Accessing Transition Metal Hydride Complexes Through Proton- Coupled Electron Transfer Processes

ACCESSING TRANSITION METAL HYDRIDE COMPLEXES THROUGH PROTON- COUPLED ELECTRON TRANSFER PROCESSES

Daniel A. Kurtz

A dissertation submitted to the faculty at the University of North Carolina at Chapel Hill in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Department of Chemistry.

Chapel Hill 2019

Approved by:

Jillian Dempsey

Gerald Meyer

Alexander Miller

Joanna Atkin

Erik Alexanian

ABSTRACT

Daniel A. Kurtz: Accessing Transition Metal-Hydride Complexes Through Proton-Coupled Electron Transfer Processes (Under the direction of Jillian Dempsey)

The combustion of fossil fuels such as coal, natural gas, and oil is the primary source of energy to power our everyday lives. These energy sources are non-renewable and their continued consumption will only accelerate the already evident warming of the planet. One of the primary products of burning fossil fuels, carbon dioxide (fossil fuels are carbon-based fuels), has increased by 24% in the past 60 years which can be explained by the increased global energy demand and subsequent acceleration of fossil fuel consumption. It is unreasonable, though, to immediately stop our use of fossil fuels as an energy source without a viable replacement.

Although it is 93 million miles away, the largest energy source available to humans is the sun. Nature has evolved to utilize the vast quantity of solar irradiation that illuminates the earth in order to perform complex photochemical reactions such as the biosynthesis of sugars. While there are researchers who are attempting to mimic the structure and function of these biological systems, the research discussed in this dissertation is focused on a more energy-relevant transformation: the conversion of solar energy into the simplest chemical fuel, hydrogen fuel.

Transition metal complexes are capable of performing all the required steps needed to convert sunlight into hydrogen fuel: absorption of solar radiation, and catalytic orchestration of the combination of the two electrons and two protons required to form H2. While known technology exists for this conversion, the efficiency of all required steps are currently not at a level that makes the system economically/commercially viable. The work discussed herein is

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related to key steps in the path from solar energy to H2. The first chapter is focused on factors that influence the formation of a transition metal hydride complex, a key intermediate in catalytic hydrogen production. Chapter 2 then describes a model system in which to study hydride formation using two different methods is described, and Chapter 3 discusses systematic variation in structural parameters on the model system and how they influence hydride formation is then discussed. Finally, presented in Chapter 4 are efforts to improve the spectral absorbance of an existing class of chromophores, followed by efforts to directly form transition metal hydrides via light absorption in Chapter 5.

ACKNOWLEDGEMENTS

First and foremost, Jill. I can’t express how thankful I am to have been able to join your lab to pursue my PhD. You’ve been an amazing mentor throughout the last four and a half years.

You’ve let me guide myself when I wanted to, and pushed me in areas where I need it. Every visitation weekend, the prospective students ask the question: “What’s Jillian like as a mentor?” and my answer is always the same: “She’s exactly the mentor you need her to be”, and you’ve been precisely that for me. I remember in our first meeting when I joined the lab, you told me that I was going to be your “no organic substrates” guy, and it seems like we’ve been on the same page and wavelength since then. I’ve gotten everything I wanted out of my PhD and more, and you’re the main reason for that. Thank you.

Next, my other PhD “parent”, Alex. You’ve become like a second mentor to me, and I greatly appreciate that. I truly feel like I’ve been an adopted member of the Miller lab in my time here (I think being the only Dempsey lab member to attend a full “paper meeting” makes that official). From partaking in some of the (many) traditions, to being on the other side of the camera for a number of group photos, and of course the awesome science and publications that

I’ve had the pleasure of working on with you and your lab. Thank you for the chemistry and basketball chats over the years.

Jerry, thank you for being my orals and thesis committee chair and for all the help over the years. Your excitement during a conversation of ours at a visitation weekend with Matt

Brady helped to spark the inspiration for Chapter 4 of my thesis, and hopefully a bright new direction for researchers down the road. I never understood your devotion to those cheese-heads

from the other side of Lake Michigan, but I did enjoy our friendly football rivalry over the years.

To Joanna and Erik, I appreciate you being on my thesis committee and helping to guide me over this final hurdle. Your time, advice, and discussion is invaluable to my degree, thank you.

As the 4th cohort of Dempsey lab, I have two very different groups of lab mates to thank.

The first group are those that left before me: Thomas for infecting me with the UNC basketball bug, Robin for handing down her wisdom and knowledge of the laser lab, Brian for being my first desk mate (sorry Eric) and teaching me how to do my first reactions in lab, Eric for being a phenomenal friend both during and after you left, Chris for sharing my interest and excitement over “neat and weird” science facts, and Kelley for the introduction of “Disney Fridays” in lab.

You all made me feel welcome in lab, and made it a fun place to work. The second group are those who will leave after me: Katherine for all the amazing baked goods over the years, Melody for the fun MCU talks at lunch whenever the new movies come out, Carolyn for all the laughs and for listening to my complaints and being comfortable enough to share yours, Michael for having such capable and meticulous hands to hand the laser lab over to, Tayliz for keeping the sass of A tower alive, and Aldo and Brittany for being bright lights in our lab’s future. I thank you all and am looking forward to keeping in touch and seeing how you all develop in your time remaining in the group and in the future.

I’ve gotten to work with some phenomenal post-doctoral mentors during my time here.

Having Noémie as my desk mate for over half of my time here was instrumental in my growth as a chemist. You and Matt really taught me to think critically about everything and to not get over- excited about data before checking everything about it. Thank you both for being great mentors and great friends, and I’m looking forward to seeing your respective groups flourish over the years. Tao and Debanjan, thank you for always being there to bounce ideas off of in lab and for

setting an example on what hard work means in the Dempsey lab. Good luck in your futures, I’m sure they’ll be great.

I’ve had the pleasure of working with some fabulous undergraduate researchers here.

Thank you to Hui-Min and Kevin who both took charge of their projects and helped me push the rhenium chromophore project through. Kevin, thanks for all of the laughs over the last 3 years, the basketball and football talks, the chemistry we’ve done together, and for challenging me to become more creative in the lab. You’ve been a phenomenal colleague to work with, and I know you’re going to go on to have a great grad school experience and career.

My friends outside of the Dempsey lab, and outside of UNC chemistry, have gotten me through some of the not so fun times in grad school. Thanks to Kyle and Nick for coming to visit me from Michigan, thanks to Earl for being an awesome roommate for the past few years, to Ali,

Andreas, and Andrew for being great friends throughout the 5 years we’ve been here together, to

Kelsey for always being there to talk about the good things and the bad even across many state lines, and thanks to all of the other people around the department that I don’t have room to list

(otherwise this section would be twice as long as it already is).

I wouldn’t be where I am today without all of my parents’ help and guidance both before

I came to grad school and during it. They’ve always wanted what’s best for me from the time I was born until the time I’m going to be defending this thesis, and I can’t thank them enough for that. They say it takes a village, and my family being about the size of one has really helped me become the person I am today. I love you all for everything you’ve done and the sacrifices you’ve made over the years, thank you.

Finally, thank you to Ann Marie. These last few years with you have been the best of my time here in Chapel Hill. You’ve been my support through everything, and I can’t thank you

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enough. From our advanced nature walks, to our trips to Michigan and Florida, to “raising”

Wynnstan, to you offering and wanting to read every piece of writing I’ve done, your constant presence in my life is something I cherish. The next few years in Boston will be tough without you there, but I’m very much looking forward to our future together. I lava you.

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TABLE OF CONTENTS

LIST OF FIGURES ...... xiv

LIST OF TABLES ...... xxvii

LIST OF SCHEMES...... xxviii

LIST OF ABBREVIATIONS AND SYMBOLS ...... xxix

CHAPTER 1: KINETIC ACIDITY OF TRANSITION METAL HYDRIDE COMPLEXES ...... 1

1.1 Introduction and Background ...... 1

1.2 Kinetic Acidity ...... 2

1.3 Intrinsic Kinetic Barriers and Self-Exchange ...... 6

1.4 Circumvention of Intrinsic Barriers – Relevance to Catalysis ...... 9

1.5 Methods of Determining Rate Constants for TM Protonation...... 11

REFERENCES ...... 13

CHAPTER 2: VERSATILE MODEL SYSTEM FOR METAL-HYDRIDE FORMATION ...... 18

2.1 Introduction ...... 18

2.2 General Results ...... 21

2.2.1 Synthesis of Complexes ...... 21

2.3 Electrochemical PCET ...... 21

2.3.1 Electrochemistry in the Absence of Acids ...... 21

2.3.1 Voltammetric Response with Added Acid ...... 24

2.3.2 Brønsted Relationship ...... 25

2.4 Photoinduced PCET ...... 28

2.4.1 Optical Characterization of PCET Intermediates ...... 28

2.4.2 Photoinduced Reactivity in the Absence of Acids ...... 31

2.4.3 Kinetic Processes with Added Acid ...... 35

2.5 Discussion ...... 37

2.5.1 Disparity Between Measured Protonation Rate Constants ...... 37

2.6 Conclusions ...... 43

REFERENCES ...... 45

CHAPTER 3: STERIC INFLUENCE ON KINETIC ACIDITY OF METAL HYDRIDES ...... 48

3.1 Introduction ...... 48

3.2 Results and Discussion ...... 50

3.2.1 Synthesis and Characterization ...... 50

3.2.2 Electrochemistry in the Absence of Acids ...... 52

3.2.3 Voltammetric Response with Added Acid ...... 54

3.2.4 Brønsted Relationship ...... 55

3.2.5 Intrinsic Barriers – Self-Exchange ...... 60

3.3 Conclusions ...... 63

3.4 Experimental Details ...... 64

REFERENCES ...... 68

CHAPTER 4: BATHOCHROMIC SHIFTS IN RHENIUM CARBONYL DYES INDUCED THROUGH DESTABILIZATION OF OCCUPIED ORBITALS ...... 70

4.1 Introduction ...... 70

4.2 Results ...... 72

4.2.1 Synthesis ...... 72

4.2.2 Structural Characterization in Solution and Solid State ...... 74

4.2.3 Photophysical Characterization ...... 81

4.2.4 Electrochemical Studies ...... 85

4.2.5 Computational Studies ...... 86

4.3 Discussion ...... 89

4.3.1 Redox properties and Electrochemically Induced Reactivity ...... 89

4.3.2 Photophysics and Excited-State Redox Properties ...... 93

4.3.3 Influence of Axial Ligands on Electronic Structure ...... 96

4.4 Conclusion ...... 99

REFERENCES ...... 101

CHAPTER 5: TOWARDS THE FORMATION OF TRANSITION METAL HYDRIDE COMPLEXES VIA LIGAND-TO-METAL CHARGE TRANSFER EXCITATION ...... 106

5.1 Introduction ...... 106

5.2 Results and Discussion ...... 109

5.2.1 Previous Synthetic Methods ...... 109

5.2.2 Alternative Synthetic Methods ...... 111

5.2.3 Differences Between Aryl and Alkyl Phosphines ...... 116

5.3 Conclusions and Future Directions ...... 118

5.4 Experimental Details ...... 118

REFERENCES ...... 120

APPENDIX A: ADDITIONAL DETAILS AND CHARACTERIZATION FOR CHAPTER 4 ...... 122

A.1 Experimental Details ...... 122

A.1.1 General Considerations...... 122

A.1.2 Electrochemical Measurements...... 122

A.1.3 Time-Resolved Photoluminescence...... 123

A.1.4 Single-crystal X-ray diffraction...... 124

A.1.5 Computational Details...... 125

A.1.6 Quantum Yield Measurement Experimental Details ...... 125

A.1.7 Synthesis ...... 125

A.2 Additional Characterization ...... 129

APPENDIX B: ADDITIONAL ELECTROCHEMICAL AND PROTON TRANSFER RATE CONSTANT DETERMINATION DATA FOR CHAPTERS 2 AND 3 ...... 134

APPENDIX C: PHOTOCHEMICAL REACTIVITY OF IRIDIUM HYDRIDE COMPLEXES ...... 204

C.1 Introduction and Background ...... 204

C.2 Results and Discussion ...... 205

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C.3 Conclusions ...... 213

REFERENCES ...... 214

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LIST OF FIGURES

Figure 1.1 Linear-free energy relationships for a series of TMHC complex deprotonation by substituted anilines. ((CO)4CoH α = 0.48, (CO)5FeH2 α = 31 0.55, (CO)5MnH α = 0.54, and Cp(CO)3WH α = 0.65)...... 4

Figure 1.2 Deprotonation of HCo(CO)4 by NEt3 to form [TEAH][Co(CO)4], and carbonyl stretching frequencies for the TMHC and its conjugate base...... 7

Figure 1.3 Potential energy vs q (A-H distance) profiles at Q = Q‡ (A-B distance at the transition state). Comparison between a carbon acid, CH4, and two 43 Eigen acids, NH3 and H2O...... 8

Figure 1.4 Comparison between direct protonation of metal center to form metal hydride and protonation at pendant amine followed by intramolecular rearrangement to form metal hydride...... 10

Figure 2.1 Square scheme showing stepwise ET-PT and PT-ET pathways and concerted pathway to form a transition metal hydride intermediate in H2 formation ...... 19

III Figure 2.2 Cyclic voltammograms of Co -dppe-NCCH3 in CH3CN (0.50 mM, blue, bottom) and in CH2Cl2 (0.48 mM, red, top). The voltammograms were recorded at 100 mV/s in 0.25 M TBAPF6...... 22

III Figure 2.3 (Top) Cyclic voltammograms of 0.48 mM of Co -dppe-NCCH3 in 10 mL of CH2Cl2 with 0.25 M TBAPF6 as supporting electrolyte at a scan rate ( of 50, 100, 250, 500, 750, and 1000 mV/s. (Bottom) Plot of the peak potential of the irreversible CoIII/II cathodic wave as a function of the natural logarithm of the inverse of scan rate. The intercept is −0.347 V 7 −1 which yields an estimated kd of 4 × 10 s ...... 24

I Figure 2.4 Determination of kPT of Co -dppe with 4-cyano-anilinium tetrafluoroborate (Left) Cyclic voltammograms of 0.5 mM of CoIII-dppe- NCCH3 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.006 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. (Right) 2 Evolution of kobs obtained with acid concentration (R = 0.997). The slope 7 −1 −1 yields kPT = 1.7 × 10 M s ...... 25

Figure 2.5 (Left) Spectrophotometric titration in 5% toluene in CH3CN starting with 0.137 mM HCoIII-dppe (blue trace ─) and titrating in increasing amounts of NEt3 (gray traces: 0.40, 0.80, 1.20, and 1.61 equivalents, respectively). End point of titration is reached upon the addition of 6.4 equivalents of NEt3 (red trace ─), when spectral profile matches that of CoI-dppe. The absorbance at 500 nm and the extinction coefficients of HCoIII-dppe and CoI-dppe were used to calculate concentration of each species as a function of added base. (Right) Linear regression of the

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equation for the acid base equilibrium between the species and the known III pKa of NEt3 was used to determine the pKa of HCo -dppe as 18.4...... 26

Figure 2.6 Brønsted plot of kPT values vs ΔpKa for the protonation reaction of CoI-dppe to form HCoIII-dppe ...... 27

II Figure 2.7. Cyclic voltammograms of Co -dppe recorded in CH3CN with 0.25 M TBAPF6. The open-circuit potential (OCP) was determined prior to recording the cyclic voltammogram, and the scan was begun at the OCP (arrow indication on the voltammogram)...... 29

Figure 2.8. (Top) UV/Vis absorption spectra of spectroelectrochemical oxidation + of Ir(ppy)3 (blue) to generate [Ir(ppy)3] (red). (Bottom) Absorption spectra of synthetically isolated cobalt complexes...... 30

Figure 2.9. Time-resolved photoluminescence quenching of [Ir(ppy)3]* decay in the presence of increasing concentration of CoII-dppe and Stern-Volmer quenching plot (inset)...... 31

Figure 2.10 Plot of second-order decay of transient signal at 586 nm after III excitation of Ir(ppy)3 at 445 nm in the presence of 0.057 mM of Co - dppe-NCCH3. Using the slope of 1/ΔOD vs time in the above region and the coefficients of the absorbing species, the kexc rate constant was determined to be 2.3 × 109 M−1s−1. The open circles are simulated data 9 −1 −1 using MatLab but using a kexc = 2.9 × 10 M s ...... 32

Figure 2.11 Plot of second-order decay of transient signal at 486 nm after II excitation of Ir(ppy)3 at 445 nm in the presence of 0.98 mM of Co -dppe. Using the slope of 1/ΔOD vs time in the above region and the coefficients of the absorbing species, the kcom rate constant was determined to be 1.1 × 109 M−1s−1. The open circles are simulated data using MatLab, which show excellent agreement with the experimental data...... 34

Figure 2.12.Transient absorption data at 486 nm after exciting 150 µM Ir(ppy)3 at 445 nm in the presence of 1.0 mM CoII-dppe in the presence of increasing concentration of 2,3,5,6-tetrafluoro-4-trifluoromethylphenol (concentration of acid in each trace shown in legend). Single exponential fits are also shown (black lines), and the inset shows resulting kobs from fits vs acid concentration...... 35

Figure 2.13. Brønsted plot of Equation 6 and structures of acids used. Linear regression gives a Brønsted slope of α = 0.4...... 36

Figure 2.14 Overlaid Brønsted plots comparing electrochemically determined kPT I rate constants and photochemically determined kPT rate constants for Co - dppe...... 37

Figure 2.15 (Left): UV/vis absorbance titration of CoII-dppe with up to 9.5 equivalents of 4-methyl-anilinium tetrafluoroborate in CH3CN (pKa = 11.4), and (right) UV/vis absorbance titration of CoII-dppe with up to 57.5 equivalents of 4-methoxy-anilinium tetrafluoroborate in CH3CN (pKa = 11.86)...... 38

Figure 2.16 Cyclic voltammograms comparing the orange product of the reaction II between Co -dppe and HBF4 etherate described in the main text (top), and independently prepared samples of HCoIII-dppe (middle) and CoIII- dppe-NCCH3 (bottom)...... 39

1 II Figure 2.17. H NMR spectrum in CD3CN of the reaction between Co -dppe and an excess of strong acid (4-(methylbenzoate)anilinium, pKa = 8.62) (top) compared with 1H NMR spectra of independently synthesized HCoIII- III dppe (middle) and Co -dppe-NCCH3 (bottom). Symbols * indicate resonances corresponding to the 4-(methylbenzoate)anilinium and its conjugate base, and # indicates resonances corresponding to solvents (CD2HCN and diethyl ether)...... 40

III III Figure 2.18 Cyclic voltammograms of Co -dppe-NCCH3 and Co -dppe- salicylate in CH3CN...... 42

III Figure 2.19 Titration of Co -dppe-NCCH3 with increasing equivalents of [TEA][F4-PhO] in CH3CN...... 43

Figure 3.1 Square scheme showing stepwise and concerted pathways to form a metal hydride complex, a key intermediate in catalytic hydrogen evolution ...... 48

Figure 3.2. Structure of the complexes of the type [CoCp(dxpx)(NCCH3)][PF6]2 and the structural permutations studied herein (dxpx is an abbreviation encompassing depe, dppe, dcpe, and dppv) ...... 51

Figure 3.3 Structural representation of HCoIII-depe (left), HCoIII-dppe (middle), and HCoIII-dcpe (right) viewed looking down on the top of the Cp ring with the hydride ligand pointing down to highlight the differing angle of access to the metal center or hydride ligand. For HCoIII-depe and HCoIII- dcpe, thermal ellipsoids are shown at the 50% probability level; for HCoIII-dppe the quality of the data was such that only a connectivity structure could be obtained at this time...... 52

Figure 3.4. Cyclic voltammograms of isolated CoIII complexes in 0.25 M TBAPF6 in CH3CN. All voltammograms are referenced to the internal standard, the Fc+/Fc redox couple...... 53

III Figure 3.5 Cyclic voltammograms of Co -dcpe-NCCH3 over time ...... 54

Figure 3.6 (Left) Spectrophotometric titration in 5% toluene in CH3CN starting with 1.17 mM HCoIII-depe (blue trace ─) and titrating in increasing

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amounts of DBU (gray traces). UV/vis absorbance spectrum of 1.17 mM CoI-depe is shown to indicate end point (red trace ─). The absorbance at 500 nm and the extinction coefficients of HCoIII-depe and CoI-depe were used to calculate concentration of each species as a function of added base. (Right) Linear regression of the equation for the acid base equilibrium between the species and the known pKa of DBU was used to III determine the pKa of HCo -depe as 23.6...... 56

Figure 3.7 1H NMR spectrum of mixture of HCoIII-dppe, HCoIII-dppv, CoI- I dppe, and Co -dppv in 20% C6D6 in CD3CN for solubility. The integrations of the Cp resonances for each species was used to calculate the equilibrium constant for the indicated reaction, and Keq and the known III III pKa of HCo -dppe was used to calculate the pKa of HCo -dppv to be 18.2...... 57

1 III Figure 3.8 Hydride region of the H NMR spectra in CD3CN of the HCo species studied herein ...... 57

Figure 3.9 Overlaid Brønsted plots for the protonation of CoI-dppe and CoI- dppv ...... 58

Figure 3.10 (Left) Overlaid Brønsted plots for the protonation of CoI-depe, CoI- dppe, and CoI-dcpe using non-bulky acids, and (right) overlaid Brønsted plots for the protonation of CoI-depe, CoI-dppe, and CoI-dcpe highlighting sterically encumbered acids...... 59

Figure 3.11 Cp region of the 2D EXSY 1H NMR spectrum of HCoIII-depe and I Co -depe used to calculate the kself rate constant ...... 61

Figure 3.12 (Left) LFER region of Brønsted plot for CoI-depe, CoI-dppe, and CoI-dcpe compared to (right) simulated kPT values using Marcus cross relation for CoI-depe, CoI-dppe, CoI-dcpe...... 63

Figure 4.1 Simplified orbital energy diagram illustrating the strategy of red- shifting absorption by A) stabilizing * LUMO levels with electron withdrawing groups or B) by destabilizing HOMO levels with modification of axial ligands...... 71

Figure 4.2 1H NMR spectrum of mixture of [Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] and [Re(deeb)(CO)2(PMe3)(NMe3)][PF6] taken after allowing [Re(deeb)(CO)3(PMe3)][PF6] and TMNO to react at room temperature...... 73

1 Figure 4.3 Aromatic region of H NMR spectra of [Re(deeb)(CO)3(PMe3)][PF6] (bottom), [Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] (middle), and reaction products of [Re(deeb)(CO)3(PMe3)][PF6] and TMNO reaction after 12 minutes at react at room temperature...... 73

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1 Figure 4.4 H NMR spectrum of 1 in CD3CN...... 75

13 Figure 4.5 C NMR spectrum of 1 in CD3CN...... 76

1 Figure 4.6 Region of H NMR of [Re(deeb)(CO)3(NCCH3)][PF6] in CD3CN taken immediately after dissolution (top) and then after 7 days (bottom), showing that the resonance corresponding to the bound CH3CN ligand disappears and 1 equivalent of free CH3CN grows in...... 76

1 Figure 4.7 H NMR spectrum of 2 in CD3CN...... 77

Figure 4.8 Carbonyl stretching region of IR spectra of complexes 1-4, collected in CH2Cl2...... 78

1 Figure 4.9 H NMR spectrum of 3 in CD3CN...... 79

13 1 Figure 4.10 C{ H} NMR spectrum of 3 in CD3CN...... 79

1 Figure 4.11 H NMR spectrum of 4 in CD3CN...... 80

Figure 4.12 Structural representation of 4 with thermal ellipsoids shown at the 50% probability level. Hydrogen atoms and a slightly disordered dichloromethane solvent molecule are omitted for clarity. Selected distances (Å): Re1-Cl1 2.4918(10), Re1-P1 2.3386(11), Re1-N1 2.180(3), Re1-N2 2.161(3), Re1-C1 1.903(4), Re1-C2 1.884(5)...... 81

Figure 4.13 Absorption spectra of complexes 1-4 (solid lines) in CH3CN at 295 K and emission spectra of complexes 1-3 (dashed lines) in 90/10 2- MeTHF/CH3CN) at 77 K. For emission spectra, exc = 365 nm for complexes 1 and 2 and exc = 455 nm for complex 3...... 82

Figure 4.14 Emission spectrum of 1 at RT in CH3CN (exc = 405 nm), and in 90:10 2-MeTHF:CH3CN at RT and at 77 K (exc = 365 nm) ...... 83

Figure 4.15 Emission spectrum of 2 at RT in CH3CN, and in 90:10 2- MeTHF:CH3CN at RT and at 77 K (exc = 365 nm) ...... 83

Figure 4.16 Time-resolved photoluminescence measurements of 1 at 295 K (left) and at 77 K (right). Measurements taken in 90/10 2-MeTHF/CH3CN...... 84

Figure 4.17 Time-resolved photoluminescence measurements of 2 at 295 K (left) and at 77 K (right). Measurements taken in 90/10 2-MeTHF/CH3CN...... 84

Figure 4.18 Time-resolved photoluminescence measurements of 3 at 77 K. Measurement taken in 90/10 2-MeTHF/CH3CN...... 85

Figure 4.19 Cyclic voltammograms of 1 mM of 1-4 in 0.25 M TBAPF6 in CH3CN. For complexes 2-4, 1 mM Fc was added and the internal

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reference Fc+/Fc couple (0 V) is shown with a dashed box. Due to an interaction with oxidized 1 and Fc+, ferrocene was not added to the solution when scanning oxidatively; the reduction feature was referenced to Fc+/Fc in a separate experiment. All scans were recorded using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode, and 3 V/s scan rate...... 86

Figure 4.20 Calculated excitations (black vertical lines) and experimental UV/Vis absorption spectra (colored lines) for complexes 1-4 (A-D, respectively) along with pictures of the calculated HOMO-1 and LUMO orbitals involved in the most intense low energy transition for each complex...... 88

Figure 4.21 Scan rate dependence of the first reductive feature of 1 and 2. Cyclic voltammograms were recorded in 0.25 M TBAPF6 in CH3CN using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode...... 90

Figure 4.22 Scan rate dependence of the first reductive feature of 3 and 4. Cyclic voltammograms were recorded in 0.25 M TBAPF6 in CH3CN using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode...... 90

Figure 4.23 Scan rate dependence of the oxidative feature of 1 and 2. Cyclic voltammograms were recorded in 0.25 M TBAPF6 in CH3CN using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode...... 91

Figure 4.24 Catalytic electrooxidation of Fc+ by complex 1, and cyclic voltammograms of 1 and Fc alone. Cyclic voltammograms were recorded in 0.25 M TBAPF6 in CH3CN using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode...... 92

Figure 4.25 Scan rate dependence of the oxidative feature of 3 and 4. Cyclic voltammograms were recorded in 0.25 M TBAPF6 in CH3CN using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode...... 93

Figure 4.26 Emission spectrum of 1 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 365 nm) and tangential fit to the high energy side to estimate GST...... 94

Figure 4.27 Emission spectrum of 2 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 365 nm) and tangential fit to the high energy side to estimate GST...... 94

Figure 4.28 Emission spectrum of 3 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 455 nm) and tangential fit to the high energy side to estimate GST...... 95

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Figure 4.29 Linear correlation between lowest energy max experimentally determined GST values of 1-3 used to estimate GST of 4...... 95

Figure 4.30 (Left) Energy diagram calculated using TD-DFT highlighting the relevant visible transitions and (right) calculated molecular orbitals for the HOMO, HOMO-1, and HOMO-2 orbitals for complexes 2 and 4 showing the degree of backbonding to the carbonyl ligands in each...... 98

Figure 5.1 Comparison between the prior use of photo-reductants and photoacids to generate TMHCs and our proposed reactivity based on photoactive transition metal complexes...... 106

31 1 Figure 5.2 P{ H} NMR spectra before and after heating ReCl3(PPh3)2(NCCH3) and dmpe in toluene ...... 110

1 Figure 5.3 H NMR of ReCl3(dmpe)(PPh3) in CDCl3 ...... 111

1 Figure 5.4 H NMR spectrum of OReCl3(depe) in d6-DMSO ...... 112

31 1 Figure 5.5 P{ H} NMR spectrum of OReCl3(depe) in d6-DMSO ...... 113

1 Figure 5.6 Labeled H NMR spectrum of Re(depe)H7 and minor impurity (depe)H2Re(µ-H4)ReH2(depe) in C6D6 ...... 113

31 1 Figure 5.7 P{ H} NMR spectrum of Re(depe)H7 and minor impurity (depe)H2Re(µ-H4)ReH2(depe) in C6D6 ...... 114

Figure 5.8 1H NMR spectrum of J-Young tube scale reaction between Re(depe)H7 and 1 equivalent of CpH in C6D6 heated to 90° C overnight (~16 hours)...... 115

Figure 5.9 J-Young tube scale reaction between Re(depe)H7 and 30 equivalents of CpH in C6D5Cl prior to heating, and heated to 135 °C for 45 minutes, 90 minutes, and ~16 hours...... 115

Figure 5.10 J-Young tube scale reaction between Re(depe)H7 and 30 equivalents of CpH in d8-toluene prior to heating, and heated to 135 °C for 30 minutes and ~16 hours...... 116

31 1 Figure A.1 P{ H} NMR spectrum of [Re(deeb)(CO)3(NCCH3)][PF6] (1) in CD3CN ...... 129

1 Figure A.2 H NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN...... 130

13 1 Figure A.3 C{ H} NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN...... 130

31 1 Figure A.4 P{ H} NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN...... 131

31 1 Figure A.5 P{ H} NMR spectrum of 3 in CD3CN...... 131

13 1 Figure A.6 C{ H} NMR spectrum of 4 in CD3CN...... 132

31 1 Figure A.7 P{ H} NMR spectrum of 4 in CD3CN...... 132

Figure A.8 Absorption and emission spectra of 1 and [Ru(bpy)3][PF6]2 standard in deoxygenated CH3CN used for quantum yield determination...... 133

Figure B.1 Determination of diffusion coefficients for [CoCp(dppe)(NCCH3)][PF6]2 ...... 134

Figure B.2 Determination of diffusion coefficients for [CoCp(depe)(NCCH3)][PF6]2 ...... 135

Figure B.3 Determination of diffusion coefficients for [CoCp(dcpe)(NCCH3)][PF6]2 ...... 136

Figure B.4 Determination of diffusion coefficients for [CoCp(dppv)(NCCH3)][PF6]2 ...... 137

Figure B.5 Determination of heterogeneous electron transfer rate constants for [CoCp(dppe)(NCCH3)][PF6]2 ...... 138

