Investigation of Thermal Products from Photochemically Generated Intermediates of a Diiron-Diphosphine Complex; Relevance to Fefe-Hydrogenase

Investigation of Thermal Products from Photochemically Generated Intermediates

of a Diiron-Diphosphine Complex; Relevance to FeFe-Hydrogenase

Anne Elizabeth Nelson

A thesis submitted to

Sonoma State University

in partial fulfillment of the requirements

for the degree of

MASTER OF SCIENCE

Interdisciplinary Studies

Committee Members:

Dr. Carmen Works, Chair*

Dr. Jennifer Whiles

Dr. Joseph Lin

05 May 2017

Anne E. Nelson

Authorization for Reproduction of Master’s Thesis

I grant permission for the print or digital reproduction of this thesis in its entirety, without further authorization from me, on the condition that the person or agency requesting reproduction absorb the cost and provide proper acknowledgment of authorship.

DATE: 05 May 2017 Anne Elizabeth Nelson Name

iii

Investigation of Thermal Products from Photochemically Generated Intermediates of a Diiron-Diphosphine Complex; Relevance to FeFe-Hydrogenase

Thesis by Anne Elizabeth Nelson

ABSTRACT

The biohydrogen cycling enzymes known collectively as hydrogenases are shaping the vision of industrial hydrogen catalysis. This family of enzymes reversibly oxidizes molecular hydrogen. Inorganic structural and functional models of the [FeFe]- hydrogenase active site provide a convenient and simple approach to studying changes that mimic the active site chemistry. Presented herein, (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) and + {μ-H(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] ) were photochemically altered to observe thermal changes specific to various solvent and solute mixtures.

In this study continuous wave photolysis was used to generate reactive intermediates in organic solutions. The resulting thermal chemistry was monitored by infrared and ultraviolet-visible spectroscopy.

It was determined that the photochemically generated CO unsaturated photoproducts of [2] and [2H]+ thermally combine with solvent or dissolved CO producing semi-stable complexes in solution.

This work supports the hypothesized irradiation-dependent CO loss and offers evidence for the correlation between solvent or solutes and stability of thermal products.

MS Program: Interdisciplinary Studies Sonoma State University Date: 05 May 2017

Dedication

For my mother

who has been to the ends of the earth for me

many, many times.

Acknowledgements

My graduate studies are a reflection of the supportive community I have encountered at Sonoma State University and the many people who helped to shape my experiences there. Firstly, I wish to thank my graduate research advisor, Professor

Carmen Works. Through relentless questioning and tough love she drew from me the scholarly confidence I desired to possess and so admired in her. I am forever grateful for the opportunities she provided me with and for her role in facilitating my pursuit of graduate studies. I also appreciate her willingness to seek funding to sustain my studies.

From the bottom of my heart, I thank my committee member and undergraduate academic advisor, Professor Jennifer Whiles for her guidance and compassion. She is a woman of true grit who understands the fundamental importance of inspiring young female scientists to invest in themselves. I would have been lost many times over without her reassurance and encouragement.

Thank you to Professor Jon Fukuto who has been instrumental in my development as a scholar. He taught me not to back down in the face of a question I could not answer, but rather to view it as an opportunity to better myself.

I am appreciative of Professor Joseph Lin for serving on my committee and challenging me to think about my project from a different prospective.

I am grateful for my lab mate and close friend, Patricia De La Torre, most notably for her significant contributions to this project and constant encouragement.

I thank past instructors, Shauna Ferdinandson, for teaching me the fundamentals of writing, and Roger Wilson, whose advanced biology class marked my decision to decision to pursue higher education in the physical sciences.

Finally, I am indebted to my family for loving me through this difficult process.

To my mother, Susan Nestor, thank you for always telling me I can do anything, sacrificing your life for mine and stressing the value of education. And to my devoted husband, Mark Nelson, who watched me fall apart again and again. His love was the glue that held me each time I put myself back together.

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Table of Contents

Chapter Page

I. Introduction to Hydrogenase Enzymes and Model Complexes

1.1 Hydrogen fuel……………………………………………………………………………………….1 1.2 The Structure and Function of Hydrogenase Enzymes………………………………..1 1.2.1 Active site of the [FeFe]-Hydrogenase………………………………………..5 1.2.2 Mechanism of Hydrogen Activation in [FeFe]-Hydrogenase…………7 1.2.3 Regulation and Oxygen Sensitivity of Hydrogenases……………………9 1.3 Hydrogenase Models……………………………………………………………………………..9 1.3.1 Photochemical Studies of (μ-dt)[Fe(CO)3]2………………………………11 1.3.2 Photochemical Studies of (μ-pdt)[Fe(CO)2(PMe3)]2……………………12

1.3.3 Photochemical Studies of (μ-pdt)[HFe2(CO)4(PMe3)2][PF6]………...14 1.4 Conclusion………………………………………………………………………………………….15

II. Preparation and Characterization of Diiron Hydrogenase Model Complexes

2.1 Instrumentation…………………………………………………………………………………...17 2.2 Materials..………………………………………………………………………………………….17 2.3 Methods...... 18

2.3.1 Synthesis of (μ-pdt)[Fe(CO)3]2 ([1])...... 18

2.3.2 Synthesis of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2])...... 18 + 2.3.3 Synthesis of {(μ-H)(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] )...... 20

2.4 Characterization of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2])...... 20 2.4.1 UV-Vis...... 20 2.4.2 FTIR...... 21 2.4.3 NMR...... 25 2.4.4 FDMS...... 28 + 2.5 Characterization of {μ-H(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] )...... 28 2.5.1 UV-Vis...... 28 2.5.2 FTIR...... 28 2.5.3 NMR...... 33 2.6 Conclusion...... 33

III. Photochemical Behavior of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) in the Infrared

3.1 Introduction...... 36 3.2 Methods...... 37 3.3 Results...... 39 3.3.1 THF/Air...... 39

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3.3.2 THF/N2...... 41 3.3.3 THF/CO...... 44 3.3.4 THF/H2...... 44 3.3.5 Summary of Spectral Changes in THF...... 44 3.3.6 MeCN/N2...... 47 3.3.7 MeCN/CO...... 47 3.3.8 Summary of Spectral Changes in MeCN...... 47 3.3.9 Toluene/N2...... 50 3.3.9.1 Toluene/N2 with Residual Phosphine and THF...... 50 3.3.9.2 Toluene/N2/Phosphine...... 50 3.3.9.3 Bulk Solution Photolysis...... 53 3.3.10 Toluene/CO...... 53 3.3.11 Toluene/H2...... 56 3.3.12 Summary of Spectral Changes in Toluene……………...... 56 3.3.13 DCM/N2………………………………………………...... 56 3.3.14 DCM/CO……………………………………….…………59 3.3.15 Summary of Spectral Changes in DCM…………………..59 3.4 Discussion………………………………………………………...... ………..59 3.4.1. Comparative Interpretation of THF and Toluene Solvent Studies..59 3.4.2. Analysis of MeCN and DCM Efficacy as Solvent Choices………63 3.5 IR Conclusion...... 64

IV. Photochemical Behavior of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) in the Optical

4.1 Introduction...... 66 4.2 Methods...... 67 4.3 Results...... 68 4.3.1 Toluene/N2...... 68 4.3.2 Toluene/CO...... 68 4.3.3 Toluene/phosphine/N2...... 68 4.3.3.1 Residual Phosphine...... 68 4.3.3.2 Added Phosphine...... 72 4.3.5 THF/N2...... 72 4.3.6 CO Detection by [(RhPPh3OAc)2  (H2O)2] ...... 76 4.4 Discussion...... 76 4.4.1 Solute Effects...... 76 4.4.2 Qualitative CO Detection...... 81 4.5 Conclusion...... 81

V. Thermal Behavior of Photochemically altered {(μ-H)(μ-pdt) [Fe(CO)2(PMe3)]2 + [PF6]} ([2H] ) in the Infrared

5.1 Introduction...... 83 5.2 Methods...... 84 5.3 Results...... 84

5.3.1 DCM/N2...... 84 5.3.2 DCM/H2...... 84 5.3.3 DCM/CO...... 87 5.3.4 Acetone/N2...... 87 5.3.5 Acetone/CO...... 91 5.4 Discussion...... 94 5.4.1 Photochemical Changes of [2H]+ in DCM Solutions...... 94 5.4.2 Thermal Reversibility of CO Release in DCM...... 95 5.4.3 Photochemical Changes of [2H]+ in Acetone Solutions...... 95 5.4.4 Thermal Reversibility of CO Release in Acetone...... 96 5.5 Conclusion...... 97

VI. A Brief Look at the Photolysis of (pdt)[Fe(CO)2(PMe3)]2 ([2]) and {(μ-H)(μ- + pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] ) by NMR

6.1 Introduction...... 99 6.2 Methods...... 100 6.3 Results...... 100 6.3.1 Photolysis of [2] With Excess PMe3...... 100 6.3.2 Thermal Decomposition of [2] With Excess PMe3...... 102 + 6.3.3 Photolysis of [2H] with N2...... 102 6.3.4 Photolysis of [2H]+ with CO...... 102 6.4 Discussion...... 104 6.4.1 Observed Behavior of [2] by 31P NMR...... 106 6.4.2 Observed Behavior of [2H]+ by 31P NMR...... 107 6.5 Conclusion...... 107

VII. Conclusion

7.1 Conclusion...... 108 References...... 111

List of Figures

Figure Page

1. Taxonomic distribution of [FeFe]-hydrogenases...... 3 2. Active site of two [FeFe]-hydrogenases...... 6 3. H-cluster of the [FeFe]-hydrogenase from Chlamydomonas reinhardtii...... 8 4. Structural models of the [2Fe]H active site...... 13 5. UV-Vis spectrum of [2] in DCM with λmax at 348 nm...... 22 6. IR spectrum of [2] in MeCN...... 23 7. IR spectrum of [2] in toluene...... 24 1 8. H NMR spectrum of [2] in d6-acetone...... 26 31 9. P NMR spectrum of [2] in d6-acetone...... 27 10. FDMS spectrum of [2]...... 29 11. UV-vis spectrum of [2H]+ in DCM...... 30 12. IR spectra of [2H]+ in three solvents...... 31 13. IR of [2H]+ in THF...... 34 1 + 14. H NMR of [2H] in CDCl3...... 35 15. Diagram of irradiation set up for IR experiments...... 40 16. Photolysis of [2] in THF/air at 365 nm for 20 minutes...... 42 17. Photolysis of [2] in THF/N2 at 365 nm for 40 minutes...... 43 18. Photolysis of [2] in THF/CO at 365 nm for 40 minutes...... 45 19. Photolysis of [2] in THF/H2 at 365 nm for 30 minutes...... 46 20. Photolysis of [2] in MeCN/N2 at 365 nm for 25 minutes...... 48 21. Photolysis of [2] in MeCN/CO at 365 nm for 40 minutes...... 49 22. Photolysis of [2] in toluene/N2 at 365 nm for 30 minutes...... 51 23. Photolysis of a reflux solution of [2]...... 52 24. Photolysis of [2] in toluene/phosphine/N2 at 365 nm for 30 minutes...... 54 25. Photolysis of [2] in toluene/CO at 365 nm for 20 minutes...... 55 26. Photolysis of [2] in toluene/H2 at 365 nm for 30 minutes...... 57 27. Photolysis of [2] in DCM/N2 at 365 nm for 15 minutes...... 58 28. Photolysis of [2] in DCM/CO at 365 nm for 25 minutes...... 60 29. UV-vis monitored photolysis of [2] in toluene/N2 at 365 nm for 100 seconds...... 70 30. UV-vis photolysis of [2] in toluene/CO for 100 seconds...... 71 31. UV-vis photolysis of [2] in toluene/N2 in the presence of residual PMe3 for 100 seconds...... 73 32. UV-vis photolysis of [2] in toluene/N2 with 50 μL excess PMe3 for 80 seconds...... 74 33. UV-vis photolysis of [2] in THF/N2 for 120 seconds...... 75 34. UV-vis photolysis of [2] in THF/N2 for 240 seconds...... 77 35. Photolysis with a CO probe...... 78 + 36. Photolysis of [2H] in DCM/N2 for 25 minutes...... 85 + 37. Photolysis of [2H] in DCM/H2 for 25 minutes...... 86 38. Photolysis of [2H]+ in DCM/CO for 20 minutes...... 88 39. Back reaction of [2H]+ in DCM/CO...... 89

+ 40. Photolysis of [2H] in acetone/N2 for 30 minutes...... 90 41. Photolysis of [2H]+ in acetone/CO for 25 minutes...... 92 42. Back reaction of [2H]+ in acetone/CO...... 93 31 43. Photolysis of [2] in d6-benzene/N2 monitored by P NMR...... 101 44. Thermal degradation of [2] in d6-benzene/N2...... 103 + 45. Photolysis of [2H] in CDCl3/N2...... 104 31 + 46. P NMR of [2H] in d6-acetone/CO mixture...... 105

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List of Schemes

Scheme Page

1. Catalytic model of hydrogen activation with respect to the redox states of the H-cluster as shown by Mulder, et al...... 10 2. Photochemically induced CO loss in [2]...... 38 3. Reaction of binuclear rhodium(II) complex with CO...... 69

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List of Tables

Table Page

1. IR extinction coefficients for peaks in [2H]+ in three solvents………………………32

xiv

List of Abbreviations

Ar argon ca. circa

CDCl3 deuterated chloroform cm-1 inverse centimeters

CO carbon monoxide d6-acetone deuterated acetone

DCM dichloromethane et al. and others

FDMS field desorption mass spectrometry

H2 dihydrogen i.e. id est

IR infrared

LUCA last universal common ancestor

M molar

MeCN acetonitrile mg milligram mL milliliter

MLCT metal to ligand charge transfer

MMCT metal to metal charge transfer mmol millimole m/z mass to nuclear charge ratio

N2 dinitrogen

NH4PF6 ammonium hexafluorophosphate

NMR nuclear magnetic resonance

O oxygen

P phosphorous pdt propanedithiol

PMe3 trimethylphosphine ppm parts per million rota-vap rotary evaporation

THF tetrahydrofuran

TRIR time resolved infrared

V volts

VT variable temperature vide infra see below

σ sigma

°C degrees Celsius

[1] (μ-pdt)[Fe(CO)3]2

[2] (μ-pdt)[Fe(CO)2(PMe3)]2

+ [2H] {μ-H(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]}

xvi 1

Chapter 1

Introduction to Hydrogenase Enzymes and Model Complexes

1.1 Hydrogen fuel

The increasing need to develop non-carbon based energy sources has inspired the investigation of alternative “biofuels” like hydrogen. Hydrogen is an ideally clean and potentially renewable fuel because its only byproduct is water:

2 H2 + O2 ⇌ 2 H2O

Equation 1.

+ Coupling H reduction by electrochemical cells for storage with H2 oxidizing catalysts for fuel is one model for efficient cycling of hydrogen energy. The most competent catalytic hydrogen production relies on platinum electrodes, making this system cost prohibitive for large-scale industrial use; additionally, common trace

1 impurities in H2 gas irreversibly inactivate the platinum. Replacing precious metal catalysts with earth abundant metals like iron may be the solution to functionalizing hydrogen-based fuel.

1.2 The Structure and Function of Hydrogenase Enzymes

+ - H2 ⇌ 2 H + 2 e

Equation 2.

Hydrogenases exist in the three domains of life (eukarya, archaea, and bacteria) with high occurrence in anaerobic bacteria. While communities of microorganisms in anoxic environments are estimated to produce and consume 200 million tons of H2

3 annually, only trace amounts exist in the atmosphere suggesting that H2 production, rather than consumption, is rate limiting.4 By remaining in close proximity, hydrogenase- bearing microorganisms further increase efficiency as one organism can utilize the H2 produced by another.

