View Showing Intramolecular H-Bonds Within Diol Molecules and O-HS Interactions in the Solid State Between Neighboring Dimers

METAL-FREE ELECTROCATALYSTS FOR OXYGEN EVOLUTION REACTION AND PHOTOCATALYSTS FOR CARBON DIOXIDE REDUCTION REACTION

Usha Pandey Kadel

A Dissertation

Submitted to the Graduate College of Bowling Green State University in partial fulfillment of the requirements for the degree of

DOCTOR OF PHILOSOPHY

May 2018

Committee:

Ksenija D. Glusac, Advisor

Liangfeng Sun Graduate Faculty Representative

Alexander N. Tarnovsky

Usha Pandey Kadel

Ksenija Glusac, Advisor

Oxygen evolving reaction (OER) and carbon dioxide reduction reaction (CO2RR) are the

two important reactions related to energy conversion and storage. Both OER and CO2RR are

multi electrons and protons transfer reactions with slower reaction kinetics. Thus catalysts are

required to mediate the reactions in one step to desired products without forming the high energy

.- intermediates, such as H2O2 in case of OER and CO2 in CO2RR. In case of OER, our previous

study shows that the flavin-based catalysis depends on the type of working electrode. The oxides

formed on the electrode surface assist the evolution of the oxygen. The disadvantage of this kind

of catalysis is the difficulty in studying the catalytic mechanism by using conventional

spectroscopic techniques. The proposed catalytic mechanism for a homogeneous system involves

four different steps: (i) pseudobase formation via a reaction of xanthylium ion with water, (ii)

proton coupled electron transfer (PCET) of pseudobase to form alkoxy radicals, (iii) coupling of

alkoxy radicals to form peroxide intermediate and, (iv) oxidation of peroxide to release oxygen

and regenerate the catalyst. The result from electrode-assisted mechanism suggests that a

homogeneous catalyst can be developed, where two cation species are covalently linked with

suitable linker.

+ The xanthylium dimer (Xan )2 containing two xanthene units covalently linked with

+ diphenyl ether was studied as the model electrocatalyst for water oxidation. The (Xan )2 readily

reacts with water to form mono and di-hydroxylated species, which is the first step of proposed

catalytic mechanism. For the formation of peroxide from the pseudobase, the two OH group

should orient in favorable geometry (In-In conformer). The conformational flexibility of (Xan– iv

OH)2 was studied by using NMR spectroscopy, X-ray crystallography and, DFT calculation

methods. Although, DFT calculation and solid-state structure show the stability of In-Out

conformer, the NMR study in solution shows that the conformers freely interconvert in NMR

time scale, indicating the desired In-In conformer to be populated in millisecond timescale,

which is required for catalysis. Moreover, the electrochemical study shows that the catalytic

+ + water oxidation occurs at lower potential in case of (Xan )2 compared to Xan . However, the

+ catalytic behavior of (Xan )2 depends on the type of working electrode

In addition, the ground state hydride donating ability (hydricity) of organic hydrides,

NADH analogues (BNAH, CN-BNAH, Me-MNAH and HEH), methylene tetrahydromethanopterin analogs (BIMH and CAFH), acridine derivatives (Ph-Acr, Me2N-AcrH,

T-AcrH, 4OH, 2OH, 3NH), and a triarylmethane derivative (6OH) were studied by using

theoretical (DFT) and experimental methods (potential –pKa and hydride transfer) in two

different solvents (acetonitrile and dimethyl sulfoxide). The results show that the hydricity

values of these organic hydrides are comparable to those of metal hydrides and most of the

hydrides are capable to reduce proton in acetonitrile. v

To my parents vi ACKNOWLEDGMENTS

I would like to express my sincere appreciation to my advisor Dr. Ksenija D. Glusac for her patience, guidance, and continuous support during my PhD. I would never have been what I am today without her support and encouragement. I would like to take an opportunity to thank my committee members Dr. R. Marshall Wilson, Dr. Alexendar N. Tarnovsky, and Dr.

Liangfeng Sun for their support, time and advice. I am also grateful to Dr. Thomas H Kinstle for his invaluable support, encouragement and advice.

I also like to thank all my past and present lab members with whom I got an opportunity to work. Special thanks go to Dr. Janitha Walpita for his kind support during my initial phase in the lab, Marija Zoric and Stefan Ilic for their very helpful support during research. I have learnt so much from you people. In addition, special thanks go to Dr. Xin Yang, Dr. Suja Shyam, Yun

Xie, Varun Singh, George Hargenrader, Aco Radujevic, Andrej Penavic, and Ravindra

Weerasooriya.

I take this opportunity to thank Dr. Peter Lu and his group, especially Dr. Bharat Dhital and

Achyut P Silwal for allowing me to use their instrument during my studies. I would like to thank the faculty and staff of the Chemistry Department for their help, guidance, and advice. I would like to thank Bowling Green State University for the financial support.

Finally, I would like to thank my parents, my brother Ubhar, my sister Urmila, my brother in law Dharma Raj, and my sister in law Kalpana for their unconditional love and support. Special thanks go to my brother in law Shreedhar for his help. I appreciate the love and support from my Kadel family during my studies. Most importantly, special thanks go to my husband Gokul for his patience, love, support and encouragement. Without Gokul, none of this would be possible. vii

TABLE OF CONTENTS

Page CHAPTER 1 INTRODUCTION ...... 1

1.1 Oxygen Evolving Reaction (OER) ...... 2

1.2 Transition Metal Based Homogeneous Water Oxidation Catalysts ...... 4

1.2.1 Ruthenium Based Catalysts ...... 4

1.2.2 Iridium Based Catalysts ...... 9

1.2.3 Manganese Based Catalysts ...... … 11

1.2.4 Cobalt Based Catalysts ...... 12

1.2.5 Iron and Copper Based Catalysts ...... 14

1.3 Metal Free Water Oxidation Catalysts ...... 16

1.4 Carbon dioxide Reaction Reactions (CO2RR) ...... 17

1.4.1 Metal Complexes ...... 19

1.4.2 Metal Free Catalysts ...... 21

1.5 Our Aim ...... 23

1.5.1 Electrocatalytic Water Oxidation Project ...... 23

1.5.2 Carbon Dioxide Reduction Project ...... 25

1.6 References ...... 26

CHAPTER 2 CONFORMATIONAL FLEXIBILITY OF XANTHENE BASED COVALENTLY LINKED DIMERS ...... 36

2.1 Covalently Liked Dimers ...... 37

2.2 Experimental Section ...... 38

2.2.1 General Methods ...... 38 viii

2.2.2 Synthesis of Compounds ...... 38

2.2.3 Computational Methods ...... 40

2.2.4 Variable Temperature NMR Spectroscopy (VT-NMR) ...... 40

2.3 Result and Discussion ...... 42

+ 2.3.1 Reaction of (Xan )2 with water ...... 42

2.3.2 Conformational Flexibility of Molecules ...... 44

+ 2.3.2.1 Conformational Study and Stability of (Xan )2 ...... 45

2.3.2.2 Conformation Study of (Xan-OH)2 ...... 48

2.3.2.3 Conformation Study of Xan+-Xan-OH ...... 54

2.4 Conclusions ...... 59

2.5 References ...... 59

CHAPTER 3 ELECTROCATALYTIC WATER OXIDATION BY USING COVALENTLY LINKED XANTHENIUM DIMERS ...... 64

3.1 Electrochemical Techniques ...... 65

3.1.1 Cyclic Voltammetry ...... 65

3.1.2 Study of Reaction Mechanism ...... 71

3.1.3 Bulk Electrolysis ...... 73

3.2 Experimental Section ...... 74

3.2.1 General Methods ...... 74

3.2.2 Cyclic Voltammetry ...... 74

3.2.3 Controlled Potential Electrolysis ...... 75

+ 3.2.4 Chemical Oxidation of (Xan )2 ...... 76

3.3 Result and Discussion ...... 76

3.3.1 UV-Vis Spectroscopy: Pseudobase Xanthylium ion Equilibrium ...... 76

3.3.2 Cyclic Voltammetry ...... 77

3.3.3 Electrode Type Dependence ...... 79

+ 3.3.4 Bulk Electrolysis of (Xan )2 ...... 82

+ 3.3.5 Chemical Oxidation of (Xan )2 ...... 83

3.4 Conclusion ...... 85

3.5 References ...... 86

CHAPTER 4 THERMODYNAMIC HYDRICITIES OF BIOMIMETIC ORGANIC HYDRIDE DONORS ...... 88

4.1 Hydricity ...... 89

4.2 Experimental Methods ...... 94

4.2.1 General Methods ...... 94

4.2.2 Cyclic Voltammetry ...... 95

4.2.3 Hydride Transfer Studies ...... 96

4.2.4 pKa Determination ...... 97

4.2.5 Computational Methods ...... 99

4.3 Result and Discussion ...... 102

4.3.1 Calculated Hydricity Values ...... 102

4.3.2 Experimental Hydricity Values ...... 105

4.3.2.1 Potential pKa Method ...... 106

4.3.2.2 Hydride Transfer Method ...... 111

4.3.3 Comparison of Hydricity Values Obtained from Calculated and Experimental Methods ...... 113

4.3.4 Structural Effect and Solvent Effect on Hydricity values ...... 114

4.3.5 Comparison with Metal-based Analogs ...... 115

4.4 Conclusions ...... 118 x

4.5 References ...... 118 xi

LIST OF FIGURES

Figure Page 2+ 1.1 (a) Structure of Ru-Hbpp catalyst [Ru2(OH)2(bpp)(tpy)2] developed by Llobert and group, (b) Structure of ruthenium complex developed by Thummel and coworkers ...... 7

1.2 (a) Fe complexes containing taml ligand reported by Collins and coworkers, (b) Cu complex reported by Lin and coworkers ...... 16

1.3 Nitrogen doped graphene showing different types of nitrogen ...... 17

2.1 General representation of dynamic NMR spectra showing the peak broadening and coalescence with the change in temperature ...... 41

+ 2.2 (a) UV-Vis absorption spectra of (XanOH)2 (blue, pH=7 aqueous solution), Xan -XanOH + (red, pH=0.95 aqueous solution), and (Xan )2 (orange, pH=0.5 aqueous solution). (b) pH- + + dependent absorption of monomer Xan (black) and dimer (Xan )2 (orange) monitored at λ=378 nm...... 44

+ 1 13 2.3 (a) Structure of (Xan )2 with labels for protons and carbons, (b), (c) H and C NMR spectra + + of (Xan )2 respectively in deuterated acetonitrile. 2D spectra of (Xan )2 in deuterated acetonitrile, (d) COSY, (e) HSQC, and (f) HMBC ...... 46

2.4 (a) Representation of crystal structure of (Xan-OH)2, (b) An alternative view showing intramolecular H-bonds within diol molecules and O-HS interactions in the solid state between neighboring dimers ...... 52

1 13 2.5 (a) Structure of (Xan-OH)2 with labels for protons and carbons, (b), (c) H and C NMR

spectra of (Xan-OH)2 respectively in deuterated acetonitrile. 2D spectra of (Xan-OH)2 in deuterated acetonitrile, (d) COSY, (e) HSQC and (f) HMBC ...... 53

2.6 1H NMR peaks of compounds in deuterated acetonitrile with triflic acid, (a) Pure (Xan-

OH)2, (b) (XanOH)2 with 0.75 equivalent of acid which shows the mixture of (XanOH)2 and + + Xan -XanOH, (c) (XanOH)2 with 1 equivalent of triflic acid that forms pure Xan -XanOH), + (d) (XanOH)2 with 2 equivalents of triflic acid that forms pure (Xan )2 ...... 55 xii

1 + 2.7 H spectra of Xan -Xan-OH in CD3CN at different concentrations: 0.03 mM (violet), 0.3 mM (blue), 3 mM (green) and 30 mM (red) ...... 56

2.8 (a) Structure of Xan+-XanOH with labels for proton and carbon, (b) 1H NMR spectra of Xan+-XanOH at 25°C (blue), 0°C (green) and െ25°C (red) in deuterated acetonitrile, (c) 1H NMR spectrum of Xan+-XanOH at െ25°C with assigned protons, 2D spectra of Xan+- XanOH at െ25°C in deuterated acetonitrile, (d) COSY (e) HSQC and (f) HMBC of Xan+- XanOH ...... 58

3.1 A typical voltammogram of the reversible electron transfer showing the parameters cathodic

peak potential (ܧ௣௖), cathodic peak current (݅௣௖), anodic peak potential (ܧ௣௔), anodic peak

current ൫݅௣௔൯ǡ and peak separation (οܧ௣) ...... 66

3.2 A typical cyclic voltammogram of the irreversible electron transfer, where the return peak is absent ...... 69

3.3 A typical cyclic voltammogram for the quasi-reversible electron transfer, where the peak separation deviates from 59 mV/n...... 70

3.4 pH dependent absorption spectra of xanthylium ions: (a) Xan+ in aqueous solution at different pH values [pH 5 (black), pH 1.5 (red), pH 1.1 (blue), pH 0.75 (magenta), and pH + 0.4 (cyan)], (b) (xan )2 in aqueous solution at different pH values [pH 7 (navy), pH 4.2 (dark yellow), pH 2.6 (pink), pH 0.95 (cyan), pH 0.44 (blue), pH 0.1 9 (red), and pH –0.5 (black)] ...... 77

+ + 3.5 Cyclic voltammograms of 1.5 mM Xan (black), (Xan )2 (red) and DPE (blue) in acetonitrile with 0.1 M TBAP as electrolyte, scan rate 100 mv/s, Electrodes: platinum working

electrode, Ag/AgNO3 reference electrode, and platinum wire counter electrode. Scan + + direction + 0.5 V → +2.5 V → –1 V → +0.5 V for Xan (black), (Xan )2 (red) and the + 0.5 V → +2.5 V → +0.5 V to DPE ...... 78

+ 3.6 (a) Charge density of (Xan )2, (b) Charge density (blue) and spin density (red) of one + electron oxidized (Xan )2 ...... 79 xiii

+ + 3.7 Cyclic voltammograms of 1.5 mM Xan (Figure A) and (Xan )2 (Figure B) in acetonitrile containing 0.1 M TBAP as the electrolyte; on (a) Pt electrode, (b) Au electrode, (c) BDD electrode, and (d) BDD electrode in both cases. (Scan rate: 100 mV/s. Scan direction: blue curve: +0.5 → –1 V → +0.5 V; red curve: +0.5 → +2.5 → –1 V → +0.5 V; black curve:

baseline scan for the full range. The Icat/Id is the ratio of currents at +2.5 V and –0.01 V for + + (Xan )2, +2.5 V and –0.14 V for Xan ...... 81

+ 3.8 Cyclic voltammograms of 1.5 mM (Xan )2 in acetonitrile using 0.1 M TBAP as electrolyte. Scan direction: red curve +0.3 V→ +2.5 V→ –1 V→ +0.3 V; blue curve, +0.3 V→ –1 V→ +0.3 V ...... 83

+ 3.9 (a) UV-Vis spectra showing the decomposition of (Xan )2 with 1000 equivalent of sodium + persulfate, (b) UV-Vis spectra of (Xan )2 in 3:1(ACN: H2O) pH 0.5 solution at 60 °C .... 84

3.10 Amount of oxygen detected by GC in the presence of 2.7 M persulfate only (black) and + 2.5 mM (Xan )2 in the presence of 2.7 M persulfate (red). The experiment was performed in pH 0.5 solution of 3:1 (water: acetonitrile) mixture at 60 °C ...... 85

4.1 pKa determination of 6OH by using indicator anion method: (a) Generation of indicator anion ሺȂሻ by the reaction of indicator (XanH) in DMSO with potassium dimsyl, (b) Beer’s law plot for determination of the extinction coefficient of indicator anion at maximum absorbance (505 nm), (c) Titration of 24 mM 6OH with indicator anion ...... 107

4.2 Cyclic voltammograms of model compound (cations) in the cathodic range: Pt working

electrode, Pt counter electrode, and non-aqueous Ag/AgNO3 reference electrode. Scan rate, 0.1 V/s (3N+, CN-BNA+, HE+, BNA+, Me-MBNA+, BIM+ and CAF+), 25 V/s (6O+, T-Acr– + + + + + , Me2N-Acr , 2O ), 2 kV/s (T-Acr ), 100 V/s (Ph-Acr , 4O ); electrolyte: 0.1 M TBAP in DMSO ...... 109

4.3 1H NMR spectra showing the reaction between BNAH and 2O+ in DMSO immediately and at equilibrium ...... 112

4.4 Comparison between the calculated and experimental hydricities for the NADH-analogs ...... 114

4.5 The comparison of energy required to regenerate the hydride forms of metal-based

Ph Ph + ([Ni(P 2N )2H] ) and metal-free (2OH) hydride donors ...... 117 xiv

LIST OF SCHEMES

Scheme Page 1.1 Proposed peroxide bond formation mechanism by (a) water nucleophilic attack (WNA), (b) Radical coupling of alkoxides (12M) ...... 4

4+ 1.2 (a) Structure of cis,cis-[(bpy)2(H2O)Ru(μ-O)Ru(H2O)(bpy)2] (blue dimer) reported by T.J Meyer in 1982, (b) Proposed mechanism of water oxidation by using blue dimer as catalyst ...... 5

1.3 (a) Single site Ru complexes developed by Thummel and coworkers, (b) Ru complexes reported by Berlinguette and coworkers, (c) Mechanism of water oxidation by single site Ru complex proposed by T.J Meyer ...... 9

1.4 (a) Iridinium aqua complexes reported by Bernhard and coworkers, (b) Iridinium complexes reported by Crabtree and coworkers, (c) Proposed mechanism of water oxidation by iridium complexes...... 10

1.5 (a) Dinuclear Mn complex reported by Crabtree and group, (b) Proposed mechanism of water oxidation by using dinuclear Mn complex ...... 12

1.6 (a) β-octafluoro CoIII xanthene hangman corrole, bearing 5,15-bis-(pentafluorophenyl) substituents reported by Nocera, (b) Proposed mechanism of water oxidation catalyzed by fluorinate Co porphyrinem ...... 14

1.7 Photocatalytic mechanism of CO2 reduction reported by Ishitani and coworkers ...... 21

+ 1.8 Mechanism of CO2 reduction by pyrH (a) Carbamate-like radical intermediate formation method, (b) Hydride transfer from the hydride formed on electrode surface. (c)

Dihydropyridine formation method which serves as hydride source for the reduction of CO2 ...... 23

1.9 (a) Electrode assisted mechanism of water oxidation by Et-Fl+, (b) Proposed mechanism of water oxidation by fully molecular catalyst ...... 25

1.10 Schematic representation of photocatalytic reduction of CO2 to methanol ...... 26

2.1 Structure of xanthene based covalently linked dimers used in the study ...... 36 xv

+ 2.2 Formation of mono-hydroxylated (Xan -Xan-OH) and di-hydroxylated (Xan-OH)2 + pseudobases from the reaction of (Xan )2 with water ...... 43

+ + 2.3 Decomposition of (Xan )2 at elevated temperature to form F-Xan ...... 48

2.4 Decomposition of Xan+-XanOH to form F-XanOH at room temperature ...... 58

+ + 3.1 Molecular Structures of compounds, (Xan )2 and Xan ...... 64

4.1 Structures of compounds; NADH analogues (BNAH, CN-BNAH, Me-MNAH and HEH), methylene tetrahydromethanopterin analogs (BIMH and CAFH), acridine derivatives (Ph-

Acr, Me2N-AcrH, T-AcrH, 4OH, 2OH, 3NH), and a triarylmethane derivative (6OH) ....89

4.2 Relationship between the thermodynamic hydricityοܩுష, bond dissociation free energy (BDFE), acidity (pKa) and reduction potentials ...... 90

4.3 Hydricity determination by using potential-pKa method ...... 91

4.4 Hydricity determination by using hydride transfer method ...... 92

4.5 Hydricity determination by using hydrogen heterolysis method ...... 92 xvi

LIST OF TABLES

Table Page + 2.1 Optimized low energy structure of (Xan )2 with ߙ ൌ ͹ͶǤ͸ι and the transition state structure + of (Xan )2 for rotation around the dihedral angle α at the IEFPCM (acetonitrile) B3LYP/6- 311G* level of theory...... 47

2.2 Structures of three (Xan-OH)2 conformers with respective energies optimized at the B3LYP/6-311G level of theory ...... 49

2.3 Selected geometry parameters for the calculated conformers and the X-ray structure of

(XanOH)2 ...... 50

2.4 Optimized low energy structure of (Xan-OH)2 and the transition state structure of (Xan-

OH)2 for rotation around the dihedral angle α at the IEFPCM (acetonitrile) B3LYP/6-311G* level of theory...... 54

ି ڄ ڄ ା 4.1 Calculated standard reduction potentials (vs. NHE) for ܴ Ȁܴ and ܴ Ȁܴ , ݌ܭ௔ values for

RH and ߂ܩுష for RH in different solvents ...... 103

4.2 Indicators and their pKa values in DMSO, model compounds with experimentally determined pka values ...... 107

଴ ଴ Ǥ 4.3 Experimentally obtained ܧೃశǡ ܧ ೃ and ݌ܭ௔ of model compounds for obtaining experimental ష ೃǤ ೃ

߂ܩுି values using “potential-pKa” method in dimethyl-sulfoxide ...... 110

4.4 Hydricity of compounds determined by using hydride transfer method ...... 112 1

CHAPTER 1 INTRODUCTION

The world energy demand is increasing rapidly with the increase in population and industrialization. It is assumed that the growth in world energy demand in coming decades will be very large, an increase of as much as 73% by 2035.1 Most of the present energy requirement is

fulfilled by fossil fuels, which are not enough to meet this growing demand, so an alternative

source of renewable energy source is required. Also, the burning of the fossil fuels produces

2 greenhouse gas such as CO2 which leads to global warming. There are various sources of

alternative energy, such as wind energy, bio-mass, geothermal, nuclear energy, solar energy etc.

Among these alternative energy sources, the solar energy is the most promising approach since the

amount of energy that the earth surface gets in one hour is sufficient for whole world to utilize for

one year.3 However, direct use of sunlight is not practical as there is fluctuation in the availability of solar energy. Therefore, conversion as well as storage of solar energy is required. This can be achieved using photoelectrochemical fuel–forming cells that consist of photoelectrodes and electrocatalysts. The role of photoelectrodes is to harvest sunlight and generate charge carriers, whereas electrocatalysts facilitate the fuel forming processes [such as oxygen evolution reaction

(OER) and the carbon dioxide reduction reaction (CO2RR)]. In case of OER, sunlight is used to

split the most abundant water source on earth’s surface into hydrogen and oxygen. Hydrogen can

be stored as fuel and oxygen can be used as oxidant. The combination of hydrogen and oxygen in

fuel cells generates electricity whenever required. In case of CO2RR, the CO2 get reduced to

different products (formic acid, carbon monoxide, methanol etc.) which can be used as fuels, the

4 best one being the liquid product that can be stored and transport easily. Both OER and CO2RR

are multielectron and proton transfer reactions with sluggish kinetics. Thus, catalysts are required

.- to avoid the higher energy intermediates such as H2O2 in case of OER and CO2 in case of CO2RR. 2

My research focuses on the investigation of metal free catalysts for both OER and CO2RR.

Our previous study on simple N(5)-ethyl flavinium ion (Et-Fl+) as electrocatalyst for water

oxidation, it was found that the catalysis depend on the working electrode suggesting the

heterogeneous nature of catalysis where the oxides formed on electrode surface plays role in the

O–O bond formation.5 The mechanism of such heterogeneous system is difficult to study so it is desirable to investigate the homogeneous catalysts for water oxidation whose mechanism can be studied by using conventional spectroscopic techniques. Moreover, the hydricity of several metal hydrides were studied for the potential application of these metal hydrides for the reduction of proton and CO2 to fuels. However, studies have not been done to study the hydricity of organic

hydrides. Thus, it is interesting to study the hydricity of several organic hydrides and know whether

they are capable to reduce proton and CO2 in different solvents. Related to the research work, homogeneous electrocatalysts that have been investigated for OER and photocatalysts/electrocatalysts for CO2RR have been reviewed in the following text respectively.

1.1 Oxygen Evolving Reaction (OER)

In nature, light is used as an energy source to oxidize water to oxygen in the enzyme–

cofactor complex Photosystem II (PSII) and the water oxidation is catalyzed by the oxygen

evolving complex, a tetramanganese cluster that contains a calcium atom.6 The water oxidation

process involves four proton-coupled electron transfer (PCET) steps that avoids the catalytic

system from accumulating high charges.6 For the efficient water oxidation, the artificial process also requires the similar PCET mechanism. The reaction for electrochemical water oxidation is represented as equation (1) at pH 0.

+ - 2H2O → O2 + 4H + 4e E°=1.23 V Vs NHE (1) 3

The standard electrode potential (E°) for water oxidation is 1.23 V vs Normal Hydrogen Electrode

(NHE).7 In addition to this thermodynamic barrier, the kinetics of formation of oxygen is very slow which leads to the requirement of electrode potential higher than 1.23 V vs NHE. This extra potential relative to standard electrode potential required to proceed the reaction is termed as an overpotential. Moreover, the water oxidation process can proceed through one or two electron transfer steps forming the high energy intermediates (equation 2,3).8 Thus, catalysts are required to achieve the formation of oxygen at a reasonable rate and overpotential avoiding these high energy intermediates.

