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Tuning the electronic properties of coordination complexes with sulfur-bridged ligands Brown, Christopher M. 2018

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TUNING THE ELECTRONIC PROPERTIES OF COORDINATION COMPLEXES WITH SULFUR-BRIDGED LIGANDS by Christopher M. Brown M.Chem., The University of Newcastle upon Tyne, 2011    A Thesis Submitted in Partial Fulfilment of the Requirements for the Degree of  DOCTOR of PHILOSOPHY in The Faculty of Graduate and Postdoctoral Studies (Chemistry)    The University of British Columbia (Vancouver) October 2018  © Christopher M. Brown, 2018   ii The following individuals certify that they have read, and recommend to the Faculty of Graduate and Postdoctoral Studies for acceptance, the dissertation entitled: Tuning the Electronic Properties of Coordination Complexes with Sulfur-Bridged Ligands  submitted by Christopher Mark Brown in partial fulfillment of the requirements for the degree of Ph.D. in Chemistry  Examining Committee: Michael Wolf Supervisor  Mark MacLachlan Supervisory Committee Member  Dan Bizzotto Supervisory Committee Member Zachary Hudson University Examiner Keng Chou University Examiner  Additional Supervisory Committee Members: Sarah Burke Supervisory Committee Member  Supervisory Committee Member   iii ABSTRACT  By altering the oxidation state at the sulfur atom of sulfur-bridged ligands we have shown that photophysical properties of coordination complexes can be changed. A new class of sulfur-bridged dipyridyl ligand has been developed for effective “two-level” tuning of the emission colour of Ir(III) complexes. Tuning can occur via the classical method of altering substituents at the pyridyl ring, and/or through oxidizing the sulfur. Changing the oxidation state of the sulfur atom serves as a switch to alter the emissive state from that of mainly 3LC character (blue-green emission) to one of 3MLCT/3LLCT character (yellow emission), evidenced from photophysical characterization and DFT calculations.  The larger bite angle of the sulfur-bridged dipyridyl ligands, in conjunction with additional coordination sites afforded by the oxygen atoms at the sulfur, has led to interesting binding modes when bound to copper(I) in heteroleptic diimine-diphosphine species. On increasing steric constraints about the copper(I) center, bimetallic adducts are isolated, with thermally activated delayed fluorescence the mechanism of photoluminescence. Of the copper(I) species reported in this body of work, one was found to exhibit remarkable thermochromic properties not-before seen in a monometallic copper(I) complex. When isolated as an amorphous powder, Cu-Me-DPSO shows a yellow solid state emission, however when heated to 180 ºC a crystalline powder is formed which shows an orange luminescence. This crystalline powder displays a reversible thermochromic solid state emission, from orange at room temperature to yellow at low temperatures. Using solid state variable temperature excitation and absorption data, this phenomenon is attributed to a change in coordination geometry about the copper atom in the excited-state.   iv With a view to further understanding the impact of sulfur oxidation state on the excited-state properties of coordination complexes, two new Ru(II)-Re(I) dyads were synthesized, for use as CO2 photocatalysts. Connected by a sulfur-bridged bisphenanthroline bridge, oxidation state at the sulfur can be changed which could modulate the electron transfer from the photo-excited Ru(II) to the catalytic Re(I) center. The synthesis, characterization and preliminary photocatalytic results are discussed.   v LAY SUMMARY  The control of light is of great value to the human race, both in utilizing the sun’s rays or generating our own light using electricity. In this work, the manipulation of light is afforded with “coordination complexes”, which consist of building blocks (called ligands) bound to a specific metal. Some of these complexes can emit light, and ligands are used to finely-tune the colour of this light radiated by these species. The ligands herein contain sulfur, a versatile element to which oxygen atoms can be added, but so far an underutilized handle in the chemistry of coordination complexes. This thesis demonstrates that the number of oxygen atoms bonded to the sulfur can have a large impact on the properties of a complex, in particular regarding the colour of light emitted, opening a new avenue of control for chemists to explore.   vi PREFACE  In all chapters I am the principle author of the work, carried out under the supervision of Prof. Michael O. Wolf. I performed all the experiments, with additional contributions and collaborations noted below. A version of Chapter 2 has been previously published as a full paper† written by myself with contributions of M. O. Wolf. Complexes Ir-DPS, Ir-DPSO and Ir-DPSO2 were synthesized by Mitchell J. Kitt. DFT calculations were performed by Zhen (Allen) Xu and the crystallographic data collected and solved by Duane Hean. NMR characterization for complex Ir-4,4'-Me-DPS was assigned by Dr. Maria B. Ezhova.  The contents of Chapter 3 are included in a manuscript in preparation, written by myself with contributions from M. O. Wolf. The project involves a collaboration with Dr. Eli Zysman-Colman and Prof. Ifor D. W. Samuel at the University of St. Andrews, Scotland. DFT calculations were performed by Zhen (Allen) Xu, crystallographic data collected and solved by Veronica Carta and time-resolved emission studies performed by Chenfei Li at the University of St. Andrews.  A version of Chapter 4 has been accepted for publication,‡ authored by myself with contributions from M. O. Wolf. Single-crystal X-ray data was collected and solved by Veronica Carta, and pXRD data collected by Anita Lam.  All experiments in Chapter 5 were conducted by myself, with the exceptions of electrochemical data, which was collected by Emile E. DeLuca, and NMR characterization for Ru-PSO2P-Re which was assigned by Dr. Maria B. Ezhova.  † Brown, C. M.; Kitt, M. J.; Xu, Z.; Hean, D.; Ezhova, M. B.; Wolf, M. O. Inorg. Chem. 2017, 56, 15110–15118. ‡ Brown, C. M.; Carta, V.; Wolf, M. O. Chem. Mater. 2018, 30, 5786–5795.   vii TABLE of CONTENTS  ABSTRACT .................................................................................................................................. iii LAY SUMMARY .......................................................................................................................... v PREFACE ..................................................................................................................................... vi TABLE of CONTENTS .............................................................................................................. vii LIST of TABLES ........................................................................................................................ xii LIST of FIGURES ..................................................................................................................... xiv LIST of SCHEMES ................................................................................................................... xxi LIST of SYMBOLS and ABBREVIATIONS ......................................................................... xxii ACKNOWLEDGEMENTS ..................................................................................................... xxx DEDICATION ......................................................................................................................... xxxi CHAPTER 1: Introduction ........................................................................................................................ 1 1.1 Light ........................................................................................................................................................1 1.2 Photophysical Processes .........................................................................................................................2 1.3 Photophysics of Transition Metal Complexes ........................................................................................5 1.3.1 Photophysics of d6 Metal Complexes ................................................................................................5 1.3.2 Iridium(III) Luminescent Complexes ................................................................................................8 1.3.3 Tuning the Emission of Bis-cyclometallated Ir(III) Complexes ......................................................10 1.3.4 Luminescent Ruthenium(II) Complexes ..........................................................................................12 1.3.5 Photophysics of Cu(I) d10 Metal Complexes ...................................................................................14 1.3.6 Thermally Activated Delayed Fluorescence (TADF) ......................................................................19   viii 1.3.7 Copper(I) TADF Emitters ................................................................................................................21 1.4 Sulfur Oxidation States Alter Photophysical Processes ........................................................................23 1.5 Goals of Thesis .....................................................................................................................................25 CHAPTER 2: Tunable Emission of Iridium(III) Complexes Bearing Sulfur-Bridged Dipyridyl Ligands ....................................................................................................................................................... 27 2.1 Introduction ...........................................................................................................................................27 2.2 Ligand Synthesis and Characterization .................................................................................................28 2.3 Synthesis of [Ir(ppy)2(N^N)][PF6] Complexes .....................................................................................29 2.3.1 Nuclear Magnetic Resonance (NMR) ..............................................................................................30 2.3.2 Single Crystal X-Ray Diffraction ....................................................................................................31 2.4 Photophysical Properties .......................................................................................................................33 2.4.1 Absorption Spectroscopy .................................................................................................................33 2.4.2 Photoluminescence Spectroscopy ....................................................................................................34 2.4.3 Photoluminescence Quantum Yields (PLQY) .................................................................................36 2.4.4 Photoluminescence Lifetimes ..........................................................................................................37 2.5 Electrochemical Properties ...................................................................................................................38 2.6 Theoretical Calculations .......................................................................................................................40 2.7 Conclusions ...........................................................................................................................................43 2.8 Experimental Details .............................................................................................................................45 2.8.1 General .............................................................................................................................................45 2.8.2 Spectroscopy ....................................................................................................................................45 2.8.3 X-ray Crystallography ......................................................................................................................46 2.8.4 Electrochemistry ..............................................................................................................................46 2.8.5 Computational Details ......................................................................................................................47 2.8.6 Synthesis ..........................................................................................................................................47 CHAPTER 3: Substituent and Oxidation State Influence on Sulfur-Bridged Luminescent Copper(I) Complexes Showing Delayed Fluorescence ...................................................................... 57   ix 3.1 Introduction ...........................................................................................................................................57 3.2 Ligand Synthesis and Characterization .................................................................................................59 3.3 Synthesis and Characterization of Cu(I) Complexes ............................................................................60 3.3.1 Nuclear Magnetic Resonance (NMR) Spectroscopy .......................................................................60 3.3.2 Single Crystal X-ray Diffraction ......................................................................................................61 3.4 Electrochemical Properties ...................................................................................................................65 3.5 Solution-state Photophysical Properties ...............................................................................................67 3.5.1 Absorption Spectroscopy .................................................................................................................67 3.5.2 Photoluminescence Spectroscopy ....................................................................................................68 3.6 Solid state Photophysical Properties .....................................................................................................69 3.6.1 Thin Film Photoluminescence .........................................................................................................69 3.6.2 Thin Film Photoluminescence Quantum Yields (PLQY) ................................................................70 3.6.3 Low Temperature Thin Film Photoluminescence ............................................................................71 3.6.4 Thin Film Photoluminescence Lifetimes .........................................................................................73 3.7 Time-resolved Photophysical Properties ..............................................................................................74 3.8 Theoretical Calculations .......................................................................................................................77 3.9 Conclusions ...........................................................................................................................................80 3.10 Experimental Details .............................................................................................................................82 3.10.1 General ........................................................................................................................................82 3.10.2 Spectroscopy ...............................................................................................................................82 3.10.3 X-Ray Crystallography ...............................................................................................................83 3.10.4 Computational Details .................................................................................................................83 3.10.5 Synthesis ......................................................................................................................................84 CHAPTER 4: Thermochromic Solid state Emission of Dipyridyl Sulfoxide Copper(I) Complexes ................................................................................................................................................................ 