Figure B.6 Determination of heterogeneous electron transfer rate constants for [CoCp(depe)(NCCH3)][PF6]2 ...... 138

Figure B.7 Determination of heterogeneous electron transfer rate constant for [CoCp(dcpe)(NCCH3)][PF6]2...... 139

Figure B.8 Determination of heterogeneous electron transfer rate constants for [CoCp(dppv)(NCCH3)][PF6]2 ...... 139

Figure B.9 Electrochemical determination of kPT for CoCp(dppe) with dimethylformamidium triflate ...... 143

Figure B.10 Electrochemical determination of kPT for CoCp(dppe) with 4-cyano- anilinium tetrafluoroborate ...... 144

Figure B.11 Electrochemical determination of kPT for CoCp(dppe) with 4-CF3- anilinium tetrafluoroborate ...... 145

Figure B.12 Electrochemical determination of kPT for CoCp(dppe) with p- toluenesulfonic acid monohydrate ...... 146

Figure B.13 Electrochemical determination of kPT for CoCp(dppe) with 4-Br- anilinium tetrafluoroborate ...... 147

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Figure B.14 Electrochemical determination of kPT for CoCp(dppe) with 4-Cl- anilinium tetrafluoroborate ...... 148

t Figure B.15 Electrochemical determination of kPT for CoCp(dppe) with 4- Bu- anilinium tetrafluoroborate ...... 149

Figure B.16 Electrochemical determination of kPT for CoCp(dppe) with 4- methylanilinium tetrafluoroborate (method 1)...... 150

Figure B.17 Electrochemical determination of kPT for CoCp(dppe) with 4- methylanilinium tetrafluoroborate (method 2)...... 151

Figure B.18 Electrochemical determination of kPT for CoCp(dppe) with N,N- dimethylanilinium tetrafluoroborate ...... 152

Figure B.19 Electrochemical determination of kPT for CoCp(dppe) with 4-MeO- anilinium tetrafluoroborate ...... 153

Figure B.20 Electrochemical determination of kPT for CoCp(dppe) with pyridinium tetrafluoroborate ...... 154

Figure B.21 Electrochemical determination of kPT for CoCp(dppe) with 2,6- lutidinium tetrafluoroborate ...... 155

Figure B.22 Electrochemical determination of kPT for CoCp(dppe) with 4-MeO- pyridinium tetrafluoroborate ...... 156

Figure B.23 Electrochemical determination of kPT for CoCp(dppe) with 2,4,6- trimethylpyridinium tetrafluoroborate ...... 157

Figure B.24 Electrochemical determination of kPT for CoCp(dppe) with 4-CF3- 2,3,5,6-F4-Phenol ...... 158

Figure B.25 Electrochemical determination of kPT for CoCp(dppe) with salicylic acid ...... 159

Figure B.26 Electrochemical determination of kPT for CoCp(dppe) with pentafluorophenol ...... 160

Figure B.27 Electrochemical determination of kPT for CoCp(dppe) with 2,3,5,6- tetrafluorophenol ...... 161

Figure B.28 Electrochemical determination of kPT for CoCp(dppe) with 2,4,6- tribromophenol ...... 162

Figure B.29 Electrochemical determination of kPT for CoCp(dppe) with benzoic acid ...... 163

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Figure B.30 Electrochemical determination of kPT for CoCp(dppe) with glacial acetic acid...... 164

Figure B.31 Electrochemical determination of kPT for CoCp(dppe) with 4-chlorophenol ...... 165

Figure B.32 Photochemical determination of kPT of CoCp(dppe) using benzoic acid ...... 166

Figure B.33 Photochemical determination of kPT for CoCp(dppe) using 2,3,5,6- tetrafluorophenol ...... 167

Figure B.34 Photochemical determination of kPT of CoCp(dppe) using triethylammonium tetrafluoroborate ...... 168

Figure B.35 Photochemical determination of kPT for CoCp(dppe) using 2,3,4,5,6- pentachlorophenol ...... 169

Figure B.36 Photochemical determination of kPT for CoCp(dppe) using salicylic acid ...... 170

Figure B.37 Photochemical determination of kPT for CoCp(dppe) using 2,3,5,6- F4-4-CF3-phenol ...... 171

Figure B.38 Electrochemical determination of kPT of CoCp(depe) with phenol ...... 172

Figure B.39 Electrochemical determination of kPT of CoCp(depe) with 4-CH3- phenol ...... 173

Figure B.40 Electrochemical determination of kPT of CoCp(depe) with 4-CF3- phenol ...... 174

Figure B.41 Electrochemical determination of kPT of CoCp(depe) with 4-Cl- phenol ...... 175

Figure B.42 Electrochemical determination of kPT of CoCp(depe) with acetic acid ...... 176

Figure B.43 Electrochemical determination of kPT of CoCp(depe) with benzoic acid ...... 177

Figure B.44 Electrochemical determination of kPT of CoCp(depe) with 2,3,5,6- tetrafluorophenol ...... 178

Figure B.45 Electrochemical determination of kPT of CoCp(depe) with benzylammonium tetrafluoroborate ...... 179

Figure B.46 Electrochemical determination of kPT of CoCp(depe) with salicylic acid ...... 180

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Figure B.47 Electrochemical determination of kPT of CoCp(depe) with 4-CH3O- pyridinium tetrafluoroborate ...... 181

Figure B.48 Electrochemical determination of kPT of CoCp(depe) with pyridinium tetrafluoroborate ...... 182

Figure B.49 Electrochemical determination of kPT of CoCp(depe) with 4-CH3- anilinium tetrafluoroborate ...... 183

t Figure B.50 Electrochemical determination of kPT of CoCp(depe) with 4- Bu- anilinium tetrafluoroborate ...... 184

Figure B.51 Electrochemical determination of kPT of CoCp(depe) with 4-Cl- anilinium tetrafluoroborate ...... 185

Figure B.52 Electrochemical determination of kPT of CoCp(depe) with 4-CF3O- anilinium tetrafluoroborate ...... 186

Figure B.53 Electrochemical determination of kPT of CoCp(depe) with 4-CN- anilinium tetrafluoroborate ...... 187

Figure B.54 Electrochemical determination of kPT of CoCp(dcpe) with 4-Cl- phenol ...... 188

Figure B.55 Electrochemical determination of kPT of CoCp(dcpe) with acetic acid ...... 189

Figure B.56 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,5,6- tetrafluorophenol ...... 190

Figure B.57 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,4,5,6- pentafluorophenol ...... 191

Figure B.58 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,5,6- tetrafluoro-4-CF3-phenol...... 192

Figure B.59 Electrochemical determination of kPT of CoCp(dcpe) with 4-CH3O- pyridinium tetrafluoroborate ...... 193

Figure B.60 Electrochemical determination of kPT of CoCp(dcpe) with pyridinium tetrafluoroborate ...... 194

t Figure B.61 Electrochemical determination of kPT of CoCp(dcpe) with 4- Bu- pyridinium tetrafluoroborate ...... 195

Figure B.62 Electrochemical determination of kPT of CoCp(dcpe) with anilinium tetrafluoroborate ...... 196

Figure B.63 Electrochemical determination of kPT of CoCp(dcpe) with 4-Br- anilinium tetrafluoroborate ...... 197

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Figure B.64 Electrochemical determination of kPT of CoCp(dppv) with acetic acid ...... 198

Figure B.65 Electrochemical determination of kPT of CoCp(dppv) with benzoic acid ...... 199

Figure B.66 Electrochemical determination of kPT of CoCp(dppv) with salicylic acid ...... 200

t Figure B.67 Electrochemical determination of kPT of CoCp(dppv) with 4- Bu- anilinium tetrafluoroborate ...... 201

Figure B.68 Electrochemical determination of kPT of CoCp(dppv) with 4-Br- anilinium tetrafluoroborate ...... 202

Figure B.69 Electrochemical determination of kPT of CoCp(dppv) with 4-CN- anilinium tetrafluoroborate ...... 203

1 Figure C.1 Concentration of H2 (▲) and HD (◆) detected by H NMR spectroscopy during the photolysis of 11.5 mM of [HIrIII]+ and 100 mM of CD3CO2D in CD3CN with 460 nm light. Inset shows the first 1 min. H2 concentration corrected for thermal population of para-H2...... 206

Figure C.2 Room temperature absorbance and emission spectrum of [HIrIII]+ in CH3CN. For emission spectrum, λexc = 428 nm...... 207

Figure C.3 Transient absorption spectrum taken at various delays (picoseconds to nanosecond) after excitation of [HIrIII]+...... 207

Figure C.4 (Top) Time-resolved photoluminescence decay of excited state of [HIrIII]+, and (middle and bottom) time-resolved transient absorption kinetic profiles at various wavelengths showing the decay of the excited- III + state transients of the excited state of [HIr ] . For all experiments, λexc = 445 nm...... 209

Figure C.5 (Left) Time-resolved photoluminescence decay at different III + concentrations of [HIr ] , and (right) Stern-Volmer plot of kobs vs III + [HIr ] concentration to determine kq and k0 ...... 209

Figure C.6 Overlay of TRPL decay of the excited state of [HIrIII]+ at 708 nm and the TA kinetics of growth of transient signal assigned to IrI probed at 625 nm...... 211

Figure C.7 (Left) Kinetic traces at various wavelengths of [HIrIII]+, dashed line is 10 µs, and (right) absorbance values at 10 µs overlaid with (IrI + [IrIII- 2+ III + NCCH3] ) – [HIr ] difference spectrum ...... 212

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Figure C.8 (Left) Acid dependence of the decay of transient signal assigned to IrI (λprobe = 650 nm) and plot of kobs vs [HOAc] showing second order dependence on acid...... 212

Figure C.9 Summary of mechanism of catalytic H2 evolution upon excitation of [HIrIII]+ in the presence of acid and STAB ...... 213

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LIST OF TABLES

Table 2.1. Thermochemical data for iridium and cobalt species...... 29

Table 2.2 Summary of photochemically determined kPT values ...... 36

Table 3.1 Summary of electrochemical properties of cobalt complexes...... 54

Table 4.1 Experimental and calculated vibrational frequencies for complexes 1-4...... 87

Table 4.2 Photophysical properties of complexes 1-4...... 95

Table 4.3 Summary of ground state and excited state redox properties of complexes 1-4...... 96

Table A.1 Radiative and non-radiative rate constants for complexes 1 and 2, and photophysical parameters used in calculations...... 133

Table B.1 Summary of electrochemically determined protonation rate constants of CoCp(dppe) utilizing acids with varying pKas...... 140

Table B.2 Summary of photochemically determined protonation rate constants of CoCp(dppe) utilizing acids with varying pKas...... 141

Table B.3 Summary of electrochemically determined protonation rate constants of CoCp(depe) utilizing acids with varying pKas...... 141

Table B.4 Summary of electrochemically determined protonation rate constants of CoCp(dcpe) utilizing acids with varying pKas...... 142

Table B.5 Summary of electrochemically determined protonation rate constants of CoCp(dppv) utilizing acids with varying pKas...... 142

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LIST OF SCHEMES

Scheme 2.1 PCET reaction showing the addition of one proton and one electron to [CoIICp(PP)]+ (where PP = chelating bisphosphine ligand) to form the analogous [HCoIIICp(PP)]+ complex ...... 20

Scheme 2.2 Synthetic summary of all relevant Co-dppe species with photographs of each isolated sample. The postulated HCoIV complex has not been experimentally observed, and is therefore grayed out ...... 22

III Scheme 2.3 EC’ mechanism of reduction of Co -dppe-NCCH3 followed by II CH3CN ligand dissociation to form Co -dppe...... 23

Scheme 2.4. Summary of kinetic processes discussed herein...... 33

Scheme 2.5 Ground state reactivity between CoII-dppe and acid as well as CoIII ligation ...... 40

Scheme 4.1 Synthesis of complexes 1-4. (i) 3 equiv AgOTf in CH3CN at 295 K, [NH4][PF6] metathesis (ii) 3 equiv AgOTf in refluxing acetone, 30 equiv PMe3 in refluxing acetone, [NH4][PF6] metathesis (iii) 1 equiv TMNO in refluxing CH3CN (iv) 15 equiv. [BnMe3N][Cl] in refluxing CH2Cl2...... 74

Scheme 5.1 LMCT excited states increase the basicity of the metal center and generate a strongly oxidizing ligand cation...... 107

Scheme 5.2 Stepwise pathways of PCET that begin from the LMCT excited state...... 108

Scheme 5.3 Synthesis of ReCl3(dmpe)(PPh3) from ReCl3(PPh3)2(NCCH3) and dmpe...... 109

Scheme 5.4 Synthesis of HReCp(dmpe)Cl from ReCl3(dmpe)(PPh3) ...... 111

Scheme 5.5 Synthesis of OReCl3(depe) from ammonium perrhenate and depe ...... 112

Scheme 5.6 Reduction of OReCl3(depe) with LiAlH4 to form Re(depe)H7 ...... 113

Scheme C.1 Photoelectrochemical hydrogen evolution using iridium Cp* bipyridine complexes ...... 205

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LIST OF ABBREVIATIONS AND SYMBOLS

α Brønsted coefficient

µA Microamp

°C Degrees Celsius cm Centimeter

CPET Concerted proton/electron transfer

CV Cyclic voltammogram

2D Two-dimensional

D Diffusion coefficient

δ Change or difference (NMR chemical shift)

Δ Change or difference

ε Molar extinction coefficient

E0* Excited-state reduction potential

E1/2 Reduction potential

EC Electrochemical event followed by chemical step

Ep Peak potential equiv. Equivalents eV Electron volts

EXSY EXchange SpectroscopY

F Faraday’s constant

ΔG Gibbs free energy

ΔG‡ Reaction barrier energy mg Milligrams h Hour

HOMO Highest occupied molecular orbital

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Hz Hertz

IR Infrared

K Kelvin kcal Kilocalorie kx Rate constant for process x

Kx Equilibrium constant for equation x

µL Microliter

LFER Linear free energy relationship

LUMO Lowest unoccupied molecular orbital mL Milliliter

λ Wavelength

M Molar

MLCT Metal-to-ligand charge transfer mM Millimolar mol Mole nm Nanometer

NMR Nuclear magnetic resonance ns Nanosecond

OCP Open circuit potential

OD Optical density

PCET Proton-coupled electron transfer pKa Logarithm of acid dissociation constant

PL Photoluminescence

PP Chelating bisphosphine ligand ppm Parts per million

xxx

PT Proton transfer

π* Pi antibonding

R Gas constant s Second ms Millisecond

µ Microsecond

T Temperature

T1 Longitudinal relaxation constant

TD-DFT Time-dependent density functional theory

TMHC Transition metal hydride complex

TOF Turnover frequency

TON Turnover number

τ Lifetime

 Angle

υ Scan rate

UV/vis Ultraviolet/visible

V Volt mV Millivolt

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CHAPTER 1: KINETIC ACIDITY OF TRANSITION METAL HYDRIDE COMPLEXES

1.1 Introduction and Background

Organometallic transition metal hydride complexes (TMHC) trace back the mid-1950s

1,2 with the first report and characterization of Cp2ReH by Wilkinson et al. Once thought to be

rare and elusive structures, identified initially in part by their characteristic upfield shifted

hydride resonance in 1H NMR spectra,3 an impressive amount of literature has been published

regarding the synthesis, characterization, and reactivity of TMHCs. Not only are TMHCs of

fundamental interest due to their unique structure and bonding, but they are common

intermediates in industrially relevant reactions such as alkene hydrogenation and

4 hydroformylation (both processes catalyzed by the well-known HRh(CO)(PPh3)3). Many

aspects of TMHCs have been well reviewed, including their relevance in homogeneous

catalysis5, acidic6 and hydridic7 thermodynamics, and non-traditional hydrides such as

polyhydride clusters.8,9 This extensive literature reflects the diverse ways that TMHCs can

react.10 Due to the extent to which the metal-hydrogen bond can be polarized, TMHCs can act as

hydride donors (two electrons and a proton), net H• donors (one electron and one proton), or

proton donors. 11 This chapter will focus on the latter reaction, either forming or breaking a

transition metal hydride bond through the addition or removal of a proton, i.e. acid/base

chemistry of TMHCs.

The thermodynamic Brønsted-Lowry acidity of TMHCs is a well-studied field in the

literature. Reported pKa values span >60 kcal/mol, and can be estimated by metrics like the

ligand acidity constant (LAC) method developed by Morris.6 These thermodynamic values are

experimentally usually determined by establishing a detectable equilibrium between a TMHC and its conjugate base with some organic acid of known pKa. Detection methods include UV/vis

12 13 14 absorption spectroscopy , IR spectroscopy , or NMR spectroscopy. Alternatively, pKa values can be determined indirectly through thermodynamic square schemes.

Just as important as the thermodynamic acidity of TMHCs is the kinetic acidity. It is worth noting that equilibrium conditions required to determine accurate thermodynamic parameters can take hours, and sometimes days to establish.15,16 Early on in the study of metal- hydrides, observations were made on how sluggish proton transfer kinetics were, even for thermoneutral or downhill proton transfers. Once a point of fundamental intrigue, implications of proton transfer kinetics involving TMHCs have extended into the field of solar fuel generation.

TMHCs are common intermediates in fuel-forming reactions such as catalytic proton reduction to form hydrogen and carbon dioxide reduction to formate.17–19 Thus, the study of kinetics of fundamental steps involving TMHCs is envisioned to improve our understanding of these fuel- forming reactions and guide future catalyst design. Therefore, this chapter intends to bridge the gap between the early literature on kinetic acidity reports of transition metal carbonyl complexes and the current understanding and efforts in the field of fuel generation catalyzed by transition metal complexes via TMHC intermediates.

1.2 Kinetic Acidity

Kinetic acidity of transition metal hydrides encompasses the formal reaction shown in equation 1; the rate constant describing a proton being removed from a transition metal-hydride complex to form its conjugate base (k1), or the rate constant describing the addition of a proton to a transition metal complex to form its conjugate acid (k-1).

There are a handful of reports very early in the study of metal hydride complexes that briefly discuss the kinetic acidity reactivity described above; however most of these reports were in a wide variety of solvents and conditions20–28, and the thermodynamic driving force for proton transfer was not discussed in the context of kinetic acidity. Norton’s 1982 report of deprotonation rate constants for a series of structurally similar transition metal carbonyl hydride complexes (M = Cr, Mo, and W) and substitutionally-different osmium carbonyl hydrides (cis-

Os(CO)4(H)2, cis-Os(CO)4(CH3)H, and Os(CO)4H) is the first to directly discuss the concept of kinetic acidity.29,30 This study, and that of Kresge, Pearson and coworkers a few years earlier25, made the interesting observation that rate constant magnitudes for the deprotonation of TMHCs more closely match the rate constants for deprotonation of carbon-based acids like (4- nitrophenyl)nitromethane and nitroparrafins and less closely matched deprotonation rates of nitrogen- or oxygen-based acids with comparable thermodynamic driving forces.

In 1987, Norton and coworkers reported the rate constants for the deprotonation of a series of different TMHCs (Cp(CO)3WH, (CO)5MnH, (CO)4FeH2, and (CO)4CoH) by a series of para-substituted aniline bases with similar steric profiles.31 Plots of the deprotonation rate

+ constant kH+ vs. driving force for the proton transfer reaction (pKeq = pKa(MH)–pKa(BH )) revealed a linear free energy relationships (LFERs) with Brønsted slopes that varied from 0.48 to

0.65 across the four hydride complexes studied (Figure 1.1). Separately, the kinetics for deprotonation of 14 different metal hydrides by the same base (aniline) were measured, and a rough linear correlation between kH+ vs -pKeq. Collectively, these data supported the conclusion that the LFER can be used as a “diagnostic kinetic acidity series”.

Figure 1.1 Linear-free energy relationships for a series of TMHC complex deprotonation by substituted anilines. ((CO)4CoH α = 0.48, (CO)5FeH2 α = 0.55, (CO)5MnH α = 0.54, and 31 Cp(CO)3WH α = 0.65).

Even more sluggish rate constants have been observed for the protonation of transition metal clusters (or the deprotonation of transition metal hydride clusters) that contain bridging hydrides that are relatively more acidic than monometallic TMHC analogs. For example, Norton

2- et al reported a series of deprotonation rate constants of [Rh13(CO)24H3] by various substituted anilines, which included the extraordinarily slow rate constant of 1.42 × 10-3 M-1 s-1 for p- toluidine.13 This is surprising because the deprotonation with p-toluidine is actually a slightly exergonic reaction.

In addition to slowing proton transfer kinetics by trapping the hydride ligand between multiple metal centers as mentioned above, there have been multiple examples of sterically encumbered TMHCs that exhibit rather slow deprotonation rates. In one example, 1H NMR spectra of combinations of Cr(tripod)2 (tripod = MeC(CH2PMe2)3) and its conjugate acid

+ 32 HCr(tripod)2 showed no line broadening as compared to their separate NMR spectra. Noting that intramolecular migration of the hydride ligand is facile on the NMR timescale, the authors

suggest that the steric bulk on the phosphine ligands is the cause for not observing intermolecular

15 proton exchange. In another example, Darensbourg et al showed that HMo(CO)2(PP)2 (PP = chelating bisphosphine ligand such as dppe = 1,2-bis(diphenylphosphino)ethane, dmpe = 1,2- bis(dimethylphosphino)ethane, and depe = 1,2-bis(diethylphosphino)ethane) complexes were deprotonated with rate constants far less than would predicted by thermodynamic arguments.

Interestingly, the rate of deprotonation of the hydride complexes by a number of amine bases was greatly enhanced when small anionic catalysts were added to solution (I–, Br–, Cl–, F–, acetate). The authors proposed a mechanism in which the steric ligand profile (two bulky bisphosphine ligands) prevented facile approach of the amine base, but allowed for small anionic

“proton shuttles” to approach the hydride ligand more easily, and catalytically transfer the proton to the more thermodynamically stable amine base. The authors also found a complex relationship between deprotonation rate constant and phosphine substituent that likely arises from a combination of steric and electronic effects.33 Another report from Norton et al described the

+ 34 thermodynamic and kinetic acidity of the polyhydride complex [H4Re(PMe2Ph)4] . While also being rather basic compared to other carbonyl hydride complexes studied previously, this complex is deprotonated more than 5 orders of magnitude slower than predict based on thermodynamic arguments alone. In this case, and others with sterically encumbered ligand sets, there is relatively little structural rearrangement between the TMHC and its conjugate base, and therefore approach by the base being limited by the flanking ligand substituents is the main reason for the dramatically lower rate constants.

Kristjansdottir and Norton reported one of the first examples of proton transfer cross exchange between two different transition metal complexes in 1991.35 Importantly, they not only were able to measure the protonation of a reduced transition metal complex using another

TMHC, but they were also able to show that a Marcus-type cross relation (similar to that used for electron transfer) was valid for proton transfers involving TMHCs. One main reason that

TMHCs adhere to this relationship is that there is little to no ion pairing of the conjugate bases or hydrogen bonding of the TMHCs in CH3CN.

+ 36 TMHCs have been shown to be quenchers of photobases, namely ReO2(py)4 *. The

+ excited state of ReO2(py)4 has been extensively studied, and is quenched via excited state proton transfer by a number of protic solvents and molecular organic acids.37 These acids, however, hydrogen bond with the oxo ligand of the complex, thus complicating diffusional quenching studies. Complexes of the type CpM(CO)3H (M = Cr, Mo, W) were not only shown to

+ be efficient quenchers of ReO2(py)4 *, but due to the fact that TMHC complexes do not engage in significant hydrogen bonding, the quenching rate constants and the previously determined proton-self exchange rates (discussed below) were used to demonstrate an example of Marcus theory describing excited-state proton transfer.

1.3 Intrinsic Kinetic Barriers and Self-Exchange

It is apparent that there are limiting kinetic factors in proton transfer reactions involving

TMHCs that lower PT rate constants as compared to those of proton transfer reactions involving common oxygen or nitrogen acids and bases with similar thermodynamic driving forces. When a nitrogen or oxygen based acid is deprotonated, the electron pair in the nitrogen-hydrogen or oxygen-hydrogen bond stays localized on the nitrogen or oxygen, and there is generally very little geometric or electronic rearrangement associated with deprotonation.38 When a transition- metal hydride complex is deprotonated, however, there can be significant geometric changes due to the lowering of the coordination number which can lead to significant electronic structure differences between the protonated and deprotonated form of the complex. These electronic and

geometric changes are distinctly observed in the deprotonation of HCo(CO)4 by NEt3; the ν(CO) values shift by ~100 cm-1 indicating a higher degree of backbonding, and the geometry changes from trigonal bipyramidal to tetrahedral upon deprotonation.39 This leads to an “intrinsic barrier” to deprotonation of a TMHC or protonation to form a TMHC, or an energy barrier that is inherent to the structure of the complex regardless of the acid/base partner.

Figure 1.2 Deprotonation of HCo(CO)4 by NEt3 to form [TEAH][Co(CO)4], and carbonyl stretching frequencies for the TMHC and its conjugate base.

These intrinsic barriers can be probed by examining the self-exchange rate constants for proton transfer between a transition-metal hydride complex and its conjugate base. Self- exchange reactions are commonly used because the free energy for the reaction is, by definition, zero and therefore isolates the intrinsic barriers to proton transfers. Self-exchange proton transfer rates have been studied for a number of TMHC/conjugate base systems.29,31,35,40

As noted previously, there are numerous similarities in the kinetic acidity of carbon- based acids and TMHCs. There are parallels between both experimental observations of rate constants and theory used to support and justify the rate constants. Creutz and coworkers used a

“weak-interaction” model of non-adiabatic transfer to theoretically model the proton self- exchange rates between a number of TMHCs and their conjugate bases.41 This model was adapted from a similar model describing outer-sphere electron transfer reactions42, and outlines a number of parameters that influence the rate constants for proton transfer between a TMHC. This model describes significant classical activation barriers to proton transfer as a function of largely

angular nuclear configuration changes of ancillary ligands, which arise from influencing the probability of the “correct” orientation of the TMHC and its conjugate base in order for proton transfer to occur. Also predicted in this model is that steric influence about the proton source

(either the THMC or the base) has a large influence on proton transfer rate constants, as the

Frank-Condon factor for proton transfer varies dramatically as a function of distance. This has been experimentally observed in a number of cases wherein bulky substituents surround the hydride ligand15, 32–34, and in our own work in the protonation of a reduced metal complex with bulky bases (see Chapters 2 and 3).

Figure 1.3 Potential energy vs q (A-H distance) profiles at Q = Q‡ (A-B distance at the transition 43 state). Comparison between a carbon acid, CH4, and two Eigen acids, NH3 and H2O.

Similar non-adiabatic treatments of carbon based acids and bases have been proposed, namely by Saveant and coworkers.43 They use quantum chemical calculations to calculate carbon-, oxygen-, and nitrogen-hydrogen distances as a function of carbon-carbon, oxygen-

– oxygen, and nitrogen-nitrogen distances in the proton self-exchange reactions of CH4/CH3 ,

– + H2O/OH , and NH4/NH3, respectively (Figure 1.3).

Due to N-H and O-H bonds being more polar than C-H bonds, the distance between the

– + self-exchanging partners for H2O/OH and NH4/NH3 is shorter and thus falls under the adiabatic

– regime whereas the C-C bond distance in the self-exchanging partners for CH4/CH3 is much greater and thereby introduces an intrinsic activation barrier for proton transfer. These findings parallel Creutz’s findings related to metal-hydride self-exchange reactions.

The influence of electronic configuration and geometric rearrangement on rate constants are observed in some electron transfer reactions as well, albeit with bond-distance changes as compared to bond-angle changes. Outer sphere self-exchange electron transfer rate constants are influenced dramatically by geometric and electronic rearrangements. The canonical example of cobalt tris-bipyridine complexes demonstrates this concept well. Electron transfer self-exchange

+ 2+ between [Co(bpy)3] and [Co(bpy)3] is relatively fast due to their similarities in geometry and electronic structure, wherein the electron transfer self-exchange rate constant for the

2+/3+ [Co(bpy)3] reaction is greatly attenuated due to the electronic difference of population of an antibonding orbital which leads to elongation of the Co-N bonds.44

1.4 Circumvention of Intrinsic Barriers – Relevance to Catalysis

As previously discussed, catalytic hydrogen evolution mediated by transition metal catalysts often progresses through the formation of a TMHC, the formation of which can be rate- limiting. Several groups have been able to circumvent these intrinsically slow bimolecular protonation events, however, with great success. The strategy involves the introduction of a nitrogen or oxygen base appended to the ligand in the second coordination sphere, termed a

“pendant base” or “proton-relay”.45 It has proven to be particularly successful in the field of

electrocatalytic hydrogen evolution, wherein there is a series of electrochemical and chemical steps between the electrode, the catalyst, and an acid source in solution that leads to hydrogen generation. In the absence of the pendent base, the first chemical step involving the protonation of reduced metal complex to form a TMHC can be slow due to the aforementioned factors.

However, in the presence of the pendant base, the intermolecular protonation of the nitrogen or oxygen is kinetically more facile as compared to the metal center, and the metal-hydride formation step is then an isomerization.45 Some of the most well-studied hydrogen evolution catalysts are nickel-based bis(P2N2) complexes where P2N2 is a chelating bisphosphine ligand with two pendant secondary amines.46,47

Figure 1.4 Comparison between direct protonation of metal center to form metal hydride and protonation at pendant amine followed by intramolecular rearrangement to form metal hydride.

Both experimental and computation methods have been used to show that the nitrogen base is the kinetically favored protonation site by exogenous acids under catalytic conditions, and that intramolecular proton transfer from the protonated amine to the metal centered is facile, thereby increasing the overall rate of catalysis as compared to similar complexes without the pendant bases (Figure 1.4).47,48 This concept is also the basis of cobalt and nickel “hangman porphyrin” catalysts for hydrogen generation, where there is a carboxylic acid of phosphonic acid moiety situated above the metal center in order to facilitate metal-hydride formation.

This difference between ligand-based protonation and metal-based protonation is also an area of research related to the protonation of TMHCs to form either transition metal hydrogen complexes or transition metal dihydride complexes. This field has been reviewed thoroughly in the literature,49–51 but to the best of our knowledge, all but one example52 of THMC protonation occurs kinetically at the hydride ligand first to form a bound hydrogen ligand even if the thermodynamic product is the dihydride ligand.53 In the absence of easily protonated oxygen or nitrogen based ligands as discussed above, this is attributed to intermolecular protonation at the metal center being kinetically slower due to electronic and geometric rearrangements that are avoided by protonation at the hydride ligand first.

