1000 mM. 0.1 M NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-50
0
50
100
150
200
-0.4 -0.3 -0.2 -0.1 0 0.1 0.2 0.3 0.4
1 mM ferrocene+ 100 mM tetramethylguanidine+ 10 mM NEC-H12, 100 mM tetramethylguanidine
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
-20
0
20
40
60
80
100
120
140
-0.6 -0.4 -0.2 0 0.2 0.4
no tetramethylguanidine5 mM10 mM20 mM50 mM100 mM200 mM500 mM1000 mM
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
54
Ferrocene and NEC-H12 were also examined with increasing amounts of 1,1,3,3-tetramethylguanidine (Figure 3.17). Although the current did increase with increased base, the increases were not dramatic, with a maximum current at the ferrocene peak potential of 89 μA (1000 mM 1,1,3,3-tetramethylguanidine, 100 mV/s), an increase of 47 μA over ferrocene alone (42 μA, 100 mV/s).
Ferrocene is also able to catalyze the oxidation of the NEC-H8/NEC-H12 mixture, with catalytic currents even greater than for NEC-H12 alone. This is expected, given the more negative oxidation potential of NEC-H8. However, ferrocene does not have any catalytic effect on NEC-H4, even though the oxidation potential for NEC-H4 is only 0.07 V more positive than that of NEC-H12. Voltammetry shows no increased current at the ferrocene peak potential, and controlled-potential electrolysis of NEC-H4 with ferrocene resulted in very little current and no consumption of NEC-H4. When 1,1,3,3-tetramethylguanidine was added, the primary products formed seemed to be from the catalytic oxidation of 1,1,3,3-tetramethylguanidine, as the products were very water-soluble and the amount of NEC-H4 was again essentially unchanged. Overall, this implies that an effective catalytic system for initial dehydrogenation of NEC-H12 may not be effective for complete dehydrogenation to NEC.
3.2.5 Bulk Electrolysis and Product Analysis of N-Ethyldodecahydrocarbazole
A single previous study27 has examined NEC-H12 in a fuel cell environment, demonstrating an open-circuit voltage of 1.4 V, which is higher than the theoretical limit for PEM fuel cells (1.23 V). This is possible because the use of fuels besides hydrogen directly in PEM fuel cells allows for anodic reaction potentials more negative than that of hydrogen oxidation. Reproducible currents were seen when the cell was run at 120 °C, with small amounts of N-ethyldecahydrocarbazole and NEC-H8 detected, both known intermediates in the non-oxidative dehydrogenation process.2
The work presented here was intended to expand on the previously published work by evaluating the effects of redox catalysts and bases on the bulk electrolysis process, both in terms of current produced and in terms of products formed. Analysis of products was performed by both GC-MS and NMR spectroscopy. NEC-H12, NEC, and the two most common intermediates (NEC-H4 and NEC-H8) all have characteristic NMR shifts and mass spectrometry peak patterns, which are detailed below (Table 3.5, Figure 3.18).
Compound 1H NMR Shift(s) of Ethyl CH2 (δ in CDCl3)
Mass Spectrometry Peaks (m/z)
NEC-H12 m, 2.43, 1H; m, 2.77, 1H 207, 192, 178, 164 NEC-H8 q, 3.68, 2H 203, 188, 174, 160 NEC-H4 q, 4.08, 2H 199, 184, 170, 156
NEC q, 4.39, 2H 195, 180, 166, 152 Table 3.5: Characteristic 1H NMR shifts of the ethyl CH2 group and major mass spectrometry peaks for N-ethylcarbazoles in varying states of hydrogenation.
55
Figure 3.18: Characteristic 1H NMR shifts of the ethyl CH2 group for N-ethylcarbazoles in varying states of hydrogenation (L-R: NEC, NEC-H4, NEC-H8, NEC-H12). The resonances for NEC, NEC-H4 and NEC-H8 appear as a quartet from coupling to the ethyl CH3 group.
Our initial studies looked at NEC-H12 oxidation in acetonitrile, using a carbon felt electrode. Controlled-potential electrolyses were performed for 20 hours and the resulting charge measured (Table 3.6). Although the addition of 1,1,3,3-tetramethylguanidine noticeably increased the charge, ferrocene addition did not. However, the addition of ferrocene did allow for the lowering of the oxidation potential by over 0.4 V, which is desirable for maximizing overall fuel cell voltage. In the case of both ferrocene and 1,1,3,3-tetramethylguanidine being added, the potential was lowered by nearly 0.3 V and the charge was increased, although it is possible that this is partly due to oxidation of 1,1,3,3-tetramethylguanidine by ferrocene.
Table 3.6: Comparison of controlled-potential electrolyses performed with 10 mM NEC-H12 using a carbon felt electrode in acetonitrile.
In the electrolysis of NEC-H12 alone and in both runs with 1,1,3,3-tetramethylguanidine, no products were detected by GC-MS, and the NMR spectra showed complex mixtures, making it difficult to determine the identity of any of the individual components. However, in the electrolysis of NEC-H12 in the presence of ferrocene, two peaks were detected by mass spectrometry that had the same pattern as seen in the other carbazoles (M, M−15, M−29, M−43), and molecular ion masses (M) of 217 and 219. These peaks are attributed to hydroxylated side products formed by reaction with trace water in the solution during electrolysis, since those molecular weights correspond to N-ethylhexahydrocarbazole and NEC-H8 plus oxygen (Figure
56
3.19). This is further supported by the appearance of a broad singlet at 5.43 ppm in both NMR spectra, characteristic of the labile proton of a hydroxyl group. This suggests that NEC-H12 will dehydrogenate in the presence of ferrocene.
Figure 3.19: Tentative products from controlled-potential electrolysis of NEC-H12 in the presence of ferrocene and controlled-current electrolysis of NEC-H12 by two and four electrons. Hydroxylation postitions are arbitrary.
Subsequent attempts to react NEC-H12 with exactly two and four electrons per molecule using controlled current electrolysis also produced small amounts of side products, including 1,1’-bicyclohexyl, hydroxylated NEC-H8 and N-ethylhexahydrocarbazole, and a small amount of NEC-H8, as well as some other products that could not be determined (Figure 3.19). One or more new products were formed after controlled current electrolysis to react NEC-H12 by the full 12 electrons, as seen by the appearance of peaks between 3.0 and 3.6 ppm and between 7.0 and 7.5 ppm in the 1H NMR spectrum. However, none of the components could be identified using GC-MS, except for the small amount of NEC-H12 that remained.
Although this is the first attempt to electrochemically dehydrogenate NEC-H12, electrochemical dehydrogenation of tetrahydrocarbazoles has been studied before. One paper48 reported that a tetrahydrocarbazole derivative tends to dimerize in acetonitrile but will form the fully dehydrogenated product in dichloromethane using a platinum electrode. Although the experiment in that report was not run in dichloromethane with a carbon electrode, the use of a carbon electrode in acetonitrile produced various dimers, with no monomeric tetrahydrocarbazole or carbazole detected. The same experiment with a platinum electrode in acetonitrile also led to no monomeric tetrahydrocarbazole or carbazole being detected (although dimers were not formed in this case).
DCM, 10 mM tetramethylguanidine platinum 0.70 84 14 14
MeCN platinum 0.45 120 13 13 Table 3.7: Comparison of controlled-potential electrolyses performed with 1 mM NEC-H12. (a) Normalized for electrode area, as determined by step-potential voltammetry of a ferrocene solution. (b) Electrolysis performed at 0.55 V for the first 24 hours and 0.85 V for the second 24 hours.
It was hypothesized that the use of a platinum electrode and/or dichloromethane as solvent in the oxidation of NEC-H12 might produce a significant amount of dehydrogenated N-ethylcarbazole, instead of the small amounts of side products seen with carbon electrodes. In addition to the specific support from literature on tetrahydrocarbazole electrooxidation, platinum is one of the
major components of a commercial fuel cell anode and is well-known as a non-electrochemical hydrogenation and dehydrogenation catalyst.
The platinum electrode used had a smaller surface area than those of the carbon felt electrodes; thus, these electrolyses had to be run longer in order to pass enough charge to see any change, and even then the amount of charge was much smaller than for a carbon electrode. To correct for this, the electrode area was measured using step-potential voltammetry in the presence of ferrocene. While the exact electrode area could not be determined due to the complexities of mass transport to a large, uneven surface, this provided a means of comparing charge at the platinum electrode with charge at the carbon felt electrodes. The “normalized charge” listed in Table 3.7 is corrected for surface area and provides a true comparison of charge. This allows us to see that platinum is actually a more efficient electrode material, allowing for the passage of more charge in the same amount of time (when corrected for surface area).
Although no clear products were detected for three of the electrolyses, the electrolysis with NEC-H12 in dichloromethane at a platinum electrode (no base) showed clear evidence of dehydrogenation. NEC-H12 itself has no 1H NMR resonances downfield of 3 ppm, but an NMR spectrum of the electrolysis products indicates multiple peaks between 4 and 7 ppm (Figure 3.20). In particular, the peaks in the aromatic region have a pattern that resembles the aromatic resonances seen for NEC and NEC-H4, although the exact shifts do not correspond to either. The resonances seen between 2.3 and 2.6 ppm, however, are likely due to unreacted NEC-H12.
Figure 3.20: 1H NMR spectrum of 1 mM NEC-H12 after 102 hours of electrolysis at a platinum electrode in dichloromethane. Green circles indicate one product, likely a long-chain alkene, while blue stars denote another, unidentified product. Resonances between 2.3 and 2.6 ppm are attributed to unreacted NEC-H12.
Upon further analysis, it was determined that there are two main products seen in the NMR spectrum. One of these, marked with green circles, likely corresponds to a long-chain alkene (or a mixture thereof). GC-MS analysis indicated formation of several long-chain alkenes: 1-dodecene, 1-tetradecene, 1-hexadecene, 1-octadecene, and 1-docosene. While these are easily separated on a GC column due to differences in molecular weight, they all have essentially the same NMR resonances in the 2–6 ppm region, and thus may appear as one compound here. The other compound is not yet identified. GC-MS analysis suggests that it may be a substituted phenol, but that does not match the NMR. Analysis is therefore ongoing. The most interesting
58
feature of these two products, however, is that they are both byproducts of NEC-H4 synthesis, and were identified as such by NMR spectra of late column fractions from NEC-H4 purification. It is suspected that these compounds may be decomposition products from NEC-H4 being kept under ambient conditions, and that NEC-H4 is being formed in solution but decomposing due to the long electrolysis time. If this is in fact the case, NEC-H12 is dehydrogenating to form NEC-H4 at a platinum electrode in dichloromethane, which is a very big step towards being able to use NEC-H12 as a fuel in a virtual hydrogen storage system.
3.3 Conclusions
The electro-oxidation of N-ethyldodecahydrocarbazole (NEC-H12), N-ethyloctahydrocarbazole (NEC-H8), and N-ethyl-1,2,3,4-tetrahydrocarbazole (NEC-H4) has been studied. It has been demonstrated that N-ethyldodecahydrocarbazole oxidative current can be increased by the addition of strong base, and that its oxidation potential can be lowered by use of ferrocene as a redox mediator. It has also been shown that partial dehydrogenation can be effected with the use of ferrocene in acetonitrile at a carbon felt electrode, or at a platinum electrode in dichloromethane. In particular, electrolysis in dichloromethane using a platinum electrode showed substantial formation of dehydrogenated products, which is a key finding for fully developing N-ethyldodecahydrocarbazole as a fuel for direct electrochemical dehydrogenation in a PEM fuel cell.
3.4 Experimental Methods and Materials
Ferrocene was obtained commercially and purified by sublimation before use. Tetra-n-butylammonium tetrafluoroborate (NBu4BF4) was purchased from Oakwood Products, Inc., and was purified by dissolving in ethyl acetate, precipitating with pentane, filtering, and drying in vacuo overnight. All other reagents used were obtained commercially and used as received. Solvents for electrochemical experiments were performed using high-purity acetonitrile or dichloromethane, dried and degassed using a Phoenix solvent drying system commercially available from JC Meyer Solvent Systems.