• + - H2O → OH + H + e E°=2.84 V Vs NHE (2)

+ - 2H2O → H2O2 + 2H + 2e E°=1.77 V Vs NHE (3)

Several research groups are investigating the molecular water oxidation catalysts. The homogeneous WOCs are favorable for mechanistic studies, due to the following advantages: (i)

Mechanism of catalytic cycle can be studied by using simple spectroscopic techniques, (ii) Several compounds can be designed as synthetic organic chemistry is the developed area of research, (iii)

The molecular properties of catalyst can be easily tuned for better performances. Although, molecular catalysts have many advantages, they are less robust thus easily undergoes degradation in harsh conditions required for water oxidation. Research has been focused on the study of stability as well as the catalytic mechanism pathways of metal based homogeneous water oxidation catalyst. One of the important step that need to be understood is the formation of O-O bond.9 The investigation so far on this step has divided the mechanism of water oxidation in two different possibilities.

4 i) Water Nucleophilic Attack (WNA): In this mechanism, water acts as nucleophile and attack on the higher valent metal oxo species, which results in the cleavage of the metal-oxo π bond and formation of O-O bond. The metal center undergoes the 2e– reduction forming the metal hydroperoxide species, the oxidation of this metal hydroperoxide species eventually release O2 with the regeneration of catalyst (Scheme 1.1a). ii) Radical coupling method (12M): In this mechanism, the two metal oxo species with oxygen centered radical couple to give the metal peroxides (M-O-O-M). The oxidation of thus formed peroxides liberate O2 and regenerate the catalyst (Scheme 1.1b).

Scheme 1.1 Proposed peroxide bond formation mechanism by (a) water nucleophilic attack

(WNA). (b) Radical coupling of alkoxides (12M).

1.2 Transition Metal Based Homogeneous Water Oxidation Catalysts

1.2.1 Ruthenium Based Catalys

The first discovery of homogeneous catalyst was made by Meyer and coworkers and the complex was dinuclear μ-oxo-bridged ruthenium complex cis,cis-[(bpy)2(H2O)Ru(μ-

4+ 10 O)Ru(H2O)(bpy)2] (Scheme 1.2a), commonly known as the “blue dimer” due to its characteristic blue color. This complex is active for water oxidation in the presence of cerium ammonium nitrate (CAN) as oxidants with a maximum turn over number (TON) of 13.211 and turn over frequency (TOF) of 0.0042 s-1.12 This study shows that the multi electron water oxidation to oxygen is possible and the presence of the μ-oxo-bridge in the complex helps to stabilize the higher

valent metal complex by electronic delocalization.13 Moreover, the catalytic rates of blue dimer

was increased up to 30 times by using electron transfer mediator.14

4+ Scheme 1.2 (a) Structure of cis,cis-[(bpy)2(H2O)Ru(μ-O)Ru(H2O)(bpy)2] (blue dimer) reported

by T.J Meyer in 1982. (b) Proposed mechanism of water oxidation by using blue dimer as

catalyst.15

The mechanism of water oxidation by blue dimer has been studied over the past 30 years.

Extensive mechanistic knowledge has been gained by using experimental as well as theoretical approaches. Meyer and coworkers proposed the mechanism that involves four PCET steps which give the higher valent intermediate (RuV- RuV). One of these higher valent Ru centers is then

attacked by a H2O molecule, generating hydroperoxido species, which undergoes intramolecular

15 oxidation by the second ruthenium center, resulting in the release of O2 (Scheme 1.2b). Hurst

and coworkers also purposed the mechanism which involves the formation of high energy

V V intermediate (Ru - Ru ) as in case of Meyer’s mechanism, but the nucleophilic attack of H2O is

not on one of the Ru center but on the α-position of bipyridine ligand.16 This mechanism was not

well accepted for O2 evolution however this provides information on the low stability of the catalyst due to decomposition of ligand framework.

Many approaches have been investigated to increase the activity of blue dimer by changing the μ-oxo bridge to some other bridging ligand, which could bring the two Ru centers to close proximity. Llobet and group reported the Ru-Hbpp catalyst without μ-oxo bridge

2+ [Ru2(OH)2(bpp)(tpy)2] [Hbpp = 2-2- (1-H- pyrozole-3,5-dilyl) dipyridine宠季孫Figure 1.1a) is季

capable of oxidizing water to oxygen. They used rigid pyrazol ligand as the bridge instead of μ-

oxo bridge, which force the two metal centers in cis conformation. The results shows that the use

of this rigid ligand increases the oxygen evolution rate three times faster than blue dimer with TON

of 17.5.17 The mechanistic study shows that the peroxide bond formation follows the 12M

mechanism in contrast to the WNA mechanism in case of blue dimer. By using the solid support

and optimizing the various conditions the TON up to 250 was obtained.18 Moreover, Lolbet and

coworkers show that the subtle change in the ligand architecture can induce the geometrical

changes in the arrangement of Ru-OH units which leads to the different mechanism of O-O bond

formation.19 Thummel and coworkers synthesized several dinuclear Ru complex containing the polypyridyl based ligand as the bridge between the two metal centers (Figure 1.1b).20 They

introduce several axial ligands and study the role of these ligands on the catalytic activity. Their

result shows that these complexes show the enhanced catalytic activity compared to blue dimer

with the TON up to 689 in the presence of CeIV as oxidant.20

2+ Figure 1.1 (a) Structure of Ru-Hbpp catalyst [Ru2(OH)2(bpp)(tpy)2] developed by Llobert and

group.19 (b) Structure of ruthenium complex developed by Thummel and coworkers.20

One of the interesting finding in case of the dinuclear ruthenium complex containing the

redox active quinone unit is that the deposition of this complex 实Ru2(OH)2(3,6-t-

2+ Bu2qui)2(btpyan)] (3,6-t-Bu2qui=3,6-di-tert-butyl-1,2 benzoquinone; btpyan =

1,8bis(2,2’:6’2”terpyridyl) anthracene) on ITO electrode enhances the O2 evolution rate with TON

as high as 33,500.21 The drawback of most of the catalysts designed for water oxidation is that the

oxidation need to be driven by a strong oxidant such as CeIV. The use of CeIV is not appropriate as

CeIII has low absorption in visible range, the photochemical generation of CeIV is not possible.

Thus, the redox potential of catalysis should be reduced in order to use the oxidants which can be

3+ 2+ photogenerated, such as [Ru(bpy)3] can be generated in visible region from [Ru(bpy)3] . The

redox potential of catalysis can be tuned by changing the ligand, the incorporation of negatively

charged ligand in the metal complexes stabilized the metal center reducing the redox potential.22

The Akermark and Sun groups synthesized the dinuclear ruthenium complex containing two

carboxylic groups attached to the metal centers and using aromatic ring at the 4,5-positions of the

pyridazine moiety as the linker between two metal centers that favors the cis geometry of the

complex.23 This complex has the lower redox potential of ~1.2 V vs NHE at neutral pH, but

3+ lowering the redox potential does not favor the use of [Ru(bpy)3] as oxidant. However, using

CeIV as oxidant the TON of > 10000 and TOF of 1.2 s-1 was achieved in acidic pH.

The significant investigation from Thummel and coworkers revealed that a single site is

sufficient for water oxidation catalysis.24 These single site Ru complexes consist a tridentate

polypyridyl type ligand, 2,6-di(1,8 naphthyridin-2-yl)pyridine with uncoordinated naphthyridine

nitrogen (Scheme 1.3a). The hydrogen bonding between the uncoordinated nitrogens with the aqua

ligand stabilized the single site metal complexes. Berlinguette and coworkers investigated the

catalytic activity of [Ru(tpy)(bpy)Cl]+ complexes with their corresponding aqua complexes

[Ru(tpy)(bpy)(OH2)]2+ (Scheme 1.3b).25 Their result shows that the aqua complexes have a higher

catalytic activity compared to chloride ligated complexes and the substituents in the periphery of

both tpy and bpy ligand frameworks influence catalytic activity of those Ru complexes. The

mechanistic approach of such single site Ru complexes was later reported by T.J. Meyer based on

both thermodynamics and kinetics evidences.26 The reported mechanism involves the formation

V 3+ IV II 2+ of [Ru =O] intermediate from electrochemical or chemical oxidation (Ce ) of [Ru -OH2]

(Scheme 2c). This intermediate rapidly reacts with water to give a peroxide complex, [RuIII-

OOH]2+, via the WNA mechanism. The one electron oxidation of this peroxide intermediate gives seven-coordinate Ru center with a bidentate peroxide species [RuIV(OO)]2+ which undergoes the

subsequent decomposition releasing oxygen and regenerating the catalyst.

Scheme 1.3 (a) Single site Ru complexes developed by Thummel and coworkers.24 (b) Ru

complexes reported by Berlinguette and coworkers.25 (C) Mechanism of water oxidation by single

site Ru complex proposed by T.J Meyer.26

1.2.2 Iridium Based Catalysts

Another class of WOCs are the iridium based catalysts which have been extensively studied for the past few years. Bernhard and coworkers first reported a series of cyclometalated iridium(III) aqua complexes (Scheme 1.4a) as WOC in 2008.27 The water oxidation ability of such

catalysts were evaluated by chemical oxidation of water using cerium ammonium phosphate and

ceric triflate (Ce(OTf)4) as oxidants. The results showed that the catalytic activity depends on the

type of the ligands that binds to the metal center. Electrochemical studies showed that the Ir(IV)/Ir(III)

oxidation waves of these compounds had a range from 1.20 to 1.74 V vs. NHE and highest TON

of 2760 was obtained. Later on, Crabtree and coworkers developed highly active mononuclear

iridium WOCs (Scheme 1.4b) by using stronger electron donating Cp* ligand (pentamethyl–

cyclopentadienyl) to improve the catalysis, however in case of compound containing the

10 pyrimidine ligand the catalytic activity is lower possibly due to the protonation of uncoordinated nitrogen in acidic medium.28

In order to get the better understanding of single site Ir complexes catalysis, their better performance and stability, Crabtree and coworkers investigated the water oxidation by using several half sandwiched Ir complexes in the presence of CeIV as oxidant.29 Also, the O–O bond formation step was studied by using computational method. The results revealed that the initial rate of O2 evolution depend on the Ir coordinated environment suggesting that the chelates did not dissociate upon reaction and the catalysts are more active in basic medium. The computational study suggest that the O–O bond formation proceed by the nucleophilic attack of H2O to higher valent IrV=O species (Scheme 1.4c).29

Scheme 1.4 (a) Iridinium aqua complexes reported by Bernhard and coworkers.27 (b) Iridinium complexes reported by Crabtree and coworkers.28 (c) Proposed mechanism of water oxidation by iridium complexes.29

The stability of the catalysts is the major challenge for the Ir- based WOCs. The oxidative

decomposition products such as iridium oxides are active for water oxidation and the observed

activity may be due to these oxides.28,30 In order to increase the stability of such Ir complexes,

several groups are putting efforts on the investigation of appropriate ligands which are stable

during the catalysis.31,32

1.2.3 Manganese Based Catalysts

Based on the knowledge gained from natural oxygen evolving complex of photosystem

(PS) II, many biomimetic manganese compounds have been investigated as the catalysts for water oxidation. Naruta and coworkers first developed the Mn porphyrin dimers as WOC,33 the

irreversible increase in current at potentials >1.4 V vs. NHE was obtained in the aqueous

acetonitrile solution which shows the catalytic activity affording up to a TON of 9.2.34 The authors

proposed that the O–O bond formation proceed through the nucleophic attack of OH– on a dimeric

V IV IV 34 Mn =O species generating a Mn –O–O–Mn species from which the O2 is liberated. The

3+ Crabtree group showed that the dimeric [(tpy)(OH2)Mn-(μ-O)2Mn(OH2)(tpy)] complex (Scheme

1.5a, tpy=2,2’:6’,2”–terpyridine) mediates water oxidation in the presence of NaOCl as oxidant

with TON of 4 and TOF of 0.0033 s-1.35 The crystal structure of the complex shows the mixed

valent dimer with one Mn center in +III state and the other in +IV state. Later studies on the same

– IV 36 complex showed that the oxidation can be driven with oxone (HSO5) or Ce as oxidants. Based

on the UV-Vis spectroscopy and electron paramagnetic experiments (EPR), the mechanism of

water oxidation was proposed in the presence of NaClO4 as oxidant. The proposed mechanism

IV, IV involves the initial oxidation of dimeric Mn complex to form a Mn2 species, which generate

MnV=O species by reacting with oxidant. This higher valent species leads to the O-O bond

12 formation via a nucleophilic attack of OH–. The formed Mn-peroxo species undergoes oxidation

37 which results in the release of O2 with the regeneration of catalyst (Scheme 1.5b).

Scheme 1.5 (a) Dinuclear Mn complex reported by Crabtree and group. (b) Proposed mechanism of water oxidation by using dinuclear Mn complex.35

Also, manganese bis(porphyrin) system was reported as catalyst for oxygen evolution.33,34

Recently, a number of other manganese compounds have been reported as water oxidation catalysts by using oxone or cerium as oxidants. Sun and coworkers reported the single site Mn complex as WOC in basic medium and experimental evidence shows that the O–O bond formation occurs by nucleophilic attack of hydroxyl on higher valent metal MnV=O intermediate.38

1.2.4 Cobalt Based Catalysts

Several mononuclear as well as binuclear cobalt complexes have been reported as the electrocatalysts for water oxidation. The work by Nocera and group shows that the hangman cobalt

corroles are OER catalysts, β-octafluoro CoIII xanthene hangman corrole, bearing 5,15-bis-

(pentafluorophenyl) substitutients (Scheme 1.6a) is the most efficient WOC with the modest

overpotential.39 Moreover, another Co corrole complex reported for water oxidation was

[Co(tpfc)(pyr)2], {tpfc = 5,10,15-tris(pentafluorophenyl)corrole,pyr=pyridine}. The catalytic

activity of this corrole complex gets enhanced with increasing phosphate concentration indicating

the PCET for O–O bond formation.40 Also, Berlinguette and coworkers reported the Co complex

2+ with pyridine based pentacoordinated ligand py5 [Co(py5)(OH2)] as a catalyst for water

oxidation.41 The advantage of using py5 ligand is to make the complex more stable to withstand the harsh water oxidation conditions. The authors reported the O–O bond formation step as the nucleophilic attack of water on a high valent CoIV–Oxo/hydroxo species which was formed after

the initial two redox steps.41 Moreover, a series of cationic Co-based porphyrin complexes were

reported as WOCs, complex with electron-deficient ligand being the most efficient catalyst.42 The

mechanistic investigation shows that the key species of the catalytic system can be CoIV–oxo

complex containing an oxidized radical porphyrin ligand (Scheme 1.6b) where the nucleophilic

42 attack by H2O takes place to form the O–O bond formation. In addition to the mononuclear

complexes, there are few dinuclear Co complexes that have been reported as the potential

electrocatalysts for water oxidation.43,44 Recently, Nocera and coworkers reported the

intramolecular radical coupling mechanism of O–O bond formation in dinuclear Co complexes

by using differential electrochemical mass spectrometry (DEMS) analysis.45

Scheme 1.6 (a) β-octafluoro CoIII xanthene hangman corrole, bearing 5,15-bis-

(pentafluorophenyl) substituents reported by Nocera.39 (b) Proposed mechanism of water oxidation catalyzed by fluorinate Co porphyrine.42

1.2.5 Iron and Copper Based Catalysts

In search of the least expensive water oxidation catalysts by using earth abundant metals, the first reported iron complex catalysts are the five iron complexes [FeIII(taml]- (taml=tetraamido macrocyclic ligand) (Figure 1.8a) using CAN as oxidant by Collins and coworkers.46 The substitutions of electron donating and electron withdrawing group at the aromatic ring or the alkyl bridge shows the change in catalytic behavior, electron withdrawing character assisting the water oxidation process. The best catalyst have the initial turnover frequency of more than 1.3 s-1, however the catalyst undergoes decomposition within few seconds.46 Fillol, Costas, and coworkers also reported a series of Fe complexes containing tetra and pentadentate ligands as the active

(IV) 47 WOCs by using Ce or NaIO4 as primary oxidants. Using NaIO4 as primary oxidant, the TOFs

of > 1050 was achieved. More recently, a pentanuclear iron complex has been reported as efficient

water oxidation catalyst with TOF of 1900 s-1. 48

The first reported homogeneous cu-based catalyst was by Mayer and coworkers in 2012

49 which is a [Cu(bpy)(OH)2] complex that shows a TON of ~30. Lin and coworkers studied 4,4’-

and 6,6’-substituted bipyridine based Cu complexes and found that the 6,6’-dihydroxy-2,2’-

bipyridine complex [Cu(bpyOH)(OH)2] (Figure 1.2b) as an efficient catalyst operating at

overpotential of ~640 mV with TON ~400.50 Meyer and coworkers also reported Cu complex

51 which contains a triglycylglycine ligand (H4tgg) as an efficient WOC. The reason behind using

52 H4tgg ligand is that this ligand creates a suitable environment for the co-ordination of Cu. Llobet

and coworkers reported that CuII complexes having the tetra-anionic tetradentate amidate ligands

with a square planar geometry show the pH dependent electrocatalytic water oxidation with

overpotential up to 170 mV.53 More recently, Crabtree and group reported CuII complex

Cu(pyalk)2 (pyalk=2-pyridyl-2-propnoate) as robust catalyst for water oxidation with TOF ~ 0.7

-1 54 II,II s . Also, dinuclear complex Cu2 complex containing 2,7-[bis(2-pyridylmethyl)aminomethyl]-

1,8-naphthyridine (bpman) ligand was reported as water oxidation catalyst in neutral solution with

TOF of ~0.6 s-1 and faradaic efficiency of ~98%.55

Figure 1.2 (a) Fe complexes containing taml ligand reported by Collins and coworkers.46 (b) Cu

complex reported by Lin and coworkers.50

1.3 Metal Free Water Oxidation Catalysts

Recently significant scientific interests have been generated for the heteroatom doped nanocarbon materials (such as N, P or S-doped graphene, nanotubes and other hybrid materials) as the metal free, earth abundant electrocatalysts for OER. Those kinds of carbon nanotubes and graphenes have large surface area and excellent carrier mobility which are desired characteristics of a successful electrocatalysts. The first reported catalyst was the nitrogen doped carbon nanomaterials with TOF of 2,600 s-1.56 Later investigation shows that the nitrogen doped 2D and

3D carbon material shows OER catalysis in alkaline medium with currents and overpotentials

56-58 similar to those of most efficient metal oxides OER catalysts (IrO2/C and RuO2). Also it was

found that the dual doping with heteroatoms such as N,P59,60 and N,S61 enhance OER catalysis.

Although experimental approach has not been reported for the mechanism of the OER, the

theoretical studies have been made, which suggest the following mechanism in alkaline medium:

(e− (1 + כOH → כ + −OH

− (H2O + e− (2 + כOH → O + כOH

(e− (3 + כOH− → OOH + כO

− (H2O + e− (4 + כ + OH → O2 + כOOH

Ǥ כ where

ሺ͵ሻ כThe rate determining step for OER is the formation of OOH

ʹǤ͸כOOH

Ǥͷͺ

Figure 1.3 Nitrogen doped graphene showing different types of nitrogen.

1.4 Carbon dioxide Reduction Reaction (CO2RR)

CO2 is the most notorious greenhouse gas, released by both natural and artificial processes.

Although CO2 is required in nature for the growth of plants as well as for many industrial

processes, the production of CO2 on earth should be balanced with what is consumed in order for

the environmental stability. Unfortunately, due to intensified human industrialization, the amount

of produced CO2 is higher that the consumption, leading to global warming. Thus, conversion of

CO2 to useful materials is required. As discussed earlier, the CO2 can be reduced (photo)

18 electrochemically to carbon based fuels such as formate, CO, methanol, etc. The development of appropriate catalysts is crucial for these conversions, in part so that the reduction pathway can

͌Ѹ avoid high-energy intermediates, such as the one-electron-reduced radical anion CO2 , whose reduction potential is Ѹ1.90 V (Equation 10) vs. NHE. The catalysts are also required for selectivity as (photo)electrochemical reduction of CO2 can proceed through two, four, six and eight electron pathways as shown in equations (5 –9). The better fuels being less explosive liquid with higher energy density that can be transported easily.4

+ – CO2 + 2H + 2e → CO + H2O, E0 = – 0.53 V (5)

+ – CO2 + 2H + 2e → HCO2H, E0 = – 0.61 V (6)

+ – CO2 + 4H + 4e → HCHO + H2O, E0 = – 0.48 V (7)

+ – CO2 + 6H + 6e → CH3OH + H2O, E0 = – 0.38 V (8)

+ – CO2 + 8H + 8e → CH4 + H2O, E0 = – 0.24 V (9)

– .– CO2 + e → CO , E0 = – 1.90 V (10)

The photoelectrochemical reduction of CO2 with a light absorbing material is attractive because sunlight could be the energy input to drive the reaction. Thus, the strategy to combine semiconductor electrodes with molecular CO2 reduction catalysts is the promising approach to utilize the solar energy and the selective reduction of CO2 to different products. The homogeneous catalysts investigated for such strategy have been reviewed briefly in the following text.

1.4.1 Metal Complexes

The first reported investigation of photoelectrochemical reduction of CO2 was by Bradley

et al. in 1982. They studied the series of tetraazomacrocyclic NiII and CoII complexes at p-Si

photocathodes in 0.1 M TBAP/MeCN.63 The more improved reduction was observed by using

II 2+ [Me6[14]aneN4Ni ] as the electron transfer mediator in 1:1 acetonitrile/water solution which

64 give CO and H2 in the ratio of 2:1 at -1 V Vs SCE. The reaction was found to be solvent

dependent, in aqueous solvents (DMSO and DMF) the Faradaic efficiency of <5-50% was

2– obtained for CO whereas in dry solvents CO and CO3 were formed in equal amounts at potential

lower than –1.3 V Vs SCE. Kubiak and coworkers reported [Re(bpy- tBut)(CO)3Cl] (bpy-tBut =

4,4′-di-tert-butyl-2,2′-bipyridine) as electrocatalyst for CO2 reduction to CO at hydrogen

terminated p-type silicon.65 The reduction was achieved at potential 600 mV more positive than at a Pt electrode with Faradaic efficiency of 97±3%. Also, several Re complexes with different ligands framework which serves both as a catalyst as well as photosensitizer were reported as

66-68 + photocatalysts for selective reduction of CO2 to CO. [ Re (bpy) (CO) 3{P(OEt) 3 } ]

(bpy=2,2’) bipyridine) was reported as the most efficient photocatalyst for the selective reduction

69 of CO2 to CO with quantum yield of 0.38.

Several groups investigated the mechanism of CO2 reduction by using fac-

I n+ 70,71 [Re (bpy)(CO)3(L)] as photocatalysts. It was found that the first step of catalytic mechanism

was the formation of one electron reduced species of Re complexes which were produced by the

triplet metal to ligand charge transfer 3(MLCT) excited state of metal complexes quenched by

tertiary amine such as triethanolamine (TEOA). However, the source of second electron transfer

and the reaction between one electron reduced species with CO2 was not fully understood. In 2007,

Ishitani and coworkers reported the complete mechanism of CO2 reduction to CO by using several

– –, – 72 Re complexes [Re(bpy)(CO)3L] (L=SCN , Cl CN). Among the studied complexes,

[Re(bpy)(CO)3SCN] was found to be the most efficient catalyst with quantum yield of 0.30. The mechanism shows that the one electron reduced species was formed (as discussed earlier), which

– give the precursor of CO2 reduction by elimination of SCN ligand (scheme 1.7). Thus formed

precursor reacts with CO2 to form CO2 adduct, the one electron reduced species play important

role as electron donor to CO2 adduct which results in CO production with the regeneration of catalyst. Moreover, photoelectrochemical reduction of CO2 was studied by using Gallium arsenide

and Gallium phosphide (p-GaAs, p-GaP) as semiconductors in the presence of Ni(cyclam)2+

(cyclam = 1,4,8,11-tetraazacyclotetradecane) as catalysts.73,74 More recently, it was reported that

nanographene–Re complex acts as an efficient electrocatalyst/photocatalyst for the selective

75 reduction of CO2 to CO. This complex reduces CO2 electrochemically to CO at potential –0.42

vs. NHE in tetrahydrofuran, which is the least negative potential reported so far for the molecular

catalysts.