89 4.1 Introduction ...........................................................................................................................................89 4.2 Ligand Synthesis and Characterization .................................................................................................91   x 4.3 Synthesis and Characterization of [Cu(P^P)(N^N)[BF4] Complexes ..................................................92 4.3.1 Nuclear Magnetic Resonance (NMR) ..............................................................................................92 4.3.2 Single Crystal X-Ray Diffraction ....................................................................................................94 4.3.3 Powder X-Ray Diffraction (pXRD) .................................................................................................97 4.3.4 Differential Scanning Calorimetry (DSC) .....................................................................................100 4.3.5 Fourier Transform Infrared Spectroscopy (FT-IR) ........................................................................101 4.4 Solution-state Photophysical Properties .............................................................................................102 4.4.1 Pro-ligand Photophysical Properties ..............................................................................................102 4.4.2 Cu(I) Complex Photophysical Properties ......................................................................................103 4.5 Cu(I) Complex Solid state Photophysical Properties ..........................................................................106 4.5.1 Photoluminescence Spectroscopy ..................................................................................................106 4.5.2 Low-Temperature Photophysical Properties ..................................................................................107 4.6 Conclusions .........................................................................................................................................113 4.7 Experimental Details ...........................................................................................................................115 4.7.1 General ...........................................................................................................................................115 4.7.2 Spectroscopy ..................................................................................................................................115 4.7.3 X-Ray Crystallography ..................................................................................................................116 4.7.4 Synthesis ........................................................................................................................................117 CHAPTER 5: Sulfur-Bridged Bimetallic Ru(II)-Re(I) Systems – Modulating CO2 Reduction with Sulfur Oxidation State ........................................................................................................................ 122 5.1 Introduction .........................................................................................................................................122 5.1.1 Photocatalytic CO2 Reduction ........................................................................................................122 5.1.2 Ru(II)-Re(I) Supramolecular Photocatalytic Systems ...................................................................125 5.2 Ligand Synthesis and Characterization ...............................................................................................127 5.3 Synthesis and Characterization of Ru(II) Complexes .........................................................................128 5.4 Electrochemical Properties .................................................................................................................130 5.5 Photophysical Properties .....................................................................................................................134   xi 5.5.1 Absorption Spectroscopy ...............................................................................................................134 5.5.2 Photoluminescence Spectroscopy ..................................................................................................136 5.6 Preliminary Photocatalytic Studies .....................................................................................................137 5.7 Conclusions and Future Work ............................................................................................................139 5.8 Experimental Details ...........................................................................................................................140 5.8.1 General ...........................................................................................................................................140 5.8.2 Spectroscopy ..................................................................................................................................140 5.8.3 Electrochemistry ............................................................................................................................141 5.8.4 Photocatalytic Setup .......................................................................................................................142 5.8.5 Synthesis ........................................................................................................................................143 CHAPTER 6: Conclusions and Future Work .................................................................................. 152 REFERENCES ................................................................................................................................... 157 APPENDIX .......................................................................................................................................... 174    xii LIST of TABLES  Table 2-1 Photophysical data of Ir(III) complexes in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature. ...................................................................................................................... 35 Table 2-2 Electrochemical data of Ir(III) complexes in deaerated CH2Cl2 solutions and referenced to Fc/Fc+. ........................................................................................................................... 40 Table 2-3 Ground-state frontier orbital energies for Ir(III) complexes. ....................................... 40 Table 2-4 The calculated adiabatic energy differences of Ir(III) complexes with regard to the experimental emission energies. ....................................................................................... 41 Table 3-1 Selected bond lengths (Å) and angles (˚) of Cu-DPS, Cu-DPSO2, Cu-Me-DPS, and Cu-Me-DPSO2. ................................................................................................................ 65 Table 3-2 Electrochemical data of Cu(I) complexes in CH2Cl2 solutions. .................................. 66 Table 3-3 Solid state photophysical properties of Cu(I) complexes as neat thin films at room temperature. ...................................................................................................................... 70 Table 3-4 ΔE(S1–T1) of Cu(I) complexes calculated from variable temperature solid state emission data. ................................................................................................................................... 72 Table 3-5 S1–T1 energy gaps of Cu(I) complexes as calculated using low temperature steady-state emission and low temperature time-resolved emission spectra. ....................................... 77 Table 3-6 TD-DFT calculated energy levels, oscillator strength and orbital transitions for Cu(I) complexes. ........................................................................................................................ 78 Table 4-1  Selected bond lengths (Å) and angles (º) of Cu-DPSO (N,N-bound), Cu-DPSO (N,O-bound) and Cu-Me-DPSO (N,O-bound). ......................................................................... 95 Table 4-2 Photophysical data for compounds in CH2Cl2 (~1 × 10–5 M). ................................... 104   xiii Table 4-3 Solid state emission data for thin films of Cu-DPSO, a-Cu-Me-DPSO and c-Cu-Me-DPSO. ............................................................................................................................. 107 Table 5-1 Electrochemical data of Ru and Re mono- and bimetallic species. ........................... 133 Table 5-2 Photophysical properties of Ru-PSP-Re, Ru-PSO2P-Re and their respective monometallic model complexes.[a] ................................................................................. 137    xiv LIST of FIGURES  Figure 1-1 Depiction of a simplified Jablonski diagram. Adapted from reference.16 .................... 3 Figure 1-2 (a) Octahedral structure of a ML6 transition metal complex. (b) Molecular orbital diagram showing the absorptive transitions for an octahedral transition metal complex. Adapted from reference.17 ................................................................................................... 6 Figure 1-3 (a) fac-Ir(ppy)3 (1) as reported by King et al.25 (b) Structure of a bis-cylometallated Ir(III) species. (c) Structure of [Ir(ppy)2(bpy)][PF6] (2). .................................................... 9 Figure 1-4 The electronic effects of substitution at the cyclometallating ligands on luminescence colour.13,30,31 ...................................................................................................................... 11 Figure 1-5 The electronic effects of substitution at the ancillary ligands on luminescence colour.13,32,33 ...................................................................................................................... 11 Figure 1-6 (a) Structure of [Ru(bpy)3]2+ (7). (b) Simplified Jablonski diagram illustrating the excited-state processes in [Ru(bpy)3]2+ (7). Adapted from references.39,40 ....................... 13 Figure 1-7 Donor–acceptor–chromophore triads based on [Ru(bpy)3]2+ by Meyer et al.46,47 ..... 14 Figure 1-8 Tetrahedral geometry of [Cu(N^N)2]+ species. .......................................................... 15 Figure 1-9  Simplified excited-state pathways of Cu(I) diimine complexes. After undergoing excitation, a Jahn-Teller distortion occurs, leaving an open site available for exciplex formation, following which the complex can relax radiatively or non-radiatively. Adapted from references.57,58 .......................................................................................................... 16 Figure 1-10 Increasing the steric bulk in the 2,9-positions of phen restricts the Jahn-Teller distortion at the Cu(I) center in the excited-state and increases emission intensity.65 ...... 17   xv Figure 1-11 A sterically congested Cu(I) photosensitizer used for hydrogen photocatalysis by Castellano et al.67 .............................................................................................................. 18 Figure 1-12 Jablonski diagram illustrating the process of thermally activated delayed fluorescence (TADF). Adapted from reference.71 .................................................................................. 20 Figure 1-13 Neutral Cu(P^P)(N^N) emitters reported by Yersin et al.75 .................................... 22 Figure 1-14 Adding steric bulk can decrease the degree of geometric distortion in the excited-state and can increase the emission intensity in Cu(I) TADF emitters in solution, as reported by Linfoot, Robertson and Yersin.76 Image reproduced with permission from the American Chemical Society, Copyright © 2014. .............................................................................. 23 Figure 1-15 Structures of sulfur-bridged dimers synthesized by our group77 showing (a) increased PLQY as the oxidation state at sulfur is increased. (b) Emission spectra in CH2Cl2 solutions for T3SOn in comparison to the parent arene, terthiophene (T3). Image reproduced with permission from the American Chemical Society, Copyright © 2013. ............................ 24 Figure 2-1 Expanded NMR spectrum of Ir-DPSO in CD2Cl2 at 25 ºC (850 MHz). ................... 30 Figure 2-2 Crystal structure of Ir-DPS (Λ conformer). Ellipsoids plotted at 50% probability and H atoms and hexane molecules removed for clarity. ........................................................ 32 Figure 2-3 Absorption spectra of (a) Pro-ligands DPSOn and 4,4'-Me-DPSOn and (b) complexes Ir-DPSOn and Ir-4,4'-Me-DPSOn. All spectra were recorded in ~ 2 ´ 10–5 M CH2Cl2 solutions at room temperature. .......................................................................................... 33 Figure 2-4 Photoluminescence spectra of Ir(III) complexes (λex = 390 nm). All spectra were recorded in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature and sparged with Ar for 25 min. ................................................................................................................................... 34   xvi Figure 2-5 Photoluminescence lifetimes of Ir(III) complexes (λex = 370 nm). All spectra were recorded in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature and sparged with Ar for 25 min. ................................................................................................................................... 37 Figure 2-6 Cyclic voltammograms of Ir(III) in dry, N2-sparged CH2Cl2 at room temperature. (a) Ir-DPSOn complexes. (b) Ir-4,4'-Me-DPSOn complexes. .............................................. 