Synthesis of N-ethyldodecahydrocarbazole: N-ethyldodecahydrocarbazole was synthesized as reported in the literature.49 N-ethylcarbazole (22.48 g, 115 mmol) was placed in a pressure reactor (Parr Instruments Company) with 5% Ru/C catalyst (10.8 g) and anhydrous heptane (450 mL). The reactor was flushed 3 times with hydrogen and then pressurized to 1000 psi hydrogen. The pressurized reactor was placed into an oil bath and the bath was heated to 150 °C while stirring. After 20 hours, the vessel was removed from the bath, allowed to cool, and depressurized. The reaction mixture was filtered through Celite with a mixture of heptane and hexanes and the solvent removed by rotary evaporation. The resulting oil was purified by column chromatography, using alumina as stationary phase and a gradient elution of 0 to 25% ethyl acetate in hexanes. Both NEC-H12 and a mixture of 8.5% NEC-H8 in NEC-H12 were isolated from the column fractions. 1H NMR (300 MHz, CDCl3): δ 2.63–2.75 (m, 1H), 2.46–2.59 (m, 1H), 2.30–2.41 (m, 1H), 1.63–1.95 (m, 4H), 1.47–1.60 (m, 5H), 1.34–1.42 (m, 1H), 1.31–1.32 (m, 2H), 1.23–1.25 (m, 2H), 1.11–1.17 (m, 4H), 1.01 (t, J = 7.2 Hz, 3H), 0.82–0.87 (m, 1H). 13C{1H} NMR (CDCl3): 70.01, 62.02, 47.95, 45.50, 40.95, 32.16, 31.58, 28.86, 25.72, 24.98, 24.04, 22.75, 20.89, 13.70. 80% yield (NEC-H12 alone).
59
Synthesis of N-ethyl-1,2,3,4-tetrahydrocarbazole: N-ethyl-1,2,3,4-tetrahydrocarbazole was synthesized using a procedure modified from a synthesis of N-ethylindole.50 1,2,3,4-tetrahydrocarbazole (8.562 g, 50 mmol) and potassium hydroxide (5.611 g, 100 mmol) were placed into a 500 mL round-bottomed flask. Dimethylsulfoxide (DMSO, 100 mL) was added and the resulting mixture stirred, forming a dark brown solution. Ethyl bromide (7.5 mL, 100 mmol) was added along with 200 mL additional DMSO, forming a lighter orange-brown solution. The reaction mixture was allowed to stir for 5 hours, at which point 100 mL water were added to the reaction flask, causing the reaction mixture to become golden-yellow and opaque. The reaction mixture was extracted with three subsequent portions of diethyl ether. The ether portions were then combined and washed 3 times with water to ensure full removal of the DMSO, washed once with brine, dried over magnesium sulfate, filtered, and concentrated to an orange liquid. If deemed necessary after an initial 1H NMR spectrum was acquired, the product was purified by column chromatography, using silica as stationary phase and hexanes as eluent, followed by re-extraction with acetonitrile. 1H NMR (400 MHz, CDCl3): δ 7.47 (d, 1H), 7.28 (d, 1H), 7.10–7.18 (m, 1H), 7.02–7.10 (m, 1H), 4.08 (q, J = 7.2 Hz, 2H), 2.73 (m, 4H), 1.87–1.96 (m, 4H), 1.32 (t, J = 7.2 Hz, 3H). UV-Vis (nm): 230, 285, 292. 83% yield.
NMR spectra were obtained at ambient temperature in CDCl3 or C6D6 using Bruker AV-300, AVB-400, AVQ-400, or DRX-500 spectrometers at the University of California, Berkeley, and referenced to the residual solvent peak. UV-visible spectra were determined with a Varian Cary 50 UV-Vis spectrophotometer. GC-MS analyses were performed using a Supelco SPB-5 column on an Agilent 5975 Series GC/MSD. o-Terphenyl was added to the GC samples as a standard for calibration.
Cyclic voltammetry and step-potential voltammetry experiments were performed using a Gamry 600 Reference potentiostat. Rotating-disk voltammetry was performed using an ALS RRDE-3A Rotator and a BioLogic Science Instruments VSP bipotentiostat. All voltammetry experiments used 0.1 M tetraethylammonium tetrafluoroborate (NEt4BF4) in acetonitrile as the supporting electrolyte. Voltammetry solutions were deoxygenated before use and kept under a blanket of inert gas (either nitrogen or argon) during analysis. A three-electrode cell was used, with a platinum mesh counter electrode and a silver ion non-aqueous reference electrode. The reference electrode was prepared by immersing a silver wire in a 0.01 M solution of silver nitrate in supporting electrolyte, and was separated from the bulk solution by a Vycor frit. Potentials are reported vs. Ag/Ag+ (0.01 AgNO3), which is 0.548 V vs. NHE and −0.087 V vs. ferrocene (Fc).51 For cyclic voltammetry and step-potential voltammetry, a glassy carbon working electrode (3 mm diameter) was used. For rotating-disk voltammetry, a working electrode with a glassy carbon disk (4 mm diameter) and a platinum ring was used, although only the disk portion was used in the experiment.
Controlled-potential and controlled-current electrolyses were performed using a PAR 276 potentiostat, using either 0.1 M NEt4BF4 in acetonitrile or 0.1 M NBu4BF4 in dichloromethane. The working electrode was composed of either platinum mesh or carbon felt (99%, Alfa Aesar) rinsed with dilute hydrochloric acid and isopropyl alcohol prior to use. The counter electrode was platinum mesh, separated from the bulk solution by a porous frit. The reference electrode was prepared with a silver wire in a 0.01 M solution of AgNO3 in acetonitrile, contained in a glass tube with a ceramic junction. For controlled-potential electrolysis, potentials were chosen to be just negative of the oxidation peak of the catalyst or substrate (in the absence of catalyst)
60
after an initial cyclic voltammogram of the solution using a platinum disk working electrode. Electrolysis products were identified using NMR spectroscopy and GC-MS analysis.
Correction factors for comparing electrodes for bulk electrolysis were estimated by step-potential voltammetry of a 1 mM ferrocene solution. The total charge was measured under identical conditions for the platinum electrode and a series of carbon felt electrodes of different sizes. These charges were normalized to platinum by dividing by the total charge for platinum, so that QPt = 1. A plot of surface area vs. charge was generated for the carbon felt electrodes, giving a linear fit of Q = (0.6264)A + 1.456. This allowed for the calculation of normalized Q for any size of carbon felt electrode. The charge for each the bulk electrolysis experiment was then divided by the corresponding charge for ferrocene.
3.5 References
1. Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H. Hydrogen Storage by Reversible Hydrogenation of Pi-Conjugated Substrates. US Pat. 7,101,530 2006.
2. Sotoodeh, F.; Zhao, L.; Smith, K. J. Kinetics of H2 Recovery from Dodecahydro-N-Ethylcarbazole over a Supported Pd Catalyst. Appl. Catal., A 2009, 362, 155–162.
3. Wang, Z.; Tonks, I.; Belli, J.; Jensen, C. M. Dehydrogenation of N-Ethylperhydrocarbazole Catalyzed by PCP Pincer Iridium Complexes: Evaluation of a Homogeneous Hydrogen Storage System. J. Organomet. Chem. 2009, 694, 2854–2857.
4. Morawa Eblagon, K.; Rentsch, D.; Friedrichs, O.; Remhof, A.; Zuettel, A.; Ramirez-Cuesta, A. J.; Tsang, S. C. Hydrogenation of 9-Ethylcarbazole as a Prototype of a Liquid Hydrogen Carrier. Int J Hydrog. Energy 2010, 35, 11609–11621.
5. Morawa Eblagon, K.; Tam, K.; Yu, K. M. K.; Zhao, S.-L.; Gong, X.-Q.; He, H.; Ye, L.; Wang, L.-C.; Ramirez-Cuesta, A. J.; Tsang, S. C. Study of Catalytic Sites on Ruthenium For Hydrogenation of N-Ethylcarbazole: Implications of Hydrogen Storage via Reversible Catalytic Hydrogenation. J. Phys. Chem. C 2010, 114, 9720–9730.
6. Sotoodeh, F.; Smith, K. J. Kinetics of Hydrogen Uptake and Release from Heteroaromatic Compounds for Hydrogen Storage. Ind. Eng. Chem. Res. 2010, 49, 1018–1026.
7. Sobota, M.; Nikiforidis, I.; Amende, M.; Zanon, B. S.; Staudt, T.; Hoefert, O.; Lykhach, Y.; Papp, C.; Hieringer, W.; Laurin, M.; et al. Dehydrogenation of Dodecahydro-N-Ethylcarbazole on Pd/Al2O3 Model Catalysts. Chem. Eur. J. 2011, 17, 11542–11552.
8. Sotoodeh, F.; Smith, K. J. Structure Sensitivity of Dodecahydro-N-Ethylcarbazole Dehydrogenation over Pd Catalysts. J. Catal. 2011, 279, 36–47.
9. Sotoodeh, F.; Huber, B. J. M.; Smith, K. J. Dehydrogenation Kinetics and Catalysis of Organic Heteroaromatics for Hydrogen Storage. Int. J. Hydrogen Energy 2012, 37, 2715–2722.
10. Sotoodeh, F.; Huber, B. J. M.; Smith, K. J. The Effect of the N Atom on the Dehydrogenation of Heterocycles Used for Hydrogen Storage. Appl. Catal., A 2012, 419-420, 67–72.
11. Wan, C.; An, Y.; Xu, G.; Kong, W. Study of Catalytic Hydrogenation of N-Ethylcarbazole over Ruthenium Catalyst. Int J Hydrog. Energy 2012, 37, 13092–13096.
12. Amende, M.; Schernich, S.; Sobota, M.; Nikiforidis, I.; Hieringer, W.; Assenbaum, D.; Gleichweit, C.; Drescher, H.-J.; Papp, C.; Steinrück, H.-P.; Görling, A.; Wasserscheid, P.; Laurin, M.; Libuda, J. Dehydrogenation Mechanism of Liquid Organic Hydrogen
61
Carriers: Dodecahydro-N-Ethylcarbazole on Pd(111). Chem. Eur. J. 2013, 19, 10854–10865.
13. Gleichweit, C.; Amende, M.; Schernich, S.; Zhao, W.; Lorenz, M. P. A.; Höfert, O.; Brückner, N.; Wasserscheid, P.; Libuda, J.; Steinrück, H.-P.; Papp, C. Dehydrogenation of Dodecahydro-N-Ethylcarbazole on Pt(111). ChemSusChem 2013, 6, 974–977.
14. Sotoodeh, F.; Smith, K. J. Analysis of H2 Release from Organic Polycyclics over Pd Catalysts Using DFT. J. Phys. Chem. C 2013, 117, 194–204.
15. Sotoodeh, F.; Smith, K. J. An Overview of the Kinetics and Catalysis of Hydrogen Storage on Organic Liquids. Can. J. Chem. Eng. 2013, 9999, 1–14.
16. Amende, M.; Gleichweit, C.; Werner, K.; Schernich, S.; Zhao, W.; Lorenz, M. P. A.; Höfert, O.; Papp, C.; Koch, M.; Wasserscheid, P.; Laurin, M.; Steinrück, H.-P.; Libuda, J. Model Catalytic Studies of Liquid Organic Hydrogen Carriers: Dehydrogenation and Decomposition Mechanisms of Dodecahydro-N-Ethyl-Carbazole on Pt(111). ACS Catal. 2014, 4, 657–665.
17. Amende, M.; Gleichweit, C.; Schernich, S.; Höfert, O.; Lorenz, M. P. A.; Zhao, W.; Koch, M.; Obesser, K.; Papp, C.; Wasserscheid, P.; Steinrück, H.-P.; Libuda, J. Size and Structure Effects Controlling the Stability of the Liquid Organic Hydrogen Carrier Dodecahydro-N-Ethylcarbazole during Dehydrogenation over Pt Model Catalysts. J. Phys. Chem. Lett. 2014, 5, 1498–1504.