72 Scheme 1.7 Photocatalytic mechanism of CO2 reduction reported by Ishitani and coworkers.

1.4.2 Metal Free Catalysts

In case of metal-free approaches doped carbon materials76-78 and small molecular model catalyst (pyridinium ion) have been studied.79 In the doped carbon catalysts, the overpotential is reported to be similar to that of Ag catalyst and is more highly stable than metal based catalyst76,77 where graphitic and pyridinic nitrogens have been reported as the active sites for catalysis.76,78 The theoretical study shows that the rate determining step in such N-doped material is the formation of

species by the first coupled electron- proton transfer to CO2 that bind to the pyridinic כCOOH nitrogen and eventually leads to the formation of CO.80 Also, recent study shows that the use of ionic liquid as electrolyte, N-doped materials undergo a high degree reduction to methane.81

Moreover, it has been reported that the boron82 and nitrogen doped diamond83 act as catalysts for

3 CO2RR to formaldehyde and acetate where the sp hybridized carbon acts as active sites for

catalysis.

In case of molecular model catalyst, several groups reported the pyridine based molecules

79,84,85 86 as potential (photo)electrocatalysts for CO2 reduction to methanol. Bocarsly group

+ reported that the protonated pyridine (pyrH ) reduce CO2 to methanol electrochemically on pd

electrode79 as well as photoelectrochemically on p-Gap photocathode.87 Three mechanistic study

+ approachs for (photo)electrochemical CO2 reduction by pyrH were suggested. The first

mechanism involves the sequence of one electron transfer method with the formation of

carbamate-like radical as intermediate (Scheme 1.8a).88 The mechanism was not well accepted as

there was discrepancy between the experimental (–0.53 V vs. SCE) and computational (–1.44 V

vs. SCE) values of reduction potential pyrH+ to form carbamate-like radical.89 The second

mechanism suggests the role of surface hydride formed on the electrode surface where CO2 is

reduced by the proton coupled hydride transfer from electrode surface activated by pyrH+.90 In this

mechanism, pyrH+ acts as Bronstrated acid which facilitates the formation of surface hydride and protonation of reduced CO2 (Scheme 1.8b). The third mechanism involves the formation of dihydropyridine (DHP) from the reduction of pyrH+ by the alternating proton transfer (PT) electron

transfer (ET) mechanism (Scheme 1.8c).91 Thus formed dihydropyridine serves as the hydride

donor for the reduction of CO2 and the aromaticity of the product formed is the driving force for

such reactions.

+ Scheme 1.8 Mechanism of CO2 reduction by pyrH (a) Carbamate-like radical intermediate

formation method. (b) Hydride transfer from the hydride formed on electrode surface. (c)

Dihydropyridine formation method which serves as hydride source for the reduction of CO2.

1.5 Our Aim

This dissertation work focuses on the investigation of metal free homogeneous catalyst for water oxidation whose catalytic mechanism can be studied by using conventional spectroscopic techniques. The electrocatalytic water oxidation capability of covalently linked xanthenium ion

+ (Xan )2 was studied and to explore the role of dimeric form a comparative study of a monomeric

+ + xanthylium ion (Xan ) was investigated. The result showed that (Xan )2 can catalyze water

oxidation at above 2.1 V Vs NHE where such catalysis was not observed in case of Xan+. In case

of fuel forming CO2RR process, the ground state hydride donating ability (hydricity) of organic

hydrides were studied by using theoretical (DFT) as well as experimental methods. The hydridicity

of thirteen different organic hydrides with different structural properties were studied in two

different solvents (acetonitrile and DMSO). The structural effects and solvent effects on hydride

donating ability of compounds were observed.

1.5.1 Electrocatalytic Water Oxidation Project

Our group is interested in organic cationic structures (R+), composed of either iminium or

oxonium ion motifs, as molecular models for metal-free OER.5,92 These organic cations were

selected because they readily react with water to form hydroxylated pseudobase derivatives

(ROH), and this process can be utilized to initiate the OER catalytic cycle. Furthermore, the

R+/ROH interconversion can be achieved in a reversible fashion, which is important for the closure

of the catalytic cycle. Finally, the structures of ROH derivatives can be tailored to exhibit favorable

thermodynamics for the concerted proton-coupled one-electron oxidation to RO. radical species,

which are in turn precursors for the important O–O bond coupling step of the OER. Our initial

study on the flavin-based iminium ion, N(5)-ethyl flavinium ion (Et-Fl+) has shown that this simple

organic compound catalyzes the electrochemical OER and that the oxidized pseudobase Et-FlOH+

intermediate is a likely catalytic intermediate.5 The water oxidation process depends on the type

of electrode used suggesting the role of working electrode in the catalytic mechanism. The surface

oxides formed on the working electrode is involved in the O–O bond formation as shown in scheme

(1.9a). Further indirect support for the importance of ROH intermediate was shown in PCET study

of pseudobases.93

Therefore, it is desirable to develop the fully organic homogeneous catalyst where the O–

O bond can be formed from the coupling of the two alkoxy radicals within the catalytic moiety.

Thus, the fully organic homogeneous catalyst should have the two iminium moieties connected with suitable linker (scheme 1.9b) The proposed mechanism involves four different steps: (i)

pseudobase formation via a reaction of iminium ion with water; (ii) proton coupled electron

transfer of pseudobase to form alkoxy radicals; (iii) coupling of alkoxy radicals to form peroxide

intermediate; (iv) oxidation of peroxide to release oxygen and regenerate the catalyst.

Scheme 1.9 (a) Electrode assisted mechanism of water oxidation by Et-Fl+. (b) Proposed

mechanism of water oxidation by fully molecular catalyst.

Based on the proposal as shown in Scheme 1.9b, the electrocatalytic water oxidation

+ capability of covalently linked xanthylium ion (Xan )2 was studied where diphenyl ether is used

+ as a linker. First, the conformational flexibility of (Xan )2 was studied and then various

electrochemical experiments were performed to investigate the water oxidation capability of

+ (Xan )2 which is discussed in detail in chapter 2 and 3.

1.5.2 Carbon Dioxide Reduction Project

In case of fuel forming CO2RR process, the steps of dye sensitized photoelectrochemical

reduction of CO2 to methanol was proposed. The aspects of proposed mechanism are as follows

(Scheme 1.10): (i) The system contains the dual light harvesting method composed of

semiconductor and dye (NAD+ analog); (ii) The organic molecule serves both as dye sensitizer

+ (NAD analog) as well as CO2 reduction catalyst (NADH analog); (iii) the rate of catalysis is

expected to increase by the two electrons and a proton coupled reduction of NAD+ analog to

NADH analog by using a single photon. Our group previously reported the results on the efficient

photoreduction of NAD+ analog dyes on p-type gallium phosphide (p-Gap).94 The other important

parameter that need to be understood in case of such solar energy driven fuel forming reactions is

the ground state hydride donating ability (hydricity) of organic hydrides. In this dissertation work,

the hydricity of organic hydrides (NADH analogs, methylene tetrahydromethanopterin analogs

and acridine derivatives) were studied by using theoretical and experimental methods which is

discussed in detail in chapter 4.

Semiconductor Semiconductor

Scheme 1.10 Schematic representation of photocatalytic reduction of CO2 to methanol.

1.6 References

(1) Schneider, M.; Froggatt, A. 2013.

(2) Armaroli, N.; Balzani, V. Angew. Chem. Int. Ed. 2007, 46, 52.

(3) Sun, S.-S.; Sariciftci, N. S. Organic photovoltaics: mechanisms, materials, and devices; CRC press, 2005.

(4) Singh, M. R.; Clark, E. L.; Bell, A. T. Proc. Natl. Acad. Sci. U.S.A. 2015, 112,

E6111.

(5) Mirzakulova, E.; Khatmullin, R.; Walpita, J.; Corrigan, T.; Vargas-Barbosa, N. M.;

Vyas, S.; Oottikkal, S.; Manzer, S. F.; Hadad, C. M.; Glusac, K. D. Nat. Chem. 2012, 4, 794.

(6) Meyer, T. J.; Huynh, M. H. V.; Thorp, H. H. Angew. Chem. Int. Ed. 2007, 46, 5284.

(7) Bard, A. J.; Parsons, R.; Jordan, J.; CRC press, 1985; Vol. 6.

(8) Rüttinger, W.; Dismukes, G. C. Chem. Rev. 1997, 97, 1.

(9) Romain, S.; Bozoglian, F.; Sala, X.; Llobet, A. J. Am. Chem. Soc. 2009, 131, 2768.

(10) Gersten, S. W.; Samuels, G. J.; Meyer, T. J. J. Am. Chem. Soc. 1982, 104, 4029.

(11) Collin, J.; Sauvage, J. P. Inorg. Chem. 1986, 25.

(12) Nagoshi, K.; Yamashita, S.; Yagi, M.; Kaneko, M. J. Mol. Catal. A: Chem.

1999, 144, 71. 28

(13) Doppelt, P.; Meyer, T. J. Inorg. Chem. 1987, 26, 2027.

(14) Concepcion, J. J.; Jurss, J. W.; Templeton, J. L.; Meyer, T. J. Proc. Natl. Acad. Sci.

U.S.A. 2008, 105, 17632.

(15) Liu, F.; Concepcion, J. J.; Jurss, J. W.; Cardolaccia, T.; Templeton, J. L.; Meyer,

T. J. Inorg. Chem. 2008, 47, 1727.

(16) Cape, J. L.; Hurst, J. K. J. Am. Chem. Soc. 2008, 130, 827.

(17) Sens, C.; Romero, I.; Rodríguez, M.; Llobet, A.; Parella, T.; Benet-Buchholz, J. J.

Am. Chem. Soc. 2004, 126, 7798.

(18) Mola, J.; Mas‐Marza, E.; Sala, X.; Romero, I.; Rodríguez, M.; Viñas, C.; Parella,

T.; Llobet, A. Angew. Chem. 2008, 120, 5914.

(19) Neudeck, S.; Maji, S.; López, I.; Meyer, S.; Meyer, F.; Llobet, A. J. Am. Chem.

Soc. 2013, 136, 24.

(20) Deng, Z.; Tseng, H.-W.; Zong, R.; Wang, D.; Thummel, R. Inorg. Chem. 2008, 47,

1835.

(21) Wada, T.; Tsuge, K.; Tanaka, K. Inorg. Chem. 2001, 40, 329.

(22) Kärkäs, M. D.; Åkermark, B. The Chemical Record 2016, 16, 940.

(23) Xu, Y.; Fischer, A.; Duan, L.; Tong, L.; Gabrielsson, E.; Åkermark, B.; Sun, L.

Angew. Chem. Int. Ed. 2010, 49, 8934.

(24) Zong, R.; Thummel, R. P. J. Am. Chem. Soc. 2005, 127, 12802.

(25) Wasylenko, D. J.; Ganesamoorthy, C.; Koivisto, B. D.; Henderson, M. A.;

Berlinguette, C. P. Inorg. Chem. 2010, 49, 2202.

(26) Concepcion, J. J.; Jurss, J. W.; Templeton, J. L.; Meyer, T. J. J. Am. Chem. Soc.

2008, 130, 16462.

(27) McDaniel, N. D.; Coughlin, F. J.; Tinker, L. L.; Bernhard, S. J. Am. Chem. Soc.

2008, 130, 210.

(28) Hull, J. F.; Balcells, D.; Blakemore, J. D.; Incarvito, C. D.; Eisenstein, O.; Brudvig,

G. W.; Crabtree, R. H. J. Am. Chem. Soc. 2009, 131, 8730.

(29) Blakemore, J. D.; Schley, N. D.; Balcells, D.; Hull, J. F.; Olack, G. W.; Incarvito,

C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 2010, 132, 16017.

(30) Savini, A.; Belanzoni, P.; Bellachioma, G.; Zuccaccia, C.; Zuccaccia, D.;

Macchioni, A. Green Chem. 2011, 13, 3360.

(31) Brewster, T. P.; Blakemore, J. D.; Schley, N. D.; Incarvito, C. D.; Hazari, N.;

Brudvig, G. W.; Crabtree, R. H. Organometallics 2011, 30, 965.

(32) Cao, R.; Ma, H.; Geletii, Y. V.; Hardcastle, K. I.; Hill, C. L. Inorg. Chem. 2009,

48, 5596.

(33) Naruta, Y.; Sasayama, M. a.; Sasaki, T. Angew. Chem. Int. Ed.1994, 33, 1839.

(34) Shimazaki, Y.; Nagano, T.; Takesue, H.; Ye, B. H.; Tani, F.; Naruta, Y. Angew.

Chem. Int. Ed. 2004, 43, 98.

(35) Limburg, J.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 1997, 119, 2761.

(36) Limburg, J.; Vrettos, J. S.; Liable-Sands, L. M.; Rheingold, A. L.; Crabtree, R. H.;

Brudvig, G. W. Science 1999, 283, 1524.

(37) Limburg, J.; Vrettos, J. S.; Chen, H.; de Paula, J. C.; Crabtree, R. H.; Brudvig, G.

W. J. Am. Chem. Soc. 2001, 123, 423.

(38) Gao, Y.; Liu, J.; Wang, M.; Na, Y.; Åkermark, B.; Sun, L. Tetrahedron 2007, 63,

1987.

(39) Dogutan, D. K.; McGuire Jr, R.; Nocera, D. G. J. Am. Chem. Soc. 2011, 133, 9178.

(40) Lei, H.; Han, A.; Li, F.; Zhang, M.; Han, Y.; Du, P.; Lai, W.; Cao, R. PCCP 2014,

16, 1883.

(41) Wasylenko, D. J.; Ganesamoorthy, C.; Borau-Garcia, J.; Berlinguette, C. P. Chem.

Commun. 2011, 47, 4249.

(42) Wang, D.; Groves, J. T. Proc. Natl. Acad. Sci. 2013, 110, 15579.

(43) Rigsby, M. L.; Mandal, S.; Nam, W.; Spencer, L. C.; Llobet, A.; Stahl, S. S. Chem.

Sci. 2012, 3, 3058.

(44) Luo, J.; Rath, N. P.; Mirica, L. M. Inorg. Chem. 2011, 50, 6152.

(45) Ullman, A. M.; Brodsky, C. N.; Li, N.; Zheng, S.-L.; Nocera, D. G. J. Am. Chem.

Soc. 2016, 138, 4229.

(46) Ellis, W. C.; McDaniel, N. D.; Bernhard, S.; Collins, T. J. J. Am. Chem. Soc. 2010,

132, 10990.

(47) Fillol, J. L.; Codolà, Z.; Garcia-Bosch, I.; Gómez, L.; Pla, J. J.; Costas, M. Nat.

Chem. 2011, 3, 807.

(48) Okamura, M.; Kondo, M.; Kuga, R.; Kurashige, Y.; Yanai, T.; Hayami, S.;

Praneeth, V. K.; Yoshida, M.; Yoneda, K.; Kawata, S. Nature 2016, 530, 465.

(49) Barnett, S. M.; Goldberg, K. I.; Mayer, J. M. Nat. Chem. 2012, 4, 498.

(50) Zhang, T.; Wang, C.; Liu, S.; Wang, J.-L.; Lin, W. J. Am. Chem. Soc. 2013, 136,

273.

(51) Zhang, M.-T.; Chen, Z.; Kang, P.; Meyer, T. J. J. Am. Chem. Soc. 2013, 135, 2048.

(52) Kim, M.; Martell, A. J. Am. Chem. Soc. 1966, 88, 914.

(53) Garrido-Barros, P.; Funes-Ardoiz, I.; Drouet, S.; Benet-Buchholz, J.; Maseras, F.;

Llobet, A. J. Am. Chem. Soc. 2015, 137, 6758.

(54) Fisher, K. J.; Materna, K. L.; Mercado, B. Q.; Crabtree, R. H.; Brudvig, G. W. ACS

Catal. 2017, 7, 3384.

(55) Su, X. J.; Gao, M.; Jiao, L.; Liao, R. Z.; Siegbahn, P. E.; Cheng, J. P.; Zhang, M.

T. Angew. Chem. Int. Ed. 2015, 54, 4909.

(56) Zhao, Y.; Nakamura, R.; Kamiya, K.; Nakanishi, S.; Hashimoto, K. Nat. Commun.

2013, 4, 2390.

(57) Tian, G. L.; Zhao, M. Q.; Yu, D.; Kong, X. Y.; Huang, J. Q.; Zhang, Q.; Wei, F.

Small 2014, 10, 2251.

(58) Yang, H. B.; Miao, J.; Hung, S.-F.; Chen, J.; Tao, H. B.; Wang, X.; Zhang, L.;

Chen, R.; Gao, J.; Chen, H. M. Sci. Adv. 2016, 2, e1501122.

(59) Li, R.; Wei, Z.; Gou, X. Acs Catal. 2015, 5, 4133.

(60) Zhang, J.; Zhao, Z.; Xia, Z.; Dai, L. Nat. Nanotechnol. 2015, 10, 444.

(61) Zhao, J.; Liu, Y.; Quan, X.; Chen, S.; Zhao, H.; Yu, H. Electrochim. Acta. 2016,

204, 169.

(62) Li, M.; Zhang, L.; Xu, Q.; Niu, J.; Xia, Z. J. Catal. 2014, 314, 66.

(63) Bradley, M.; Tysak, T. J. electroanal. chem. interfacial electrochem. 1982, 135,

153.

(64) Bradley, M. G.; Tysak, T.; Graves, D. J.; Viachiopoulos, N. A. J. Chem. Soc.,

Chem. Commun. 1983, 349.

(65) Kumar, B.; Smieja, J. M.; Kubiak, C. P. J. Phys. Chem. C 2010, 114, 14220.

(66) Koike, K.; Hori, H.; Ishizuka, M.; Westwell, J. R.; Takeuchi, K.; Ibusuki, T.;

Enjouji, K.; Konno, H.; Sakamoto, K.; Ishitani, O. Organometallics 1997, 16, 5724.

(67) Kurz, P.; Probst, B.; Spingler, B.; Alberto, R. Eur. J. Inorg. Chem. 2006, 2006,

2966.

(68) Hawecker, J.; Lehn, J. M.; Ziessel, R. Helv. Chim. Acta 1986, 69, 1990.

(69) Hori, H.; Johnson, F. P.; Koike, K.; Ishitani, O.; Ibusuki, T. J. Photochem.

Photobiol., A 1996, 96, 171.

(70) Gibson, D. H.; Yin, X. Chem. Commun. 1999, 1411.

(71) Hayashi, Y.; Kita, S.; Brunschwig, B. S.; Fujita, E. J. Am. Chem. Soc. 2003, 125,

11976.

(72) Takeda, H.; Koike, K.; Inoue, H.; Ishitani, O. J. Am. Chem. Soc. 2008, 130, 2023.

(73) Petit, J.-P.; Chartier, P.; Beley, M.; Deville, J.-P. J. electroanal. chem. interfacial electrochem. 1989, 269, 267.

(74) Petit, J.-P.; Chartier, P.; Beley, M.; Sauvage, J.-P. New J. Chem. 1987, 11, 751.

(75) Qiao, X.; Li, Q.; Schaugaard, R. N.; Noffke, B. W.; Liu, Y.; Li, D.; Liu, L.;

Raghavachari, K.; Li, L. J. Am. Chem. Soc. 2017, 139, 3934.

(76) Wu, J.; Liu, M.; Sharma, P. P.; Yadav, R. M.; Ma, L.; Yang, Y.; Zou, X.; Zhou, X.-

D.; Vajtai, R.; Yakobson, B. I. Nano Lett. 2015, 16, 466.

(77) Kumar, B.; Asadi, M.; Pisasale, D.; Sinha-Ray, S.; Rosen, B. A.; Haasch, R.;

Abiade, J.; Yarin, A. L.; Salehi-Khojin, A. Nat. Commun. 2013, 4.

(78) Sharma, P. P.; Wu, J.; Yadav, R. M.; Liu, M.; Wright, C. J.; Tiwary, C. S.;

Yakobson, B. I.; Lou, J.; Ajayan, P. M.; Zhou, X. D. Angew. Chem. 2015, 127, 13905.

(79) Seshadri, G.; Lin, C.; Bocarsly, A. B. J. Electroanal. Chem. 1994, 372, 145.

(80) Wu, J.; Yadav, R. M.; Liu, M.; Sharma, P. P.; Tiwary, C. S.; Ma, L.; Zou, X.; Zhou,

X.-D.; Yakobson, B. I.; Lou, J. ACS nano 2015, 9, 5364.

(81) Sun, X.; Kang, X.; Zhu, Q.; Ma, J.; Yang, G.; Liu, Z.; Han, B. Chem. Sci. 2016, 7,

2883.

(82) Yang, L.; Jiang, S.; Zhao, Y.; Zhu, L.; Chen, S.; Wang, X.; Wu, Q.; Ma, J.; Ma, Y.;

Hu, Z. Angew. Chem. 2011, 123, 7270.

(83) Liu, Y.; Chen, S.; Quan, X.; Yu, H. J. Am. Chem. Soc 2015, 137, 11631.

(84) Boston, D. J.; Xu, C.; Armstrong, D. W.; MacDonnell, F. M. J. Am. Chem. Soc.

2013, 135, 16252.

(85) Lim, C.-H.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B. J. Phys. Chem. Lett. 2015,

6, 5078.

(86) Keith, J. A.; Carter, E. A. Chem. Sci. 2013, 4, 1490.

(87) Barton, E. E.; Rampulla, D. M.; Bocarsly, A. B. J. Am. Chem. Soc. 2008, 130, 6342.

(88) Lim, C.-H.; Holder, A. M.; Musgrave, C. B. J. Am. Chem. Soc. 2012, 135, 142.

(89) Keith, J. A.; Carter, E. A. J. Am. Chem. Soc. 2012, 134, 7580.

(90) Ertem, M. Z.; Konezny, S. J.; Araujo, C. M.; Batista, V. S. J. Phys. Chem. Lett.

2013, 4, 745.

(91) Lim, C.-H.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B. J. Am. Chem. Soc. 2014,

136, 16081.

(92) Yang, X.; Walpita, J.; Mirzakulova, E.; Oottikkal, S.; Hadad, C. M.; Glusac, K. D.

ACS Catl. 2014, 4, 2635.

(93) Walpita, J.; Yang, X.; Khatmullin, R.; Luk, H. L.; Hadad, C. M.; Glusac, K. D. J.

Phys. Org. Chem. 2016, 29, 204.

(94) Ilic, S.; Brown, E. S.; Xie, Y.; Maldonado, S.; Glusac, K. D. J. Phys. Chem. C.

2016, 120, 3145. 36

CHAPTER 2 CONFORMATIONAL FLEXIBILITY OF XANTHENE BASED

COVALENTLY LINKED DIMERS

Reproduce in part with permission from Zoric, M.R., Pandey Kadel, U., Korvinson, K.A., Luk,

H.L., Nimthong-Roldan, A., Zeller, M., Glusac, K.D. J. Phys. Org. Chem. 2016, 29(10), 505-513.

In the proposed catalytic cycle in chapter 1 (Scheme 1.9b), the second step involves the

PCET to form the alkoxy radicals, that couple to give the peroxide intermediate. These alkoxy radicals are highly reactive and undergo fast decomposition reactions, such as hydrogen atom abstraction, β elimination etc. Thus, to achieve the efficient coupling between the alkoxy radicals before they undergo decomposition, the two –OH groups of the pseudobase should orient towards each other in a favorable geometry. In this chapter, the conformational flexibility of covalently linked dimer consisting two xanthenium moieties connected through a diphenyl ether linker

+ + (Xan )2 and its mono-hydroxylated (Xan -XanOH) and di-hydroxylated (Xan-OH)2 derivatives

(Figure 2.1) was studied by using the NMR spectroscopy and density functional theory calculation methods.

Scheme 2.1 Structure of xanthene based covalently linked dimers used in the study. 37

2.1 Covalently Linked Dimers

Covalently-linked dimers with rigid spacers (CDs) are useful for various studies such as exciton coupling1, energy transfer efficiency,2-4 and the singlet fission efficiency.5 Research has been done to study the hindered rotation in CDs by using a range of the 2D NMR techniques to obtain information about the rotamer structures.6-14 Also, the variable temperature NMR lineshape analysis has been used to gain insight into rotamer interconversion kinetics.15-17 These findings led to the development of chiral biaryl catalysts that were utilized in a number of asymmetric reactions, such as nitro-aldol condensations,16 Michael additions, epoxidations,16 and hetero Diels-Alder reactions.18

Moreover, CDs are used for the oxygen reduction as well as water oxidation catalysis. The catalytic reduction of molecular oxygen to water is catalyzed by porphyrin dimers that exhibit two metal-porphyrin units in a co-facial arrangement and at the appropriate distance to efficiently

“bite” molecular oxygen.16,19 Similarly, the water oxidation reaction can be accelerated by two catalytic porphyrin units linked either covalently20,21 or noncovalently.20 In the case of water oxidation, the role of the CD catalyst is to bring two oxygen atoms from water molecules into an ideal geometry for O−O bond formation. While rigid spacers provide good control of interatomic distances, a certain degree of flexibility was found to be vital for accommodation of dynamic changes during the catalytic cycle.20

2.2 Experimental Section

2.2.1 General Methods

All reagents and starting materials were purchased from commercial sources and used as

received. 1H, 13C, COSY, HSQC, and HMBC NMR spectra at room temperature were recorded using a spectrometer operating at a field of 11.7 T (500 MHz for 1H and 125 MHz for 13C, Bruker-

Spectrospin 500 Ultra Shield). Low temperature NMR spectra were collected using two different

spectrometers operating at a field of 17.6 T (700 MHz for 1H and 175 MHz for 13C, Varian vnmrs

700 Ytterbium) and 9.4 T (400 MHz for 1H and 100 MHz for 13C, Oxford NMR 400. Spectra are

described through chemical shifts, multiplicity (s, singlet; d, doublet; t, triplet; q, quartet; m,

multiplet), coupling constant in hertz (Hz), and number of protons. Single crystal X-ray structures

were collected on area detector diffractometers at 100 K using monochromatic Mo or Cu KD

radiation. UV-Vis absorption spectra were taken on a HP8453 UV-vis spectrophotometer. MALDI

spectra were recorded on a Bruker Drucker Daltonics Omniflex mass spectrometer.