The diversity of hydrogenases is exemplified by their evolution and widely distributed taxonomy, and is illustrated in Figure 1. Phylogenic reconstructions date some hydrogenases back to the last universal common ancestor (LUCA), indicating an essential role for this enzyme in the metabolism of early life.5 Additionally, organisms employ different types of hydrogenases to satisfy various energetic demands or dispose of metabolic waste. These primary processes that require hydrogenase activity include methanogenesis, nitrogen fixation, fermentation, reduction of sulfur species, and photosynthesis, to name a few.5 The presence of multiple hydrogenases within a single organism, while seemingly redundant, signifies an ability to rapidly respond to changing environments.6

Figure 1. Taxonomic distribution of [FeFe]-hydrogenases (A) and [NiFe]-hydrogenases (B)5. Red lines indicate 1 or more homologous enzymes. All divergent lines can be traced to the last universal common ancestor (LUCA).

Hydrogenases are metalloenzymes rich in Fe-S clusters. They are classified by the metals in their active sites and known generally as the [NiFeSe]-, [NiFe]-, and [FeFe]- hydrogenases. A fourth class known as the [Fe]-only hydrogenase contains no Fe-S clusters and couples H2 uptake to methenyltetrahydromethanopterin reduction as an intermediary step in methanogenesis (methane is produced when methanogenic

7 archaebacteria metabolize CO2 for growth).

Hydrogenases are central to the energy metabolism of microorganisms and are capable of driving unfavorable reactions through electron bifurcation, offering reducing

5 equivalents, and providing hydrogen as a final electron acceptor (H2 production). The

[NiFe]- and [FeFe]-hydrogenases have been isolated from numerous sources and have been the focus of many structural and functional studies.5, 8 Interestingly, these two classes likely evolved independently (see Figure 1) but employ the same unusual ligands and accomplish similar chemistry. [FeFe]-Hydrogenases are evolutionarily younger than the [NiFe] class and show 100 fold more activity.9 Although bidirectional, hydrogenases show preference for either hydrogen oxidation or proton reduction. Indeed this is evidenced by the [FeFe]-hydrogenase isolated from Desulfovibrio desulfuricans which can consume 28,000 molecules of H2 per second versus 9,000 molecules of H2 produced in the same amount of time.2a While the active sites of the [FeFe] and [NiFe] classes share similar features, the [FeFe] class enzymes generally contain fewer auxiliary [Fe-S] clusters, which makes crystal structures and spectroscopic data (i.e. Mössbauer, EPR, IR) easier to interpret. The [FeFe]-hydrogenases will be the focus of the remainder of this chapter.

1.2.1 Active site of the [FeFe]-Hydrogenase

Structural studies of the [FeFe]-hydrogenase date back to 1946 when O. Warburg observed that H2 formation in growing cultures of Clostridium butyricum was inhibited by CO and this inhibition was reversed with light.10 Warburg suggested an active site consisting of Fe complexed by S ligands, which has since been confirmed by the crystal structures of the [FeFe]-hydrogenases from Clostridium pasteurianum11 and

Desulfovibrio desulfuricans (Figure 2).2c

Enzymes of the [FeFe] class vary in architecture but are among the least complex of the hydrogenase family, particularly those from algal and clostridial sources, making them apt for maturation studies and functional assays. The active site is sequestered in a hydrophobic pocket and relies on a relay of Fe-S clusters or ferredoxin to shuttle electrons to or from the protein surface.8a

The catalytic subunit, termed the H-domain, contains the [6Fe-6S] H-cluster that is comprised of a [4Fe-4S] cubane subcluster linked by a cysteine thiolate to a novel

[2Fe-2S] (distinguished as [2Fe]H) subcluster, which is the active site (Figure 3). The distorted octahedral Fe atoms of the unique organometallic active site are coordinated to two terminally bound CO ligands, two terminal (and biologically unusual) CN- ligands, and a third semi-bridging CO ligand. A labile water molecule solvates the Fe distal to the

[4Fe-4S] subunit. The diiron core of the active site is bound by a bridging di(thiomethyl)amine. The long debated amine in the apex of this bridge has not been resolved by crystal structure, but computational studies, ERP, and functional assays

Figure 2. Active site of two [FeFe]-hydrogenases (A) from C. pasteurianum11 selected amino acid residues in the polypeptide environment of the [2Fe] subcluster (Fe, rust; S, yellow; C, black; O, red; N, blue). The [4Fe-4S] subcluster and associated ligand are included (gray) for perspective (C, Cys; F, Phe; K, Lys; S, Ser; M, Met); and (B) from D. desulfurians2c with Fe-Fe bridge modeled as 1,3 propanedithiol and each Fe coordinated to putative CO and CN- ligands. ‘X’ represents a water derived solvent molecule (Fe, orange; S, green; C, yellow; N, blue; O, red).

7 support this assignment.12 Three highly conserved cysteine containing motifs are essential for coordination of the H-cluster5 but a fourth cysS-Fe bond (shown in Figure 3) provides the only direct covalent link between the active site and the rest of the protein.11

1.2.2 Mechanism of Hydrogen Activation in [FeFe]-Hydrogenase

+ While hydrogenases can accept both H and H2 as substrates, the [FeFe] class shows preference toward proton reduction to form H2. The H-cluster has been identified in a number of redox states as characterized by the cubane unit ([4Fe-4S]H) and the respective proximal (Fep) and distal (Fed) irons of the [2Fe]H subunit. Recent work by

Mulder et al. shows the first evidence of a metal hydride (using a combination of

Mössbauer, DFT, and EPR monitored redox titration) and its role in the catalytic cycle.13

2+ II I As shown in Scheme 1, in the resting Hox state, the H-cluster ([4Fe-4S] -Fe -Fe ) may become protonated by solvent and Fep accepts an electron to form the one electron reduced state (Hred); intramolecular proton transfer from the bridgehead amine to Fed results in a terminal hydride ([4Fe-4S]1+-FeII-FeII).13 A second protonation event at the hydride forms H2, which is removed via gas channels, and physiological electron donors regenerate Hred.

2+ II I Conversely, the Hox state ([4Fe-4S] -Fe -Fe ) supports direct binding of H2 at Fed

(Scheme 1). The pendant amine acts as an internal base and heterolytic cleavage of H2

5 results in protonation of the amine and a terminal hydride bound to [Fe]d. Hydride electrons may be transported via [4Fe-4S]H to additional Fe-S clusters or ferredoxin depending on the hydrogenase.

The enzyme may also exist as the Hox-CO inhibited form when a labile water molecule (shown as OH- in Figure 3) is replaced by a CO ligand. The strong Fe-CO

cys NH

[Fe4S4] S S S OH Fe Fe OC p d CN C O CO NC

Figure 3. H-cluster of the [FeFe]-hydrogenase from Chlamydomonas reinhardtii. Fep and Fed indicate proximal and distal (respectively) relative to the [Fe4S4] cluster, and OH indicates a labile water species. The Fep-S bond originates from the adjacent cysteine that covalently links the [2Fe]H active site to the protein.

9 interaction blocks hydrogen metabolism, but curiously, the active enzyme is regenerated upon white light irradiation.14

1.2.3 Regulation and Oxygen Sensitivity of Hydrogenases

Oxygen, CO, and hydrogen compete for the active site in [FeFe]-hydrogenases.

The majority of hydrogenases are oxygen intolerant. While CO inhibition is reversible, oxygen bound to the [2Fe]H catalytic unit irreversibly disrupts the [4Fe-4S]H subcluster, inactivating enzyme. Although most hydrogenases are housed in anaerobic bacteria, a particularly interesting division of the [FeFe] class is the oxygen tolerant variety found in photosynthetic green algae. Molecular oxygen negatively regulates hydrogenase synthesis in anaerobic microorganisms, while molecular hydrogen activates hydrogenase expression in aerobic and photosynthetic bacteria.6 In eukaryotic green algae the presence of H2 indicates compromised photosynthesis or lack of O2 and initiates upregulation of hydrogenase genes. The algae will adapt to light-dependent hydrogen metabolism which uses photosystem I (PSI) and Ferredoxin PetF to remove low potential electrons in response to anaerobic stress.15

1.3 [FeFe]-Hydrogenase Models

Organometallic models of the [2Fe]H active site of [FeFe]-hydrogenase have been well established and studied in a variety of contexts.12a, b, 16 These structural and functional analogues have made numerous advances in the last two decades as potential hydrogen evolving catalysts, particularly since publication of the X-ray crystal structures of two bacterial [FeFe]-hydrogenases.2c, 11 Models of the active site vary in complexity but can be simplified by the general formula (μ-Rdt)[Fe2L6] (Figure 4A). This minimal model is an ideal synthetic scaffold that readily undergoes a variety of modifications such as ligand substitutions and bridgehead alterations to enhance desired electrochemical

Scheme 1. Catalytic model of hydrogen activation with respect to 13 the redox states of the H-cluster as shown by Mulder, et al. Hox = oxidized cluster (red); Hox[H2] = H2 bound to the oxidized cluster (red); Hhyd = terminal hydride bound to the oxidized cluster (blue); Hsred = super (2 electron) reduced cluster (green); Hred = one electron reduced cluster (purple).

11 properties. While much emphasis has been placed on electrochemical reduction of protons, these models are also photochemically active.

1.3.1 Photochemical Studies of (μ-pdt)[Fe(CO)3]2

Among the active site models, one of the most studied is (μ-pdt)[Fe(CO)3]2 ([1])

(pdt = 1,3-propanedithiol) and is shown in Figure 4B. In this model the diiron moiety is bridged by a propanedithiol ligand (SCH2CH2CH2S), and each Fe coordinates three terminal CO ligands. This model bears structural similarity to the [2Fe]H subunit and mimics the Hox-CO inhibited enzyme and each iron is formally Fe(I). As with the active enzyme, irradiation of [1] stimulates release of a CO ligand.16i While the resulting pentacarbonyl species structurally represents the Hox active site, it cannot act as a functional model for H2 oxidation. Access to the Fe(II) oxidation state is essential to catalytic activity in the models but [1] and its photoproducts lack the basicity required to support formation of a bridging hydride,17 thus excluding it from hydrogen turnover assays. Although, [1] has been useful for establishing general photochemical behavior of hydrogenase models.

Initially chosen for its structural simplicity, work with [1] pioneered photochemical and mechanistic studies aimed at elucidation of details related to solvent and irradiation wavelength in hydrogenase models. Irradiation of [1] in organic solutions at 355 nm has been shown to release a CO ligand, which is either replaced by a solvent

16c, 16e, 16i, molecule in coordinating solvents or remains unsaturated as (μ-pdt)[Fe2(CO)5].

18 When this system is irradiated in the presence of excess dissolved CO, the photogenerated (μ-pdt)[Fe2(CO)5] forms a solvento species that exchanges with dissolved CO to reform [1] with solvent dependent kinetics in a fast process; also the unsaturated (μ-

pdt)[Fe2(CO)5] excited state can decay into a second product that reforms with CO in a slow, CO-independent process.16i

1.3.2 Photochemical Studies of (μ-pdt)[Fe(CO)2(PMe3)]2

As mentioned above, the synthetic scaffold [1] readily undergoes direct modification. Replacement of two CO ligands yields the diphosphine substituted (μ- pdt)[Fe(CO)2(PMe3)]2 ([2]) (Figure 4C). Both physical and computational studies have investigated events immediately surrounding photo-induced CO loss. Theoretical calculations by Johnson et al. predict five energetically feasible ground-state isomers of

19 [2], and 20 different CO-loss photoproducts with respect to the two FeL3 subunits.

Parallel experiments conducted by the same group monitored spectral changes in the IR on the picosecond timescale. Following laser flash photolysis at 355 nm, spectra were recorded at various time-points up to 700 ps. A short-lived excited state species (formed when [2] undergoes CO loss) appeared immediately but decayed into a different ground state isomer. Transient difference spectra in acetonitrile (MeCN) show a pronounced growth at 1865 cm-1 in the first 100 ps after photolysis that decays back to baseline.

Heptane solutions, by contrast, showed a distinct growth at 1997 cm-1 that persistently increased up to 700 ps post photolysis (marking the end of the experiment).

A weak growth at 1870 cm-1 appears at ca. 50 ps but decays by 200 ps. These data suggest that MeCN and heptane form weakly coordinating, short-lived solvent adducts as evidenced by the growths near 1870 cm-1.

By modeling different ratios of specific isomers Johnson determined that in some instances the labilized CO did not escape the solvent cage, thus the parent complex

R S S S S OC L L CO Fe Fe Fe Fe L L OC CO L L OC CO A) B)

+ = S S S S S S PMe3 CO PMe3 CO NC CN Fe Fe Fe Fe Fe Fe CO CO OC CO OC OC H CO OC PMe3 OC PMe3 OC

C) D) E)

Figure 4. Structural models of the [2Fe]H active site. A) General formula of synthethic models (μ-Rdt)[Fe(L)3]2 B) (μ-pdt) [Fe(CO)3]2 ([1]) C) (μ-pdt)[Fe(CO)2(PMe3)]2 )[2]) D) (μ- + + 2- pdt)[HFe2(CO)4(PMe3)2] ([2H] ) E) (μ-pdt)[Fe(CO)2(CN)]2

reformed, but as a different isomer due to the pseudo rotation of respective FeL2 subunit.19-20 This creates a complex mixture of parent isomers, excited state species, reconfigured ground state isomers and photoproducts. These numerous arrangements complicate experimental spectra and explain discrepancies with computed spectra.

Further evidence of these solution mixtures is supported by variable temperature

(VT) NMR studies done by Zhao and coworkers. Upon cooling to -60 °C different conformers can be spectroscopically isolated.21 VT NMR in the CO region of the 13C spectra also revealed intramolecular ligand exchanges in the reflecting the rotational fluidity of the ligands, but site exchanges are localized to the respective FeL3 units (with

21-22 L = CO or PMe3).

1.3.3 Photochemical Studies of (μ-pdt)[HFe2(CO)4(PMe3)2][PF6]

Ligand substitutions are central to catalytic studies of hydrogenase models.

While [1] does not support protonation of the Fe-Fe bond, replacing 2 CO ligands with a

- better electron donor, such as CN or PMe3, makes protonation possible, facilitating formation of a bridging hydride (Figure 4D).16d, 21 Upon protonation, each Fe undergoes a

I II formal change in oxidation from Fe to Fe , a prerequisite for hydrogen oxidation or reduction by [2Fe]H models. While there is no evidence of a bridging hydride in biological systems, it is essential to the activation of hydrogen in the models. In addition

21 to assisting H2 binding, the bridging hydride acts as an internal base consistent with the proposed role of the amine in the natural enzyme.

The reduction potential of the [FeII FeII] core for the [2H]+ 0/- couple is less negative (-1.4 V) than that of the corresponding neutral [2]0/- (-2.3 V) as referenced against ferrocene+/0 in MeCN solution.23 Milder reduction potentials allow substituted hydride analongs of [1] to be linked with photosensitizers that produce H2 upon UV

15 irradiation under acidic conditions. Interestingly, the diphosphine substituted hydride derivatives do not require sensitizers (like [Ru(bipy)3]), in contrast to the vast majority of photo-hydrogen-evolving catalysts.23 With growing interest in these models, investigation of their photo-reactivity is gaining attention, but thermal behavior of photochemically generated intermediates remains a weak point in this field. Ultra-fast studies offer information about the immediate behavior of these systems in response to photolysis, but long-term stability must be considered as well if these models are to have catalytic applications.