1.5 Methods of Determining Rate Constants for TM Protonation

A number of techniques have been employed to quantify the kinetics of acid/base reactions involving TMHCs.

Due to the electronic rearrangements discussed above, TMHCs often exhibit distinct

UV/vis absorbance spectra as compared to their deprotonated conjugate bases. This optical difference makes stopped-flow techniques coupled to a UV/vis absorbance spectrometer a convenient method to monitor proton transfer rates between TMHCs and organic acids and bases in solution. In a typical setup, the metal complex in one syringe and the organic acid or base in another syringe are rapidly mixed using a pneumatic drive into a mixing chamber, wherein optical density is measured as a function of time.

A variety of 1D 1H NMR spectroscopic techniques have been used to measure proton transfer rates involving TMHCs. Line broadening of proton resonances has been used to determine rate constants of deprotonation of TMHCs by aniline bases that occur on the NMR timescale. For instance, the increasing excess line width of the cyclopentadienyl resonance for

CpW(CO)3H as a function of added aniline was used to calculate the rate constant for deprotonation. Variable temperature NMR has also been used to calculate proton transfer rate constants. In one example, the shape and broadening of the hydride resonance of cis-

Os(CO)4(CH3)H in the presence of NEt3 as a function of temperature was simulated and compared to experimental spectra in order to obtained a proton transfer rate constant. Finally, spin saturation labeling has been used to measure the rate constant for proton self-exchange

– between Os(CO)4H2 and its conjugate base [Os(CO)4H] wherein one of the hydride resonances was irradiated and the effect on the intensity of the other hydride resonance was observed.

A lesser used technique for studying intermolecular proton exchange (or intermolecular exchange in general) is 2D EXSY NMR (also called 2D NOESY in the literature). Proton self- exchange rate constants can be measured between a TMHC and its conjugate based with careful selection of NMR parameters including relaxation and mixing times. We have recently shown the utility of this technique by correlating proton self-exchange rate constants with propensity to undergo concerted proton-electron transfer (CPET) with a pair of tungsten hydride complexes. A comprehensive review on 2D EXSY NMR can be found elsewhere.54

The following Chapters 2 and 3 of this thesis will dive into a model system for further studying parameters that influence transition metal hydride formation. A cobalt-based model system is presented, the kinetics of protonation is studied using a variety of electrochemical and photochemical means, and then investigations into how electronic structure and steric influence affects the kinetic acidity of hydride formation is described.

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CHAPTER 2: VERSATILE MODEL SYSTEM FOR METAL-HYDRIDE FORMATION

Portions of this chapter adapted with permission from Elgrishi, N.; Kurtz, D. A.; Dempsey, J. L. Reaction Parameters Influencing Cobalt Hydride Formation Kinetics: Implications for Benchmarking H2-Evolution Catalysts. J. Am. Chem. Soc. 2017, 139 (1), 239–244. Copyright (2018) American Chemical Society.

2.1 Introduction

Alleviating societal reliance on carbon-based fossil fuels is necessary in order to mitigate

1–3 global climate change. Hydrogen (H2) fuel is an attractive alternative chemical fuel due to its

high energy density and benign emission upon combustion.4 The current technology dominating

hydrogen production worldwide is steam reformation, which involves heating methane/water

(CH4/H2O) mixtures to high temperatures in the presence of nickel catalysts to produce

hydrogen-rich syngas.5 While the process is currently relatively efficient (up to 90% efficiencies

have been achieved), a large energy input is needed to achieve the temperatures necessary to

drive the reaction and the process is not carbon neutral, an important factor in a sustainable

energy economy.

An alternative approach is to produce H2 from water. There are a number of approaches

to achieving efficient hydrogen evolution from water.6–10 Under standard conditions, the splitting

of water into hydrogen and oxygen requires a thermodynamic input of 1.23 V of energy. While

this may not seem energetically insurmountable (a low energy photon of 1008 nm infrared light

is thermodynamically sufficient), there are kinetic challenges inherent to the multi-step reaction

that result in an overpotential for hydrogen production.11 The energy required for water splitting

can be provided by either a potentiostat in the form of electrical energy or from the sun as solar

energy. Due to the aforementioned kinetic challenges, however, a catalyst is generally required in order to mediate the transformations of both water oxidation to O2 and water reduction to H2.

Both half reactions of water splitting involve the transfer of multiple protons and electrons in processes known as proton-coupled electron transfers (PCETs). Transition metal complexes are often used as catalysts for both water oxidation and reduction due to their ability to adeptly facilitate multi-electron and multi-proton transfer events. Often, catalytic proton reduction mediated by transition metal complexes proceeds through a metal-hydride intermediate.12–15 The formation of metal-hydride complexes by PCET can occur by multiple mechanisms as shown in

Figure 2.1: via one of two stepwise paths (electron transfer followed by proton transfer, ET-PT, or vice versa, PT-ET) or via concerted-proton electron transfer (CPET).

Figure 2.1 Square scheme showing stepwise ET-PT and PT-ET pathways and concerted pathway to form a transition metal hydride intermediate in H2 formation

As discussed in Chapter 1, MH formation via PCET is often slow, attenuated by kinetically sluggish protonation of the reduced metal complex. These PCET rates are slow enough that MH formation is often the rate limiting step for the catalytic hydrogen production, impacting the overall efficiency and rate of fuel generation.12–14,16 It is clear, then, that understanding the parameters that influence metal hydride formation is essential for future catalyst design.

Scheme 2.1 PCET reaction showing the addition of one proton and one electron to [CoIICp(PP)]+ (where PP = chelating bisphosphine ligand) to form the analogous [HCoIIICp(PP)]+ complex

While hydride formation via PCET has been discussed in the literature to an extent, the overwhelming literature is in the context of catalytic system wherein the metal hydride intermediate is not isolable and the formation of which is therefore convoluted with other steps in the catalytic cycle. Parameters such as TON and TOF are often the focus of such studies, and thus there is a dearth of systematic studies on the elementary steps involving the formation of transition metal hydride complexes through PCET reactions. Additionally, there is a limited number of model systems that have studied metal hydride formation using both electrochemical and photochemical initiated PCETs. We therefore sought a model system in which we could study the one electron, one proton addition to a transition metal complex to form a stable metal hydride complex by both electro- and photochemical methods. The system studied in this work

2+ uses the platform [CoCp(dppe)(NCCH3)] (dppe = 1,2-bis(diphenylphosphino)ethane). This ligand platform was reported to support a stable, isolable cobalt(III) hydride species,

[HCoCp(dppe)]+, enabling the study of the kinetics of the hydride formation step decoupled to further reactivity such as hydrogen evolution.17,18 Additionally, by varying the substituents on the phosphine ligand, we have be able to realize further trends in reactivity based on steric profile about the metal center (see Chapter 3). We report herein both the electrochemically and photochemically initiated step-wise PCET reaction to form [HCoCp(dppe)]+. Trends in reactivity are conserved across the two different methods; however, quantitative deviations between the methods are observed due to complicating side-reactivity.

2.2 General Results

2.2.1 Synthesis of Complexes

The synthesis of the cobalt-containing species discussed in this chapter are summarized

III in Scheme 2.2. The synthesis of [CoCp(dppe)(NCCH3)][PF6]2 (Co -dppe-NCCH3),

I III CoCp(dppe) (Co -dppe), and [HCoCp(dppe)][PF6] (HCo -dppe) have been described in detail previously.18,19 CoIII-dppe can alternatively be obtained via oxidation of CoI-dppe by two equivalents of ferrocenium hexafluorophosphate (Fc[PF6]), which resulted in an initial color change from dark red to green, followed by a color change to intense orange. Removal of solvent, dissolution in minimal CH2Cl2 followed by addition of the red solution to diethyl ether

II while stirring resulted in the precipitation of an orange solid. [CoCp(dppe)][PF6] (Co -dppe)

19 I was synthesized by oxidation of the previously isolated Co -dppe by either AgPF6 or Fc[PF6]

II in CH3CN to afford Co -dppe as a green powder. Alternatively, equimolar equivalents of

18,19 III I previously isolated Co -dppe-NCCH3 and Co -dppe can be combined in CH3CN to afford

CoII via a comproportionation reaction.

2.3 Electrochemical PCET

2.3.1 Electrochemistry in the Absence of Acids

In order to fully understand the PCET formation of HCoIII-dppe, it is important to thoroughly characterize and understand the electrochemical behavior of the parent complex,

III III Co -dppe-NCCH3. Cyclic voltammograms containing solutions of Co -dppe-NCCH3 contain

+ two diffusion-controlled, reversible redox features at –0.53 V and –0.93 V vs Fc /Fc in CH3CN

(Figure 2.2)

Scheme 2.2 Synthetic summary of all relevant Co-dppe species with photographs of each isolated sample. The postulated HCoIV complex has not been experimentally observed, and is therefore grayed out

III Figure 2.2 Cyclic voltammograms of Co -dppe-NCCH3 in CH3CN (0.50 mM, blue, bottom) and in CH2Cl2 (0.48 mM, red, top). The voltammograms were recorded at 100 mV/s in 0.25 M TBAPF6.

The diffusion coefficient (3.4 × 10−6 cm2/s), determined via variable scan rate experiments and the Randles-Sevcik equation is comparable with that of other molecular cobalt complexes.16,20 The heterogeneous electron transfer rate constants for both the CoIII and CoII

21 reduction were determined via working curves. Interestingly, the rate constant of ks = 0.051

cm/s for the CoIII/II couple is an order of magnitude faster than for cobaloximes.20 The larger rate

2 II II constant of ks = 0.11 cm /s for Co reduction is similar to that of Co reduction in hangman porphyrin molecules as well as that of ferrocene.16,22

I In order to protonate Co -dppe, the CH3CN ligand in the parent complex must dissociate.

Electrochemical experiments in CH2Cl2 provided insight into the CH3CN dissociation step: if

III/II II/I CH3CN was lost upon dissolution in CH2Cl2, then both the Co and Co couples would

I III/II II/I remain reversible; if CH3CN was lost upon reduction to Co , then both the Co and Co

II waves would become chemically irreversible; and if CH3CN was lost upon reduction to Co , then only the CoIII/II couple would become chemically irreversible. The third and final scenario is experimentally observed: the CoIII/II couple is irreversible while the CoII/I couple remains electrochemically and chemically reversible (Figure 2.2). The irreversible nature of the CoIII/II

II feature indicates that the coordinated CH3CN ligand is lost upon reduction to the Co oxidation state via an EC mechanism (Scheme 2.3).

III Scheme 2.3 EC’ mechanism of reduction of Co -dppe-NCCH3 followed by CH3CN ligand dissociation to form CoII-dppe

From variable scan rate data, the rate constant (4 × 107 s−1) for dissociation of the ligated

CH3CN was determined in CH2Cl2 (Figure 2.3). The rapid dissociation and the 0.42 V

III/II II/I I difference in the Co and Co waves indicates that the bound CH3CN dissociates before Co generation; this intramolecular chemical step is therefore not expected to influence hydride formation kinetics.

2.3.1 Voltammetric Response with Added Acid

III Cyclic voltammograms of Co -dppe-NCCH3 exhibit a marked difference when acids sources are added prior to electrode polarization. The CoII/I reduction loses reversibility and reductive peak position shifts to more positive potentials as the concentration or strength of the acid is increased. The EC mechanism in cyclic voltammetry can be described by the following equation:

where EP is the cathodic peak potential, E1/2 is the potential of the reversible electron transfer in the absence of substrate, R is the gas constant, T is the temperature, F is Faraday’s constant, υ is the scan rate of the cyclic voltammogram, kPT is the second order rate constant for

protonation of the CoI species and [HA] is the acid concentration.23 According to these equations, the Ep shifts positively by 30 mV/decade with log([substrate]) and by –30 mV/decade with log(υ). An example set of data is shown in Figure 2.4.

I Figure 2.4 Determination of kPT of Co -dppe with 4-cyano-anilinium tetrafluoroborate (Left) III Cyclic voltammograms of 0.5 mM of Co -dppe-NCCH3 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.006 M of acid. The CVs are + referenced to the Fc /Fc couple at 0 V and were recorded at 100 mV/s. (Right) Evolution of kobs 2 7 −1 −1 obtained with acid concentration (R = 0.997). The slope yields kPT = 1.7 × 10 M s

These changes are indeed consistent with an EC mechanism, wherein an electrochemical step is followed by an irreversible chemical step. Specifically, CoII-dppe is reduced to CoI-dppe followed by protonation by the acid to form HCoIII-dppe. Independent isolation and electrochemical analysis of HCoIII-dppe show that it can be reduced at potentials much more negative of those used in this study. It is clear, then, that according to these equations that the protonation rate constant can be easily obtained by observing the reductive peak potential of the

CoII/I wave evolve as a function of either acid concentration of scan rate.

2.3.2 Brønsted Relationship

Rate constants for the protonation of CoI-dppe were determined using 23 different

III molecular acids with pKa values ranging from 6.1 to 25.4 (Table B.1). The pKa of HCo -dppe was determined via spectrophotometric titration using triethylamine (NEt3) as a base (Figure

2.5).

Figure 2.5 (Left) Spectrophotometric titration in 5% toluene in CH3CN starting with 0.137 mM III HCo -dppe (blue trace ─) and titrating in increasing amounts of NEt3 (gray traces: 0.40, 0.80, 1.20, and 1.61 equivalents, respectively). End point of titration is reached upon the addition of I 6.4 equivalents of NEt3 (red trace ─), when spectral profile matches that of Co -dppe. The absorbance at 500 nm and the extinction coefficients of HCoIII-dppe and CoI-dppe were used to calculate concentration of each species as a function of added base. (Right) Linear regression of the equation for the acid base equilibrium between the species and the known pKa of NEt3 was III used to determine the pKa of HCo -dppe as 18.4.

III With the pKa of HCo -dppe in hand, we are able to accurately plot the kPT data in a

III Brønsted plot, wherein the x-axis is the difference in pKa between HCo -dppe and the acid used

(ΔpKa) and the y-axis is the log of the kPT rate constant. There are three distinct regions of the

Brønsted plot, discussed below.

First, acids in the ΔpKa range of roughly 7 to –2, there is a linear-free energy relationship

(LFER); as the acid strength is increased, the log of the kPT increases in a linear fashion (red markers, Figure 2.6) in accordance with classic Brønsted acid/base theory.24 The slope of the

LFER (Brønsted coefficient, α) is 0.55, which is relatively similar to α values for endergonic deprotonation reactions of transition metal carbonyl hydride complexes.25 The “ideal” Brønsted coefficient is 0.5 for proton transfers involving transition metal hydride complexes with relatively large intrinsic barriers for protonation (ΔG‡).26

I Figure 2.6 Brønsted plot of kPT values vs ΔpKa for the protonation reaction of Co -dppe to form HCoIII-dppe

The next group of acids all belong to the group of ΔpKa < –2, or more exergonic reactions. Interestingly, the kPT values in this region are acid-pKa independent, with an average

(plateau) rate constant of 1.7 × 107 M−1 s−1 (green markers, Figure 2.6). This plateau value is well below the estimated diffusion limit of ~1 × 107 M−1 s−1, estimated using the Debye-

Smoluchowski equation.19 In general, proton transfers involving nitrogen or oxygen based acids and bases will reach a maximum rate constant very near the diffusion limit. The fact that there is a large discrepancy between the maximum protonation rate constant and the diffusion limit in this study suggests that there is an intrinsic kinetic limitation to protonation (discussed further in

Chapter 3).

The final group of acids lie off of both the LFER line and the plateau line, thereby not belonging to either regime. The commonality between these acids is the presence of steric hinderance around the proton source, and thus we group the remaining acids into a “sterically encumbered” region (blue markers, Figure 2.6), albeit with no relationships drawn within the region as the steric influence is not identical amongst the acids.

2.4 Photoinduced PCET

In order to support and supplement our previous electrochemical analysis, we turned to time-resolved spectroscopic methods in order to optically detect the intermediates in the

III formation of HCo -dppe via PCET and map out the reaction kinetics further. Utilizing Ir(ppy)3

(ppy = 2-phenylpyridine, IrIII) as a strong photoreductant,27 we can photochemically generate

CoI-dppe in solution via flash-quench techniques. CoI-dppe is subsequently protonated to form the corresponding HCoIII-dppe complex with appropriate acids, and reactivity optically monitored on the ns-ms timescale utilizing transient absorption spectroscopy.

2.4.1 Optical Characterization of PCET Intermediates

Isolation and Characterization of Reactants, Anticipated Intermediates and Products. In order to photochemically trigger the formation of CoI-dppe via flash-quench methods, we used the isolated one electron-reduced complex, CoII-dppe. The electrochemistry of isolated CoII-

19 III dppe matches the reported behavior of Co -dppe-NCCH3, and the open circuit potential is

III/II +/0 II/I between the Co (E1/2 = –0.51 V vs. Fc ) and Co (E1/2 = –0.93 V) couples (Figure 2.7). The

UV/vis absorbance spectrum of the green CoII-dppe exhibits multiple absorbance bands in the range of 350-450 nm, and a broad, weak band centered around 580 nm (  140 M−1cm−1).

The cobalt complexes isolated in previous studies also display unique optical signatures in their UV/vis absorbance spectra (Figure 2.8). The red CoI-dppe exhibits relatively intense optical transitions around 375 nm and 480 nm (  4000 and 1500 M−1 cm−1, respectively), while the yellow HCoIII-dppe has only a weak absorption band around 380 nm (  1650 M−1cm−1).

III The UV/vis absorbance spectrum of orange Co -dppe-NCCH3 has two distinct absorption bands around 415 nm and 500 nm (  2000 and 740 M−1cm−1, respectively).

Table 2.1. Thermochemical data for iridium and cobalt species. Reaction ΔG (V vs Fc+/Fc) IrIV + e−  IrIII 0.39 IrIII*  IrIII –2.51 IrIV + e−  IrIII* –2.11 III − II Co -dppe-NCCH3 + e  Co -dppe + NCCH3 –0.51 CoII-dppe + e−  CoI-dppe –0.93

II/I III Based on the Co reduction potential (E1/2 = –0.93 V), we identified Ir(ppy)3 (Ir ) as a suitable photoreductant (E0* = –2.11 V vs. Fc+/Fc) for the flash-quench generation of CoI-dppe from the isolated CoII-dppe. While ruthenium tris-bipyridine derivatives are commonly utilized

in flash-quench-based studies, thermochemical analysis revealed that both oxidative and reductive quenching pathways were viable and likely using CoII-dppe as a quencher. The stronger photoreductant (and weaker photooxidant), IrIII is thus more ideal for this study. This cyclometalated complex is commonly used in organic light emitting diodes and photoredox catalysis studies, and its synthesis and reduction potentials in CH3CN have been reported (Table

2.1).28 However, to the best of our knowledge the optical properties of IrIII and its one electron

+ IV oxidized analog, [Ir(ppy)3] (Ir ) have not been previously reported in CH3CN.

Spectroelectrochemical oxidation of the yellow IrIII at ~0.5 V vs Fc+/Fc (Figure 2.8) cleanly generates IrIV, as evidenced by a bleach of the main visible transitions of IrIII from 350-500 nm, accompanied by the growth of a broad absorption band centered around 600 nm as well as higher energy transitions around 320 nm with clear isosbestic points at 346 and 486 nm. The electrochemical oxidation is fully reversible and the spectral signatures of IrIII are recovered upon application of a potential cathodic of the IrIV/III redox couple.

Figure 2.8. (Top) UV/Vis absorption spectra of spectroelectrochemical oxidation of Ir(ppy)3 + (blue) to generate [Ir(ppy)3] (red). (Bottom) Absorption spectra of synthetically isolated cobalt complexes.

2.4.2 Photoinduced Reactivity in the Absence of Acids

III Upon excitation, the photoreductant Ir exhibits strong photoluminescence (max = 525

5 −1 nm) with an excited state lifetime of 1.6 µs (Scheme 2.4 Equation 1, k0 = 6.07  10 s ). In the presence of CoII-dppe, the photoluminescence of IrIII is quenched (Figure 2.9) and Stern-

Volmer analysis gives a near diffusion limited second-order quenching rate constant (kq = 1.03 

1010 M−1 s−1, Scheme 2.4 Equation 2). This large quenching rate constant is unsurprising as the oxidative quenching reaction between IrIII* and CoII-dppe has a driving force of nearly 1.2 V

(Table 2.1).

Figure 2.9. Time-resolved photoluminescence quenching of [Ir(ppy)3]* decay in the presence of increasing concentration of CoII-dppe and Stern-Volmer quenching plot (inset).

To observe optical changes in cobalt-containing species following photoinduced electron transfer, we probed at the IrIII/IrIV isosbestic point (λ = 486 nm). A transient signal corresponding to the formation CoI-dppe is observed at 486 nm, and persists on the μs timescale, consistent with oxidative quenching of IrIII* by CoII-dppe. The signal decays with second- order, equal concentration kinetics as expected. However, we did not observe the concomitant

IV formation of Ir (λmax = 650 nm), consistent with thermochemical analysis that indicates that

IV II Ir is capable of oxidizing the excess Co -dppe in solution (ΔGET = −20.8 kcal/mol) (Scheme

2.4 Equation 4).

II IV The rate constant for oxidation of Co -dppe by Ir (kexc) was estimated by monitoring the decay of the transient signal at 586 nm corresponding to back-electron transfer reaction for

III III* 9 −1 −1 the photoinduced reduction of Co -dppe-NCCH3 by Ir (kexc = 2.3  10 M s ) (Figure

2.10). This rate constant is an estimate because the kinetics involved with dissociation of the

III acetonitrile ligand of Co -dppe-NCCH3 upon reduction likely influence the simple electron transfer reactions discussed above.

Figure 2.10 Plot of second-order decay of transient signal at 586 nm after excitation of Ir(ppy)3 III at 445 nm in the presence of 0.057 mM of Co -dppe-NCCH3. Using the slope of 1/ΔOD vs time in the above region and the coefficients of the absorbing species, the kexc rate constant was determined to be 2.3 × 109 M−1s−1. The open circles are simulated data using MatLab but using a 9 −1 −1 kexc = 2.9 × 10 M s .

Scheme 2.4. Summary of kinetic processes discussed herein.

Recognizing that the photoinduced electron transfer from IrIII* to CoII-dppe initially yields IrIV and CoI-dppe, and the IrIV reacts with the excess CoII-dppe in solution to produce

CoIII-dppe, we realized that the long-lived transient signal at 486 nm on the μs time scale corresponds not exclusively to CoI-dppe absorption, but to absorption from equimolar amounts

I III III of immediately generated Co -dppe and Co -dppe (or Co -dppe-NCCH3 depending on the rate of CH3CN ligation), as both these species have larger molar extinction coefficients than

II −1 −1 I −1 −1 III Co -dppe at 486 nm (ε486 nm = 1460 M cm for Co -dppe, ε486 nm = 870 M cm for Co -

−1 −1 II dppe-NCCH3, and ε486 nm = 180 M cm for Co -dppe). As discussed above, the

III I II comproportionation of Co -dppe-NCCH3 and Co -dppe cleanly generates Co -dppe, giving rise to decay of the signal at 486 nm on the µs timescale. A second-order rate constant is

9 −1 −1 obtained from the second-order, equal concentration fit to this signal (kcom = 1.1  10 M s ,

Scheme 2.4 Equation 5, Figure 2.11).

Figure 2.11 Plot of second-order decay of transient signal at 486 nm after excitation of Ir(ppy)3 at 445 nm in the presence of 0.98 mM of CoII-dppe. Using the slope of 1/ΔOD vs time in the above region and the coefficients of the absorbing species, the kcom rate constant was determined to be 1.1 × 109 M−1s−1. The open circles are simulated data using MatLab, which show excellent agreement with the experimental data.

Kinetic simulations were performed numerically with a series of differential equations that were derived from the reaction model described in Scheme 2.4. Vectors containing initial concentrations of all species involved, the time interval for analysis, and rate constants were provided as inputs to solve the initial value problem. Concentration–time traces for each species generated as the output from an ordinary differential equation solver were converted to simulated transient absorption difference spectra via the Beer-Lambert law and summation of absorbance of each species at the wavelength of interest. The rate constants that were determined via independent experiments described above (k0, kq, kexc, kcom) were held constant throughout the simulations, and the remaining rate constants (k) were adjusted iteratively to optimize the match between the simulation and the dataset (Figure 2.10 and Figure 2.11). Simulations fit best with a

9 −1 −1 kexc value of 2.9  10 M s , differing slightly from the experimentally determined value described above (2.3  109 M−1 s−1).

2.4.3 Kinetic Processes with Added Acid

As noted above, the transient signal at 486 nm decays with 2nd order, equal concentration kinetics on the µs timescale arising from the comproportionation reaction between CoI-dppe and

CoIII-dppe species. When acids are added to solution, the signal decay is accelerated and fits well to a single exponential kinetics model with >10 equiv. (Figure 2.12 and Figure B.32 –

Figure B.37). The rate of decay is dependent on the acid concentration, suggesting the CoI-dppe reacts with the proton sources to form HCoIII-dppe (Scheme 2.4 Equation 6). Observed rate constants (kobs) from the single exponential fit are linearly correlated with the concentration of each acid, and linear regression yields the second-order rate constant for CoI-dppe protonation

(kPT).

Figure 2.12.Transient absorption data at 486 nm after exciting 150 µM Ir(ppy)3 at 445 nm in the presence of 1.0 mM CoII-dppe in the presence of increasing concentration of 2,3,5,6-tetrafluoro- 4-trifluoromethylphenol (concentration of acid in each trace shown in legend). Single exponential fits are also shown (black lines), and the inset shows resulting kobs from fits vs acid concentration.

Protonation rate constants were determined for a series of acids ranging from pKa = 21.51 to

16.62 (Figure B.32 – Figure B.37 and Table 2.2). As anticipated, kPT values increase with

increasing acid strength. In the pKa range accessible in this study, we observe a linear free energy relationship (LFER) for nearly all acids (Figure 2.13). The Brønsted slope of the LFER is 0.40, notably smaller than the Brønsted slope of the proton transfer rate constants determined using electrochemical methods (α = 0.55).19 Additionally, the absolute value of the rate constants measured here are slightly lower than those measured using electrochemical methods (Figure

2.14). Similar to our previous study, we observe noticeable attenuation of the kPT rate constant with sterically bulky acids, such as triethylammonium tetrafluoroborate ([HNEt3][BF4]); these acids are thus not included in the free energy correlation.

Table 2.2 Summary of photochemically determined kPT values

−1 −1 Acid pKa ΔpKa kPT (M s ) log(kPT) benzoic acid 21.51 3.11 4.2 × 105 5.62 2,3,5,6-tetrafluorophenol 20.12 1.72 1.4 × 106 6.15 triethylammonium tetrafluoroborate 18.82 0.42 3.4× 105 5.53 2,3,4,5,6-pentachlorophenol 18.02 −0.38 6.2 × 106 6.79 salicylic acid 16.7 −1.7 3.5× 107 7.54 2,3,5,6-tetrafluoro-4-trifluoromethylphenol 16.62 −1.78 4.4× 107 7.64

Figure 2.13. Brønsted plot of Equation 6 and structures of acids used. Linear regression gives a Brønsted slope of α = 0.4.

Kinetic simulations were performed as described above, holding all previously discussed electron transfer rate constants constant and using the kPT rate constants determined from single exponential decay kinetics. In contrast to the kinetics simulations in the absence of acid described above, the kinetics simulations with acids did not closely match the experimental data.

Minor iterative changes to the kPT rate constants did not improve the overlay significantly. This suggested additional acid-related ground-state or excited-state reactivity.

Figure 2.14 Overlaid Brønsted plots comparing electrochemically determined kPT rate constants I and photochemically determined kPT rate constants for Co -dppe.

2.5 Discussion

2.5.1 Disparity Between Measured Protonation Rate Constants

Ground State Reactions of Co in the Presence of Acids. In our previous studies utilizing electrochemical methods, we were able to extend the Brønsted plot for proton transfer rate

19 constants to acids as strong as dimethylformamidinium triflate (pKa in CH3CN = 6.1). When attempting to probe photoinduced PCET reactivity in this study, however, a significant ground state reaction between the CoII-dppe and stronger acids in solution was readily detected visually,

as the solution changed from green to orange. To further probe the ground state interactions and their influence on the quantification of proton transfer rate constants, reactions between CoII- dppe and acids were monitored via UV/Vis absorption spectroscopy (Figure 2.15).

Qualitatively, the solution color changes from green to orange with fewer equivalents of stronger acids as compared to weaker acids, indicating a pKa-dependent process.

Figure 2.15 (Left): UV/vis absorbance titration of CoII-dppe with up to 9.5 equivalents of 4- methyl-anilinium tetrafluoroborate in CH3CN (pKa = 11.4), and (right) UV/vis absorbance titration of CoII-dppe with up to 57.5 equivalents of 4-methoxy-anilinium tetrafluoroborate in CH3CN (pKa = 11.86).

In order to characterize the reaction products observed optically, one equivalent of HBF4

29 II etherate (pKa ≈ 1.5) was added to a solution of Co -dppe, resulting in a color change from green to orange. Addition of diethyl ether led to precipitation of a bright orange powder. A cyclic voltammogram of the isolated solid exhibits reduction features consistent with the CoIII/II and

II/I III II I Co waves observed for isolated Co -dppe-NCCH3, Co -dppe, and Co -dppe species (–0.51

+ V and –0.93 V vs. Fc /Fc, Figure 2.16), along with another, irreversible reduction at Epc = –1.88

V vs. Fc+/Fc which matches the reduction feature of HCoIII-dppe (Figure 2.16).

Figure 2.16 Cyclic voltammograms comparing the orange product of the reaction between CoII- dppe and HBF4 etherate described in the main text (top), and independently prepared samples of III III HCo -dppe (middle) and Co -dppe-NCCH3 (bottom).

These data are consistent with the identification of the orange precipitate as a mixture of

III III 1 yellow HCo -dppe and red Co -dppe-NCCH3. H NMR spectroscopy was used to further characterize the products of this ground state reaction between CoII-dppe and acid. While CoII- dppe is paramagnetic and has very broad 1H NMR resonances, the addition of excess strong acid

(4-(methylbenzoate)anilinium, pKa = 8.62) gives rise to two sharp diamagnetic Cp resonances with equivalent integrations (Figure 2.17). Comparison to authentic samples of CoIII-dppe-

III NCCH3 and HCo -dppe confirms the identity of the reaction products.