18. Brayton, D. F.; Jensen, C. M. Solvent Free Selective Dehydrogenation of Indolic and Carbazolic Molecules with an Iridium Pincer Catalyst. Chem. Commun. 2014, 50, 5987–5989.
19. Brayton, D. F.; Beaumont, P. R.; Fukushima, E. Y.; Sartain, H. T.; Morales-Morales, D.; Jensen, C. M. Synthesis, Characterization, and Dehydrogenation Activity of an Iridium Arsenic Based Pincer Catalyst. Organometallics 2014, 33, 5198–5202.
20. Dutta Chowdhury, A.; Weding, N.; Julis, J.; Franke, R.; Jackstell, R.; Beller, M. Towards a Practical Development of Light-Driven Acceptorless Alkane Dehydrogenation. Angew. Chem. Int. Ed. 2014, 53, 6477–6481.
21. Morawa Eblagon, K.; Tsang, S. C. E. Structure-Reactivity Relationship in Catalytic Hydrogenation of Heterocyclic Compounds over Ruthenium Black-Part A: Effect of Substitution of Pyrrole Ring and Side Chain in N-Heterocycles. Appl. Catal., B 2014, 160–161, 22–34.
22. Gleichweit, C.; Amende, M.; Bauer, U.; Schernich, S.; Höfert, O.; Lorenz, M. P. A.; Zhao, W.; Müller, M.; Koch, M.; Bachmann, P.; Wasserscheid, P.; Libuda, J; Steinrück, H.-P.; Papp, C. Alkyl Chain Length-Dependent Surface Reaction of Dodecahydro-N-Alkylcarbazoles on Pt Model Catalysts. J. Chem. Phys. 2014, 140, 204711-1–204711-9.
24. Yang, M.; Dong, Y.; Fei, S.; Ke, H.; Cheng, H. A Comparative Study of Catalytic Dehydrogenation of Perhydro-N-Ethylcarbazole over Noble Metal Catalysts. Int. J. Hydrogen Energy 2014, 39, 18976–18983.
25. Dutta Chowdhury, A.; Julis, J.; Grabow, K.; Hannebauer, B.; Bentrup, U.; Adam, M.; Franke, R.; Jackstell, R.; Beller, M. Photocatalytic Acceptorless Alkane Dehydrogenation: Scope, Mechanism, and Conquering Deactivation with Carbon Dioxide. ChemSusChem 2015, 8, 323–330.
62
26. Mehranfar, A.; Izadyar, M.; Esmaeili, A. A. Hydrogen Storage by N-Ethylcarbazol as a New Liquid Organic Hydrogen Carrier: A DFT Study on the Mechanism. Int. J. Hydrogen Energy 2015, 40, 5797–5806.
27. Ferrell III, J. R.; Sachdeva, S.; Strobel, T. A.; Gopalakrishan, G.; Koh, C. A.; Pez, G.; Cooper, A. C.; Herring, A. M. Exploring the Fuel Limits of Direct Oxidation Proton Exchange Membrane Fuel Cells with Platinum Based Electrocatalysts. J. Electrochem. Soc. 2012, 159, B371–B377.
28. Wasmus, S.; Küver, A. Methanol Oxidation and Direct Methanol Fuel Cells: A Selective Review. J. Electroanal. Chem. 1999, 461, 14–31.
29. Liu, H.; Song, C.; Zhang, L.; Zhang, J.; Wang, H.; Wilkinson, D. P. A Review of Anode Catalysis in the Direct Methanol Fuel Cell. J. Power Sources 2006, 155, 95–110.
30. Bianchini, C.; Shen, P. K. Palladium-Based Electrocatalysts for Alcohol Oxidation in Half Cells and in Direct Alcohol Fuel Cells. Chem. Rev. 2009, 4183–4206.
31. Annen, S. P.; Bambagioni, V.; Bevilacqua, M.; Filippi, J.; Marchionni, A.; Oberhauser, W.; Schönberg, H.; Vizza, F.; Bianchini, C.; Grützmacher, H. A Biologically Inspired Organometallic Fuel Cell (OMFC) That Converts Renewable Alcohols into Energy and Chemicals. Angew. Chem. Int. Ed. 2010, 49, 7229–7233.
32. Bonitatibus, P. J.; Rainka, M. P.; Peters, A. J.; Simone, D. L.; Doherty, M. D. Highly Selective Electrocatalytic Dehydrogenation at Low Applied Potential Catalyzed by an Ir Organometallic Complex. Chem. Commun. 2013, 49, 10581–10583.
33. Rainka, M. P.; Peters, A.; Soloveichik, G. Base Effects on Electrochemical Oxidation of Indoline. Int. J. Hydrogen Energy 2013, 38, 3773–3777.
34. Peters, A. J.; Rainka, M. P.; Krishnan, L.; Laramie, S.; Dodd, M.; Reimer, J. A. Electrochemical Characterization of Hydrogen-Bonding Complexation between Indoline and Nitrogen Containing Bases. J. Electroanal. Chem. 2013, 691, 57–65.
35. Driscoll, P. F.; Deunf, E.; Rubin, L.; Arnold, J.; Kerr, J. B. Electrochemical Redox Catalysis for Electrochemical Dehydrogenation of Liquid Hydrogen Carrier Fuels for Energy Storage and Conversion. J. Electrochem. Soc. 2013, 160, G3152–G3158.
36. Compton, R. G.; Banks, C. E. Understanding Voltammetry, 2nd ed.; Imperial College Press: London, UK, 2011.
37. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons: New York, NY, 2001.
38. “iR drop” refers to the shift in voltage seen when either current or solution resistance becomes sufficiently large. It can be explained by the equation V = IR, describing the voltage (V), current (I), and resistance (R) in a circuit. At very large scan rates, the resulting current is also quite large, leading to voltage shifts due to a large value for IR.
39. Valencia, D. P.; González, F. J. Estimation of Diffusion Coefficients by Using a Linear Correlation between the Diffusion Coefficient and Molecular Weight. J. Electroanal. Chem. 2012, 681, 121–126.
40. Savéant, J.-M. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry; John Wiley & Sons: Hoboken, NJ, 2006.
41. Scifinder, version 2015.1; Chemical Abstracts Service: Columbus, OH, 2015; RN 146900-30-3 (accessed July 27, 2015); calculated using Advanced Chemistry Development (ACD/Labs) Software, version 11.02; ACD/Labs 1994-2015.
63
42. Tsierkezos, N. G. Cyclic Voltammetric Studies of Ferrocene in Nonaqueous Solvents in the Temperature Range from 248.15 to 298.15 K. J. Solution Chem. 2007, 36, 289–302.
43. Tsushima, M.; Tokuda, K.; Ohsaka, T. Use of Hydrodynamic Chronocoulometry for Simultaneous Determination of Diffusion Coefficients and Concentrations of Dioxygen in Various Media. Anal. Chem. 1994, 66, 4551–4556.
44. Kulkarni, C. L.; Scheer, B. J.; Rusling, J. F. Potential-Sweep and Pulse Voltammetric Investigations of the Anodic Dimerization of Tetrahydrocarbazoles. J. Electroanal. Chem. 1982.
45. Streitwieser, A.; Kim, Y.-J. Ion Pair Basicity of Some Amines in THF: Implications for Ion Pair Acidity Scales. J. Am. Chem. Soc. 2000, 122, 11783–11786.
46. Gierczyk, B.; Wojciechowski, G.; Brzezinski, B.; Grech, E.; Schroeder, G. Study of the Decarboxylation Mechanism of Fluorobenzoic Acids by Strong N-Bases. J. Phys. Org. Chem. 2001, 14, 691–696.
47. The anodic peak potential of tetramethylguanidine is around 0.9 V, nearly 0.4 V positive of the peak potential for NEC-H12 (0.48 V). Although tetramethylguanidine does start to oxidize as negative as 0.4 V, a solution of 0.1 M tetramethylguanidine only produces about 50 μA of current at 0.48 V, which does not account for the vast majority of the current increase seen.
48. Reymond, G.; Vasilevskis, J.; Toome, V. Preparative Electrochemistry I. A Novel Electrochemical Oxidation of a 1,2,3,4-Tetrahydrocarbazole to a Carbazole Derivative. Heterocycles 1982, 19, 2345–2347.
49. Araujo, C. M.; Simone, D. L.; Konezny, S. J.; Shim, A.; Crabtree, R. H.; Soloveichik, G. L.; Batista, V. S. Fuel Selection for a Regenerative Organic Fuel Cell/Flow Battery: Thermodynamic Considerations. Energy Environ. Sci. 2012, 5, 9534–9542.
50. Zhang, L.; Peng, C.; Zhao, D.; Wang, Y.; Fu, H.-J.; Shen, Q.; Li, J.-X. Cu(II)-Catalyzed C–H (SP3) Oxidation and C–N Cleavage: Base-Switched Methylenation and Formylation Using Tetramethylethylenediamine as a Carbon Source. Chem. Commun. 2012, 48, 5928–5930.
51. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
64
Part II:
“Click”-Functionalized Metallocorrole Complexes for Oxygen Reduction
65
Chapter 4 Oxygen Reduction Catalysis by First-Row Transition Metal Corrole Complexes
4.1 Introduction
Transition metal complexes of macrocycles, such as porphyrins and corroles, are of interest for their ability to catalyze oxygen reduction and potentially replace platinum in PEM fuel cells.1,2 There are several examples3–10 of cobalt corroles and a one example11 of an iron corrole that are active oxygen reduction catalysts.
Although macrocylic oxygen reduction catalysts are inspired by biological examples,12 in practice they must be much more robust in order to maintain durability in a PEM fuel cell environment. One of the drawbacks of using small molecule catalysts is the tendency for catalyst aggregation and the kinetic losses due to slow diffusion to the electrode. This is particularly of concern with large, slow-diffusing molecules like corroles. The goal of the work detailed in Chapters 4 and 5 was to synthesize and characterize corrole complexes that could be attached to an electrode surface, allowing for good incorporation into a PEM fuel cell.
4.2 Results and Discussion
4.2.1 Voltammetric Characterization of Copper, Iron and Cobalt Corrole Complexes
The primary corrole ligand used for these studies is 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole ((C6F5)2(p-OMePh)corrole). It was synthesized using the Gryko “A2B” method of corrole synthesis, which allows for facile introduction of one different meso substituent on the corrole ring,13 such as a single substituent that would allow for surface attachment. The pentafluorophenyl substituents were chosen due to their electron-withdrawing properties. These substituents help stabilize electron-rich late transition metals, and literature precedent9 indicates that corrole complexes with electron withdrawing substituents have more positive onset potentials for oxygen reduction, due to their more positive metal reduction potentials. The 4-methoxyphenyl substituent was used as a “placeholder” while ligands with surface attachment moieties were developed (Chapter 5).
Copper, iron and cobalt complexes of the 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole ligand were prepared using modified literature procedures for the preparation of similar copper,14 iron,15 and cobalt8 corrole complexes (Figure 4.1). While both cobalt and iron corroles have demonstrated activity for oxygen reduction, no copper complexes have been reported as oxygen reduction catalysts. Our copper complex was used for initial studies to gain familiarity with the system, in order to conserve the less-stable cobalt and iron complexes. Although not anticipated to be active for oxygen reduction, it proved to be the easiest of the three to synthesize. The copper complex was also used for initial Huisgen azide-alkyne “click” studies (Chapter 5), since there was concern of transmetallation with the copper catalyst required for the “click” reaction.
66
Figure 4.1: Copper, iron and cobalt corrole complexes studied in this work. Diethyl ether ligands not shown on 4-2a for clarity.