A significant study by Zhao et al. examined a series of light driven H/D exchange reactions catalyzed by [2H]+ and other similar models. Upon irradiation with white light or prolonged exposure to ambient light (over the course of days) solutions of [2H]+

II II II II 21 exchanged [Fe -H-Fe ] with [Fe -D-Fe ] and formed HD from H2 and D2. These reactions only proceeded after exposure to light and rate of HD formation (monitored by

1H and 2H NMR) increased as the intensity of the irradiation source increased. While this work signifies the importance of light to generating the active catalyst in solution, characterization of the light was not investigated. Because different wavelengths excite different bonds, using a specific irradiation wavelength allows for probing of a particular interaction within the molecule, which has yet to be done on this timescale.

Intramolecular CO site exchanges did not occur in [2H]+ consistent with higher energy barriers associated with the octahedral geometry21- a property that perhaps contributes to the greater thermal stability with respect to [2].

1.4 Conclusion

Hydrogenase enzymes provide an essential function in the metabolism of microorganisms, particularly anaerobic bacteria, oxidizing or reducing hydrogen as

16 required to meet energy demands. Biohydrogen cycling microbiomes inspired synthetic models of the [FeFe]- hydrogenase active site, which show great promise as environmentally clean and renewable hydrogen catalysts. These potential catalysts are conveniently activated by light. While hydrogenase enzymes do not require direct irradiation to catalyze H2 oxidation, it is a desirable route to access the Hox active state in model complexes. The extreme oxygen sensitivity remains a difficulty in utilizing both the enzymes and models. Much of the work in this field has focused on the electrochemical properties of the enzyme mimetic molecules, but clear details regarding the photochemistry are needed to fully understand catalysis. The work presented herein pertains to the preparation of two [FeFe]-hydrogenase model complexes and their thermal behavior in solution following photoactivation.

Chapter 2

Preparation and Characterization of Diiron Hydrogenase Model

Complexes

2.1 Instrumentation

Ultraviolet-visible spectra were obtained using an Agilent 8453 diode array spectrophotometer. Infrared spectra were collected on a Nicolet Avatar 370 DTGS FTIR.

An Agilent 400 MHz NMR was used to acquire 1H and 31P nuclear magnetic resonance spectra, and chemical shifts are reported in ppm. The instrument used to collect mass spectrometry data was an Agilent 7890 Gas Chromatograph equipped with a Waters GCT

Premier high resolution Time-of-flight mass spectrometer using a field desorption (FD) ion source, with an upper limit of detection of m/z 4000. All instruments were used at room temperature.

2.2 Materials

Triiron dodecacarbonyl and 1,3-propanedithiol were purchased from Acros

Organics and used as received. All solvents were also obtained from Acros Organics.

Anhydrous solvents, including toluene, tetrahydrofuran (THF), dichloromethane (DCM), and acetonitrile (MeCN), were received in nitrogen sparged containers which were sealed

18 with septa. Spectroscopy grade acetone (not dry) was used as well. Deuterated acetone was acquired in ampoules. Deuterated chloroform was not dry and used as received or degassed by purging as needed. Trimethylphosphine was bought from Sigma-Aldrich as a

1 molar (1 M) solution in THF or toluene. Ammonium hexafluorophosphate salt

(NH4PF6) was obtained from Sigma-Aldrich. Concentrated hydrochloric acid (HCl, 8 M) was purchased from Fischer Scientific.

2.3 Methods

2.3.1 Synthesis of (μ-pdt)[Fe(CO)3]2 ([1])

24 Following published methods, (μ-pdt)[Fe(CO)3]2 ([1]) was prepared. In short, using standard Schlenk and vacuum line techniques, ca. 1.0 g of Fe3(CO)12 (1.99 mmol) was dissolved in 50 mL toluene and 0.40 mL of 1,3-propanedithiol (pdt) (4 mmol) was added by syringe. The solution was refluxed under a N2 atmosphere for 1.5 hours.

Reaction completion was determined by comparing IR bands in the metal-CO region to previously reported bands.18a In toluene, peaks were observed at 2073, 2032, 2000, 1988, and 1975 cm-1. Upon cooling, the air stable solution was filtered on the bench top using a

Büchner funnel to remove residual starting material. Excess solvent was removed by rota-vap. Two successive silica gel columns (toluene followed by hexane) were run to further purify the product. Final removal of solvent by rota-vap produced lusterful red- orange crystals. Average yields ranged from 55-65 %.

2.3.2 Synthesis of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2])

Following published methods17, 230 mg of [1] (0.58 mmol) was dissolved in 15 mL toluene in a degassed 3 neck round bottom flask fit with a reflux condenser. To the heated solution, 2.3 mL of PMe3 (23 mmol, 1 M in THF) was added via syringe. The

reaction mixture was refluxed under a N2 atmosphere for 24 hours. Reaction progress was monitored by IR, and a color change from red-orange to deep brownish-red was observed. Bands characteristic of the reaction mixture were seen at 1981, 1944, and 1901 cm-1. Excess unreacted starting material was removed by cannula filtration. Solvent was removed by rota-vap and the resulting red-brown crystals were dried overnight in vacuo to ensure complete removal of reaction toluene. Average yields ranged from 80-85 %.

While the synthesis and characterization of [2] has been established in literature 17,

21 many difficulties were encountered and resolved. Firstly, the 24-hour reflux proved problematic because the condenser did not remain cold overnight, resulting in total solvent evaporation. Switching from hexane to toluene did not resolve the issue, even if the reflux was kept at low heat. Using a circulating water pump with a thermostatic unit resolved the problem of solvent evaporation.

It was originally assumed that [2] would possess similar thermal stability to the precursor [1]. However, while [1] can be kept for several months and is air-stable in solid form, it has been determined that [2] has a much shorter shelf-life and should be synthesized in small batches as needed. Although, there is a fine line when scaling down because while 0.5 mmol [1] yields the desired product, using 0.05 mmol [1] does not, and the product remains sticky and impure. Solid samples of [2] maintain integrity for at most one month if stored at -4 °C under N2 and minimally exposed to oxygen. A sample stored in an argon glove box degraded in half the time despite precautions to minimize oxygen exposure. Ideally, samples of [2] should be used within one week, with all precautions taken to minimize oxygen exposure and thermal decomposition.

Degradation of [2] was observable when the crystals changed from a luster-full deep red-purple to a dull brownish-purple. Unfortunately, peaks in the IR and UV-vis

20 often gave little indication of sample integrity, thus the most reliable way to gauge sample quality was by examining the single resonance in the 31P NMR signal which appeared at 22.89 ppm in d6-acetone (see 2.4.3).

+ 2.3.3 Synthesis of {(μ-H)(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] )

Literature methods also guided the synthesis of μ-H(μ-pdt) [Fe(CO)2(PMe3)]2[PF6]

([2H]+) with little adaptation.17 In an Ar glove box 62 mg of [2] (1.28 mmol) was transferred to a 3 neck round bottom flask. The reaction vessel was further manipulated on the bench-top under a N2 atmosphere. The solid was dissolved with 7 mL of degassed

MeOH. To the stirring solution 1.2 mL of 8 M HCl (10 mmol) was added by syringe.

Upon addition of the acid, a color change was noted from red-brown to red-orange, and a black precipitate formed. The mixture was stirred for one hour then cannulated to remove solid impurities. The solution volume was minimized by reduced pressure. A saturated aqueous solution of NH4PF6 was added dropwise to precipitate out an orange solid. The solid was washed alternately with ether and water three times. Excess solvent was cannulated away and the solid was dried in vacuo with around a 90 % yield. Complex

[2H]+ appears to be more stable than [2] both as a solid and in solution, but must still be treated as air sensitive to prolong its shelf-life.

2.4 Characterization of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2])

2.4.1 UV-Vis

UV-vis spectra for characterization purposes were collected using a standard quartz cuvette with a 1 cm pathlength. Figure 5 shows the UV-vis spectrum of [2] with a λmax at ca. 348 nm (depending on solvent), which is red-shifted 23 nm from [1] (data not shown), as well as a weaker shoulder around 520 nm. The prominent band at 348 nm (ε = 14,051

M-1cm-1) was previously assigned as a metal-to-ligand charge transfer, while the shoulder

(ε = 1,516 M-1cm-1) was assigned as a metal-to-metal charge transfer. 19

2.4.2 FTIR

In the IR, compounds were monitored and distinguished from one another by the positions of their peaks in the metal-carbonyl stretching range (approximately 1850 cm-1 to 2150 cm-1). Compound [2] displayed three peaks centered at 1900, 1942, and 1979 cm-1 in more polar solvents such as MeCN (Figure 6), THF, and DCM. In toluene, the highest frequency peak blue shifts 5 wavenumbers to 1981 cm-1 and an additional peak is observed at 1971 cm-1 (Figure 7). These peaks were consistent with literature values for

[2];17 such solvent effects have been observed by others19 and also with the starting material [1].18a Additional peaks indicate a mixture of unreacted starting material and product, which is easily identified by the presence of a peak at 2073 cm-1 not present in

[2].

Using Beer’s law, the extinction coefficients were calculated for the three peaks in

THF for solutions of known concentrations and found to be 421 M-1cm-1 (1981 cm-1);

3,630 M-1 cm-1 (1944 cm-1); 2,090 M-1 cm-1 (1901 cm-1).

3.0

2.5

2.0

a b

r 1.5

b A

1.0

0.5

0.0 300 400 500 600 Wavelength (nm)

Figure 5. UV-Vis spectrum of [2] in DCM with λmax at 348 nm

Figure 6. IR spectrum of [2] in MeCN showing peaks at 1979, 1942, and 1900 cm-1

0.8

0.6

c n

a 0.4

b A

0.2

0.0

2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm )

Figure 7. IR spectrum of [2] in toluene showing peaks at 1981, 1971, 1942, and 1900 cm-1

2.4.3 NMR

Monitoring the 1H resonances for [2] proved problematic due to the paramagnetic

Fe-Fe core. While spectra of [2] are obtainable (Figure 8), they exhibit peak broadening and poor resolution, making assignments in this region challenging. Ultimately, proton peaks in d6-acetone were assigned based on their integration values and matched as closely as possible with peaks reported by Xhao et al.17 Methylene protons in the apex of the bridgehead appeared as a shallow doublet at 1.4 ppm (SCH2CH2CH2S); the equivalent methyl groups on the phosphine ligands came in as a broad singlet at 1.5 ppm

(P(CH3)3); and a small broad peak at 1.8 ppm was attributed to the two methylenes adjacent to the sulfur atoms (SCH2CH2CH2S).

Unlike the 1H peaks, the proton decoupled 31P spectrum provides more definite information. Freshly synthesized samples of [2] exhibit a single peak at 22.89 ppm when dissolved in d6-acetone (Figure 9). Although one might expect the spectrum to reflect the unique environments of the phosphine ligands, a single peak is consistently observed.

This is attributed to the fluxionality of the propanedithiol bridgehead and a mixture of conformational isomers in solution which present as an average.21

When a sample of [2] has degraded the single phosphorous resonance broadens slightly and a second smaller peak at ca. 38 ppm appears. This new peak indicates inequivalent phosphorus environments, and while it is assumed the compound has oxidized,25 degradation products of [2] have not been characterized. Similarly, a failed synthesis exhibits the same secondary phosphorus peak when oxygen has contaminated the reaction.

1.5

1.8 1.4

1 Figure 8. H NMR spectrum of [2] in d6-acetone. Labeled peaks indicate the following protons: (SCH2CH2CH2S) at 1.8 ppm, (P(CH3)3) at 1.5 ppm, (SCH2CH2CH2S) at 1.4 ppm.

31 Figure 9. P NMR spectrum of [2] in d6-acetone showing a single resonance at 22.89 ppm

2.4.4 FDMS

Time of flight FDMS was preformed to further verify successful synthesis of [2].

The positive ionization technique revealed a primary mass at 482 m/z normalized to

100% (Figure 10). Five lesser peaks were observed between 480 and 485 m/z. The distribution of these peaks was compared to a simulated spectrum, which was calculated based upon the molecular formula of [2] using a free program available through

Scientific Instrument Services, Inc. The two spectra were in agreement with respect to peak number, position and relative intensities.

+ 2.5 Characterization of {μ-H(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] )

2.5.1 UV-Vis

The protonated complex [2H]+ absorbs in the UV with a peak at 285 nm and a shoulder at 325 nm (Figure 11). The peak at 285 nm is largely overshadowed by intense absorbance, which may be due to solvent. Attempts to resolve the peak at 325 nm by increasing concentration were unsuccessful. Transitions have not been assigned for this species and no literature comparison exists.

2.5.2 FTIR

+ The IR spectrum of [2H] (Figure 12) is characterized in DCM by two distinct νCO bands at 2032 cm-1 and 1990 cm-1. These two peaks are present in acetone (2030 cm-1,

1989 cm-1) and MeCN (2030 cm-1, 1988 cm-1) as well. The extinction coefficients for the two primary peaks were the greatest in DCM (3,666 M-1cm-1, 2,718 M-1cm-1 respectively). Extinction coefficients for all three solvents are reported in Table 1.

Figure 10. FDMS spectrum of [2] with primary mass at 481.92 m/z

+ Figure 11. UV-vis spectrum of [2H] in DCM

3000

)

- M

( DCM

t 2000

e i

c Acetone

ﬁ f

e o

C MeCN

x E

1000

2200 2150 2100 2050 2000 1950 1900 1850 -1 Wavenumbers (cm )

Figure 12. IR spectra of [2H]+ in three solvents: DCM (green), acetone (red), MeCN (blue). Peaks located at approximately 2030, 1990 cm-1 (exact values vary with solvent).

Solvent Peak/ cm-1 ε/ M-1 cm-1

DCM 2032 3666

1990 2718

MeCN 2030 3257

1989 2331

Acetone 2030 3037

1988 2056

Table 1. IR extinction coefficients for peaks in [2H]+ in three solvents

The charged nature of the hydrido species prohibits use of nonpolar solvents such as toluene and benzene. Solubility of [2H]+ is poor in the moderately polar solvent THF, and the best absorbing peak in the IR will not exceed an absorbance of ca. 0.4 (Figure

13). As seen with [1] and [2], solvent interactions with [2H]+ are reduced in less polar solvents; thus, in addition to two primary peaks at 2031 cm-1 and 1985 cm-1, two secondary, albeit weak, peaks appear at 1944 cm-1, 1901 cm-1.

2.5.3 NMR

Successful synthesis of [2H]+ was corroborated by the presence of a small triplet

1 centered at -14.7 ppm in the H NMR (CDCl3) which is characteristic of the bridging hydride (FeHFe) (Figure14).21 The downfield region of the 1H spectrum contained additional, unidentifiable peaks, similar to [2] (Figure 8). As discussed above (2.4.3), monitoring this region of the spectrum is difficult and did not provide any information pertinent to the present study.

The 31P spectrum of [2H]+ in d6-acetone is nearly identical to the spectrum of [2], but shifts upfield slightly (from 22.89 ppm) centering at 21.07 ppm (data not shown).

2.6 Conclusion

Following established synthesis routes, with minor adaptations to improve yield and purity, compounds [1], [2], and [2H]+ were successfully synthesized and characterized by a variety of spectroscopic methods. Special care should be taken to prepare the phosphine derivatives only as needed for immediate experiments, and samples should ideally be used within one week. Compounds must be stored as solids in the absence of oxygen and kept cold to prolong structural integrity.

Figure 13. IR of [2H]+ in THF showing peaks at 2031, 1985, 1844, 1901cm-1

-14.71 ppm

1 + Figure 14. H NMR of [2H] in CDCl3. Inset: expanded view of (FeHFe) bridging hydride triplet centered at -14.71 ppm

Chapter 3

Photochemical Behavior of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) in the

Infrared

3.1 Introduction

Although the photochemical activity of the neutral (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) has been established and investigated by others,19-20 [2] exhibits no catalytic activity.21 As such, much of the work involving [FeFe]-hydrogenase models focuses on the cationic

[2H]+ and similar analogs (adaptations of [1] containing a bridging hydride with alterations to the S to S linked bridgehead and/or ligand substitutions). Furthermore, among the photochemically centered studies that do focus on [2], many are computationally based or on the ultrafast timescale, leaving an opportunity for wavelength specific studies which monitor changes on a much longer (minute) timescale.