39 Figure 2-7 (a) Schematic diagram of the adiabatic energy difference (∆E) between S0 and T1, and the emission energy (Eem) from T1. (b) Unpaired-electron spin density contours for Ir-DPS and Ir-DPSO2. .................................................................................................................. 42 Figure 3-1 ORTEP representations of single crystal structures of (a) Cu-DPS; (b) Cu-Me-DPS; (c) Cu-DPSO2; (d) Cu-Me-DPSO2. Ellipsoids are plotted at the 50% probability level. Hydrogen atoms and solvent molecules are removed for clarity. ..................................... 63 Figure 3-2 Absorption spectra of (a) pro-ligands and (b) Cu(I) complexes in CH2Cl2 solutions. All spectra were recorded in ~ 2 ´ 10–5 M CH2Cl2 solutions at room temperature. ............... 67 Figure 3-3 Emission spectra of Cu(I) complexes in CH2Cl2 solutions at room temperature. Concentrations ~ 2 ´ 10–5 M. λex = 370 nm. ..................................................................... 68 Figure 3-4 Solid state emission spectra of Cu(I) complexes as thin films drop-cast from MeOH on quartz. λex = 370 nm. ......................................................................................................... 69 Figure 3-5 Emission lifetime measurements of the four Cu(I) complexes. (a) Cu-DPS and Cu-DPSO2 measured in a 50 ns time regime. (b)  Cu-Me-DPS and Cu-Me-DPSO2 measured in a 50 ns time regime. (c) Cu-DPS and Cu-DPSO2 measured in a 5 µs time regime. (d) Cu-Me-DPS and Cu-Me-DPSO2 measured in a 5 µs time regime. All films drop-cast from MeOH on to quartz and spectra collected in air. Prompt is the decay profile of the   xvii NanoLED excitation source (λex = 370 nm) with a pulse width of 1.2 ns.Error! Bookmark not defined. Figure 3-6 Solid state emission spectra of (a) Cu-DPS and Cu-DPSO2; (b) Cu-Me-DPS and Cu-Me-DPSO2 at room temperature (solid line) and 77 K (dashed line). Thin films drop-cast from MeOH on quartz. λex = 370 nm. ............................................................................... 71 Figure 3-7 Time-resolved solid state photoluminescence spectra of (a) Cu-DPS; (b) Cu-DPSO2; (c) Cu-Me-DPS; (d) Cu-Me-DPSO2 at 77 K. Grey traces represent emission from the S1 state, and pink traces emission from the T1 state. Black lines indicate the determination of E0–0. Samples prepared as neat powders. .......................................................................... 76 Figure 3-8 Frontier orbital plots illustrating the HOMO and LUMO distributions for Cu(I) complexes ......................................................................................................................... 79 Figure 4-1 1H NMR spectra of Cu-DPSO and Cu-Me-DPSO in CD3CN at 25 ºC (400 MHz). 92 Figure 4-2 Low temperature 1H NMR spectra of Cu-Me-DPSO in CD3CN from 25 ºC to –40 ºC (400 MHz). ........................................................................................................................ 93 Figure 4-3 Solid state structures of (a) Cu-DPSO (N,N and N,O coordination modes, same unit cell). (b) Cu-Me-DPSO (N,O coordination mode). Ellipsoids are plotted at the 50% probability level. Hydrogen atoms and solvent molecules are removed for clarity. ........ 94 Figure 4-4 (a) pXRD of Cu-DPSO compared to the simulated powder pattern generated from the single crystal structure (N,N and N,O bound). (b) pXRD spectra of amorphous a-Cu-Me-DPSO (grey) heated to 180 ºC (teal) and then cooled to room temperature, retaining crystalline species c-Cu-Me-DPSO (pink) ...................................................................... 97 Figure 4-5 (a) pXRD of Cu-DPSO during heating to 180 ºC. (b) pXRD of Cu-DPSO during cooling to –180 ºC. (c) pXRD comparing c-Cu-Me-DPSO (pink) to the diffraction pattern   xviii generated from the single crystal data of N,O-bound Cu-Me-DPSO (black). a-Cu-Me-DPSO: amorphous state. c-Cu-Me-DPSO: crystalline state. (d) pXRD cycle of c-Cu-Me-DPSO cooled from room temperature to –180 ºC and back to room temperature. .......... 98 Figure 4-6 DSC data of (a) Cu-DPSO and (b) Cu-Me-DPSO. Both samples were scanned over a temperature range of 25–180 ºC with three cycles. ........................................................ 100 Figure 4-7 FT-IR spectra of a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) as neat solids at 25 ºC. ............................................................................................................................... 101 Figure 4-8 (a) UV-vis absorption spectra of DPSO (blue trace) and Me-DPSO (pink trace). (b) Normalized excitation (dashed line) and emission (solid line) spectra of DPSO (blue trace) and Me-DPSO (pink trace). λex = 310 nm. Solutions of CH2Cl2 at 25 ºC, concentrations ~1 × 10–5 M. ......................................................................................................................... 102 Figure 4-9 (a) Absorption and emission spectra of Cu-DPSO (blue) and Cu-Me-DPSO (pink) in sparged (Ar) CH2Cl2 solutions (concentrations ~3 × 10–5 M). Emission spectra are corrected for absorbance at the excitation wavelength (lex = 380 nm). (b) Absorption spectra of a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) in CH2Cl2 at 25 ºC (concentrations ~5 × 10–5 M). ......................................................................................... 103 Figure 4-10 Excitation spectra of Cu-DPSO and Cu-Me-DPSO in CH2Cl2 at 25 ºC (concentrations ~3 × 10–5 M). ......................................................................................... 105 Figure 4-11 (a) Room temperature solid state emission of Cu-DPSO (blue), a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) drop-cast from MeOH on quartz. (lex = 380 nm); (b) Photograph of the three complexes under UV hand lamp irradiation. (lex = 365 nm) ... 106 Figure 4-12 Variable-temperature solid state emission spectra of (a) Cu-DPSO and (b) a-Cu-Me-DPSO from –196 to 27 ºC. Thin films dropcast on quartz from MeOH. (λex = 380 nm)107   xix Figure 4-13 (a) Variable temperature solid state emission spectra of c-Cu-Me-DPSO from –196 to 27 ºC (λex = 380 nm). (b) Variable temperature, solid state emission lifetimes of c-Cu-Me-DPSO at –196 and 25 ºC. To maintain accuracy, a 50 µs time range was used for –196 ºC and 20 µs for 25 ºC. Longest time component for each complex reported (λex = 370 nm). Thin films dropcast on quartz from MeOH. ................................................................... 108 Figure 4-14 (a) Photographs showing the thermochromic behaviour of thin films of c-Cu-Me-DPSO in the ground-state and under UV lamp excitation (lex = 365 nm) at different temperatures. (b) The proposed excited-state conformational change for the thermochromic emission of c-Cu-Me-DPSO. ......................................................................................... 110 Figure 4-15 (a) Variable temperature UV-vis absorption spectra of c-Cu-Me-DPSO from –196 to 27 ºC. (b) Variable temperature excitation spectra of c-Cu-Me-DPSO measuring the HE emission at 556 nm. (c) Variable temperature excitation spectra of c-Cu-Me-DPSO measuring the LE emission at 658 nm. (d) Variable temperature excitation spectra of Cu-DPSO measuring the emission at 524 nm. Thin films dropcast on quartz from MeOH. 111 Figure 4-16 Solid state emission spectra of the crystallization process from a-Cu-Me-DPSO to c-Cu-Me-DPSO. At 87 ºC (black trace) the thin film shows HE emission, and at 107 ºC (initial) the LE band has already started to grow in. Over the next 30 min the HE band diminishes and the LE band increases in intensity (pink trace). ..................................... 113 Figure 5-1 Depiction of a supramolecular photocatalyst assembly for CO2 reduction. ............ 124 Figure 5-2 Selected Ru(II)-Re(I) systems.209-211 ........................................................................ 126 Figure 5-3 (a) 1H NMR shifts for the bridging ligand in Ru-PSO2P-Re. Black text designates signals from one diastereotopic species, red text for the other. Chemical shifts with only one number indicates overlapping, indistinguishable species. (b) HSQC NMR spectrum   xx (CD3CN, 850 MHz, 25 ºC) showing two diastereotopic carbon signals correlating to one broad singlet. ................................................................................................................... 130 Figure 5-4 Cyclic voltammograms of (a) Ru(II) monometallic complexes; (b) Re(I) monometallic complexes; (c) Ru(II)-Re(I) bimetallic species. All spectra taken in dry, N2-sparged CH3CN solutions at room temperature. .......................................................................... 132 Figure 5-5 Absorption spectra of PSP and PSO2P in CH2Cl2 solutions at room temperature. . 134 Figure 5-6 Absorption spectra of (a) Ru-PSP-Re (blue) and its monometallic model complexes Ru-PSP (grey) and Re-PSP (purple). (b) Ru-PSO2P-Re (pink) and model complexes Ru-PSO2P (grey) and Re-PSO2P (teal). CH3CN solutions of concentration ~3 × 10–6 M at room temperature. ........................................................................................................... 135 Figure 5-7 (a) Emission spectra of Ru-PSP (light grey), Ru-PSO2P (dark grey), Ru-PSP-Re (blue) and Ru-PSO2P-Re (pink). Spectra were normalized to the absorbance at the excitation wavelength (λex = 450 nm). (b) TCSPC data for Ru-containing monometallic and bimetallic species (λex = 453 nm). CH3CN solutions of concentration ~3 × 10–6 M at room temperature. ........................................................................................................... 136 Figure 5-8 Gas chromatograms of (a) Ru-PSP-Re and (b) Ru-PSO2P-Re. ............................. 138 Figure 5-9 (a) Photocatalytic setup under irradiation. (b) LED housing used for photocatalytic experiments. (c) Emission profile of the LED source (peak wavelength = 503 nm, FWHM = 30 nm). ......................................................................................................................... 142    xxi LIST of SCHEMES  Scheme 2-1 Synthesis of Iridium(III) Complexes Ir-DPSOn and Ir-4,4'-Me-DPSOn (where n = 0, 1 or 2). ........................................................................................................................... 29 Scheme 3-1 Synthesis of Cu(I) Complexes .................................................................................. 59 Scheme 4-1 Synthesis of Cu(I) species Cu-DPSO and Cu-Me-DPSO. ...................................... 91 Scheme 5-1 Synthesis of Ru(II)-Re(I) complexes Ru-PSP-Re and Ru-PSO2P-Re. ................ 128    xxii LIST of SYMBOLS and ABBREVIATIONS  1MLCT   singlet metal-to-ligand charge transfer 3MLCT   triplet metal-to-ligand charge transfer 3LLCT   triplet ligand-to-ligand charge transfer 1H    proton 13C    carbon-13 31P    phosphorus-31 4,4′-Me-DPS   bis(4-methylpyridin-2-yl)sulfane 4,4′-Me-DPSO  2,2′-sulfinylbis(4-methylpyridine) 4,4′-Me-DPSO2  2,2′-sulfonylbis(4-methylpyridine) Å    Angström (10–10 m) AC    alternating current ATR    attenuated total reflection a.u.    arbitrary units bpy    2,2′-bipyridine BNAH    1-benzyl-1,4-dihydronicotinamide Δ    crystal field splitting ºC    degrees Celsius C^N    cyclometallating ligand CCD    charge-coupled device cm    centimeter cm–1    wavenumber   xxiii COSY    correlation spectroscopy CRT    cathode ray tube CS    charge-separated CT    charge transfer CV    cyclic voltammetry Δ    difference d    doublet dbp    2,9-dibutyl-1,10-phenanthroline DC    direct current D–C–A   donor–chromophore–acceptor dfppy    2-(2,4-difluorophenyl)pyridine) DFT    density functional theory dma    dimethylamine dmbpy   4,4′-dimethyl-2,2′-biyridyl dmgH    dimethylglyoxime (anionic form) dmp    2,9-dimethyl-1,10-phenanthroline DMT    N,N-dimethyl-p-toluidine dnpp    2,9-dineopentyl-1,10-phenanthroline dop    2,9-dioctyl-1,10-phenanthroline DPS    di(priding-2-yl)sulfane DPSO    2,2′-sulfinyldipyridine DPSO2   2,2′-sulfonyldipyridine dsbp    2,9-di(sec-butyl)-1,10-phenanthroline   xxiv dsbtmp   2,9-di(sec-butyl)-3,4,7,8-tetramethyl-1,10-phenanthroline DSSC    dye-sensitized solar cell E0–0    zero–zero energy Eem    emission energy EHOMO    HOMO energy ELUMO    LUMO energy Eox*    excited-state oxidation potential Ered*    excited-state reduction potential ΔE(S1–T1)   energy gap between S1 and T1 states EDG    electron-donating group ESI-MS   electrospray ionization mass spectrometry Et2O    diethyl ether     EtOH    ethanol eq.    equivalent eV    electron volts EWG    electron-withdrawing group Exciplex   excited-state complex f    oscillator strength fac    facial Fc    ferrocene fs    femtosecond (10–15 s) FT-IR    Fourier transform infrared spectroscopy FWHM   full width at half maximum   xxv g    grams GC    gas chromatography GC-MS   gas chromatography-mass spectrometry h    hours HOMO   highest occupied molecular orbital HMBC   heteronuclear multiple-bond correlation spectroscopy HR-ESI MS   high resolution electrospray ionization mass spectrometry  HSQC    heteronuclear single-quantum correlation spectroscopy Hz    hertz IC    internal conversion IQE    internal quantum efficiency IR    infrared ISC    intersystem crossing J–T    Jahn–Teller K    Kelvin kISC    rate of intersystem crossing knr    rate of non-radiative decay kr    rate of radiative decay kHz    kilohertz (103 Hz) λ    wavelength λem    wavelength of emission λex    wavelength of excitation λmax    wavelength at peak maximum   xxvi L^X    ancillary ligand LC    ligand-centered LCD    liquid crystal display LED    light-emitting diode LEEC    light-emitting electrochemical cell LLCT    ligand-to-ligand charge transfer LMCT   ligand-to-metal charge transfer LUMO   lowest unoccupied molecular orbital µs    microsecond m    multiplet m/z    mass-to-charge ratio M    molar MC    metal-centered  Me-DPS   bis(6-methylpyridin-2-yl)sulfane Me-DPSO   6,6′-sulfinylbis(2-methylpyridine) Me-DPSO2   6,6′-sulfonylbis(2-methylpyridine) MeCN    acetonitrile MeOH   methanol MHz    megahertz (106 Hz) min    minutes mL    milliliter  MLCT   metal-to-ligand charge transfer mol    mole   xxvii MS    mass spectrometry mV    millivolts N^N    diimine ligand N^P    imine-phosphine ligand nm    nanometer NMR    nuclear magnetic resonance NOESY   nuclear Overhauser effect spectroscopy ns    nanosecond (10–9 s) OERS    one-electron reduced species OLED    organic light-emitting diode ORTEP   Oak Ridge Thermal Ellipsoid Plot P^P    diphosphine ligand PDT    photodynamic therapy phen    1,10-phenanthroline PLQY    photoluminescence quantum yield POP    bis(2-(diphenylphospanyl)phenyl) ether ppy    2-phenylpyridine ps    picosecond (10–12 s) PS    photosensitizer PSP    di(1,10-phenanthrolin-5-yl)sulfane PSO2P    5,5′-sulfonylbis(1,10-phenanthroline) PTI    Photon Technology International pXRD    powder X-ray diffraction   xxviii py    pyridine pz2BH2   bis(pyrazol-1-yl)borohydrate pz2Bph2   bis-(pyrazol-1-yl)biphenylborate) ΦCO    quantum yield of CO q    quintet RISC    reverse intersystem crossing ROESY   