2.2.2 Synthesis of Compounds

22 (XanOH)2: (XanOH)2 was synthesized according to a published procedure. Diphenyl

ether (1.00 g, 6.0 mmol) was dissolved in 15 mL of anhydrous THF. The solution was purged with

argon for 30 min. n-Butyllithium solution in hexane (2.5 mol/L, 4.9 mL, and 12.3 mmol) was

added dropwise over a period of 5 min. The resulting mixture was stirred for 4 hours at room

temperature and the xanthone (2.35 g, 12.0 mmol) suspension in 30 mL of anhydrous THF was

added in portion via cannula over a period of 10 min. The reaction mixture was left to stir overnight

which formed the white precipitate. The resulting white precipitate was filtered, treated with 39 saturated aqueous ammonium chloride, extracted with dichloromethane, washed with water, and dried over sodium sulfate. The solvent was evaporated under reduced pressure to give 1.82 g (3.2

1 mmol, 54%) of the desired compound as a white solid. H NMR (500 MHz, CD3CN): δH 4.30

(2H, s), 5.98 (2H, d, J = 8.0 Hz), 6.96-7.07 (10H, m), 7.10-7.16 (4H, m), 7.23-7.27 (6H, m), 7.41

13 (2H, d, J = 7.1); C NMR (125 MHz, CD3CN): δC 70.1, 116.6, 117.1, 120.9, 123.8, 123.1, 123.9,

127.6, 127.9, 128.6, 129.0, 129.4, 129.5, 138.4, 154.7.

+ - + 22 (Xan )2(ClO4 )2: (Xan )2 was synthesized according to a published procedure. A 15 mL vial was charged with (XanOH)2 (0.30 g, 0.50 mmol) and concentrated perchloric acid (70%, 8 mL). The resulting mixture was stirred overnight. The red precipitate that formed was filtered off, washed with anhydrous ethyl ether (30 mL), and dried in vacuum to give 0.36 g (0.49 mmol, 98%)

1 of a deep red solid. H NMR (500 MHz, CD3CN): δH 7.34 (2H, d, J = 8.6 Hz), 7.39 (2H, dd, J =

1.7, 7.6), 7.49 (2H, td, J = 1.0, 7.6), 7.60 (4H, ddd, J = 1.0, 7.0, 8.2), 7.74-7.80 (6H, m), 8.33 (4H,

13 dd, J = 1.0, 8.8 Hz, H2), 8.50 (4H, ddd, J = 1.6, 6.9, 8.7); C NMR (125 MHz, CD3CN): δC 120.3,

121.2, 123.7, 125.4, 126.5, 130.6, 132.6, 133.2, 135.6, 146.4, 154.1, 159.6, 173.1.

+ + Xan-XanOH: Xan -Xan-OH was prepared by the titration of (XanOH)2 with trifluoromethanesulfonic acid (triflic acid). The formation of the Xan+-Xan-OH was monitored by proton NMR, but the product was never isolated. Disappearance of the OH proton peak was

1 observed upon addition of one equivalent of acid. H NMR (700 MHz, CD3CN, -25ºC): δH 4.38

(1H, s, OH proton), 5.31 (1H, dd, J = 1.2, 8.1 Hz,), 6.32-6.37 (1H, m, ), 6.38 (1H, dd, J = 1.8, 7.8

Hz), 6.48 (1H, d, J = 8.6 Hz), 6.71 (1H, dd, J = 1.2, 8.1 Hz), 6.79 (1H, ddd, J = 1.8, 6.9, 8.4 Hz),

6.86-6.95 (1H, m), 7.00-7.10 (2H, m), 7.21-7.30 (2H, m), 7.32 (1H, dd, J = 1.7, 7.6 Hz), 7.35-7.42

(2H, m), 7.45 (1H, td, J = 1.8, 7.9 Hz), 7.61-7.72 (1H, m), 7.97 (1H, dd, J = 1.7, 8.3), 8.03 (1H, 40 dd, J = 1.5, 8.6 Hz), 8.16 (1H, dd, J = 1.8, 7.8 Hz), 8.20 (1H, d, J = 8.7 Hz), 8.23-8.29 (1H, m),

8.33 (1H, ddd, J = 1.5, 6.8, 8.5 Hz), 8.57 (1H, d, J = 8.7 Hz), 8.81 (1H, ddd, J = 1.5, 6.8, 8.5 Hz);

13 C NMR (175 MHz, CD3CN, -25ºC): δC 69.5 , 116.3, 116.4, 118.2, 118.3, 118.4, 122.0 , 122.3,

122.5, 124.8, 125.4, 125.6, 125.7, 127.0, 127.1, 128.6, 128.7, 130.4, 130.9, 131.2, 131.3, 132.1,

132.2, 133.9, 134.8, 135.1, 136.5, 141.9, 146.7, 147.9, 150.6, 151.4, 153.2, 156.5, 159.7, 161.4,

174.5.

2.2.3 Computational Methods

All calculations were performed at the Ohio Supercomputer Center using Gaussian 09 software.23 Density Functional Theory (DFT) was used to optimize geometries at ground state.

The calculations employed the B3LYP24,25 hybrid functional and the 6-311G* basis set. Geometry optimizations were performed in the gas phase. The analysis of the vibrational frequencies for the optimized structures showed the absence of imaginary frequencies for the ground state structures, and the presence of only one imaginary frequency for transition state geometries. The energies were calculated in acetonitrile as a solvent using the IEFPCM26 solvation model.

2.2.4 Variable Temperature NMR Spectroscopy (VT-NMR)

Organic compounds generally adopt more than one conformation or configuration in a solution. This change in the molecular dynamics can be monitored by using the VT-NMR when the interconversion process is slow on the NMR timescale. VT- NMR spectroscopy has been applied extensively in kinetic studies of dynamic processes.27-29 The energy range covered by this technique is limited by the NMR time scale, signal resolution and the practical NMR temperature range, which ultimately depends on the melting and boiling points of solvents. 41

A simple example is shown in figure 2.1, where two nuclei (AB) show sharp signal at low

temperature and undergo chemical exchange resulting in line broadening and coalescence with

increase in temperature. At coalescence temperature, the interconversion rate constant is given as,

గο௏ ܭ ൌ (1) ்௖ ξଶ

where,

οܸ= the difference in chemical shift in Hz of two signals without interconversion.

ܶܿ= coalescence temperature (K)

-1 ܭ்௖= the rate constant (s )

Figure 2.1 General representation of dynamic NMR spectra showing the peak broadening and

coalescence with the change in temperature. 42

By using the Eyring equation, the gibbs free energy of interconversion can be calculated

by using the rate constant at any temperature (equation 2),

כοಸ ௄ഁ் ି ܭൌ ݁ ೃ೅ (2) ௛

where,

ܭఉ = Boltzmann constant

ܶ= absolute temperature

݄ = Planck’s constant

ܴ = the universal gas constant

2.3 Results and Discussion

+ 2.3.1 Reaction of (Xan )2 with Water

The first step of the proposed catalytic water oxidation reaction (Scheme 1.9b) is the

+ pseudobase formation via reaction of (Xan )2 with water. For comparison, the reaction of monomer

Xan+ with water was also studied. The reaction was monitored by using UV-Vis absorption

+ + spectroscopy in the pH range of 0 to 7. The results show that the cations (Xan )2 and Xan readily

react with water to form pseudobases (XanOH)2 and Xan-OH respectively. The pka of Xan-OH

was calculated to be 1.3 (Figure 2.1b), which is in agreement with already reported value.30,31 The

two pka values were obtained in case of (XanOH)2, pKa1=0.2 and pKa2=2.6 which were assigned

as the formation of mono-hydroxylated and di-hydroxylated species respectively (Scheme 2.2). 43

+ Scheme 2.2 Formation of mono-hydroxylated (Xan -Xan-OH) and di-hydroxylated (Xan-OH)2

+ pseudobases from the reaction of (Xan )2 with water.

+ The UV-Vis absorption spectrum of (Xan )2 shows that the first hydroxylation occurs at

low pH (pH=0.2) to form mono-hydroxylated derivative (Xan+-XanOH). This lower pH

+ + hydroxylation of (Xan )2 compared to monomer Xan (pH=1.3) suggests the lower stability of

+ (Xan )2 compared to its monocation analogue due to the repulsive interaction between the cationic moieties in dimer. However, the second hydroxylation of Xan+-XanOH occurs at higher pH

(pH=2.6) compared to Xan+, which suggests the higher stability of Xan+-XanOH due to the

cationic stabilization by electron donating Xan-OH moiety. Also, the pKa value of (Xan-OH)2

might be affected slightly by the statistical effect of having two OH groups.32 44

(a) 1.0 (b) 1.5

1.0 + + (Xan ) Xan

2 0.5 + + Xan -XanOH (Xan )2 (Xan-OH) Absorption 0.5 2

Absorption at 378 nm

0.0 0.0 300 400 500 01234567 Wavelength(nm) pH

+ Figure 2.2 (a) UV-Vis absorption spectra of (XanOH)2 (blue, pH=7 aqueous solution), Xan -

+ XanOH (red, pH=0.95 aqueous solution), and (Xan )2 (orange, pH=0.5 aqueous solution). (b) pH-

+ + dependent absorption of monomer Xan (black) and dimer (Xan )2 (orange) monitored at λ=378 nm.

2.3.2 Conformational Flexibility of Molecules

The second and third steps of the proposed catalytic cycle (scheme 1.9b) showed the PCET of pseudobase to form alkoxy radicals, which coupled to give peroxide intermediate. Due to highly unstable nature of alkoxy radicals, they undergo decomposition reactions such as H atom

33 abstraction, β-elimination etc. Thus, the two –OH groups of (Xan-OH)2 should point toward each other with suitable distance appropriate for peroxide bond formation. The conformational study of compounds is helpful to gain insight into the flexibility and right geometrical orientation of molecules for the formation of required peroxide bond. The conformational study of all the dimeric

+ + compounds (Xan-OH)2, (Xan )2 and Xan -XanOH was made by using the NMR spectroscopy and density functional theory, which is discussed in the following text for each compound. 45

+ 2.3.2.1 Conformational Study and Stability of (Xan )2

1 13 + The H and C NMR spectra of (Xan )2 were collected in deuterated acetonitrile at room temperature. The 1H and 13C NMR peak assignments were made based on the 2D NMR techniques

(HSQC, HMBC, and COSY). The 1H NMR spectrum shows 8 different H-atoms and 13C NMR spectrum shows 13 different C-atoms (Figure 2.3 a, b). This result suggests that either the molecule is locked in one conformation with C2V symmetry or freely interconverts in the NMR time scale

+ with some possible degree of freedom. (Xan )2 molecule can rotate through the four different dihedral angles (α, α’, β, β’) defines as C6-C7-C8-C13 (α), C6’-C7’-C8’-C13’(α’), C12-C13-O-C13’(β) and C12’-C13’-O-C13(β’). If the molecule is locked in the C2V symmetry, then there should be high steric repulsion between the protons 12 and 12’ which make the molecule unstable at that conformation. This information led to the conclusion that at room temperature the molecule is not locked in one conformation but freely rotates along the dihedral angles β/β’ within the NMR timescale. This conclusion is futher justified by the already reported reults for substituted diphenyl ethers, the barrier of interconversion between the local minima is within 10 kcal/mol which is 46 faster than the NMR timescale at room temperature.34

(d) (e) (f)

+ 1 13 Figure 2.3 (a) Structure of (Xan )2 with labels for protons and carbons. (b), (c) H and C NMR

+ + spectra of (Xan )2 respectively in deuterated acetonitrile. 2D spectra of (Xan )2 in deuterated acetonitrile, (d) COSY, (e) HSQC, and (f) HMBC. 47

Due to the C2v symmetry of xanthenium moiety in the molecule, NMR spectroscopy method was not useful to calculate the rotation barrier through the dihedral angles α, and α’.

+ Therefore, a DFT calculation method was used to optimize the structure of (Xan )2. The structure

with minimum energy was obtained with α=74.6° (Table 2.1), which is consistent with the result

shown with substituted biphenyls previously.35,36 The transition state of interconversion through

rotation of α was calculated at B3LYP/6-311G* level of theory, which shows that the transition

state molecule adopts structure with α=-172.2° (Table 2.1). The barrier of interconversion was

calculated to be ~23 kcal/mol. Such a high barrier of interconversion gives the half-life of around

107 min, thus a molecule cannot rotate through the dihedral angle α, and α’ in the NMR timescale

at room temperature.

+ Table 2.1 Optimized low energy structure of (Xan )2 with α=74.6° and the transition state structure

+ of (Xan )2 for rotation around the dihedral angle α at the IEFPCM (acetonitrile) B3LYP/6-311G*

level of theory.

Structure

Optimized structure Transition state structure

Relative energy 0 22.86 (kcal/mol) 48

+ Moreover, the (Xan )2 undergoes an intramolecular Friedel-Crafts alkylation to yield a fused F-Xan+ derivative under prolonged heating (Scheme 2.3). The Friedel-Crafts alkylation process was confirmed by the obtained X-ray structure of F-Xan+ in an attempt to get the X-ray

+ structure of (Xan )2. This type of Friedel-Crafts reactivity has been reported previously for a wide range of benzylic cations,37-39 and involves a reaction of electrophilic benzylic cations (in our case, the Xan+ moiety) towards aromatic compounds that contain electron-donating groups (in our case, a DPE moiety).

+ + Scheme 2.3 Decomposition of (Xan )2 at elevated temperature to form F-Xan .

2.3.2.2 Conformational Study of (Xan-OH)2

In the case of (XanOH)2, the relative arrangement of the two –OH groups can lead to three possible conformers namely In-In, In-out and out-out conformers (Table 2.2). For the propose of catalytic water oxidation, the In-In conformer with two –OH groups close to each other is desired for the O-O bond formation, which was already been demonstrated by Sun in catalytic water

40 oxidation by Ru-based non-covalent dimers. The most stable conformer of (XanOH)2 was 49 obtained experimentally by using X-ray crystallography in solid phase and by NMR spectroscopy in solution phase. Also, DFT calculation method was used to find the lowest energy conformer.

The DFT calculation structures obtained at the B3LYP/6-311G* level of theory show the

In-Out conformer to be the lowest energy conformer (Table 2.2). The In-In and Out-Out conformers are higher in energy by 1.94 kcal/mol and 2.58 kcal/mol respectively compared to In-

Out conformer. This theoretical calculation result indicates that only 3.6 % of (XanOH)2 are expected to adopt the In-In conformation in solution, which is the favorable conformation for catalytic activity.

Table 2.2 Structures of three (Xan-OH)2 conformers with respective energies optimized at the

B3LYP/6-311G level of theory.

Optimized

structures

In-Out In-In Out-Out

Relative energies 0 1.94 2.58 (kcal/mol)

The calculated structure of the In-In conformer exhibits one intramolecular hydrogen bond involving the two –OH groups and the oxygen atom on the bridging diphenyl-ether moiety (Table

2.3). The calculated OH…O bond distances is 1.92 Å and the OH…O angles is 162.8°, which is within the criteria generally used for the assignment of moderately strong hydrogen bonds.41 Even 50 though the energy stabilization due to the presence of moderate intramolecular hydrogen bonding is expected to be in the 4-15 kcal/mol range for each bond,41 the gas phase energy of the In-In conformer is higher than that of the In-Out conformer. Thus, it appears that most of the stabilizing effects due to the presence of the hydrogen bonds are compensated by other unfavorable geometric factors. The comparison of the dihedral angles related to the diphenyl-ether linker (DPE) shows that these angles are not in the optimal range for the In-In conformer. For example, the optimal

DPE angle values previously reported for other DPE derivatives are usually in the 25-50° range.42

While the DPE angles for the In-Out and Out-Out conformers (C12-C13-O-C13’ and C12’-C13’-O-C13,

Table 2.3) are within this optimal, the In-In conformer contains two DPE angles that are outside this optimal range (-79.2°, 6.1°). Thus, the stabilizing effect of the hydrogen bonds in the In-In conformer is compensated by the need to place the DPE angles into a less favorable geometry.

Table 2.3 Selected geometry parameters for the calculated conformers and the X-ray structure of

(XanOH)2.

Dihedral Angles Hydrogen Bonds

O14-H14-O14’ O14’-H14’-O15 C12-C13- C12’-C13’- C13-C8- C13’-C8’- Leng Angle Leng Angle O15-C13’ O15-C13 C7-O14 C7’-O14’ th th

1.93 162.8 In-In -79.2º 6.1º -23.4º 34.2º N/A N/A Å º

1.91 140.4 In-Out -47.5º -45.3º -178.5º 41.8º N/A Å º 51

Out- -52.3º -52.3º -174.0º -174.0º N/A N/A Out

1.89 146.8 X-ray -43.33º -50.11º -174.60º -2.29º N/A Å 1º

The solid-state crystal structure of (Xan-OH)2 adopts the in-out conformer (Figure 2.4), which is similar to that of the calculated in-out conformer. The only huge difference between the

X-ray structure and the calculated In-Out conformer is the dihedral angle C13’-C8’-C7’-O14’ which is -2.29° in X-ray structure whereas 41.8° in In-Out conformer. This difference in dihedral angle is possibly due to the presence of four intermolecular hydrogen bonds between the outward facing

–OH group of one molecule and the S electron density of the phenyl rings of neighboring molecules in crystal structure which results in the small dihedral angle (Figure 2.4b), whereas such intermolecular interaction is not accounted in the gas phase single molecule calculations.

(a) (b)

Figure 2.4 (a) Representation of crystal structure of (Xan-OH)2. (b) An alternative view showing intramolecular H-bonds within diol molecules and O-HS interactions in the solid state between neighboring dimers.

1 To investigate the preferred conformation of (XanOH)2 in solution, the NMR spectra ( H,

13C, COSY, HSQC, HMBC) were collected in deuterated acetonitrile. The 13C NMR shows 17 nonequivalent carbon atoms (Figure 2.5 c) suggesting a C2 point group with a C2 axis running through the oxygen atom of the DPE moiety. If all relevant dihedral angles in (XanOH)2 rotated freely on the NMR timescale, the NMR spectrum would exhibit only 13 C-atoms due to C2v symmetry. Thus, this simple inspection of the NMR spectrum reveals that rotation around one or more dihedral angles in (XanOH)2 is restricted.

(d) (e) (f)

1 13 Figure 2.5 (a) Structure of (Xan-OH)2 with labels for protons and carbons. (b), (c) H and C

NMR spectra of (Xan-OH)2 respectively in deuterated acetonitrile. 2D spectra of (Xan-OH)2 in

deuterated acetonitrile, (d) COSY, (e) HSQC and (f) HMBC.

The possible degree of freedom of the (Xan-OH)2 molecule can be defined in terms of four

different dihedral angles (α,α’,β, β’) defines as C6-C7-C8-C9 (α/ α’), C12-C13-O-C13’(β) and C12’-

C13’-O-C13(β’).The conditions that can be taken in to account on the rotation of these dihedral angles can be summarized as four different possibilities; (i) if α/α’ and β/β’ were all free to rotate, the NMR spectrum would have C2v symmetry; (ii) if α/α’ and β/β’ were all frozen at the NMR timescale, the NMR spectrum would have either C2 (for the In-In and Out-Out conformers) or C1 symmetry (for the In-Out conformer); (iii) if α/α’ were frozen and β/β’ were free to rotate, the

NMR spectrum would have either C2v (for In-In and Out-Out conformers) or Cs (for In-Out) symmetry; (iv) if α/α’ were free and β/β’ were frozen, the NMR spectrum would have C2 symmetry. The obtained NMR spectra of (Xan-OH)2 in deuterated acetonitrile shows the C2

symmetry which lead us to conclude that the dihedral angles related to the DPE moiety (β /β’) are 54 frozen on the NMR timescale and α/α’ are freely rotating. Moreover, the DFT calculation of rotation around the dihedral angle α gives the transition state with energy 11.13 kcal/mol (Table

2.3). The rate of inter-conversion between different conformers based on this transition state energy was calculated to be 4.3104 s-1 which suggests that the catalytically preferred In-In conformation can be reached readily at the microsecond timescale.

Table 2.4 Optimized low energy structure of (Xan-OH)2 and the transition state structure of (Xan-

OH)2 for rotation around the dihedral angle α at the IEFPCM (acetonitrile) B3LYP/6-311G* level of theory.

Optimized structures

Optimized structure Transition state structure

Relative energies 0 11.13 (kcal/mol)

2.3.2.3 Conformational Study of Xan+-Xan-OH

As discussed in section 2.3.1, the hydroxylation of dictation forms mono-hydroxylated and di-hydroxylated pseudobases respectively as shown by UV-Vis absorption spectra. Here, to further get insight into the intermediate (mono hydroxylated derivative), we performed the NMR spectroscopy experiments. We monitored the formation of intermediate by using 1H NMR spectra 55

in deuterated acetonitrile. In this case, the intermediate was formed by partial dehydroxylation of

the (Xan-OH)2 with triflic acid (CF3SO3H). As shown in figure 2.6, addition of 0.75 equivalent of

+ acid (Figure 2.6b) formed the mixture of (Xan-OH)2 and Xan -XanOH with two peaks in aliphatic

region (4.35 ppm and 4.15 ppm). The peak at 4.35 ppm was assigned as the –OH peak of (Xan-

+ OH)2, whereas the peak at 4.15 ppm was assigned as the –OH peaks of Xan -XanOH. Also, in the

aromatic region some extra broad peaks were formed in addition to (Xan-OH)2 peaks. The pure

intermediate was formed after adding 1 equivalent of triflic acid (Figure 2.6c) and converted to

+ pure (Xan )2 with 2 equivalents of acid (Figure 2.6d).

Figure 2.6 1H NMR peaks of compounds in deuterated acetonitrile with triflic acid. (a) Pure (Xan-

+ OH)2. (b) (XanOH)2 with 0.75 equivalent of acid which shows the mixture of (XanOH)2 and Xan -

+ XanOH. (c) (XanOH)2 with 1 equivalent of triflic acid that forms pure Xan -XanOH). (d)

+ (XanOH)2 with 2 equivalents of triflic acid that forms pure (Xan )2. 56

The peaks of the intermediate showed broadening in the aromatic region suggesting that

the solution of Xan+-XanOH undergo dynamic processes at the NMR timescale. The dynamic process can be intermolecular or intramolecular. The concentration dependent study in the concentration range of 0.03 mM to 30 mM (Figure 2.7) showed the same broadening feature of the NMR peaks indicating that the process is intermolecular. Since the barriers for rotation around dihedral angles associated with xanthenol (C5’-C6’-C14-C15) and xanthylium moieties (C5-C6-C7-

C8) were previously found to be too small or too large to cause NMR broadening at room

temperature, the observed dynamics are likely due to the rotation of the two dihedral angles

associated with the DPE linker (angles C2-C1-O-C1’ and C2’-C1’-O-C1) structure.

1 + Figure 2.7 H spectra of Xan -Xan-OH in CD3CN at different concentrations: 0.03 mM (violet),

0.3 mM (blue), 3 mM (green) and 30 mM (red). 57

The broadening of peaks due to intermolecular dynamics were studied using the variable temperature NMR Spectroscopy (VT-NMR). At lower temperature (-25°C), the peaks were sharpened and well separated. The 1H and 13C carbon atoms were assigned based on this low temperature NMR peaks where various 2D NMR techniques were applied (Figure 2.8). The VT-

NMR spectra show that at -25°C the 1H NMR peaks for proton 9 and 9’ are well separated and the coalescence of those peaks were achieved at room temperature (Figure 2.8b). This shows that the molecule is in frozen conformation at lower temperature with C1 symmetry and converts to freely rotating structure with Cs symmetry at room temperature. Based on this coalescence temperature, the barrier of conversion was calculated to be ΔG#=15 kcal/mol.

Figure 2.8 (a) Structure of Xan+-XanOH with labels for proton and carbon. (b) 1H NMR spectra of Xan+-XanOH at 25°C (blue), 0°C (green) and -25°C (red) in deuterated acetonitrile. (c) 1H NMR spectrum of Xan+-XanOH at -25°C with assigned protons. 2D spectra of Xan+-XanOH at -25°C in deuterated acetonitrile, (d) COSY (e) HSQC and (f) HMBC of Xan+-XanOH.

The intermediate was stable only for a few hours at room temperature. It undergoes an intramolecular Friedel-Crafts reaction to form F-XanOH (Scheme 2.4) as previously observed in

+ case of (Xan )2 at elevated temperature.