While ultrafast studies are important in their own right, extended time scales reveal macroscopic information about the robustness and lasting behavior of the complex in solution. Such knowledge is pertinent to the advancement of industrial scale catalysis with these models.

The studies presented below employ continuous wavelength photolysis for extended periods of time (several minutes) as a means of generating reactive intermediates and photoproducts for the mechanistic investigation of the solution phase chemistry of [2]. Probing the photo- and thermal chemistry of [2] aids in validating a mechanism proposed for the catalytically relevant [2H]+.

The objective of the studies presented in this chapter was to spectroscopically characterize the solution phase chemistry of [2] using a specific irradiation wavelength.

In consideration of the well characterized photochemistry of [1] which loses a CO ligand16c, 16i, 18a similar hypotheses were drawn regarding the photolability of a CO ligand, formation of a solvento species, and the thermal regeneration of the parent compound in the presence of excess dissolved CO. As shown in Scheme 2 it is hypothesized that light generates some excited state, ultimately leading to the release of a single CO (or other) ligand. The resulting open coordination site (represented by ‘X’) allows for the binding of a solvent molecule (or other ligand in solution) or remains as an open coordination site.

This hypothesis was tested using THF and MeCN. The behavior of [2] in non- coordinating solvents was modeled with toluene and DCM.

3.2 Methods

Samples for IR experiments were prepared by degassing the solid (ca. 2 mg) in a

5 mL round bottom flask sealed with a rubber septum. The flask was evacuated and back filled with the appropriate gas (N2, CO, or H2) for three cycles. A separate flask

38 containing 1-2 mL of the desired solvent was saturated with the corresponding gas by means of purging. For this method of solute displacement, CO concentrations are

26 assumed to be ca. 7 mM and ca. 3 mM for H2 concentrations. These concentrations were

* S S S S Me P Me P 3 CO light 3 CO CO Fe Fe Fe Fe OC CO OC CO

OC PMe3 OC PMe3

S S Me3P Fe Fe OC CO

solvent 'X' OC PMe3

S S + CO Me3P X Fe Fe OC CO Thermal Degradation OC PMe3

Scheme 2. Photochemically induced CO loss in [2]. Light generates an excited state leading to CO loss. The unsaturated photoproduct can accept a coordinating solute or solvent molecule (‘X’) that exchanges with CO to reform some isomer of the parent. All paths eventually result in thermal decomposition.

calculated with Henry’s law based on values reported at 293 K. However, solubility of these solutes is not explicitly reported in the literature for all solvents used in the present study, and room temperature was highly variable from day to day. Therefore, these approximations qualitatively illustrate solution mixtures with respect to ratios of [2] and participating solutes.

Using a gastight Hamilton syringe, the solvent was transferred to the flask containing the solid to dissolve it such that an absorbance of between 0.6-1.0 in the metal-CO region was achieved and then syringed into a N2 purged IR cell.

Samples were collected using a gas tight, sealed IR cell (International Crystal

Laboratories) with CaF2 windows (pathlength 0.2 mm). The injection ports on the cell were stoppered with rubber septa and further wrapped with parafilm to reduce evaporation and gas exchange. Samples were photolyzed directly in the cell using a handheld Hg lamp with a short wavelength of 254 nm and a long wavelength of 365 nm.

All studies were preformed using the 365 nm setting. The lamp was positioned at close range to the cell (5-7 cm) to ensure complete exposure (Figure 15). Typical photolysis protocol was continuous exposure in five minute intervals over the course of 20 to 30 minutes total. Spectra were collected immediately following each photolysis interval and/or at regular intervals post-photolysis to monitor thermal changes. Photolysis times and intervals were adjusted to suit specific experiments (see Results 3.3).

3.3 Results

3.3.1 THF/Air

Hg lamp IR cell

5 cm

Figure 15. Diagram of irradiation set up for IR experiments. Depicted is the portable Hg lamp positioned at a distance of approximately 5 cm from the sealed IR cell containing the solution of interest.

When samples of [2] were photolyzed in THF in the presence of oxygen, a single peak at 2001 cm-1 grew in, while the parent peaks decreased in intensity (Figure 16).

Post-photolysis, all peaks continued to decrease in intensity due to expedited thermal degradation, which is a trend observed regardless of solvent. Dark controls (non- photolyzed aliquots of the solution that were covered with a black cloth) revealed less extensive absorbance decreases of the three parent peaks. Facile oxidation of the Fe-Fe core or formation of partially reduced oxygen species that affect the thiolate bridgehead likely contributes to the thermal instability of [2] both as a solid and in solution.25

Because oxygen expedites the degradation of [2], air studies were not pursued and have been excluded from further analysis.

3.3.2 THF/N2

Experiments performed under a N2 atmosphere were useful as a point of reference for further experiments in which CO and H2 were introduced. Also, exhaustive photolysis was performed to gauge how long the compound could withstand irradiation. A 0.72 mM solution of [2] was saturated with N2 and photolyzed in 10 minute intervals for 40 minutes (see methods 3.2). Under these anoxic conditions, a photo-growth at 2001 cm-1 appeared in the first 10 minutes and reached an approximate absorbance capacity of 0.02

(Figure 17). The next 30 minutes of photolysis had little effect on further growth of this peak. A second, broad peak appeared at 1868 cm-1. The growth of this peak was comparable in absorbance to the 2001 cm-1 peak. The parent peaks at 1900 cm-1 and 1944 cm-1 decreased in intensity but to a lesser extent than when oxygen was present, and

42 spectral changes became less over time. The peak at 1979 cm-1 decreased slightly though remained relatively unchanged.

0.1

1.2 0.0

S B

A -0.1 ∆

-0.2 1.0 -0.3 2100 2050 2000 1950 1900 1850 1800 -1 cm

0.8

o s

b 0.6 A

0.4

0.2

0.0 2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 16. Photolysis of [2] in THF/air at 365 nm for 20 minutes (5 min per line)

Inset: difference spectrum shows total changes during photolysis 1.0 0.2

0.1 S

B 0.0

A ∆

0.8 -0.1

-0.2 2100 2050 2000 1950 1900 1850 1800 -1 cm

0.6

b A

0.4

0.2

0.0

2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 17. Photolysis of [2] in THF/N2 at 365 nm for 40 minutes (10 min per line) Inset: difference spectrum shows total changes during photolysis

3.3.3 THF/CO

Mimicking the above photolysis protocol (see 3.3.2) when the solvent was sparged with CO (ca. 7 mM)26 prior to photolysis spectral changes (Figure 18) were grossly similar to the THF/N2 data but with two distinctions. Firstly, the growth at 2001 cm-1 formed more gradually and spacing between spectral lines was more apparent than the THF/N2 conditions (Figure 17). The concentration of [2] here (1.6 mM) was approximately double that of the previous experiment and the peak at 2001 cm-1 grew proportionally. Secondly, although the parent peaks experienced consistent bleaching, the growth at 1868 cm-1 exhibited a minimal increase in absorbance.

3.3.4 THF/H2

A sample of [2] was dissolved in a THF/H2 mixture and photolyzed in two 15 minute intervals for a total exposure of 30 minutes (Figure 19). Concentration of H2 was estimated to be 3 mM,26 about double the concentration of [2], which was calculated to be 1.5 mM. Changes in the presence of H2 looked similar to THF/CO in which the three parent peaks decrease in intensity while a peak at 2001 cm-1 grew in and, to a lesser extent, a peak at 1868 cm-1 did as well.

3.3.5 Summary of Spectral Changes in THF

There are two main features in the photospectrum of [2] in THF. The first is a photo-growth at 2001 cm-1 representing a photoproduct characterized as the loss of a CO ligand. The growth of this peak is attenuated in the presence of excess dissolved CO. The second feature is a growth at 1868 cm-1, and is the coordination of a solvent molecule to the unsaturated photoproduct.

0.2

0.1

0.0

S B

A -0.1 1.5 ∆ -0.2

-0.3

-0.4 2100 2050 2000 1950 1900 1850 1800 -1 cm

1.0

b A

0.5

0.0

2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 18. Photolysis of [2] in THF/CO at 365 nm for 40 minutes (10 min per line) Inset: difference spectrum shows total changes during photolysis

0.1

0.0

B A

1.5 ∆ -0.1

-0.2

2100 2050 2000 1950 1900 1850 1800 -1 cm

1.0

b A

0.5

0.0

2100 2050 2000 1950 1900 1850 1800 Wavenumbers (cm-1)

Figure 19. Photolysis of [2] in THF/H2 at 365 nm for 30 minutes (15 min per line)

Inset: difference spectrum shows total changes during photolysis

3.3.6 MeCN/N2

A sample of [2] was dissolved in N2 saturated MeCN and photolyzed in five minute intervals for 25 minutes total. Spectral changes were characterized by decreasing parent peaks and a 2001 cm-1 centered photo-growth (Figure 20). As a coordinating ligand, spectral changes representing a solvento species were anticipated. This would appear as a growth at 1868 cm-1 (molecules capable of coordination consistently affect specifically this region of the spectrum) yet in MeCN there is only the slightest indication in this region that a solvent product may be forming. The spectral bleaching of the parent peaks become less spaced and begin to overlap, and while this can indicate a stable photoproduct, lack of an obvious photo-induced spectral growth and immediate precipitate formation suggests that photolysis of [2] is not compatible in MeCN.

3.3.7 MeCN/CO

A minimal amount of [2] was dissolved in an MeCN/CO solution and photolyzed in five minute intervals for 20 minutes, followed by an additional 20 minutes of continuous photolysis. The IR spectrum showed no growth of a peak at 2001 cm-1, or in any other region of the metal-carbonyl stretch (Figure 21). While the parent peaks decrease in absorbance as a result of photolysis, initial decrease is less and further changes are more gradual than those observed in the MeCN/N2 solution. After 20 minutes of photolysis no further changes were observed.

3.3.8 Summary of Spectral Changes in MeCN

In MeCN, samples of [2] showed slightly unexpected changes. Upon photolysis, parent peaks show immediate decreases in absorbance but little change overall. The CO

0.8 100

50 S

B 0

A ∆

-50

-3 0.6 -100x10 2100 2050 2000 1950 1900 1850 1800 -1

a b

r 0.4

b A

0.2

0.0

2100 2050 2000 1950 1900 1850 1800 Wavenumbers (cm-1)

Figure 20. Photolysis of [2] in MeCN/N2 at 365 nm for 25 minutes (5 min per line) Inset: difference spectrum shows total changes during photolysis

1.5

0.15

0.10

0.05

S B

A 0.00 ∆

-0.05

-0.10

-0.15 2100 2050 2000 1950 1900 1850 1800 -1 cm

1.0

b A

0.5

0.0 2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 21. Photolysis of [2] in MeCN/CO at 365 nm for 40 minutes: 5 min per line for first 20 min, then 20 min continuous Inset: difference spectrum shows total changes during photolysis

50 reporter ligands do not indicate formation of a solvento species. The CO-loss peak at

2001 cm-1 in MeCN behaves similarly to the THF solutions (3.3.2, 3.3.3).

3.3.9 Toluene/N2

Samples of [2] photolyzed in a toluene/N2 solution behave similarly to the coordinating solvents described above (see Results 3.3.2). The parent peaks bleach and a small peak at 2001 cm-1 appears (Figure 22). After 25 minutes of interval photolysis no further changes were observed. No growth at 1868 cm-1 appeared, however, occasionally a peak did present in this region and varied in shape, prompting further experiments, which are described below.

3.3.9.1 Toluene/N2 with Residual Phosphine and THF

Immediately following the synthesis of [2] (but prior to filtration or solvent removal) a 50 μL aliquot was drawn up from the reaction vessel and diluted with toluene to achieve a desired absorbance of approximately 1.0 (1.68 mM). The solution was photolyzed in five minute intervals for 20 minutes total. In addition to the typical spectral changes described above, the peak centered at 1868 cm-1 (broad) appeared (Figure 23). A sample of [2] that has not been completely dried prior to being re-dissolved in toluene

(i.e. the sample is gummy and contains residual phosphine and THF from the reflux) shows identical spectral changes.

3.3.9.2 Toluene/N2/Phosphine

To further distinguish whether the changes were due to THF or excess phosphine ligand, a thoroughly dried sample of [2] was dissolved in N2 saturated toluene with no

THF. The solution was intentionally spiked with 20 μL of phosphine solution (1 M in toluene) and photolyzed in five minute intervals for thirty minutes. To exhaust all photo-

1.0

0.2

0.1

0.0

B A ∆ -0.1

0.8 -0.2

-0.3 2100 2050 2000 1950 1900 1850 1800 -1 cm

0.6

b A 0.4

0.2

0.0

2100 2050 2000 1950 1900 1850 1800 Wavenumbers (cm-1)

Figure 22. Photolysis of [2] in toluene/N2 at 365 nm for 30 minutes (5 min per line)

Inset: difference spectrum shows total changes during photolysis

Fe-PMe Tol N2 @365nm 52

2.0

0.2

0.1 S

B 0.0

A ∆

-0.1 1.5 -0.2 2100 2050 2000 1950 1900 1850 1800 -1

c n

a 1.0

b A

0.5

0.0

2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 23. Photolysis of a reflux solution of [2]. A solution containing toluene, THF, PMe3, and N2 was photolyzed at 365 nm for 25 min (5 minutes per line). Inset: difference spectrum shows total changes during photolysis

53 changes the sample was photolyzed for an additional 50 min. This photolysis brought about extensive bleaching of all three of the parent peaks (Figure 24). The 1901 cm-1 parent peak gave way to two new, sharp peaks centered at 1868 cm-1 and 1880 cm-1, passing through a discrete isobestic point. Interestingly, the small photogrowth normally observed in toluene at 2001 cm-1 did not form well, looking rather broad and shallow.

3.3.9.3 Bulk Solution Photolysis

The unique spectrum produced from the above conditions (see 3.3.9.2) as shown in Figure 24 showed potential as a stable photoproduct that could be isolated and characterized. Dry samples of [2] were dissolved in 2 mL toluene/N2 and 100 μL PMe3

(1M in toluene) was added. The entire sample was irradiated for 10, 20 or 30 minutes in a quartz cuvette (1 cm pathlength). An aliquot of the photolyzed solution was sampled by

IR seeking to verify phosphine coordination after which solvent was removed from the entire stock by rotary evaporation or reduced pressure on the Shlenk line. The dried photoproduct was taken up in toluene and sampled by IR again. The spectrum (not shown) was the same as Figure 22 at 30 minutes and the overall changes are represented by the difference spectrum in Figure 22.

3.3.10 Toluene/CO

When [2] was photolyzed in a toluene/CO mixture in five minute intervals (20 minutes total exposure) the parent peaks decreased in absorbance and no photo-growths appeared (Figure 25). Several attempts were also made to photolyze [2] and then regenerate the parent complex thermally in the dark. Varying total photolysis times, intervals of photolysis, concentrations of [2], and concentrations of CO all yielded the

1.4

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0.8

s b

A 0.6

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0.2

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2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 24. Photolysis of [2] in toluene/phosphine/N2 at 365 nm for 30 minutes (5 min per line), then 20 min continuous, followed by 30 min continuous (80 min total). Inset: difference spectrum shows total changes during photolysis

0.2

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-0.3

-0.4 2100 2050 2000 1950 1900 1850 1800 -1 cm

1.0

b A

0.5

0.0 2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 25. Photolysis of [2] in toluene/CO at 365 nm for 20 minutes (5 min per line) Inset: difference spectrum shows total changes during photolysis

56 same result; although [2] is stable enough in solution to observe photochemically induced changes (particularly so the presence of excess phosphine), we hypothesize that irreversible thermal degradation prevents regeneration of the parent complex in solution.