rotating frame nuclear Overhauser effect spectroscopy rpm     revolutions per minute S0    ground-state S1    first excited singlet state S2    second excited singlet state s    seconds S    sulfide SCXRD   single crystal X-ray diffraction SO    sulfoxide SO2    sulfone SOC    spin-orbit coupling SWV    squarewave voltammetry t     triplet T    Tesla T1    first excited triplet state TADF    thermally activated delayed fluorescence TCI    Tokyo Chemical Industry Co., Ltd.   xxix TCSPC   time-correlated single photon counting TD-DFT   time-dependent density functional theory Tg    glass transition temperature TEOA    triethanolamine TLC    thin layer chromatography tmbpy    4,4′,6,6′-tetramethyl-2,2′-biyridyl TOF    turn over frequency TON    turn over number TONCO   turnover number for carbon monoxide UBC    University of British Columbia UV    ultraviolet V    volts v/v    volume/volume Vis    visible W    watts XRD    X-ray difraction    xxx ACKNOWLEDGEMENTS Firstly I would like to thank my boss, Mike Wolf, for his support and guidance during this endeavor. The trust you put in me, in addition to the creative freedom afforded with my work, has helped me grow not just as a scientist, but as a person. From the bottom of my heart, thank you. I have been fortunate to work with a number of great collaborators to whom I am very grateful. In particular I would like to thank Dr. Saeid Kamal at LASIR  his vital help regarding excited-state lifetimes; Milan Coschizza, whose electrical engineering talents babied the fluorimeter through some tough times; Veronica Carta, for always enthusiastically running my (sometimes) crystalline species; and of course Dr. Maria Ezhova, whose patience and tremendous talent has shown me that NMR is not as scary as it seems. During my fifth year I had the pleasure of working with an undergraduate student, Mitchell Kitt, whose unrelenting drive and determination laid the foundation for the work in Chapter 2. Thank you. I’d like to thank all the members of the Wolf group (past and present) for making this such an enjoyable period of my life. Special thanks to Marek Majewski, Ashlee Howarth and Yang Cao for your guidance when I started, and for continuing to support me in the years after your graduation. To the lads of the office (BRFL), thanks for the good times in and out of the lab. Janet, Emile, Duane, Allen, Cathrin, Val, Kyle, Lev, Damon, Amber and Mitch, it’s been a blast. To Rebecca, without you these last few years in Vancouver would not have been the same. You are an inspiration to me, and your unwavering support has been invaluable. I can’t wait for what the future holds for us. And finally, my greatest appreciation goes out to my family. Who would have thought after dropping me off in Newcastle on a damp day in 2006 I would still be at university and half-way around the world? Thank you for your support, I love you all dearly.    xxxi DEDICATION    Mum, Dad, Ellie and Libby       1 CHAPTER 1: Introduction  1.1 Light In the fifth century BC, Greek philosopher Empedocles proposed that four elements – fire, earth, air and water – made up the composition of everything in the universe. He suggested that the goddess Aphrodite crafted the human eye out of these four elements.1  She then lit the fire, causing it to radiate from the eye, allowing us to see. While this view was at the time generally accepted, in ~ 50 BC Lucretius wrote in On the Nature of Things:2 “The light and heat of the sun; these are composed of minute atoms which, when they are shoved off, lose no time in shooting right across the interspace of air in the direction imparted by the shove.”  This surprisingly accurate interpretation was derided until around 1000 years later, when al-Haytham argued that sight is due to a light beam from the sun reflecting off an object and entering our eyes.3 Pioneering work by Gassendi, Newton, Hooke and Huygens – to name but a few – advanced our understanding, but it is the modern day view of Maxwell’s theory of electromagnetic radiation that is now widely accepted.4,5 Light is intertwined intimately with the natural world and the human race. From the day-night cycle that is linked to our circadian rhythm, to the photosynthetic machinery of plants, algae and cyanobacteria, to the Aurora Borealis; light is not just a vital source of energy, but also a thought-provoking and emotive entity. At the onset of the Information Age, a broad range of applications for light are helping to drive civilization: harnessing solar irradiation is seen to be an alternative to reduce our dependence   2 on fossil fuels;6 luminescence-based bioimaging can help detect tumours;7 quantum computing could perform calculations much quicker than binary digital computers based on transistors.8 Of particular interest are photofunctional materials for illumination applications. In recent decades display technologies have advanced from cathode ray tube (CRT) devices to liquid-crystal displays (LCD),9 to electroluminescent materials such as light-emitting diodes (LEDs), organic light-emitting diodes (OLEDs)10 or light-emitting electrochemical cells (LEECs).11 OLEDs in particular, have been successfully commercialized, with a market predicted to be valued at 49 billion US$ by 2023.12 The advantage for these newer technologies is that the sub-pixels themselves emit light, negating the need for a backlight and thus forming thinner devices that are more energy-efficient. Iridium(III) complexes have been prepared for uses in both OLEDs and LEECs due to their high photoluminescence quantum yields (PLQY) and highly tuneable emission colour.13,14 Iridium however is a precious metal with low Earth abundance and so cheaper, more sustainable alternatives are sought, such as emitters based on copper(I) complexes. The properties of these materials are dictated by their excited-state dynamics, which for a chosen metal center can be tuned by altering the electronics at the chelating ligands. In this work a number of new complexes with sulfur-functionalized ligands are reported, illustrating that subtle electronic changes can greatly impact the photophysical properties of a coordination complex.  1.2 Photophysical Processes Photophysical processes occur following excitation of a system by electromagnetic radiation (photons). The subsequent events induce a change of states through radiative and non-radiative transitions, resulting in a physical outcome with no chemical change. These   3 photophysical processes were illustrated by Polish physicist Aleksander Jablonski in 1933, with a simplified Jablonski diagram given in Figure 1-1.15 The diagram depicts the different electronic states of a system and their relative energies. Singlet electronic states are denoted by S0, S1, S2 etc. and triplet states as T1, T2, etc. and within each of these states are vibrational levels which are expressed in order of increasing energy by v = 0, v =1, v = 2, etc. Radiative and non-radiative transitions can occur between these energy levels and are depicted in Figure 1-1 as solid and dashed lines, respectively.   Figure 1-1 Depiction of a simplified Jablonski diagram. Adapted from reference.16   In order to absorb light and induce an electronically-excited state, the incoming photon must have an energy equal to the difference between the two electronic states to be able to promote an electron from the ground-state (S0) to the higher energy level (S1, S2, etc.). This absorption   4 process occurs on a very fast timescale of 10–15 s. The resultant excited-state is inherently unstable and thus will relax back to the ground-state via a number of possible relaxation processes:17 • Vibrational relaxation (10–14 to 10–11 s): occurs within an electronic state between a vibrationally-excited state (v ≥ 1) and v = 0. Excess energy is dissipated as heat. • Internal conversion (IC) (10–14 to 10–11 s): occurs between isoenergetic states of the same spin multiplicity, with IC between excited states (e.g. S2 à S1) occurring much faster than IC between S1 and S0. Excess energy is dissipated as heat. • Intersystem Crossing (ISC) (10–8 to 10–3 s): the spin-forbidden transition between two electronic states with different spin multiplicity, such as S1 à T1. • Fluorescence (10–9 to 10–7 s): the spin-allowed radiative transition between states of the same multiplicity to the ground-state (e.g. S1 à S0). Results in the emission of a photon. • Phosphorescence (10–3 to 102 s): the spin-forbidden radiative transition between states of different multiplicity to the ground-state (e.g. T1 à S0). Results in the emission of a photon. Generally only observed in organic molecules at low temperatures where non-radiative decay is minimized; or when utilizing a heavy atom to induce spin-orbit coupling (SOC) promotes the spin-forbidden transition.  As discussed above, the electronic processes involved occur on a very fast timescale and can be considered as vertical transitions and explained in terms of the Franck-Condon principle. The efficiency of the absorption and radiative processes are therefore dependent on the overlap between the square of the wavefunctions of the excited-state and the ground-state.     5 1.3 Photophysics of Transition Metal Complexes Luminescent transition metal complexes have applications in a wide variety of fields, such as solar energy conversion,6 light-emitting devices,14 photocatalysis,18 molecular electronics19 and as bioimaging probes,20 to name a few. This is due to the high photostability, long luminescence lifetimes and large Stokes shifts exhibited by transition metal complexes, in addition to the high degree of tunability that can be afforded with these materials. In this section, the photophysical properties of the metals utilized in this work, namely copper, iridium and ruthenium, are discussed.  1.3.1 Photophysics of d6 Metal Complexes  Transition metal complexes with a d6 electronic configuration, such as Ir(III) and Ru(II) give a bright phosphorescence at room temperature both in solution and in the solid state due to fast intersystem crossing that is promoted by the strong SOC induced by the heavy-atom. These d6 complexes adopt an octahedral geometry (Figure 1-2(a)) and the photophysical processes involved can be understood using a molecular orbital diagram (Figure 1-2(b)).    6  Figure 1-2 (a) Octahedral structure of a ML6 transition metal complex. (b) Molecular orbital diagram showing the absorptive transitions for an octahedral transition metal complex. Adapted from reference.17   In a free transition metal ion, the five d-orbitals are degenerate. However, once the metal is bound as an octahedral complex the crystal field of the ligands splits these into triply degenerate t2g (non-bonding) and doubly degenerate eg* molecular orbitals (Figure 1-2(b)).21 The extent of this splitting (Δ) is determined by the crystal field strength of the ligands and the central metal ion, and thus changes at the ligand, the geometry or the metal ion can have a substantial impact on the luminescence properties.22 The distribution of electrons between the t2g and eg energy levels is dictated by the magnitude of Δ. In a system with a large Δ (low spin), according to Hund’s rule it is energetically more favourable to pair the electrons in the t2g level. If the magnitude of Δ is less than that of the pairing energy then the electrons are distributed to give the maximum number of unpaired electrons, resulting in a high spin system. High spin systems however are not favourable   7 for emission, and so only low spin metal complexes where all six d electrons are paired in the t2g orbitals (for a d6 system) are discussed below.  The classification for the excited-states is derived from the various orbital configurations and the transitions which occur between them, and are visualized in Figure 1-2(b). In a d6 strong field complex all spins are paired in its non-excited-state, denoted as a ground-state singlet (S0). Upon irradiating with light of sufficient energy one of these paired electrons is promoted to an empty orbital of a higher energy to give an excited-state. Three different types of excited-state are seen: metal-centered (MC) states resulting from a d–d transition; ligand-centered states (LC) arising from a π–π* transition; and charge transfer (CT) states. The CT states can occur from either the promotion of a d electron to the π* antibonding orbital of the ligand to give a metal-to-ligand charge transfer (MLCT), or from promotion of a π bonding orbital electron from the ligand to an empty d orbital on the metal, yielding a ligand-to-metal charge transfer (LMCT) state.22 In each instance the excited-state can be singlet or triplet in nature, with the triplet state lower in energy than its corresponding singlet state. The MC d–d states are very weakly absorbing because, even if spin-allowed, they are Laporte forbidden (t2g à eg). As such they are often too weak to be seen in comparison to the more intense CT transitions, which are more strongly allowed and have favourable characteristics such as higher molar extinction coefficients and shorter emission lifetimes.17  Emission from d6 complexes always occurs from the lowest excited-state and, because of efficient SOC due to the heavy metal center, the dominant emission is from a triplet state. Thus when designing a luminescent transition complex a number of factors must be considered:22  1. Avoid d–d excited states. These transitions are inherently unstable and so CT or LC excited-states are preferable.   8 2. Lowest d–d states should be higher in energy than the emitting state. This prevents thermal population from the emitting state. 3. Have a high degree of SOC. This increases the probability of emission and allows the radiative decay to be competitive with non-radiative decay. 4. Encourage mixing between LC states and CT states. Purely π–π* phosphorescence is too long-lived to be efficient, and so mixing with the more-allowed CT state increases the emission probability. 5. The emitting state should not be too low in energy. The energy gap law dictates that radiationless processes are more efficient as the emitting state approaches the ground-state.23 Though of course when designing an emitter the desired colour dictates the energy required.  1.3.2 Iridium(III) Luminescent Complexes  Phosphorescent emitters have been intensively researched due to their ability to harvest both the electrogenerated singlet and triplet excitons in an electroluminescent device, leading to a theoretical internal quantum efficiency (IQE) of 100%.24 The study of iridium systems is a relatively mature field and Ir(III) complexes fall into one of two categories: tris- cyclometallated neutral complexes; and bis-cyclometallated ionic complexes bearing an ancillary ligand. In 1985, King et al. described the synthesis of fac-Ir(ppy)3 (1, ppy = 2-phenylpyridine, Figure 1-3(a)),25 a neutral cyclometallated complex which, in 1999, was used as a green phosphorescent dopant in an OLED that still retains the highest IQE of an electroluminescent material using an Ir(III) complex.26    9  Figure 1-3 (a) fac-Ir(ppy)3 (1) as reported by King et al.25 (b) Structure of a bis-cylometallated Ir(III) species. (c) Structure of [Ir(ppy)2(bpy)][PF6] (2).  