The cyclic voltammogram of copper(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole ((C6F5)2(p-OMePh)corroleCu, 4-1) in dichloromethane displays a reversible reduction at −0.23 V (Figure 4.2). This process is assigned as a metal-centered reduction from copper (III) to copper (II), based on comparison with the free-base ligand and with potentials reported for analogous complexes (Table 4.1).16 Previous studies of copper corroles have found that both copper 5,10,15-tris(pentafluorophenyl)corrole ((C6F5)3corroleCu) and copper 5,10,15-tris(4-methoxyphenyl)corrole ((p-OMePh)3corroleCu) are copper(III),16 despite the fact that the copper(II)-corrole radical state is often only slightly higher in energy than the copper(III)-neutral corrole state.14,16,17
A quasi-reversible oxidative peak is seen at 0.73 V, which is assigned as a ligand-based oxidation. Corrole ligands are trianionic, which likely accounts for the two additional oxidation peaks seen, although they are difficult to resolve due to solvent oxidation. No ligand reduction is seen, due to the potentials for 4-1 being cathodically shifted such that any ligand reduction would be outside the solvent window. Voltammograms of 4-1 in acetonitrile are similar at the metal center, although ligand processes vary slightly, especially at very positive potentials.
67
Figure 4.2: 1 mM (C6F5)2(p-OMePh)corroleCu. 0.1 NBu4BF4 in CH2Cl2, 100 mV/s, glassy carbon electrode.
Table 4.1: Comparison of redox processes seen in copper corrole complexes. Potentials reported as V vs. Ag/Ag+. For the literature examples, potentials were originally reported in V vs. SCE and were converted to Ag/Ag+ by subtracting 298 mV.18 Peaks marked (a) were poorly resolved; potentials estimated.
In contrast to the relative stability of 4-1, reaction of 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole with iron led to two separate species: iron(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole bis(diethyl ether) ((C6F5)2(p-OMePh)corroleFe(Et2O)2, 4-2a) and bis(iron(IV) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole) (μ-oxo) (((C6F5)2(p-OMePh)corroleFe)2(μ-O), 4-2b). 4-2a was more sensitive than its 5,10,15-tris(perfluorophenyl)corrole analogue,15 and was cleanly isolated either by careful workup under air- and water-free conditions or through reduction of 4-2b with hydrazine. Similarly, 4-2b can be formed from 4-2a by recrystallization under wet benchtop conditions.19
4-2a also displays three reversible-appearing ligand oxidations, much like 4-1. However, these oxidations are shifted cathodically, to 0.29, 0.84 and 1.23 V vs. Ag/Ag+, meaning the ligand is easier to oxidize in the case of 4-2a. This is consistent with its greater overall susceptibility to oxidation. There is a fourth oxidation at 0.01 V, which is assigned to the oxidation of iron(III) to iron(IV). There are also two reductions, a quasi-reversible one at −0.76 V and a reversible one at −1.88 V. These are not readily assigned, as the literature is vague on the redox processes present in iron(III) corrole complexes with solvent axial ligands. Comparison with previous work11 suggests that the process at −0.76 V is reduction to iron(II), and the process at −1.88 V is either
-20
0
20
40
60
80
-2 -1.5 -1 -0.5 0 0.5 1 1.5 2
(C6F5)2(p-OMePh)corroleCu
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
68
reduction to iron(I) or ligand-centered. These assignments are based in part on the fact that the potential differences between corrole ligand oxidations and reductions are usually at least 2 V15 and electrochemical formation of iron(V) corrole is unlikely.20
Figure 4.3: 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2. 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Surprisingly, the voltammogram of 4-2b is almost identical to that of 4-2a, with the exception of the peak assigned to FeIII/IV at 0.01 V (Table 4.2, Figure 4.4). Although an FeIV/III reduction process has been reported for a similar bis(iron corrole) μ-oxo complex,15 no reduction to iron(III) is seen for 4-2b, even when scanning cathodically first. It has been suggested11 that iron(IV) can undergo a reduction directly to an iron(II)-corrole radical anion species, which might occur for 4-2b at −0.74 V, nearly the same potential as the reduction to iron(II) for 4-2a. An alternative explanation is that 4-2a is unstable under our voltammetry conditions, and the voltammogram seen in Figure 4.3 is actually of a mixture of 4-2a and 4-2b.
Complex Other Reductions
FeIV/III Reduction
FeIII/IV Oxidation
Corrole Oxidations
4-2a −1.88, −0.76 --- 0.01 0.29, 0.84, 1.23 4-2b −1.93, −0.74 not seen --- 0.29, 0.80, 1.28 (C6F5)3corroleFe(IV)Cl11 −1.05 0.11 --- not reported ((C6F5)3corroleFe)2(μ-O)15 --- −0.17 --- 0.58, 0.95
Table 4.2: Comparison of redox processes seen in iron corrole complexes. Potentials reported as V vs. Ag/Ag+. For the literature examples, potentials were originally reported in V vs. SCE and were converted to Ag/Ag+ by subtracting 298 mV.18
-100
-50
0
50
100
-2.5 -2 -1.5 -1 -0.5 0 0.5 1 1.5
(C6F5)2(p-OMePh)corroleFe
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
69
Figure 4.4: 0.54 mM ((C6F5)2(p-OMePh)corroleFe)2(μ-O). 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode. The most negative reduction, at −1.93 V, is somewhat visible here but was seen more clearly in other experiments that scanned to more negative potentials.
Figure 4.5: 1 mM (C6F5)2(p-OMePh)corroleCo. 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-50
0
50
100
150
200
-2 -1.5 -1 -0.5 0 0.5 1 1.5 2
((C6F5)2(p-OMePh)corroleFe)2( -O)
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-20
0
20
40
60
80
-2 -1.5 -1 -0.5 0 0.5 1 1.5
(C6F5)2(p-OMePh)corroleCo
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
70
Cobalt(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole ((C6F5)2(p-OMePh)corroleCo, 4-3) was also found to be somewhat sensitive to reaction conditions and workup, and reduction with hydrazine was also used to cleanly isolate the cobalt(III) species.
Voltammetry of 4-3 (Figure 4.5) displays four cleanly reversible peaks: two oxidations and two reductions. Comparison with previously reported similar complexes6 allows for assignment of the first reduction to the CoIII/II transition and the second reduction to CoII/I. The first oxidation is assigned to the formation of cobalt(III) with a corrole radical cation, and the second oxidation is also ligand-based. The data tabulated below (Table 4.3) indicates good agreement between redox potentials and the presence of either electron-donating or electron-withdrawing meso substituents on the corrole ring, with more electron-withdrawing substituents shifting the metal reductions anodically. The reduction potentials observed for 4-3 lie between cobalt tris(pentafluorophenyl)corrole, which has three electron-withdrawing substituents, and cobalt 5,15-(mesityl)-10-(pentafluorophenyl)corrole, which has two electron-donating substituents and one electron-withdrawing substituent.
Table 4.3: Comparison of redox processes seen in cobalt corrole complexes. Potentials reported in as V vs. Ag/Ag+. Analysis of 4-3 was performed in acetonitrile. Literature examples were performed in dichloromethane with potentials originally reported in V vs. SCE and were converted to Ag/Ag+ by subtracting 298 mV.18
4.2.2 Solution-Phase Oxygen Reduction Reactivity
Complexes 4-1, 4-2a, 4-2b and 4-3 were examined for their oxygen reduction activity in solution. In each case, oxygen was bubbled into a solution of metallocorrole. Non-aqueous, non-protic solvents support reduction of oxygen to superoxide and [O2]−2, but cannot promote the formation of water in the absence of protons. In order to better simulate the conditions found in a PEM fuel cell, acetic acid was added to each solution to give a final concentration of 100 mM acid.
As expected, 4-1 did not promote oxygen reduction beyond the reduction seen on the bare glassy carbon electrode. The onset potential for reduction at the bare electrode, defined as the point at which the current reaches −100 μA, is −0.89 V vs. Ag/Ag+ in acidic acetonitrile and −0.98 V vs. Ag/Ag+ in acidic dichloromethane. For 4-1, the reduction onset potential in dichloromethane is −0.96 V, representing an insignificant improvement over the bare electrode (Figure 4.6). 4-2b also does not significantly promote oxygen reduction, displaying a reduction onset at −0.83 V vs. Ag/Ag+ (Figure 4.7).
71
Figure 4.6: 100 mM acetic acid in oxygen-saturated solution, with (green) and without (grey, dashed) 1 mM (C6F5)2(p-OMePh)corroleCu. 0.1 NBu4BF4 in CH2Cl2, 100 mV/s, glassy carbon electrode.
Figure 4.7: 100 mM acetic acid in oxygen-saturated solution, with (brown) and without (grey, dashed) 0.54 mM ((C6F5)2(p-OMePh)corroleFe)2(μ-O). 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
-500
-400
-300
-200
-100
0
100
200
-2 -1.5 -1 -0.5
(C6F5)2(p-OMePh)corroleCu, O2, acetic acidO2 and acetic acid alone
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-1000
-800
-600
-400
-200
0
200
400
-2 -1.5 -1 -0.5
((C6F5)2(p-OMePh)corroleFe)2( -O), O2, acetic acidO2 and acetic acid alone
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
72
Figure 4.8: 100 mM acetic acid in oxygen-saturated solution, with (red) and without (grey, dashed) 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2. 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Figure 4.9: 100 mM acetic acid in oxygen-saturated solution, with (blue) and without (grey, dashed) 1 mM (C6F5)2(p-OMePh)corroleCo. 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
In contrast, both 4-2a and 4-3 show activity towards oxygen reduction catalysis, with onset potentials of −0.56 V (Figure 4.8) and −0.34 V (Figure 4.9) respectively. In the case of 4-3, the
-1000
-800
-600
-400
-200
0
200
-2 -1.6 -1.2 -0.8 -0.4
(C6F5)2(p-OMePh)corroleFe, O2, acetic acidO2 and acetic acid alone
Potential (V vs. Ag/Ag+)
Cur
rent
(A
)
-1000
-800
-600
-400
-200
0
200
-2 -1.6 -1.2 -0.8 -0.4
(C6F5)2(p-OMePh)corroleCo, O2, acetic acidO2 and acetic acid alone
Cur
rent
(A
)
Potential (V vs. Ag/Ag+)
73
onset potential of oxygen reduction corresponds to the CoIII/II transition, which has been previously shown to correlate with oxygen reduction, as the metal(II) oxidation state is the form that binds oxygen.9,11 For 4-2a, the proposed reduction to iron(II) occurs at −0.76 V vs. Ag/Ag+, which is 200 mV more negative than the oxygen reduction potential seen here. However, it has also been proposed that an internal corrole-to-metal(III) electron transfer can occur, reducing the metal center to metal(II) and enabling oxygen binding.11 To further support the dependence of catalysis on the presence of the metal in 4-2a and 4-3, the free corrole ligand was also examined with oxygen and acid and found to have no activity towards oxygen reduction.
4.2.3 Analysis of Oxygen Reduction Reactivity by Rotating Ring-Disk Voltammetry
Oxygen reduction in aqueous acid proceeds by two different possible pathways. One is the two-electron reduction to hydrogen peroxide. This can be followed by further reduction to water by two additional electrons, depending on the nature of the catalyst used. The other pathway is the direct four-electron reduction to water. The four electron reduction is more desirable, as it produces more current and avoids the production of hydrogen peroxide, which can damage the fuel cell membrane (Scheme 4.1).
Scheme 4.1: Pathways for oxygen reduction to hydrogen peroxide or water. “ads” indicates that the species is adsorbed on the electrode surface.21
Rotating ring-disk voltammetry (RRDV) is a standard electrochemical technique used for characterizing the activity of an oxygen reduction catalyst. It uses a three-electrode setup similar to that used in cyclic voltammetry, except that the working electrode consists of a disk surrounded by a ring, with the two components being electrically independent.
In our studies, each catalyst was incorporated into an ink using high-surface area carbon powder and a small amount was deposited on the disk portion of the ring-disk electrode. A motor is used to rotate the working electrode at defined speeds, allowing for the measurement of current under well-defined mass transport conditions. A bipotentiostat is used to sweep the disk potential, as in cyclic voltammetry, while the ring potential is held constant. The current from the disk results from the reduction of oxygen by the catalyst ink, while the current from the ring allows for the detection of hydrogen peroxide.