3.3.11 Toluene/H2

A sample of [2] dissolved in a toluene/H2 mixture was photolyzed for 30 minutes in five minute intervals to observe changes specific to H2. Under these conditions the spectrum (Figure 26) resembles that of the toluene/CO mixture (see Figure 25). A slight increase centered at 1868 cm-1 appears adjacent to the 1901 cm-1 parent peak and, while no conclusion can be definitively made, it hints at the possibility of H2 coordination.

Others have assigned IR absorptions at 2300-2900 cm-1 to the H-H stretch27 and while a broad absorbance grows in this region during photolysis in H2 containing solutions, it also appears in the presence of N2, offering no further evidence of H2 binding.

3.3.12 Summary of Spectral Changes in Toluene

Spectral changes observed in toluene mixtures consistently reveal a growth at

2001 cm-1 that is affected by CO in a manner consistent with the behavior of THF/CO mixtures (see 3.3.3) and no or little regeneration is observed. In the presence of a coordination capable solute the photochemically created pentacoordinate Fe within [2] accepts a new coordination partner, which is visualized as a growth at 1868 cm-1. This growth does not appear toluene/N2 mixtures.

3.3.13 DCM/N2

A solution of [2] in DCM/N2 was photolyzed in five minute intervals for 15 minutes (Figure 27). The difference spectrum reveals a small but defined growth at 2001

0.8

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B 0.00

A ∆

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-0.10 2100 2050 2000 1950 1900 1850 1800 -1

c n

a 0.4

b A

0.2

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2100 2050 2000 1950 1900 1850 1800 Wavenumbers (cm-1) Figure 26. Photolysis of [2] in toluene/H2 at 365 nm for 30 minutes (5 min per line) Inset: difference spectrum shows total changes during photolysis

0.6

100

50 S

B 0 A

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-50

-3 -100x10 2100 2050 2000 1950 1900 1850 1800 -1 cm

0.4

a b

r 0.3

b A

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0.1

0.0 2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 27. Photolysis of [2] in DCM/N2 at 365 nm for 15 minutes (5 min per line) Inset: difference spectrum shows total changes during photolysis

59 cm-1 following the first five minutes of irradiation, and some bleaching of the parent peaks.

3.3.14 DCM/CO

The spectral changes of [2] in a DCM/CO mixture (Figure 28) mimic those observed in toluene/CO (see Figure 25). When photolyzed for 25 minutes in five minute increments the parent peaks bleach most noticeably in the first 10 minutes. Further bleaching occurs but the spectral lines lie more closely together. The volatility of DCM prevents photolysis beyond 20-25 minutes, despite precautions to keep the IR cell well sealed.

3.3.15 Summary of Spectral Changes in DCM

Photolysis changes in DCM mixtures are consistent with changes observed in toluene (see 3.3.12) showing a single growth at 2001 cm-1 upon photolysis, which is impeded by CO.

3.4 Discussion

3.4.1. Comparative Interpretation of THF and Toluene Solvent Studies

In all experiments described above spectra exhibit specific changes at 2001 cm-1, though the changes in this region vary in shape and intensity. Based on the observed spectral changes, the growth at 2001 cm-1 is assigned as a tricarbonyl species associated with the loss of a CO ligand. In support of the hypothesis, the isobestic conversion between the 1979 cm-1 parent peak and the 2001 cm-1 photo-growth suggest a uniform photoproduct in which only one CO ligand is removed. Formation of a photoproduct consisting of multiple photolabilized CO ligands would likely present more complex

0.7

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b A 0.3

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2100 2050 2000 1950 1900 1850 1800 -1 Wavenumbers (cm ) Figure 28. Photolysis of [2] in DCM/CO at 365 nm for 25 minutes (5 min per line) Inset: difference spectrum shows total changes during photolysis

spectral changes and such were not observed. The 2001 cm-1 peak grows in most completely in the presence of N2 but does not appear to be affected by solvent.

As was expected of a coordinating solvent, release of a CO ligand allows for the binding of a THF ligand. This coordination event was observed as a photo-growth centered at 1868 cm-1 that appears under all conditions. This peak exhibits the most pronounced changes in the presence of dissolved N2. With no competing solutes THF (or another ligand) is free to occupy the photochemically provided open coordination site of the unsaturated Fe. Solvated photoproduct species of [2] have been observed by others in this region as well.20

When photolysis is performed in the presence of excess dissolved CO, slightly different trends occur; a binding competition between the bulk solvent and dissolved CO informs the spectroscopy. The photo-growth at 1868 cm-1 shows some (but not significant) growth, while the peak at 2001 cm-1 grows in initially but exhibits little further change. These results are consistent with the initial photo-induced loss of a CO ligand, allowing for THF binding and competitive binding of dissolved CO to partially reform the parent species. The effect is less efficient formation of the solvent adduct. The toluene/CO spectra corroborate this analysis, showing parallel behavior of the 2001 cm-1 peak in toluene/CO versus toluene/N2 mixtures. These studies are analogous to ultrafast studies in which formation of a solvento species competes with geminate recombination of the initial CO lost.20

Attempts to thermally reform the parent species in solution in the presence of excess CO were unsuccessful in the absence of photolysis. While dissolved CO will recombine with the unsaturated photoproduct during active periods of irradiation, this process does not continue to occur thermally. Furthermore, while changes specific to

62 irradiation can be monitored they are undoubtedly affected by thermal changes, which begin immediately in solution and continue post photolysis. Once [2] is dissolved it will change irreversibly into unknown thermal decomposition products making complete regeneration of starting material impossible. This decomposition process is slow relative to the changes that occur as a result of photolysis, but samples of [2] thermally degrade regardless of photolysis.

Spectra obtained in the presence of dissolved H2 show features similar to the above THF/CO studies (see 3.3.3). In THF the muted growth of the 1868 cm-1 peak suggests that H2 also competes with THF for the same binding site. Toluene/H2 studies reveal analogous growth in this region that otherwise remains flat in the presence of N2.

The IR photospectrum of excess phosphine conditions show the greatest potential for isolation of a thermally stable photoproduct. The unwavering isobestic character seen in Figure 24 clearly indicates that a single photoproduct forms by coordination of a third phosphine ligand. Addition of a third, trimethylphosphine ligand sterically restricts the numerous isomers thought to be in solution,19, 20-21 explaining why formation of this photoproduct is perfectly isobestic, whereas THF coordination appears in the same region but spectral crossings drift slightly.

Isolation of a triphosphine coordinated photoproduct has not yet been successful.

The photogenerated triphospine adduct appears only weakly under the conditions in

3.3.9.3 because the volatile PMe3 ligand (vapor pressure = 49 kPa at 20 °C) escapes into the cuvette leading to less intense spectral changes. A sample was dried on the Shlenk line to eliminate oxygen exposure and only showed evidence of the unsaturated [2]. To reduce the drying time, subsequent samples were dried by rotary evaporation, but with the same result. Given the consistent, strong spectral changes seen in Figure 24 following

63 intense periods of irradiation, it seems unlikely that reduced pressure would remove a coordinated PMe3. Two possible explanations are offered: 1) an insufficient amount of the triphosphine photoproduct is generated initially, which would be addressed by reducing headspace in the cuvette, or 2) in the absence of direct UV exposure the photoproduct thermally releases the third phosphine ligand, which could be resolved with a comparison to post-photolysis thermal changes of a photo-generated triphosphine species.

3.4.2. Analysis of MeCN and DCM Efficacy as Solvent Choices

Samples dissolved in MeCN showed obvious physical decomposition marked by the formation of precipitates. Interestingly, when [2] was photolyzed in a N2 saturated solution the only new feature of the spectrum was formation of the 3-CO species, seen as a growth at 2001 cm-1 (the three parent peaks decreased as well). No additional coordination event was observed, which would have presented as an 1868 cm-1-centered growth. Given the small size, linearity, and ability to act as a good Lewis base, spectral changes resulting from coordination of MeCN were anticipated. One possible explanation is that the extinction coefficient of the peak representative of this binding event is small and obscured by thermal degradation in this solvent. Although, this seems unlikely since time resolved IR (TRIR) by two other groups shows a short lived (200 ps) growth near

1870 cm-1 upon photolysis of [2] in MeCN.19-20 This peak completely decays by 440 ps, while a peak near 2001 cm-1 appears and persists up to 700 ps, which was the last reported spectrum. In addition, others described loss of catalytic activity of [2H]+ in

MeCN due to thermal decomposition and CO/MeCN ligand substitutions.21 In light of these pieces of evidence, MeCN is not a viable solvent choice for mechanistic studies of

[2] and [2H]+.

The inconclusive results of the MeCN/N2 studies compromise the analysis of

MeCN/CO experiments. However, based on the results of THF/CO experiments, it appears that the photochemically provided coordination site at the Fe core responds to the presence of dissolved CO. While the photospectrum in MeCN/N2 shows a peak at 2001 cm-1, the same region is relatively featureless under excess CO conditions and corroborates the rebinding of CO. Here, again, minimal photoactivity and lack of evidence supporting a solvento species suggests that thermal degradation in solution prevents formation of any stable photoproducts.

While solubility of [2] in DCM is excellent, the photochemical behavior is inconsistent. This may be attributed in part to the high volatility of the solvent, which in turn produces false spectral changes (i.e. absorbance increases due to evaporation).

Although DCM may be an unreliable solvent, photo-activity of the 2001 cm-1 growth is consistent with photolabilization of a CO ligand (see 3.4.1) and effect of CO on this feature is analogous to all other solvents.

3.5 IR Conclusion

The wavelength specific photo-induced loss of a CO ligand is observed as a single growth at 2001 cm-1, which is consistent in all solvents. Photolysis in the presence of dissolved CO lessens the growth of this peak, supporting the hypothesis of formation of a

3-CO photoproduct via loss of a single CO ligand. Further changes due to coordinating solvents or solutes cause a growth at 1868 cm-1.

MeCN and DCM are not reliable solvents for mechanistic studies of the neutral compound [2] because it remains unclear if spectral changes are primarily the result of thermal degradation rather than photolysis. However, absorbance decrease of the parent peaks is to be expected in all solvents due to the general instability of [2] in solution.

No photoproduct has been isolated yet due to issues of decomposition in solution, but the tri-substituted phosphine species shows promise and will be investigated in the future.

Chapter 4

Photochemical Behavior of (μ-pdt)[Fe(CO)2(PMe3)]2 ([2]) in the Optical

4.1 Introduction

The clear absorbance peak of [2] makes it an ideal species for studies in the UV and visible range, but few have shown interest in tracking such photo-spectral changes.

Optical spectra provide additional support for analysis of changes tracked by IR. While transitions have been previously assigned in [2] as metal to ligand charge transfer

(MLCT, 348 nm) and metal to metal charge transfer (MMCT, ca. 500 nm)19 the spectrum of [2H]+ does not have a presence in the literature of these complexes because of its less defined shape.

In addition to tracking spectral changes of [2], a dirhodium(II) complex has been cleverly employed as a colorimetric assay to detect CO in solution and therefore confirm the loss of CO from [2] and potentially [2H]+. The photochemically liberated CO of [2] coordinates the probe complex causing a change in color and a spectral shift in support of the hypothesized CO loss from [2] and is described below.

4.2 Methods

To prepare samples for UV-vis studies a standard quartz cuvette (1 cm pathlength) containing a minimal amount of solid [2] was fit with a rubber septum and gently degassed then back filled with N2 (or CO or H2) for three cycles. To the cuvette 2 mL of solvent was added by syringe and final concentrations were near 0.08 mM. For studies requiring dissolved N2, Sure/Seal solvents were used as received. When samples required dissolved CO or H2, the N2 was replaced with the appropriate solute by bubbling the corresponding gas through the solvent at approximately 1 mL per minute. The samples were magnetically stirred while being irradiated by a portable Hg lamp at 365 nm. The set up was similar to the one depicted in Figure 15 for IR experiments. Duration of exposure was adjusted to suit the specific goals of each experiment and 10 second intervals were typical (see Results, 4.3), which is much shorter than studies in the IR.

An alternate method of sample preparation was also explored. Dissolved gasses were removed from liquid samples by three freeze-pump-thaw cycles in a specialized

Schlenk flask containing both a bulb and a quartz cuvette. The desired gas was then allowed to diffuse into the solution while magnetically stirring. The sample was poured from the bulb to the cuvette of the flask and spectra could then be collected without the risk of oxygen contamination or evaporation.

Initially, UV-vis studies were carried out using the Schlenk flask described above because samples could be monitored for lengths of time without evaporating. Samples prepared in this manner were in solution for up to 30 minutes prior to photolysis, leading to increased thermal decomposition as evidenced by precipitate formation and decreased absorbance monitored at the λmax (348 nm). Furthermore, all attempts to monitor regeneration of [2] in the presence of excess CO were unsuccessful, eliminating the need

68 for experiments longer than 30 minutes. In light of the aforementioned reasons the more convenient approach of direct cuvette sample preparation was the preferred method.

For CO detection experiments the above protocol was employed with the addition of [(RhPPh3OAc)2  (H2O)2] (PPh3 = triphenylphosphine; OAc = acetate) which is a binuclear rhodium(II) complex with axial water ligands and was synthesized according to

28 literature methods. In the presence of CO this sensitive Rh(II) dimer exchanges H2O for

CO (shown in Scheme 3) as evidenced by a hypsochromic shift. The Rh(II) dimer was in excess relative to [2].

4.3 Results

4.3.1 Toluene/N2

A sample of [2] was photolyzed in a toluene/N2 mixture in 10 second intervals for

100 seconds total. Photolysis produced consistent decrease of the parent compound at

λmax (348 nm) and a new species appeared as a blue-shifted growth with an isobestic crossing (Figure 29). The small shoulder at 520 nm decreased in absorbance.

4.3.2 Toluene/CO

Spectral changes of [2] were recorded in a CO-saturated toluene solution.

Irradiation was performed in 10 second intervals for 100 seconds. The photospectrum shown in Figure 30 shows similar decreasing trends to Figure 29 but lacks isobestic behavior and photo-growths.

4.3.3 Toluene/phosphine/N2

4.3.3.1 Residual Phosphine

A wet sample of [2] (i.e. the compound was not completely dried and contained

Scheme 3 Reaction of binuclear rhodium(II) complex with CO where L is water.28 Shading corresponds to color of solutions in CHCl3.

Figure 29. UV-vis monitored photolysis of [2] in toluene/N2 at 365 nm for 100 seconds (10 sec per line) showing bleach of λmax at 348 nm and a growth near 300 nm.

Figure 30. UV-vis photolysis of [2] in toluene/CO for 100 seconds (10 sec per line). Spectrum shows bleaching of λmax at 348 nm.

72 residual unbound phosphine ligand) was tested to observe its spectroscopic behavior in a toluene/N2 mixture and is shown in Figure 31. Upon photolysis, evenly spaced bleaches of the parent compound mimicked those of the toluene/N2 spectrum (Figure 29). The trough near 300 nm grew in slightly (pink lines) followed by a jump after 20 seconds of photolysis (blue lines). At 60 seconds a third trend was observed (green lines) when spectral spacing became narrow again, converging at 326 nm. A new isobestic point appeared at 409 nm, which was not seen under any other experimental conditions. The result was a small growth of the MMCT band at 520 nm.

4.3.3.2 Added Phosphine

A freshly synthesized sample of [2], which was completely dry, was dissolved in a solution of toluene/N2 that contained 50 μL excess PMe3 (1M in toluene). The mixture was photolyzed in a cuvette in 10 second intervals for 80 seconds and changes are shown in Figure 32. The spectrum resembles Figure 31, including the growth at 520 nm, but the isobestic crossings are not as clean.