Scheme 2.4 Decomposition of Xan+-XanOH to form F-XanOH at room temperature. 59

2.4 Conclusions

The study of conformational flexibility of xanthene based covalently linked dimers was made based on the NMR spectroscopy, X-ray crystallography and DFT calculation methods. The rotational barrier around the xanthylium moieties were calculated to be ~23 kcal/mol whereas the xanthenol moieties rotate with less energy barrier of 11 kcal/mol. The rotational barriers around the diphenyl ether moieties of the dimer molecules depend on the substituents on the ortho position

+ + with the trend (Xan )2 > (Xan -XanOH) > (Xan-OH)2. These results show that the conformational flexibility of linker plays an important role in investigating the molecule with desired geometry for water oxidation catalysis. Moreover, the result from this study shows that the desired In-In conformer of (Xan-OH)2 can be readily accessed at room temperature which is beneficial for water oxidation catalysis.

2.5 References

(1) Osuka, A.; Maruyama, K. J. Am. Chem. Soc. 1988, 110, 4454.

(2) Jensen, K. K.; van Berlekom, S. B.; Kajanus, J.; Mårtensson, J.; Albinsson, B. J.

Phys. Chem. A. 1997, 101, 2218.

(3) Osuka, A.; Tanabe, N.; Kawabata, S.; Yamazaki, I.; Nishimura, Y. J. Org. Chem.

1995, 60, 7177.

(4) Lindquist, R. J.; Lefler, K. M.; Brown, K. E.; Dyar, S. M.; Margulies, E. A.; Young,

R. M.; Wasielewski, M. R. J. Am. Chem. Soc. 2014, 136, 14912. 60

(5) Korovina, N. V.; Das, S.; Nett, Z.; Feng, X.; Joy, J.; Haiges, R.; Krylov, A. I.;

Bradforth, S. E.; Thompson, M. E. J. Am. Chem. Soc. 2016, 138, 617.

(6) Gil, R. R. Angew. Chem. Int. Ed. 2011, 50, 7222.

(7) Wormald, M. R.; Petrescu, A. J.; Pao, Y.-L.; Glithero, A.; Elliott, T.; Dwek, R. A.

Chem. Rev. 2002, 102, 371.

(8) Seco, J. M.; Quinoá, E.; Riguera, R. Chem. Rev. 2004, 104, 17.

(9) Wagner, G.; Wüthrich, K. In nature.com 1978.

(10) Böckmann, A. Angew. Chem. Int. Ed. 2008, 47, 6110.

(11) Stewart, I. C.; Benitez, D.; O’Leary, D. J.; Tkatchouk, E.; Day, M. W.; Goddard

III, W. A.; Grubbs, R. H. J. Am. Chem. Soc. 2009, 131, 1931.

(12) Steinberg, A.; Froimowitz, M.; Parrish, D. A.; Deschamps, J. R.; Glaser, R. J. Org.

Chem. 2011, 76, 9239.

(13) Pastor, A.; Martínez-Viviente, E. Coord. Chem. Rev. 2008, 252, 2314.

(14) Sebaoun, L.; Maurizot, V.; Granier, T.; Kauffmann, B.; Huc, I. J. Am. Chem. Soc.

2014, 136, 2168.

(15) Lunazzi, L.; Mazzanti, A.; Minzoni, M. J. Org. Chem. 2007, 72, 2501.

(16) Port, A.; Moragas, M.; Sanchez-Ruiz, X.; Jaime, C.; Virgili, A.; Alvarez-Larena,

A.; Piniella, J. J. Org. Chem. 1997, 62, 899. 61

(17) Lunazzi, L.; Mancinelli, M.; Mazzanti, A. J. Org. Chem. 2008, 74, 1345.

(18) Pu, L. Chem. Rev. 1998, 98, 2405.

(19) Collman, J. P.; Denisevich, P.; Konai, Y.; Marrocco, M.; Koval, C.; Anson, F. C.

J. Am. Chem. Soc. 1980, 102, 6027.

(20) Jiang, Y.; Li, F.; Zhang, B.; Li, X.; Wang, X.; Huang, F.; Sun, L. Angew. Chem.

2013, 125, 3482.

(21) Shimazaki, Y.; Nagano, T.; Takesue, H.; Ye, B. H.; Tani, F.; Naruta, Y. Angew.

Chem. Int. Ed. 2004, 43, 98.

(22) Suzuki, T.; Nishida, J. i.; Ohkita, M.; Tsuji, T. Angew. Chem. Int. Ed. 2000, 112,

1874.

(23) Frisch, M.; Trucks, G.; Schlegel, H. B.; Scuseria, G.; Robb, M.; Cheeseman, J.;

Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. Inc., Wallingford, CT 2009, 200.

(24) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev. B. 1988, 37, 785.

(25) Becke, A. D. Phys. Rev. A. 1988, 38, 3098.

(26) Tomasi, J.; Persico, M. Chem. Rev. 1994, 94, 2027.

(27) Casarini, D.; Lunazzi, L.; Pasquali, F.; Gasparrini, F.; Villani, C. d. J. Am. Chem.

Soc. 1992, 114, 6521. 62

(28) Casarini, D.; Foresti, E.; Gasparrini, F.; Lunazzi, L.; Misiti, D.; Macciantelli, D.;

Villani, C. J. Org. Chem. 1993, 58, 5674.

(29) Dell'Erba, C.; Gasparrini, F.; Grilli, S.; Lunazzi, L.; Mazzanti, A.; Novi, M.; Pierini,

M.; Tavani, C.; Villani, C. J. Org. Chem. 2002, 67, 1663.

(30) Wada, M.; Kirishima, K.; Oki, Y.; Miyamoto, M.; Asahara, M.; Erabi, T. Bull.

Chem. Soc. Jpn. 1999, 72, 779.

(31) Wan, P.; Yates, K.; Boyd, M. K. J. Org. Chem. 1985, 50, 2881.

(32) Woolley, E. M.; Hepler, L. G.; S. Roche, R. Can. J. Chem. 1971, 49, 3054.

(33) Kochi, J. K. J. Am. Chem. Soc. 1962, 84, 1193.

(34) Hu, J.; Eriksson, L.; Bergman, Å.; Kolehmainen, E.; Knuutinen, J.; Suontamo, R.;

Wei, X. Chemosphere 2005, 59, 1033.

(35) Arulmozhiraja, S.; Selvin, P. C.; Fujii, T. J. Phys. Chem. A 2002, 106, 1765.

(36) Dorofeeva, O. V.; Novikov, V. P.; Moiseeva, N. F.; Yungman, V. S. J. Mol. Struct. thermochem. 2003, 637, 137.

(37) Chateauneuf, J. E.; Nie, K. Adv. Environ. Res. 2000, 4, 307.

(38) Stadler, D.; Bach, T. Chem. Asian J. 2008, 3, 272.

(39) Mayr, H.; Kempf, B.; Ofial, A. R. Acc. Chem. Res. 2003, 36, 66. 63

(40) Duan, L.; Wang, L.; Inge, A. K.; Fischer, A.; Zou, X.; Sun, L. Inorg. Chem. 2013,

52, 7844.

(41) Steiner, T. Angew. Chem. Int. Ed. 2002, 41, 48.

(42) Schaefer, T.; Penner, G. H.; Takeuchi, C.; Tseki, P. Can. J. Chem. 1988, 66, 1647. 64

CHAPTER 3 ELECTROCATALYTIC WATER OXIDATION BY USING

COVALENTLY LINKED XANTHEINIUM DIMER.

Our previous study shows that the flavin based catalysis depends on the type of working electrode and that the oxides formed on the metal surface assist the oxygen evolution.1 The disadvantage of this kind of catalysis is the difficulty in studying the catalytic mechanism. The result from electrode assisted mechanism suggests that a homogeneous catalyst can be developed,

+ where two cation species are covalently linked with suitable linker. The xanthylium dimer (Xan )2 containing two xanthene units covalently linked with diphenyl ether was studied as the model

+ electrocatalyst for water oxidation. In this chapter, the electrocatalytic water oxidation by (Xan )2 is evaluated using various electrochemical techniques. To get insight into the role of the dimeric form in such catalysis, a comparative study with the monomer cation (Xan+) is performed.

+ + Scheme 3.1. Molecular Structures of compounds, (Xan )2 and Xan .

The electrocatalytic water oxidation process was investigated by using electrochemical methods, and some electrochemical techniques are introduced here. 65

3.1 Electrochemical Techniques

Electrochemistry is a branch of chemistry, which studies the interaction between the

electricity with the chemical changes in the solution. It provides the better understanding of redox

properties of certain species by evaluating the thermodynamics and kinetics parameters. Such

thermodynamic and kinetics parameters can be studied by using several voltammetric techniques,

some of which are discussed here.

3.1.1 Cyclic Voltammetry

Cyclic voltammetry is a widely used electroanalytical technique in different areas of chemistry. It is used mostly to study the redox process, understand the reaction intermediates and to know the stability of the reaction’s products. In this technique, generally the potential on the working electrode is varied with time, and change in the current is monitored. The result is a current-potential curve in which potential can be correlated with the thermodynamics of the process, while kinetics can be evaluated by the current. A typical cyclic voltammogram shows two potential peaks, cathodic (ܧ௣௖) and anodic (ܧ௣௔) for two current peaks (݅௣௖,݅௣௔) (Figure 3.1).

Depending on the shape of the CV curve, the electrochemical process can be classified into three

different processes:

(a) Reversible process: If the electron transfer process is fast relative to other processes, such as

diffusion, the process is electrochemically reversible.

(ݔ ൅ ݊݁ ܴ݁݀ (3.1ܱ 66

For such reversible reaction, the peak separation between the anodic and cathodic potential should be ͲǤͲͷͻʹȀ݊ V, i.e. ~60 mV for one electron transfer, where ݊ is the number of electrons transfered.

οܧ݌ ൌ ȁܧ݌ܿȂ ܧ݌ܽȁ ൌ ʹǤ͵Ͳ͵ܴܶȀ݊ܨ (3.2) where, ܴ= gas constant

ܶ = temperature,

ܨ = Faraday constant.

Figure 3.1 A typical voltammogram of the reversible electron transfer showing the parameters cathodic peak potential (ܧ௣௖), cathodic peak current (݅௣௖), anodic peak potential (ܧ௣௔), anodic peak current ൫݅௣௔൯ǡ and peak separation (οܧ௣).

The standard redox potential ሺܧιሻ for a reversible couple is given by,

ι ܧ ൌሺܧ௣௖ ൅ܧ௣௔ሻȀʹ (3.3) 67

For a diffusion controlled reversible reaction, the scan rate (ݒ) is related to peak current by the

Randles–Sevcik expression at 25 °C.2

ଷ ଵ ଵ ହ (୮ ൌ ʹǤ͸ͺ͸ ൈ ͳͲ ଶ଴ଶݒଶ (3.4

2 where, ݅௣ is the peak current in ampere, ܣ is the electrode area (cm ), ܦ is the diffusion coefficient

2 – 1 –3 ଴ is the concentration in molcm , ݊ is the number of electron, and ݒ is the scan rate inܥ ,( cm s)

Vs–1. According to this equation, the peak current increases with the square root of the scan rate and bulk concentration. The rate of such scan rate dependence electron transfer process depends

3 on the mass transport rate. The faradaic peak current (݅௙௖) increases linearly with the square root of scan rate whereas the charging current ሺ݅௖ሻ increases linearly with the scan rate, which results in the increase in the background noise changing the shape of the CV curve. Thus, important information could not be abstracted from such data as the ratio of ݅௙௖Ȁ݅௖decreases. Due to this reason, the high scan rate experiments were usually performed by using microelectrodes to minimize the background noise.4

In case of the adsorption controlled reversible electrode reaction, the peak current is given by

ଶ ଶ (3.5) כൌ Ȟ ୮ Ͷ ଴ where,

ܨ = Faraday’s constant,

ܴ = a gas constant (J/mol∙K), 68

ܶ = temperature (K), (mol/cm2), and

כ Ȟ଴ = the amount of analyte adsorbed on the electrode surface.

(b) Irreversible process: In case of the irreversible process, the rate of the electron transfer is slower than the mass transport. As a result, the cyclic voltammogram shows no return peak (Figure 3.2), and the oxidation potential (equation 3.1) occurs at more positive potential compared to standard potential (ܧιሻ. This situation arises when the electron transfer process is complicated by a followed up chemical reaction. Also, the separation between the forward and reverse peak is large and the return peak is absent孱 For one electron process (equation 3.1), the peak current for the irreversible

process is given by;

భ భ భ ହ (మݒమ (3.6ܦ଴ܥܣ௣ ൌ ʹǤ͸ͺ͸ ൈ ͳͲ ݊ሺߙ݊ఈሻమ݅

where,

ߙൌ transfer coefficient,

݊ߙ = number of electron transfer,

ܣ = electrode area (cm2),

ܦ = diffusion coefficient (cm2 s–1),

–3 ܥ଴ = the concentration in mol cm . 69

Figure 3.2 A typical cyclic voltammogram of the irreversible electron transfer, where the return

peak is absent.

In this irreversible system, the forward peak shifted to more anodic potential and the value of that shift at 25 °C is about ͵Ͳൗ for a ten-fold increase of the scan rate.3 Also, due to the ߙ݊ఈ

sluggish electron transfer, the peak height is lower compared to the corresponding reversible

system.3 The peak potential of irreversible process is given by:2

ଵ ݒ ଶܨ ι ߙ݊݇ ܴܶ ܧ ൌܧι െ ቎ͲǤ͹ͺ െ ݈݊ ൅݈݊൬ ఈ ൰ ቏ ௣ ߙ݊ܨ ଵ ܴܶ ܦଶ

(3.7)

where,

ܧι = the standard potential, 70

݇ι = the standard rate constant, and the other parameters have the same meaning as defined above in equation (3.6).

(c) Quasi-reversible process: The quasi-reversible process is the intermediate of the reversible and irreversible processes. In this process, the current is controlled by both the mass transport and charge transfer. The electron transfer rate constant ሺ݇ιሻ range of such quasi-reversible process is

reported to beͳͲିଵ ൐݇ι ൐ͳͲିହcm/s.2 The shape of a cyclic voltammogram is a function of

ݒܨ݊ ι݇ ൗߨܽܦ , where ܽൌ ൗǤܴܶ The process becomes reversible with the decrease in scan rate, i.e.

݇ι ݇ι at higher value of ൗߨܽܦ and the process becomes irreversible at the low value of ൗߨܽܦ. Also,

the separation of peak potential is larger compared to reversible process.

Figure 3.3 A typical cyclic voltammogram for the quasi-reversible electron transfer, where the

peak separation deviates from 59 mV/n.

3.1.2 Study of Reaction Mechanism

In the electrochemical process, generally chemical reactions (CR) precede or succeed the electrons transfer process (ET). The occurrence of such chemical reactions directly affects the available concentration of electroactive species around electrode, which results in the change in the shape of cyclic voltammograms. This change in shape of cyclic voltammograms can be useful to elucidate the reaction pathways and provide the information about the reactive intermediates.

For example, consider a reaction (equation 3.8) in which the ET is followed by CR, generally referred as EC mechanism.

ܣ൅݊݁ ܤ՜ܥ (3.8)

The cyclic voltammogram of such reaction will have small reverse peak, which is due to the reduced concertation of product (B) by chemical reaction. If the CR is fast, then all of the product will get converted to C and there will be no return peak.

The other scenario is a homogeneous CR precede a reversible electron transfer process, which is refer as CE mechanism as shown in equation 3.9;

ܣ ൅ ݊݁ (3.9) where, C represents the non-electroactive species whereas A and B represents the electroactive species. In this mechanism, the electroactive species (A) are formed from the chemical reaction, thus it is important to know the formation of A species in cyclic voltammetry timescale. The timescale of cyclic voltammetry is given by the parameters:

ݒܨ݊ ൗܴܶ for a reversible process

ݒܨߙ݊ఈ ൗܴܶ for a quasi-reversible and irreversible processes. 72

This relationship shows that the timescale of cyclic voltammogram depends on the scan rate.3 The limit of chemical complication depends either on equilibrium constant (K) or the kinetics

of forward and reverse reactions (ܭ௙and ܭ௥ሻǤ The three different possibilities are described as,

ݒܨ݊ ൗሻǤܴܶ When K is large, ET is not affected ا ௥ܭ ௙ ൅ܭ) ,i) In case of slow chemical reaction) by CR and reversible cyclic voltammogram response appears. When K is small, the CV response appears reversible, however the amount of current response is small compared to the initial concentration of electro-inactive species (C) present in solution.

ݒܨ݊ ൗሻǤܴܶ When K is large, the CV response is ب ௥ܭ ௙ ൅ܭ) ,ii) In case of fast chemical reaction)

reversible but the measured standard potential ܧ଴shifted to more negative potential value. When

K is small, the CV shows a sigmoidal shape due to the small equilibrium concentration of [A].

Also, there is no change in the height of the peak current with scan rate. Such limiting peak current

ሺ݅௅ሻ is given by

భ భ మ మ ݅௅ ൌ݊ܨܣܦ஺ܥ஼ܭ൫݇௙ ൅݇௥൯ (3.10)

where, ܥ஼ is the concentration of electro-inactive species in solution, ܦ is the diffusion coefficient

of species A. For the known value of K, the forward and backward rate constants can be calculated

by using equation 3.10.

ݒܨ݊ ௥ ؆ ൗሻܴܶ , the electrode process is notܭ ௙ ൅ܭ) ,iii) In case of intermediate chemical reaction)

affected by chemical reaction. However, forward current decrease much more with scan rate than

the reverse current.

The other possible scenario is the catalytically regeneration of oxidized species. Consider a reaction, 73

ܣ൅݊݁ ܤ

ܤ൅ܼ ܣ൅ܥ (3.11)

Here, Z is a powerful electro-inactive oxidant and its concentration should be higher than the concentration of electro active species, A. If kf is small, the following chemical reaction is not

possible, and the reversible CV response appears. If kf is large or scan rate is reduced, then there

is regeneration of the reagent species (A), which results in the higher current for the forward

reaction compared to simple reversible reaction. The increase in peak current with decrease in scan

rate reaches the limiting value due to the equal speeds of electrode process to consume the species

A. In such limiting cases, the forward peak changes to S-shaped and the limiting current ሺ݅௅ሻ is

given as;

భ భ మ మ ݅௅ ൌ݊ܨܣܦ஺ܥ஺݇௙ (3.12)

Alternatively, the peak current can be calculated by using the known value of kinetic current ሺ݅௞ሻ

and diffusion current (݅ௗ) as;

భ ௞ ோ் ݅ Ȁ݅ ൌ ʹǤʹͶʹ ቀ ೑ ቁమ (3.13) ௞ ௗ ௡ி௩

This limiting condition is used to calculate the catalytic rate constants for water oxidation

catalysts.5,6

3.1.3 Bulk Electrolysis

Bulk electrolysis is also referred as control potential electrolysis. This technique employs a three-electrode system controlled by a potentiostat. The potential of working electrode is held at 74

constant value and the current is measured at various time intervals. According to Faraday’s law,

the total current passed during the electrolysis is calculated by integrating the current;7

௧ (ݐ (3.14݀ܫ ൌ׬଴ܳ

(is the current at timeሺݐሻ. The chargeሺܳ ܫ where, ܳ is the charge passed during the electrolysis and

is related to number of electron transferredሺ݊), Faraday’s constantሺܨሻ, and number of moles of

analyte (N) as;

ܳ ൌ ݊ܨܰ (3.15)

This method can be used to calculate the number of electrons transferred during the electrolysis.

3.2 Experimental Section

3.2.1 General Methods

All reagents and starting materials were purchased from commercial sources and used as

received. 1H and 13C NMR spectra at room temperature were recorded using a spectrometer

operating at a field of 11.7 T (500 MHz for 1H and 125 MHz for 13C), Bruker-Spectrospin 500

Ultra Shield UV-Vis absorption spectra were taken on a HP8453 UV-Vis spectrophotometer.

Synthesis of compounds used in this study were already mentioned in Chapter 2.

3.2.2 Cyclic Voltammetry

季Cyclic voltammetry (CV) was performed using a BASI epsilon potentiostat in a VC-2 voltammetry cell (Bioanalytical Systems-BASI) using a platinum working electrode (1.6 mm 75

diameter, MF-2013, BASI), fluorine-doped tin oxide FTO (area 4.5 cm2, Hartford Glass), boron

doped diamond BDD (area 2 cm2, Fraunhofer USA) working electrodes, platinum wire auxiliary

electrode (MW-4130, BASI), and non-aqueous Ag/Ag+ reference electrode (MF-2062, BASI).

Acetonitrile (anhydrous, 99.8%) and Tetrabutylammonium perchlorate (TBAP) was purchased

from Sigma-Aldrich and used without further purification.

3.2.3 Controlled Potential Electrolysis

Controlled Potential Electrolysis (CPE) was performed in a custom-built two-

compartment gas-tight electrochemical cell under argon atmosphere. One arm of the cell consists

of Pt mess working electrode (Sigma Aldrich), Ag/AgCl aqueous reference electrode (BASI),

oxygen sensor (HIOXY-R, Ocean Optics), and argon gas inlet. The second arm is consisted of a

Pt wire auxiliary electrode and a gas outlet. Electrolysis was carried out using an EC Epsilon

potentiostat (BASI) at +2.1 V vs. Ag/AgCl in a 0.1 M phosphate buffer (3:1 ACN: water) at 0.5

pH value. Ultra-high pure water was used for the electrolysis, where the water was purified using

a Barnstead Nano pure purification system. Prior to each experiment, the sensor was calibrated

using a two-point reading (20.9 % O2 in the air and 0 % O2 in the argon-purged cell). The oxygen

sensing experiment was performed for both the baseline solution (0.1 M phosphate buffer pH 0.5)

+ + and solution with compounds (1 mM solution of (Xan )2 and Xan in 0.1 M phosphate buffer pH

0.5). The oxygen sensor probe was placed in the headspace and data were collected at 10-second

intervals. Before electrolysis was initiated the O2 signal was monitored for 10 minutes to ensure

there was no leakage of O2 from the outside. After ensuring there were no leaks, the electrolysis was started and continued for 30 min. Conversion of O2 percentage into μmol was obtained from 76

the known volumes of the solution (VS = 130 mL) and the headspace (VH = 55mL) using Henry’s

Law.

+ 3.2.4 Chemical Oxidation of (Xan )2

To investigate the nature of catalysis (homogeneous or heterogeneous) chemical oxidation

+ of (Xan )2 was made by using sodium persulfate as oxidant in 3:1 ACN: water (pH 0.5) solution

at 60 °C. The oxygen measurement was performed by analyzing the headspace of the reaction

mixture vessel by using gas chromatogram (Shimadzu GC-8A) operated with ultra-high purity

argon carrier gas and 5 Å molecular sieve column. The detector was first calibrated with the known

+ amount of oxygen gas. 22 mL solution containing 2.5 mM (Xan )2 and 2.7 M persulfate in 3:1

ACN: water (pH 0.5) was placed in a vial and heated to 60 °C. The oxygen content of the headspace

(1 mL) was then measured by injecting 100 μL sample using a Hamilton airtight syringe.

Contamination of headspace with oxygen is quantified by measuring nitrogen in the headspace.

3.3 Result and Discussion

3.3.1 UV-Vis Spectroscopy: Pseudobase Xanthylium ion Equilibrium

The proposed catalytic cycle (Scheme 1.9b), the first step is the reaction of xanthylium cations with water to form pseudobase intermediates. This process can be monitored by using UV-

+ + Vis spectroscopy. In our study, both Xan and (Xan )2 ions generate psedobases with water (Figure

3.4). As already discussed in chapter 2, the pka value of hydroxylation of Xan+ was calculated as

+ 1.3, while (Xan )2 showed two pKa values, the first and second hydroxylation occurred at pKa 0.2

and pka 2.6 respectively. 77

3 4

3 2

2 Absorbance Absorbance + Xan 1 (Xan-OH)2 + (Xan )2 1 Xan-OH

0 0 200 300 400 500 600 700 200 300 400 500 600 700 Wavelength (nm) Wavelenght (nm)

Figure 3.4 pH dependent absorption spectra of xanthylium ions: (a) Xan+ in aqueous solution at

different pH values [pH 5 (black), pH 1.5 (red), pH 1.1 (blue), pH 0.75 (magenta), and pH 0.4

+ (cyan)]. (b) (xan )2 in aqueous solution at different pH values [pH 7 (navy), pH 4.2 (dark yellow), pH 2.6 (pink), pH 0.95 (cyan), pH 0.44 (blue), pH 0.1 9 (red), and pH –0.5 (black)].

3.3.2 Cyclic Voltammetry

+ + Cyclic voltammograms (CV) were obtained for both xanthylium ions Xan and (Xan )2. In

+ case of (Xan )2, the two xanthylium ions were connected by diphenyl ether (DPE) as a linker. The

CV of DPE was also obtained to see the electrochemical properties of linker. 78

-20 O

P$ -40

-60 O O Current ( Current -80 O

-100 O

-120 2.5 2.0 1.5 1.0 0.5 0.0 -0.5 -1.0 Potential Vs Ag/AgNO3

+ + Figure 3.5 Cyclic voltammograms of 1.5 mM Xan (black), (Xan )2 (red) and DPE (blue) in

acetonitrile with 0.1 M TBAP as electrolyte, scan rate 100 mv/s. Electrodes: platinum working

electrode, Ag/AgNO3 reference electrode, and platinum wire counter electrode. The grey line represents the background current. Scan direction + 0.5 V → +2.5 V → –1 V → +0.5 V for Xan+

+ (black), (Xan )2 (red) and the + 0.5 V → +2.5 V → +0.5 V to DPE.