Unlike neutral iridium(III) species, ionic complexes feature desirable characteristics due to their charged nature, such as enhanced solubility, in addition to ease of design, synthesis and purification, which has piqued interest industrially with some complexes showing promise in LEECs.11,27 In this section, bis-cyclometallated Ir(III) complexes of the structure [Ir(C^N)2(L^X)]n (Figure 1-3(b)) will be discussed, whereby C^N is a cyclometallating ligand and L^X is the ancillary ligand. The charge of the complex can be modulated with the choice of ancillary ligand by choosing neutral, monoanionic or dianionic species (N^N, O^O or C^C, respectively).27 Of the transitions discussed in Section 1.3.1, those seen in Ir(III) complexes are predominantly of mixed MLCT/LLCT and LC character. Once in the singlet excited-state, ISC can occur due to the strong SOC enabled by the heavy iridium atom, yielding a mixed triplet excited-state of 3MLCT and 3LC nature.28 Commonly in bis-cyclometallated Ir(III) complexes the cyclometallating ligand is linked to the 3LC state and the ancillary ligand with the 3MLCT state. In a typical cationic complex such as [Ir(ppy)2(bpy)][PF6] (2, bpy = 2,2′-bipyridine, Figure 1-3(c)), its highest occupied molecular orbital (HOMO) is a mixture of Ir dπ orbitals and phenyl π orbitals   10 of the cyclometallating ppy ligands, while the lowest unoccupied molecular orbital (LUMO) lies on the ancillary bpy ligand.13 As the HOMO and LUMO are often located on different ligands in cationic Ir(III) complexes, it is therefore possible to finely-tune the emission colour by independently modifying the C^N and L^X ligands.29  1.3.3 Tuning the Emission of Bis-cyclometallated Ir(III) Complexes As discussed above, altering the electronics of the cyclometallating or ancillary ligands can result in a rational tuning of emission colour of bis-cyclometallated Ir(III) complexes to cover the visible spectrum. Of the C^N ligands used, ppy and its various derivatives are amongst the most common, and the effects of altering these ligands is highlighted in Figure 1-4. On the introduction of an electron-withdrawing group (EWG), such as fluorine, it is possible to stabilize the HOMO and thus blueshift the emission. For example, [Ir(ppy)2(bpy)][PF6] (2) gives yellow emission at 585 nm in acetonitrile solution;13 however, Park et al. showed that on fluorinating the phenyl ring of the ppy ligands in [Ir(dfppy)2(bpy)][PF6] (3, dfppy = 2-(2,4-difluorophenyl)pyridine)) a blue-shift of the emission maximum of over 30 nm is seen.30 Conversely, the emission can be bathochromically shifted through the addition of electron-donating groups (EDG) to the ppy ligands. Hasan and co-workers demonstrated that when substituting the phenyl ring of ppy with methoxy groups in the 3- and 4-positions (4) gives a strongly red-shifted emission with a maximum of 730 nm in acetonitrile solution.31    11  Figure 1-4 The electronic effects of substitution at the cyclometallating ligands on luminescence colour.13,30,31  In comparison to modification at the C^N ligands, altering the electronics at the L^X offers a more convenient and flexible approach to altering the luminescence colours of cationic Ir(III) complexes because the LUMO exclusively lies on the ancillary ligand. In Figure 1-5 [Ir(ppy)2(bpy)][PF6] (2) is again used as the standard. In substituting the bpy with EDGs such as dimethylamino (dma) groups to give [Ir(ppy)2(dma-bpy)][PF6] (5) it is possible to blue-shift the emission by destabilizing the LUMO.32 Using the same rationale, employing EWGs will stabilize the LUMO and red-shift the luminescence. Rodriquez-Redono et al. introduced carboxylate groups to the 4- and 4′-positions of the bpy in complex 6 and were able to achieve deep-red emission with a maximum at 687 nm.33   Figure 1-5 The electronic effects of substitution at the ancillary ligands on luminescence colour.13,32,33   12  1.3.4 Luminescent Ruthenium(II) Complexes Over the last 50 years, ruthenium-based coordination compounds have garnered a great deal of research interest due to their desirable photophysical properties. Ruthenium complexes have been synthesized for numerous applications including: the study of light-induced electron and energy transfer (such as artificial photosynthesis);34,35 bioimaging;36 photodynamic therapy (PDT);37 and as sensitizers for dye-sensitized solar cells (DSSCs).38 Ru(II) polypyridyl complexes, in particular [Ru(bpy)3]2+ (7, Figure 1-6(a)) and its many derivatives, have been the focus of many photophysical studies,38-40 due to their strong absorption in the visible spectrum, high luminescence, and only small changes in geometry following excitation. [Ru(bpy)3]2+ undergoes the same transitions as discussed in Section 1.3.1 and highlighted in Figure 1-2(b). A 1MLCT state is formed when the complex is irradiated with light at λmax = 452 nm, with the photoexcited electron localized on one of the bpy ligands.41 Using ultrafast time-resolved absorption anisotropy measurements, McCusker et al. reported that the excited electron is initially delocalized between all three ligands, but quickly localizes to only one of the ligands in ~60 fs in acetonitrile solution.42 Following excitation to the 1MLCT, a fast (τ ≈ 100 fs)  ISC to a 3MLCT excited-state occurs with a near-unity quantum yield of formation.43 Deactivation of this excited-state can then occur radiatively with the emission of a photon, or through a non-radiative pathway.40 This deactivation shows both solvent44 and temperature dependence.39 In H2O the emission quantum yield is 0.04 with an excited-state lifetime of 0.6 µs, but the lifetime increases to 5 µs at 77 K. The temperature dependence is due to a thermal population of a MC excited-state that then proceeds to decay non-radiatively. The excited-state lifetime is therefore a sum of the radiative rate of decay (kr), non-radiative decay (knr) and the rate   13 of the population of the MC state (kMC). A simplified Jablonski diagram illustrating this process is shown in Figure 1-6(b).   Figure 1-6 (a) Structure of [Ru(bpy)3]2+ (7). (b) Simplified Jablonski diagram illustrating the excited-state processes in [Ru(bpy)3]2+ (7). Adapted from references.39,40   The excited-state energy of Ru(II)-polypyridyl complexes can be altered by varying the electronics at the ligands. Substituents that affect the electron density utilizing EDGs, EWGs or altering the degree of conjugation can tune the π* acceptor levels of the ligands and give systematic alterations to both the redox and photophysical properties of the complexes.45    14  Figure 1-7 Donor–acceptor–chromophore triads based on [Ru(bpy)3]2+ by Meyer et al.46,47  This highly tuneable nature, in addition to well-established excited-state properties has led ruthenium complexes to become integral in the study of charge-separated (CS) states. The use of donor–chromophore–acceptor (D–C–A) triads based on [Ru(bpy)3]2+ triads has been instrumental in the understanding of artificial photosynthesis.48,49 Meyer et al. have shown that D–C–A triads (8 and 9) can undergo a photoinduced, intramolecular electron transfer and generate CS states (Figure 1-7).46,47 The transfer of electrons involved is qualitatively analogous to the electron cascade in Photosystem II, the first protein complex involved in the light-dependent reactions of oxygenic photosynthesis.   1.3.5 Photophysics of Cu(I) d10 Metal Complexes While iridium and ruthenium complexes make excellent candidates for the study of fundamental photophysical processes, they are both precious metals and therefore more abundant   15 alternatives are sought. Thanks to the pioneering work by the McMillin group in the 1970s,50 copper(I) complexes have emerged as promising candidates due to the ability to finely-tune their photophysical properties, in addition to a greater abundance, lower toxicity and lower cost.51 The potential for Cu(I) complexes in commercial applications is high, such as use in electroluminescent materials like OLEDs52 and LEECs,53 as temperature sensors54 or as photocatalysts.55   Figure 1-8 Tetrahedral geometry of [Cu(N^N)2]+ species.  [Cu(N^N)2]+ complexes (N^N = diimine ligand) have been seen as an alternative to extensively researched [Ru(bpy)3]2+ (10) due to MLCT absorptions bands located in the visible region, in addition to emission in the red region of the spectrum. Three MLCT bands are seen in the visible spectral region.56 Band I is seen as a shoulder above 500 nm (S1 ß S0), the most prominent band, Band II, has a maximum between 430–480 nm (S2 ß S0) and Band III around 390–420 nm, though it is often hidden by the onset of Band II. The summation of these bands defines the shape of the visible absorptions and the spectral intensities are related to the symmetry of the complex.57 Typically the coordination number about the copper atom is lower than the d6 complexes discussed earlier in this chapter with a pseudo-tetrahedral ligating environment (10, Figure 1-8), though a range of distortions can exist in the ground-state from nearly-tetrahedral (D2d) symmetry to flattened-tetrahedral (D2) depending on the steric constraints imposed by the   16 ligands.58 An additional distortion is also attained in the excited-state due to changes in electronic distribution upon irradiation.   Figure 1-9  Simplified excited-state pathways of Cu(I) diimine complexes. After undergoing excitation, a Jahn-Teller distortion occurs, leaving an open site available for exciplex formation, following which the complex can relax radiatively or non-radiatively. Adapted from references.57,58  The excited-state dynamics of Cu(I) diimine complexes are illustrated in Figure 1-9. Upon photoexcitation one electron is transferred from the Cu(I) center to one of the diimine ligands, oxidizing the metal to a transient Cu(II) species and reducing the ligand to a radical-anion species.59 The excitation process alters the complex from a closed shell 3d10 configuration in the ground-state to an open shell 3d9 configuration in the excited-state. The molecule then undergoes a Jahn–Teller distortion whereby a fast structural rearrangement, in which the pseudo-tetrahedral geometry observed in the ground-state is flattened, decreasing the angle between the two ligand   17 planes.60,61 Singlet–triplet ISC then occurs on a picosecond timescale to give the 3MLCT state. In this flattened geometry, an additional axial coordination site is opened up that is available for nucleophilic attack by solvent, counteranion, or Lewis bases present in solution, resulting in a 5-coordinate structure. Increased coordination number at the copper center is known to quench the excited-state. McMillin et al. have shown in a number of reports that the addition of a Lewis base yields a 5-coordinate exciplex (excited-state complex)59,62-64 where the excited-state is stabilized, thus leading to a decrease in the energy gap between the ground- and excited-states which in turn promotes non-radiative decay and reduces the emissivity of the complex.   Figure 1-10 Increasing the steric bulk in the 2,9-positions of phen restricts the Jahn-Teller distortion at the Cu(I) center in the excited-state and increases emission intensity.65  The geometry of [Cu(N^N)2]+ can be altered by modulating the sterical hindrance at the ligand, which is often a variation of 1,10-phenanthroline (phen, 11). Altering the bulk at the 2,9-positions of the phen can limit the degree of distortion in the excited-state and thus decrease the   18 exciplex formation and increase the emission intensity. McMillin and co-workers found that the addition of alkyl groups at these positions restricted the Jahn-Teller distortion and increase the photoluminescence quantum yield (PLQY) and emission lifetimes (Figure 1-10).65 [Cu(phen)2]+ (11) shows virtually no emission in CH2Cl2 solution,66 but the series [Cu(dmp)2]+ (12, dmp = 2,9-dimethyl-1,10-phenanthroline), [Cu(dbp)2]+ (13, dbp = 2,9-dibutyl-1,10-phenanthroline) and [Cu(dop)2]+ (14, dop = 2,9-dioctyl-1,10-phenanthroline) have PLQYs of 4 ´ 10–4, 9 ´ 10–4 and 10 ´ 10–4, and emission lifetimes of 90, 150 and 155 ns, respectively. Introducing even more steric bulk with branching chains further inhibits the Jahn-Teller distortion. For [Cu(dnpp)2]+ (15, dnpp = 2,9-dineopentyl-1,10-phenanthroline) and [Cu(dsbp)2]+ (16, dsbp = 2,9-di(sec-butyl)-1,10-phenanthroline) PLQYs and emission lifetimes were increased to 16 ´ 104 and 260 ns, and 45 ´ 104 and 400 ns, respectively.   Figure 1-11 A sterically congested Cu(I) photosensitizer used for hydrogen photocatalysis by Castellano et al.67    In 2013, Castellano et al. designed a Cu(I) photosensitizer for use as a hydrogen evolution catalyst (Figure 1-11).67 [Cu(dsbtmp)2][PF6] (17, dsbtmp = 2,9-di(sec-butyl)-3,4,7,8-tetramethyl-  19 1,10-phenanthroline) was used together with the Co(dmgH)2(py)Cl (18) water reduction catalyst and a N,N-dimethyl-p-toluidine (19, DMT) sacrificial donor in 1:1 H2O:CH3CN, and the photosensitizer was found to be stable after 5 days of visible light irradiation, due to the inhibited Jahn-Teller distortion about the copper center. Turnover number (TON) was 35 (H2/Co) on average and turnover frequency (TOF) was 5 h–1, though it was found that the catalyst, not the photosensitizer limited the performance as with other systems utilizing a Co-dmg platform.  1.3.6 Thermally Activated Delayed Fluorescence (TADF) In an OLED device the excitons generated are of 25% singlet character and 75% triplet character. In order to have high efficiencies it is key that all singlet and triplet excited-states generated are harvested and utilized for luminescence.68 Third row transition metals, such as Ir(III), are well suited for these applications due to the efficient SOC that they possess between the lowest triplet state and higher-lying singlet states, as discussed in Section 1.2. These triplet emitters, however, are high-cost precious metals which may not be industrially scalable and so cheaper, more abundant alternatives are sought.69 As mentioned above, copper(I) species are possible alternatives for electroluminescent materials, but initially it appears as if they were not very suitable candidates. As a first row transition metal, copper suffers from a much weaker SOC and thus transitions between the excited triplet state and singlet ground state are principally forbidden. This leads to long phosphorescence decay times (>100 µs)10 which in a device can lead to strong saturation effects and side reactions which reduce the device stability,70 in addition to the structural non-radiative deactivation of the excited-state discussed in Section 1.3.5. However, these deficiencies may be overcome by the thermally activated delayed fluorescence (TADF) properties that Cu(I) complexes can possess.   