4-1, 4-2a and 4-3 were all examined using RRDV. Although 4-1 was not anticipated to display catalytic activity for oxygen reduction, it was used as a means of validating the experimental setup, as it was the easiest complex to prepare and the least sensitive to water.
74
4-1 starts to reduce oxygen at 0.15 V vs. NHE (Figure 4.10),22 making it a slightly better catalyst than an ink of just carbon powder (−0.16 V vs. NHE). However, it did not display any dependence on rotation rate until nearly −0.2 V vs. NHE. Since there is very little dependence on mass transport of oxygen to the electrode surface, it is likely that the catalyst has very slow kinetics, such that there is always ample oxygen available at the electrode surface. Additional analysis indicates that 4-1 favors two-electron reduction to hydrogen peroxide over four-electron reduction to water.23
Figure 4.10: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleCu adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode. 0.5 M aqueous H2SO4, 10 mV/s.
4-2a has a much more positive onset potential for oxygen reduction, at 0.63 V vs. NHE (Figure 4.11). It also displays a dependence on rotation rate, allowing for the calculation of the number of electrons involved in the reaction using the Koutecký-Levich equation (Equation 4.1).
Equation 4.1
A plot of 1/i vs. 1/(ω1/2) at any given potential should be linear, with a slope of 1/((0.201)nFD2/3ν−1/6C), where i is the current (A), ω is the rotation rate (rpm), n is the number of electrons transferred, F is Faraday’s constant (96485 C/mol), A is the electrode area (cm2), D is the diffusion coefficient of the analyte (cm2/s), ν is the kinematic viscosity of the solution (cm2/s), and C is the analyte concentration (mol/mL). Assuming the other values are known, this can be used to find n, which indicates the amount of two-electron vs. four-electron reduction occurring.
-400
-300
-200
-100
0
100
200
-0.2 0 0.2 0.4 0.6 0.8 1
400 rpm600 rpm800 rpm1000 rpm1296 rpm
Potential (V vs. NHE)
Cur
rent
(A
)
75
Figure 4.11: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleFe(Et2O)2 adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode. 0.5 M aqueous H2SO4, 10 mV/s.
Figure 4.12: Koutecký-Levich analysis of oxygen reduction by (C6F5)2(p-OMePh)corroleFe(Et2O)2.
In our analysis, A is 0.126 cm2, D is the diffusion coefficient of oxygen in 0.5 M sulfuric acid (1.4 x 10−5 cm2/s), ν is 0.01 cm2/s, and C is the concentration of oxygen in oxygen-saturated 0.5
-500
-400
-300
-200
-100
0
100
-0.2 0 0.2 0.4 0.6 0.8 1
100 rpm400 rpm600 rpm800 rpm1000 rpm1296 rpm
Potential (V vs. NHE)
Cur
rent
(A
)
-10
-8
-6
-4
-2
0
0 0.02 0.04 0.06 0.08 0.1 0.12
0.2 V0.1 V0 V-0.1 V-0.2 V
1/( , rpm)1/2
1/(i,
mA
)
76
M sulfuric acid (1.10 x 10−6 mol/mL).21,24 This allowed for the calculation of n at various potentials from the Koutecký-Levich slopes (Figure 4.12, Table 4.4).
Another calculation of the extent of two-electron vs. four-electron reduction can also be obtained by comparing the current responses at the disk and at the ring, as the ring is set to a potential where it will detect the amount of hydrogen peroxide formed. The number of electrons is given by
Equation 4.2
where iD is the disk current, iR is the ring current, and N is the collection efficiency, which is a parameter unique to each electrode that refers to what percentage of the products generated at the disk will reach the ring. N for our system is 0.42.
The percentage of hydrogen peroxide is given by
Equation 4.3
V vs. NHE n (Koutecký-Levich) n (H2O2) % H2O2 0.2 3.55 3.54 23.2 0.1 3.45 3.63 18.3 0 3.70 3.72 13.9
−0.1 3.86 3.77 11.5 −0.2 3.97 3.83 8.6
Table 4.4: Number of electrons and % hydrogen peroxide calculated for the reduction of oxygen by (C6F5)2(p-OMePh)corroleFe(Et2O)2 at different potentials.
Both Koutecký-Levich analysis and hydrogen peroxide formation indicate that 4-2a is capable of nearly complete four-electron reduction to water, with four-electron reduction favored more at more negative potentials (Table 4.4). Our results are consistent with the single previous study of oxygen reduction by iron corroles,11 which also concluded that iron corroles favored four-electron reduction. This suggests that iron corrole complexes are worthy of further study for this particular application.
4-3 displays the most positive onset potential of the corrole complexes studied, at 0.75 V vs. NHE (Figure 4.13). However, it reduces oxygen by a two-electron process, forming up to 95% hydrogen peroxide.23 This can also be seen in Figure 4.14, which indicates nearly twice as much current formed with 4-2a as with 4-3. This is also consistent with previous reports of cobalt corrole complexes, which indicate that cobalt favors two-electron reduction except in the case of very electron-withdrawing corrole ligands or bimetallic complexes.3,6,9,11
77
Figure 4.13: Rotation-rate dependence of oxygen reduction by (C6F5)2(p-OMePh)corroleCo adsorbed on carbon powder and deposited on a glassy carbon rotating disk electrode. 0.5 M aqueous H2SO4, 10 mV/s.
Figure 4.14: Comparison of oxygen reduction by blank carbon ink (gold) and copper(III), iron(III), and cobalt(III) corrole complexes (green, red, and blue respectively). 0.5 M aqueous H2SO4, 10 mV/s, 400 rpm.
-300
-200
-100
0
100
0 0.2 0.4 0.6 0.8 1
400 rpm600 rpm1000 rpm1500 rpm1800 rpm2400 rpm
Potential (V vs. NHE)
Cur
rent
(A
)
-300
-200
-100
0
100
-0.2 0 0.2 0.4 0.6 0.8 1
blank, N2blank, O2Cu corroleFe corroleCo corrole
Potential (V vs. NHE)
Cur
rent
(A
)
78
4.3 Conclusions
Copper (III), iron(III), and cobalt(III) complexes of 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole, as well as a bisiron(IV) biscorrole μ-oxo species, have been synthesized and characterized electrochemically. Both the iron(III) and cobalt(III) species, as expected, are active for oxygen reduction, with the iron species favoring four-electron reduction and the cobalt species favoring two-electron reduction. The behavior measured for oxygen reduction by these complexes provides a useful benchmark for comparison with similar complexes covalently attached to an electrode surface.
4.4 Experimental Methods and Materials
Ferrocene was obtained commercially and purified by sublimation before use. Tetra-n-butylammonium tetrafluoroborate (NBu4BF4) was obtained commercially and purified by dissolving in ethyl acetate, precipitating with pentane, filtering, and drying in vacuo overnight. All other reagents were obtained commercially and used as received. Non-aqueous electrochemical experiments were performed in high-purity acetonitrile or dichloromethane, dried and degassed using a Phoenix solvent drying system commercially available from JC Meyer Solvent Systems.
All corrole ligands and complexes were prepared by Dr. Heather Buckley. Copper(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole, iron(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole bis(diethyl ether) and bis(iron(IV) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole) (μ-oxo) were prepared according to previously published procedures.25 Cobalt(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole was prepared according to a modified literature procedure.8,19
Cyclic voltammetry experiments were performed using a Gamry 600 Reference potentiostat. All voltammetry experiments used 0.1 M tetraethylammonium tetrafluoroborate (NEt4BF4) in acetonitrile or 0.1 M NBu4BF4 in dichloromethane as the supporting electrolyte. Voltammetry solutions were saturated with either nitrogen or oxygen and kept under a blanket of that same gas during analysis. A three-electrode cell was used, with a glassy carbon working electrode (3 mm diameter), a platinum mesh counter electrode and a silver ion non-aqueous reference electrode. The reference electrode was prepared by immersing a silver wire in a 0.01 M solution of silver nitrate in 0.1 M NEt4BF4 in acetonitrile, and was separated from the bulk solution by a Vycor frit. The reference electrode potential was checked periodically by comparison with ferrocene, and measured potentials were corrected as necessary. Potentials are reported vs. Ag/Ag+ (0.01 AgNO3), which is 0.548 V vs. NHE and −0.087 V vs. ferrocene (Fc).18
Rotating ring-disk voltammetry was performed using an ALS RRDE-3A Rotator and a BioLogic Science Instruments VSP bipotentiostat. The electrodes used consisted of a working electrode with a glassy carbon disk (4 mm diameter) and a platinum ring (collection efficiency N=0.42), a platinum mesh counter electrode, and a silver/silver chloride reference electrode. The reference electrode was prepared by electrocoating (10 min at 3 V) the surface of a silver wire with a layer of silver chloride using a saturated potassium chloride solution. The coated wire was then immersed in saturated potassium chloride and was kept separate from the bulk solution by a Vycor frit. 0.5 M sulfuric acid in water (Millipore, 18 MΩ) was used as the electrolyte. The disk
79
potential was swept at 10 mV/s while the ring potential was held at 1 V. Potentials are reported vs. NHE, which is −0.197 V vs. Ag/AgCl (saturated KCl).26 Koutecký-Levich analysis for iron(III) 5,15-bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole bis(diethyl ether) was performed on currents that had been corrected by subtracting out the current from the catalyst in an oxygen-free environment.
To prepare the catalyst inks physisorbed on the glassy carbon surface, each metallocorrole complex was dissolved in a minimal amount of dichloromethane. Cabot XC-72R high surface area carbon powder was added to this solution and the entire mixture was stirred under atmosphere until the DCM was completely evaporated. 5% Nafion in isopropyl alcohol was added to the catalyst-carbon powder and the entire mixture was sonicated until a homogeneous mixture was created (approximately 30 minutes). The rotating ring-disk electrode surface was polished using 0.05 μm alumina, rinsed with distilled water, and allowed to dry. A small amount of ink was then deposited on the electrode using a micropipette and dried at room temperature overnight.
4.5 References
1. Collman, J. P.; Marrocco, M.; Denisevich, P.; Koval, C.; Anson, F. C. Potent Catalysis of the Electroreduction of Oxygen to Water by Dicobalt Porphyrin Dimers Adsorbed on Graphite Electrodes. J. Electroanal. Chem. Interfacial Electrochem. 1979, 101, 117–122.
2. Guilard, R.; Kadish, K. M. Some Aspects of Organometallic Chemistry in Metalloporphyrin Chemistry: Synthesis, Chemical Reactivity, and Electrochemical Behavior of Porphyrins with Metal-Carbon Bonds. Chem. Rev. 1988, 88, 1121–1146.
3. Kadish, K. M.; Frémond, L.; Ou, Z.; Shao, J.; Shi, C.; Anson, F. C.; Burdet, F.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Cobalt(III) Corroles as Electrocatalysts for the Reduction of Dioxygen: Reactivity of a Monocorrole, Biscorroles, and Porphyrin-Corrole Dyads. J. Am. Chem. Soc. 2005, 127, 5625–5631.
4. Kadish, K. M.; Shao, J.; Ou, Z.; Frémond, L.; Zhan, R.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Electrochemistry, Spectroelectrochemistry, Chloride Binding, and O2 Catalytic Reactions of Free-Base Porphyrin-Cobalt Corrole Dyads. Inorg. Chem. 2005, 44, 6744–6754.
5. Kadish, K. M.; Frémond, L.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Cobalt(IV) Corroles as Catalysts for the Electroreduction of O2: Reactions of Heterobimetallic Dyads Containing a Face-to-Face Linked Fe(III) or Mn(III) Porphyrin. J. Inorg. Biochem. 2006, 100, 858–868.
6. Kadish, K. M.; Shen, J.; Frémond, L.; Chen, P.; El Ojaimi, M.; Chkounda, M.; Gros, C. P.; Barbe, J.-M.; Ohkubo, K.; Fukuzumi, S.; Guilard, R. Clarification of the Oxidation State of Cobalt Corroles in Heterogeneous and Homogeneous Catalytic Reduction of Dioxygen. Inorg. Chem. 2008, 47, 6726–6737.