4.3.5 THF/N2

The spectrum in THF/N2 exhibits similar trends as compared to toluene/N2. As the parent peak bleaches the first 50 seconds of photolysis produces only a minimal growth in the trough at 300 nm (lightest blue lines in Figure 33). Over the next 40 seconds of photolysis this valley continues to grow with increased spectral spacing

(bright blue lines) until 100 seconds when spectral lines in the growth region nearly overlap (darkest blue lines). There is no apparent isobestic behavior and the charge transfer band at 520 nm decreases.

2.0

1.5

a b

r 1.0

b A

0.5

0.0 300 350 400 450 500 Wavelength (nm)

Figure 31. UV-vis photolysis of [2] in toluene/N2 in the presence of residual PMe3 for 100 seconds (10 sec per line). As λmax decreases absorbance increases near 300 and 500 nm.

Figure 32. UV-vis photolysis of [2] in toluene/N2 with 50 μL excess PMe3 for 80 seconds (10 sec per line). The λmax decrease is associated with an increase near 300 and 500 nm.

Figure 33. UV-vis photolysis of [2] in THF/N2 for 120 seconds (10 sec per line) showing a decrease at 348 nm and an increase near 300 nm.

Using the same conditions of sample preparation, the experiment was modified such that irradiations lasted 30 seconds and were collected for four minutes (Figure 34).

During the first two minutes of irradiation all lines intersect at 316 nm and a shallow growth appears in the valley at 305 nm. Beyond two minutes isobestic character is lost and a general increase in this region continues.

4.3.6 CO Detection by [(RhPPh3OAc)2  (H2O)2]

A minimal amount of [2] was dissolved in DCM/N2 with excess [(RhPPh3OAc)2 

(H2O)2]. As a solid [(RhPPh3OAc)2  (H2O)2] is a deep purple powder but turns the typically pale yellow [2]/DCM solution green. This solution was photolyzed for 70 seconds, collecting spectra after each 10 second interval (Figure 35). During this time the solution became light orange. The probe absorbs intensely below 400 nm but has a smaller absorption band at 600 nm. Photolysis caused a 9 nm blueshift in this band and a growth at 518 nm appeared. Previous controls were done to ensure that the probe was not photochemically active.

4.4 Discussion

4.4.1 Solute Effects

The decomposition of [2] in solution is expedited by photolysis, which may be related to solvent. This observation has not been quantified but is based on the visual appearance of brown-orange precipitates, which seem to form most readily in DCM and

MeCN (data not shown). The spectra also change inconsistently in these solvents and show little more than a global decrease in absorbance.

2.0

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a b

r 1.0

b A

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0.0 300 350 400 450 500 550 600 Wavelength (nm)

Figure 34. UV-vis photolysis of [2] in THF/N2 for 240 seconds (30 sec per line). Spectrum shows absorbance decrease at λmax and a growth near 300 nm.

Figure 35. Photolysis with a CO probe. A solution of [2] was made with DCM/CO and excess [(RhPPh3OAc)2(H2O)2] and photolyzed for 70 sec (10 sec per line). Spectrum shows hypsochromic shift from 600 to 591 nm and growth of the valley at 518 nm indicating formation of [(RhPPh3OAc)2(H2O)2-nCOn].

The large extinction coefficient of [2] in the UV (ca. 14,000 M-1cm-1)19 requires low concentrations (ca. 70 μM) to stay within the desired absorbance range of 1.0-1.5.

This is on the higher side for UV-vis absorbance but beneficial because sample concentration dictates how long the compound can withstand photolysis. While an IR sample continues to show clean spectral changes for 20 to 30 minutes (see Results, 3.3), a

UV sample is far more susceptible to degradation. Not only must the total time of photolysis be reduced to less than five minutes, but the intervals between each spectral scan must also be brief (10-30 seconds) in order to capture an accurate representation of the photochemical changes. Spectra of overphotolyzed samples show drifting, unusual absorbance decreases, and broadening of regions that initially showed isobestic behavior.

Spectral changes in the UV-vis in toluene/N2 solutions clearly show that the loss of the parent species is associated with conversion to a new photoproduct (as indicated by a single isobestic point at 314 nm), which is thought to be the unsaturated tricarbonyl derivative of [2]. These changes resemble the photochemical behavior of [1] which also releases a CO ligand upon irradiation.16i In contrast, the dicyano-substituted model in

- which the two PMe3 ligands of [2] are replaced with CN ligands (Figure 4E), exhibits very different photochemistry in the optical range showing no isobestic behavior and no photo-induced growths. Unlike [1] and [2] this compound is believed to release a CN- ligand.16b

As shown in Figure 30, photolysis in a toluene/CO mixture results in an absorbance decrease of the MLCT band at 348 nm, but no growth at 300 nm, which is comparable to IR changes under similar conditions (compare Figures 22, 25 with Figures

29, 30). These changes are best explained by the photo-induced loss of CO from [2], subsequent rearrangement of the unsaturated FeL2 subunit, followed by re-coordination

80 of dissolved CO. The re-saturated [2] prevents the photo-growth and any isobestic behavior as the mixture of species in solution becomes increasingly complex. However, an appreciable amount of the parent complex is simultaneously lost leading to the λmax decrease.

The presence of excess phosphine in solution affects spectral changes in a different manner. Initially, the photospectrum reflects CO loss with only the slightest growth in the valley similar to toluene/N2 and THF/N2 (Figures 29, 33, 34). Soon after, a large jump occurs (blue lines in Figure 31) presumably reflecting coordination of PMe3, which is limited by mixing of the solution. This behavior is also observed in Figure 32 but less drastically because of the higher PMe3 concentration. Energetic barriers associated with steric accommodation of the third PMe3 ligand restricts the FeL2 subunit, which is reflected by the nearly isobestic behavior at 326 nm.

Coordination of a third PMe3 ligand is the only observed event that causes the

MMCT band at 520 nm to increase (Figures 31 and 32). Strong σ donation from the three phosphines increases basicity about the Fe-Fe core and increases electronic communication across the Fe-Fe bond. It is unclear if this band decreases under other experimental conditions as a result of diminishing parent complex concentrations or if it reflects electronic isolation of the respective Fe(I) subunits.

THF coordination is affected by the extent of photolysis intervals. During 10 second irradiations a solvento species that is [(pdt)Fe2CO3(PMe3)2THF] forms gradually

(Figure 33). However, 30 second irradiation intervals produce a photospectrum very similar to Figure 29 in which no solvento species forms but spectral lines intersect isobestically. Thus, it is possible that the intensity of lengthier photolysis induces some longer-lived excited state that must relax before the complex can accept a solvent ligand.

Quantum yields, which relate changes in molar absorptivity, irradiation time, and lamp power, may be useful in determining these effects related to duration of exposure.

4.4.2 Qualitative CO Detection

Spectral changes associated with the highly sensitive (CO specific) binuclear

Rh(II) probe offer support of CO photo-dissociation from [2]. Chromogenic changes are discernible by the naked eye at the lower limit of 35 ppm CO but chemosensitivity goes as low as 1 ppm.22 The probe does not show light induced spectral changes in the absence of CO, therefore, spectral changes in Figure 35 are cause by the exchange of an axial H2O ligand on the probe with CO released from [2] (which is the only source of CO in solution).

The sensitivity of this detection system potentially lends itself to more quantitative assignment of CO release associated with the color changes. For example,

CHCl3 solutions of [(RhPPh3OAc)2  (H2O)2] change to orange then yellow upon coordination of one and two CO ligands, respectively. While [2] and [2H]+ are believed to lose a single CO ligand, these tests may further verify this hypothesis, as well as define parameters of photolysis such that specifically one or two CO ligands could be differentially labilized from the parent.

4.5 Conclusion

Monitoring optical spectral changes remains a challenge for both [2] and [2H]+.

The inconsistent behavior of [2] appears to be due in part to concentration and duration of photolysis. It is hoped that a semi-stable photoproduct that is a trisubstituted phosphine complex may be generated and immediately characterized by MS as validation of CO loss. The Rh(II) dimer detects CO when in solution with [2] but thermal (dark) release of

CO should be considered as well. Although [2H]+ was expected to behave similarly to [2]

82 in the UV, because of the broad sloping shape of the optical spectrum of [2H]+ (Figure

11) and lack of spectral crossings, photospectral changes have not been thoroughly explored. However, the absorbance of the spectrum of [2H]+ is not monitored in the CO detection assay meaning the probe can be used for this compound as well.

Chapter 5

Thermal Behavior of Photochemically altered {(μ-H)(μ-

+ pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] ) in the Infrared

5.1 Introduction

Upon protonation of [2], the catalytically relevant model (μ-H)(μ-S(CH2)3S)

+ + [Fe(CO)2(PMe3)]2[PF6] ([2H] ) can competently reduce H to H2. While electrochemical methods produce the greatest quantity of H2, synthetic models also produce H2 photochemically (see 1.3). The [2Fe]H models are synthesized as the inhibited form of the enzyme (Hox-CO), which necessitates an open coordination site prior to photochemical production of H2. In addition to hydrogen evolution, the inherent stability provided by the bridging hydride makes thermal regeneration of the [2H]+ complex possible in the presence of CO, a feature important to understanding the behavior of this model system.

This property has been established in the hexacarbonyl scaffold [1]16i and the enzyme.14

This chapter describes the photochemical behavior of [2H]+ in two solvents and highlights the limited thermal reversibility of CO-loss in the presence of dissolved CO.

5.2 Methods

The sample preparation for IR studies of [2] described in Methods, 3.2 was applied to [2H]+ without further modification. Times and intervals of photolysis are specific to each experiment and described individually in Results (see 5.3). Because the cationic nature of [2H]+ prohibits use of nonpolar organic solvents such as toluene, DCM was used as a non-coordinating model and acetone was used to model coordinating solvents.

5.3 Results

5.3.1 DCM/N2

+ A solution of [2H] (ca. 5 mg) in DCM/N2 was photolyzed for 25 min in five minute intervals to examine the overall effect of photolysis (Figure 36). The characteristic parent peaks in this solvent (2032, 1990 cm-1) decreased initially though changed little beyond 10-15 minutes of photolysis. In addition, two lower energy νCO bands appeared at roughly 1959 and 1932 cm-1. These growths were small relative to the intense absorbance from the parent complex and showed no further change after 10 minutes of photolysis. The trough between the two parent peaks grew subtly with light exposure.

5.3.2 DCM/H2

+ Changes seen in the IR when [2H] was dissolved in DCM/H2 (H2 concentration is estimated to be ca. 3 mM, see 3.2) are shown in Figure 37 and look nearly identical to the DCM/N2 solution described above (see 5.3.1). The only difference being a minor spacing of spectral lines with respect to the growth region (1960-1930 cm-1) but the final spectrum exhibits the same low absorbing peaks at 1959 and 1932 cm-1.

0.2 1.4 0.1

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2200 2100 2000 1900 1800 -1 cm 1.0

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+ Figure 36. Photolysis of [2H] in DCM/N2 for 25 minutes (5 min per line) Inset: difference spectrum reflects total changes during photolysis

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0.0

2200 2100 2000 1900 1800 Wavenumbers (cm-1) + Figure 37. Photolysis of [2H] in DCM/H2 for 25 minutes (5 min per line) Inset: difference spectrum reflects total changes during photolysis

5.3.3 DCM/CO

Spectral changes during photolysis of [2H]+ were collected in the presence of dissolved CO (Figure 38). Using CO purged DCM (containing ca. 7 mM CO), [2H]+ was photolyzed in 5 minute intervals. Over 20 minutes of photolysis the parent peaks maintain a decreasing trend with even spacing between spectra. Also, as the parent complex changes, the growth region (1960-1930 cm-1) shows broad, less defined absorbance increases centered at 1942 cm-1. This peak is adjacent to a shoulder (ca. 1969 cm-1) that transitions from the parent peak isobestically and is clearly visualized in the difference spectrum.

In addition to characterizing total photolysis changes, samples of [2H]+ were photolyzed in a DCM/CO mixture and thermal behavior was monitored in the dark.

Under these conditions the solution was initially irradiated for 15 minutes continuously to ensure sufficient photoproduct formation. These changes are represented by the black trace in the Figure 39 inset difference spectrum. By 20 minutes the parent peaks at 2030 and 1990 cm-1 showed full recovery (Figure 39, red trace). However, the photo-growths at 1960 and 1932 cm-1, which were expected to decrease, hardly changed, and the trough between the two parent peaks (a region of subtle growth during photolysis) remained unchained during thermal monitoring.

5.3.4 Acetone/N2

+ In a solution of acetone and N2, [2H] was photolyzed for a total of 30 minutes in five minute intervals. IR spectra were collected after each interval (Figure 40) and total changes over the course of the experiment are shown in the difference spectrum (Figure

40 inset). After the first five minutes of exposure the parent peaks at 2030 and 1989 cm-1

2.0 0.2

0.1

1.8 0.0

S B

A -0.1 ∆

-0.2 1.6 -0.3

-0.4 2200 2100 2000 1900 1800 1.4 -1 cm

1.2

n a

b 1.0

b A

0.8

0.6

0.4

0.2

0.0

2200 2150 2100 2050 2000 1950 1900 1850 1800 Wavenumbers (cm-1) Figure 38. Photolysis of [2H]+ in DCM/CO for 20 minutes (5 min per line) Inset: difference spectrum reflects total changes during photolysis

1.4 0.2

0.1 S

B Back reaction with CO

A ∆ 0.0 Photolysis changes 1.2 -0.1

-0.2 2200 2100 2000 1900 1800 -1 cm 1.0

e 0.8

b A 0.6

0.4

0.2

0.0

2200 2150 2100 2050 2000 1950 1900 Wavenumbers (cm-1) Figure 39. Back reaction of [2H]+ in DCM/CO. Depicted is 15 min continuous photolysis (blue) and regeneration of saturated species at 20 min (red) with the initial spectrum in black. Inset: Difference spectra shows photolysis (black) and thermal regeneration (tan)

2.0

0.2 1.8

0.0

B A

∆ -0.2 1.6 -0.4

-0.6 2200 2100 2000 1900 1800 1.4 -1 cm

1.2

n a

b 1.0

b A

0.8

0.6

0.4

0.2

0.0

2200 2100 2000 1900 1800 Wavenumbers (cm-1) + Figure 40. Photolysis of [2H] in acetone/N2 for 30 minutes (5 min per line) Inset: Difference spectrum reflects total changes during photolysis

91 bleach by almost 50% while a photogrowth appears at 1942 cm-1. This lower frequency growth is fairly sharp and the only significant feature of the photospectrum, although minor changes are present as well. The entire spectrum appears to redshift slightly in response to photolysis, and this shift is fully established within the first five minutes of exposure. The CO stretches show no difference between the first photolysis interval after five minutes and the final exposure after 30 minutes.

5.3.5 Acetone/CO

As described in 5.3.3, photolysis in the presence of dissolved CO was performed in two ways for samples of [2H]+ in acetone solutions. Figure 41 displays transient photolysis changes when 4 mg of [2H]+ was dissolved in CO purged acetone and photolyzed in consecutive five minute intervals for 30 minutes total. Similar to

-1 acetone/N2 mixtures, the parent peaks centered at 2030 and 1989 cm bleach intensely, but unlike N2, these two peaks continue to decrease in intensity in the presence of dissolved CO. The red-shifting parent peaks give way to a sharp growth at 1942 cm-1. As this growth develops, the first two photospectral lines (representing 10 minutes exposure) intersect the waning parent peak through the same point. By 15 minutes a shoulder at

1972 cm-1 develops but decreases without further intersection.

The objective of a second experiment was to generate a photoproduct and observe thermal regeneration of the parent species. This was achieved by photolyzing a fresh sample of [2H]+ for 20 minutes (Figure 42, blue trace) then thermal changes were monitored in the dark. Spectra were collected at 10 minutes (Figure 42, fuchsia trace) and

20 minutes (Figure 42, purple trace). The difference spectrum in the Figure 42 inset clearly shows the same 1942 cm-1 growth and shoulder described above (black trace).