+ + The cathodic scan of cyclic voltammograms of Xan and (Xan )2 show a single peak at –

0.14 V and –0.01 V respectively. In case of Xan+, the one electron reduction peak is reversible

suggesting the stability of electrochemically formed radical (Xan.). The quasi-reversible peak is

+ obtained for (Xan )2, which suggests that the electrochemically formed radical is unstable and

+ + undergoes some decomposition reaction. In anodic regions, both Xan and (Xan )2 show the similar

catalytic current at above 2.1 V. The cyclic voltammogram of DPE shows the oxidation peak at

1.68 V. The previous study on DPE shows that DPE undergoes the electrochemical polymerization

8 + during oxidation, which was also observed in our case. The comparison of CV of DPE and (Xan )2 79

+ show that the oxidation of DPE moiety in (Xan )2 occurs at more positive potential (peak at 2.2 V

in figure) compared to DPE (peak at 1.68 V). This shift in oxidation potential might be due to the

+ presence of electron withdrawing xanthylium moieties in (Xan )2. However, the first electron

+ oxidation occurs at the DPE moiety in (Xan )2, which was also confirmed by the theoretical

calculation results as shown in Figure 3.6.

+ Figure 3.6 (a) Charge density of (Xan )2, (b) Charge density (blue) and spin density (red) of one

+ electron oxidized (Xan )2.

3.3.3 Electrode Type Dependence

From our previous study on Et-Fl+, the water oxidation catalysis depend on the type of

working electrode.1 The catalysis was only observed in case of active electrodes (Pt, GC), but not

in electrochemically inactive fluorine doped tin oxide (FTO) electrode. Thus, to investigate the

electrode dependence nature, the experiments were performed in Pt, Au, FTO and BDDE. 80

(A)

(a) Pt (b) Au 0 )

2 -35

mm I /I =3.52 -70 cat d P$ Icat/Id=5.87 -105 30

-30 Current density ( density Current Icat/Id=5.17 -60

-90 2.4 1.8 1.2 0.6 0.0 -0.6 2.4 1.8 1.2 0.6 0.0 -0.6 Potential Vs Ag/AgNO 3

(B)

30 (a) Pt (b) Au 0

2 -30 mm

I /I =3.63 Icat/Id=4.25

P$ cat d -60

-90 30 (c) BDD (d) FTO 0 Current density( -30 Icat/Id=4.58

-60

-90 2.4 1.8 1.2 0.6 0.0 -0.6 2.4 1.8 1.2 0.6 0.0 -0.6 Potential vs Ag/AgNO 3 81

+ + Figure 3.7 Cyclic voltammograms of 1.5 mM Xan (Figure A) and (Xan )2 (Figure B) in acetonitrile containing 0.1 M TBAP as the electrolyte; on (a) Pt electrode, (b) Au electrode, (c)

BDD electrode, and (d) BDD electrode in both cases. (Scan rate: 100 mV/s. Scan direction: blue curve: +0.5 → –1 V → +0.5 V; red curve: +0.5 → +2.5 → –1 V → +0.5 V; black curve: baseline

+ scan for the full range. The Icat/Id is the ratio of currents at +2.5 V and –0.01 V for (Xan )2, +2.5 V and –0.14 V for Xan+.

These electrodes mentioned generated hydroxylated surface species (S–OH; s- electrode’s surface) species. These S-OH species on Pt and Au electrodes are considered as active groups, as they are prone to undergo further oxidation to generate S–O species.9,10 In contrast, BDD and FTO electrodes are considered to be inactive, as they don’t show an additional oxidation step to form

S-O groups.

The oxidation and reduction peaks depend on the electrode material. The similar oxidation behaviors were observed in case of Pt, Au, and BDD electrodes, the current increases above 2.1 V

+ + in (Xan )2 and above 2.3 V in Xan . However, no oxidation was observed in FTO electrode, which was consistent with the result obtained for Et-Fl+ on FTO electrode.1 These results from FTO electrode suggests that the compounds interact with the electrode surface prior to charge transfer.

The electrode dependence oxidation process was further investigated by using icat/id ratio (Figure

3.7). The highest icat/id values (where Icat is the current during catalysis, Id is the peak current for reversible electron transfer in the absence of a followed up chemical reaction) were obtained for

Pt, Au and BDD electrodes whereas it was impossible to get the values FTO electrode. These results suggest that efficient interaction occur with Pt, Au and BDD electrodes and the least efficient interaction is with FTO electrode. 82

The strong interaction was observed in case of BDDE, which was considered as electrochemically inactive electrode. This contradictory result forced us to think whether the oxidation process is homogeneous. To figure out this dilemma, we performed chemical oxidation which is discussed in section 3.3.5.

+ 3.3.4 Bulk Electrolysis of (Xan )2

+ The catalytic peak observed at above 2.1 V in (Xan )2 was further verified by using bulk

electrolysis experiment (Figure 3.8 b).季The oxygen evolution was monitored during the controlled

+ + potential electrolysis at 2.1 V (vs Ag/Ag ) in aqueous solutions of (Xan )2 on Pt mess working

+ electrode. The amount of O2 increased with time in the absence of (Xan )2, which is due to the

+ water oxidation by the electrode itself (black curves in the figure 3.8 b). In the presence of (Xan )2,

the electrolysis of solution produced a significantly higher amount of O2. Thus, this oxygen

+ evolution confirmed that the peak above 2.1 V was due to the catalytic water oxidation by (Xan )2.

The inset figure (Figure 3.8 a) shows the amount of current increase at 2.5 V with the increase in

water concentration. As the catalytic peak is observed only above 2.3 V in case of Xan+, the similar

+ bulk electrolysis at 2.1 V in the presence of Xan did not show any evolution of O2 as expected.

However, the bulk electrolysis experiments performed at 2.5 V in the presence of compounds did

not result in the significant evolution of oxygen. This is possibly due to the presence of other

competing oxidation processes at higher potentials. This kind of behavior was also observed in

+ 1 case of Et-Fl . 83

+ Figure 3.8 Cyclic voltammograms of 1.5 mM (Xan )2 in acetonitrile using 0.1 M TBAP as electrolyte. Scan direction: red curve +0.3 V→ +2.5 V→ –1 V→ +0.3 V; blue curve, +0.3 V→ –

1 V→ +0.3 V. The grey line represents the background current. Electrodes: platinum working electrode, platinum wire counter electrode and Ag/AgNO3 reference electrode, scan rate 0.1 V/s; inset figure shows the percentage increase in current at 2.5 V with the concentration of pH 0.5 water (Iw represents current in the presence of water and Inw represents current in the absence of

+ water); (B) Oxygen evolution during the electrolysis of 1mM (Xan )2 in 0.1 M phosphate buffer pH= 0.5 (red) (30% faradaic efficiency), the black line represent the amount of oxygen evolved in

+ + the absence of (Xan )2 (86 % faradic efficiency); applied potential: +2.1 V vs Ag/Ag .

+ 3.3.5 Chemical Oxidation of (Xan )2

+ To investigate further the nature of water oxidation catalysis by (Xan )2 (a homogeneous

+ or a heterogeneous process), we performed the chemical oxidation of (Xan )2 by using sodium 84 persulfate (Na2S2O8) as an oxidizing agent. Persulfate is usually activated by using UV light or heat to produce sulfate radicals (SO4–•),11 which acts as a strong oxidizing agent with potential

~2.5 V vs NHE. First, the UV-Vis experiments were performed to confirm that the persulfate can

+ + acts as an oxidizing agent for (Xan )2. The result shows that the (Xan )2 get decomposed with time

+ in the presence of persulfate at 60 °C (Figure 3.9 a). It is also important to mention that (Xan )2 is stable at 60 °C in the absence of persulfate (Figure 3.9 b).

(a) (b)

2 2

30 min 1 hr 30 min + 1.5 hr (Xan ) 5 hr 2 2 hr 3 hr Absorbance Absorbance 4 hr

0 0 200 300 400 500 600 700 800 200 300 400 500 600 700 800 Wavelength (nm) Wavelength (nm)

+ Figure 3.9 (a) UV-Vis spectra showing the decomposition of (Xan )2 with 1000 equivalent of

+ sodium persulfate. (b) UV-Vis spectra of (Xan )2 in 3:1(ACN: H2O) pH 0.5 solution at 60 °C.

The oxygen measurement was made by using gas chromatography. The result shows that

+ the O2 evolved even in the absence of (Xan )2, which is due to the persulfate decomposition in

12 + acidic solution. However, the amount of O2 evolved was not significant in the presence of (Xan )2

(Figure 3.10). This result from chemical oxidation supported the fact that the working electrode plays a key role in water oxidation catalysis. 85

ol) 15 P0

No. of moles ( 5

10 20 30 40 50 60 Time (min)

Figure 3.10 Amount of oxygen detected by GC in the presence of 2.7 M persulfate only (black)

+ and 2.5 mM (Xan )2 in the presence of 2.7 M persulfate (red). The experiment was performed in pH 0.5 solution of 3:1 (water: acetonitrile) mixture at 60 °C.

3.4 Conclusion

+ + The electrochemical behavior of xanthium ions (Xan and (Xan )2) were studied by using electrochemical techniques in order to investigate their ability to oxidize water. The cyclic

+ + voltammograms of both Xan and (Xan )2 in acetonitrile are similar and the electrochemical

+ process depend on the working electrode. The huge catalytic peak in case of (Xan )2 was assigned to the electrochemical process that lead to water oxidation, which was confirmed by controlled

+ potential electrolysis. However, the result obtained from chemical oxidation of (Xan )2 in aqueous solution in the presence of sodium persulfate further verify the involvement of electrode surface on water oxidation catalysis. Moreover, the catalytic peak in Xan+ was observed at above + 0.2 V 86

+ positive potential compared to (Xan )2, which did not show the O2 evolution during bulk electrolysis. This might be because of the competitive side reactions that occur at higher potentials.

3.5 References

(1) Mirzakulova, E.; Khatmullin, R.; Walpita, J.; Corrigan, T.; Vargas-Barbosa, N. M.;

Vyas, S.; Oottikkal, S.; Manzer, S. F.; Hadad, C. M.; Glusac, K. D. Nat. Chem. 2012, 4, 794.

(2) Wang, J. Analytical electrochemistry; John Wiley & Sons, 2006.

(3) Zanello, P. Inorganic electrochemistry: theory, practice and application; Royal

Society of Chemistry, 2007.

(4) Wipf, D. O.; Kristensen, E. W.; Deakin, M. R.; Wightman, R. M. Anal. Chem. 1988,

60, 306.

(5) Chen, Z.; Concepcion, J. J.; Luo, H.; Hull, J. F.; Paul, A.; Meyer, T. J. J. Am. Chem.

Soc. 2010, 132, 17670.

(6) Wasylenko, D. J.; Palmer, R. D.; Schott, E.; Berlinguette, C. P. Chem. Commun.

2012, 48, 2107.

(7) Bard, A. J.; Faulkner, L. R.; Leddy, J.; Zoski, C. G. Electrochemical methods: fundamentals and applications; wiley New York, 1980; Vol. 2.

(8) Xu, J.; Wan, L.; Li, Y.; Liu, H.; Pu, S.; Shen, L. J. Polym. Sci., Part A: Polym.

Chem. 2007, 45, 5932. 87

(9) Rossmeisl, J.; Logadottir, A.; Nørskov, J. K. Chem. Phys. 2005, 319, 178.

(10) Panizza, M.; Cerisola, G. Chem. Rev. 2009, 109, 6541.

(11) Huang, K.-C.; Couttenye, R. A.; Hoag, G. E. Chemosphere 2002, 49, 413.

(12) Kolthoff, I.; Miller, I. J. Am. Chem. Soc. 1951, 73, 3055. 88

CHAPTER 4 THERMODYNAMIC HYDRICITIES OF BIOMIMETIC ORGANIC

HYDRIDE DONORS

In this chapter, the hydride donating abilities (Hydricity) of several organic compounds were studied. The hydricity of thirteen different organic hydrides with different structural properties; NADH analogues (BNAH, CN-BNAH, Me-MNAH and HEH), methylene tetrahydromethanopterin analogs (BIMH and CAFH), acridine derivatives (Ph-Acr, Me2N-AcrH,

T-AcrH, 4OH, 2OH, 3NH), and a triarylmethane derivative (6OH) were studied by using experimental and theoretical methods. The study was conducted in two different solvents, acetonitrile and DMSO. The calculated hydricity values, obtained using density functional theory, showed a reasonably good match with the experimental values, obtained using “potential-pKa” and “hydride-transfer” methods. The solvent effect and the structural effects were observed in the hydride donating abilities of these compounds.

H H H H CN H H H CONH2 H H EtOOC COOEt N N N CH Ph N N CH2Ph 2 BNAH CN-BNAH Me-MNAH HEH Ph-AcrH

N N OO H N H H H OO O O N N N N

Me2N-AcrH T-AcrH 4OH 2OH

R N O O O O H H O N N N H N N R R O H O N H O N N R=C8H17 3NH 6OH BIMH CAFH 89

Scheme 4.1 Structures of compounds; NADH analogues (BNAH, CN-BNAH, Me-MNAH and

HEH), methylene tetrahydromethanopterin analogs (BIMH and CAFH), acridine derivatives (Ph-

Acr, Me2N-AcrH, T-AcrH, 4OH, 2OH, 3NH), and a triarylmethane derivative (6OH).

4.1 Hydricity

Hydricity is defines as the Gibb’s free energy required to release a hydride ion ( ିሻ from

a molecule (Equation a). Hydricity is also referred as a hydride donating ability of a molecule.1 It

is denoted by οܩுష and clearly indicates that the compounds with lower οܩுషvalues are better

hydride donors whereas the compounds with higher οܩுష values are better hydride acceptor.

ା ି െ ՜ ൅ οܩுష (a)

Hydricity strongly depends on the solvents.2 This may be due the difference in solvation of species, most importantly the difference in solvation of free hydride ion. Mostly, hydricity values of compounds are studied in acetonitrile. Hydricity values of different compounds are comparable when all the compounds are studied in same solvent.

Experimentally, hydricity values can be determined by using three widely used approaches;

(i) Potential-pKa method, (ii) Hydride transfer method, and (iii) H2 heterolysis method. Hydricity

is a state function and it can be calculated by using any pathways of proton and electron transfer

as shown in Scheme (4.2). Acidity values, BDFE values and reduction potentials values can be

converted to free energy and combined to give the hydricity values. 90

൅ ൅ ൅

Scheme 4.2 Relationship between the thermodynamic hydricityοܩுష, bond dissociation free

energy (BDFE), acidity (pKa) and reduction potentials.

(i) Potential pKa method: In this method, the two electron reduction potentials of conjugate

hydride acceptor (ܴା) and acidity of compounds (ܴെܪ) (Scheme 4.2, equation c) are used to

determine the hydricity (Scheme 4.2).3 The experimentally obtained values of free energies and

the free energy of two electron reduction of ܪା to ܪି(Scheme 4.2, equation f) are combined to

give the hydricity of a compound. The reduction potentials for equations d and e are usually

determined by using cyclic voltammetry technique. The ܧభvalue obtained from cyclic మ

voltammetry is the good approximation ofܧι, if the electrochemical wave is reversible.4 The

obtained value of ܧι is converted to Gibb’s free energy by using the relation as shown in equation

(b), where ݊ is the number of electrons, ܨ is Faraday’s constant, and ܧι is standard reduction

potential.

οܩ ൌ െ݊ܨܧι ൌ െʹ͵ǤͲ͸ܧι (b)

In case of irreversible electrochemical waves, other electrochemical technique such as redox

potentiometric can be used.5

ି ା ι ܴെܪ ܴ ൅ܪ οܩଵ ൌ ͳǤ͵͸Ͷ݌ܭ௔ (c)

ି Ȉ Ȃ ι ι ܴ ܴ ൅݁ οܩଶ ൌ ʹ͵ǤͲ͸ܧೃȈ (d) ೃȂ 91

Ȉ ା Ȃ ι ι ܴ ܴ ൅݁ οܩଷ ൌ ʹ͵ǤͲ͸ܧೃశ (e) ೃȈ

ା Ȃ Ȃ ι ܪ ൅ʹ݁ ܪ οܩସ (f)

Ȃ ି ܴെܪ൅ʹ݁ ܪ οܩுష (g)

Scheme 4.3 Hydricity determination by using potential-pKa method.

(ii) Hydride transfer method: In hydride transfer method, the equilibrium constant of the reaction between the hydride donor of unknown hydricity (ܴെܪ) with the hydride acceptor (ܣ) of known hydricity or Vice versa is used (Scheme 4.4). The equilibrium constant (ܭ௘௤ሻ is converted to Gibb’s free energy by using the relation as shown in equation (h), where ܴ is the universal gas constant and ܶ is the temperature in Kelvin.

οܩ ൌ െܴ݈ܶ݊൫ܭ௘௤൯ ൌ െͳǤ͵͸Ͷ݈݋݃൫ܭ௘௤൯ (h)

NMR spectroscopy and UV-Vis spectroscopy are the most commonly used techniques to monitor the reaction equilibrium. The equilibrium of the reaction should be confirmed by measuring the equilibrium constants of both forward and backward reactions. To use this method, the difference in hydricity values between the hydride donor and hydride acceptor molecules should be within 3 kcal/mol.

ା ି ι ܴെܪ൅ܣ ܴ ൅ܪܣ οܩହ ൌ െͳǤ͵͸Ͷሺܭ௘௤ሻ (i)

ି ି ι ܪܣ ൅ܪ οܩுି݋݂ܪܣ (j)

ା ି ι ܴ െ ܪ ܴ ൅ܪ οܩுି݋݂ܴ െ ܪ (k) 92

Scheme 4.4 Hydricity determination by using hydride transfer method.

(iii) Hydrogen heterolysis method: In ܪଶ heterolysis method, the equilibrium constant is

ା determined for the reaction of hydride acceptor (ܴ ሻ, a base and ܪଶ (Scheme 4.5, equation l) to

form a hydrideሺܴ െ ܪሻ of unknown hydricity. This method determines the hydricity of unknown

6 hydride based on the hydricity of ܪଶ (Scheme 4.5).

ା ା ι ܴെܪ൅ܪെܤܽݏ݁ ܴ ൅ܪଶ οܩ଺ ൌ െͳǤ͵͸Ͷሺܭ௘௤ሻ (l)

ା ା ι ܪ ൅ܤܽݏ݁ ܪെܤܽݏ݁ οܩ଻ ൌ െͳǤ͵͸Ͷ݌ܭ௔ (m)

ା ି ι (n) ݏ݅ݏ݋݈ݕݎଶ݄݁ݐ݁ܪ଼ܩο ܪ൅ ܪ ଶܪ

ା ି ι ܴെܪ ܴ ൅ܪ οܩுି݋݂ܴ െ ܪ (o)

Scheme 4.5 Hydricity determination by using hydrogen heterolysis method

The equilibrium constant ൫ܭ௘௤൯ is determined usually by using NMR spectroscopy method.

The ݌ܭ௔ values of organic acids (equation m) used in these kinds of studies are mostly available for commonly used solvents such as acetonitrile, DMSO, etc.7-9

In nature, the ground state hydrides are common such as reduced nicotine adenine dinucleotide (NADH)10 which reduce carbonyl compounds,11 carbon dioxide,12 imines13 etc. by

hydride transfer, and the process is driven by the aromatic stabilization of the oxidized product.

Inspired by these natural process, hydride donating ability of several metal hydrides have been

studied for potential application of these hydrides on fuel forming reductions reactions.14-16

Experimental as well as computational approaches were made to study the hydride transfer abilities

of several metal hydrides14,17 and it was found that the hydricity strongly depends on solvents.2 93

The compounds with hydricity values below 76.6 kcal/mol and 44 kcal/mol are capable of

reducing protons and CO2, respectively, in acetonitrile and compounds with hydricity values below

6 60.7 kcal/mol and 42 kcal/mol are required respectively for proton and CO2 reduction in DMSO.

The hydricity of several metal hydrides were mostly studied in acetonitrile and the range of

hydricity varies from 25-120 kcal/mol. The hydricity of metal hydrides depends on the type of

metal,6 ligand bite angles,18 and the substituents group in the ligands.3,6

Although hydricity of metal hydrides have been mostly studied, limited investigations have

been done in case of organic hydrides using NADH analogs,19 hydroqinones,20 and

triarylmethanes.21Among these studied organic hydrides, NADH analogs are better hydride donors with the hydricity values of (59-70) kcal/mol 19 followed by hydroquinones with hydricity values

(58-125) kcal/mol22,23 and triarylmethanes with hydicity values (75-120) kcal/mol.21,24 From these

studied organic compounds, it was found that the electron donating group increases the hydricity

of compounds by stabilizing the (R+) cation formed after the hydride transfer.

Here, we studied the hydride donating abilities of thirteen organic hydrides; NADH analogs

+ (BNAH, CN-BNAH, HEH and Me-MNAH), methylene tetrahydromethanopterin H4MPT

analogs (BIMH and CAFH), acridine derivatives (Ph-AcrH, Me2N-AcrH, T-AcrH, 4OH, 2OH,

3NH) and a triarlymethane derivative (6OH). The study was done in two different solvents,

acetonitrile and DMSO by using experimental and theoretical methods. Experimentally, Potential-

pka method (Scheme 4.3) and Hydride transfer method (Scheme 4.4) were used whereas DFT

calculations were used for theoretical studies.

4.2 Experimental Methods

4.2.1 General Methods

All chemicals were purchased from commercial suppliers and used without further

purification unless notified. NMR spectra were recorded on a Bruker Avance III 500 MHz system.

Steady-state UV/Vis absorption spectra were recorded on a Varian Cary 50 UV-Vis

spectrophotometer. The compounds 1-Benzyl-1,4-dihydronicotinamide (BNAH) and 10-Methyl-

9-phenylacridinium perchlorate (Ph-Acr+) were purchased from TCI America. Fluorene (FlH),

triphenylmethane (TPM), diphenylyldiphenylmethane (DPE) and 1M Super-Hydride were

+ + 25 + 25 + 25 + 26 + 25 + 25 purchased from Sigma. NAD analogs (6O , 4O, 2O , 3N, T-Acr , Me2N-Acr ,

BNA+,27 CN-BNA+,27 Me-MNA+,28 HE+,29 BIM+30 and CAF+31), NADH analogs (6OH,32 2OH,33

Ph-AcrH,34 CN-BNAH,27 BIMH,30 CAFH31), indicator 9-phenylxanthene (XanH35) and nickel-

36 complex ([Ni(dmpe)2](PF6)2 ) were synthesized according to the previously published

procedures.

Synthesis of N,N-dimethyl-4-(10-methyl-9,10-dihydroacridin-9-yl)aniline (Me2N-AcrH):

+ Me2N-Acr (412 mg, 1 mmol) was dissolved in 5 mL ethanol and cooled in an ice bath. Sodium borohydride (150 mg, 4 mmol, 4 eq) was then added and color changed to yellow. The reaction mixture was warmed to room temperature and stirred for additional 4 hours. Resulting solution was filtered and washed with dichloromethane. The filtrate was extracted with dichloromethane, organic extracts were combined and solvent evaporated. The yellow oil was dissolved in ethanol and precipitated by addition of water. The yellow precipitated was filtered, washed with cold water

1 and dried under vacuum to yield 115 g (37%) of pure product. H- NMR (CD3CN, 500 MHz): 95

7.30-7.25 (4H, m), 7.05 (2H, d), 7.00-6.90 (4H, m), 6.60 (2H, d), 5.14 (1H, s), 3.41 (3H, s), 2.81

(6H, s).

Synthesis of N,N-dimethyl-4-((10-methyl-9,10-dihydroacridin-9-yl)ethynyl)aniline (T-

AcrH): T-Acr+ (120 mg, 0.27 mmol) was dissolved in 6 mL ethanol and cooled in an ice bath.

Sodium borohydride (62 mg, 1.62 mmol, 6 eq) was added and left to stir for 2 hours. Deep blue colored disappeared over time. The reaction mixture was then filtered, filtrate disposed and precipitate washed with dichloromethane. Dichloromethane solution was evaporated and yielded

1 in 30 mg of brownish product (33%). H- NMR (CD3CN, 500 MHz): 7.67 (2H, d), 7.38 (2H, d),

7.33 (2H, t), 7.10-7.05 (4H, m), 6.73 (2H, d), 5.00 (1H, s), 3.46 (3H, s), 2.98 (6H, s).