7. Kadish, K. M.; Frémond, L.; Shen, J.; Chen, P.; Ohkubo, K.; Fukuzumi, S.; El Ojaimi, M.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Catalytic Activity of Biscobalt Porphyrin-Corrole Dyads toward the Reduction of Dioxygen. Inorg. Chem. 2009, 48, 2571–2582.
8. Dogutan, D. K.; Stoian, S. A.; McGuire, Jr., R.; Schwalbe, M.; Teets, T. S.; Nocera, D. G. Hangman Corroles: Efficient Synthesis and Oxygen Reaction Chemistry. J. Am. Chem. Soc. 2011, 133, 131–140.
80
9. Schechter, A.; Stanevsky, M.; Mahammed, A.; Gross, Z. Four-Electron Oxygen Reduction by Brominated Cobalt Corrole. Inorg. Chem. 2012, 51, 22–24.
10. Wang, Z.; Lei, H.; Cao, R.; Zhang, M. Cobalt Corrole on Carbon Nanotube as a Synergistic Catalyst for Oxygen Reduction Reaction in Acid Media. Electrochim. Acta 2015, 171, 81–88.
11. Collman, J. P.; Kaplun, M.; Decréau, R. A. Metal Corroles as Electrocatalysts for Oxygen Reduction. Dalton Trans. 2006, 554–559.
12. Wang, B. Recent Development of Non-Platinum Catalysts for Oxygen Reduction Reaction. J. Power Sources 2005, 152, 1–15.
13. Koszarna, B.; Gryko, D. T. Efficient Synthesis of Meso-Substituted Corroles in a H2O-MeOH Mixture. J. Org. Chem. 2006, 71, 3707–3717.
14. Luobeznova, I.; Simkhovich, L.; Goldberg, I.; Gross, Z. Electronic Structures and Reactivities of Corrole−Copper Complexes. Eur. J. Inorg. Chem. 2004, 1724–1732.
15. Simkhovich, L.; Mahammed, A.; Goldberg, I.; Gross, Z. Synthesis and Characterization of Germanium, Tin, Phosphorus, Iron, and Rhodium Complexes of Tris(pentafluorophenyl)corrole, and the Utilization of the Iron and Rhodium Corroles as Cyclopropanation Catalysts. Chem. Eur. J. 2001, 7, 1041–1055.
16. Ou, Z.; Shao, J.; Zhao, H.; Ohkubo, K.; Wasbotten, I. H.; Fukuzumi, S.; Ghosh, A.; Kadish, K. M. Spectroelectrochemical and ESR Studies of Highly Substituted Copper Corroles. J. Porphyr. Phthalocyanines 2004, 8, 1236–1247.
17. Ghosh, A.; Wondimagegn, T.; Parusel, A. B. J. Electronic Structure of Gallium, Copper, and Nickel Complexes of Corrole. High-Valent Transition Metal Centers versus Noninnocent Ligands. J. Am. Chem. Soc. 2000, 122, 5100–5104.
18. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
19. Buckley, H. L. Unpublished Work. 20. Steene, E.; Wondimagegn, T.; Ghosh, A. Electrochemical and Electronic Absorption
Spectroscopic Studies of Substituent Effects in Iron(IV) and Manganese(IV) Corroles. Do the Compounds Feature High-Valent Metal Centers or Noninnocent Corrole Ligands? Implications for Peroxidase Compound I and II. J. Phys. Chem. B 2001, 105, 11406–11413.
21. Song, C.; Zhang, J. Electrocatalytic Oxygen Reduction Reaction. In PEM Fuel Cell Electrocatalyst and Catalyst Layers. Fundamentals and Applications; Zhang, J., Ed.; Springer-Verlag London: London, UK, 2008; pp 89–134.
22. Onset of oxygen reduction is defined as the point at which the reduction current reaches −20 μA at 400 rpm.
23. McNicholas, B. J. Oxygen Reduction Catalysis by First-Row Transition Metal Corrole Complexes. B.S. Thesis, University of California, Berkeley, Berkeley, CA, May 2014.
24. Hsueh, K.-L.; Gonzalez, E. R.; Srinivasan, S. Electrolyte Effects on Oxygen Reduction Kinetics at Platinum: A Rotating Ring-Disc Electrode Analysis. Electrochim. Acta 1983, 28, 691–697.
25. Buckley, H. L.; Rubin, L. K.; Chromiński, M.; McNicholas, B. J.; Tsen, K. H. Y.; Gryko, D. T.; Arnold, J. Corroles That “Click”: Modular Synthesis of Azido- and Propargyl-Functionalized Metallocorrole Complexes and Convergent Synthesis of a Bis-Corrole Scaffold. Inorg. Chem. 2014, 53, 7941–7950.
81
26. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons: New York, NY, 2001.
82
Chapter 5 “Clickable” Propargyl- and Azido-Modified Metallocorrole Complexes: Biscorrole
Characterization and Covalent Attachment
5.1 Introduction
Metallocorrole complexes have a variety of demonstrated uses in oxygen reduction catalysis,1–9 small molecule activation,10,11 environmental sensing,12–14 and medical applications.15–19 This presents a need to develop methods for incorporating such complexes into devices and drugs. For example, while metallocorroles could be used in PEM fuel cells as an alternative to platinum, heterogeneous catalysts benefit from relative robustness, despite the sensitivity of platinum to poisoning by carbon monoxide and other contaminants.20 In contrast, small molecule catalysts such as metallocorroles require anchoring in a device in order to prevent problems of catalyst aggregation and slow diffusion to the electrode.
While physisorption of molecules onto an electrode surface is possible, covalent attachment is a more robust method of attachment. Huisgen azide-alkyne cycloaddition,21,22 or “click” chemistry, has been demonstrated as a clean, quantitative reaction for attaching molecules to surfaces23–27 or to biomolecules.28–32 In both cases, modifications have been made with azide groups, enabling clean attachment of alkyne-functionalized molecules in situ.33–37 Herein is presented the characterization of new “clickable” corroles and “clicked” biscorroles, and preliminary studies of attachment to an azide-modified glassy carbon electrode.
5.2 Results and Discussion
5.2.1 Preparation of “Clickable” Metallocorroles and Homo- and Hetero-bimetallic Biscorroles
5,15-Bis(pentafluorophenyl)-10-(4-methoxyphenyl)corrole, 5,15-bis(pentafluorophenyl)-10-(4-propargyloxyphenyl)corrole, and 5,15-bis(pentafluorophenyl)-10-(3-azidophenyl)corrole ligands were prepared using the Gryko “A2B” method,38 which allows for facile introduction of one different meso substituent on the corrole ring. 4-Methoxyphenyl was chosen as a model “B” substituent that would enable comparison of metal redox processes between complexes that differed only in their peripheral ligands. These ligands were then metallated with both copper and iron (Scheme 5.1).
Initially, attempts were made to “click” small test substrates to the unmetallated ligands; however, this was unsuccessful due to metallation of the ligands by the copper catalyst used. Instead, the propargyl- and azide-functionalized metallocorroles were reacted with small azide and alkyne test substrates by Huisgen azide-alkyne cycloaddition (Scheme 5.2). Although copper complexes were used for initial studies due to the presence of copper in the “click” reaction and concerns of transmetallation, no substitution of copper into the iron corroles was seen.
83
Scheme 5.1: Syntheses of (C6F5)2(p-OMePh)corroleCu (5-1), (C6F5)2(p-OMePh)corroleFe(Et2O)2 (5-2), (C6F5)2(p-O(CH2CCH)Ph)corroleCu (5-3), (C6F5)2(p-O(CH2CCH)Ph)corroleFe(Et2O)2 (5-4), and (C6F5)2(m-CH2N3)Ph)corroleCu (5-5).
84
Scheme 5.2: (C6F5)2(p-O(CH2CCH)Ph)corroleCu (5-3) reacting with a small test substrate to form (C6F5)2(p-O(CH2(C2HN3(CH2CH2OCH2CH2OH)))Ph)corroleCu (5-6).
Two bimetallic biscorrole complexes were also prepared, using the propargyl- and azido-functionalized metallocorroles 5-3, 5-4, and 5-5 (Scheme 5.3).
Scheme 5.3: Formation of biscorrole complexes 5-7 and 5-839 using Husigen azide-alkyne cycloaddition.
5.2.2 Electrochemical Characterization
Three copper corrole complexes with different peripheral substituents (5-1, 5-3, and 5-6) were screened to evaluate the effects of peripheral ligand changes on the redox activity at the metal center (Figure 5.1). No substantial changes were seen, suggesting that attachment to a surface is also not likely to substantially shift the metal redox potential. This enabled screening of possible
85
catalysts by solution-phase and physisorption methods (Chapter 4) without worry that the catalytic activity would change dramatically due to a shift in redox potential upon covalent attachment to a surface.
Figure 5.1: 1 mM (C6F5)2(p-OMePh)corroleCu (dark green), (C6F5)2(p-O(CH2CCH)Ph)corroleCu (olive green), or (C6F5)2(p-O(CH2(C2HN3(CH2CH2OCH2CH2OH)))Ph)corroleCu (lime green). 0.1 NBu4BF4 in CH2Cl2, 100 mV/s, glassy carbon electrode.
Biscorrole complexes 5-7 and 5-8 were characterized by cyclic voltammetry. The voltammogram of 5-7 shows a single copper reduction peak at −0.21 V vs. Ag/Ag+ in dichloromethane (Figure 5.2, left). This reduction occurs at essentially the same potential with approximately double the current as is observed for its monocorrole counterpart (Figure 5.2, right), and corresponds to the CuIII/II process. 5-7 also has a quasireversible oxidation at 0.74 vs. Ag/Ag+ with an oxidative shoulder at 0.64 V (100 mV/s) that merges with the main peak at higher scan rates. These peaks correspond to ligand oxidation, and are again very similar to the corresponding ligand oxidation for copper 5,15-(pentafluorophenyl)-10-(4-methoxyphenyl)corrole (4-1), at nearly the same potential but double the current. The presence of a single peak for the CuIII/II transition and the nearly identical ligand oxidations for both the monocorrole and biscorrole species suggest that there is no interaction or electronic communication between the two copper centers. This is consistent with what is observed by NMR spectroscopy.40
Figure 5.2: 1 mM copper-copper biscorrole complex (left). 1 mM (C6F5)2(p-OMePh)corroleCu (right). 0.1 NBu4BF4 in CH2Cl2, 100 mV/s, glassy carbon electrode.
Cyclic voltammetry of 5-8 in acetonitrile displays an oxidation and a reduction consistent with the presence of both copper and iron (Figure 5.3, left). The reduction at −0.19 vs. Ag/Ag+ corresponds to the CuIII/II couple, and the oxidation at 0.04 V corresponds to FeIII/IV. Again, these are essentially the same peaks observed for the two corresponding mononuclear species (Figure 5.3, right). Similar to what was seen with 5-7, the voltammogram of 5-8 displays oxidative peaks corresponding to ligand oxidation. Additionally, there is no indication of interaction between the two metal centers, based on the fact that the metal oxidations and reductions are observed at basically the same potentials as seen in the monometallic analogues. It should be noted here that the peaks recorded for 5-8 have very low current due to the poor solubility of this compound in both acetonitrile and dichloromethane.
Figure 5.3: <1 mM copper-iron biscorrole complex (left). 1 mM (C6F5)2(p-OMePh)corroleCu (4-1, green, right) and 1 mM (C6F5)2(p-OMePh)corroleFe(Et2O)2 (4-2a, red, right). 0.1 NEt4BF4 in acetonitrile, 100 mV/s, glassy carbon electrode.
Table 5.1: Assignment of oxidation and reduction processes for the bimetallic biscorroles and comparison with related compounds. Potentials reported as V vs. Ag/Ag+. For the literature examples, potentials were originally reported in V vs. SCE and were converted to Ag/Ag+ by subtracting 298 mV.43 (a) Peaks were poorly resolved; potentials estimated. (b) Irreversible oxidations; reported as the oxidative peak potential.