1 II Figure 2.17. H NMR spectrum in CD3CN of the reaction between Co -dppe and an excess of 1 strong acid (4-(methylbenzoate)anilinium, pKa = 8.62) (top) compared with H NMR spectra of III III independently synthesized HCo -dppe (middle) and Co -dppe-NCCH3 (bottom). Symbols * indicate resonances corresponding to the 4-(methylbenzoate)anilinium and its conjugate base, and # indicates resonances corresponding to solvents (CD2HCN and diethyl ether).

Eq. 7 describes the stoichiometry of the above reactivity with a generic acid HA (Scheme

2.5). This net reaction is a combination of the disproportionation free energy of CoII-dppe to

I III I III form Co -dppe and Co -dppe-NCCH3, and the protonation of Co -dppe to form HCo -dppe.

III III Thus, the equilibrium concentration of Co -dppe-NCCH3 and HCo -dppe formed upon mixing CoII-dppe and acid can be calculated based on the initial concentrations of CoII-dppe and HA added to each cuvette as well as the acid pKa. These calculations for the acids used in

III this study (pKa values between 21.51 and 16.62) revealed that the concentration of HCo -dppe

III and Co -dppe-NCCH3 prior to laser excitation in the photoinduced PCET reactions described above are on the same order of magnitude as the concentration of photogenerated CoI-dppe.

Scheme 2.5 Ground state reactivity between CoII-dppe and acid as well as CoIII ligation

With a clearer understanding of the ground state reactions between CoII-dppe and acid, kinetics simulations were revised to take this ground state reactivity into account, specifically by employing initial concentration vectors that more accurately describe the initial concentrations of species in solution. With these revisions to the ground state conditions, the kinetics simulations data using the same rate constants as described above better matched the experimental data but still showed marked deviations. Additional clues to the sources of the deviations were discovered by 1H NMR. Upon mixing CoII-dppe and salicylic acid, an additional set of resonances was

III III observed that did not correspond to the Co -dppe-NCCH3 or HCo -dppe. Separately, the

III addition of sodium salicylate to Co -dppe-NCCH3 yielded the same new set of resonances, which we assign to [CoCp(dppe)(salicylate)]+ (CoIII-dppe-salicylate), wherein the parent

CH3CN ligand is displaced. The appearance of this new species when salicylic acid is added to

CoII shows that the conjugate base (A–) formed in Scheme 2.5 Equation 7 can compete with

III CH3CN ligation to form species of the type Co -dppe-A (Scheme 2.5 Equation 8).

Each of the acids used in the present study (Table 2.2) have different conjugate bases, and thus will have differing equilibrium constants for the ligation of the conjugate base to CoIII- dppe-NCCH3 (Scheme 2.5 Equation 8 equilibrium of ligation, or KL). Additionally, it is expected that each CoIII-dppe-A species will have a different CoIII/II reduction potential, influenced by the electronic structure of the ligand, and heterogeneous electron transfer rate constant for the CoIII/II reduction event. For example, when CoIII-dppe-salicylate is generated by

III the addition of excess sodium salicylate to Co -dppe-NCCH3, cyclic voltammetry shows

III marked differences between this species and Co -dppe-NCCH3. The addition of excess sodium salicylate causes the CoIII/II redox feature to significantly broaden, indicating electrochemical quasi-reversibility and relatively slower electron transfer kinetics (Figure 2.18).

III III Figure 2.18 Cyclic voltammograms of Co -dppe-NCCH3 and Co -dppe-salicylate in CH3CN.

Slow CoIII/II electron transfer kinetics is not uncommon.20,30 The wave is very broad even at low scan rates, which hinders the measurement of an exact E1/2 value; however, it appears to be qualitatively cathodically shifted from the analogous wave in the cyclic voltammogram of the

III parent Co -dppe-NCCH3. The relevance of the differing KL values, reduction potentials of

CoIII-dppe-A species, and CoIII/II ET kinetics of those species is discussed below.

In the absence of acid, the decay of the transient signal at 486 nm on the µs timescale is

III I the comproportionation reaction between Co -dppe-NCCH3 and Co -dppe (Scheme 2.4

Equation 5). While the kcom rate constant has been measured in the absence of other possible

III ligands for Co -dppe-NCCH3 (Figure 2.11), it is safe to assume that the comproportionation rate between CoI-dppe and CoIII-dppe-A species will vary slightly due to a difference in driving force and intrinsic CoIII/II electron transfer kinetics. Due to the improvement of the kinetic simulations from the inclusion of the minor ground state reactivity between CoII and acids described above, we hypothesize that the remaining disparity between simulated kinetic data and experimental data stems from CoIII ligation chemistry, its thermodynamic and kinetic

III implications, and subsequent optical changes as compared to Co -dppe-NCCH3. For example,

preliminary work measuring the CH3CN ligand displacement with tetraethylammonium 2,3,5,6- tetrafluorophenolate ([TEA][F4-PhO]) show an increased absorbance at 486 nm as compared to the parent CH3CN ligated species (Figure 2.19). Efforts are ongoing in our lab to measure the

+ III/II redox properties of this new species (hypothesized to be [CoCp(dppe)(F4-PhO)] ), its Co electron transfer properties, and incorporate those data into the kinetic simulations in the expectation that the agreement with experimental TA data will be improved.

III Figure 2.19 Titration of Co -dppe-NCCH3 with increasing equivalents of [TEA][F4-PhO] in CH3CN.

2.6 Conclusions

In summary, we have demonstrated a robust platform in which to study the elementary

PCET formation of a metal-hydride complex. We demonstrated that the rate constants for the stepwise electron transfer-proton transfer events can be measured explicitly using both electrochemical and time-resolved spectroscopic techniques.

Electrochemical proton transfer rate constants were determined over a range of 19+ pKa values in CH3CN, which revealed a complicated relationship between rate constant and proton transfer driving force. Weaker acids follow a LFER between rate constant and driving force, with a transition to a “pKa-independent” relationship when moving to stronger acids. This plateau in

proton transfer rate constant with increasing driving force has important implications in catalytic

III hydrogen evolution systems; if Co -dppe-NCCH3 were an electrocatalyst for proton reduction with rate-limiting hydride formation, the choice of strongly acidic 4-cyano-anilinium as an acid source as compared to salicylic acid would result in the introduction of a ~570 mV increase in overpotential with a negligible enhancement of catalytic rate. This study also begins to reveal the effect that sterics plays in the protonation of metal complexes, wherein the acids 4-methylanilinium and N,N-dimethylanilinium have nearly the same pKa, but the kPT rate constant to form

HCoIII-dppe varies by two orders of magnitude. More extensive studies on steric effects on protonation of metal complexes can be found in Chapter 3.

Photoinduced electron transfer initiated by excitation of a strong photoreductant followed by ground state proton transfer by acids was used in order to spectroscopically identify intermediates in the step-wise formation of HCoIII-dppe from CoII-dppe. Reaction intermediates in this photon-triggered PCET can be definitively identified through comparison of transient spectra to the optical signatures of the isolated intermediates for this ET-PT reaction, allowing comprehensive analysis of the reaction kinetics and extraction of rate constants for elementary reaction steps. Proton transfer rate constants for a limited range of acids were able to be determined using photochemical methods, which qualitatively agree with those determined using electrochemical methods. The lack of close agreement stems from additional side reactivity due to having to use a further reduced starting material in the photochemical studies, (namely CoII-

III dppe as opposed to Co -dppe-NCCH3).

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(24) Eigen, M. Proton Transfer, Acid-Base Catalysis, and Enzymatic Hydrolysis. Part I: Elementary Processes. Angew. Chemie Int. Ed. English 1964, 3 (1), 1–19.

(25) Edidin, R. T.; Sullivan, J. M.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition-Metal Complexes. 4. Kinetic Acidities toward Aniline and Their Use in Identifying Proton-Transfer Mechanisms. J. Am. Chem. Soc. 1987, 109 (13), 3945– 3953.

(26) Weberg, R. T.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition- Metal Complexes. 6. Interstitial Hydrides. J. Am. Chem. Soc. 1990, 112 (3), 1105–1108.

(27) Prier, C. K.; Rankic, D. A.; MacMillan, D. W. C. Visible Light Photoredox Catalysis with Transition Metal Complexes: Applications in Organic Synthesis. Chem. Rev. 2013, 113 (7), 5322–5363.

(28) Dedeian, K.; Djurovich, P. I.; Garces, F. O.; Carlson, G.; Watts, R. J. A New Synthetic Route to the Preparation of a Series of Strong Photoreducing Agents: Fac-Tris-Ortho- Metalated Complexes of Iridium(III) with Substituted 2-Phenylpyridines. Inorg. Chem. 1991, 30 (8), 1685–1687.

(29) Paenurk, E.; Kaupmees, K.; Himmel, D.; Kütt, A.; Kaljurand, I.; Koppel, I. A.; Krossing, I.; Leito, I. A Unified View to Brønsted Acidity Scales: Do We Need Solvated Protons? Chem. Sci. 2017, 8 (10), 6964–6973.

(30) Macartney, D. H.; Sutin, N. Electron-Exchange Rates of Polypyridine Complexes: Electron-Transfer Reactions Involving the Tris(Polypyridine)Nickel(II/III) Couple in Acidic Aqueous Media. Inorg. Chem. 1983, 22 (24), 3530–3534.

CHAPTER 3: STERIC INFLUENCE ON KINETIC ACIDITY OF METAL HYDRIDES

3.1 Introduction

Just like a chain is only as strong as its weakest link, catalysis is only as fast as its slowest

step.1 Therefore, both the nature of the rate-limiting step in a catalytic cycle and the factors that

influence that step are vital to improving the overall efficiency. In fuel-forming reactions such as

hydrogen production or carbon dioxide reduction to formic acid, the formation of transition

metal-hydride bonds is a common elementary step that can have a significant impact on the

overall rates of product generation.2–5 A transition metal hydride complex is often formed

through a proton-coupled electron transfer process, via either a stepwise pathway (proton transfer

followed by electron transfer, PT-ET, or vice-verse, ET-PT) or a single concerted step (CPET).5

Figure 3.1 Square scheme showing stepwise and concerted pathways to form a metal hydride complex, a key intermediate in catalytic hydrogen evolution

While there have been numerous reports describing the application of electrochemical

techniques to quantify the performance of first-row transition metal hydrogen evolution

catalysts,4,6–9 limited work has focused on the specific parameters that influence the formation of

metal-hydride reaction intermediates. Chapter 2 described a systematic study of how acid strength and acid structure influenced the rate constant of protonation (kPT) of an electrochemically generated CoCp(dppe) to form a stable cobalt(III) hydride complex.10 We observed a linear free-energy relationship for relatively weak acids and noted that sterically encumbered acids result in dramatically slower protonation when compared to acids with similar pKa values and no steric bulk near the acidic proton. Intriguingly, we also observed a distinct plateau of kPT values with stronger acids that extends over 14.5 kcal/mol of driving force for protonation. The second order rate constants were 3 orders of magnitude lower than the calculated diffusion limited rate constant. These results underscored the complexity of a seemingly simple hydride formation reaction.

It has been known for decades that protonation/deprotonation reactions involving metal complexes are more similar to those involving carbon acids than oxygen- or nitrogen-based acids.11–15 Large geometric and structural rearrangements can occur during acid/base reactions involving metal complexes, imparting large reorganization energies and thereby impacting the rate constants of these reactions. We hypothesized that these large reorganization energies could limit proton transfer in metal-hydride species. A straightforward way to probe the correlations between reorganization of metal-hydride complexes and their rate constants for protonation/deprotonation is to change the ligand steric profile. To that end, we have synthesized a series of derivatives of the parent cobalt cyclopentadienyl bisphosphine complex that differ by the phosphine substituents and backbone structure. Electrochemical determination of kPT values for each complex with a variety of acids reveal that phosphine substituents have a significant effect on protonation kinetics. We find that while the electronics of the complexes are altered as a function of the R groups on the phosphine, comparison of kPT values on a relative energy scale

reveal dramatic differences in proton transfer rate constants that arise from differing intrinsic barriers to protonation imparted by the ligand sterics. Proton transfer self-exchange rates constants measured using 2D 1H NMR were used to justify the intrinsic barriers for protonation to form the HCoIII complexes, which show the dramatic effect that sterics play on the rate constants for hydride formation

3.2 Results and Discussion

3.2.1 Synthesis and Characterization

In order to study the influence of steric parameters on the formation of cobalt(III) hydride complexes, we synthesized a series of complexes with varying degrees of steric differences in both the ligand backbone and the phosphine substituents (Figure 3.2). Two different structural

III changes were made to the parent complex [CoCp(dppe)(NCCH3)][PF6]2 (Co -dppe, dppe = 1,2-

III bis(diphenylphosphino)ethane). [CoCp(dppv)(NCCH3)][PF6]2 (Co -dppv, dppv = cis-1,2- bis(diphenylphosphino)ethylene) has an unsaturated (vinyl) backbone in the phosphine ligand

III instead of an ethylene linker. [CoCp(depe)(NCCH3)][PF6]2 (Co -depe, depe = 1,2-

III bis(diethylphosphino)ethane) and [CoCp(depe)(NCCH3)][PF6]2 (Co -dcpe, 1,2- bis(dicyclohexylphosphino)ethane) have the same ethylene backbone as the parent CoIII-dppe but different substituents on the phosphine

The orange cobalt(III) complexes (CoIII-dppe, CoIII-depe, CoIII-dcpe, CoIII-dppv) were all synthesized through routes described previously involving a common [CoCp(SMe2)3][PF6]2

(SMe2 = dimethylsulfide) starting material. The yellow cobalt(III) hydride complexes were prepared through in situ generation of the CoI complex via ligand exchange followed by protonation with NH4PF6 (see experimental section). The cobalt(I) complexes were synthesized via deprotonation of the isolated cobalt(III) hydride complexes using NaH.

Figure 3.2. Structure of the complexes of the type [CoCp(dxpx)(NCCH3)][PF6]2 and the structural permutations studied herein (dxpx is an abbreviation encompassing depe, dppe, dcpe, and dppv)

X-ray quality crystals of HCoIII-depe, HCoIII-dppe, and HCoIII-dcpe were grown from the vapor diffusion of diethyl ether into solutions of the hydride complexes in CH3CN. There is relatively little deviation of the Co-Cp or Co-P bond distances across the series. The most dramatic difference between the series is the accessibility of the hydride ligand to incoming

bases. Two of the R groups on each bisphosphine ligand (one on each phosphorous) flank the hydride ligand on either side.

Figure 3.3 Structural representation of HCoIII-depe (left), HCoIII-dppe (middle), and HCoIII- dcpe (right) viewed looking down on the top of the Cp ring with the hydride ligand pointing down to highlight the differing angle of access to the metal center or hydride ligand. For HCoIII-depe and HCoIII-dcpe, thermal ellipsoids are shown at the 50% probability level; for HCoIII-dppe the quality of the data was such that only a connectivity structure could be obtained at this time.

3.2.2 Electrochemistry in the Absence of Acids

For each complex, the cyclic voltammogram recorded in 0.25 M TBAPF6 acetonitrile solution exhibits two, reversible waves assigned to CoIII/II and CoII/I reductions, respectively

(Figure 3.4 and Table 3.1). All reductions are chemically and electrochemically reversible and

10 III have ks values similar to those of the parent dppe complex (Table 3.1). When Co -dcpe is left

III/II sitting in CH3CN under N2, the Co wave broadens and becomes more quasi-reversible over several hours (Figure 3.5), possibly due to slow isomerization of the cyclohexyl substituents on the dcpe ligand. As CH3CN ligand dissociation was found to occur immediately following

III reduction of Co -dppe, the stereochemistry of the cyclyohexyl group may influence CH3CN association/dissociation and therefore influence the reversibility of the CoIII/II wave in the dcpe complex. The CoII/I wave does not appear to be influenced by this slow isomerization, however, as that redox couple does not change over the same time scale as the CoIII/II wave (Figure 3.5).

III Alternatively, the as-synthesized Co -dcpe may not have a CH3CN coordinated due to the steric

bulk of the cyclohexyl groups, and dissolution in CH3CN in electrochemistry experiments slowly establishes an equilibrium between species with and without nitrile bound. Further experiments are ongoing to probe this reactivity using NMR spectroscopy.

The reduction potentials of CoIII-dppe and CoIII-dppv are very similar in energy, indicating a mild electronic withdrawing effect of the alkene linker as compared to the alkane backbone. This mild anodic shift in reduction potentials when changing from a vinyl backbone to an ethylene backbone has been observed in nickel complexes previously.16 The CoII/I couples for both alkyl phosphine complexes are more negative than that of the dppe complex, consistent with alkyl phosphines being more donating than aryl phosphines.16 The same is true for the

CoIII/II couple for CoIII-depe, but the CoIII/II couple for CoIII-dcpe is actually more anodic than that of the CoIII-dppe complex.

III Figure 3.4. Cyclic voltammograms of isolated Co complexes in 0.25 M TBAPF6 in CH3CN. All voltammograms are referenced to the internal standard, the Fc+/Fc redox couple.

III Figure 3.5 Cyclic voltammograms of Co -dcpe-NCCH3 over time

Table 3.1 Summary of electrochemical properties of cobalt complexes.

Diffusion Coefficient Potential (V vs. Fc+/Fc) k (cm/s) Complex (D, cm2/s) s CoIII/II CoII/I CoIII/II CoII/I CoIII/II CoII/I Co-depe –0.71 –1.21 4.4 × 10–6 6.7 × 10–6 0.022 0.36 Co-dppe –0.51 –0.93 3.4 × 10–6 5.2 × 10–6 0.051 0.11 Co-dcpe –0.39 –1.25 - 7.0 × 10–6 - 0.19 Co-dppv –0.50 –0.88 5.9 × 10–6 7.8 × 10–6 0.19 0.50

3.2.3 Voltammetric Response with Added Acid

When acids of appropriate pKa are added to solutions of the cobalt(III) complexes, the

CoII/I wave becomes irreversible and the reductive peak potential shifts anodically. The peak position evolves as a function of both added acid concentration and scan rate, consistent with an

EC process in which an irreversible chemical (C) step follows electron transfer (E). Based on our studies of CoIII-dppe, we assign the chemical step following reduction of CoII as protonation of the CoI species to form the corresponding HCoIII complex. The peak potential shifts as a function of both acid concentration ([HA]) and scan rate (υ) according to:

푅푇 푅푇 푘푃푇[퐻퐴]푅푇 퐸푝 = 퐸1 − (0.78) + ln ( ) ⁄2 퐹 2퐹 퐹휐 where R, T, and F are the gas constant, temperature (in K), and Faraday’s constant, respectively,

E1/2 is the potential of the one-electron reversible electrochemical event in the absence of a following chemical step, and kPT the rate constant for protonation. The evolution of the peak potential with acid concentration or scan rate allows ready determination of kPT (Figure 2.4).

3.2.4 Brønsted Relationship

Proton transfer rate constants were determined for a variety of cobalt/acid combinations

III using acids that span 22 pKa units in acetonitrile; 9 different acids for Co -dppv, 22 acids for

CoIII-depe, 14 acids for CoIII-dcpe, and 23 acids for previously published data for CoIII-dppe.

The rate constants for the same acids with different cobalt complexes are different, indicating that parameters about the cobalt complexes influence the rate constant for protonation. These differences arise from a combination of electronic structure differences causing differences in driving force for proton transfer (dictated by the difference in pKa values of the hydride complex and the acid) and differences in reactivity dictated by the ligand steric profile.

III III To quantify the pKa values of HCo -depe and HCo -dcpe, the hydride complexes were isolated and spectrophotometric titrations were performed with appropriate organic bases to establish equilibria between HCoIII-dxpx, the corresponding CoI-dxpx, and the organic acid and base (example shown in Figure 3.6)

Figure 3.6 (Left) Spectrophotometric titration in 5% toluene in CH3CN starting with 1.17 mM HCoIII-depe (blue trace ─) and titrating in increasing amounts of DBU (gray traces). UV/vis absorbance spectrum of 1.17 mM CoI-depe is shown to indicate end point (red trace ─). The absorbance at 500 nm and the extinction coefficients of HCoIII-depe and CoI-depe were used to calculate concentration of each species as a function of added base. (Right) Linear regression of the equation for the acid base equilibrium between the species and the known pKa of DBU was III used to determine the pKa of HCo -depe as 23.6.

III III The pKa of HCo -dcpe is 22.6 and the pKa HCo -depe is 23.6, indicating the species

III are significantly less acidic than the previously reported HCo -dppe (pKa = 18.4), as predicted for the more strongly donating alkyl phosphines in comparison to aryl phosphines. The pKa of

HCoIII-dppv was determined via NMR equilibrium by added HCoIII-dppv to CoI-dppe and allowing equilibrium to be established, and then integrating the Cp resonances for the 4 species

III in solution (Figure 3.7). The calculated pKa of HCo -dppv is 18.2, very similar to that of

HCoIII-dppe, indicating a mild electron withdrawing effect of the alkene backbone, consistent with the reduction potentials discussed previously. Interestingly, these electronic affects are also reflected in the hydride resonance in the 1H NMR of each complex (Figure 3.8); the alkyl phosphine complexes exhibit hydride resonances more upfield (-16.7 ppm for HCoIII-depe and -

16.5 ppm for HCoIII-dcpe) than the aryl phosphines (-15.3 ppm for HCoIII-dppe and -15.6 ppm for HCoIII-dppv).

Figure 3.7 1H NMR spectrum of mixture of HCoIII-dppe, HCoIII-dppv, CoI-dppe, and CoI- dppv in 20% C6D6 in CD3CN for solubility. The integrations of the Cp resonances for each species was used to calculate the equilibrium constant for the indicated reaction, and Keq and the III III known pKa of HCo -dppe was used to calculate the pKa of HCo -dppv to be 18.2.

1 III Figure 3.8 Hydride region of the H NMR spectra in CD3CN of the HCo species studied herein

Brønsted plots are presented as kPT vs. ΔpKa, where ΔpKa is the difference between the

III pKa of the added acid and the pKa of the HCo complex (Table B.1, Table B.3, Table B.4,

Table B.5, Figure B.10 - Figure B.31, and Figure B.38 - Figure B.69). Taking into account the

electronic influence of the phosphines by the use of this scaled x-axis, we can realize trends in reactivity due to the steric factors of the ligands alone. This analysis reveals Brønsted plots that are extraordinarily similar in shape across each phosphine ligand with 3 distinct regions for each

(Figure 3.10); a linear-free energy relationship (LFER) region, a class of acids that contain steric bulk about the proton source, and a region where protonation rate constants plateau at mildly exergonic proton transfers (~ΔpKa = 3).

First, we observe very little difference in the protonation rate constants for CoI-dppe and

CoI-dppv (Figure 3.9). This suggests that the conjugated backbone does not significantly influence kPT.

Figure 3.9 Overlaid Brønsted plots for the protonation of CoI-dppe and CoI-dppv

Each of the Brønsted plots contain a LFER region with relatively similar Brønsted slopes

(α). There is no significant trend in α values across the series: αdepe = 0.51, αdppe = 0.55, and αdcpe

= 0.45. These values are very similar to other reported Brønsted slopes for the protonation of a transition metal hydride complex.15 As in our previous work with CoIII-dppe, protonation reactions for each of the cobalt complexes studied here with sterically bulky acids are slower

than expected for the given pKa, as compared to other acids with similar pKa values in the LFER or plateau region (discussed below). These acids include pentabromophenol, triethylammonium tetrafluoroborate, and a series of para-substituted 2,6-dimethylpyridinium salts. Although there appears to be qualitative free-energy relationships within the group of bulky acids, we refrain from drawing quantitative conclusions due to the differing steric profile within the group.

Figure 3.10 (Left) Overlaid Brønsted plots for the protonation of CoI-depe, CoI-dppe, and CoI- dcpe using non-bulky acids, and (right) overlaid Brønsted plots for the protonation of CoI-depe, CoI-dppe, and CoI-dcpe highlighting sterically encumbered acids.

For each complex, protonation rate constants become independent of acid strength when the difference between the pKa of the hydride complex and the acid exceeds ca. 3 pKa units. This leads to a distinct plateau in the Brønsted plot. While each of the 4 complexes exhibits this distinct plateau in the respective Brønsted plots, the maximum rate constant for proton transfer differs for each phosphine ligand depe, dppe, and dcpe (though the plateau observed is indistinguishable for dppe and dppv). Qualitatively, the plateau rate constants correlate with the steric bulk of the phosphine substituent; the largest plateau rate constant is for the ethyl substituted phosphine ligand (3.5 × 107 M–1 s–1), followed by the phenyl substituted ligand (1.7 ×

107 M–1 s–1), and the cyclohexyl substituted ligand (7.1 × 104 M–1 s–1) exhibits the lowest plateau rate constant. It is clear then that the phosphine substituents impart significant influence on the rate constants for protonation of the CoI-dxpe complexes, implying that these maximum protonation rate constants are limited by properties intrinsic to each of the cobalt complex.

3.2.5 Intrinsic Barriers – Self-Exchange

As discussed in Chapter 1, “intrinsic barriers” to protonation of reduced metal complexes to form metal hydride complexes has been discussed in the literature previously by Norton and

11,17–19 coworkers. Self-exchange rate constants for proton transfer (kself) can be used to gain insight into these intrinsic barriers. We hypothesized that trends in kself values for the cobalt complexes would track with the trends in plateau rate constants for hetero-proton transfer rate constants.

While 1H NMR techniques such as variable temperature NMR and line broadening have

15 previously been used to determine kself rate constants, no appreciable amount of broadening was observed in the Cp resonances of mixtures of the studied HCoIII and CoI complexes up to

70 °C, indicating the exchange was slower than could be measured using those techniques. We recently reported utilizing 2D-EXSY 1H NMR techniques in order to measure self-exchange rate constants for proton transfer (kself) for a class of tungsten hydride complexes, and variable temperature 2D-EXSY experiments were used to extract activation parameters for the thermoneutral self-exchange reaction.20 We therefore sought to use similar methods in order to better understand the various plateau rate constants for the cobalt complexes studied herein. 2D-

EXSY 1H NMR can measure rate constants on slower timescales (0.01-100 s–1), but has rarely been used to measure rate constants for atom-transfer reactions in general.21 For self-exchange reaction involving HCoIII and its conjugate base, CoI, kinetic information can be obtained using

appropriate mixing times and the relative integrals of the off-diagonal EXSY peaks in the 2D

NMR. To minimize the effects of longitudinal relaxation, which are rather large in these cobalt complexes (T1 < 0.3 sec), the smallest mixing time that gave appreciable off-diagonal EXSY peaks was used, typically 50-100 ms.21

Figure 3.11 Cp region of the 2D EXSY 1H NMR spectrum of HCoIII-depe and CoI-depe used to calculate the kself rate constant

1 III I 2D EXSY H NMR was used to measure the kself rate constant for HCo -dppe/Co -dppe

III I and HCo -depe/Co -depe in CD3CN:C6D6 (50:50) (example in Figure 3.11). As predicted, the

depe –1 –1 kself value for the ethyl-substituted complexes (kself = 61 M s ) was higher than that of the

dppe –1 –1 I phenyl-substituted complexes (kself = 3.1 M s ). Due to poor solubility of Co -dcpe in nearly every attempted solvent mixture, 2D 1H NMR was not able to be used to determine the

dcpe kself rate constant. Efforts are currently underway to measure this final rate constant using stopped-flow rapid mixing with UV/vis detection to obtain this value, which is expected to be

depe dppe orders of magnitude lower than that of the value of kself and kself .

It’s clear from the crystal structures of HCoIII-depe, HCoIII-dppe, and HCoIII-dcpe that the hydride ligand (and metal center in the case of the CoI species) is most easily accessible in the case of the depe ligand and the least accessible in the case of dcpe ligand. While the crystal

structure only gives us a snap-shot of the substituent orientation, previous reports of palladium complexes with chelating phosphines have calculated interior “pocket angles” () that reflect access to the metal center while allowing for free rotation of the P-C bonds.22 Although the absolute values of the reported pocket angles will not be the same in the case of our cobalt complexes due to the presence of the Cp ligand, the trend of dcpe < dppe < depe is consistent with our observed trend in reactivity towards incoming acids.

As discussed in Chapter 1, Creutz and Sutin have proposed a “weak-interaction” model of proton self-exchange with transition metal hydride complexes in which the activation barrier for proton transfer depends in large part on the probability of the appropriate orientation of the ancillary ligands for proton transfer to occur.23 The more sterically encumbered the metal center is by ancillary ligands, the higher the intrinsic barrier is for proton self-exchange, and by extension by the Marcus cross relation, the higher the barrier for protonation by exogenous acids.

Our experimental proton self-exchange rates are consistent with this type of “weak-interaction” model proposed by Creutz and Sutin. With all structural features equal in CoI-depe, CoI-dppe, and CoI-dcpe, the substituents on the phosphine group are the largest influence on the activation barrier for proton transfer, leading to lower rate constants for proton transfer (both self-exchange and protonation by molecular acids).

The Marcus cross relation has been shown to be valid for proton transfer reactions involving TMHCs.17 We therefore used the equation

푘퐴퐵 = √퐾푒푞푘퐴퐴푘퐵퐵

where kAB is the proton transfer rate constant to form the metal hydride complex (analogous to kPT), Keq is the equilibrium constant for proton transfer, kAA is the proton transfer self-exchange rate constant for organic acids used in this study (taken to be diffusion limited, or ~1 × 1010 M-1

-1 s ), and kBB is the proton transfer self-exchange rate constant for the TMHC and its conjugate base (kself discussed earlier). Comparison between simulated kPT values using the Marcus cross relation reveals excellent agreement with the LFER region of the Brønsted plots for the depe and

depe dppe dppe complexes using kself and kself (Figure 3.12). The simulated data for the dcpe

dcpe -1 -1 complexes is using an estimated kself value of 0.001 M s so that the simulated data matches

dcpe the experimental data. As discussed previously, direct measurement of kself is ongoing.

Figure 3.12 (Left) LFER region of Brønsted plot for CoI-depe, CoI-dppe, and CoI-dcpe I I compared to (right) simulated kPT values using Marcus cross relation for Co -depe, Co -dppe, CoI-dcpe.

3.3 Conclusions

In this study, we show that phosphine substituents on a series of cobalt cyclopentadienyl bisphosphine complexes greatly influences the rate constant for protonation of the reduced CoI species to form the analogous HCoIII species. Alkyl-substituted bisphosphine ligands do impart different electronic structure effects as compared to aryl-substituted bisphosphine ligands, and accounting for those differences by overlaying Brønsted plots on a relative pKa scale, we were able to compare protonation rate constant effects due to ligand sterics alone. While the overall shapes of the Brønsted plots were very similar to one another, the kPT values for the ethyl substituted complex were the largest, followed by the phenyl substituted complex, and finally the

cyclohexyl substituted complex. This indicates an intrinsic barrier to protonation that is unique to each complex, which is supported by the relative self-exchange proton transfer rate constants of the depe and dppe complexes. These findings shed light onto the ligand steric sensitivity on proton transfer rate constants to form metal hydride complexes, which supports theoretical work from Creutz et al over 30 years ago.23 Identifying and quantifying factors that influence the formation of transition metal hydride complexes such as the steric effects discussed here is vital to the guide of future catalysts for hydrogen evolution.