20  Figure 1-12 Jablonski diagram illustrating the process of thermally activated delayed fluorescence (TADF). Adapted from reference.71  At its heart, TADF is a process in which otherwise “wasted” triplets can be harvested thermally to give a delayed fluorescence, thus increasing the efficiency of a system. A simplified Jablonski diagram illustrating the photophysical processes of TADF is given in Figure 1-12. First, one can imagine the system at a very low temperature, such as 4 K. Following irradiation with light, the Cu(I) complex is excited to a higher energy singlet state which then very quickly undergoes a fast IC to the S1 excited-state. An efficient ISC (S1 à T1) then occurs with relatively short lifetimes of ~10 ps,61,72 followed by a bright, long-lived phosphorescence (T1 à S0). Consequently, a prompt fluorescence (S1 à S0) is not seen.71,73 As the temperature is increased, the S1 state can be re-populated depending on both the thermal energy available and the energy separation between the lowest energy singlet and triplet states, ∆E(S1–T1). The re-population occurs via a reverse intersystem crossing (RISC) and while TADF has been detected with a ∆E(S1–T1) of ≈ 0.43 eV in the red fluorescent dye eosin,74 for practical applications this value should not exceed 0.12 eV.71 Following the RISC, a strong S1-  21 based fluorescence is observed with a long lifetime due to the population and depopulation of the T1 state.  In order to design a successful TADF emitter, there are three parameters that must be addressed:71 1. ∆E(S1–T1): the energy gap between the lowest energy singlet and triplet states should be as small as possible. Cu(I) complexes are well suited as they possess MLCT states resulting in discrete charge separation between excited and non-excited electrons causing a small singlet-triplet splitting. 2. τ(S1): the fluorescence decay time τ(S1) should be as short as possible. Unfortunately, a short τ(S1) is often related to a large ∆E(S1–T1). 3. τ(T1): the phosphorescence decay time τ(T1) should also be as short as possible in order to utilize an effective decay pathway in addition to the TADF.  1.3.7 Copper(I) TADF Emitters Heteroleptic copper(I) complexes bearing diphosphine and diimine ligands, of the structure [Cu(P^P)(N^N)]n (where P^P = diphosphine and n = 0 or 1), have been shown to be effective TADF emitters. As discussed in Section 1.3.5, Cu(I) complexes undergo a significant geometric reorganization upon excitation and that with a judicious use of sterically demanding ligands the distortion can be minimized, reducing the excited-state quenching and thus increasing the efficiency of the radiative decay. This also holds true for the design of copper(I) TADF emitters.    22  Figure 1-13 Neutral Cu(P^P)(N^N) emitters reported by Yersin et al.75  In 2011, Czerwieniec, Yu and Yersin reported strongly blue/white luminescent neutral Cu(I) complexes utilizing POP (bis(2-(diphenylphospanyl)phenyl) ether) as the diphosphine and bis(pyrazol-1-yl)borate ligands as the diimine ligands (Figure 1-13).75 At room temperature the emission of the complexes was nearly completely  fluorescence from the S1 à S0 transition with the S1 state thermally populated from the T1 state, as determined from emission lifetime measurements. In CH2Cl2 solutions, Cu(POP)(pz2BH2) (20) and Cu(POP)(pz2Bph2) (21) (where pz2BH2 = bis(pyrazol-1-yl)borohydrate and pz2Bph2 = bis-(pyrazol-1-yl)biphenylborate) showed emission maxima at 525 nm and 498 nm, and PLQYs of 0.09 and 0.08, respectively. The low emission quantum yields are attributed to the greater rearrangement of molecular geometry afforded in solution. As a neat crystalline powder, however, where molecular distortion is disfavoured, λmax values of 436 and 464 nm were seen in addition to much higher PLQYs of 0.45 and 0.90, respectively.   Introducing steric constraints, in a similar vein to the work by McMillin65 highlighted in Figure 1-14, can aid the emissivity of a Cu(I) TADF  luminophore. In 2014, Linfoot, Robertson and Yersin reported a Cu(I) system where adding steric bulk to a dmbpy (4,4′-dimethyl-2,2′-biyridyl) diimine ligand in complex 22 to give a tmbpy (4,4′,6,6′-tetramethyl-2,2′-biyridyl) diimine   23 ligand in 23 not only increased the PLQY in the solid state from 0.09 to 0.74, but also in EtOH solutions from <0.01 to 0.06.   Figure 1-14 Adding steric bulk can decrease the degree of geometric distortion in the excited-state and can increase the emission intensity in Cu(I) TADF emitters in solution, as reported by Linfoot, Robertson and Yersin.76 Image reproduced with permission from the American Chemical Society, Copyright © 2014.  1.4 Sulfur Oxidation States Alter Photophysical Processes In recent years, our group has shown that electronic coupling between two identical molecules can be controlled by altering the oxidation state at a sulfur-bridge linking the two molecules. In 2013, we reported a system in which two identical chromophores bridged by a sulfur atom can undergo a systematic increase in PLQY by oxidizing the sulfur from sulfide (S), to sulfoxide (SO) and to sulfone (SO2) (Figure 1-15).77 In the case of bridged terthiophene dimer T3S, oxidizing to sulfone compound T3SO2 increased the emission intensity by an order of magnitude. In addition, a solvent dependence of the PL spectra, whereby a bathochromic shift was seen in more polar solvents, hinted towards an increase in CT character.    24  Figure 1-15 Structures of sulfur-bridged dimers synthesized by our group77 showing (a) increased PLQY as the oxidation state at sulfur is increased. (b) Emission spectra in CH2Cl2 solutions for T3SOn in comparison to the parent arene, terthiophene (T3). Image reproduced with permission from the American Chemical Society, Copyright © 2013.    Further study of the T3 system by our group, in conjunction with the Bardeen group at University of California Riverside, elucidated the cause of this phenomenon through the use of femtosecond transient absorption in tandem with computational calculations.78 In changing the oxidation state of the bridging sulfur, the electronic coupling of the chromophoric system is tuned. Electron lone pairs in the S and SO bridges of T3S and T3SO screen the interaction between the π electrons of the two terthiophene chromophores. Full oxidation to SO2 in T3SO, however, leads to polarized bonds that decrease the screening, thus stabilizing the CT contributions. We have   25 further demonstrated this oxidation-dependent intramolecular charge transfer in anthracene-containing donor–acceptor–donor systems79 and oligothiophenes.80  1.5 Goals of Thesis Photo-active organometallic complexes have applications as emissive materials and in photocatalysis. By modifying the electronics of the ligands using substituent control one can affect the HOMO and LUMO energies and alter the excited-state properties, such as luminescence colour. The primary focus of this thesis lies on the synthesis and spectroscopic characterization of organometallic complexes bearing diimine sulfur-bridged ligands. In addition to substituent control, the sulfur center provides a convenient approach to adjust the electron density with chemical oxidation. By varying the oxidation state of the sulfur atom (S, SO, SO2) the excited-state electronics of the complex can be altered, adding a degree of “two-level tuning” that can be achieved with a number of inexpensive, easy-to-handle oxidants. In Chapter 2 a series of new sulfur-bridged ligands is introduced. These ligands are analogous to 2,2′-bipyridine with the two pyridyl rings bound in the 2,2′-positions by a sulfur atom. Both the pyridyl and 4-methylpyridyl derivatives are prepared as sulfide, sulfoxide and sulfone derivatives. The spectroscopic properties of these compounds are investigated, both as pro-ligands and as ligands in a series of bis-cyclometallated Ir(III) complexes. The effect of these ligands on the excited-state is quantified using a combination of ground-state and excited-state spectroscopic techniques, in addition to computational calculations. It is found that complexes featuring the S- and SO-bridged ligands emit from a 3LC state, but, on ligand oxidation to SO2, the emission is   26 dominated by a 3LLCT/3MLCT state, therefore demonstrating an oxidation state-sensitive luminescence. Chapter 3 highlights a series of heteroleptic Cu(I) complexes ligated by sulfur-bridged dipyridyl ligands of S and SO2 oxidation states with steric bulk in the 6,6′-positions, designed as possible TADF emitters. A combination of increased diimine bite angle and additional coordination sites at the S=O oxygen afforded interesting bimetallic species when bulkier substituents were used, with excited-state properties probed at ambient temperature and at 77 K, including time-resolved spectroscopy. Chapter 4 discusses Cu(I) sulfoxide species which illustrate how very subtle changes in ligand design can greatly impact the excited-state properties. A reversible, solid state thermochromic emission is observed in one of the complexes, in addition to phase-specific luminescence as a powder. A correlation between the structural properties (via crystallographic data) and the temperature-dependent excited-state behaviour is discussed. The modulation of electron transfer in a bimetallic system is addressed in Chapter 5. Ru(II)-Re(I) supramolecular photocatalysts can be used in the reduction of CO2 to more useful products, and the nature of the electron transfer from the Ru(II) photosensitizer to the Re(I) catalyst is imperative to the process. Here the two metallic centers are bound via a sulfur-linked bisphenanthroline bridge, with the photocatalytic properties of the S and SO2 oxidation states explored.    27 CHAPTER 2: Tunable Emission of Iridium(III) Complexes Bearing Sulfur-Bridged Dipyridyl Ligands  2.1  Introduction Since the discovery of the incandescent carbon filament lamp by Thomas Edison in 1879, artificial light has had a great impact on the human race, releasing us from the light-dark cycle our ancestors were bound by. The global electricity consumption for grid-based lighting sits at roughly 19%81 and thus there has been a recent push to increase the efficiency of modern-day lighting as efficiency gains in this sector would greatly contribute to decreasing worldwide energy use.82 Over the past two decades solid state lighting (SSL) has emerged as the source for higher efficiency light generation with reduced heat output.83 Widely used SSL is based on inorganic light-emitting diodes (LEDs), with organic light-emitting diodes (OLEDs)84,85 and, more recently, light-emitting electrochemical cells (LECs)86,87 showing promise for the future.  Iridium(III) complexes are known for their high efficiency and wide scope of tunability in both OLEDs27,88-90 and LECs,91-94 yet there remains a lack of wide band-gap blue emitters for both these purposes.27 Blue emission typically requires the addition of fluorine substituents on the cyclometallated C^N ligands, but these groups are often unstable under bias, leading to degradation and lower device longevity.95 In [Ir(C^N)2(N^N)]+ complexes the HOMO contains contributions from the Ir(III) dπ orbitals and the phenyl π-orbitals of the C^N cyclometallating ligand, and the LUMO is usually localized on the N^N ancillary ligands.94 This localized nature leads to effective colour tuning of the HOMO–LUMO gap. Adding electron-donating substituents   28 to the N^N ligand destabilizes the LUMO and blue-shifts the emission. Additionally, altering the N^N ligands can help stabilize the HOMO to increase the HOMO–LUMO gap even further. Our group has previously reported that the degree of sulfur oxidation in sulfur-bridged conjugated organic molecules has significant effects on the electronic properties of these systems.77,78 It has been shown that by increasing the oxidation state of a bridging sulfur in a thiophene-based system from sulfide (S), to sulfoxide (SO), to sulfone (SO2) it is possible to influence intramolecular charge transfer and emission intensity. In this body of work it is shown that the emission colour of an Ir(III) complex can be tuned through altering the oxidation state at the sulfur of the N^N ancillary ligands.  This approach is advantageous as only one parent ligand is needed, with simple oxidation yielding the other two variants, considerably simplifying synthesis.  2.2 Ligand Synthesis and Characterization The symmetrical sulfide-bridged pro-ligands DPS and 4,4'-Me-DPS were prepared via a nucleophilic aromatic substitution reaction using thiourea,96 starting with either 2-bromopyridine or 2-chloro-4-(methyl)pyridine (Scheme 2-1). The desired sulfoxide and sulfone compounds were then obtained by oxidation of the corresponding diaryl sulfides. Oxidation to the sulfoxide products was performed using one of two different pathways:97,98 Compound DPSO was prepared through the addition of 30% H2O2 to DPS dissolved in glacial acetic acid, while 4,4'-Me-DPSO was formed by the addition of m-CPBA to 4,4'-Me-DPS at 0 ºC. Sulfone pro-ligands DPSO2 and 4,4'-Me-DPSO2 were synthesized through oxidation of the appropriate diaryl sulfide with 30% H2O2 using a NbC catalyst.99 Moderate to high yields of the desired products were obtained in all synthetic steps and the six pro-ligands were characterized using NMR spectroscopy, mass   29 spectrometry and infrared spectroscopy. 1H and 13C signals were assigned using COSY, NOESY, HSQC and HMBC experiments.  Scheme 2-1 Synthesis of Iridium(III) Complexes Ir-DPSOn and Ir-4,4'-Me-DPSOn (where n = 0, 1 or 2).   2.3 Synthesis of [Ir(ppy)2(N^N)][PF6] Complexes The [Ir(ppy)2Cl]2 dimer was prepared by a microwave-assisted reaction of IrCl3·xH2O with 2-phenylpyridine.100 The dimer was isolated as a yellow solid in high yield, requiring no further purification. Complexes [Ir(ppy)2(N^N)][PF6] (N^N = DPSOn or 4,4'-Me-DPSOn, where n = 0, 1   30 or 2) were synthesized via an intermediate solvento complex, formed by the reaction of the iridium dimer with AgPF6 in MeOH, followed by the addition of the appropriate N^N ligand yielding the desired compounds as yellow powders in all cases (Scheme 2-1).11 The HR-ESI mass spectra show a product peak corresponding to [M−PF6]+ exhibiting a characteristic iridium isotope pattern.   2.3.1 Nuclear Magnetic Resonance (NMR)  Figure 2-1 Expanded NMR spectrum of Ir-DPSO in CD2Cl2 at 25 ºC (850 MHz).  1H and 13C NMR spectra of the Ir(III) complexes were assigned using COSY, NOESY, HSQC and HMBC experiments. The sulfoxide component of complexes Ir-DPSO and Ir-4,4'-Me-DPSO causes the two cyclometallating ligands, and thus the two pyridyl rings of the ancillary ligand, to become inequivalent.  