5.2.3 Covalent Attachment to Glassy Carbon Electrodes
Preliminary studies have been conducted attempting to covalently attach 5-3 to an azide-functionalized glassy carbon electrode surface. Azide modification of glassy carbon electrodes has been previously demonstrated,34 and ethynylferrocene has been attached and studied by cyclic voltammetry as a proof of concept.25 However, this would be the first example of a macrocyclic complex covalently tethered to a surface using azide-alkyne cycloaddition.
Glassy carbon disks were modified to create an azide functionalized surface using a literature procedure, and reacted with 5-3 in the presence of copper iodide and triethylamine. Voltammetry was then performed on the disks to assess the presence of the copper corrole on the surface. Initial experiments were done without any attempt to exclude water or oxygen from the system, thus, a large oxygen reduction peak can be seen in the voltammograms below.
A voltammogram of a copper corrole-modified glassy carbon disk in acetonitrile is shown below (Figure 5.4, green trace). A reduction peak was seen at −0.64 V vs. ferrocene (−0.56 vs. Ag/Ag+), which was attributed to the CuIII/II transition, although it occurs at a more negative potential than would be expected. A corresponding oxidation peak was not seen, however. For surface-attached species, the oxidation and reduction peaks would be anticipated to occur at the same potential. It is possible that the oxidation of copper is masked by the oxidation of superoxide, which is the oxidation peak marked as “oxygen reduction” in the figure. After the initial voltammetric studies, the disk was placed in a solution of ethylenediamine tetraacetic acid (EDTA) for 10 minutes. The hope was that EDTA would strip at least some of the copper out of the corrole ligands, leading to a smaller reduction peak and confirming that copper had been present previously. However, the decrease in the putative copper reduction current is small enough to be inconclusive.
88
Figure 5.4: Azide-modified glassy carbon disk reacted with (C6F5)2(p-O(CH2CCH)Ph)corroleCu, initially (green) and after exposure to an aqueous saturated EDTA solution for 10 minutes (pink, dashed). 0.1 NBu4PF6 in acetonitrile, 1000 mV/s.
Figure 5.5: Azide-modified glassy carbon disk reacted with (C6F5)2(p-O(CH2CCH)Ph)corroleCu, initially (green) and after exposure to an aqueous saturated solution of disodium diethyldithiolcarbamate for 10 minutes (gold, dashed). 0.1 NBu4PF6 in acetonitrile, 1000 mV/s.
-1000
-500
0
500
1000
1500
-3 -2 -1 0 1 2
Copper corrole on electrode surfaceAfter EDTA exposure
Potential (V vs. Fc)
Cur
rent
(A
)
Cu reduction?
oxygen reduction
-400
-200
0
200
400
600
-3 -2 -1 0 1 2
Copper corrole on electrode surfaceAfter diethyldithiolcarbamate exposure
oxygen reduction
Cu reduction?
Potential (V vs. Fc)
Cur
rent
(A
)
89
A similar voltammogram of a different disk, also in acetonitrile, shows a more significant decrease in the current attributed to copper reduction (Figure 5.5). In this case, cyclic voltammetry was performed on the disk before and after it was exposed to diethyldithiolcarbamate for 10 minutes. A more significant change in the putative copper reduction is seen here, although copper oxidation is still not seen. No conclusions can be drawn as of yet, and further study and optimization of this process is ongoing.
5.3 Conclusions
Azide- and alkyne-functionalized corrole ligands have been prepared, and bimetallic biscorrole complexes have been synthesized via Huisgen azide-alkyne cycloaddition chemistry. The biscorrole complexes have been characterized by cyclic voltammetry and electronic independence of the metal centers has been demonstrated. Additional studies further indicate that peripheral ligand changes do not substantially affect the redox potential of the metal centers in the metallocorrole complexes studied. Preliminary studies suggest that copper (5,15-bis(pentafluorophenyl)-10-(4-propargyloxyphenyl)corrole) can be attached to an azide-modified glassy carbon surface, but additional control studies, especially those excluding oxygen, must be performed before any firm conclusions can be reached.
5.4 Experimental Methods and Materials
Ferrocene was obtained commercially and purified by sublimation before use. Tetra-n-butylammonium tetrafluoroborate (NBu4BF4) was obtained commercially and purified by dissolving in ethyl acetate, precipitating with pentane, filtering, and drying in vacuo overnight. All other reagents used were obtained commercially and used as received. Cyclic voltammetry was performed in high-purity acetonitrile or dichloromethane, dried and degassed using a Phoenix solvent drying system commercially available from JC Meyer Solvent Systems.
All corrole ligands and complexes were synthesized by Dr. Heather Buckley according to published procedures.40
Cyclic voltammetry experiments were performed using a Gamry 600 Reference potentiostat. All voltammetry experiments used 0.1 M tetraethylammonium tetrafluoroborate (NEt4BF4) in acetonitrile or 0.1 M NBu4BF4 in dichloromethane as the supporting electrolyte. Voltammetry solutions were degassed with nitrogen and kept under a nitrogen blanket during analysis. A three-electrode cell was used, with a glassy carbon working electrode (3 mm diameter), a platinum mesh counter electrode and a silver ion non-aqueous reference electrode. The reference electrode was prepared by immersing a silver wire in a 0.01 M solution of silver nitrate in 0.1 M NEt4BF4 in acetonitrile, and was separated from the bulk solution by a Vycor frit. The reference electrode potential was checked periodically by comparison with ferrocene, and measured potentials were corrected as necessary. Potentials are reported vs. Ag/Ag+ (0.01 AgNO3), which is 0.548 V vs. NHE and −0.087 V vs. ferrocene (Fc).43
Azide-modified glassy carbon disks were prepared44 according to previously published procedures.25 The modified disks were placed in a dichloromethane solution containing copper(III) 5,15-bis(pentafluorophenyl)-10-(4-propargyloxyphenyl)corrole, copper iodide, and triethylamine, and allowed to react in a nitrogen-filled glovebox. After 1 hour, the disks were
90
removed from the glovebox and sonicated in dichloromethane for 10 minutes to remove physisorbed material. They were then rinsed with methanol and dried under nitrogen. Voltammetry of the functionalized disks was performed with a platinum disk counter electrode and a silver wire pseudo-reference electrode, referenced to ferrocene. 0.1 M tetra-n-butylammonium hexafluorophosphate (NBu4PF6) in acetonitrile was used as the supporting electrolyte. Copper stripping tests were performed with saturated aqueous solutions of ethylenediaminetetraacetic acid (EDTA) and disodium diethyldithiolcarbamate.
5.5 References
1. Kadish, K. M.; Frémond, L.; Ou, Z.; Shao, J.; Shi, C.; Anson, F. C.; Burdet, F.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Cobalt(III) Corroles as Electrocatalysts for the Reduction of Dioxygen: Reactivity of a Monocorrole, Biscorroles, and Porphyrin-Corrole Dyads. J. Am. Chem. Soc. 2005, 127, 5625–5631.
2. Kadish, K. M.; Shao, J.; Ou, Z.; Frémond, L.; Zhan, R.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Electrochemistry, Spectroelectrochemistry, Chloride Binding, and O2 Catalytic Reactions of Free-Base Porphyrin-Cobalt Corrole Dyads. Inorg. Chem. 2005, 44, 6744–6754.
3. Collman, J. P.; Kaplun, M.; Decréau, R. A. Metal Corroles as Electrocatalysts for Oxygen Reduction. Dalton Trans. 2006, 554–559.
4. Kadish, K. M.; Frémond, L.; Burdet, F.; Barbe, J.-M.; Gros, C. P.; Guilard, R. Cobalt(IV) Corroles as Catalysts for the Electroreduction of O2: Reactions of Heterobimetallic Dyads Containing a Face-to-Face Linked Fe(III) or Mn(III) Porphyrin. J. Inorg. Biochem. 2006, 100, 858–868.
5. Kadish, K. M.; Shen, J.; Frémond, L.; Chen, P.; El Ojaimi, M.; Chkounda, M.; Gros, C. P.; Barbe, J.-M.; Ohkubo, K.; Fukuzumi, S.; Guilard, R. Clarification of the Oxidation State of Cobalt Corroles in Heterogeneous and Homogeneous Catalytic Reduction of Dioxygen. Inorg. Chem. 2008, 47, 6726–6737.
6. Kadish, K. M.; Frémond, L.; Shen, J.; Chen, P.; Ohkubo, K.; Fukuzumi, S.; El Ojaimi, M.; Gros, C. P.; Barbe, J.-M.; Guilard, R. Catalytic Activity of Biscobalt Porphyrin-Corrole Dyads toward the Reduction of Dioxygen. Inorg. Chem. 2009, 48, 2571–2582.
7. Dogutan, D. K.; Stoian, S. A.; McGuire, Jr., R.; Schwalbe, M.; Teets, T. S.; Nocera, D. G. Hangman Corroles: Efficient Synthesis and Oxygen Reaction Chemistry. J. Am. Chem. Soc. 2011, 133, 131–140.
8. Schechter, A.; Stanevsky, M.; Mahammed, A.; Gross, Z. Four-Electron Oxygen Reduction by Brominated Cobalt Corrole. Inorg. Chem. 2012, 51, 22–24.
9. Wang, Z.; Lei, H.; Cao, R.; Zhang, M. Cobalt Corrole on Carbon Nanotube as a Synergistic Catalyst for Oxygen Reduction Reaction in Acid Media. Electrochim. Acta 2015, 171, 81–88.
10. Aviv, I.; Gross, Z. Iron Porphyrins Catalyze the Synthesis of Non-Protected Amino Acid Esters from Ammonia and Diazoacetates. Chem. Commun. 2006, 43, 4477–4479.
11. Mahammed, A.; Mondal, B.; Rana, A.; Dey, A.; Gross, Z. The Cobalt Corrole Catalyzed Hydrogen Evolution Reaction: Surprising Electronic Effects and Characterization of Key Reaction Intermediates. Chem. Commun. 2014, 50, 2725–2727.
12. Barbe, J.-M.; Canard, G.; Brandès, S.; Jérôme, F.; Dubois, G.; Guilard, R. Metallocorroles as Sensing Components for Gas Sensors: Remarkable Affinity and Selectivity of Cobalt(III) Corroles for CO vs. O2 and N2. Dalton Trans. 2004, 1208–1214.
91
13. Santos, C. I. M.; Oliveira, E.; Fernández-Lodeiro, J.; Barata, J. F. B.; Santos, S. M.; Faustino, M. A. F.; Cavaleiro, J. A. S.; Neves, M. G. P. M. S.; Lodeiro, C. Corrole and Corrole Functionalized Silica Nanoparticles as New Metal Ion Chemosensors: A Case of Silver Satellite Nanoparticles Formation. Inorg. Chem. 2013, 52, 8564–8572.
14. Santos, C. I. M.; Oliveira, E.; Barata, J. F. B.; Faustino, M. A. F.; Cavaleiro, J. A. S.; Neves, M. G. P. M. S.; Lodeiro, C. New Gallium(III) Corrole Complexes as Colorimetric Probes for Toxic Cyanide Anion. Inorg. Chim. Acta 2014, 417, 148–154.
15. Agadjanian, H.; Ma, J.; Rentsendorj, A.; Valluripalli, V.; Hwang, J. Y.; Mahammed, A.; Farkas, D. L.; Gray, H. B.; Gross, Z.; Medina-Kauwe, L. K. Tumor Detection and Elimination by a Targeted Gallium Corrole. Proc. Natl. Acad. Sci. 2009, 106, 6105–6110.
16. Haber, A.; Aviram, M.; Gross, Z. Variables That Influence Cellular Uptake and Cytotoxic/Cytoprotective Effects of Macrocyclic Iron Complexes. Inorg. Chem. 2012, 51, 28–30.
17. Haber, A.; Ali, A. A. Y.; Aviram, M.; Gross, Z. Allosteric Inhibitors of HMG-CoA Reductase, the Key Enzyme Involved in Cholesterol Biosynthesis. Chem. Commun. 2013, 49, 10917–10919.