4.2.2 Cyclic Voltammetry

Cyclic voltammetry was performed using a BASi epsilon potentiostat in a VC-2

voltammetry cell (Bioanalytical Systems) using platinum working electrode (1.6 mm diameter,

MF-2013, Bioanalytical Systems), a nonaqueous Ag/Ag+ reference electrode (MF-2062,

Bioanalytical Systems) and a platinum wire (MW-4130, Bio-analytical Systems) as a counter

electrode. The spectroscopic grade solvent DMSO and the electrolyte tetrabutylammonium

perchlorate (TBAP) were purchased from Sigma Aldrich and used as received. Fast scan rate cyclic

voltammetry was performed using CHI 600 C potentiostat and platinum working electrode (CHI-

107, CH instruments, 10 μm diameter). In case of T-Acr+, the second standard reduction potential

was obtained by oxidation of T-Acr-, which was in-situ prepared from T-AcrH and potassium-

dymsil.37 Electrochemical potentials were converted to NHE by adding 0.548 V to the

experimental potentials.38 96

4.2.3 Hydride Transfer Studies

The hydricities of selected model compounds were obtained by determining the equilibrium constant for the hydride transfer to an appropriate hydride acceptor with known hydride affinity. To ensure that the equilibrium constant can be reached, the reference compounds were selected so that their hydricities are within 3 kcal/mol of the hydridicities of our model compounds (as estimated from DFT calculations). The equilibrium concentration ratios of reactants and products were obtained using 1H NMR spectroscopy. The following steps were performed to ensure that the equilibrium was reached: the progress of the reaction was monitored until the integration of NMR peaks stopped to change. Then, an additional amount of one of the products was added and the reaction was monitored again until the equilibrium was reached.

Deuterated acetonitrile and DMSO were used as solvents. All reaction mixtures were prepared in the glove box using dry reagents and air-tight NMR tubes.

Equilibrium of NADH and 2O+: BNAH (3.80 mg, 0.0177 mmol) and 2O+ (8.10 mg, 0.0175 mmol) were dissolved in 0.6 mL of deuterated acetonitrile or DMSO. The equilibrium constant was reached after 14 days in acetonitrile yielding Keq=9.61 whereas the equilibrium was reached after 19 days in DMSO yielding Keq =1.68. The hydricity of 2OH in acetonitrile was obtained from

19 Keq and the reported hydricity of BNAH (59 Kcal/mol) as reference. In case of DMSO, the hydricity of 2OH (58.3 Kcal/mol) was calculated by using potential-pKa method and the obtained value was used as reference to calculate the hydricity of BNAH in DMSO. The 2OH hydricity was

60 kcal/mol in acetonitrile and the hydricity of BNAH was 57.7 kcal/mol in DMSO.

Equilibrium of BNAH and HE+: BNAH (3.80 mg, 0.0177 mmol) and HE+ (5.50 mg, 0.0179 mmol) were dissolved in 0.6 ml deuterated acetonitrile or DMSO. In acetonitrile, the equilibrium 97

was reached after 15 days, yielding Keq= 87.52. In DMSO, the equilibrium was reached after 49

days, yielding Keq= 1.53. The hydricity of HEH in acetonitrile was obtained from Keq and the

reported hydricity of BNAH (59 Kcal/mol) as reference.19 The hydricity of HEH in case of DMSO

was obtained from Keq and hydricity of BNAH (57.7 Kcal/mol) as reference. The HEH hydricity

was 61.5 kcal/mol in acetonitrile and 58.2 kcal/mol in DMSO.

+ + + + Equilibrium of [Ni(dmpe)2H] and 3N or BIM . [Ni(dmpe)2H] was prepared in situ by

addition of 1M Super-Hydride (20 PL, 0.020 mmol) to a solution of [Ni(dmpe)2](PF6)2 (16.2 mg,

0.025 mmol) in 0.6 mL deuterated acetonitrile. To this solution was then added 3N+ (12.7 mg,

+ + 0.025 mmol) or BIM (8.1 mg, 0.026 mmol). The Keq = 3.33 was obtained after 15 days for 3N

+ and Keq= 0.69 was obtained for BIM after 11 days. From these equilibrium constants and reported

+ 33 hydricity of [Ni(dmpe)2H] (49.9 kcal/mol), the hydricity value was calculated as 49.2 kcal/mol

fo 3NM and 50.1 kcl/mol for BIMH in acetonitrile.

4.2.4 pKa Determination

The pKa values of the NADH analogs were determined using the indicator anion method in DMSO.35 Under inert atmosphere, indicators anions ሺܫ݊Ȃሻ were generated by the reaction of

+ - indicator solutionሺܫ݊ܪ) with the lyate ion (potassium dimsyl, K CH3SOCH2 ). The

ܫ݊Ȃconcentrations were determined using recorded absorbance and ܫ݊Ȃ extinction coefficients.

The colored ܫ݊Ȃsolutions formed were then quenched by addition of small aliquots of model hydrides ሺܴെܪሻ solutions in DMSO. The equilibrium constant of the reaction between the model

ܴെܪ and ܫ݊Ȃwas calculated by monitoring the change in the concentration ofܫ݊Ȃ.

ܴെܪ൅ܫ݊ି ܴି ൅݈݊ܪ (p) 98

ሾோషሿሾூ௡ுሿ ܭ ൌ (q) ௘௤ ሾோିுሿሾூ௡షሿ

where,

ܴെܪ = model hydride

ܴି= conjugate base of model hydride

ܫ݊ି = indicator anion

ܫ݊ܪ = indicator

The equilibrium constant is related to the pKa of model hydride and indicator as;

ሾோషሿሾுశሿ ܴെܪ ܴି ൅ܪା ܭ ൌ (r) ௔ ሾோିுሿ

ሾூ௡షሿሾுశሿ ݈݊ܪ ܫ݊ି ൅ܪା ܭ ൌ (s) ௕ ሾூ௡ுሿ

ܭ௔ ܭ௘௤ ൌ ൗ (t) ܭ௕

The pKa values for the model hydrides were determined using the known indicator pKa

value and experimentally obtained equilibrium constants of the reactions between indicator anions

and the hydrides. Indicators were chosen to be within two pKa units from the hydrides and

indicator absorbed in visible spectrum where the other species were transparent.35 In case of the

Ȃ – – overlapping absorptions of ܫ݊ and deprotonated hydride R (Me2N-Acr ), the absorption of Me2N-

- Acr was subtracted using its extinction coefficient at λmax for indicator In. The pKa values of

indicators used in this study are:35 triphenylmethane (TPM, pKa = 30.6) for 4OH, 99 diphenylyldiphenylmethane (DPE, pKa = 29.4) for Me2N-AcrH, 9-phenylxanthene (XanH, pKa =

27.9) for Ph-AcrH and 6OH, fluorene (FlH, pKa = 22.6) for 2OH.

4.2.5 Computational Methods

Geometry optimization: All calculations were performed using Gaussian 09 package39 with the resources of the Ohio Supercomputer Center. The geometries of relevant species (R+, R., R- and R-H) were optimized at the ZB97X-D/6-311G(d) level of theory with the conductor-like polarizable continuum model (CPCM) for solvents (acetonitrile and dimethyl sulfoxide).40-42 The frequency calculation was performed to confirm the absence of imaginary frequencies. The output files from the frequency calculations provided the values thermal corrections to free energies

௦௢௟ + . - (߂ܩ௖௢௥௥) for R , R , R and R-H. The structures optimized at the ZB97X-D/6-311G(d) level were then used to perform a single-point energy calculation at the ZB97X-D/6-

௦௢௟ + . - 311++G(2df,p)/CPCM(ACN or DMSO) level and the electronic energies (ࣟ଴ ) of R , R , R and

R-H were obtained from these output files.

Hydricity calculations. The computational method was adopted from the previously published study.34 The hydricity of a model compound R-H is defined as the thermodynamic driving force ሺοܩுିሻ for the following reaction:

+ – R–H o R + H οܩுି ൌܩோା ൅ܩு௬ௗ െܩோିு

+ The οܩுି values were derived from the calculated energies of R–H, R and the hydride ion: 100

൯ כܩ௦௢௟ ൅߂ܩ௚௔௦ ൅߂ܩ൯ ൅൫ࣟ௚௔௦ ൅߂ כܩ௦௢௟ ൅߂ܩൌ൫ࣟ௦௢௟ ൅߂ ܩο ுି ଴ ௖௢௥௥ ௢՜ ோା ଴ ௖௢௥௥ ௛௬ௗ ௢՜ ு௬ௗ

൯ כܩ௦௢௟ ൅߂ܩെ൫ࣟ௦௢௟ ൅߂ ଴ ௖௢௥௥ ௢՜ ோିு

௦௢௟ ௚௔௦ ௦௢௟ where ࣟ଴ and ࣟ଴ represent electronic energies in solvated and gas-phases, ߂ܩ௖௢௥௥ and

௚௔௦ ௦௢௟ ߂ܩ௖௢௥௥are thermal correction to the Gibbs free energy in solvated and gas-phases, ߂ܩ௛௬ௗ is

כ solvation free energy for the hydride anion and ߂ܩ଴ ՜ is a standard state correction (the value is

௞௖௔௟ ൌ ൅ͳǤͺͻͳ for all species that do not have gaseous standard state).43,44 Electronic כܩ߂ ௢՜ ௠௢௟

energies and thermal corrections to the Gibbs free energy were obtained as described in the

௚௔௦ geometry optimization section. To derive ܩு௬ௗ, the electronic energy (ࣟ଴ ൌ െ͵͵ͳǤͳͶ݈݇ܿܽȀ

௚௔௦ ݉݋݈ሻ and the thermal correction (߂ܩ௖௢௥௥ ൌ െ͸Ǥʹͺ݈݇ܿܽȀ݉݋݈ሻ were obtained for gas-phase using

௦௢௟ – the ZB97X-D/6-31+G(d,p) level of theory. The solvation energy ߂ܩ௛௬ௗ for H was obtained from

the thermochemical cycle connecting gas phase and solution phase one-electron reduction, as

expressed in the following equation:

௦௢௟ ௦௢௟௩ ௚௔௦ ௦௢௟௩ ߂ܩ = Δܩ ಹ - Δܩ ಹ + Δܩ , ௛௬ௗ ቀ ିቁ ቀ ିቁ ሺுሻ ಹ ಹ

௚௔௦ ௦௢௟௩ where, Δܩ ಹ and Δܩ ಹ represent the Gibbs free energy changes for the one electron reduction ቀ ିቁ ቀ ିቁ ಹ ಹ

ு of hydrogen atom in the gas phase and the solution, respectively. ߂ܩ ቀ ቁ is the negative of the ௚௔௦ ுష

ு 45 ௦௢௟௩ electron affinity of hydrogen atom (߂ܩ ቀ ቁ =-17.39 kcal/mol ). Δܩ ಹ was obtained from ௚௔௦ ுష ቀ ିቁ ಹ

଴ ଴ 46 46 the experimental one-electron potentials ܧౄ : using ܧౄ = −0.60 V and −0.55 V for acetonitrile ି ି ౄ ౄ

௦௢௟௩ and dimethyl sulfoxide, Δܩ ಹ values were estimated to be −84.88 kcal/mol and −86.04 kcal/mol ቀ ିቁ ಹ 101

௦௢௟௩ for acetonitrile and dimethyl sulfoxide, respectively. Δܩሺுሻ represents the solvation energy of hydrogen atom and this value was computed using CPCM/6-311++G(2df,p) and found to be −0.1 kcal/mol in both solvents. Using this procedure, the computed values for ܩு௬ௗ were −404.8 kcal/mol and −406.0 kcal/mol for acetonitrile and dimethyl sulfoxide, respectively.

o . o . − Reduction potential calculations: The first (E R+/R ) and second (E R /R ) reduction

. potentials for our model compounds were derived from the calculated driving forces ('GR+/R and

. − 'GR /R ), as follows:

כ ௦௢௟ כ ௦௢௟ ൯ ܧ൯ െ൫ࣟ ൅߂ ܧൌ ൫ࣟ ൅߂ ڄ శ ܩο ோ Ȁோ ଴ ௢՜ ோǤ ଴ ௢՜ ோశ

כ ௦௢௟ כ ௦௢௟ ൯ ܧ൯ െ൫ࣟ ൅߂ ܧష ൌ ൫ࣟ ൅߂ ڄ ܩο ோ Ȁோ ଴ ௢՜ ோష ଴ ௢՜ ோǤ where electronic energies were obtained by performing single point calculations using the

B3LYP47 - D3BJ48 and ZB97X-D341 exchange correlation functionals with the Def2-TZVP49 basis set using the SMD continuum solvation model50 on fully optimized structures obtained using the

BP8651 -D3BJ/Def2-SVP49 model chemistry with ORCA52. The entropic contributions were assumed to be similar for all species, which resulted in their mutual cancellation. The οܩ values

οீ were then used to calculate the standard reduction potentials (ܧൌെ ). The calculated values ௡ி were referenced to NHE by subtracting 3.92 V43 from computed absolute potentials.53

o . − Accuracies of E R /R reduction potentials were systematically improved compared to available experiment when a counter ion (K+) was included in both the ܴǤ and ܴିstates, i.e. using reduction potentials modeled as [ܴǤ-X]+ and [ܴି-X]0. It seemed that adding a counter ion stabilizes the anion relative to the neutral radical, and this resulted in error cancellation. 102

First reduction potential was calculated using ZB97X-D3 calculations, while the second reduction potential is calculated with ZB97X-D3 and the counter ion.

4.3 Result and Discussion:

4.3.1 Calculated Hydricity Values

The hydricities ሺοܩுିሻ of model compounds ܴെܪ (equation a) are calculated from the absolute Gibbs energies of reactant and product states in the appropriate solvation model:

οܩுି ൌ ܩோା ൅ܩ௛௬ௗ െܩோିு (u)

While the Gibbs energies of solvated R+ and R-H species can be calculated reasonably well using the standard DFT methodology and solvation models, the calculation of absolute Gibbs free energy for the solvated hydride ion (ܩ௛௬ௗ) represents a challenge. One way to overcome this drawback is to calculate the thermodynamic parameters for a hydride transfer reaction between R-H and a reference hydride acceptor (such as acridinium cation or p-benzoquinone) whose hydride affinity

54,55 is known from the experiment. Alternatively, the ܩ௛௬ௗvalue can be obtained as a fitting

18,56-58 parameter from the experimental hydricities and calculated Gibbs energies ܩோାand ܩோିு.

Unfortunately, ܩ௛௬ௗ values derived from these studies are not consistent (for example, ܩ௛௬ௗ values in acetonitrile were reported to be –400.7 kcal/mol,18 –404.7 kcal/mol58 and –412.7 kcal/mol.57

In collaboration with the Krylov group, we previously calculated the hydricity of an acridine-based hydride donor and the obtained value was in excellent agreement with the 103

34 experimental hydricity. In our approach, the absolute Gibbs energy ܩ௛௬ௗ was obtained as the

௚௔௦ ௦௢௟ sum of the gas-phase energy ܩ௛௬ௗ and the solvent contribution߂ܩ௛௬ௗ:

௚௔௦ ௦௢௟ ܩ௛௬ௗ ൌܩ௛௬ௗ ൅߂ܩ௛௬ௗ (v)

௚௔௦ ௦௢௟ The gas phase energy ܩ௛௬ௗ was calculated using DFT, while the solvation energy ߂ܩ௛௬ௗwas derived from the experimental one-electron reduction potential of hydrogen atom in a solvent of interest59 and the calculated gas-phase electron affinity of H-atom (as detailed in the computational section). The ܩ௛௬ௗvalues obtained in this way are −404.8 kcal/mol (in ACN) and −406.0 kcal/mol

(in DMSO). Here, this computational methodology was utilized to calculate the hydricities of our model hydrides in two solvents (ACN and DMSO, Table 4.1).

ି ڄ ڄ ା Table 4.1. Calculated standard reduction potentials (vs. NHE) for ܴ Ȁܴ and ܴ Ȁܴ , ݌ܭ௔ values for RH and ߂ܩுష for RH in different solvents.

+ . 1 . - 2 3 4 4 E1 (R /R ) E2 (R /R ) pKa (RH) ΔGH-(RH) ΔGH-(RH) Compound DMSO ACN

6OH 0.08 −1.27 30.4 84.1 85.8

4OH −0.61 −1.48 37.6 73.2 75.1

PhAcrH −0.25 −1.17 26.1 72.8 74.9

Me2N-AcrH −0.30 −1.20 25.6 70.3 72.2

CN-BNAH −0.69 −1.42 33.1 66.5 68.5

T-AcrH −0.09 −1.02 14.9 64.7 66.6

2OH −0.58 −1.40 27.0 61.1 62.9 104

HEH −1.04 −1.50 36.0 60.6 62.5

BNAH −0.94 −1.84 38.4 58.3 60.3

CAFH −1.87 −1.65 45.8 51.3 53.2

Me-MNAH −1.24 −1.63 34.1 50.3 52.2

BIMH −1.51 −1.69 38.4 48.6 50.3

3NH −1.07 −1.75 30.8 46.9 48.7

1 Calculated using ZB97X-D3/ Def2-TZVP using the SMD continuum solvation

model.

2 Calculated using ZB97X-D3/ Def2-TZVP using the SMD continuum solvation

model with a counter ion (K+).

3Calculated using the equation w.

4Calculated using ZB97X-D/6-311++G(2df,p)/CPCM (DMSO or ACN).

The calculations were also used to estimate the standard reduction potentials and pKa

+ . values of relevant species (Table 4.1). While E1(R /R ) values acquired using sole electronic

. energies showed a reasonable match with experimental values, the E2(R /R¯) values using standard

procedures did not match experiment well (mean unsigned error = 0.17 V). Since our original

calculated E2 values were consistently too negative compared to experiment, we speculated this was because the calculated free energies of R¯ states were systematically too unstable regardless of different exchange correlation functionals, continuum solvation methods, and basis set sizes.

As a simple correction and following previous work,60 we added a positively charges counter ion,

+, . K into the calculations on the R¯ and R states, and the resulting E2 values agreed with available 105

experimental data much better (mean unsigned error = 0.08 V). Adding an analogous counter ion,

Cl¯, to the states needed for the E1 calculations did not improve the agreement of calculated vs.

experiment. The obtained calculated reduction potentials are then used to estimate the pKa values

for the model hydride donors (equation w). However, the calculated pKa values (Table 4.1) are

not very accurate.

ಹశ ష ௱ீ ିଶଷǤ଴଺൭ாబ ାாబ Ǥ ൱ି௱ீಹ ಹష ೃశ ೃ ೞ೚೗ೡ೐೙೟ ష ݌ܭோு ൌ ೃǤ ೃ (w) ௔ ଵǤଷ଺ସ

4.3.2 Experimental Hydricity Values

Two experimental approaches namely “potential-pKa” method and “hydride transfer”

methods were used to determine the hydricity of compounds in two different solvents ACN and

DMSO. The “potential-pKa” method requires the standard reduction potentials and pKa values of

ι model compounds to determine hydricity as shown in Scheme 4.3. The value for οܩଵ (Scheme 4.3,

equation c) was calculated by determining the pKa value of deprotonation of R-H model

ι ι compounds, whereas οܩଶand οܩଷ (Scheme 4.3, equation d,e) were calculated by using the

଴ ଴ Ǥ + standard reduction potentials of first ቆܧೃశቇand second electron reduction ሺܧ ೃ ሻ of R ష ೃǤ ೃ

ι compounds. οܩସrepresents the gibbs free energy change for the two-electron reduction of proton.

This value is calculated by using the standard reduction potential of proton in the respective solvent

used to calculate the hydricity. The reported values are 69.9 kcal/mol in DMSO and 54.4 kcal/mol

in ACN.46 In case of the “hydride transfer method” the hydricity of model compounds were

calculated by using the equilibrium constant (ܭ௘௤) of reaction between the hydride donor (R-H)

and the appropriate hydride acceptor (A) whose hydricity is already known, as shown in Scheme 106

4.4. The equilibrium constant is used to calculate the gibbs free energy of hydride transfer between

the donor (R-H) and acceptor (A), equation (j).

4.3.2.1 Potential-pKa Method

This method is used to calculate the hydricity of model compounds in DMSO whose pKa

can be calculated experimentally. The pKa of the model compounds R-H can be experimentally

determined only if R-H is more acidic than the solvent (for DMSO, pKa=35, which limits the pKa

determination for compounds with pKa values lower than 32).8 The calculated values of pKa in

DMSO shows that the model compounds CN-BNAH, Me-MNAH, HEH, 4OH, BNAH, BIMH,

CAFH are higher than that of DMSO-limit, indicating that their hydricities are not likely to be

଴ ଴ Ǥ determined using the “potential-pKa” method. Also, the reduction potentials ܧೃశand ܧ ೃ need to ష ೃǤ ೃ

be less negative than the cathodic electrochemical window of the solvent. In case of DMSO the

electrochemical window is ~ -1.9 V vs NHE under our experimental condition (TBAP as

electrolyte and platinum working electrode).

pKa determination: The pKa of the R-H model compounds were calculated in DMSO by

using spectrophotometric indicator anion method developed by Bordwell.35 The indicators with

known pKa values were deprotonated using dimsyl anion and then reacted with R-H model

compounds. The concentration of the indicator anion was calculated from the observed absorbance

and determined extinction coefficient (using Beer’s law, Figure 4.1 a, b). Thus formed colored

indicator anion was quenched by adding aliquots of R-H model compounds ( Figure 4.1c for

6OH).To obtain the high accuracy, the acidity difference between the indicator and the R-H model

compounds should be within 2 pka units.35 The pKa values of the model compounds were 107

estimated from known indicator pKa values and experimentally obtained equilibrium constants of

the reaction between indicator anion and the model compounds (Table 4.2)

a) b) - Xan 2.0 45 μM 1.6 67 μM 89 μM 1.4 1.5 110 μM Xan- 132 μM 1.2

1.0 1.0 Absorbance 0.8 Abs @ 505nm 0.6 0.5 - -1 -1 E (Xan ) =14005 M cm

-6 0.0 60 80 100 120x10 [Xan-] (M) 300 400 500 600 Wavelength (nm) c) 0.8 Xan- 0.6

0.4 - Xan 6O- 10 μL 6OH Absorbance 0.2 20 μL 6OH 30 μL 6OH - 6O 0.0 300 400 500 600 Wavelength (nm)

Figure 4.1 pKa determination of 6OH by using indicator anion method: (a) Generation of indicator

anion (Xan–) by the reaction of indicator (XanH) in DMSO with potassium dimsyl. (b) Beer’s law plot for determination of the extinction coefficient of indicator anion at maximum absorbance (505 nm); (c) Titration of 24 mM 6OH with indicator anion.

Table 4.2 Indicators and their pKa values in DMSO (left two column), model compounds with experimentally determined pka values (right two columns).

Indicator pKa of Model compounds (R-H) Determined

Indicator pka

Triphenylmethane 30.6 4OH 30.4

Diphenyldiphenyl 29.4 Me2N-AcrH 29.2

methane 108

9-phenylxanthene 27.9 Ph-AcrH 28.3

6OH 26.9

Fluorene 22.6 2OH 23.4

Cyclic voltammetry was used to calculate the standard reduction potentials. At low scan rate (100 mV/s), the reversible peaks were obtained for the first electron reduction of 6O+ and

+ + + + + acridine based models (Ph-Acr , Me2N-Acr , 4O , 2O and 3N ) (Figure 4.2). This reversible behavior of peaks indicates a good chemical stability of the corresponding radicals. However, the first reduction potentials of pyridinium (CN-BNA+, HE+, BNA+ and ME-BNA+) and imidazolium

(BIM+ and CAF+) models are chemically irreversible and shifted to more negative potentials

(Figure 4.2). Such irreversibility has been previously reported by others and assigned to the dimerization of the radicals.61-64 The lower reactivity of acridine-based radicals over the pyridine- based structures is likely due to higher delocalization of the unpaired spin in the acridine-based radicals.62 However, in case of the acridine based T-Acr+, the first reduction peak is also irreversible at low scan rate (100 mv/s) and requires a high scan rate of 2 kV/s (Figure 4.2) to become fully reversible. The chemical instability of imidazolium radicals has been previously attributed to either their dimerization64 or to the loss of H-atom and formation of carbene analogs64-

67. Although the scan rate up to 10 kV/s was used, chemical reversibility for the one-electron reduction of pyridinium and imidazolium models could not be achieved. Thus, this irreversibility of peaks prevented us from obtaining the standard reduction potentials for all of the model compounds under study. The reversible peaks were obtained only for the compounds 6O+, T-Acr+,

+ + + + + Ph-Acr , Me2N-Acr , 4O , 2O and 3N (Figure 4.1, left). 109

+ 6O 6 μA·s/mm2·V + + CN-BNA T-Acr

+ + HE Ph-Acr

+ + BNA Me2N-Acr + + Me-BNA 4O

+ 2O + BIM + 3N + CAF

0.5 0.0 -0.5 -1.0 -1.5 -2.0 0.0 -0.5 -1.0 -1.5 -2.0 Potential vs. NHE (V) Potential vs. NHE (V)

Figure 4.2 Cyclic voltammograms of model compound (cations) in the cathodic range: Pt working electrode, Pt counter electrode, and non-aqueous Ag/AgNO3 reference electrode. Scan rate, 0.1

+ + + + + + + + – V/s (3N , CN-BNA , HE , BNA , Me-MBNA , BIM and CAF ), 25 V/s (6O , T-Acr , Me2N-

Acr+, 2O+), 2 kV/s (T-Acr+), 100 V/s (Ph-Acr+, 4O+); electrolyte: 0.1 M TBAP in DMSO. The second reduction peak of T-Acr+ was obtained from the oxidation of T-Acr– which was formed by the deprotonation of TAcr-H in the presence of dimsyl ion (pKa= 35).8

The second reduction peaks were obtained only for the model compounds whose first

+ + + + + + reduction peaks were reversible (6O , Ph-Acr , Me2N-Acr , 4O , 2O and 3N ). The second reduction peaks were irreversible for all of these compounds at low scan rates (100 mV/s) which may be due to the protonation of the generated anion to form NADH analogs.68 The results from the high scan rate show that the reactivity of R– anions correlate well with pKa values of the corresponding R-H analogs. For example, the second reduction peak of 6O+ becomes reversible at relatively low scan rates (25 V/s), which is consistent with relatively low basicity of 6O– anion 110

(the pKa of 6OH is 26.9). On the other hand, the reversibility for 4O+ requires the scan rates of

100 V/s and the pKa of 4OH is 30.3. Due to the sluggish electron transfer kinetics in case of T-

Acr+, the reversible second reduction peak could not be achieved even at high scan rate of 10 kV/s.