This is due to the proximity of the sulfoxide oxygen to one of the ppy ligands, changing its environment relative to the other ppy. The NMR spectrum for Ir-DPSO   31 exhibiting this inequivalence is shown in Figure 2-1. The two different phenylpyridine ligands were distinguished through NOESY cross peaks. Due to the presence of a chiral sulfur center and a chiral metal center it would be expected to see a mixture of diastereomers in the NMR spectrum, however only one species is seen. It is postulated that hydrogen bonding between the oxygen atom and H8 could be providing a stabilizing interaction that preferentially forms one stereoisomer. The room temperature NMR spectra of all six complexes have clear, defined peaks with no evidence for exchange or dissociation occurring in CD2Cl2; however, in CD3CN, a stronger coordinating NMR solvent, it was observed for complexes Ir-DPSO, Ir-DPSO2 and Ir-4,4'-Me-DPSO, Ir-4,4'-Me-DPSO2 that the ancillary ligand is replaced by solvent, resulting in Ir(ppy)2(CD3CN)2][PF6]. It is known that upon increasing the oxidation state at the sulfur from sulfide to sulfoxide to sulfone, the electron withdrawing strength of the substituent increases.101 It is likely that the increased electron-withdrawing character reduces the σ-donating ability of the pyridyl moieties. This may also explain the lower isolated yields found in the complexes with ligands of higher sulfur oxidation state. Complexes Ir-DPSO and Ir-4,4'-Me-DPSO both have the possibility of N,N coordination through both pyridyl moieties and N,O binding through one pyridyl group and the oxygen of the sulfoxide. NMR experiments show that in both cases the ancillary ligand is bound N,N as there are no NOE cross peaks observed between protons H4 and H34, the presence of which would indicate free rotation of a pendant pyridyl ring.  2.3.2 Single Crystal X-Ray Diffraction Single crystals of Ir-DPS were grown by vapour diffusion of hexanes into a CH2Cl2 solution of the complex. Ir-DPS crystallizes in the space group P-1 with the asymmetric unit   32 containing two complex molecules and their respective counter-anions (Figure 2-2). The hexane solvent molecule contained within the structure is disordered and has been removed using the SQUEEZE method.102   Figure 2-2 Crystal structure of Ir-DPS (Λ conformer). Ellipsoids plotted at 50% probability and H atoms and hexane molecules removed for clarity.  As shown in Figure 2-2, coordination occurs through the pyridyl nitrogen atoms of the diimine ligand (DPS), each adopting a trans configuration with respect to the corresponding coordinating carbon atom of the cyclometallating phenyl ring. Bond angles and lengths discussed in this section are in respect to the Λ conformation of Ir-DPS. The complex exists in an octahedral geometry with both ppy cyclometallating ligands showing deviation from planarity between the rings (angles between ring planes are 5.8 and 13.15º respectively). The tetrahedral geometry about the sulfur atom results in non-planarity of the DPS ligand pyridyl rings with an angle of 40.3º between each ring plane and results in a N1–Ir–N2 bite angle of 89.99º. This allows the complex   33 to adopt a more perfect octahedron when compared to [Ir(ppy)2(bpy)][PF6]13 which has a smaller bite angle for the bpy ligand of 76.20º. The bond Ir–N bond lengths are comparable between the two species, with Ir-DPS having slightly longer bond lengths (Ir–N1, Ir–N2 = 2.198, 2.173 Å, respectively) than [Ir(ppy)2(bpy)][PF6] (Ir–N1, Ir–N2 = 2.129, 2.136 Å respectively), indicating a slight reduction in σ-donation from the ancillary ligand in Ir-DPS. Complexes containing ligand DPS have previously been synthesized, with crystal structures reported for species containing ruthenium,103-105 rhodium106 and platinum and palladium.107,108 In the case of [Ru(DPS)2(N,S-DPS)][PF6],105 a homoleptic Ru(II) complex chelated with ligand DPS, two ligands bind N,N while the third binds N,S resulting in a pendant pyridyl ring. This is not observed in Ir-DPS, likely due to the cyclometalling ppy ligands having less bulk than DPS.  2.4 Photophysical Properties 2.4.1 Absorption Spectroscopy  Figure 2-3 Absorption spectra of (a) Pro-ligands DPSOn and 4,4'-Me-DPSOn and (b) complexes Ir-DPSOn and Ir-4,4'-Me-DPSOn. All spectra were recorded in ~ 2 ´ 10–5 M CH2Cl2 solutions at room temperature.   34 The absorption spectra of sulfide, sulfoxide and sulfone pro-ligands DPSOn and 4,4'-Me-DPSOn (n = 0, 1 or 2) are shown in Figure 2-3(a). Features at higher energies than 325 nm in all pro-ligands correspond to π–π* transitions. The spectra of the sulfide compounds DPS and 4,4'-Me-DPS contain a weak low energy feature at 375 nm. Absorption spectra of all six Ir(III) complexes were recorded in CH2Cl2 with concentrations of ~2 × 10−5 M (Figure 2-3(b)). The spectra are similar with all complexes showing intense absorption bands lying in the UV region with maxima in the range 250–270 nm, which are assigned to spin-allowed π–π* transitions of the ligands. Lower intensity bands are observed between 350 and 450 nm which correspond to spin-allowed metal-to-ligand (1MLCT) and ligand-to-ligand (1LLCT) transitions. Weaker intensity tails above 450 nm comprise spin-forbidden 3MLCT, 3LLCT and ligand centered (3LC) transitions.94  2.4.2 Photoluminescence Spectroscopy  Figure 2-4 Photoluminescence spectra of Ir(III) complexes (λex = 390 nm). All spectra were recorded in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature and sparged with Ar for 25 min.    35 The normalized photoluminescence spectra of all six Ir(III) complexes were recorded in CH2Cl2 solution after deaerating with Ar for 25 min and are shown in Figure 2-4(c). Emission maxima are summarized in Table 2-1. Emission of Ir(III) complexes with the general structure [Ir(N^N)2(N^N)][PF6] occurs from the lowest-lying triplet state (T1), which due to the heavy atom effect, typically contains a mixture of contributions from 3MLCT, 3LLCT and 3LC states.109 In general, when the contribution is higher in charge-transfer character, the emission profile becomes broader and less structured.   Table 2-1 Photophysical data of Ir(III) complexes in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature.  λem [nm] [a] τem [μs] [b] PLQY [a] kr [106 s–1] knr [106 s–1] Ir-DPS 478, 510, 548 0.360 0.04 0.10 2.7 Ir-DPSO 478, 510, 548 0.163 (3.6%), 0.024 (93.2%), 0.001 (3.2%) <0.01 0.25 [c] 41.4 [c] Ir-DPSO2 552 0.231 0.03 0.14 4.2 Ir-4,4'-Me-DPS 478, 510, 548 0.828 0.08 0.10 1.1 Ir-4,4'-Me-DPSO 478, 510, 548 0.898 (13.2%) , 0.060 (83.9%), 0.044 (2.9%) 0.03 0.43 [c] 16.2 [c] Ir-4,4'-Me-DPSO2 537 0.065 0.04 0.66 14.7 [a] λex = 390 nm. [b] λex = 370 nm. [c] kr and knr for sulfoxide complexes calculated using the major lifetime component.  Sulfide (Ir-DPS, Ir-4,4'-Me-DPS) and sulfoxide (Ir-DPSO, Ir-4,4'-Me-DPSO) complexes show identical, fine-structured luminescence profiles with maxima at 478, 510 and 548   36 nm, indicating a large 3LC character of the emissive state. This emission is independent of excitation wavelength. Solvents with higher polarity can often stabilize charge transfer states resulting in broadened, red-shifted emission, hence photo-luminescence spectra of Ir-DPS were obtained in solvents of increasing polarity (CH2Cl2, MeOH and CH3CN) to probe the nature of the emission.  No change in emission profile was observed in these experiments with the fine structure intact, supporting the assignment of an emitting 3LC state of π–π* character, localized on the ppy cyclometallating ligands. This shape also bears a distinct resemblance to the emission band of [Ir(ppy)2(CO)(Cl)]110, further suggesting a dominant 3LC contribution to the emissive triplet state. In contrast, sulfone complexes Ir-DPSO2 and Ir-4,4'-Me-DPSO2 show broad, featureless photoluminescence spectra indicative of 3MLCT character with both species exhibiting bathochromically shifted emission bands compared to their sulfide and sulfoxide counterparts. A small blue-shift of 15 nm is observed on going from Ir-DPSO2 (552 nm) to Ir-4,4'-Me-DPSO2 (537 nm) which is attributed to the methyl substituents in the 4- and 4′-positions of the ancillary ligand in Ir-4,4'-Me-DPSO2, resulting in destabilization of the LUMO due to the higher electron density on the N^N ligand.111  2.4.3 Photoluminescence Quantum Yields (PLQY) Photoluminescence quantum yields (PLQY, Table 2-1) in deaerated CH2Cl2 are in the range ~0.01–0.08, slightly lower than that of [Ir(ppy)2(bpy)][PF6] in deaerated CH3CN (0.14).112 These low values are attributed to the increased ligand flexibility of the ancillary ligands about the sulfur atom. The excited-state energy can presumably be dissipated more readily through ligand motion when compared to more rigid ancillary ligands such as 2,2'-bipyridine. The low PLQYs for sulfoxide-containing complexes Ir-DPSO and Ir-4,4'-Me-DPSO (<0.01 and 0.03,   37 respectively) could be due to a deactivating pathway opened up by excited-state sulfoxide inversion. The general trend of higher PLQY for the methyl-functionalized complexes Ir-4,4'-Me-DPSOn relative to Ir-DPSOn is attributed to the increased σ-donation caused by this methyl substitution destabilizing a short-lived, non-emissive triplet metal-centered state (3MC) and increases the energy separation between this 3MC state and the emissive 3LC state based on the cyclometalling ligands.113,114  2.4.4 Photoluminescence Lifetimes  Figure 2-5 Photoluminescence lifetimes of Ir(III) complexes (λex = 370 nm). All spectra were recorded in ~2 ´ 10–5 M CH2Cl2 solutions at room temperature and sparged with Ar for 25 min.  Excited-state lifetimes (τ) in deaerated CH2Cl2 are less than 1 µs for all complexes (Figure 2-5 and Table 2-1). The sulfone complexes Ir-DPSO2 and Ir-4,4'-Me-DPSO2 each display a mono-exponential decay. The lifetimes are shorter than those of sulfide complexes Ir-DPS and  Ir-4,4'-Me-DPS, consistent with an emission attributed to 3MLCT/3LLCT states rather than from a 3LC state.115-117 This is further supported as a lower 3LC character is indicated by a higher kr.94,117   38 Both sulfoxide complexes Ir-DPSO and Ir-4,4'-Me-DPSO display a tri-exponential decay with the major component of each emission residing in the 10’s of nanoseconds range.  The presence of the sulfoxide group breaks the symmetry in the complex (see 2.3.1) and it is thought that this break in symmetry causes the multi-exponential emission lifetimes.  2.5 Electrochemical Properties Cyclic voltammetric data for [Ir(N^N)(N^N)]+ (N^N = DPSOn or 4,4'-Me-DPSOn, where n = 0, 1 or 2) are visualized in Figure 2-6 and summarized in Table 2-2. All complexes show irreversible oxidation processes. For both series, upon increasing sulfur oxidation state from sulfide through sulfone, the first oxidation of the complex is shifted to higher potential. This is attributed to the electron-withdrawing nature of the sulfoxide and sulfone groups reducing the electron density on the iridium center, thereby stabilizing the HOMO. Additional peaks may be due to oxidation at the sulfur center on the ancillary ligand.    39  Figure 2-6 Cyclic voltammograms of Ir(III) in dry, N2-sparged CH2Cl2 at room temperature. (a) Ir-DPSOn complexes. (b) Ir-4,4'-Me-DPSOn complexes.  As the sulfur oxidation state is increased, the first reduction occurs at a more positive potential when measured against the sulfide complexes (–1.26 V, –1.73 V for Ir-DPSO2 and Ir-4,4'-Me-DPSO2 respectively compared to –2.16 V, –2.07 V for Ir-DPS and Ir-4,4'-Me-DPS respectively), consistent with stabilization of the LUMO. Additional reductive peaks seen in Ir-DPSO, Ir-DPSO2, Ir-4,4'-Me-DPSO and Ir-4,4'-Me-DPSO2 are attributed to reduction of the respective sulfoxide and sulfone gap. The smaller electrochemical ∆E for sulfone complexes Ir-DPSO2 and Ir-4,4'-Me-DPSO2 is consistent with the red-shift of the emission compared to the respective sulfide and sulfoxide complexes.    40 Table 2-2 Electrochemical data of Ir(III) complexes in deaerated CH2Cl2 solutions and referenced to Fc/Fc+. Complex 𝑬1/2 𝒐𝒙  [V]  𝑬1/2𝒓𝒆𝒅 [V]  ∆𝑬1/2	[V] [a] Ir-DPS 1.06 –2.16 3.22 Ir-DPSO 1.16 –1.59 2.75 Ir-DPSO2 1.17 –1.26 2.43 Ir-4,4'-Me-DPS 0.92 –2.40 3.32 Ir-4,4'-Me-DPSO 1.02 –2.16 3.18 Ir-4,4'-Me-DPSO2 1.13 –1.73 2.86 [a]  ∆𝐸 = 𝐸& '()* − 𝐸& '(,-.   2.6 Theoretical Calculations The energy and electron density contours were calculated for the HOMO and lowest unoccupied molecular orbitals (LUMO, LUMO+1 and LUMO+2) for all six Ir(III) complexes and are shown in Table 2-3, with frontier orbital diagrams illustrated in Appendix Chapter 6:Listing A.1Chapter 6:Listing A.2.  Table 2-3 Ground-state frontier orbital energies for Ir(III) complexes.  Energy [eV]  Ir-DPS Ir-DPSO Ir-DPSO2 Ir-4,4'-Me-DPS Ir-4,4'-Me-DPSO Ir-4,4'-Me-DPSO2 HOMO –5.90 –6.00 –6.05 –5.86 –5.95 –6.01 LUMO –2.07 –2.25 –2.70 –1.97 –2.11 –2.53 LUMO+1 –2.01 –2.14 –2.24 –1.94 –2.06 –2.11 LUMO+2 –1.99 –2.04 –2.08 –1.92 –1.94 –2.01    41 It was found that both increasing the oxidation state of sulfur and introducing methyl groups have a negligible impact on the HOMO energy and electron density distribution of the complexes. The HOMO is mainly located at the iridium and ppy ligands, with functionalization at the ancillary N^N ligand not contributing to this orbital. These trends are also consistent with the oxidation potentials of these complexes. The LUMOs of these complexes are located at the ancillary N^N ligand with a small contribution from iridium (except in Ir-4,4'-Me-DPS), and oxidation of the sulfur stabilizes the orbitals centered on the N^N ligand. In the case of the unsubstituted series, the LUMO energy of Ir-DPSO2 is 0.63 eV lower than Ir-DPS. The addition of methyl groups in the 4- and 4'- positions of the ancillary ligand destabilizes the LUMO by 0.1–0.17 eV. The LUMO, LUMO+1 and LUMO+2 of Ir-4,4'-Me-DPS all have similar energies and contributions from both the ppy ligands and the N^N ligand. The combination of these effects leads to HOMO–LUMO energy gaps that decrease on oxidation of sulfur, and increase with the introduction of methyl groups. These trends follow those seen in the electrochemical data, however they contradict the experimental emission data. To gain more insight into the emitting states, the geometry of the lowest-energy triplet excited state (T1) must be probed.  Table 2-4 The calculated adiabatic energy differences of Ir(III) complexes with regard to the experimental emission energies.  Ir-DPS Ir-DPSO Ir-DPSO2 Ir-4,4'-Me-DPS Ir-4,4'-Me-DPSO Ir-4,4'-Me-DPSO2 ΔE [eV] 2.66 2.65 2.58 2.66 2.65 2.66 Eem [eV] 2.35 2.34 2.24 2.35 2.34 2.28 λem calc. [nm] 528 530 554 528 530 543    42 To elucidate the emitting state, the geometry of the T1 state was optimized using the spin-unrestricted UB3LYP method. Table 2-4 summarizes the adiabatic energy difference between S0 and T1 and emission energy (Eem) calculated as the vertical energy difference between T1 and S0 at the optimized geometry of T1 (Figure 2-7(a)). The calculations find almost identical emission energies for Ir-DPS, Ir-DPSO, Ir-4,4'-Me-DPS and Ir-4,4'-Me-DPSO. This is in accordance with the experimental data that on oxidizing the ligands from sulfide to sulfoxide (in addition to the introduction of methyl groups in the case of Ir-4,4'-Me-DPS and Ir-4,4'-Me-DPSO) the emission energy is unaffected since the emitting state is centered on the ppy ligands. The calculations also correctly predict the red-shifted emission on full oxidation to sulfone, in addition to the blue-shifted emission of Ir-4,4'-Me-DPSO2 relative to Ir-DPSO2 after substitution of methyl groups in the 4- and 4'- positions of the ancillary ligand.   