3.4 Experimental Details

III Synthesis of [CoCp(depe)(NCCH3)][PF6]2 (Co -depe) To a stirring solution of

[CoCp(SMe2)3][PF6]2 (202 mg, 0.34 mmol) in 5 mL of CH3CN, depe (80 µL, 0.34 mmol) was added slowly with an accompanying color change from pink to dark orange. The solution was stirred at 295 K for 48 hours, filtered, then evaporated to dryness. Dichloromethane (10 mL) was added to the residue, followed by sonication until a free flowing powder was obtained. The powder was collected, redissolved in minimal CH3CN, and the solution was added dropwise to stirred diethyl ether to obtain a bright orange powder. The orange powder was dried under reduced pressure and was obtained in 60% yield.

III Synthesis of [CoCp(dcpe)(NCCH3)][PF6]2 (Co -dcpe) To a stirring solution of

[CoCp(SMe2)3][PF6]2 (87.9 mg, 0.14 mmol) in 5 mL of CH3CN, dcpe (62 mg, 0.15 mmol) was added with an accompanying color change from pink to dark orange. The solution was stirred at

295 K overnight, after which is was filtered. The filtrate was added dropwise to diethyl ether, and the resulting solid was collected and dried via filtration. The solid was then suspended in chloroform, sonicated, filtered and washed with chloroform, then dried under high vacuum to give CoIII-dcpe as an orange solid.

III Synthesis of [CoCp(dppv)(NCCH3)][PF6]2 (Co -dppv) To a stirring solution of

[CoCp(SMe2)3][PF6]2 (200 mg, 0.33 mmol) in 10 mL of CH3CN, dppv (132 mg, 0.33 mmol) was added with an accompanying color change from pink to dark orange/brown. The solution was stirred at 295 K overnight, after which the suspension was filtered and the filtrate was dried. To the filtrate, chloroform (10 mL) was added and the suspension was sonicated until there was a free flowing powder. The solid was collected and washed thoroughly with chloroform until the filtrate ran colorless. The solids were dissolved in minimal CH3CN and added dropwise to diethyl ether. The resulting bright orange powder was collected, and dried to afford CoIII-dppv in 73% yield.

III Synthesis of [HCoCp(depe)][PF6] (HCo -depe) CoCp(CO)2 (0.25 mL, 1.87 mmol) and depe (0.44 mL, 1.88 mmol) were combined in 20 mL of toluene in a bomb flask. The bomb flask was put on the Schlenk line under an N2 atmosphere and heated at 105 °C for 30 minutes, wherein the color changed from light red to dark red with bubbling for the first 10 minutes. After the reaction flask was cooled to room temperature, the bomb flask was brought back into the glovebox and the reaction mixture was diluted with 20 mL of methanol. Solid NH4PF6 (618 mg,

3.79 mmol) was added to the mixture and it was stirred overnight, wherein the color changed from dark red to dark orange with a small amount of yellow precipitate. The entire mixture was then added to 200 mL of diethyl ether with vigorously stirring to yield a pale yellow powder. The yellow solid was filtered, rinsed with ether, and vapor diffusion of diethyl ether into a

III concentrated solution of the solid in CH3CN yielded large orange crystals of HCo -depe.

III Synthesis of [HCoCp(dcpe)][PF6] (HCo -dcpe) CoCp(CO)2 (0.23 mL, 1.72 mmol) and dcpe (0.704 mg, 1.67 mmol) were combind in 10 mL of toluene in a bomb flask. The bomb flask was put on the Schlenk line under an N2 atmosphere and heated at 100 °C for 3 hours, wherein

the color changed from light red to dark red with bubbling for the first 10 minutes. After the flask was cooled to room temperature, it was brought back into the glovebox and NH4PF6 (0.285 mg, 1.75 mmol) was added to the toluene solution, followed by 10 mL of methanol. The suspension was stirred vigorously in the glovebox for 16 hours, which resulted in the precipitation of a yellow solid. The pale yellow solid was filtered, rinsed with diethyl ether, and vapor diffusion of diethyl ether into a concentrated solution of the solid in CH3CN yielded orange crystals of HCoIII-dcpe.

III Synthesis of [HCoCp(dppv)][PF6] (HCo -dppv) CoCp(CO)2 (0.10 mL, 0.75 mmol) and dppv (0.304 mg, 0.77 mmol) were combined in 10 mL of toluene in a bomb flask. The bomb flask was put on the Schlenk line under an N2 atmosphere and heated at 100 °C for 3 hours, wherein the color changed from light red to dark red with bubbling for the first 10 minutes. After the flask was cooled to room temperature, it was brought back into the glovebox and NH4PF6

(0.263 mg, 1.75 mmol) was added to the toluene solution, followed by 10 mL of 1-propanol. The suspension was stirred vigorously in the glovebox for 1 hours, which resulted in the precipitation of a yellow solid. The pale yellow solid was filtered, rinsed with diethyl ether, and vapor diffusion of diethyl ether into a concentrated solution of the solid in CH3CN yielded orange crystals of HCoIII-dppv in 74% yield.

General procedure for generation of CoI-dxpx species: To a suspension/solution of

HCoIII-dxpx in THF, an excess of NaH (10-100 equivalents) was added. The suspension was stirred overnight at room temperature in the glovebox, which was accompanied by a color change from yellow to red. The mixture was dried under vacuum, and re-suspended in toluene.

The mixture was stirred overnight, followed by filtration to get rid of excess NaH, unreacted

III I HCo -dxpx, and NaPF6. The filtrate was then dried under vacuum to afford the Co -dxpx product.

REFERENCES

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(2) Dempsey, J. L.; Brunschwig, B. S.; Winkler, J. R.; Gray, H. B. Hydrogen Evolution Catalyzed by Cobaloximes. Acc. Chem. Res. 2009, 42 (12), 1995–2004.

(3) Helm, M. L.; Stewart, M. P.; Bullock, R. M.; DuBois, M. R.; DuBois, D. L. A Synthetic -1 Nickel Electrocatalyst with a Turnover Frequency Above 100,000 s for H2 Production. Science (80-. ). 2011, 333 (6044), 863–866.

(4) Bediako, D. K.; Solis, B. H.; Dogutan, D. K.; Roubelakis, M. M.; Maher, A. G.; Lee, C. H.; Chambers, M. B.; Hammes-Schiffer, S.; Nocera, D. G. Role of Pendant Proton Relays and Proton-Coupled Electron Transfer on the Hydrogen Evolution Reaction by Nickel Hangman Porphyrins. Proc. Natl. Acad. Sci. 2014, 111 (42), 15001–15006.

(5) Dempsey, J. L. Proton-Coupled Electron Transfer: Metal Hydrides Find the Sweet Spot. Nat. Chem. 2015, 7 (2), 101–102.

(6) Connor, G. P.; Mayer, K. J.; Tribble, C. S.; McNamara, W. R. Hydrogen Evolution Catalyzed by an Iron Polypyridyl Complex in Aqueous Solutions. Inorg. Chem. 2014, 53 (11), 5408–5410.

(7) Lewandowska-Andralojc, A.; Baine, T.; Zhao, X.; Muckerman, J. T.; Fujita, E.; Polyansky, D. E. Mechanistic Studies of Hydrogen Evolution in Aqueous Solution Catalyzed by a Tertpyridine-Amine Cobalt Complex. Inorg. Chem. 2015, 54 (9), 4310– 4321.

(8) Elgrishi, N.; Rountree, K. J.; McCarthy, B. D.; Rountree, E. S.; Eisenhart, T. T.; Dempsey, J. L. A Practical Beginner’s Guide to Cyclic Voltammetry. J. Chem. Educ. 2017, acs.jchemed.7b00361.

(9) Elgrishi, N.; McCarthy, B. D.; Rountree, E. S.; Dempsey, J. L. Reaction Pathways of Hydrogen-Evolving Electrocatalysts: Electrochemical and Spectroscopic Studies of Proton-Coupled Electron Transfer Processes. ACS Catal. 2016, 6 (6), 3644–3659.

(10) Elgrishi, N.; Kurtz, D. A.; Dempsey, J. L. Reaction Parameters Influencing Cobalt Hydride Formation Kinetics: Implications for Benchmarking H2-Evolution Catalysts. J. Am. Chem. Soc. 2017, 139 (1), 239–244.

(11) Weberg, R. T.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition- Metal Complexes. 6. Interstitial Hydrides. J. Am. Chem. Soc. 1990, 112 (3), 1105–1108.

(12) Walker, H. W.; Kresge, C. T.; Ford, P. C.; Pearson, R. G. Rates of Deprotonation and PKa Values of Transition Metal Carbonyl Hydrides. J. Am. Chem. Soc. 1979, 101 (24), 7428– 7429.

(13) Kresge, A. J. What Makes Proton Transfer Fast. Acc. Chem. Res. 1975, 8 (10), 354–360.

(14) Kramarz, K. W.; Norton, J. R. Slow Proton-Transfer Reactions in Organometallic and Bioinorganic Chemistry. In Progress in Inorganic Chemistry, VOl. 42; 2007; pp 1–65.

(15) Edidin, R. T.; Sullivan, J. M.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition-Metal Complexes. 4. Kinetic Acidities toward Aniline and Their Use in Identifying Proton-Transfer Mechanisms. J. Am. Chem. Soc. 1987, 109 (13), 3945– 3953.

(16) Berning, D. E.; Miedaner, A.; Curtis, C. J.; Noll, B. C.; Rakowski DuBois, M. C.; DuBois, D. L. Free-Energy Relationships between the Proton and Hydride Donor Abilities of + [HNi(Diphosphine)2] Complexes and the Half-Wave Potentials of Their Conjugate Bases. Organometallics 2001, 20 (9), 1832–1839.

(17) Kristjansdottir, S. S.; Loendorf, A. J.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition-Metal Complexes. 9. A Sterically Hindered Cationic Hydride, + [H4Re(PMe2Ph)4] . Inorg. Chem. 1991, 30 (23), 4470–4471.

(18) Jordan, R. F.; Norton, J. R. Kinetic and Thermodynamic Acidity of Hydrido Transition- Metal Complexes. 1. Periodic Trends in Group VI Complexes and Substituent Effects in Osmium Complexes. J. Am. Chem. Soc. 1982, 104 (5), 1255–1263.

(19) Jordan, R. F.; Norton, J. R. Proton-Transfer Reactions in Organometallic Chemistry; 1982; Vol. 80521, pp 403–423.

(20) Huang, T.; Rountree, E. S.; Traywick, A.; Bayoumi, M.; Dempsey, J. L. Switching between Stepwise and Concerted Proton-Coupled Electron Transfer Pathways in Tungsten Hydride Activation. Submitted.

(21) Nikitin, K.; O’Gara, R. Mechanisms and Beyond: Elucidation of Fluxional Dynamics by Exchange NMR Spectroscopy. Chem. - A Eur. J. 2019, 1–40.

(22) Koide, Y.; Bott, S. G.; Barron, A. R. Alumoxanes as Cocatalysts in the Palladium- Catalyzed Copolymerization of Carbon Monoxide and Ethylene: Genesis of a Structure−Activity Relationship. Organometallics 1996, 15 (9), 2213–2226.

(23) Creutz, C.; Sutin, N. Intrinsic Barriers to Proton Exchange between Transition-Metal Centers: Application of a Weak-Interaction Model. J. Am. Chem. Soc. 1988, 110 (8), 2418–2427.

CHAPTER 4: BATHOCHROMIC SHIFTS IN RHENIUM CARBONYL DYES INDUCED THROUGH DESTABILIZATION OF OCCUPIED ORBITALS

Chapter adapted with permission from Kurtz, D. A.; Brereton, K. R.; Ruoff, K. P.; Tang, H. M.; Felton, G. A. N.; Miller, A. J. M.; Dempsey, J. L. Bathochromic Shifts in Rhenium Carbonyl Dyes Induced through Destabilization of Occupied Orbitals. Inorg. Chem. 2018, 57 (9), 5389– 5399. Copyright (2018) American Chemical Society.

4.1 Introduction

Rhenium(I) carbonyl complexes are versatile dyes (chromophores) that have been

1−5 extensively used in applications such as photosensitization of light-driven CO2 reduction,

long-range electron transfer through proteins,6−10 and singlet oxygen generation.11,12 Most

rhenium dyes do not exhibit strong light absorption across the visible spectrum, with transitions

from filled Re d-orbitals to empty ligand π* orbitals generally requiring ultraviolet or blue light.

Efforts to red-shift the absorbance of rhenium carbonyl complexes have focused on modifying

the diimine ligand with electron-withdrawing groups or introducing additional conjugation,

which lowers the energy of the metal-to-ligand charge transfer (MLCT) transition by stabilizing

the π* orbitals (Figure 4.1A).13,14 These modifications can have moderate effects on the energy

of the relevant MLCT transition, such as a red-shift of ∼100 nm when switching from -NH2

15 groups to -NO2 groups on the 4 and 4′ positions on bipyridine. Although extensive synthetic

modification can be used to access bidentate ligands with widely varied electronic properties,

there are practical limits to the extent of LUMO stabilization before the ligand will no longer

effectively bind the metal center.

Figure 4.1 Simplified orbital energy diagram illustrating the strategy of red-shifting absorption by A) stabilizing * LUMO levels with electron withdrawing groups or B) by destabilizing HOMO levels with modification of axial ligands.

An alternative, underexplored strategy to tune the absorption profile of rhenium chromophores is to modify the axial ligands cis to the diimine ligand in order to destabilize the d-orbitals (Figure 4.1B). In the parent complex, the occupied d-character orbitals are highly stabilized through backbonding interactions with the facially coordinated carbonyl ligands;16 destabilization of these orbitals thus lowers the energy of the MLCT transition. While previous reports have demonstrated that the axial ligand identities affect the energy of the MLCT transitions,17 a detailed study of the relationship between ligand electronic attributes and experimental optical properties has not been reported. We report herein a hybrid approach to absorbance tuning of rhenium dyes involving destabilization of the metal d-orbitals and stabilization of the bipyridine π* orbitals. A series of ester-substituted bipyridine rhenium

− complexes incorporating axial ligands with different donating strength (CO, CH3CN, Cl , and

PMe3) provides an opportunity to correlate changes in ligand bonding properties with the electronic structure of the complex.

The ester-substituted bipyridine ligand, in addition to facilitating attachment of these complexes on metal oxide thin films in future studies, lowers the π* orbital energy to redshift the absorbance profile.18,19 Through electrochemical and spectroscopic measurements, we quantified

the relationship between the ligand donor strength and the HOMO/LUMO energetics. We utilized computational methods to further elucidate the electronic structure of these complexes and rationalize the bathochromic shifts in the MLCT transitions as a function of axial ligand identity.

4.2 Results

4.2.1 Synthesis

The syntheses of the four target complexes are summarized in Scheme 4.1.

[Re(deeb)(CO)3(NCCH3)][PF6] (1, deeb = 4,4′-diethylester-2,2′-bipyridine) was obtained as a

15 yellow solid via halide abstraction from previously reported Re(deeb)(CO)3Cl (2) using silver triflate (AgOTf) in CH3CN, followed by salt metathesis with ammonium hexafluorophosphate

([NH4][PF6]) (84% yield). While selective displacement of a single carbonyl ligand in rhenium(I) tricarbonyl complexes is often thermodynamically unfavorable,5,20 the presence of a tertiary phosphine ligand trans to a carbonyl ligand can enable substitution via photodissociation or oxidative decarbonylation.21−23 A suitable tertiary phosphine complex intermediate,

[Re(deeb)(CO)3(PMe3)][PF6], was accessed in 76% yield by halide abstraction from 2 in acetone followed by addition of excess PMe3 and salt metathesis with [NH4][PF6]. When an orange solution of [Re(deeb)(CO)3(PMe3)][PF6] in CH3CN was treated with 1 equiv of trimethylamine

N-oxide (TMNO), a color change to deep red was observed within minutes. In situ 1H NMR spectroscopy (Figure 4.2 and Figure 4.3) indicated decarbonylation followed by coordination of

CH3CN. The solution was then refluxed under N2 overnight to help avoid formation of an unwanted trimethylamine complex.22

1 Figure 4.2 H NMR spectrum of mixture of [Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] and [Re(deeb)(CO)2(PMe3)(NMe3)][PF6] taken after allowing [Re(deeb)(CO)3(PMe3)][PF6] and TMNO to react at room temperature.

Workup and recrystallization provided [Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] (3) as a deep red solid in 60% yield. The neutral dicarbonyl complex Re(deeb)(CO)2(PMe3)Cl (4) was obtained by ligand substitution of cationic 3. Excess benzyltrimethylammonium chloride ([BnMe3N]Cl) was added to a solution of 3 in CH2Cl2, which was then refluxed for 24 h. Removal of the excess salts by an aqueous workup and precipitation with pentane yielded 4 in 65% yield as a dark purple powder.

4.2.2 Structural Characterization in Solution and Solid State

The structures of the four new complexes reported here were characterized by multinuclear NMR spectroscopy, infrared (IR) spectroscopy, and single-crystal X-ray diffraction.

NMR spectroscopy in CD3CN provides insight into the solution symmetry and geometry. IR spectroscopy reports on the number of carbonyl ligands and provides a measure of the electron-

richness of the metal center. IR spectra were recorded in CH2Cl2, which has fewer solvent

-1 1 features in the region of interest (2100 − 1800 cm ) than CH3CN. The H NMR spectrum of 1

(Figure 4.4) features three aromatic resonances and two aliphatic resonances of the deeb ligand, consistent with a Cs-symmetric complex. The bound CH3CN ligand is initially observed in both the 1H NMR (δ 2.02) and in the 13C NMR (Figure 4.5, δ 123.8, 3.9) spectra; however, within 7 days of dissolution in CD3CN, the bound CH3CN ligand resonance diminishes in intensity as a resonance that corresponds to free CH3CN grows in (Figure 4.6). This indicates a slow ligand substitution reaction in which the large excess of CD3CN slowly replaces the bound CH3CN.

1 Figure 4.4 H NMR spectrum of 1 in CD3CN.

13 Figure 4.5 C NMR spectrum of 1 in CD3CN.

The carbonyl region in the IR spectrum of 1 (Figure 4.8) shows one sharp absorption peak and one broad absorption peak attributed to two overlapping stretches, for a total of three

CO stretching vibrational modes with A′(1), A″, and A′(2) symmetry. The 1H NMR spectrum of chloride tricarbonyl 2 (Figure 4.7) indicates retention of Cs symmetry.

1 Figure 4.7 H NMR spectrum of 2 in CD3CN.

The IR spectrum of 2 exhibits three distinct peaks for the same three CO vibrational modes as 1 (Figure 4.8). The CO stretches of 2 are shifted to lower energy than those in 1, reflecting the reduced overall charge of the complex and increased electron density at the metal center.

Figure 4.8 Carbonyl stretching region of IR spectra of complexes 1-4, collected in CH2Cl2.

The 1H NMR spectrum of complex 3 (Figure 4.9) exhibits a doublet for the methyl

2 protons of the PMe3 ligand (δ 1.15, JPH = 9.5 Hz) along with the aromatic resonances assigned

4 to the symmetric deeb ligand. Additionally, there is a JPH coupling of 0.7 Hz between PMe3 ligand and the two protons ortho to the nitrogen on the deeb ligand. Resonances for the bound

1 13 CH3CN ligand are seen in both the H NMR (δ 2.04) and the C NMR (Figure 4.10, δ 124.7,

3.9) spectra. A single CO resonance in the 13C NMR appears as a doublet weakly coupled to the

2 PMe3 ligand ( JPC = 7.2 Hz), confirming overall Cs symmetry with a cis-phosphinocarbonyl configuration.

1 Figure 4.9 H NMR spectrum of 3 in CD3CN.

13 1 Figure 4.10 C{ H} NMR spectrum of 3 in CD3CN.

The IR spectrum of 3 shows two peaks in the carbonyl region for the cis CO ligands, assigned to a symmetric stretching mode with A′ symmetry and an asymmetric stretching mode with A″ symmetry. The CO stretches of 3 shift to lower energy relative to 1 and 2, as expected for a dicarbonyl complex with enhanced backbonding (due to less competition).

1 The H NMR spectrum of complex 4 (Figure 4.11) supports assignment of Cs symmetry.

2 Phosphorus−carbon coupling is observed between the carbonyl ligands and PMe3 ( JPC = 7.8

Hz), and a long-range phosphorus−hydrogen coupling is observed between PMe3 and the protons

4 ortho to the nitrogen on the deeb ligand ( JPH = 0.7 Hz). The IR spectrum of 4 exhibits the same pattern of peaks as 3, with both vibrational modes lower in energy due to both the charge neutrality and substitution by the more donating Cl− ligand.

1 Figure 4.11 H NMR spectrum of 4 in CD3CN.

Dark purple block-shaped crystals of 4 suitable for X-ray diffraction study were grown from vapor diffusion of pentane into an acetone solution of 4. The solid-state structure of complex 4 (Figure 4.12) confirms the expected pseudo-Cs symmetry with a trans configuration of the chloride and PMe3 ligands. The structure and bond lengths are in line with previously reported complexes of similar geometry.20,22,24,25

4.2.3 Photophysical Characterization

The electronic absorption and emission spectra of rhenium dyes 1-4 are shown in Figure

4.13. The UV/vis absorption spectra of 1 and 2 in CH3CN solution are dominated by a broad

MLCT band (λmax = 365 nm for 1 and 410 nm for 2), typical of many reported rhenium(I) tricarbonyl complexes.15 In contrast, the absorption spectra of 3 and 4 each exhibit two features of similar intensity in the visible region (λmax = 365 and 465 nm for 3 and 406 and 533 nm for 4).

In addition to these two major transitions, a distinct, low energy shoulder is observed in both spectra. The absorbance features shift bathochromically from complex 1 to 4. The red shift leads

−1 −1 −1 complex 4 to absorb across the entire visible spectrum (ε406 = 4710 M cm , ε533 = 5440 M cm−1, Table 4.2). Full spectral assignments and factors contributing to this impressive absorption profile are discussed below. Such significant bathochromic shifts are rarely observed in rhenium(I) bipyridine carbonyl complexes.20,26

- M

Figure 4.13 Absorption spectra of complexes 1-4 (solid lines) in CH3CN at 295 K and emission spectra of complexes 1-3 (dashed lines) in 90/10 2-MeTHF/CH3CN) at 77 K. For emission spectra, exc = 365 nm for complexes 1 and 2 and exc = 455 nm for complex 3.

Photoluminescence spectra were recorded in two different solvents for complexes 1, 2, and 3. (No emission was observed for 4 at 295 or 77 K.) Complexes 1 and 2 both display strong emission in CH3CN at 295 K (Figure 4.14 and Figure 4.15), while 3 displays extremely weak emission under the same conditions. Photoluminescence measurements at 77 K were recorded in a 90/10 2-MeTHF/CH3CN solution (2-MeTHF = 2-methyltetrahydrofuran) that combines the ideal glassing properties of 2-MeTHF with the solubilizing properties of acetonitrile. The emission maxima blue-shift in frozen 2-MeTHF/CH3CN compared to the spectra at 295 K in either CH3CN or 2-MeTHF/CH3CN (Figure 4.14 and Figure 4.15) as predicted by the rigidochromic effect.27 The emission maximum shifts to longer wavelength moving from complex 1 to 3 (Figure 4.13).

Figure 4.14 Emission spectrum of 1 at RT in CH3CN (exc = 405 nm), and in 90:10 2- MeTHF:CH3CN at RT and at 77 K (exc = 365 nm)

Figure 4.15 Emission spectrum of 2 at RT in CH3CN, and in 90:10 2-MeTHF:CH3CN at RT and at 77 K (exc = 365 nm)

The excited-state lifetimes (τ) were measured using time-resolved photoluminescence spectroscopy at 295 K in CH3CN (1 and 2, Figure 4.16 and Figure 4.17, respectively) and at 77

K in frozen 90/10 2-MeTHF/CH3CN (1, 2, and 3, Figure 4.16, Figure 4.17, and Figure 4.18, respectively). The lifetimes for complexes 1 (τ = 150 ns) and 2 (τ = 11 ns) at 295 K are similar to

n+ − − − − previously reported complexes of the type [Re(deeb)(CO)3(L)] (L = CN , Cl , Br , I , n = 0; L

= pyridine, n = 1).15 At 77 K, the lifetimes of 1 and 2 increase by at least an order of magnitude

(Figure 4.16), and comparisons to 3 (τ = 0.41 μs) become possible (Table 4.2, below). The lifetimes at 77 K decrease as the complex emission maximum red-shifts.

Figure 4.16 Time-resolved photoluminescence measurements of 1 at 295 K (left) and at 77 K (right). Measurements taken in 90/10 2-MeTHF/CH3CN.

Figure 4.17 Time-resolved photoluminescence measurements of 2 at 295 K (left) and at 77 K (right). Measurements taken in 90/10 2-MeTHF/CH3CN.

Figure 4.18 Time-resolved photoluminescence measurements of 3 at 77 K. Measurement taken in 90/10 2-MeTHF/CH3CN.

4.2.4 Electrochemical Studies

Cyclic voltammetry was used to investigate the ground-state redox properties of 1−4.

+ Each complex displays a one-electron reduction (E°′red) between -1.2 and -1.5 V vs Fc /Fc in

0.25 M TBAPF6 in CH3CN. The reduced species is proposed to involve electron delocalization onto the deeb ligand, in accordance with prior assignments of ligand-centered reductions in related complexes. Moving from complex 1 to 4, E°′red occurs at more negative potentials (Table

4.3, below). Cyclic voltammograms of 1−4 also contain a one-electron oxidation feature (E°′ox) assigned to a rhenium-based oxidation (Figure 4.19 and Table 4.3, below). The cathodic shift of

E°′ox (1.13 V difference between 1 and 4) across the four complexes is substantially larger in magnitude than the shifts in deeb-based E°′red (0.25 V shift from 1 to 4).

Figure 4.19 Cyclic voltammograms of 1 mM of 1-4 in 0.25 M TBAPF6 in CH3CN. For complexes 2-4, 1 mM Fc was added and the internal reference Fc+/Fc couple (0 V) is shown with a dashed box. Due to an interaction with oxidized 1 and Fc+, ferrocene was not added to the solution when scanning oxidatively; the reduction feature was referenced to Fc+/Fc in a separate experiment. All scans were recorded using a platinum working electrode, glassy carbon counter electrode, and Ag wire pseudo-reference electrode, and 3 V/s scan rate.

4.2.5 Computational Studies

Density Functional Theory (DFT) was employed to further probe the structural and spectroscopic features of complexes 1−4. The optimal computational protocol for the ground- state singlet (S0) geometries was developed for complex 4 by comparing three different functionals (M06, B3LYP, and PBE), each coupled with the LANL2DZ-ECP basis set for Re and 6-31G* for all other atoms. These three functionals were chosen based on literature precedent for calculations on similar rhenium dyes.28,29 All calculations were performed using an implicit SMD solvation model. The SMD model for dichloromethane was used for the calculation of IR stretching frequencies in order to compare with experimental spectra. Similarly, the SMD model for acetonitrile was used for Time-Dependent Density Functional Theory (TD-

DFT) calculations to best reflect experimental spectra collected in that solvent. The functional

B3LYP provided electronic transition energies for complexes 1 and 4 that agreed most closely with the experimentally observed UV/vis absorption spectra in acetonitrile (see below). The

B3LYP geometry optimization of complex 4 also converged at a geometry that closely resembles the crystallographic parameters, so this functional was used going forward. The vibrational frequencies for complexes 1−4 were calculated using a dichloromethane implicit solvation model. Table 4.1 shows the experimental values for the carbonyl stretching frequencies alongside calculated values that have been corrected using the reported empirical scaling factor of 0.9614 for B3LYP.30,31 All of the experimental trends are nicely reproduced in the calculated

CO stretching frequencies.

Table 4.1 Experimental and calculated vibrational frequencies for complexes 1-4.

IR stretches (in CH Cl , cm-1) Complex 2 2 Experimental Calculated 1 2042, 1945a 2026, 1937, 1932 2 2025, 1926, 1904 2009, 1917, 1900 3 1943, 1873 1930, 1867 4 1920, 1848 1912, 1846 aTwo overlapping stretches gives rise to single, broad absorption

TD-DFT was used to calculate the first 50 excitations for complexes 1−4 using the

B3LYP functional in acetonitrile solvent. Figure 4.20 shows the calculated excitation energies

(black vertical lines) and depicts the HOMO−1 and LUMO orbitals for all four complexes. The oscillator strength (right axis) calculated for a transition represents the probability of the occurrence of that transition and is proportional to the molar extinction coefficient (left axis). For all complexes, the dominant low-energy feature observed in the experimental spectra (colored lines) can be attributed to excitation from the Re-based HOMO−1 to the bipyridyl π*-based

LUMO. For complex 1, excitation from the HOMO−2 to the LUMO also contributes to the

dominant absorption feature. The analogous transition has a negligible oscillator strength in complexes 2−4. Less intense transitions from the HOMO (another Re-centered d-character orbital) to the LUMO account for the shoulder observed on the low-energy edge of the absorption spectra.

4.3 Discussion

The present series of complexes can be compared in pairs organized by apical ligand: acetonitrile vs chloride (1 vs 2, and 3 vs 4), and carbonyl vs PMe3 (1 vs 3, and 2 vs 4). These different ligands lead to relatively large spectroscopic and electrochemical changes.

4.3.1 Redox properties and Electrochemically Induced Reactivity

The deeb-based reduction features for complexes 1−3 are reversible at all scan rates studied (0.025 - 50 V/s, Figure 4.21 for 1 and 2, Figure 4.22 left for 3). For complex 4, however, the primary reduction is not chemically reversible at slow scan rates, only becoming fully reversible at scan rates above ∼0.5 V/s (Figure 4.22 right). The anodic feature observed at fast scan rates diminishes as the scan rate is lowered and a new oxidation peak grows in. The new anodic feature aligns with the reoxidation feature of reduced 3 (Figure 4.22 left). This observation suggests that chloride dissociation occurs after the one-electron reduction of 4, generating the CH3CN-ligated species 3. This EC mechanism stands in contrast to the reported reactivity of Re(dmbpy)(CO)3Cl (dmbpy = 4,4′-dimethyl-2,2′-bipyridine), which only dissociates chloride on this time scale upon a two electron reduction.32 Indeed, we also observe no evidence for chloride loss upon one-electron reduction of the tricarbonyl chloride complex 2. Therefore,

• our observation of oxidation of [Re(deeb)(CO)2(PMe3)(NCCH3)] on the return trace after the one-electron reduction of 4 indicates that chloride loss from reduced 4 is more facile and suggests that replacing the π-accepting CO in 2 with the σ-donating PMe3 ligand in 4 leads to a more electron rich metal center that facilitates chloride dissociation upon reduction.