18. Liang, X.; Mack, J.; Zheng, L. M.; Shen, Z.; Kobayashi, N. Phosphorus(V)-Corrole: Synthesis, Spectroscopic Properties, Theoretical Calculations, and Potential Utility for in Vivo Applications in Living Cells. Inorg. Chem. 2014, 53, 2797–2802.
19. Haber, A.; Gross, Z. Catalytic Antioxidant Therapy by Metallodrugs: Lessons from Metallocorroles. Chem. Commun. 2015, 51, 5812–5827.
20. Cheng, X.; Shi, Z.; Glass, N.; Zhang, L.; Zhang, J.; Song, D.; Liu, Z. S.; Wang, H.; Shen, J. A Review of PEM Hydrogen Fuel Cell Contamination: Impacts, Mechanisms, and Mitigation. J. Power Sources 2007, 165, 739–756.
21. Huisgen, R. The Concerted Nature of 1,3-Dipolar Cycloadditions and the Question of Diradical Intermediates. J. Org. Chem. 1976, 41, 403–419.
22. Kolb, H. C.; Finn, M. G.; Sharpless, K. B. Click Chemistry: Diverse Chemical Function from a Few Good Reactions. Angew. Chem. Int. Ed. 2001, 40, 2004–2021.
23. Jain, S.; Reiser, O. Immobilization of Cobalt(II) Schiff Base Complexes on Polystyrene Resin and a Study of Their Catalytic Activity for the Aerobic Oxidation of Alcohols. ChemSusChem 2008, 1, 534–541.
24. Gomila, A.; Le Poul, N.; Cosquer, N.; Kerbaol, J.-M.; Noël, J.-M.; Reddy, M. T.; Jabin, I.; Reinaud, O.; Conan, F.; Le Mest, Y. Self-Induced “Electroclick” Immobilization of a Copper Complex onto Self-Assembled Monolayers on a Gold Electrode. Dalton Trans. 2010, 39, 11516–11518.
25. Stenehjem, E. D.; Ziatdinov, V. R.; Stack, T. D. P.; Chidsey, C. E. D. Gas-Phase Azide Functionalization of Carbon. J. Am. Chem. Soc. 2013, 135, 1110–1116.
26. Jee, J. E.; Cheong, J. L.; Lim, J.; Chen, C.; Hong, S. H.; Lee, S. S. Highly Selective Macrocycle Formations by Metathesis Catalysts Fixated in Nanopores. J. Org. Chem. 2013, 78, 3048–3056.
27. Gobbo, P.; Mossman, Z.; Nazemi, A.; Niaux, A.; Biesinger, M. C.; Gillies, E. R.; Workentin, M. S. Versatile Strained Alkyne Modified Water-Soluble AuNPs for Interfacial Strain Promoted Azide–Alkyne Cycloaddition (I-SPAAC). J. Mater. Chem. B 2014, 2, 1764–1769.
28. Prescher, J. A.; Dube, D. H.; Bertozzi, C. R. Chemical Remodelling of Cell Surfaces in Living Animals. Nature 2004, 430, 873–877.
92
29. Yu, M.; Price, J. R.; Jensen, P.; Lovitt, C. J.; Shelper, T.; Duffy, S.; Windus, L. C.; Avery, V. M.; Rutledge, P. J.; Todd, M. H. Copper, Nickel, and Zinc Cyclam−Amino Acid and Cyclam−Peptide Complexes May Be Synthesized with “ Click ” Chemistry and Are Noncytotoxic. Inorg. Chem. 2011, 50, 12823−12835.
30. Uttamapinant, C.; Tangpeerachaikul, A.; Grecian, S.; Clarke, S.; Singh, U.; Slade, P.; Gee, K. R.; Ting, A. Y. Fast, Cell-Compatible Click Chemistry with Copper-Chelating Azides for Biomolecular Labeling. Angew. Chem. Int. Ed. 2012, 51, 5852–5856.
31. Beahm, B. J.; Dehnert, K. W.; Derr, N. L.; Kuhn, J.; Eberhart, J. K.; Spillmann, D.; Amacher, S. L.; Bertozzi, C. R. A Visualizable Chain-Terminating Inhibitor of Glycosaminoglycan Biosynthesis in Developing Zebrafish. Angew. Chem. Int. Ed. 2014, 53, 3347–3352.
32. Gao, X.; Hannoush, R. N. Method for Cellular Imaging of Palmitoylated Proteins with Clickable Probes and Proximity Ligation Applied to Hedgehog, Tubulin, and Ras. J. Am. Chem. Soc. 2014, 136, 4544–4550.
33. Collman, J. P.; Devaraj, N. K.; Chidsey, C. E. D. “Clicking” Functionality onto Electrode Surfaces. Langmuir 2004, 20, 1051–1053.
34. Devadoss, A.; Chidsey, C. E. D. Azide-Modified Graphitic Surfaces for Covalent Attachment of Alkyne-Terminated Molecules by “Click” Chemistry. J. Am. Chem. Soc. 2007, 129, 5370–5371.
35. McCreery, R. L. Advanced Carbon Electrode Materials for Molecular Electrochemistry. Chem. Rev. 2008, 108, 2646–2687.
36. Laughlin, S. T.; Agard, N. J.; Baskin, J. M.; Carrico, I. S.; Chang, P. V.; Ganguli, A. S.; Hangauer, M. J.; Lo, A.; Prescher, J. A.; Bertozzi, C. R. Metabolic Labeling of Glycans with Azido Sugars for Visualization and Glycoproteomics. Methods in Enzymology 2006, 415, 230–250.
37. Lang, K.; Chin, J. W. Cellular Incorporation of Unnatural Amino Acids and Bioorthogonal Labeling of Proteins. Chem. Rev. 2014, 114, 4764–4806.
38. Koszarna, B.; Gryko, D. T. Efficient Synthesis of Meso-Substituted Corroles in a H2O-MeOH Mixture. J. Org. Chem. 2006, 71, 3707–3717.
40. Buckley, H. L.; Rubin, L. K.; Chromiński, M.; McNicholas, B. J.; Tsen, K. H. Y.; Gryko, D. T.; Arnold, J. Corroles That “Click”: Modular Synthesis of Azido- and Propargyl-Functionalized Metallocorrole Complexes and Convergent Synthesis of a Bis-Corrole Scaffold. Inorg. Chem. 2014, 53, 7941–7950.
41. Ou, Z.; Shao, J.; Zhao, H.; Ohkubo, K.; Wasbotten, I. H.; Fukuzumi, S.; Ghosh, A.; Kadish, K. M. Spectroelectrochemical and ESR Studies of Highly Substituted Copper Corroles. J. Porphyr. Phthalocyanines 2004, 8, 1236–1247.
42. Simkhovich, L.; Mahammed, A.; Goldberg, I.; Gross, Z. Synthesis and Characterization of Germanium, Tin, Phosphorus, Iron, and Rhodium Complexes of Tris(pentafluorophenyl)corrole, and the Utilization of the Iron and Rhodium Corroles as Cyclopropanation Catalysts. Chem. Eur. J. 2001, 7, 1041–1055.
93
43. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
44. Azide-modification studies performed with the assistance of Thomas Cook.
94
Appendix Computational Evaluation of Virtual Hydrogen Storage Compounds and
Comparison with Direct Electrode Oxidation Potentials
A.1 Effect of Anode Reaction Potential on Overall Cell Voltage
Virtual hydrogen storage compounds are designed to deliver protons and electrons directly in a fuel cell, such that the overall fuel cell voltage is no longer dependent on the voltage of the hydrogen oxidation reaction. This allows for the cell voltage to be optimized by tuning the potential of the electrochemical dehydrogenation reaction of the virtual hydrogen storage compound. The overall voltage of a cell is controlled by the voltages of the cathodic and anodic half reactions (Equation A.1).
Equation A.1
In order to obtain maximum cell voltage, the anodic reaction should occur at as negative a potential as possible. The Gibbs’ free energy for a reaction can be directly related to its reduction potential (Equation A.2), meaning a more negative ∆G° for a given half-reaction translates to a more negative oxidation potential.
Equation A.2
A.2 Computational Evaluation of Substrates
In order to maximize fuel cell voltage by lowering the redox potential of the anodic reaction, substrates and catalysts with very negative oxidation potentials are needed. Computational predictions were used to initially screen substrates, and those predicted to fulfill the necessary criteria were then evaluated further. A protocol for the prediction of redox potentials of organic compounds in acetonitrile has previously been developed1 and was used as the basis for this work. Twenty-five commercially available compounds were chosen according to structural features that would be expected to result in low oxidation potentials: nitrogen or other heteroatoms in a ring, conjugated π systems, or electron-donating substituents (selected structures, Figure A.1). The gas-phase ionization potential for each substrate and the solvation energies for each substrate and its radical cation were calculated.2 These numbers were then used to find the thermodynamic one-electron oxidation potentials in solution (Equation A.3).
Equation A.3
IP is the gas-phase ionization potential (eV), T∆S is the gas-phase entropy change at 298 K (kcal/mol), ∆Gsolv,ox is the free energy of solvation for the radical cation (kcal/mol) and ∆Gsolv,red is the free energy of solvation for the uncharged species (kcal/mol). The free energy change for the NHE half-reaction is 4.44 eV.1 Computed oxidation potentials vs. NHE were converted to
95
potentials vs. 0.01 M Ag/Ag+ by subtracting 0.548 V.3 Experimental data were also obtained for ten of these compounds for comparison. Selected computed and experimental potentials are listed in Table A.1.
Figure A.1: Structures for selected proposed virtual hydrogen storage compounds.
Table A.1: Computed and experimental first oxidation potentials for proposed virtual hydrogen storage compounds
The absolute correlation between computed and experimental potentials was poor, likely due to overpotentials in the experimental data. However, the trends among computed and experimental values correlate reasonably well, indicating that this is a useful screening tool.
96
A.3 Experimental Methods and Materials
Tetraethylammonium tetrafluoroborate (NEt4BF4) was obtained commercially and recrystallized from ethanol. All other reagents used were obtained commercially and used as received. Solutions for electrochemical analysis were made using HPLC-grade acetonitrile and 0.1 M NEt4BF4 as the supporting electrolyte.
Cyclic voltammetry experiments were performed using a Gamry 600 Reference potentiostat. Voltammetry solutions were deoxygenated before use and kept under a blanket of inert gas (either nitrogen or argon) during analysis. A three-electrode cell was used, with a glassy carbon working electrode (3 mm diameter), a platinum mesh counter electrode, and a non-aqueous silver ion reference electrode. The reference electrode was prepared by immersing a silver wire in a 0.01 M solution of silver nitrate in supporting electrolyte, and was kept separate from the bulk solution by a Vycor frit. Potentials are reported vs. Ag/Ag+ (0.01 AgNO3), which is 0.548 V vs. NHE and −0.087 V vs. ferrocene (Fc).3
Computations were completed with Gaussian 094 at the Molecular Graphics and Computation Facility, University of California, Berkeley. Calculations of gas-phase ionization potentials were performed at the B3LYP/6-31+G(d) level. Solvation energies were calculated at the B3LYP/6-31+G(d,p) level with the IEFPCM model, using acetonitrile as the solvent, Bondi radii, and Van der Waals surface with alpha = 1.2.
A.4 References
1. Fu, Y.; Liu, L.; Yu, H.-Z.; Wang, Y.-M.; Guo, Q.-X. Quantum-Chemical Predictions of Absolute Standard Redox Potentials of Diverse Organic Molecules and Free Radicals in Acetonitrile. J. Am. Chem. Soc. 2005, 127, 7227–7234.
2. Ochterski, J. W. Thermochemistry in Gaussian. Gaussian, Inc. [Online] 2000, http://www.gaussian.com/g_whitepap/thermo.htm (accessed Apr 2, 2011).
3. Pavlishchuk, V. V; Addison, A. W. Conversion Constants for Redox Potentials Measured versus Different Reference Electrodes in Acetonitrile Solutions at 25°C. Inorg. Chim. Acta 2000, 298, 97–102.
4. Gaussian 09, Revision B.01, Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, Ö.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian, Inc., Wallingford, CT, 2009.