0.2

1.4 0.1

0.0

S -0.1

B A

∆ -0.2 1.2 -0.3

-0.4

-0.5 2200 2100 2000 1900 1800 -1

1.0 cm e

c 0.8

b A

0.6

0.4

0.2

0.0 2150 2100 2050 2000 1950 1900 1850 Wavenumbers (cm-1) Figure 41. Photolysis of [2H]+ in acetone/CO for 25 minutes (5 min per line)

Inset: Difference spectrum reflects total changes during photolysis

1.2

0.2

Back reaction with CO 0.0

S 20 min photolysis

1.0 B

A ∆ -0.2

-0.4

0.8 2200 2100 2000 1900 1800 -1

a b

r 0.6

b A

0.4

0.2

0.0 2100 2050 2000 1950 1900 Wavenumbers (cm-1) + + Figure 42. Back reaction of [2H] in acetone/CO. A solution of [2H] in acetone/CO (black) was photolyzed 20 min (blue) and thermal changes were monitored in the dark at 10 min (fuchsia) and 20 min (purple). Inset: difference spectra representing total changes for photolysis (black) and thermal reaction with CO (tan).

Overlaid in tan is the difference spectrum representing post-photolysis thermal changes monitored in the dark. The reformed parent compound shows ca. 40 % peak recovery of both starting peaks. As the 1942 cm-1 photogrowth bleaches it begins to form two small, broad peaks similar to those seen in DCM solutions (Figures 35 and 38). Under these conditions thermal reversibility is incomplete.

5.4 Discussion

5.4.1 Photochemical Changes of [2H]+ in DCM Solutions

When [2H]+ is photolyzed in DCM solutions two low absorbing, broad peaks appear at lower frequencies and the bottom portions of the parent peaks redshift slightly under all conditions. The exigent hydride bridge geometrically constrains [2H]+ but photochemical CO liberation alleviates this barrier in the unsaturated Fe, thus explaining the red-shifting character of the spectra.

DCM/N2 photolysis models a fully non-coordinating environment such that spectral changes represent formation of a pentacarbonyl-[2H]+ and is not complicated by coordination of another solute. Thus, this reference experiment is used to interpret

+ changes observed in DCM/H2 and DCM/CO. While [2H] converts to a preferred photoproduct in DCM/N2 almost immediately, H2 and CO compete with this process.

Spectral changes in the presence of H2 that deviate from the N2 reference spectrum are

2 likely due to η -H2 coordination (side on binding through the H2 σ bond) for which there is precedence in metal hydrides.21, 27, 29 The photospectra might exhibit more pronounced changes were H2 in greater concentration.

Interval photolysis of [2H]+ in a mixture of DCM/CO revealed interesting spectral changes in the form of two growths at 1969 and 1942 cm-1 that are blue-shifted 10 cm-1 from the peaks seen in the N2 reference photospectrum (see Figure 38). The band seen at

1969 cm-1 is a shoulder that shows isobestic conversion as it grows and it gives way to the 1942 cm-1 band. Following CO release, the [2H]+ photoproduct undergoes photo- induced isomerization before accepting another CO ligand.20 The new orientation of the re-saturated [2H]+ presents as a higher frequency growth at 1942 cm-1 (relative to 1932

-1 cm in DCM/N2), a band that is also seen when acetone coordinates (vide infra).

5.4.2 Thermal Reversibility of CO Release in DCM

Partial recovery of the parent compound [2H]+ is possible in the presence of CO and shown in Figure 39. This process occurs in conjunction with photolysis (see 5.4.1) and also in the dark following exposure to 365 nm light. While [2H]+ can exchange CO ligands, the primary isomer in solution (of the original parent molecule) is altered, but the implications of this with respect to catalysis (if any) are not yet understood. While this data is largely qualitative, more intensive quantitative work remains to be done. While the baselines of the spectra align at zero, DCM is prone to evaporation, particularly when photolysis exceeds 20 minutes. To eliminate the possibility of confusing CO resaturation of [2H]+ with concentration increases due to evaporation a more rigorous experimental technique should be attempted to verify observed changes. Some of which were attempted and discussed in the next section (see 5.4.3).

Kinetic experiments monitoring the back reaction of CO with photolyzed [1] inspired similar efforts with [2H]+.16i The unfortunate UV-vis spectrum of [2H]+ necessitated kinetics by IR, but concentrations of compound and CO were essentially equal, eliminating use of pseudo first order conditions. No conclusive data was obtained but such information would be of value in examining catalytic reactivity and regenerative properties.

5.4.3 Photochemical Changes of [2H]+ in Acetone Solutions

Unlike other solvent/solute mixtures, five minutes of irradiation sufficiently generates a photoproduct with a well-defined spectrum that mirrors those of longer photolysis runs (Figure 40). Photolysis of [2H]+ in acetone is a favorable system for mechanistic studies because of the dynamic interaction in which accepting and releasing the acetone ligand is a seemingly facile process. During interval photolysis experiments the growth centered at 1942 cm-1 indicates acetone coordination to the unsaturated Fe. In contrast to changes observed in DCM, dissolved CO does not affect the relative shape of this peak, as it still forms after the first exposure and remains sharp. But as photolysis progresses this peak continues to increase in absorbance and the parent peaks shift as they decrease, indicating ongoing photoproduct formation (see Figure 41). As the parent peaks bleach they, too, display pronounced spacing as more photoproduct is formed. This delayed photoproduct formation is consistent with CO and acetone competition for the

+ open coordination site. Also, [2H] lacks the pseudo rotational freedom of the FeL3 until a CO ligand is photolabilized. This CO release allows for the unsaturated FeL2 unit to reconfigure such that by the time another CO ligand coordinates the molecule is of a new orientation, which is why the parent peaks bleach more drastically when CO, rather than

N2, is present.

5.4.4 Thermal Reversibility of CO Release in Acetone

+ The robustness of [2H] makes thermal regeneration of the Hox-CO model state possible in acetone. CO readily replaces an acetone ligand during post-photolysis dark periods and the spectra reveal definite recovery of starting peaks at 2030 and 1989 cm-1.

The valley between the two parent peaks around 2005 cm-1 deepens with light exposure, and while it increases thermally, it remains offset by about 5 cm-1, suggesting that the recovered [2H]+ adopts a different isomeric form than the original complex. Additionally,

97 complete regeneration of the starting material is not to be expected as some portion of the compound likely degrades during periods of intense UV exposure and CO is not in excess.

Finally, maximum regeneration was not achievable because of experimental limitation. As discussed previously (see 5.4.2), samples irradiated directly in the IR cell that remain in the cell during thermal monitoring eventually evaporate leading to false absorbance increases. While this does not appear to be the case with acetone, it may have affected the DCM/CO regeneration experiment, which was much longer (ca. 40 minutes) than any continuous exposure experiment in that solvent (15-20 minutes). Strategies to address these issues of have been unsuccessful. For example, a stock solution (3 mL) could be photolyzed and aliquots sampled. However, standard round bottoms are not of optical quality and cannot be reliably photolyzed at a specific wavelength. A second approach is to dissolve solid [2H]+, freeze-pump-thaw, and back fill with CO. This can be done in specialized glassware that has a bulb, a cuvette and a syringe port. While this solves the problem of how to photolyze, it would require high pressures of CO (which was not attempted) to maintain a sufficient concentration of dissolved CO.

5.5 Conclusion

Spectral changes in the IR following UV photolysis of [2H]+ were studied under different solvent and solute combinations. Experiments performed in DCM solutions should be revisited and revised. To understand the full scope of the regenerative properties of [2H]+ in the presence of CO, experiments must be designed to account for the volatility of organic solvents and solutes. The solvent dependent photo-reactivity of

[2H]+ described in this chapter is a qualitative view of the processes that occur in solution. The quantitative aspects, such as quantum yields and kinetics, would greatly

98 benefit this research, as they would further validate proposed mechanisms and aid in matters of optimal irradiation wavelength and solvent choice, which may be relevant to catalysis.

Chapter 6

A Brief Look at the Photolysis of (pdt)[Fe(CO)2(PMe3)]2 ([2]) and {(μ-

+ H)(μ-pdt)[Fe(CO)2(PMe3)]2[PF6]} ([2H] ) by NMR

6.1 Introduction

Nuclear Magnetic Resonance (NMR) is a powerful technique, useful for characterizing and distinguishing molecules that differ by only one atom such as [2] and

[2H]+. The theory of multiple solution phase isomers of [2] has been well received by the scientific community, and with various possible conformations of [2], one would expect

NMR spectra to reflect such unique environments, particularly with the 31P probe (the 31P nucleus has 100% natural abundance). And while VT NMR reveals distinct isomers of

[2] below -60 °C, the 31P spectrum remains relatively unperturbed at ambient temperatures, even upon protonation of the diiron core. Instead, the observed spectrum represents an average of all conformers in solution.21

In addition to this averaging phenomenon, difficulties arising from the diiron core of [2] and [2H]+ affect resolution, and the already broadened peaks worsen with

100 photolysis. This issue can be circumvented by combining information from the 1H and 2H spectra (as was done by Zhao, et al.), but a 2H probe was not available for the present study and is not as diagnostic when considering [2]. Instead, the behavior of [2] can be monitored by 31P in the presence of excess dissolved phosphine. While the information presented in this chapter does not sufficiently stand alone, it offers guidance for prospective experiments that will hopefully be addressed in the future.

6.2 Methods

To monitor photochemical changes in [2] by 31P NMR, a standard NMR tube was fit with a septum and alternately degassed and backfilled (gently, three times) with the gas corresponding to the desired solute (N2 or CO) in a manner similar to the IR procedures above (see Methods 3.2). Gas concentrations (both in the head space and dissolved) were not used in any calculations. Approximately 15 mg of [2] was dissolved in 0.5 mL of the appropriate deuterated solvent (previously purged) and transferred to the

NMR tube via syringe. The sample was irradiated directly in the NMR tube at 365 nm using a portable Hg light. The sample was inverted periodically to mix. As an aside, use of deuterated solvents is not strictly necessary to collect 31P spectra because peaks are relative to an external 85 % phosphoric acid standard. However, for convenience deuterated solvents were used allowing for collection of 1H spectra with the same sample.

6.3 Results

6.3.1 Photolysis of [2] With Excess PMe3

A sample of [2] (ca. 15 mg) was photolyzed in d6-benzene/N2 in the presence of

20 μL trimethylphosphine (1M in toluene) which was ca. 6-fold excess. Figure 43 shows the progression of this photolysis over 80 minutes. The final spectrum in the figure (top line) was taken 24 hours after the last irradiation and was kept in the dark. Initially the

101

31 Figure 43. Photolysis of [2] in d6-benzene/N2 monitored by P NMR. Solution contains 20 μL PMe3 (1M in toluene). Initial spectrum (bottom) shows [2] at 22.98 ppm and free phosphine at 31.39 ppm, which intensifies with light exposure.

102 unphotolyzed sample exhibited a prominent peak at 22.98 ppm characteristic of [2], and a smaller peak at 31.39 ppm. By 15 minutes this peak showed significant growth, whereas the peak corresponding to [2] appeared unchanged. At 60 minutes the peak at 31.39 ppm overshadowed [2] and new minor peaks appeared. After 80 minutes of photolysis and 24 hours resting in the dark, the primary 31P resonance belonged to the unshifted peak at

31.39 ppm, but a substantial amount of [2] remained at 22.98 ppm. A third peak at 9.47 ppm was seen as well.

6.3.2 Thermal Decomposition of [2] With Excess PMe3

Following 6.3.1, [2] was dissolved in d6-benzene/N2 in the presence of 100 μL

PMe3 in a 1M in toluene solution but the NMR tube was wrapped in foil and not photolyzed. The arbitrarily decided ratio of phosphine to [2] was roughly 30:1. Initially the same peaks at 31.55 ppm and 22.98 ppm could be seen in Figure 44. After 4 hours the smaller resonance at 31.55 ppm diminished noticeably and the peak that represents [2] was seemingly unchanged. The spectrum collected after four days had a primary resonance at 36.15 ppm and three smaller peaks at 53.59, 27.29, and 19.17 ppm. Peaks at

31.55 and 22.98 ppm completely disappeared.

+ 6.3.3 Photolysis of [2H] with N2

To match the timescale of IR experiments, [2H]+ was irradiated in five minute intervals for 15 minutes (Figure 45). The solution contained N2 degassed CDCl3. The only feature of the 31P spectrum was a single resonance at 20.82 ppm.

6.3.4 Photolysis of [2H]+ with CO

+ In CO-degassed d6-acetone, [2H] was photolyzed in five minute intervals for 15 minutes (Figure 46). The first and last spectra were identical and showed a single

103

Figure 44. Thermal degradation of [2] in d6-benzene/N2 was followed by 31P NMR over four days. Mixture contains 100 μL PMe3 (1M in toluene). Initial spectrum (bottom) shows [2] at 22.98 ppm and a tri-substituted phosphine complex at 31.55 ppm.

104

+ Figure 45. Photolysis of [2H] in CDCl3/N2. A single peak is seen by 31P NMR at 20.82 ppm. Total time of irradiation was 15 min.

105

31 + Figure 46. P NMR of [2H] in d6-acetone/CO mixture. A sample + of [2H] was photolyzed in a d6-acetone/CO mixture over 15 minutes. All spectra of the 31P NMR show a single peak at 21.07 ppm.

106 phosphorous peak at 21.07 ppm.

6.4 Discussion

6.4.1 Observed Behavior of [2] by 31P NMR

Initially seen is Figure 43 is the prominent resonance at 22.98 ppm belonging to the parent complex [2]. Additionally, the presence of free PMe3 is seen at 31.39 ppm. As photolysis progresses the free phosphine peak increases and the parent peak decreases.

While this may appear as photo-induced phosphine loss, the thermal control (Figure 44) suggests otherwise (see below). Some amount of [2] appears to withstand the extended photolysis and thermal period, which may be a combination of high compound concentration and poor mixing during irradiation. Eventually a peak at 9.47 ppm appears that is likely some fragment of the original compound.

The initial thermal spectrum (bottom line in Figure 44) shows analogous behavior to Figure 43 owing to the immense excess of free PMe3 in solution. Within 4 hours however, this species reduces significantly and is completely gone by day four. The high volatility and incomplete seal of the NMR tube allow for evaporation of PMe3. After four days in the dark there remains no trace of starting material but the several smaller resonances may be remaining fragments of [2] that have not precipitated out of solution; the large peak at 36.15 ppm is likely some phosphine oxide (which would appear in this region) as the NMR tube was only sealed with a septum and parafilm. This peak would be analogous to the impurity at 38 ppm in CDCl3 (see 2.4.3).

The photolysis-expedited thermal degradation of [2] increases the concentration of free phosphine in solution and continues post-photolysis as seen in Figure 43. The seal of the NMR tube holds well enough over the 24 hour monitoring period explaining the lasting presence of the peak at 31.39 ppm under these conditions.

107

6.4.2 Observed Behavior of [2H]+ by 31P NMR

The limited investigation of the photochemical behavior of [2H]+ in the 31P NMR shows no changes after irradiation for 15 minutes (Figures 45 and 46). Given the responsiveness of the compound in the IR (see 5.3) this might be attributed to relative concentration effects whereby 15 minutes is insufficient to populate an appreciable amount of the CO-unsaturated parent. Photolysis of [2H]+ was not attempted in the presence of excess PMe3 because others reported hydride deprotonation that resulted in the neutral [2]21 but the upfield region of the 1H spectrum (see Figure 14 inset) did not change upon photolysis under experimental conditions (data not shown) indicating that photolytic cleavage of the bridging hydride did not occur.