The oxidation of 1 is chemically reversible at all scan rates studied, but electrochemically quasi-reversible (Figure 4.23), as indicated by the peak-to-peak separation, which increases from

100 mV at a scan rate of 25 mV/s to 360 mV at 50 V/s. The quasi-reversibility is attributed to a relatively slow heterogeneous electron transfer rate constant for the ReII/I couple. Interestingly,

when an equimolar amount of Fc is present in solution, current passed at the onset of the ReII/I couple (∼1.4 V) increases, and the waveform becomes irreversible (Figure 4.24). In a separate experiment, an irreversible oxidation of ferrocene was detected with an onset of ∼1.6 V, presumably decomposing upon oxidation in accordance to previous reports in other solvents

33 34 (AlCl3/1-butylpyridinium chloride mixtures and SO2 ). The catalytic current observed for mixtures of 1 and Fc suggest that 1+ mediates the oxidative decomposition of Fc+.

The oxidation of 2 (E°′ox ≈ 1.03 V) approaches reversibility at very fast scan rates (∼100

V/s, Figure 4.23), but is chemically irreversible at slower scan rates. Upon oxidation to the

Re(II) species, a number of tricarbonyl chloride rhenium complexes have been reported to undergo oxidative disproportionation via a chloride-bridged inner-sphere electron transfer mechanism; similar reactivity likely gives rise to the irreversibility observed for 2 at lower scan rates.35

The oxidation of 3 at 0.76 V is chemically reversible at all scan rates studied. Similar to

1, it is electrochemically quasi-reversible at higher scan rates, likely due to a slow heterogeneous electron transfer rate constant (Figure 4.25).

Interestingly, the oxidation of chloride complex 4 is chemically reversible at all scan rates studied, and the redox couple shows no increased peak-to-peak separation at fast scan rates

(Figure 4.25). This is in contrast to the established follow-up chemical steps noted above for rhenium chloride complexes, including 2. This differing reactivity suggests that the steric hindrance of the PMe3 may prevent the formation of the chloride-bridged dimer needed for inner sphere electron transfer. Alternatively, as the ReII/I reduction potential of 4 is 700 mV more cathodic than that of 2, the oxidized species may be more stable and not as reactive toward the aforementioned inner sphere disproportionation pathway.

4.3.2 Photophysics and Excited-State Redox Properties

The Stokes shift and luminescence lifetimes for these rhenium complexes (Table 4.2) are consistent with prior assignments for related complexes of the emissive excited state as a MLCT triplet state.20,36 The potentials for oxidation and reduction of the excited state can be calculated using a thermochemical cycle involving the ground-state reduction potentials and the excited- state energy, ΔGST, defined here as the energy difference between the ground vibrational states of the singlet ground state and the triplet excited state. Values of ΔGST were obtained for complexes 1, 2, and 3 based on the x-intercept of a linear, tangential fit to the high-energy side of the low-temperature photoluminescence spectra (Figure 4.26, Figure 4.27, and Figure 4.28,

37−40 respectively). The measured ΔGST values of 2.44, 2.20, and 1.76 eV for 1, 2, and 3, respectively, are consistent with a decrease in the energy of the HOMO/LUMO gap as the substituents become more donating (Table 4.2). A linear correlation between the measured ΔGST values and the lowest energy absorption λmax values is observed for complexes 1−3 (Figure

4.29). Although no emission was observed for complex 4, a ΔGST value can be predicted by

extrapolation from the linear correlation of Figure 4.29 (1.80 eV). TD-DFT calculated ΔGST values exhibit the same trends within the series of complexes, with reasonable 0.2 − 0.4 eV agreement with the experimental values.

Figure 4.26 Emission spectrum of 1 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 365 nm) and tangential fit to the high energy side to estimate GST.

Figure 4.27 Emission spectrum of 2 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 365 nm) and tangential fit to the high energy side to estimate GST.

Figure 4.28 Emission spectrum of 3 at 77 K in 90:10 2-MeTHF:CH3CN (exc = 455 nm) and tangential fit to the high energy side to estimate GST.

Figure 4.29 Linear correlation between lowest energy max experimentally determined GST values of 1-3 used to estimate GST of 4.

Table 4.2 Photophysical properties of complexes 1-4.

GST (eV)  (ns) max (nm) Complex  Expt. Calc. 295 K 77 K Abs. (, M-1 cm-1) Emission (77K) r 1 2.62 2.44 150 4600 365 (4750) 525 0.049 2 2.35 2.20 13 2700 410 (4820) 585 0.013b 365 (5270) 3 2.05 1.76 – 410 663 – 465 (6150) 406 (4710) 4 1.80a 1.42 – – – – 533 (5440) aValue extrapolated, see Figure 4.8

bValue taken from Reference 18

The excited-state photophysical data and electrochemical properties were used to calculate the excited-state redox properties (eqs 1 and 2), summarized in Table 4.3.41 These values indicate that these complexes are potent excited-state reductants and oxidants. Whereas the ReII/ReI* potential remains relatively constant for 1−4, the ReI*/ReI(deeb•−) potential varies by more than 1 V as a function of the ancillary ligand, with complex 1 being an extremely potent photooxidant. Excited-state quenching studies are the focus of current work on the photochemistry of these complexes.

I* I  I I  Eʹ(Re /Re (deeb )) = Eʹ(Re /Re (deeb )) + GST Eq. 1

II I* II I Eʹ(Re /Re ) = Eʹ(Re /Re )  GST Eq. 2

Table 4.3 Summary of ground state and excited state redox properties of complexes 1-4.

Eʹ (V vs. Fc+/Fc) Eʹ* (V vs. Fc+/Fc) Complex ReI/ReI(deeb) ReII/I ReI*/Re I(deeb) ReII/ReI* 1 –1.22 1.45 1.40 –1.17 2 –1.34 1.03 1.01 –1.32 3 –1.33 0.76 0.72 –1.29 4 –1.47 0.30 0.33 –1.50

4.3.3 Influence of Axial Ligands on Electronic Structure

The spectroscopic and electrochemical data presented above clearly demonstrate that the axial ligands strongly influence the ground-state electronic structure of the complexes.

Computational methods provide insight into the impact of PMe3 ligation on electronic structure.

A detailed orbital energy diagram can be built from the calculated transitions and the relative energies of orbitals relevant to the visible transitions. All the transitions discussed below are comprised of charge transfers from Re d-character orbitals (HOMO, HOMO−1, and HOMO−2) to π*-character orbitals on the deeb ligand (LUMO and LUMO+1). The TD-DFT calculations

reveal that the visible region transitions with the highest oscillator strength for all complexes involve electronic excitation from the HOMO−1 to the LUMO. The HOMO → LUMO transitions are calculated to be at slightly lower energy than this prominent band and with much lower oscillator strengths, corresponding to small shoulders on the low-energy side of the main absorbance bands of all four complexes (Figure 4.13).

The main absorption band in the UV/vis absorption spectrum of tricarbonyl chloride complex 2 (λmax = 410 nm) is red-shifted with respect to the analogous absorption band in the spectrum of tricarbonyl acetonitrile complex 1 (λmax = 365 nm), consistent with other studies comparing analogous neutral chloride and cationic acetonitrile complexes.15 The magnitude of the red-shift (∼0.4 eV) is reproduced nicely by TD-DFT. The energetic basis for the red-shift is found in ground-state DFT calculations that show a 0.4 eV destabilization of the HOMO−1 levels and only 0.1 eV change in energy of the LUMO orbitals upon replacing chloride with acetonitrile. The π-donation by the Cl− ligand into the metal-based orbitals in 2 (absent in 1), is thought to be the origin of the HOMO−1 destabilization. The same acetonitrile/chloride substitution leads to the absorbance features of 4 being red-shifted with respect to those of 3 by

∼0.35 eV. DFT calculations again indicate a 0.4 eV destabilization of HOMO−1 levels and negligible changes to the LUMO energy upon replacing acetonitrile in 3 with chloride in 4.

The substitution of a CO ligand for a PMe3 ligand also impacts the electronic structure properties, as evidenced by the absorbance differences observed for 1 vs 3 and 2 vs 4. The UV/ vis absorption spectrum of 3 is significantly red-shifted from that of 1. Replacing the strongly π- acidic CO ligand with a PMe3 ligand destabilizes the Re d-orbitals significantly, leading to a 0.8 eV increase in the HOMO−1 level of 3 relative to 1. This HOMO−1 destabilization is consistent with the ∼0.75 eV difference in the lower energy λmax values in the UV/vis absorption spectra of

1 and 3. Similarly, a 0.6 eV HOMO−1 destabilization is calculated upon moving from 2 to 4, corresponding nicely to the ∼0.7 eV difference between the lower energy λmax values. The combined data provides an explanation for the improved red light absorption properties of 3 and

4 relative to 1 and 2: as shown in Figure 4.30 (left), Re d-orbital destabilization leads to significant perturbation of the HOMO, HOMO−1, and HOMO−2 energies across the series 1-4, while the LUMO and LUMO−1 energies remain about the same.

Another interesting observation gleaned from the orbital energies displayed in Figure

4.30 (left) is the differing extent of destabilization of the individual d-orbitals across the series of complexes. The three lowest-energy d-orbitals in complex 2 are all (a) lower in energy than the corresponding orbitals in complex 4 and (b) closer in energy to each other than the same orbitals in complex 4. The HOMO and HOMO−1 have appropriate symmetry to engage in π- backbonding with all three CO ligands in complexes 1 and 2 (2 in Figure 4.30, right). Therefore, when a phosphine ligand is substituted for the CO, the amount of π-backbonding is substantially decreased and the HOMO and HOMO−1 are not stabilized to the same extent in 4 as they are in

2. In contrast, the symmetry of HOMO−2 only allows for backbonding interactions with the two equatorial CO ligands. The HOMO−2 is thus less perturbed when the axial CO ligand is replaced with PMe3, giving rise to a larger net spacing between HOMO and HOMO−2 in 4.

The molecular orbital picture of Figure 4.30 also helps rationalize the electrochemical properties of complexes 1-4 (Table 4.3). The metal-based d-orbitals are strongly affected by ligand substitution, so the ReII/I potentials are dramatically affected by ligand substitution. On the other hand, the LUMO is not greatly affected by ligand substitution, leading to the deeb-based reduction potentials remaining roughly the same over the series.

4.4 Conclusion

A series of rhenium diimine carbonyl complexes with axial carbonyl, chloride, acetonitrile, or phosphine ligands was prepared and characterized by multinuclear NMR spectroscopy, infrared spectroscopy, and X-ray diffraction crystallography (for complex 4).

Cyclic voltammetry was used to determine both the one-electron, metal-based oxidation potentials and one-electron, deeb-based reduction potentials of each complex. The axial ligand is observed to have a greater influence on the oxidation feature than the reduction feature. The electronic absorption spectra of the complexes show an increased visible region absorption profile as acetonitrile is substituted by chloride and as carbonyl is substituted by PMe3. These bathochromic shifts are also apparent in the photoluminescence spectra of the complexes and are accompanied by a decrease in excited-state lifetime. Computational methods were used to assign each of the MLCT transitions observed in the UV/vis absorption spectra. These DFT results reveal that the aforementioned red-shifting in electronic absorption and emission spectra is a consequence of destabilization of Re d-orbitals. The destabilization could arise from changes in

donor properties, such as the substitution of a strongly σ-donating phosphine for a π-accepting carbonyl ligand, or from changes in overall charge leading to energy lowering of the d-orbitals.

While rhenium carbonyl complexes are generally used as photo-oxidants, the excited- state reduction potential of 4 (E°′*(ReII/ReI*) = −1.5 V vs Fc+/Fc) and its broad absorption profile in the visible region provides potential avenues for new reactivity. Collectively, the excited-state and ground-state redox properties and wide spectral differences across this series of chromophores encourage more detailed studies of their rich photochemistry in order to reveal how Re photosensitizers can be harnessed for new applications.

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CHAPTER 5: TOWARDS THE FORMATION OF TRANSITION METAL HYDRIDE COMPLEXES VIA LIGAND-TO-METAL CHARGE TRANSFER EXCITATION

5.1 Introduction

Chapters 1 and 2 have established the importance and possible utility of utilizing solar energy in the formation of transition metal hydride complexes. This chapter describes ongoing efforts to promote metal hydride formation via an unprecedented and exciting mechanism.

Historically, promoting hydride formation with light has been accomplished via excitation of a photoreductant or photoacid in the presence of a transition metal complex.1–4

Indeed, several examples of hydride formation with photoreductants and/or photoacids have been reported. However, in these examples the photoexcited reactant simply acts as an outer-sphere reductant or proton source and the metal complex remains in the electronic ground state throughout the reaction (similar to the reactivity described in the latter half of Chapter 2).

Figure 5.1 Comparison between the prior use of photo-reductants and photoacids to generate TMHCs and our proposed reactivity based on photoactive transition metal complexes.

We envision an alternative mechanism in which PCET reactivity can be promoted by photoexcitation of the metal complex itself to yield a transition metal hydride complex (TMHC).

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While photon-promoted reactivity of TMHCs to release hydrogen fuel has recently been reported,2,5,6 excited-state PCET reactivity of a transition metal complex to form a TMHC has yet to be established. This unprecedented light-induced reactivity promises to create a new paradigm for solar fuel production in which the catalyst itself acts to harvest photon energy and mediate the PCET reactivity that converts energy poor substrates to energy rich fuels.

We sought to identify an excited-state electronic structure that would result in a more reduced metal upon photoexcitation in order to increase the basicity of the metal center to promote protonation to form a metal hydride. Ligand-to-metal charge transfers result in precisely that; a shift in electron density from an ancillary ligand to the metal center upon light absorption

(Scheme 5.1). While excited-state PCET (ES-PCET) involving LMCT excited states has not been reported, reports of excited-state PT (ES-PT) and excited-state ET (ES-ET) support the hypothesis of the proposed reactivity in well-designed systems. Metal centered photobasic

(increased basicity upon photoexcitation) reactivity, or the protonation of the LMCT excited

4- 3- state of Mo2X8 (X = Cl, Br) by hydrohalic acids to form Mo2X8H has been reported and the resulting Mo-H bond formation was rationalized by recognizing that the LMCT excited state increases electron density at the metal centers.7 Additionally, LMCT excited states of ReII

II 2+ (homoleptic, tris-diphosphine rhenium(II), [Re (PP)3] ) complexes have previously been characterized as highly oxidizing and can be quenched with mild reductants.8–11

Scheme 5.1 LMCT excited states increase the basicity of the metal center and generate a strongly oxidizing ligand cation.

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We have designed a coordination complex that we predict will have the appropriate ground and excited state electronic structure to promote ES-PCET (Scheme 5.2). The cyclopentadienyl ligand is electrochemically robust and not easily protonated, the chelating

II 2+ phosphine will ideally provide similar orbital configurations to the established [Re (PP)3] , and the R groups on the phosphine and the ancillary ligand provides a means to tune the steric profile and electronic structure of the complexes in order to realize trends in reactivity.

We hypothesize that the target complexes will exhibit LMCT excited states that can be protonated and reduced in their excited states to form TMHCs. This is rationalized by recognizing that LMCT excited states involve the transfer of charge density from ligand to the metal center, increasing the basicity of the metal center and yielding a strongly oxidizing ligand radical cation (Scheme 5.1). Thus, in the presence of a mild reductant and an acid, the photoexcited coordination complex is anticipated to undergo reduction and protonation via an electron transfer-proton transfer (ET-PT) or proton transfer-electron transfer (PT-ET) reaction pathway (Scheme 5.2).

Scheme 5.2 Stepwise pathways of PCET that begin from the LMCT excited state.

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Various members of the Dempsey lab are focusing on using 1,2- bis(diphenylphosphino)alkane ligands (dppm = 1,2-bis(diphenylphosphino)methane, dppe = 1,2- bis(diphenylphosphino)ethane, and dppp = 1,2-bis(diphenylphosphino)propane) to probe the effect of the chain length of the phosphine ligand on the reactivity of the resulting hydride complexes. The preliminary results described herein will describe the use of alkyl-substituted phosphine ligands.

5.2 Results and Discussion

5.2.1 Previous Synthetic Methods

Prior to studying light-induced PCET reactivity of the proposed [ReIICp(PP)(X/L)]0/+ complexes, a thorough thermodynamic understanding of the system is required. Determination of pKa values and reduction potentials of the PCET intermediates are needed in order to guide rational choice of ground state reductants and acids to use in ES-PCET studies.12 Therefore, we need to synthesize the direct PCET products first, complexes of the type [HReIIICp(PP)(X/L)]0/+.

Other researchers in the Dempsey lab have made great strides in the synthesis of complexes of this type where PP is dppm, dppe, or dppp. Similar synthetic routes were attempted for 1,2-bis(dialkylphosphino)ethane complexes first. The common precursor,

ReCl3(PPH3)2(NCCH3), was combined with 1,2-bis(dimethylphosphino)ethane (dmpe) in toluene in a J-Young tube (Scheme 5.3).

Scheme 5.3 Synthesis of ReCl3(dmpe)(PPh3) from ReCl3(PPh3)2(NCCH3) and dmpe.

Heating for 5 minutes results in a change from the yellow, partially soluble starting material to a green insoluble material. The 31P{1H} NMR spectrum before heating shows only

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free dmpe, and after heating shows no free dmpe and only free PPh3 (Figure 5.2). Removing the

1 toluene under high vacuum and dissolution in CDCl3 reveals a set of H NMR resonances at 13.9 ppm and 8.7 ppm that integrate in a 6:9 ratio, which are assigned to the ortho protons and the combined para and meta protons on a bound PPh3 ligand (Figure 5.3).

31 1 Figure 5.2 P{ H} NMR spectra before and after heating ReCl3(PPh3)2(NCCH3) and dmpe in toluene

The liberation of one equivalent of PPh3 and the loss of the free dmpe ligand, as indicated by the 31P{1H} NMR spectra, along with the 1H NMR resonances of the resulting green solution allows us to tentatively assign the product as ReCl3(dmpe)(PPh3). When ReCl3(PPh3)2(NCCH3) is combined with 1 equivalent of dmpe in CDCl3, nearly identical spectra are obtained, but the reaction goes at room temperature in 5 minutes.

ReCl3(dmpe)(PPh3) was then combined with 2 equivalents of Co(Cp*)2 in the presence of

30 equivalents of CpH in difluorobenzene (desired reactivity shown in Scheme 5.4), optimized procedures for the synthesis of HReCp(dppe)Cl. While there was very minor conversion to hydride-containing species as evident by the 1H NMR spectrum, the conversion was very poor compared to the aryl phosphine complexes.

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1 Figure 5.3 H NMR of ReCl3(dmpe)(PPh3) in CDCl3

Additionally, under the same conditions but applying heat for 20 minutes at 90° C, the

31P{1H} NMR spectrum shows 8-9 resonances, indicating a larger extent of undesired reactivity as compared to the aryl phosphine complexes. While further optimization of the reaction conditions is possible, we also attempted to pursue alternative routes for the syntheses of hydride complexes containing alkyl substituted chelating phosphine ligands.

Scheme 5.4 Synthesis of HReCp(dmpe)Cl from ReCl3(dmpe)(PPh3)

5.2.2 Alternative Synthetic Methods

Due to complications in the above synthesis of HReCp(dmpe)Cl, alternate synthetic routes were pursued. The first step of this alternate synthetic route involves combining ammonium perrhenate ([NH4][ReO4]) with one equivalent of 1,2-bis(diethylphosphino)ethane

(depe) in an 89:11 HCl/ethanol mixture and heating to 95 °C for 6 hours (Scheme 5.5). The

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solution turned from colorless to blue to purple over the 6 hours, and cooling the solution to -35

°C overnight yielded bright blue crystals of OReCl3(depe) in 55% yield.

Scheme 5.5 Synthesis of OReCl3(depe) from ammonium perrhenate and depe

The 1H and 31P{1H} NMR (Figure 5.4 and Figure 5.5) spectra support the facial chloride isomer shown in Scheme 5.5. Specifically, if the chloride ligands were meridional, the 31P{1H}

NMR spectrum would exhibit two phosphorous resonances due to asymmetry of the depe ligand, which is not observed. The asymmetry in the 1H NMR spectrum arises from half of the depe protons pointing up towards the oxo ligand and the other half of the protons pointing down towards the chloride ligand.

1 Figure 5.4 H NMR spectrum of OReCl3(depe) in d6-DMSO

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31 1 Figure 5.5 P{ H} NMR spectrum of OReCl3(depe) in d6-DMSO

Scheme 5.6 Reduction of OReCl3(depe) with LiAlH4 to form Re(depe)H7

The OReCl3(depe) was then reduced by excess LiAlH4 in diethyl ether at low temperatures (Scheme 5.6).

1 Figure 5.6 Labeled H NMR spectrum of Re(depe)H7 and minor impurity (depe)H2Re(µ- H4)ReH2(depe) in C6D6

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Upon quenching and workup, Re(depe)H7 was obtained as a golden solid in 70% yield.

The 1H and 31P{1H} NMR spectra (Figure 5.6 and Figure 5.7) show good purity, with ~10% of an additional hydride impurity tentatively assigned to (depe)H2Re(µ-H4)ReH2(depe) based on integration, the hydride splitting pattern and literature precedent.13

31 1 Figure 5.7 P{ H} NMR spectrum of Re(depe)H7 and minor impurity (depe)H2Re(µ- H4)ReH2(depe) in C6D6

Re(depe)H7 was then combined with cyclopentadiene (CpH) in a variety of solvents at different temperatures. The desired net reaction is the loss of 3 H2 equivalents, and the coordination of CpH followed by oxidative addition across the sp3 C-H bond to give

ReCp(depe)H2.

1 equivalent of CpH with respect to Re(depe)H7 in C6D6 was heated at 90 °C overnight in a J-Young tube, which resulted in very minor conversion to a species that exhibits a hydride resonance at -13.9 ppm and a singlet at 4.71 ppm, which integrate to a ratio of 2:5 (Figure 5.8).

This indicates the successful formation of the desired product, ReCp(depe)H2.

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1 Figure 5.8 H NMR spectrum of J-Young tube scale reaction between Re(depe)H7 and 1 equivalent of CpH in C6D6 heated to 90° C overnight (~16 hours).

The same reaction was carried out with 30 equivalents of CpH in d5-chlorobenzene, and heated to 135 °C for various time points (Figure 5.9). The resonances for ReCp(depe)H2 grow in over minutes to hours, however they are accompanied by multiple other hydride resonances indicating that other byproducts or intermediates are being formed.

Figure 5.9 J-Young tube scale reaction between Re(depe)H7 and 30 equivalents of CpH in C6D5Cl prior to heating, and heated to 135 °C for 45 minutes, 90 minutes, and ~16 hours.

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Heating the reaction mixture overnight results in complete consumption of the starting material Re(depe)H7, but the various other resonances remain at relatively the same concentration throughout. Running the same reaction (30 equivalents of CpH) in d8-toluene at

105 °C resulted in a similar outcome: near-complete consumption of the starting material after heating overnight with Re(depe)H7 as the major product and a number of other hydride- containing species at low concentration. Further optimization of the time, temperature, and solvent choice for this reaction promises to yield pure ReCp(depe)H2.

Figure 5.10 J-Young tube scale reaction between Re(depe)H7 and 30 equivalents of CpH in d8- toluene prior to heating, and heated to 135 °C for 30 minutes and ~16 hours.

5.2.3 Differences Between Aryl and Alkyl Phosphines

We aim to study the differences in PCET reactivity between complexes of the type

[ReIICp(PP)(X/L)]0/+ with aryl-substituted versus alkyl substituted chelating phosphine ligands.

The electronic and possibly steric differences between the two classes of ligands are, however, influencing the reactivity of their respective syntheses as well.

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The first major difference is in the rate of reactivity between the precursor

ReCl3(PPh3)2(NCCH3) and dppe vs dmpe; in toluene, dppe and dppm require ~1-2 hours of heating at reflux to fully metallate to form ReCl3(dppe)(PPh3), whereas dmpe metallates to form

ReCl3(dmpe)(PPh3) in CDCl3 at room temperature in ~5 minutes. Additionally, the solubility and colors of the products are dramatically different; ReCl3(dppe)(PPh3) is yellow-orange and readily soluble in toluene, whereas ReCl3(dmpe)(PPh3) is not soluble in toluene and is dark green when dissolved in CDCl3. These observations indicate a dramatic difference in the electronic structures imposed by the phosphine ligands with varying substituents.

Cyclic voltammograms of the ReCl3(PP)(PPh3) precursors were collected in order to guide the choice of reductant to use in the synthesis of HReCp(PP)Cl complexes. Cyclic voltammograms of ReCl3(dppe)(PPh3) and ReCl3(dppm)(PPh3) both exhibit a reduction feature that is chemically and electrochemically reversible at fast scan rates (>5 V/s) but chemically irreversible at lower scan rates, interpreted as a relatively slow ligand loss process that occurs upon reduction (studies are still ongoing in our lab). The reversible couples allow for the

+ determination of accurate E1/2 values: for ReCl3(dppe)(PPh3) E1/2 = -1.28 V vs Fc /Fc in 1,2-

+ difluorobenzene (E1/2 = -1.30 V vs Fc /Fc in THF) and for ReCl3(dppm)(PPh3) E1/2 = -1.22 V vs

+ + Fc /Fc in 1,2-difluorobenzene (E1/2 = -1.06 V vs Fc /Fc in CH3CN). In the cyclic voltammogram of ReCl3(dmpe)(PPh3), however, the first reduction feature is chemically irreversible even at fast

+ scan rates, and the reductive peak potential (Ep,c) is at -1.64 V – -1.81 V vs. Fc /Fc, which is much more cathodic than the reduction features of the aryl phosphine complexes. Since the first reduction feature for the complexes is anticipated to be metal-based,14 this trend in reduction potential supports the greater electron-donating ability of the alkyl phosphine ligands as compared to the aryl phosphine ligands as the complex with dmpe is more difficult to reduce.

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Further experiments are ongoing in order to probe both the reactivity of ReCl3(PP)(PPh3) complexes under electrochemically reducing conditions, and the reactivity of reduced Re(I) species with cyclopentadiene towards the efficient synthesis of [HReCp(PP)(X/L)]0/+ complexes.

5.3 Conclusions and Future Directions

Presented in this chapter are our efforts toward the synthesis of [HReCp(PP)(X/L)]0/+ complexes with 1,2-bis(dialkylphosphino)ethane ligands. Synthetic routes through

ReCl3(PP)(PPh3) intermediates that work well when PP = diarylphosphines have proved more difficult when PP = dialkyl phosphines, presumably due to the more negative reducing conditions required. Therefore, alternative synthetic routes through polyhydride intermediates were pursued, which resulted in higher yields of desired rhenium cyclopentadiene bis(alkylphosphino)ethane complexes. The complex ReCp(depe)H2 will be combined with

CHCl3 in an attempt to form HReCp(depe)Cl, the targeted rhenium(III) hydride from which thermodynamic parameters can be experimentally determined. Additionally, H• abstraction from

HReCp(depe)Cl will lead to the targeted rhenium(II) complexes which are hypothesized to exhibit LMCT transitions. With these complexes in hand, photoinduced PCET reactivity will be assessed utilizing various steady-state and time-resolved spectroscopic techniques in order to add an exciting new mechanism to the field of excited-state PCET.

5.4 Experimental Details

Synthesis of OReCl3(depe)

Ammonium perrhenate (231 mg, 0.86 mmol) was added to ethanol (50 mL) and concentrated HCl (5 mL) and the mixture was sparged with nitrogen. Depe (0.500 mL, 2.1 mmol) was added under N2 to the mixture, and the solution was stirred at reflux for 6 hours, during which the color changed slowly from colorless to light blue and then to dark blue. The

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mixture was placed in the freezer at -30° C overnight, resulting in the crystallization of a light blue solid. The solid was filtered, rinsed with ethanol, and dried under high vacuum to yield 444 mg of OReCl3(depe) (55% yield).

Synthesis of Re(depe)H7

OReCl3(depe) (237 mg, 0.46 mmol) was added to 20 mL of ether in the glovebox.

Lithium aluminum hydride (179 mg, 4.7 mmol) was added in small portions over ~5 minutes to the stirring suspension of OReCl3(depe) in ether, then the reaction was stirred at room temperature for 1.5 hours. The resulting mixture was frozen to 77 K using liquid nitrogen, 300

µL of H2O was added, and the reaction was swirled constantly while it warmed to room temperature. Excess anhydrous magnesium sulfate was added, the solution was filtered, and the

filtrate was dried under high vacuum to yield 184 mg (70% yield) of golden Re(depe)H7 solid.

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REFERENCES

(1) Perutz, R. N.; Procacci, B. Photochemistry of Transition Metal Hydrides. Chem. Rev. 2016, 116 (15), 8506–8544.

(2) Teets, T. S.; Nocera, D. G. Photocatalytic Hydrogen Production. Chem. Commun. 2011, 47 (33), 9268.

(3) Esswein, A. J.; Nocera, D. G. Hydrogen Production by Molecular Photocatalysis. Chemical Reviews. 2007, pp 4022–4047.

(4) Dempsey, J. L.; Winkler, J. R.; Gray, H. B. Mechanism of H2 Evolution from a Photogenerated Hydridocobaloxime. J. Am. Chem. Soc. 2010, 132 (47), 16774–16776.

(5) Pitman, C. L.; Miller, A. J. M. Molecular Photoelectrocatalysts for Visible Light-Driven Hydrogen Evolution from Neutral Water. ACS Catal. 2014, 4 (8), 2727–2733.

(6) Chambers, M. B.; Kurtz, D. A.; Pitman, C. L.; Brennaman, M. K.; Miller, A. J. M. Efficient Photochemical Dihydrogen Generation Initiated by a Bimetallic Self-Quenching Mechanism. J. Am. Chem. Soc. 2016, 138 (41), 13509–13512.