Thus, the standard reduction potential for this process was obtained by electrochemical oxidation

of T-Acr– anion, formed by the deprotonation of T-AcrH (Figure 4.2). This anion oxidation method

was also used for the compounds that lacked the second reduction peaks, but the experiment was

not successful because the pKa values of R-H compounds were well above the pKa of the solvent

under study.

଴ ଴ Ǥ Table 4.3 Experimentally obtained ܧೃశǡ ܧ ೃ and ݌ܭ௔ of model compounds for obtaining ష ೃǤ ೃ

experimental ߂ܩுି values using “potential-pKa” method in dimethyl-sulfoxide.

଴ ଴ ܧ శ ܧ Ǥ ோ ோ ߂ܩுି ோǤ ோష ோு Compound Pܭ௔

(V vs NHE) (V vs NHE) (kcal/mol)

6OH +0.24 -1.24 26.9 83.5

4OH -0.38 -1.41 30.4 70.2

PhAcrH -0.29 -1.23 28.3 73.5

Me2N-AcrH -0.30 -1.42 29.2 70.1

T-AcrH -0.22 -1.07 N/A N/A

2OH -0.50 -1.39 23.4 58.2

3NH -0.80 -1.62 N/A N/A

111

4.3.2.2 Hydride Transfer Method

Due to the irreversible reduction behavior of R+ models or due to the low pKa values of corresponding R-H models, the hydricities of all of the hydrides presented in Scheme 4.1 could not be obtained using the potential-pKa method. Thus hydride –transfer method was used to determine the hydricities of the remaining compounds. NMR spectroscopy was used to determine the equilibrium constant (ܭ௘௤) of the reaction between the ܴെܪ compounds of unknown hydricity and appropriate hydride acceptors ሺܣሻ of known hydricity values in two solvents: acetonitrile and dimethyl sulfoxide. The gibbs free energy of hydride transfer between the donor/acceptor molecules and hydricity of the reference molecule is used to calculate the hydricity of unknown (Scheme 4.4). As an example, the NMR spectra for the reaction between reaction between BNAH and 2O+ in DMSO is shown in figure 4.3 where 2O+ was used as the reference with known hydricity value calculated from potential-pKa method (Table 4.3). In this particular

+ case, the reaction between BNAH and 2O reach equilibrium after 19 days yielding Keq =1.68.

Similar successful experiments were performed to calculate the hydricity of some other model compounds in both solvents (DMSO and ACN) which is listed in Table 4.4. Due to the unwanted side reactions that occurred during the time required to achieve equilibrium, this method was not successful for all of the model compounds. 112

Figure 4.3 1H NMR spectra showing the reaction between BNAH and 2O+ in DMSO immediately

(red) and at equilibrium (blue).

Table 4.4 Hydricity of compounds determined by using hydride transfer method.

Compound Referent Hydride (οܩுష in kcal/mol) οܩுష (kcal/mol) Solvent

BNAH 2OH (58.2) 57.5 DMSO

HEH BNAH (57.5) 58 DMSO

BNAH CpRe(NO)(CO)(CHO)1 (55) 591 ACN

2OH BNAH (59) a 60.3 ACN 113

HEH BNAH (59) a 61.5 ACN

CN-BNAH BNAH (59) a 631 ACN

+ b BIMH [Ni(dmpe)2H] (49.9) 50.1 ACN

+ b 3NH [Ni(dmpe)2H] (49.9) 49.2 ACN

aRef.19 bRef.69

4.3.3 Comparison of Hydricity Values Obtained from Calculated and Experimental Methods

The hydricity values obtained experimentally by using potential-pka method and hydride transfer method show the good agreement with calculated values (Figure 4.3). The calculated values of hydricities are consistently higher than the experimental values by up to 3 kcal/mol. The reason for higher calculated hydricity values may be due to the uncertainties associated with the treatment of the hydride ion solvation in calculations. Similar match between experiment and theory has been reported previously for metal-based hydrides.17,18 114

Figure 4.4 Comparison between the calculated and experimental hydricities for the NADH- analogs. aValue obtained from our previous work (ref. 34). b Values obtained from ref. 19

4.3.4 Structural Effect and Solvent Effect on Hydricity Values

From the above obtained hydricity values by the experimental and calculation methods, it was found that the hydride donor abilities of R-H derivatives depend on the structural properties of model hydrides such as the hydricity is drastically improved by the stabilization of the formed cation, R+ after the hydride transfer, either by aromatization or charge delocalization. For example, the hydricities of model compounds 6OH, 4OH, 2OH and 3NH can be compared. As the planarity

(and charge delocalization) of the molecular framework increases from 6O+ to 3N+, the hydricities of the corresponding hydride donors range from 86 kcal/mol (calculated for 6OH in ACN) to 49 kcal/mol (calculated for 3NH in ACN). Also, it was found that the positive charge in R+ can be 115

stabilized through the inductive effect of electron-donating groups which was observed for the

NADH-analogs: methyl group in Me-MNAH (οܩுష = 50 kcal/mol), amide group in BNAH

(οܩுష= 59 kcal/mol), and cyano group in CN-BNAH (οܩுష= 63 kcal/mol). Moreover, it was

observed that the hydricity of Me2N-AcrH (οܩுష=70.1 kcal/mol) is lower relative to the derivative

without donating group (PhAcrH, οܩுష= 73.5 kcal/mol). Among the studied model hydrides,

imidazole based hydrides (BIMH and CAFH) have low hydricities (calculated ACN values in the

50-53 kcal/mol range). The high hydride donor ability of imidazole derivatives has been previously

credited to the specific conformation and an anomeric effect,70 where neighboring nitrogen centers

destabilize the C-H bond in R-H by donating their lone pairs to its antibonding orbital.

Moreover, the hydride donor abilities were found to be dependent on the solvent. From our

study (experimental and calculation) by using ACN and DMSO as solvent, it was found that the

compounds possess better hydricity values (by ~ 2 kcal/mol) in DMSO compared to ACN. The

difference in values with solvents can be explained in terms of the difference in the dielectric

constants (ε) of solvents. The dielectric constant of DMSO is slightly higher than dielectric

constant of ACN (ε(DMSO) = 47 vs. ε(ACN) = 38), which results in slightly better solvation of

formed charged species (R+ and H¯). Such kind of solvent effects have already been reported for

Ni-based hydride.2

4.3.5 Comparison with Metal-based Analogs

Transition metal hydrides have been identified as important intermediates in a variety of catalytic fuel-forming and other redox reactions in ground and excited state.14-16,71-74 On the other

hand, the metal-free hydrides have not been widely used for this purpose, despite the abundance

11-13,75-77 of enzymatic catalysis by NADH, FADH2 and other metal-free hydrides. Thus, it is 116

interesting to compare the thermodynamic hydricities of our model compounds with those reported

for metal-based hydrides. In general, the hydricities of metal-based hydrides span a wide range of

values (reported acetonitrile hydricities are in the 26-120 kcal/mol range69,78). The metal-free

hydrides of Scheme 4.1 exhibit values that are somewhat higher, with the calculated values in the

49-86 kcal/mol range, and seem to have hydricities similar to the first raw transition metal

hydrides, such as Co, Ni and Fe-based compounds (acetonitrile hydricities are in the 32-73

kcal/mol range69). Thus, both types of compounds are sufficiently strong hydride donors for the relevant fuel forming reactions. For example, the hydride affinity of protons in acetonitrile is ~ –

76.6 kcal/mol,79 indicating that most metal-free hydride donors in Scheme 4.1 are thermodynamically capable of driving the hydrogen evolution reaction.

However, metal-free hydride donors have not been used in fuel-forming reactions as often

as compared to their metal-based analogs. One possible explanation might be related to the

differences in the activation barriers associated with the relevant hydride transfer processes.

Another reason for lower use of metal-free hydride catalysts might be related to the closure of

catalytic cycle, which involves the two-electron, proton-coupled reduction of R+ to recover the

active R-H hydride form. To exemplify this point, Figure 4.5 presents an energy diagram for two

Ph Ph + hydride donors of similar hydricities: a metal-based [Ni(P 2N )2H] complex, whose hydricity in

80 acetonitrile is οܩுష = 59.3 kcal/mol, and 2OH, whose hydricity is οܩுష = 60.3 kcal/mol. The

first reduction potentials of the corresponding precursors, E(M2+/M+) and E(R+/R.), are relatively

similar and affordably small (around – 0.5 V vs. NHE). Previous studies of metal-based

compounds have shown that E (M2+/M+) shifts to more negative values as the hydride donor ability

of the corresponding donor increases.69,81-84 Our metal-free analogs scale in the same way, as

exemplified by the similarities in the first reduction peaks for M2+/M+ and R+/R.. However, a 117

striking difference was observed in the values for the second reduction potentials, E(M+/M0) and

E(R./R–): while the metal-based compound undergoes the second reduction at a relatively low potential (–0.5 V vs. NHE), the metal-free model compound requires a significantly more negative potential (–1.4 V vs. NHE) to inject the second electron.

Figure 4.5 The comparison of energy required to regenerate the hydride forms of metal-based

Ph Ph + 80 ([Ni(P 2N )2H] ) and metal-free (2OH) hydride donors.

Such a large energy requirement for the second reduction step prohibits the application of metal-free hydride donors in catalysis. Several approaches can be used to lower the standard reduction potentials in metal-free systems. One involves the coupling of the first electron transfer step with a proton transfer to generate RH. +, which will then be reduced at a less negative potential.

Such proton-coupled reduction has been selected by nature as a way to regenerate NADH.

+ + Specifically, NAD reduction is mediated by FADH2, which is formed from FAD through two

proton-coupled reductions.85 Another method involves a design of model compounds that exhibit 118 small or even inverted differences in reduction potentials.86 For example, it was shown for some organic compounds that the structural changes that accompany the first electron reduction can result in the lowering of their LUMO orbital energies and associated second reduction potentials.87

We observe similar effects in the case of BIM+, which exhibits the smallest energy difference between the calculated E(R+/R.) and E(R./R–) potentials (Table 4.1), likely brought about by the rotation of the phenyl ring upon one electron reduction.

4.4 Conclusions

The hydride donating ability of thirteen different organic hydrides were studied in acetonitrile and DMSO to explore their applicability in fuel-forming catalysis. The results show that hydricity values of these organic hydrides are comparable to those of metal hydrides and most of the hydrides are capable to reduce proton in acetonitrile. However, the second reduction potential of these metal-free hydrides are higher in values, prohibiting the catalyst recovery. Also, the hydricity values of these compounds depend on the solvent, lower hydricity values were obtained for all compounds in DMSO compared to acetonitrile. The structure of the compounds also plays the vital role in the hydricity; such as aromaticity, planarity, and presence of electron donating group. For example, benzoimidazole structure (BIMH) has the lowest hydricity value whereas triphenylmethane structure (6OH) has highest hydricity value.

4.5 References

(1) Ellis, W. W.; Ciancanelli, R.; Miller, S. M.; Raebiger, J. W.; Rakowski DuBois,

M.; DuBois, D. L. J. Am. Chem. Soc. 2003, 125, 12230. 119

(2) Tsay, C.; Livesay, B. N.; Ruelas, S.; Yang, J. Y. J. Am. Chem. Soc. 2015, 137,

14114.

(3) Berning, D. E.; Noll, B. C.; DuBois, D. L. J. Am. Chem. Soc. 1999, 121, 11432.

(4) Bard, A. J.; Faulkner, L. R.; Leddy, J.; Zoski, C. G. Electrochemical methods: fundamentals and applications; wiley New York, 1980; Vol. 2.

(5) Pitman, C. L.; Brereton, K. R.; Miller, A. J. J. Am. Chem. Soc. 2016, 138, 2252.

(6) Curtis, C. J.; Miedaner, A.; Ellis, W. W.; DuBois, D. L. J. Am. Chem. Soc. 2002,

124, 1918.

(7) Izutsu, K.; Pure, C. o. E. C. I. U. o.; Chemistry, A. Acid-base dissociation constants in dipolar aprotic solvents; Blackwell Scientific Publications Oxford, 1990; Vol. 35.

(8) Bordwell, F. G. Acc. Chem. Res. 1988, 21, 456.

(9) Kütt, A.; Leito, I.; Kaljurand, I.; Sooväli, L.; Vlasov, V. M.; Yagupolskii, L. M.;

Koppel, I. A. J. Org. Chem. 2006, 71, 2829.

(10) Pollak, N.; Dölle, C.; Ziegler, M. Biochemical Journal 2007, 402, 205.

(11) Hummel, W. Trends in biotechnology 1999, 17, 487.

(12) Reda, T.; Plugge, C. M.; Abram, N. J.; Hirst, J. Proc. Natl. Acad. Sci. 2008, 105,

10654. 120

(13) Schrittwieser, J. H.; Velikogne, S.; Kroutil, W. Adv. Synth. Catal. 2015, 357,

1655.

(14) Bullock, R. M.; Appel, A. M.; Helm, M. L. Chem. Commun. 2014, 50, 3125.

(15) Rakowski Dubois, M.; Dubois, D. L. Acc. Chem. Res. 2009, 42, 1974.

(16) DuBois, D. L.; Berning, D. E. Appl. Organomet. Chem. 2000, 14, 860.

(17) Qi, X.-J.; Fu, Y.; Liu, L.; Guo, Q.-X. Organometallics 2007, 26, 4197.

(18) Nimlos, M. R.; Chang, C. H.; Curtis, C. J.; Miedaner, A.; Pilath, H. M.; DuBois,

D. L. Organometallics 2008, 27, 2715.

(19) Ellis, W. W.; Raebiger, J. W.; Curtis, C. J.; Bruno, J. W.; DuBois, D. L. J. Am.

Chem. Soc. 2004, 126, 2738.

(20) Cheng, J.-P.; Handoo, K. L.; Xue, J.; Parker, V. D. J. Org. Chem. 1993, 58, 5050.

(21) Cheng, J.-P.; Handoo, K. L.; Parker, V. D. J. Am. Chem. Soc.1993, 115, 2655.

(22) Zhu, X.-Q.; Wang, C.-H.; Liang, H.; Cheng, J.-P. J. Org. Chem. 2007, 72, 945.

(23) Cheng, J.; Handoo, K. L.; Xue, J.; Parker, V. D. J. Org. Chem. 1993, 58, 5050.

(24) Zhang, X.-M.; Bruno, J. W.; Enyinnaya, E. J. Org. Chem. 1998, 63, 4671.

(25) Ilic, S.; Brown, E. S.; Xie, Y.; Maldonado, S.; Glusac, K. D. J. Phys. Chem. C

2016, 120, 3145. 121

(26) W Laursen, B.; C Krebs, F. Chem. Eur. J. 2001, 7, 1773.

(27) Paul, C. E.; Gargiulo, S.; Opperman, D. J.; Lavandera, I.; Gotor-Fernández, V.;

Gotor, V.; Taglieber, A.; Arends, I. W.; Hollmann, F. Org. Lett. 2012, 15, 180.

(28) Huang, S.; Wong, J. C.; Leung, A. K.; Chan, Y. M.; Wong, L.; Fernendez, M. R.;

Miller, A. K.; Wu, W. Tetrahedron Lett. 2009, 50, 5018.

(29) Ogata, Y.; Takagi, K.; Tanabe, Y. Perkin Trans. 1979, 1069.

(30) Lee, C. K.; Lee, I.-S. H. Heterocycles 2009, 78, 425.

(31) Hori, M.; Kataoka, T.; Shimizu, H.; Imai, E.; Matsumoto, Y. Chem. Pharm. Bull.

1985, 33, 3681.

(32) Carey, F. A.; Tremper, H. S. J. Am. Chem. Soc. 1968, 90, 2578.

(33) Guin, J.; Besnard, C.; Pattison, P.; Lacour, J. Chem. Sci. 2011, 2, 425.

(34) Yang, X.; Walpita, J.; Zhou, D.; Luk, H. L.; Vyas, S.; Khnayzer, R. S.; Tiwari, S.

C.; Diri, K.; Hadad, C. M.; Castellano, F. N. J. Phys. Chem. B 2013, 117, 15290.

(35) Matthews, W. S.; Bares, J. E.; Bartmess, J. E.; Bordwell, F.; Cornforth, F. J.;

Drucker, G. E.; Margolin, Z.; McCallum, R. J.; McCollum, G. J.; Vanier, N. R. J. Am. Chem.

Soc. 1975, 97, 7006.

(36) Robinson, S. J. C.; Zall, C. M.; Miller, D. L.; Linehan, J. C.; Appel, A. M. Dalton

Trans. 2016, 45, 10017. 122

(37) Cheng, J.-P.; Lu, Y.; Zhu, X.; Mu, L. J. Org. Chem. 1998, 63, 6108.

(38) Pavlishchuk, V. V.; Addison, A. W. Inorg. Chim. Acta. 2000, 298, 97.

(39) 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 Jr., J. A.; Peralta, J. E.; Ogliaro, F.; Bearpark,

M. J.; Heyd, J.; Brothers, E. N.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.;

Raghavachari, K.; Rendell, A. P.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.;

Millam, N. J.; 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, USA, 2009.

(40) Chai, J.-D.; Head-Gordon, M. J. Chem. Phys. 2008, 128, 084106.

(41) Chai, J.-D.; Head-Gordon, M. Phys. Chem. Chem. Phys. 2008, 10, 6615.

(42) Barone, V.; Cossi, M. J. Phys. Chem. A 1998, 102, 1995.

(43) Kelly, C. P.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. B 2007, 111, 408.

(44) Muckerman, J. T.; Skone, J. H.; Ning, M.; Wasada-Tsutsui, Y. Biochim. Biophys.

Acta 2013, 1827, 882. 123

(45) Lykke, K.; Murray, K.; Lineberger, W. Phys. Rev. A 1991, 43, 6104.

(46) Handoo, K. L.; Cheng, J. P.; Parker, V. D. J. Am. Chem. Soc. 1993, 115, 5067.

(47) Becke, A. D. J. Chem. Phys. 1993, 98, 5648.

(48) Grimme, S.; Ehrlich, S.; Goerigk, L. J. Comput. Chem. 2011, 32, 1456.

(49) Weigend, F.; Ahlrichs, R. Phys. Chem. Chem. Phys. 2005, 7, 3297.

(50) Marenich, A. V.; Cramer, C. J.; Truhlar, D. G. J. Phys. Chem. B 2009, 113, 6378.

(51) Perdew, J. P. Phys. Rev. B 1986, 33, 8822.

(52) Neese, F. Wiley Interdiscip Rev Comput Mol Sci 2012, 2, 73.

(53) Isse, A. A.; Gennaro, A. J. Phys. Chem. B 2010, 114, 7894.

(54) Shi, J.; Huang, X.-Y.; Wang, H.-J.; Fu, Y. J. Chem. Inf. Model. 2011, 52, 63.

(55) Zhu, X.-Q.; Wang, C.-H.; Liang, H.; Cheng, J.-P. J. Org. Chem. 2007, 72, 945.

(56) Lim, C.-H.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B. J. Am. Chem. Soc.

2014, 136, 16081.

(57) Muckerman, J. T.; Achord, P.; Creutz, C.; Polyansky, D. E.; Fujita, E. Proc. Natl.

Acad. Sci. U.S.A. 2012, 109, 15657

(58) Kovács, G.; Pápai, I. Organometallics 2006, 25, 820. 124

(59) Parker, V. Acta Chem. Scand. 1992, 46, 1133.

(60) Keith, J. A.; Grice, K. A.; Kubiak, C. P.; Carter, E. A. J. Am. Chem. Soc. 2013,

135, 15823.

(61) Hapiot, P.; Moiroux, J.; Saveant, J. M. J. Am. Chem. Soc. 1990, 112, 1337.

(62) Anne, A.; Hapiot, P.; Moiroux, J.; Savéant, J.-M. J. Electroanal. Chem. 1992,

331, 959.

(63) Hermolin, J.; Levin, M.; Kosower, E. M. J. Am. Chem. Soc. 1981, 103, 4808.

(64) de Robillard, G.; Devillers, C. H.; Kunz, D.; Cattey, H.; Digard, E.; Andrieu, J.

Org. Lett. 2013, 15, 4410.

(65) Gorodetsky, B.; Ramnial, T.; Branda, N. R.; Clyburne, J. A. Chem. Comm. 2004,

1972.

(66) Ogawa, K. A.; Boydston, A. J. Chem. Lett. 2014, 43, 907.

(67) Han, X.; Hao, W.; Zhu, X.-Q.; Parker, V. D. J. Org. Chem. 2012, 77, 6520.

(68) Koper, N.; Jonker, S.; Verhoeven, J.; Van Dijk, C. Recl. Trav. Chim. Pays-Bas

1985, 104, 296.

(69) Wiedner, E. S.; Chambers, M. B.; Pitman, C. L.; Bullock, R. M.; Miller, A. J.;

Appel, A. M. Chem. Rev. 2016, 116, 8655.

(70) Brunet, P.; Wuest, J. D. J. Org. Chem. 1996, 61, 2020. 125

(71) Thoi, V. S.; Sun, Y.; Long, J. R.; Chang, C. J. Chem. Soc. Rev. 2013, 42, 2388.

(72) Miller, A. J.; Labinger, J. A.; Bercaw, J. E. Organometallics 2011, 30, 4308.

(73) Pitman, C.; Finster, O.; Miller, A. Chem. Comm. 2016, 52, 9105.

(74) Barrett, S. M.; Pitman, C. L.; Walden, A. G.; Miller, A. J. J. Am. Chem. Soc.

2014, 136, 14718.

(75) Stueckler, C.; Hall, M.; Ehammer, H.; Pointner, E.; Kroutil, W.; Macheroux, P.;

Faber, K. Org. Lett. 2007, 9, 5409.

(76) Stuermer, R.; Hauer, B.; Hall, M.; Faber, K. Curr. Opin. Chem. Biol. 2007, 11,

203.

(77) Shima, S.; Pilak, O.; Vogt, S.; Schick, M.; Stagni, M. S.; Meyer-Klaucke, W.;

Warkentin, E.; Thauer, R. K.; Ermler, U. Science 2008, 321, 572.

(78) DuBois, D. L. Inorg Chem 2014, 53, 3935.

(79) Wayner, D. D.; Parker, V. D. Acc. Chem. Res. 1993, 26, 287.

(80) Fraze, K.; Wilson, A. D.; Appel, A. M.; Rakowski DuBois, M.; DuBois, D. L.

Organometallics 2007, 26, 3918.

(81) Berning, D. E.; Miedaner, A.; Curtis, C. J.; Noll, B. C.; Rakowski DuBois, M. C.;

DuBois, D. L. Organometallics 2001, 20, 1832. 126

(82) Yang, J. Y.; Bullock, R. M.; Shaw, W. J.; Twamley, B.; Fraze, K.; DuBois, M. R.;

DuBois, D. L. J. Am. Chem. Soc. 2009, 131, 5935.

(83) Galan, B. R.; Schöffel, J.; Linehan, J. C.; Seu, C.; Appel, A. M.; Roberts, J. A.;

Helm, M. L.; Kilgore, U. J.; Yang, J. Y.; DuBois, D. L. J. Am. Chem. Soc. 2011, 133, 12767.

(84) Yang, J. Y.; Smith, S. E.; Liu, T.; Dougherty, W. G.; Hoffert, W. A.; Kassel, W.

S.; DuBois, M. R.; DuBois, D. L.; Bullock, R. M. J. Am. Chem. Soc. 2013, 135, 9700.

(85) Carrillo, N.; Ceccarelli, E. A. Eur. J. Biochem. 2003, 270, 1900.

(86) Evans, D. H. Chem. Rev. 2008, 108, 2113.

(87) Evans, D. H.; Busch, R. W. J. Am. Chem. Soc. 1982, 104, 5057.