6.5 Conclusion

Although NMR photolysis of [2] and [2H]+ presents challenges inherent to transition metal complexes that is further compounded by issues of thermal instability,

31P NMR has the potential to advance long-term studies of the di- and tri-phosphine complexes. The experiments in this section are incomplete and must be repeated and refined, but in combination with other probes, specifically 13C, unique environments and photoproducts of the primary species in solution will be better understood. Comparing photochemical changes under various conditions to thermal processes will give further insight into the specific function of the irradiation wavelength.

108

Chapter 7

Conclusion

7.1 Conclusion

Hydrogenase enzymes represent a potential solution to the global demand for clean and renewable energy. The inevitable depletion of fossil fuels necessitates a solution that is sustainable, cost-effective, and scalable. The ability of hydrogenases to both produce and consume H2 is an attractive quality because it not only offers a route to

H2 fuel production, but it also serves as a platform for H2 oxidation in model systems which is necessary to make use of hydrogen fuel. Significant progress has been made in the last three decades toward advancing our understanding of hydrogenases, and new studies continue to strengthen this field of study.

The applications of hydrogenase-based energy are obvious but the development of such systems remains far off because of the inherent limitations of the enzymes’ extreme

O2 sensitivity and the relative inefficiency of the synthetic catalysts. Nevertheless, a better understanding of the natural enzyme has aided in the development and design of organometallic active site models, which aim to address these limitations.

109

The numerous and varied synthetic models of the [FeFe]-hydrogenase active site share a common motif that is a diiron core with two or more CO ligands. In general, these compounds are photoactive and irradiation induces loss of a CO ligand. Of particular interest are the phosphine substituted models, such as [2] and [2H]+, which were chosen for their anticipated similarity to the cyanide containing active site of the native enzymes.

However, the synthetic diphosphine derivatives behave differently than the dicyano derivatives. The electronic donation from the phosphine ligands has a stabilizing effect on the FeFe core that allows [2H]+ to act as a catalyst in the absence of a photosensitizer.

Photochemical behavior of the non-H2-producing [2] has been overshadowed by the catalytic abilities of [2H]+. Furthermore, most studies of [2] examine events immediately following photolysis and do not consider the longer thermal effects associated with CO ligand loss. Because photolysis is required to generate the structural mimic of the enzymatic Hox state, it is important to know how that photoproduct interacts with solutes and solvent molecules which may provide information about the lasting behavior and optimal conditions relevant to other active site model complexes.

The work described in this study shows that, following the photolysis of [2], the coordination site vacated by CO is filled by a coordinating solvent molecule or other small molecule in solution, or it may remain vacant. This species, as well as the parent complex, are not robust, and in solution within a few hours, even in the absence of photolysis. The sensitivity of [2] prevents long-term photolysis with 365 nm light at low concentrations, but less intense light may allow for longer monitoring. This would be applicable to UV-vis studies, which necessitate low concentrations, but may also be useful in the IR. Although regeneration of [2] via a back reaction with CO was not possible under the above experimental conditions (or perhaps obscured), the spectra

110 indicate that CO competes with coordinating solvents during photolysis. Thus, a lower frequency wavelength may induce the desired CO release without over-photolyzing the complex such that resaturation by CO may be possible in [2].

Although [2] and [2H]+ appear very similar, the stability provided by the bridging hydride in [2H]+ and the higher oxidation state of the Fe atoms leads to surprisingly different properties. Like [2], photolysis of [2H]+ also induces CO loss, but the thermal products can combine with dissolved CO to reform the parent compound after extended periods in the dark. Although the thermal behavior of [2H]+ (post-photolysis) could not be monitored beyond about 30 minutes, exposure to white light allows for monitoring over the course of several days (a study by Zhao and coworkers17), indicating that the photochemistry is highly dependent on the irradiation wavelength. Photolysis “activates” the model complex and depending on the process of interest the irradiation wavelength can be selected. That is to say, lower intensity light brings about slower changes but allows for longer studies, while higher intensity light reveals more immediate processes but also degrades the compound faster.

While the specific compounds examined in this study are unlikely candidates for the replacement of Pt catalysts, their investigation informs the entire field of hydrogenase model studies, and directs the future of this research by offering information relevant to optimizing conditions for solution phase photochemical activity of these models.

111

References

1. Tye, J. W.; Hall, M. B.; Darensbourg, M. Y., Better than platinum? Fuel cells energized by enzymes. Proceedings of the National Academy of Sciences of the United States of America 2005, 102 (47), 16911-16912.

2. (a) Frey, M., Hydrogenases: hydrogen-activating enzymes. Chembiochem 2002, 3 (2), 8; (b) Bullock, R. M., Catalysis without precious metals. John Wiley & Sons: 2011; (c) Nicolet, Y.; Piras, C.; Legrand, P.; Hatchikian, C. E.; Fontecilla-Camps, J. C., Desulfovibrio desulfuricans iron hydrogenase: the structure shows unusual coordination to an active site Fe binuclear center. Structure 1999, 7 (1), 13-23.

3. Thauer, R. K.; Klein, A. R.; Hartmann, G. C., Reactions with molecular hydrogen in microorganisms: evidence for a purely organic hydrogenation catalyst. Chemical reviews 1996, 96 (7), 3031-3042.

4. Belaich, J.-P.; Bruschi, M.; Garcia, J.-L., Microbiology and biochemistry of strict anaerobes involved in interspecies hydrogen transfer. Springer Science & Business Media: 2012; Vol. 54.

5. Peters, J. W.; Schut, G. J.; Boyd, E. S.; Mulder, D. W.; Shepard, E. M.; Broderick, J. B.; King, P. W.; Adams, M. W., [FeFe]-and [NiFe]-hydrogenase diversity, mechanism, and maturation. Biochimica et Biophysica Acta (BBA)-Molecular Cell Research 2015, 1853 (6), 1350-1369.

6. Vignais, P. M.; Billoud, B., Occurrence, classification, and biological function of hydrogenases: an overview. Chemical reviews 2007, 107 (10), 4206-4272.

7. Weiss, D. S.; Thauer, R. K., Methanogenesis and the unity of biochemistry. Cell 1993, 72 (6), 819-822.

8. (a) Vignais, P. M.; Billoud, B.; Meyer, J., Classification and phylogeny of hydrogenases. FEMS microbiology reviews 2001, 25 (4), 455-501; (b) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y., Structure/function relationships of [NiFe]- and [FeFe]-hydrogenases. Chemical reviews 2007, 107 (10), 4273-4303.

9. Happe, T.; Kaminski, A., Differential regulation of the Fe-hydrogenase during anaerobic adaptation in the green alga Chlamydomonas reinhardtii. European Journal of Biochemistry 2002, 269 (3), 1022-1032.

10. Warburg, O. H., Schwermetalle als Wirkungsgruppen von Fermenten. 1946.

112

11. Peters, J. W.; Lanzilotta, W. N.; Lemon, B. J.; Seefeldt, L. C., X-ray Crystal Structure of the Fe-Only Hydrogenase (CpI) from Clostridium pasteurianum to 1.8 Angstrom Resolution. Science 1998, 282 (5395), 1853-1858.

12. (a) Lawrence, J. D.; Li, H.; Rauchfuss, T. B.; Bernard, M.; Rohmer, M. Ä., Diiron Azadithiolates as Models for the Iron-Only Hydrogenase Active Site: Synthesis, Structure, and Stereoelectronics. Angewandte Chemie 2001, 113 (9), 1818-1821; (b) Fan, H.-J.; Hall, M. B., A capable bridging ligand for Fe-only hydrogenase: density functional calculations of a low-energy route for heterolytic cleavage and formation of dihydrogen. Journal of the American Chemical Society 2001, 123 (16), 3828-3829; (c) Silakov, A.; Wenk, B.; Reijerse, E.; Lubitz, W., 14N HYSCORE investigation of the H-cluster of [FeFe] hydrogenase: evidence for a nitrogen in the dithiol bridge. Physical Chemistry Chemical Physics 2009, 11 (31), 6592-6599; (d) Berggren, G.; Adamska, A.; Lambertz, C.; Simmons, T.; Esselborn, J.; Atta, M.; Gambarelli, S.; Mouesca, J.-M.; Reijerse, E.; Lubitz, W., Biomimetic assembly and activation of [FeFe]-hydrogenases. Nature 2013, 499 (7456), 66-69.

13. Mulder, D. W.; Guo, Y.; Ratzloff, M. W.; King, P. W., Identification of a Catalytic Iron-Hydride at the H-cluster of [FeFe]-hydrogenase. Journal of the American Chemical Society 2016.

14. Chen, Z.; Lemon, B. J.; Huang, S.; Swartz, D. J.; Peters, J. W.; Bagley, K. A., Infrared studies of the CO-inhibited form of the Fe-only hydrogenase from Clostridium pasteurianum I: examination of its light sensitivity at cryogenic temperatures. Biochemistry 2002, 41 (6), 2036-2043.

15. Stripp, S. T.; Happe, T., How algae produce hydrogen- news from the photosynthetic hydrogenase. Dalton Transactions 2009, (45), 9960-9969.

16. (a) Georgakaki, I. P.; Miller, M. L.; Darensbourg, M. Y., Requirements for Functional Models of the Iron Hydrogenase Active Site: D2/H2O Exchange Activity in {(μ-SMe)(μ-pdt)[Fe (CO) 2 (PMe3)] 2+}[BF4-]. Inorganic chemistry 2003, 42 (8), 2489- 2494; (b) Hunt, A.; Barrett, J.; McCurry, M.; Works, C., Photochemical reactivity of a binuclear Fe (I)‚ÄìFe (I) hydrogenase model compound with cyano ligands. Polyhedron 2016, 114, 306-312; (c) Bingaman, J. L.; Kohnhorst, C. L.; Van Meter, G. A.; McElroy, B. A.; Rakowski, E. A.; Caplins, B. W.; Gutowski, T. A.; Stromberg, C. J.; Webster, C. E.; Heilweil, E. J., Time-resolved vibrational spectroscopy of [FeFe]-hydrogenase model compounds. The Journal of Physical Chemistry A 2012, 116 (27), 7261-7271; (d) Gloaguen, F.; Lawrence, J. D.; Rauchfuss, T. B.; Benard, M.; Rohmer, M.-M., Bimetallic carbonyl thiolates as functional models for Fe-only hydrogenases. Inorganic chemistry 2002, 41 (25), 6573-6582; (e) Kaziannis, S.; Santabarbara, S.; Wright, J. A.; Greetham, G. M.; Towrie, M.; Parker, A. W.; Pickett, C. J.; Hunt, N. T., Femtosecond to microsecond photochemistry of a [FeFe] hydrogenase enzyme model compound. The Journal of Physical Chemistry B 2010, 114 (46), 15370-15379; (f) Tard, C.; Pickett, C. J., Structural and functional analogues of the active sites of the [Fe]-,[NiFe]-, and [FeFe]-

113 hydrogenases. Chemical reviews 2009, 109 (6), 2245-2274; (g) Wang, N.; Wang, M.; Wang, Y.; Zheng, D.; Han, H.; Ahlquist, M. S.; Sun, L., Catalytic activation of H2 under mild conditions by an [FeFe]-hydrogenase model via an active μ-hydride species. Journal of the American Chemical Society 2013, 135 (37), 13688-13691; (h) Zhao, X.; Chiang, C.-Y.; Miller, M. L.; Rampersad, M. V.; Darensbourg, M. Y., Activation of Alkenes and H2 by [Fe]-H2ase Model Complexes. Journal of the American Chemical Society 2003, 125 (2), 518-524; (i) Marhenke, J.; Pierri, A. E.; Lomotan, M.; Damon, P. L.; Ford, P. C.; Works, C., Flash and Continuous Photolysis Kinetic Studies of the Iron‚ÄìIron Hydrogenase Model (μ-pdt)[Fe (CO) 3] 2 in Different Solvents. Inorganic chemistry 2011, 50 (23), 11850-11852.

17. Zhao, X.; Georgakaki, I. P.; Miller, M. L.; Yarbrough, J. C.; Darensbourg, M. Y., H/D exchange reactions in dinuclear iron thiolates as activity assay models of Fe-H2ase. Journal of the American Chemical Society 2001, 123 (39), 9710-9711.

18. (a) Brown-McDonald, J.; Berg, S.; Peralto, M.; Works, C., Photochemical studies of iron-only hydrogenase model compounds. Inorganica Chimica Acta 2009, 362 (2), 318-324; (b) Ridley, A. R.; Stewart, A. I.; Adamczyk, K.; Ghosh, H. N.; Kerkeni, B.; Guo, Z. X.; Nibbering, E. T.; Pickett, C. J.; Hunt, N. T., Multiple-timescale photoreactivity of a model compound related to the active site of [FeFe]-hydrogenase. Inorganic chemistry 2008, 47 (17), 7453-7455.

19. Johnson, M.; Thuman, J.; Letterman, R. G.; Stromberg, C. J.; Webster, C. E.; Heilweil, E. J., Time-resolved infrared studies of a trimethylphosphine model derivative of [FeFe]-hydrogenase. The Journal of Physical Chemistry B 2013, 117 (49), 15792- 15803.

20. Kania, R.; Frederix, P. W.; Wright, J. A.; Ulijn, R. V.; Pickett, C. J.; Hunt, N. T., Solution-phase photochemistry of a [FeFe] hydrogenase model compound: Evidence of photoinduced isomerisation. The Journal of chemical physics 2012, 136 (4), 044521.

21. Zhao, X.; Georgakaki, I. P.; Miller, M. L.; Mejia-Rodriguez, R.; Chiang, C.-Y.; Darensbourg, M. Y., Catalysis of H2/D2 scrambling and other H/D exchange processes by [Fe]-hydrogenase model complexes. Inorganic chemistry 2002, 41 (15), 3917-3928.

22. Lyon, E. J.; Georgakaki, I. P.; Reibenspies, J. H.; Darensbourg, M. Y., Carbon Monoxide and Cyanide Ligands in a Classical Organometallic Complex Model for Fe- Only Hydrogenase. Angewandte Chemie International Edition 1999, 38 (21), 3178-3180.

23. Wang, W.; Rauchfuss, T. B.; Bertini, L.; Zampella, G., Unsensitized photochemical hydrogen production catalyzed by diiron hydrides. Journal of the American Chemical Society 2012, 134 (10), 4525-4528.

114

24. Seyferth, D.; Womack, G. B.; Gallagher, M. K.; Cowie, M.; Hames, B. W.; Fackler Jr, J. P.; Mazany, A., Novel anionic rearrangements in hexacarbonyldiiron complexes of chelating organosulfur ligands. Organometallics 1987, 6 (2), 283-294.

25. Dey, S.; Rana, A.; Crouthers, D.; Mondal, B.; Das, P. K.; Darensbourg, M. Y.; Dey, A., Electrocatalytic O2 Reduction by [Fe-Fe]-Hydrogenase Active Site Models. Journal of the American Chemical Society 2014, 136 (25), 8847-8850.

26. Luehring, P.; Schumpe, A., Gas solubilities (hydrogen, helium, nitrogen, carbon monoxide, oxygen, argon, carbon dioxide) in organic liquids at 293.2 K. Journal of Chemical and Engineering Data 1989, 34 (2), 250-252.

27. Crabtree, R. H., The organometallic chemistry of the transition metals. John Wiley & Sons: 2009.

28. Moragues, M. a. E.; Esteban, J.; Ros-Lis, J. V.; Martinez-Manez, R.; Marcos, M. D.; Martinez, M.; Soto, J.; Sancenon, F. l., Sensitive and selective chromogenic sensing of carbon monoxide via reversible axial CO coordination in binuclear rhodium complexes. Journal of the American Chemical Society 2011, 133 (39), 15762-15772.

29. Kubas, G. J., Fundamentals of H2 binding and reactivity on transition metals underlying hydrogenase function and H2 production and storage. Chemical reviews 2007, 107 (10), 4152-4205.