(7) Trogler, W. C.; Erwin, D. K.; Geoffroy, G. L.; Gray, H. B. Production of Hydrogen by Ultraviolet Irradiation of Binuclear Molybdenum(II) Complexes in Acidic Aqueous Solutions. Observation of Molybdenum Hydride Intermediates in Octahalodimolybdate(II) Photoreactions. J. Am. Chem. Soc. 1978, 100 (4), 1160–1163.

(8) Lee, Y. F.; Kirchhoff, J. R. Absorption and Luminescence Spectroelectrochemical Characterization of a Highly Luminescent Rhenium(II) Complex. J. Am. Chem. Soc. 1994, 116 (8), 3599–3600.

(9) Del Negro, A. S.; Seliskar, C. J.; Heineman, W. R.; Hightower, S. E.; Bryan, S. A.; Sullivan, B. P. Highly Oxidizing Excited States of Re and Tc Complexes. J. Am. Chem. Soc. 2006, 128 (51), 16494–16495.

(10) Messersmith, S. J.; Kirschbaum, K.; Kirchhoff, J. R. Luminescent Low-Valent Rhenium Complexes with 1,2-Bis(Dialkylphosphino)Ethane Ligands. Synthesis and X-Ray Crystallographic, Electrochemical, and Spectroscopic Characterization. Inorg. Chem. 2010, 49 (8), 3857–3865.

(11) Adams, J. J.; Arulsamy, N.; Sullivan, B. P.; Roddick, D. M.; Neuberger, A.; Schmehl, R. H. Homoleptic Tris-Diphosphine Re(I) and Re(II) Complexes and Re(II) Photophysics and Photochemistry. Inorg. Chem. 2015, 54 (23), 11136–11149.

(12) Lennox, J. C.; Kurtz, D. A.; Huang, T.; Dempsey, J. L. Excited-State Proton-Coupled Electron Transfer: Different Avenues for Promoting Proton/Electron Movement with Solar Photons. ACS Energy Lett. 2017, 2 (5), 1246–1256.

(13) Fanwick, P. E.; Root, D. R.; Walton, R. A. Reactions of the Dirhenium(II) Complexes

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Re2Cl4(PR3)4 with Lithium Aluminum Hydride. A Convenient Synthetic Route to the Dirhenium Octahydride Complexes Re2H8(PR3)4. Inorg. Chem. 1989, 28 (16), 3203–3209.

(14) Roncari, E.; Mazzi, U.; Seeber, R.; Zanello, P. Electrochemical Investigation on the Phosphine Rhenium Complexes [ReCl3(PMe2Ph)3] and [ReCl4(PMe2Ph)2] in an Aprotic Medium. J. Electroanal. Chem. Interfacial Electrochem. 1982, 132, 221–231.

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APPENDIX A: ADDITIONAL DETAILS AND CHARACTERIZATION FOR CHAPTER 4

A.1 Experimental Details

A.1.1 General Considerations.

Syntheses were performed under N2 using standard Schlenk line techniques unless otherwise noted. All solvents used for synthetic procedures were purchased from Thermo Fischer

Scientific and were used without purification. Rhenium pentacarbonyl chloride (ACROS

Organics, 98%), silver trifluoromethanesulfonate (AgOTf), ammonium hexafluorophosphate

([NH4][PF6], Oakland Chemical, 99%), trimethyl phosphine (PMe3) (Aldrich, 1.0M in toluene), trimethylamine N-oxide dihydrate (TMNO, Alfa Aesar, 98+%), and benzyltrimethylammonium chloride ([BnMe3N][Cl], Aldrich, 97%) were used without further purification. 4,4′-diethylester-

43 2,2′-bipyridine (deeb) was synthesized according to previous procedures. Acetonitrile-d3

(99.8% D) was purchased from Cambridge Isotope Laboratories. 2-Methyltetrahydrofuran

(99+%, Extra Dry) was purchased from Acros Organics.

1H, 13C, and 31P{1H} NMR spectra were collected on a Bruker 600 MHz spectrometer at

295 K. Chemical shifts are reported relative to residual protio solvent signals. UV/Vis absorption spectra were collected using an Agilent Cary 60 UV/Vis absorbance spectrophotometer.

Photoluminescence measurements were recorded using an LED excitation source (Ocean Optics

LLS, either 365 or 455 nm), fiber coupled to a 90 sample compartment (Ocean Optics COV-

UV-FL), fiber coupled to an Ocean Optics USB2000+ spectrophotometer detector.

A.1.2 Electrochemical Measurements.

Electrochemistry was performed in a N2-filled glovebox with a WaveDriver (Pine

Research) potentiostat using a 3 mm diameter platinum working electrode, a 3 mm diameter glassy carbon counter electrode, and a silver wire pseudo-reference electrode. A 20-mL

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scintillation vial was used as an electrochemical cell, fitted with a custom-made Teflon cap to hold the three electrodes. The electrode leads in the glovebox were connected to the potentiostat with a custom shielded electrode cable feedthrough. All scans were referenced to the Fc+/Fc couple at 0 V. Acetonitrile (Fisher Scientific, HPLC grade, >99.9%) for electrochemical experiments was dried and degassed using a Pure Process Technology solvent purification system. Ferrocene was present in each scan unless otherwise noted. Ohmic drop was minimized using a high electrolyte concentration (0.25 M tetrabutylammonium hexafluorophosphate,

TBAPF6), through minimization of the distance between the working and reference electrodes, and through manual iR compensation.44 Platinum electrodes (CH Instruments, 3 mm diameter disk) were polished with 0.05 μm alumina powder (CH Instruments, contained no agglomerating agents) Milli-Q water slurries, rinsed, and ultrasonicated briefly in Milli-Q water to remove residual polishing powder. The silver wire pseudo-reference electrode was submerged in a glass tube containing electrolyte (0.25 M TBAPF6 in acetonitrile, CH3CN) and separated from the solution with a porous glass Vycor tip. The working electrode was pretreated with cyclical scans from approximately 2 to −2 V (the exact value varied in accordance with the silver wire pseudo- reference) at 250 mV/s in 0.25 M TBAPF6 until cycles were superimposable (typically achieved within three cycles).

A.1.3 Time-Resolved Photoluminescence.

Samples for time-resolved photoluminescence experiments were prepared in an N2 filled glovebox. ~2-5 mg of each sample was dissolved in ~1 mL of a 90/10 2-MeTHF/CH3CN solution and transferred to a NMR tube. The cap was secured with electrical tape and parafilm. A cold-finger Dewar with a quartz cavity at the bottom was used for both the 295 K lifetime experiments, and filled with liquid N2 for the lifetime experiments at 77 K.

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Time resolved photoluminescence experiments were performed using a custom-build laser flash photolysis system.45 Laser excitation (5-7 ns FWHM, 10 Hz, Q-switched) was provided by the third harmonic of a Nd:YAG laser (Spectra-Physics, Inc., model Quanta-Ray

LAB-170-10) that pumped an OPO (basiScan, GWU Lasertechnik) to access tunable excitation

(415–800 nm). Laser power at the sample cuvette was attenuated by the use of a half waveplate

(WPMH10M-355, ThorLabs) and polarizer (GL10-A, ThorLabs). A glass window was used to deflect a small portion of excitation beam to a Si diode detector (DET10A, ThorLabs), triggering the oscilloscope to start data collection. Timing of the laser was controlled by a digital delay generator (9514+ Pulse Generator, Quantum Composers). Single wavelength emission data

(monitored near the em,max of each complex) was obtained using a double slit monochromator

(Spectral Products CM112) outfitted with a Hamamatsu R928 photomultiplier tube (PMT). The signal intensity was attenuated with a neutral density filter, and scattered excitation light was filtered with a color filter wheel containing various long pass and short pass filters. The signal was amplified by a 200 MHz wideband voltage amplifier (DHPVA-200, Electro Optical

Components) and processed using a digitizer (CompuScope 12502, GaGeScope) controlled by custom software (MATLAB). A dark current (detection without laser excitation) was subtracted from the raw luminescence data, and further analyzed using Igor Pro 6.37 (Wavemetrics).

A.1.4 Single-crystal X-ray diffraction.

Single-crystal X-ray diffraction data were collected on a Bruker APEX-II CCD diffractometer at 100 K with Cu Kα radiation (λ = 1.54175 Å). The structures were solved using

Olex246 with the ShelXT47 structure solution program using intrinsic phasing and refined with the ShelXL47 refinement program using least-squares minimization. A disordered solvent molecule, CH2Cl2 in the asymmetric unit was refined over two positions, each with 50%

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occupancy. SADABS-2016/2 (Bruker,2016/2) was used for absorption correction with Abs T max = 0.7536 and Abs T min = 0.5210. wR2(int) was 0.0792 before and 0.0529 after correction.

A.1.5 Computational Details.

Density Functional Theory (DFT) and Time-Dependent Density Functional Theory (TD-

DFT) calculations were done using Gaussian 09.48 Geometry optimizations were performed using the B3LYP hybrid functional with the LANL2DZ ECP49,50 basis set for Re and 6-31G* basis set for all heteroatoms. The implicit SMD solvation model was employed for all calculations (acetonitrile or dichloromethane solvent). The M06 and PBE functionals were also tested for select species, but B3LYP generally showed the best agreement with experimental parameters.

A.1.6 Quantum Yield Measurement Experimental Details

The relative quantum yield of 1 was determined using methods described previously.1

The standard used for reference was [Ru(bpy)3][PF6]2 in deoxygenated CH3CN, conditions under which the absolute quantum yield was reported recently to be 0.095.2 Emission spectra used for quantum yield determination were recorded using an Edinburgh FLS920 spectrometer with luminescence first passing through a 425 nm long-pass color filter, then a single grating (1800 l/mm, 500 nm blaze) Czerny-Turner monochromator (13.05 nm bandwidth) and detected by a peltier-cooled Hamamatsu R2658P photomultiplier tube. Samples were excited at 407 nm using light output from a housed 450 W Xe lamp / single grating (1800 l/mm, 250 nm blaze) Czerny-

Turner monochromator combination with 12.6 nm bandwidth.

A.1.7 Synthesis

Synthesis of Re(deeb)(CO)3Cl (Complex 2)

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Re(CO)5Cl (335 mg, 0.93 mmol) and deeb (274 mg, 0.92 mmol) were combined in 50 mL of toluene. The mixture was heated to reflux and stirred for 2 hours, resulting in a color change from white to dark orange. Once cooled to 295 K, 100 mL of pentane was added, and an orange precipitate was collected. The solid was washed with pentane and dried under reduced

1 pressure to yield a fine orange powder in 84% yield. H NMR (600 MHz, CD3CN) δ 9.19 (dd, J

= 5.7, 0.8 Hz, 2H), 8.95 (dd, J = 1.7, 0.8 Hz, 2H), 8.07 (dd, J = 5.6, 1.7 Hz, 2H), 4.48 (q, J = 7.1

13 1 Hz, 4H), 1.44 (t, J = 7.1 Hz, 7H). C{ H} NMR (151 MHz, CD3CN) δ 198.4, 189.9, 164.1,

157.1, 155.0, 142.0, 127.6, 124.5, 63.6, 14.3. Elemental Analysis, Calcd for C19H16ClN2O7Re: C,

37.66; H, 2.66; N, 4.62; Found C, 37.39; H, 2.57; N, 4.65.

Synthesis of [Re(deeb)(CO)3(NCCH3)][PF6] (Complex 1)

Re(deeb)(CO)3Cl (50.3 mg, 0.083 mmol) was dissolved in 15 mL of degassed CH3CN.

AgOTf (82.5 mg, 0.32 mmol, 3.9 equiv) was then added under positive N2 pressure. The solution was protected from light by foil and stirred at reflux for 3.5 hours. The solution changed color from orange to yellow during the reflux, and a AgCl precipitate formed. The AgCl was removed by filtration and the solvent was removed from the filtrate by rotary evaporation. The yellow- orange residue was dissolved in 2 mL of CH3CN, to which was added 15 mL of a saturated, aqueous solution of [NH4][PF6] followed by 40 mL of H2O. A yellow solid precipitated from solution and was collected on a frit, washed with copious H2O and diethyl ether, and dried in vacuo (84% yield). (600 MHz, CD3CN) δ 9.20 (dd, J = 5.7, 0.8 Hz, 2H), 9.01 (dd, J = 1.7, 0.8

Hz, 2H), 8.16 (dd, J = 5.7, 1.6 Hz, 2H), 4.50 (q, J = 7.1 Hz, 4H), 2.02 (s, 3H), 1.45 (t, J = 7.1 Hz,

13 1 6H). C{ H} NMR (151 MHz, CD3CN) δ 194.4, 190.6, 164.0, 157.7, 156.1, 142.9, 128.2, 124.8,

123.8, 63.9, 14.3, 3.9. Elemental Analysis, Calcd for C21H19F6N3O7PRe: C, 33.34; H, 2.53; N,

5.55; Found C, 33.64; H, 2.46; N, 5.62.

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Synthesis of [Re(deeb)(CO)3(PMe3)][PF6]

Re(deeb)(CO)3Cl (298 mg, 0.49 mmol) was dissolved in 50 mL of degassed acetone.

AgOTf (379 mg, 1.47 mmol, 3.0 equiv) was then added under positive N2 pressure. The solution was protected from light by foil and stirred at reflux for 2 hours, accompanied by the precipitation of AgCl and a solution color change from orange to yellow. After the reaction, the

AgCl was removed by filtration using a fine frit and to the filtrate was added 1.0 M PMe3 in toluene (14.8 mL, 14.8 mmol, 30 equiv) via syringe. The solution was sparged with N2, and stirred at reflux under N2 for 16 hours. The orange solution was concentrated and to the residue was added 5 mL of MeOH. The solution was filtered and added dropwise to a solution of ~1g of

[NH4][PF6] in 100 mL H2O while vigorously stirring. The mixture was sonicated for 10 minutes, and the resulting yellow-orange powder was collected by filtration. After rinsing with copious

H2O and diethyl ether, the solid was suspended in 5 mL MeOH and sonicated for 10 minutes.

The solid was filtered again and dried under reduced pressure to give the dark yellow product

1 (76% yield). H NMR (600 MHz, CD3CN) δ 9.20 (ddd, J = 5.7, 0.9, 0.9 Hz, 2H), 9.03 (dd, J =

1.7, 0.8 Hz, 2H), 8.12 (dd, J = 5.7, 1.7 Hz, 2H), 4.50 (q, J = 7.1 Hz, 4H), 1.45 (t, J = 7.1 Hz, 6H),

13 1 1.10 (d, J = 9.4 Hz, 9H). C{ H} NMR (151 MHz, CD3CN) δ 195.4 (d, J = 7.7 Hz), 188.5 (d, J

= 59.6 Hz), 163.9, 157.0, 155.9, 142.3, 128.3, 125.3, 63.8, 14.3, 13.2 (d, J = 32.5) Hz. 31P{1H}

NMR (243 MHz, CD3CN) δ -28.55 (s, 1P), -144.66 (hept, J = 706 Hz, 1P, PF6). Elemental

Analysis, Calcd for C22H25F6N2O7P2Re: C, 33.38; H, 3.18; N, 3.54; Found C, 33.54; H, 31.0; N,

3.54.

Synthesis of [Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] (Complex 3)

[Re(deeb)(CO)3(PMe3)][PF6] (147 mg, 0.19 mmol) was dissolved in 20 mL of degassed

CH3CN. TMNO (22.0 mg, 0.39 mmol, 1.1 equiv) was dissolved in 5 mL of CH3CN and 1 mL of

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MeOH, and then injected into the reaction mixture. Within 5 minutes of injection, the solution color changed from yellow-orange to dark red. The solution was then stirred at reflux for 2 hours. After removing the solvent, the residue was dissolved in minimal CH3CN; addition of

H2O and sonication resulted in the precipitation of a dark red solid. The solid was collected,

1 washed with excess H2O and dried under reduced pressure to give 3 in 60% yield. H NMR (600

MHz, CD3CN) δ 9.17 (ddd, J = 5.7, 0.7, 0.7 Hz, 2H), 8.96 (dd, J = 1.7, 0.8 Hz, 2H), 8.08 (dd, J =

5.7, 1.6 Hz, 2H), 4.49 (q, J = 7.1 Hz, 4H), 2.04 (s, 3H), 1.44 (t, J = 7.2 Hz, 6H), 1.15 (d, J = 9.5

13 1 Hz, 9H). C{ H} NMR (151 MHz, CD3CN) δ 202.4 (d, J = 7.2 Hz), 164.2, 157.33, 154.7, 141.6,

31 1 127.9, 124.6, 63.7, 16.7 (d, J = 36.3 Hz), 14.35, 3.92. P{ H} NMR (243 MHz, CD3CN) δ -

22.25 (s, 1P), -144.66 (hept, J = 706 Hz, 1P, PF6). Elemental Analysis, Calcd for

C23H28F6N3O6P2Re: C, 34.33; H, 3.51; N, 5.22; Found C, 34.61; H, 3.54; N, 5.16.

Synthesis of Re(deeb)(CO)2(PMe3)Cl (Complex 4)

[Re(deeb)(CO)2(PMe3)(NCCH3)][PF6] (121 mg, 0.15 mmol) was dissolved in 100 mL of dichloromethane (CH2Cl2), to which [BnMe3N][Cl] (734 mg, 4.5 mmol, 30 equiv) was added.

The solution was sparged with N2 for 20 minutes, and then stirred at reflux for 24 hours which was accompanied by a color change from dark red to dark purple. The excess [BnMe3N][Cl] was removed by filtration, and the solvent was removed by rotary evaporation. The residue was dissolved in 20 mL CH2Cl2 and impurities were extracted with H2O (3 x100 mL). The CH2Cl2 layer was then dried with anhydrous MgSO4, filtered, and the solvent was removed. The residue was dissolved in minimal (5 mL) of CH2Cl2, to which 100 mL of pentane was added, resulting in the precipitation of a dark purple solid. The solid was filtered, washed with pentane and dried under reduced pressure to give 65% yield. Single dark purple block crystals of 4 suitable for X- ray diffraction study were grown from vapor diffusion of pentane into an acetone solution of 4.

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1 H NMR (600 MHz, CD3CN) δ 9.15 (ddd, J = 5.7, 0.7, 0.7 Hz, 2H), 8.88 (dd, J = 1.7, 0.8 Hz,

2H), 7.97 (dd, J = 5.7, 1.7 Hz, 2H), 4.46 (q, J = 7.1 Hz, 4H), 1.43 (t, J = 7.1 Hz, 6H), 1.19 (d, J =

13 1 9.3 Hz, 9H). C{ H} NMR (151 MHz, CD3CN) δ 207.3 (d, J = 7.8 Hz), 164.4, 157.1, 153.5,

31 1 140.4, 127.0, 124.0, 63.4, 18.3 (d, J = 35.3 Hz), 14.36. P{ H} NMR (243 MHz, CD3CN) δ -

23.11.

A.2 Additional Characterization

31 1 Figure A.1 P{ H} NMR spectrum of [Re(deeb)(CO)3(NCCH3)][PF6] (1) in CD3CN

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1 Figure A.2 H NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN.

13 1 Figure A.3 C{ H} NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN.

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31 1 Figure A.4 P{ H} NMR spectrum of [Re(deeb)(CO)3(PMe3)][PF6] in CD3CN.

31 1 Figure A.5 P{ H} NMR spectrum of 3 in CD3CN.

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13 1 Figure A.6 C{ H} NMR spectrum of 4 in CD3CN.

31 1 Figure A.7 P{ H} NMR spectrum of 4 in CD3CN.

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Figure A.8 Absorption and emission spectra of 1 and [Ru(bpy)3][PF6]2 standard in deoxygenated CH3CN used for quantum yield determination.

Table A.1 Radiative and non-radiative rate constants for complexes 1 and 2, and photophysical parameters used in calculations.

-1 -1 Complex r  (295 K, ns) kr (s ) knr (s ) 1 0.049 150 3.3 x 105 6.3 x 106 2 0.013 13 1.0 x 106 7.6 x 106

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APPENDIX B: ADDITIONAL ELECTROCHEMICAL AND PROTON TRANSFER RATE CONSTANT DETERMINATION DATA FOR CHAPTERS 2 AND 3

Figure B.1 Determination of diffusion coefficients for [CoCp(dppe)(NCCH3)][PF6]2 Top: cyclic voltammograms of 0.5 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 supporting electrolyte were recorded at 0.05, 0.1, 0.25, 0.5, 0.75, 1, 2 and 3 V/s. The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom: plots of the peak current (corrected for capacitive current) as a function of 1/2 for the CoIII/II (left) and CoII/I (right) couples. From the slopes, an average diffusion coefficient of 3.4 × 10-6 cm2/s was calculated from the CoIII/II wave, and 5.2 × 10-6 cm2/s for the CoII/I wave. The reoxidation of the CoI to CoII was only studied up to 1 V/s as determining an accurate value for the capacitive current proved more difficult.

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Figure B.2 Determination of diffusion coefficients for [CoCp(depe)(NCCH3)][PF6]2 Top: cyclic voltammograms of 0.2 mM of [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 supporting electrolyte were recorded at 0.05, 0.075, 0.1, 0.125, 0.15, 0.2, 0.25, 0.3, 0.35, 0.4, 0.5, 0.6, 0.7, 0.8 and 1 V/s. The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom: plots of the peak current (corrected for capacitive current) as a function of υ1/2 for the CoIII/II (left) and CoII/I (right) couples. From the slopes, an average diffusion coefficient of 4.4×10-6 cm2/s was calculated from the CoIII/II wave, and 6.7×10-6 cm2/s for the CoII/I wave.

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Figure B.3 Determination of diffusion coefficients for [CoCp(dcpe)(NCCH3)][PF6]2 Top: cyclic voltammograms of 0.2 mM of [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 supporting electrolyte were recorded at 0.05, 0.075, 0.1, 0.125, 0.15, 0.2, 0.25, 0.3, 0.4, 0.5, 0.6 and 0.8 V/s. The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom: plot of the peak current (corrected for capacitive current) as a function of υ1/2 for CoII/I couple. From the slope, an average diffusion coefficient of 7.0×10-6 cm2/s was calculated for the CoII/I wave. The CoIII/II wave was not reliable due to the slow isomerization between cyclohexyl group conformations over the time scale of the CV.

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Figure B.4 Determination of diffusion coefficients for [CoCp(dppv)(NCCH3)][PF6]2 Top: cyclic voltammograms of 0.193 mM of [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 supporting electrolyte were recorded at 0.05, 0.1, 0.15, 0.2, 0.25, 0.35, 0.5, 0.75 and 1 V/s. The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom: plots of the peak current (corrected for capacitive current) as a function of υ1/2 for the CoIII/II (left) and CoII/I (right) couples. From the slopes, an average diffusion coefficient of 5.9×10-6 cm2/s was calculated from the CoIII/II wave, and 7.8×10-6 cm2/s for the CoII/I wave.

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Figure B.5 Determination of heterogeneous electron transfer rate constants for [CoCp(dppe)(NCCH3)][PF6]2 Trumpet plot of the CoIII/II (red circles) and CoII/I (blue circles) waves overlaid with the working -5 2 curve (dashed curves). The working curves used the parameters D = 1×10 cm /s and ks = 1 cm/s. The offset (∆) between the data and the x-axis of the best fitting working curve is used to III/II determine ks. The obtained ks are 0.051 cm/s (∆ = 3.05) and 0.11 cm/s (∆ = 2.4) for the Co and CoII/I wave respectively.

Figure B.6 Determination of heterogeneous electron transfer rate constants for [CoCp(depe)(NCCH3)][PF6]2 Trumpet plot of the CoIII/II (blue circles) and CoII/I (red circles) waves overlaid with the working -5 2 curve (dashed curves). The working curves used the parameters D = 1×10 cm /s and ks = 1 cm/s. The offset (∆) between the data and the x-axis of the best fitting working curve is used to III/II determine ks. The obtained ks are 0.022 cm/s (∆ = 3.1) and 0.36 cm/s (∆ = -0.7) for the Co and CoII/I wave respectively.

138

Figure B.7 Determination of heterogeneous electron transfer rate constant for [CoCp(dcpe)(NCCH3)][PF6]2. Trumpet plot of the CoII/I waves overlaid with the working curve (dashed curves). The working -5 2 curves used the parameters D = 1×10 cm /s and ks = 1 cm/s. The offset (∆) between the data and the x-axis of the best fitting working curve is used to determine ks. The obtained ks is 0.19 cm/s (∆ = 1.3) for the CoII/I wave. The CoIII/II wave was not reliable due to the slow isomerization between cyclohexyl group conformations over the time scale of the CV.

Figure B.8 Determination of heterogeneous electron transfer rate constants for [CoCp(dppv)(NCCH3)][PF6]2 Trumpet plot of the CoIII/II (blue circles) and CoII/I (red circles) waves overlaid with the working -5 2 curve (dashed curves). The working curves used the parameters D = 1×10 cm /s and ks = 1 cm/s. The offset (∆) between the data and the x-axis of the best fitting working curve is used to III/II determine ks. The obtained ks are 0.19 cm/s (∆ = 1.2) and 0.50 cm/s (∆ = 0.5) for the Co and CoII/I wave respectively.

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Table B.1 Summary of electrochemically determined protonation rate constants of CoCp(dppe) utilizing acids with varying pKas.

-1 -1 Acid pKa of acid kPT rate constant (M s ) dimethylformamidium triflate 6.1 1.1 × 107 4-cyano-anilinium tetrafluoroborate 7 1.7 × 107 7 4-CF3-anilinium tetrafluoroborate 8.03 1.7 × 10 p-toluenesulfonic acid 8.6 1.8 × 107 4-Br-anilinium tetrafluoroborate 9.43 2.5 × 107 4-Cl-anilinium tetrafluoroborate 9.7 2.4 × 107 4-tBu-anilinium tetrafluoroborate 11.1 2.1 × 107 4-methylanilinium tetrafluoroborate 11.4 1.7 × 107 N,N-dimethylanilinium tetrafluoroborate 11.43 8.3 × 104 4-MeO-anilinium tetrafluoroborate 11.86 1.9 × 107 pyridinium tetrafluoroborate 12.53 2.7 × 107 2,6-lutidinium tetrafluoroborate 14.13 1.4 × 104 4-MeO-pyridinium tetrafluoroborate 14.23 1.2 × 107 2,4,6-trimethylpyridinium tetrafluoroborate 14.98 5.5 × 103 6 4-CF3-2,3,5,6-F4-phenol 16.62 7.7 × 10 salicylic acid 16.7 2.3 × 107 triethylammonium tetrafluoroborate 18.82 7.1 × 102 pentafluorophenol 20.11 2.3 × 105 2,3,5,6-tetrafluorophenol 20.12 1.7 × 105 2,4,6-tribromophenol 20.35 1.9 × 104 benzoic acid 21.51 1.5 × 104 acetic acid 23.51 3.2 × 103 4-chlorophenol 25.44 1.5 × 102

140

Table B.2 Summary of photochemically determined protonation rate constants of CoCp(dppe) utilizing acids with varying pKas.

-1 -1 Acid pKa of acid kPT rate constant (M s ) benzoic acid 21.51 4.2 × 105 2,3,5,6-tetrafluorophenol 20.12 1.4 × 106 triethylammonium tetrafluoroborate 18.82 3.4× 105 2,3,4,5,6-pentachlorophenol 18.02 6.2 × 106 salicylic acid 16.7 3.5× 107 7 2,3,5,6-tetrafluoro-4-CF3-phenol 16.62 4.4× 10

Table B.3 Summary of electrochemically determined protonation rate constants of CoCp(depe) utilizing acids with varying pKas.

-1 -1 Acid pKa of acid kPT rate constant (M s ) phenol 29.14 8.7 × 103 4 4-CH3-phenol 27.5 1.8 × 10 phenol 26.6 8.7 × 103 5 4-CF3-phenol 25.54 7.0 × 10 4-Cl-phenol 25.44 1.2 × 105 5 4-CF3-phenol 24.9 7.0 × 10 acetic acid 23.51 5.6 × 106 benzoic acid 21.51 3.6 × 107 2,3,5,6-tetrafluorophenol 20.12 4.4 × 107 benzylammonium tetrafluoroborate 16.91 3.8 × 107 salicylic acid 16.7 5.6 × 107 7 4-CH3O-pyridinium tetrafluoroborate 14.23 3.2 × 10 pyridinium tetrafluoroborate 12.53 3.5 × 107 7 4-CH3-anilinium tetrafluoroborate 11.4 3.6 × 10 4-tBu-anilinium tetrafluoroborate 11.1 2.9 × 107 4-Cl-anilinium tetrafluoroborate 9.7 4.5 × 107 7 4-CF3O-anilinium tetrafluoroborate 9.28 3.2 × 10

141

4-CN-anilinium tetrafluoroborate 7 1.5 × 107

Table B.4 Summary of electrochemically determined protonation rate constants of CoCp(dcpe) utilizing acids with varying pKas.

-1 -1 Acid pKa of acid kPT rate constant (M s ) 4-Cl-phenol 25.44 2.60 × 102 acetic acid 23.51 8.20 × 103 2,3,5,6-tetrafluorophenol 20.12 8.20 × 104 2,3,4,5,6-pentaphenol 20.11 1.10 × 105 5 2,3,5,6-tetrafluoro-4-CF3-phenol 16.62 1.30 × 10 4 4-CH3O-pyridinium tetrafluoroborate 14.23 5.30 × 10 pyridinium tetrafluoroborate 12.53 5.90 × 104 4-tBu-anilinium tetrafluoroborate 11.1 5.20 × 104 anilinium tetrafluoroborate 10.62 4.90 × 104 4-Br-anilinium tetrafluoroborate 9.43 6.60 × 104

Table B.5 Summary of electrochemically determined protonation rate constants of CoCp(dppv) utilizing acids with varying pKas.

-1 -1 Acid pKa of acid kPT rate constant (M s ) acetic acid 23.51 1.60 × 103 benzoic acid 21.51 2.60 × 103 salicylic acid 16.7 2.20 × 106 7 4-CH3O-pyridinium tetrafluoroborate 14.23 2.20 × 10 4-tBu-anilinium tetrafluoroborate 11.1 2.20 × 107 4-Br-anilinium tetrafluoroborate 9.43 1.30 × 107 4-CN-anilinium tetrafluoroborate 7 2.10 × 107

142

Figure B.9 Electrochemical determination of kPT for CoCp(dppe) with dimethylformamidium triflate Top: cyclic voltammograms of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.045 M of acid at multiple scan rates (50, 75, 100, 200, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the CoII/I wave 2 (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.997. Bottom right: evolution of the 7 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 1.1 × 10 M s .

143

Figure B.10 Electrochemical determination of kPT for CoCp(dppe) with 4-cyano-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.5 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.006 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution II/I of the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 2 2 30 mV/dec, with R = 0.99. Bottom right: evolution of kobs obtained with acid concentration (R 7 -1 -1 = 0.997). The slope yields kPT = 1.7 × 10 M s .

144

Figure B.11 Electrochemical determination of kPT for CoCp(dppe) with 4-CF3-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.054 M of acid at multiple scan rates (50, 100, 200, 300, 400, 500 and 600 mV/s). The CVs are referenced to the Fc+/Fc couple at II/I 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate 2 (). The slope is -32 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at 7 -1 -1 each scan rate with scan rate. The average yields kPT = 1.7 × 10 M s

145

Figure B.12 Electrochemical determination of kPT for CoCp(dppe) with p-toluenesulfonic acid monohydrate Top: cyclic voltammograms of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.045 M of acid at multiple scan rates (50, 75, 100, 200, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the CoII/I wave 2 (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.996. Bottom right: evolution of the 7 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 1.8 × 10 M s

146

Figure B.13 Electrochemical determination of kPT for CoCp(dppe) with 4-Br-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.039 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution of II/I the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 26 2 2 mV/dec, with R = 0.999. Bottom right: evolution of kobs obtained with acid concentration (R = 7 -1 -1 0.999). The slope yields kPT = 2.5 × 10 M s .

147

Figure B.14 Electrochemical determination of kPT for CoCp(dppe) with 4-Cl-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.5 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.003 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution II/I of the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 2 2 29 mV/dec, with R = 0.996. Bottom right: evolution of kobs obtained with acid concentration (R 7 -1 -1 = 0.999). The slope yields kPT = 2.4 × 10 M s .

148

t Figure B.15 Electrochemical determination of kPT for CoCp(dppe) with 4- Bu-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.025 M of acid at multiple scan rates (25, 50, 75, 100, 200, 300, 400, 500, 600, 700 and 800 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -35 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields 7 -1 -1 kPT = 2.1 × 10 M s

149

Figure B.16 Electrochemical determination of kPT for CoCp(dppe) with 4-methylanilinium tetrafluoroborate (method 1) Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.017 M of acid at multiple scan rates (50, 100, 200, 300, 400, 500, 600 and 700 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of II/I 2 the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.999. Bottom right: 7 evolution of the obtained kPT at each scan rate with scan rate. The average yields: kPT = 1.7 × 10 M-1 s-1.

150

Figure B.17 Electrochemical determination of kPT for CoCp(dppe) with 4-methylanilinium tetrafluoroborate (method 2) Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 1.7 – 16.5 mM of acid at 100 mV/s (black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the II/I cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 29 2 mV/dec, with R = 0.99. Bottom right: evolution of the observed rate constant kobs at with the 2 7 -1 -1 concentration of acid. The slope of the linear fit (R = 0.99) yields: kPT = 1.7 × 10 M s

151

Figure B.18 Electrochemical determination of kPT for CoCp(dppe) with N,N-dimethylanilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.045 M of acid at multiple scan rates (50, 100, 200, 300, 400, 500, 600 and 700 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of II/I 2 the Co wave (Ep,c) with scan rate (). The slope is -35 mV/dec, with R = 0.998. Bottom right: 4 evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 8.3 × 10 M-1 s-1.

152

Figure B.19 Electrochemical determination of kPT for CoCp(dppe) with 4-MeO-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.5 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.0099 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution II/I of the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 2 2 43 mV/dec, with R = 0.998. Bottom right: evolution of kobs obtained with acid concentration (R 7 -1 -1 = 0.98). The slope yields kPT = 1.9 × 10 M s .

153

Figure B.20 Electrochemical determination of kPT for CoCp(dppe) with pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.21 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.024 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution II/I of the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 2 2 25 mV/dec, with R = 0.997. Bottom right: evolution of kobs obtained with acid concentration (R 7 -1 -1 = 0.996). The slope yields kPT = 2.7 × 10 M s

154

Figure B.21 Electrochemical determination of kPT for CoCp(dppe) with 2,6-lutidinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.046 M of acid at multiple scan rates (50, 100, 200, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the + II/I Fc /Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R2 = 0.999. Bottom right: evolution of the 4 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 1.4 × 10 M s

155

Figure B.22 Electrochemical determination of kPT for CoCp(dppe) with 4-MeO-pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.020 M of acid at multiple scan rates (100, 200, 300, 400, 500, 600 and 700 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the CoII/I wave 2 (Ep,c) with scan rate (). The slope is -34 mV/dec, with R = 0.999. Bottom right: evolution of the 7 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 1.2 × 10 M s

156

Figure B.23 Electrochemical determination of kPT for CoCp(dppe) with 2,4,6- trimethylpyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.049 M of acid at multiple scan rates (50, 100, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the CoII/I wave 2 (Ep,c) with scan rate (). The slope is -35 mV/dec, with R = 0.999. Bottom right: evolution of the 3 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 5.5 × 10 M s

157

Figure B.24 Electrochemical determination of kPT for CoCp(dppe) with 4-CF3-2,3,5,6-F4-Phenol Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.029 M of acid at multiple scan rates. (50, 75, 100, 200, 300, 40 and 500 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the CoII/I wave 2 (Ep,c) with scan rate (). The slope is -37 mV/dec, with R = 0.998. Bottom right: evolution of the 6 -1 -1 obtained kPT at each scan rate with scan rate. The average yields kPT = 7.7 × 10 M s .

158

Figure B.25 Electrochemical determination of kPT for CoCp(dppe) with salicylic acid Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.046 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of the II/I 2 Co wave (Ep,c) with acid concentration (CA). The slope is 28 mV/dec, with R = 0.996. Bottom 2 right: evolution of kobs obtained with acid concentration (R = 0.998). The slope yields kPT = 2.3 × 107 M-1 s-1.

159

Figure B.26 Electrochemical determination of kPT for CoCp(dppe) with pentafluorophenol Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.18 M of acid at multiple scan rates. (50, 75, 100, 200, 300, 400, 500, 600 and 700 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of II/I 2 the Co wave (Ep,c) with scan rate (). The slope is -36 mV/dec, with R = 0.999. Bottom right: 5 evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 2.3 × 10 M-1 s-1.

160

Figure B.27 Electrochemical determination of kPT for CoCp(dppe) with 2,3,5,6- tetrafluorophenol Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.054 M of acid at multiple scan rates. (50, 75, 100, 200, 300, 400, 500, 600 and 700 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of II/I 2 the Co wave (Ep,c) with scan rate (). The slope is -36 mV/dec, with R = 0.999. Bottom right: 5 evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.7 × 10 M-1 s-1.

161

Figure B.28 Electrochemical determination of kPT for CoCp(dppe) with 2,4,6-tribromophenol Top: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.046 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: evolution of II/I the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). The slope is 29 2 2 mV/dec, with R = 0.999. Bottom right: evolution of kobs obtained with acid concentration (R = 4 -1 -1 0.999). The slope yields kPT = 1.9 × 10 M s .

162

Figure B.29 Electrochemical determination of kPT for CoCp(dppe) with benzoic acid Top: cyclic voltammograms of 0.5 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 10 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.00099 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. Bottom left: II/I evolution of the cathodic peak potential of the Co wave (Ep,c) with acid concentration (CA). 2 The slope is 29 mV/dec, with R = 0.990. Bottom right: evolution of kobs obtained with acid 2 4 -1 -1 concentration (R = 0.998). The slope yields kPT = 1.5 × 10 M s .

163

Figure B.30 Electrochemical determination of kPT for CoCp(dppe) with glacial acetic acid Top: cyclic voltammograms of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.045 M of acid at multiple scan rates (100, 200, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc II/I couple at 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate (). The slope is -34 mV/dec, with R2 = 0.997. Bottom right: evolution of the obtained 3 -1 -1 kPT at each scan rate with scan rate. The average yields kPT = 3.2 × 10 M s .

164

Figure B.31 Electrochemical determination of kPT for CoCp(dppe) with 4-chloro-phenol Left: cyclic voltammograms of 0.2 mM of [CoCp(dppe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0 – 0.64 M of acid. The CVs are referenced to the Fc+/Fc couple at 0 V and were recorded at 100 mV/s. It has to be noted that given at such high acid concentrations, signs of reaction in the CoIII/II wave were observed. II/I Middle: evolution of the cathodic peak potential of the Co wave (Ep,c) with acid concentration 2 (CA). The slope is 52 mV/dec, with R = 0.999. Right: evolution of kobs obtained with acid 2 2 -1 -1 concentration (R = 0.988). The slope yields an estimated kPT = 1.5 × 10 M s .

165

Figure B.32 Photochemical determination of kPT of CoCp(dppe) using benzoic acid (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of benzoic acid. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression gives a second order 5 -1 -1 protonation rate constant (kPT) of 4.2 × 10 M s .

166

Figure B.33 Photochemical determination of kPT for CoCp(dppe) using 2,3,5,6-tetrafluorophenol (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of 2,3,5,6-tetrafluorophenol. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression gives a second order 6 -1 -1 protonation rate constant (kPT) of 1.4 × 10 M s .

167

Figure B.34 Photochemical determination of kPT of CoCp(dppe) using triethylammonium tetrafluoroborate (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of triethylammonium tetrafluoroborate. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression 5 -1 -1 gives a second order protonation rate constant (kPT) of 3.4 × 10 M s .

168

Figure B.35 Photochemical determination of kPT for CoCp(dppe) using 2,3,4,5,6- pentachlorophenol (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of 2,3,4,5,6-pentachlorophenol. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression gives a second 6 -1 -1 order protonation rate constant (kPT) of 6.2 × 10 M s .

169

Figure B.36 Photochemical determination of kPT for CoCp(dppe) using salicylic acid (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of salicylic acid. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression gives a second order 7 -1 -1 protonation rate constant (kPT) of 3.7 × 10 M s .

170

Figure B.37 Photochemical determination of kPT for CoCp(dppe) using 2,3,5,6-F4-4-CF3-phenol (Top) Transient absorption data for the transient signal at 486 nm with increasing concentrations of 2,3,5,6-F4-4-CF3-phenol. The decay was first to a first order exponential decay in the range shown (residuals are shown). (Bottom) The observed rate constants from the exponential decays are displayed vs acid concentration, and the slope of the linear regression gives a second order 7 -1 -1 protonation rate constant (kPT) of 4.4 × 10 M s .

171

Figure B.38 Electrochemical determination of kPT of CoCp(depe) with phenol Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.047 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 8.7 × 103 M-1 s-1.

172

Figure B.39 Electrochemical determination of kPT of CoCp(depe) with 4-CH3-phenol Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.234 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.8 × 104 M-1 s-1.

173

Figure B.40 Electrochemical determination of kPT of CoCp(depe) with 4-CF3-phenol Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.155 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 7.0 × 105 M-1 s-1.

174

Figure B.41 Electrochemical determination of kPT of CoCp(depe) with 4-Cl-phenol Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.108 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.996. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.2 × 105 M-1 s-1.

175

Figure B.42 Electrochemical determination of kPT of CoCp(depe) with acetic acid Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.087 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -34 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 5.6 × 106 M-1 s-1.

176

Figure B.43 Electrochemical determination of kPT of CoCp(depe) with benzoic acid Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.050 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -37 mV/dec, with R = 0.997. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.6 × 107 M-1 s-1

177

Figure B.44 Electrochemical determination of kPT of CoCp(depe) with 2,3,5,6-tetrafluorophenol Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.049 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields: 7 -1 -1 kPT = 4.4 × 10 M s

178

Figure B.45 Electrochemical determination of kPT of CoCp(depe) with benzylammonium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.046 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.996. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.8 × 107 M-1 s-1

179

Figure B.46 Electrochemical determination of kPT of CoCp(depe) with salicylic acid Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.055 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -36 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 5.6 × 107 M-1 s-1

180

Figure B.47 Electrochemical determination of kPT of CoCp(depe) with 4-CH3O-pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.054 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.994. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.2 × 107 M-1 s-1

181

Figure B.48 Electrochemical determination of kPT of CoCp(depe) with pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.049 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.5 × 107 M-1 s-1

182

Figure B.49 Electrochemical determination of kPT of CoCp(depe) with 4-CH3-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.030 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R = 0.997. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.6 × 107 M-1 s-1

183

t Figure B.50 Electrochemical determination of kPT of CoCp(depe) with 4- Bu-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.036 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 2.9 × 107 M-1 s-1

184

Figure B.51 Electrochemical determination of kPT of CoCp(depe) with 4-Cl-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.022 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 4.5 × 107 M-1 s-1

185

Figure B.52 Electrochemical determination of kPT of CoCp(depe) with 4-CF3O-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.036 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 3.2 × 107 M-1 s-1

186

Figure B.53 Electrochemical determination of kPT of CoCp(depe) with 4-CN-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(depe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.038 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.996. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.5 × 107 M-1 s-1

187

Figure B.54 Electrochemical determination of kPT of CoCp(dcpe) with 4-Cl-phenol Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.790 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 2.6 × 102 M-1 s-1

188

Figure B.55 Electrochemical determination of kPT of CoCp(dcpe) with acetic acid Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.310 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -29 mV/dec, with R = 0.995. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 8.2 × 103 M-1 s-1

189

Figure B.56 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,5,6-tetrafluorophenol Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.063 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -29 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 8.2 × 104 M-1 s-1

190

Figure B.57 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,4,5,6- pentafluorophenol Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.057 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.997. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.1 × 105 M-1 s-1

191

Figure B.58 Electrochemical determination of kPT of CoCp(dcpe) with 2,3,5,6-tetrafluoro-4-CF3- phenol Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.036 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.3 × 105 M-1 s-1

192

Figure B.59 Electrochemical determination of kPT of CoCp(dcpe) with 4-CH3O-pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.036 M of acid at multiple scan rates (50, 60, 75, 90, 100, 125, 150, 200, 300, 400 and 500 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 5.3 × 104 M-1 s-1

193

Figure B.60 Electrochemical determination of kPT of CoCp(dcpe) with pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.047 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 5.9 × 104 M-1 s-1

194

t Figure B.61 Electrochemical determination of kPT of CoCp(dcpe) with 4- Bu-pyridinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.050 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -33 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 5.2 × 104 M-1 s-1

195

Figure B.62 Electrochemical determination of kPT of CoCp(dcpe) with anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.054 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 4.9 × 104 M-1 s-1

196

Figure B.63 Electrochemical determination of kPT of CoCp(dcpe) with 4-Br-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.2 mM [CoCp(dcpe)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.026 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -29 mV/dec, with R = 0.999. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 6.6 × 104 M-1 s-1

197

Figure B.64 Electrochemical determination of kPT of CoCp(dppv) with acetic acid Top: cyclic voltammograms of 0.2 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.241 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -30 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 1.6 × 103 M-1 s-1

198

Figure B.65 Electrochemical determination of kPT of CoCp(dppv) with benzoic acid Top: cyclic voltammograms of 0.2 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.436 M of acid at multiple scan rates (50, 75, 100, 125, 150, 200, 250, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential II/I 2 of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R = 0.998. Bottom right: evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 2.6 × 103 M-1 s-1

199

Figure B.66 Electrochemical determination of kPT of CoCp(dppv) with salicylic acid Top: cyclic voltammograms of 0.2 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.436 M of acid at multiple scan rates (50, 75, 100, 125, 150, 175, 300, 400, 500 and 600 mV/s from black to blue). The CVs are referenced to the Fc+/Fc couple at 0 V. Bottom left: evolution of the cathodic peak potential of II/I 2 the Co wave (Ep,c) with scan rate (). The slope is -31 mV/dec, with R = 0.996. Bottom right: 6 evolution of the obtained kPT at each scan rate with scan rate. The average yields kPT = 2.2 × 10 -1 -1 M s

200

t Figure B.67 Electrochemical determination of kPT of CoCp(dppv) with 4- Bu-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.08 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.0064 M of acid at multiple scan rates (50, 100, 250, and 500 mV/s from black to blue). The CVs are referenced to the Fc+/Fc II/I couple at 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R2 = 0.996. Bottom right: evolution of the obtained 7 -1 -1 kPT at each scan rate with scan rate. The average yields kPT = 2.0 × 10 M s

201

Figure B.68 Electrochemical determination of kPT of CoCp(dppv) with 4-Br-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.08 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.0040 M of acid at multiple scan rates (50, 100, 250, and 500 mV/s from black to blue). The CVs are referenced to the Fc+/Fc II/I couple at 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate (). The slope is -32 mV/dec, with R2 = 0.992. Bottom right: evolution of the obtained 7 -1 -1 kPT at each scan rate with scan rate. The average yields kPT = 1.3 × 10 M s

202

Figure B.69 Electrochemical determination of kPT of CoCp(dppv) with 4-CN-anilinium tetrafluoroborate Top: cyclic voltammograms of 0.08 mM [CoCp(dppv)(NCCH3)][PF6]2 in 5 mL of CH3CN with 0.25 M TBAPF6 as supporting electrolyte in the presence of 0.0050 M of acid at multiple scan rates (50, 250, 750, and 2500 mV/s from black to blue). The CVs are referenced to the Fc+/Fc II/I couple at 0 V. Bottom left: evolution of the cathodic peak potential of the Co wave (Ep,c) with scan rate (). The slope is -33 mV/dec, with R2 = 0.999. Bottom right: evolution of the obtained 7 -1 -1 kPT at each scan rate with scan rate. The average yields kPT = 2.1 × 10 M s

203

APPENDIX C: PHOTOCHEMICAL REACTIVITY OF IRIDIUM HYDRIDE COMPLEXES

Parts of this Appendix adapted with permission from Chambers, M. B.; Kurtz, D. A.; Pitman, C. L.; Brennaman, M. K.; Miller, A. J. M. Efficient Photochemical Dihydrogen Generation Initiated by a Bimetallic Self-Quenching Mechanism. J. Am. Chem. Soc. 2016, 138 (41), 13509–13512. Copyright (2016) American Chemical Society.

The purpose of this Appendix is to summarize time-resolved experiments related to mechanistic investigations of hydrogen evolution from iridium pentamethylcyclopentadienyl bipyridine hydride complexes as a part of an ongoing collaboration with the Miller lab at UNC

Chapel Hill and the Castellano group at North Carolina State University.

C.1 Introduction and Background

Photocatalytic hydrogen evolution involving transition metal complexes requires the combination of two protons and two electrons driven by the energy in solar photons. As described in previous chapters, these systems generally involve separate photosensitizers, catalysts, and sometimes redox meditators to coordinate this intricate transformation.1,2 The separation of components can hinder efficiency, and thus a more ideal system would be one in which a single complex both absorbs light and acts as a catalyst for hydrogen evolution. In this pursuit, the Miller lab at UNC have developed a photoelectrocatalytic system for hydrogen evolution wherein a transition metal hydride complex can be generated electrochemically, then

3 photolyzed in the presence of acid to generate H2 (Scheme C.1). In addition to being the first report of molecular photoelectrocatalysis for H2 evolution, related complexes have been studied in the context of photochemical water-gas shift reactivity,4,5 formic acid dehydrogenation,6 proton transfer processes7, and even dark-catalysis like transfer hydrogenation reactions.8–10

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C.2 Results and Discussion

The origin of the involvement of the Dempsey lab in this collaboration stems from photoelectrocatalytic data collected by Dr. Catherine Pitman whilst studying hydrogen evolution from the complex [IrCp*(bpy)Cl]Cl under aqueous conditions.

Scheme C.1 Photoelectrochemical hydrogen evolution using iridium Cp* bipyridine complexes

While studying photoelectrocatalytic efficacy of [IrCp*(bpy)Cl]Cl, chronoamperograms were recorded under steady-state illumination and the maximum catalytic rate constant (kcat) were determined.3 While the most common mechanism for hydrogen evolution would include a first order dependence on the catalyst, an interesting second-order concentration with respect to

11 iridium catalyst was observed for the formation of H2 under these conditions. While these initial studies were performed in aqueous buffered solution, studies shortly thereafter showed

III + that the isolated hydride complex [HIrCp*(bpy)][PF6], [HIr ] , was photoactive in CH3CN as

12 III + well. Blue light irradiation of [HIr ] in CH3CN in the presence of a number of organic acids

(pyridinium tetrafluoroborate, triethylammonium tetrafluoroborate, and acetic acid) cleanly generates H2, with no reactivity observed on similar time scales in the dark. Thermodynamic parameters were calculated, and it was shown that, thermodynamically, [HIrIII]+ is both a stronger photoacid and photohydride donor in the excited state as compared to the ground state.

Interestingly, however, there is no difference in the rate of H2 formation for the three acids studied (pyridinium tetrafluoroborate, triethylammonium tetrafluoroborate, and acetic acid),

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which span a pKa range of 11 pKa units. This lack of dependence on acid strength indicates that acid dependence protonation occurs after the rate limiting step.

Kinetic labeling experiments were performed with excess deuterium labeled acetic acid

III + and [HIr ] (Figure C.1). At very early photolysis time points, H2 was formed at higher concentrations than HD, which is further support of a bimetallic mechanism as H2 could only be formed from two [HIrIII]+ molecules interacting directly with each other prior to reaction with

CD3COOD.

1 Figure C.1 Concentration of H2 (▲) and HD (◆) detected by H NMR spectroscopy during the III + photolysis of 11.5 mM of [HIr ] and 100 mM of CD3CO2D in CD3CN with 460 nm light. Inset shows the first 1 min. H2 concentration corrected for thermal population of para-H2.

In order to investigate the excited state of [HIrIII]+, a variety of steady state and time resolved spectroscopic methods were utilized. The ground state absorption spectrum of yellow

[HIrIII]+ exhibits an asymmetric absorption band that is centered at 428 nm, which has been assigned to the multiple MLCTs from the Ir(III) dπ orbitals to the π* orbital of the bipyridine ligand (Figure C.2).13 Excitation into this band results in a broad, weak emission at room temperature centered at 708 nm. The large stokes shift of this emission feature leads to the assignment of a triplet emissive state. Indeed, recent computational work supports this assignment.13

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III + Figure C.2 Room temperature absorbance and emission spectrum of [HIr ] in CH3CN. For emission spectrum, λexc = 428 nm.

Ultrafast transient absorption spectroscopy experiments performed by Chelsea Taliaferro at NCSU have shed some more light on the nature of the excited state.14 Upon excitation at 400 nm, transient features grow in on the timescale of 10’s of picoseconds to yield the spectrum as shown in Figure C.3. The large transient feature centered at 360 nm is assigned on π  π* transition of the reduced bpy•–, and the two features at 480 nm and 510 nm are assigned to π* 

π* transitions of the reduced bpy•–. These assignments are based on the similar energy and

•– 15,16 extinction coefficients as compared to previously reported bpy complexes.

Figure C.3 Transient absorption spectrum taken at various delays (picoseconds to nanosecond) after excitation of [HIrIII]+.

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There is a striking similarity of the excited state spectrum of [HIrIVCp*(bpy•–)]+* to the

•– 17 ground state absorbance spectrum of Na[Fe(bpy )(bpy)2]. Interestingly, in the absorbance

•– spectrum of Na[Fe(bpy )(bpy)2], there is not a distinct “trough” in between the main absorbance features at ~360 nm and at ~500 nm, and instead sustained, featureless absorbance with molar extinction coefficients of ~10,000 M-1 cm-1. If similar featureless absorbance at that energy is present in the absorbance spectrum of [HIrIVCp*(bpy•–)]+*, it would explain the absence of a noticeable ground-state bleach at ~420 nm. Of note, the UV/vis transient absorption spectrum of

IV •– + [HIr Cp*(bpy )] * measured in CH3CN qualitatively matches the partial spectrum measured by

Fukuzumi et al. in methanol.7 The first order rate constant measured for intersystem crossing to form the thermally equilibrated excited state in methanol is 1.4 × 1010 s-1, which is qualitatively

IV •– + consistent with the growth of [HIr Cp*(bpy )] * in CH3CN.

Time-resolved photoluminescence was used to measure the radiative lifetime of emission from [HIrIVCp*(bpy•–)]+*. Observation of emission decay at 708 nm after excitation at 445 nm revealed a relatively short lifetime of ~80-100 ns. Single wavelength time-resolved transient absorption kinetic measurements were recorded probing at 370 nm and 480 nm, and the lifetime of decay of the transient signals were very similar to that of the photoluminescence decay lifetime (Figure C.4).

The photoluminescence lifetime was invariant with respect to acetic acid concentration.

Interestingly, however, the TRPL lifetime was dependent on the concentration of [HIrIII]+; as the concentration increased the lifetime decreased (Figure C.5, left).

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Figure C.4 (Top) Time-resolved photoluminescence decay of excited state of [HIrIII]+, and (middle and bottom) time-resolved transient absorption kinetic profiles at various wavelengths showing the decay of the excited-state transients of the excited state of [HIrIII]+. For all experiments, λexc = 445 nm.

Further experimentation revealed a linear dependence of lifetime on concentration of chromophore, indicating a dynamic “self-quenching” mechanism wherein the excited state

[HIrIVCp*(bpy•–)]+* is quenched by a ground state molecule of [HIrIII]+. Stern-Volmer analysis

9 -1 -1 gives a second order quenching rate constant of kq = 3.5 × 10 M s and an intrinsic, theoretical

6 -1 decay rate constant in the absence of quencher of k0 = 9.7 × 10 s (Figure C.5, right).

Figure C.5 (Left) Time-resolved photoluminescence decay at different concentrations of III + III + [HIr ] , and (right) Stern-Volmer plot of kobs vs [HIr ] concentration to determine kq and k0

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The net reaction between two molecules of [HIrIII]+ and one photon of light to produce

I 2+ H2 yields one equivalent of IrCp*(bpy) (Ir ) and one equivalent of [IrCp*(bpy)(NCCH3)]

III 2+ ([Ir -NCCH3] ). In the absence of acetic acid, relatively rapid decomposition of the laser samples was observed. This is attributed to previously reported reactivity between generated IrI

III 2+ and [Ir -NCCH3] activating the acetonitrile solvent to form a bridging amino-acyl species, which absorbs stronger than the [HIrIII]+ and interferes with transient absorption measurements.12 Therefore, laser experiments were performed in the presence of excess acetic acid that protonates IrI and increase the longevity of the samples. Additionally, due to the

III 2+ + III + buildup of the ligated [Ir -NCCH3] or [IrCp*(bpy)(OAc)] ([Ir -OAc] ) during the laser experiment, sodium triacetoxyborohydride (Na[HB(OAc)3], STAB) was added in excess in order

III + III 2+ III + to regenerate [HIr ] . The reaction between STAB with both [Ir -NCCH3] and [Ir -OAc] was probed independently, and it was found that the nitrile complex converted to [HIrIII]+ faster than the acetate complex.

While these experiments provided clear evidence that the [HIrIII]+ excited state was directly involved in bimetallic reactivity, the mechanism of quenching needed to be ascertained.

Triplet-triplet annihilation (reactivity between two triplet excited states) is a plausible mechanism for self-quenching, but was discounted out due to invariance of the excited-state lifetime on intensity of excitation pulse. Deuterated [DIrCp*(bpy)][PF6] was prepared, and while its intrinsic lifetime was larger as compared to [HIrIII]+, the lack of KIE (1.0(1)) indicates that there is no bond-breaking occurring in the quenching step. This, along with the absence of evidence of radical products, discounts the possibility of excited state Ir-H bond homolysis which has been observed previously. Adding TBAPF6 supporting electrolyte to the laser samples resulted in a decrease the in the lifetime, indicating more efficient self-quenching, and supports a

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quenching mechanism of excited-state electron transfer. Satisfyingly, with this knowledge in hand, previously performed photolysis experiments were performed in the presence of added electrolyte, and the quantum yield for hydrogen production (ΦH2) was doubled.

Probing the excited-state absorbance of the excited state of [HIrIII]+ at 370 nm and 480 nm after excitation on longer timescales reveals that the transient does not return to baseline after excited-state, but instead persists on the timescale of hundreds of microseconds. Additionally, probing out in the red at 625 nm (where there is no excited state absorbance) reveals the growth of a transient signal that persists on the same timescale as the long-lived transient signals at 370 nm and 480 nm (Figure C.6).

Figure C.6 Overlay of TRPL decay of the excited state of [HIrIII]+ at 708 nm and the TA kinetics of growth of transient signal assigned to IrI probed at 625 nm.

The rate constant for the growth of the transient signal matches the rate constant for the decay of the excited state very well. Absorbance values were then measured at 10 µs at a number of wavelengths across the visible region (Figure C.7, left). These absorbance values were

I III 2+ overlaid with a difference spectrum of the generated Ir plus [Ir -NCCH3] minus ground state

[HIrIII]+, which shows excellent agreement with the transient absorption data (Figure C.7, right).

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Figure C.7 (Left) Kinetic traces at various wavelengths of [HIrIII]+, dashed line is 10 µs, and I III 2+ III + (right) absorbance values at 10 µs overlaid with (Ir + [Ir -NCCH3] ) – [HIr ] difference spectrum

Additionally, the decay of the transient signal (λprobe = 650 nm) accelerated with increasing acetic acid concentration (Figure C.8, left).

I Figure C.8 (Left) Acid dependence of the decay of transient signal assigned to Ir (λprobe = 650 nm) and plot of kobs vs [HOAc] showing second order dependence on acid.

We therefore assign the next observable spectroscopic intermediate following excited- state self-quenching as IrI, followed by observation if its protonation to reform [HIrIII]+. The decay of the IrI signal in the presence of >10 equiv of acetic acid fits nicely to a single exponential decay, and a plot of kobs vs [HOAc] shows a second order dependence on acid

(Figure C.8, right). This has been observed in our lab previously, and is attributed to

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dimerization of HOAc prior to proton delivery.18 The third order rate constant for the protonation of IrI with HOAc is 4.5 × 105 M-2 s-1. This protonation of IrI to form [HIrIII]+, accompanied

III 2+ III + with the hydride transfer from STAB to [Ir -NCCH3] to form [HIr ] effectively “resets” the system for the next laser shot.

III + Figure C.9 Summary of mechanism of catalytic H2 evolution upon excitation of [HIr ] in the presence of acid and STAB

C.3 Conclusions

In summary, we have described and characterized an efficient mechanism for hydrogen generation from a molecular iridium hydride complex, HIrIII. Upon the absorption of a photon of light, the excited state of the complex is self-quenched via electron transfer to form two very reactive proposed species that couple to generate hydrogen and the next spectroscopically observable intermediate, IrI. This IrI species is protonated by excess acid in solution to reform the ground state HIrIII. This mechanism is unprecedented for molecular hydrogen evolution, and opens up exciting new avenues for catalyst design.

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