Figure 2-7 (a) Schematic diagram of the adiabatic energy difference (∆E) between S0 and T1, and the emission energy (Eem) from T1. (b) Unpaired-electron spin density contours for Ir-DPS and Ir-DPSO2.  The unpaired-electron spin density plots for T1 of each complex are shown in Appendix A.3. Unpaired-electron spin density plots comparing complexes Ir-DPS and Ir-DPSO2 are shown   43 in Figure 2-7(b). The sulfide and sulfoxide complexes exhibit similar spin density distributions which are mainly centered on one of the ppy ligands, with some contribution from the iridium center. This indicates that the T1 states of the sulfide and sulfoxide complexes are mainly ligand-centered, with a small contribution of MLCT character. The unpaired electron spin density of Ir-DPSO2 and Ir-4,4'-Me-DPSO2 is spread over the complex, with ~0.5 e– on the iridium, 0.5 e– on the ppy ligands and 1 e– on the ancillary ligand, which matches the topology of MLCT transitions, supporting the assigned MLCT emitting state of sulfone complexes.  To further understand the nature of the lowest-energy triplet state, time-dependent DFT (TD-DFT) calculations were employed for all six complexes at the optimized geometry of S0 (Appendix A.4). The first triplet states of Ir-DPS, Ir-DPSO, Ir-4,4'-Me-DPS and Ir-4,4'-Me-DPSO result from the mono-excitation of the HOMO to the first orbital centered on the ppy ligands; the LUMO for Ir-4,4'-Me-DPS, the LUMO+1 for Ir-4,4'-Me-DPSO and the LUMO+2 for Ir-DPS and Ir-DPSO. The first T1 states of Ir-DPSO2 and Ir-4,4'-Me-DPSO2 are mainly contributed from HOMO–LUMO excitation and have MLCT character. The addition of methyl groups to the ancillary ligand decreases the contribution of the MLCT transition. These results again demonstrate the LC emitting nature of the sulfide and sulfoxide complexes and the MLCT emitting nature of the sulfone complexes.  2.7 Conclusions Six new [Ir(N^N)(N^N)]+ cationic complexes have been synthesized consisting of 2-phenylpyridine cyclometallating ligands (N^N) and sulfur-bridged ancillary ligands (N^N). Pro-ligands DPSOn (n = 0, 1 or 2) contain two pyridyl rings connected in the 2,2'- positions by a sulfide (DPS; S), sulfoxide (DPSO; SO) and sulfone (DPSO2; SO2), and pro-ligands 4,4'-Me-DPSOn (n   44 = 0, 1 or 2)  retained a similar structure, except with methyl substituents in the 4 and 4'-positions. These ligands were designed to gradually red-shift the emission upon increasing sulfur oxidation state.  The complexes were characterized by spectroscopic and mass spectrometry techniques, along with the single crystal structure of Ir-DPS. The sulfide (Ir-DPS, Ir-4,4'-Me-DPS) and sulfoxide complexes (Ir-DPSO, Ir-4,4'-Me-DPSO) show blue-green emission with identical, fine-structured profiles (maxima at 478, 510 and 548 nm). The vibrational structure indicates significant 3LC contributions to the lowest lying triplet state, and the absence of a red-shifted broadening in solvents of increasing polarity confirms a lack of charge transfer nature. Sulfone complexes (Ir-DPSO2, Ir-4,4'-Me-DPSO2) exhibit broad, red-shifted emission bands, indicative of a pronounced 3MLCT/3LLCT state. Complex Ir-4,4'-Me-DPSO2, with methyl substituents in the 4,4'-positions, shows a slight blue-shift (15 nm) compared to Ir-DPSO2. PLQYs in deaerated CH2Cl2 were low (<0.01–0.08), slightly slower than that of [Ir(ppy)2(bpy)][PF6] (0.14) in deaerated CH3CN. Here we have shown a new “two-level” method to tuning the emission colour of Ir(III) complexes utilizing methyl- substitution of a sulfur-bridged ancillary ligand, coupled with varying oxidation state at the sulfur center. Increasing the substituents’ electron-donating ability at the 4,4'-positions of the sulfide-bridged ancillary ligand, coupled with the addition of electron-withdrawing groups at the cyclometallating 2-phenylpyridine ligands could result in a desired blue-shift of the emission colour.    45 2.8 Experimental Details 2.8.1 General All experiments were conducted in an inert atmosphere using standard Schlenk techniques unless otherwise stated. Solvents used were of reagent grade and used without any further purification. HPLC grade solvents were used for analyses. Thiourea was purchased from TCI and 2-bromopyridine and 2-chloro-4-(methyl)pyridine purchased from Sigma Aldrich. All compounds were used as received. [Ir(ppy)2Cl]2 was prepared as per the literature.100  2.8.2 Spectroscopy 1H NMR and 13C NMR spectra were collected using either a Bruker AV-400 or AV-850 spectrometer and referenced to the residual protonated solvent peak. NMR solvents (from either Cambridge Isotope Laboratories or Aldrich) were used as received. Electrospray ionization mass spectrometry data was obtained using a Bruker Esquire LC ion trap mass spectrometer. Infrared spectroscopy was performed on an attenuated total reflection (ATR) crystal using a Perkin-Elmer Frontier FT-IR spectrometer. Electronic absorption spectra were recorded on a Varian-Cary 5000 UV-Vis-near-IR spectrophotometer. Fluorescence data were collected on a Photon Technology International (PTI) QuantaMaster 50 fluorimeter fitted with an integrating sphere, double excitation monochrometer and utilizing a 75 W Xe arc lamp as the source. Emission lifetime data were collected using a Horiba Yvon Fluorocube TCSPC apparatus. A 370 nm NanoLED source pulsing at a repetition rate of 50–100 kHz was used for excitation. Broadband emission was monitored by a CCD detector at wavelengths >450 nm using a low pass filter. Data were fitted using the DAS6 Data Analysis software package. All measurements were recorded at room   46 temperature. Sample solutions were maintained under a blanket of Ar for the duration of the measurements in 1 cm2 quartz cells (Starna Cells) fitted with a rubber septum.  2.8.3 X-ray Crystallography Single crystal X-ray data was collected using a Bruker APEX DUO diffractometer with graphite monochromated Mo Kα radiation (λ = 0.71073 Å) at 90 K. Raw frame data were processed using APEX2.118 The program SAINT+, version 6.02118 was used to reduce the data and the program SADABS was used to make corrections to the empirical absorptions. Space group assignments were made using XPREP118 on all compounds. In all cases, the structures were solved in the WinGX Suite119 of programs by direct methods using SHELXS-97120 and refined using full-matrix least-squares/difference Fourier techniques on F using SHELXL-97120 All non-hydrogen atoms were refined anisotropically. Diagrams and publication material were generated using ORTEP-3,121 and PLATON.122  2.8.4 Electrochemistry Solution state electrochemical data was collected on a CHI660D potentiostat with a 3-electrode configuration using a glassy carbon working electrode, Ag/AgCl reference electrode and Pt wire counter electrode. 0.1 M [NBu4][PF6] in CH2Cl2 was degassed with N2 and used as the electrolyte. 0.3 mM solutions of each complex were prepared and a scan rate of 100 mV/s was applied in all cases. Potentials were referenced to Fc/Fc+ with the addition of ferrocene at the end of each experiment.     47 2.8.5 Computational Details Density functional theory (DFT) calculations were carried out using the Gaussian 09 Rev.D01 suite of programs. The B3LYP functional with 6-31G** basis set (for C, H, S and O atoms)123,124 and the LANL2DZ effective-core pseudopotential (for iridium)125 were employed to optimize the geometry of the ground states (S0) and lowest-energy triplet states (T1) of all complexes, with no symmetry restrictions. Phosphorescence energy is estimated as the energy difference the minimum of T1 and the energy of S0 at the T1 optimized geometry. All calculations were performed in the presence of the solvent (CH2Cl2). Solvent effects were considered within the self-consistent reaction field (SCRF) theory using the polarized continuum model (PCM) approach.126,127 Time-dependent DFT (TD-DFT) calculations were also performed to understand the nature of T1 states.128,129  2.8.6 Synthesis di(pyridin-2-yl)sulfane (DPS)  Synthesis of DPS was adapted from the literature.96 2-bromopyridine (0.5 mL, 0.154 mol, 1.0 eq.) and thiourea (5.86 g, 0.077 mol, 0.5 eq.) were placed under a N2 atmosphere and then dissolved in EtOH (500 mL). The resulting solution was heated at reflux for 60 h, cooled and the EtOH removed. The crude product was dissolved in CH2Cl2 (20 mL), washed with water (3 ´ 20 mL), brine (20 mL) and then dried using MgSO4. The solution was filtered and solvent removed under reduced pressure. The resulting solid was purified by column chromatography (silica, 100% CH2Cl2 followed by 9:1 CH2Cl2:MeOH) yielding a yellow oil (11.6 g, 0.062 mol, 80%). NMR spectroscopic data agree with the literature values.105,130 1H NMR (400 NSN  48 MHz, CD2Cl2) δ 8.50 (ddd, J = 4.8, 2.0, 0.9 Hz, 1H), 7.63 (ddd, J = 7.9, 7.5, 2.0 Hz, 1H), 7.43 (dt, J = 8.0, 1.0 Hz, 1H), 7.17 (ddd, J = 7.5, 4.8, 1.1 Hz, 1H).  2,2'-sulfinyldipyridine (DPSO) DPS (0.173 g, 0.919 mmol, 1.0 eq.) was dissolved in 2 mL glacial acetic acid, to which a 30% solution of H2O2 (0.35 mL, 3.68 mmol, 15.0 eq.) was added dropwise. The reaction mixture was stirred at 25 ºC for 42.5 h, followed by basification with 6 M NaOH, forming a yellow suspension which turned white after stirring for 30 min. Brine (50 mL) was added and the crude product extracted with CH2Cl2 (3 ´ 10 mL) followed by drying of the combined organics over MgSO4 and removal of the CH2Cl2 in vacuo. Purification was performed by column chromatography (silica, 95:5 CH2Cl2:MeOH), resulting in an off-white powder (0.107 g, 0.524 mmol, 57%). 1H NMR (400 MHz, CD2Cl2) δ 8.57 (d, J = 4.6 Hz, 2H, H4), 7.97 (d, J = 7.9 Hz, 2H, H1), 7.88 (td, J = 7.7, 1.6 Hz, 2H, H2), 7.39 – 7.31 (m, 2H, H3). 13C{1H} NMR (CD2Cl2, 100 MHz) δ 164.84 (C6), 150.23 (C4), 138.40 (C2), 125.47 (C3), 119.96 (C1). HR-ESI MS: m/z calcd. for C10H8N2ONaS: 227.0255; Found: 227.0255 [M+Na]+. IR (neat): ṽ (σ SO) 1037 cm−1.  2,2'-sulfonyldipyridine (DPSO2) DPS (0.743 g, 3.95 mmol, 1.0 eq.) was dissolved in 12 mL EtOH, to which NbC (0.048 g, 0.460 mmol, 0.25 eq.) was added, followed by a 30% solution of H2O2 (2.85 mL, 28.5 mmol, 5.0 eq.). The reaction mixture was heated at 60 ºC for 21 h. The mixture was cooled to room temperature, saturated Na2S2O3 solution (50 mL) added to quench any unreacted peroxides and the product extracted with CH2Cl2 (3 ´  15 mL).  The NSNO1234 56NSNO1234 56O  49 combined organics were washed with brine (20 mL), dried over MgSO4 and the solvent evaporated, yielding a white powder that required no further purification (0.730 g, 3.31 mmol, 84%). 1H NMR (400 MHz, CD2Cl2) δ 8.63 (ddd, J = 4.7, 1.8, 0.9 Hz, 2H, H4), 8.29 (d, J = 7.9 Hz, 2H, H1), 8.00 (td, J = 7.8, 1.7 Hz, 2H, H2), 7.52 (ddd, J = 7.7, 4.7, 1.0 Hz, 2H, H3). 13C{1H} NMR (CD2Cl2, 100 MHz) δ 157.46 (C6), 150.73 (C4), 138.47 (C2), 127.77 (C3),	124.18 (C1). HR-ESI MS: m/z calcd. for C10H8N2O2S: 220.0306; Found: 220.0307 [M]+. IR (neat): ṽ (σ SO2) 1170 and 1309 cm−1.  bis(4-methylpyridin-2-yl)sulfane (4,4'-Me-DPS) 2-Chloro-4-(methyl)pyridine (5.71 g, 44.8 mmol, 2.0 eq.) and thiourea (1.618 g, 23.3 mmol, 0.95 eq.) were dissolved in EtOH (500 mL) under a N2 atmosphere and heated to reflux for 150 h. After cooling to room temperature, the EtOH was removed, the crude product dissolved in CH2Cl2 (10 mL) and washed using water (3 ´ 10 mL), brine (10 mL), dried with MgSO4 and filtered.  Purification was performed by column chromatography (silica, 100% CH2Cl2 then 9:1 CH2Cl2:MeOH) and after removal of the solvents in vacuo a yellow oil was isolated (2.32 g, 10.7 mmol, 48%). 1H NMR (400 MHz, CD2Cl2) δ 8.37 – 8.30 (m, 2H, H4), 7.23 (dt, J = 1.6, 0.8 Hz, 2H, H1), 6.98 (ddd, J = 5.1, 1.6, 0.8 Hz, 2H, H3), 2.29 (d, J = 0.8 Hz, 6H, HMe). 13C{1H} NMR (101 MHz, CD2Cl2) δ 156.62 (C2), 149.71 (C4), 148.51 (C6), 126.45 (C1), 122.87 (C3), 20.69 (CMe). HR-ESI MS: m/z calcd. for C12H12N2S: 216.0721; Found: 217.0797 [M+H]+.    NSN1234 56H3CMeCH3  50 2,2'-sulfinylbis(4-methylpyridine) (4,4'-Me-DPSO)  4,4'-Me-DPS (0.300 g, 1.39 mmol, 1.0 eq.) was dissolved in CH2Cl2 (7 mL) at 0 ºC. A solution of m-CPBA (0.264 g, 1.40 mmol, 1.0 eq.) in CH2Cl2 (6 mL) was cooled to 0 ºC and added dropwise over 30 min to the diaryl sulfide. The reaction mixture was stirred at 0 ºC for 6 h and then extracted with 10% NaOH (2 ´ 20 mL), 5% HCl (2 ´ 20 mL) and 10% NaHCO3 (20 mL). The organic layer was dried over MgSO4, filtered and the solvent removed under reduced pressure yielding a white solid that required no further purification (0.249 g, 1.07 mmol, 77%). 1H NMR (400 MHz, CD2Cl2) δ 8.41 (d, J = 4.9 Hz, 2H, H4), 7.78 (dt, J = 1.8, 0.8 Hz, 2H, H1), 7.18 – 7.11 (m, 2H, H3), 2.41 (s, 6H, HMe). 13C{1H} NMR (101 MHz, CD2Cl2) δ 164.85 (C6), 150.57 (C2), 150.08 (C4), 126.58 (C3), 120.69 (C1), 21.60 (CMe). HR-ESI MS: m/z calcd. for C12H12N2ONaS: 255.0568; Found: 255.0575 [M+Na]+. IR (neat): ṽ (σ SO) 1045 cm−1.  2,2'-sulfonylbis(4-methylpyridine) (4,4'-Me-DPSO2)  4,4'-Me-DPS (0.590 g, 0.273 mmol, 1.0 eq.) was dissolved in 25 mL EtOH, to which NbC (0.077 g, 0.683 mmol, 0.25 eq.) was suspended and 30% H2O2 (4.1 mL, 40.9 mmol, 5.0 eq.) was added dropwise. The reaction mixture was heated to 60 ºC and stirred for 16 h. After cooling to room temperature saturated Na2S2O3 solution (80 mL) was added, the product extracted with CH2Cl2 (3 ´ 25 mL), the organics combined, dried over MgSO4 and filtered. After removing the solvent in vacuo a white powder was isolated that required no further purification (0.566 g, 2.28 mmol, 83%). 1H NMR (400 MHz, CD2Cl2) δ 8.45 (d, J = 4.9 Hz, 1H, H4), 8.12 (s, 1H, H1), 7.32 (d, J = 5.6 Hz, 1H, H3), 2.49 (s, 4H, HMe). 13C{1H} NMR (101 MHz, CD2Cl2) δ 157.47 (C2), 150.76 (C6), 150.51 NSNO1234 56H3CMeCH3NSNO1234 56H3CMeCH3O  51 (C4), 128.62 (C3), 124.96 (C1), 21.55 (CMe). HR-ESI MS: m/z calcd. for C12H12N2O2S: 248.0619; Found: 249.0706 [M+H]+. IR (neat): ṽ (σ SO2) 1157 and 1313 cm−1.   [Ir(ppy)2(DPS)][PF6] (Ir-DPS) [Ir(ppy)2Cl]2 (0.111 g, 0.104 mmol, 1.0 eq.) and AgPF6 (0.104 g, 0.415 mmol, 4.0 eq.) were added to a round-bottomed flask and suspended in MeOH (20 mL). The reaction mixture was stirred for 12 h at room temperature, during which AgCl formed as a grey precipitate. The reaction mixture was passed over Celite to remove the AgCl and the filtrate volume reduced. To this DPS (0.039 g, 0.208 mmol, 2.0 eq.) was added, the reaction mixture left to stir for 12 h at room temperature, following which the solvent was removed in vacuo. The crude solid was purified using column  chromatography (silica, 99:1 CH2Cl2:MeOH) giving a yellow crystalline powder (0.080 g, 0.096 mmol, 92%). 1H NMR (400 MHz, CD2Cl2) δ 8.17 (ddd, J = 5.9, 1.6, 0.8 Hz, 2H, H8), 8.13 (ddd, J = 5.7, 1.8, 0.7 Hz, 2H, H4), 8.03 (dd, J = 8.0, 1.2 Hz, 2H, H11), 8.01 – 7.89 (m, 4H, H1,10), 7.88 (td, J = 7.7, 1.8 Hz, 2H, H2), 7.72 (dd, J = 7.7, 1.3 Hz, 2H, H14), 7.23 – 7.14 (m, 4H, H3,9), 7.07 (td, J = 7.5, 1.2 Hz, 2H, H15), 6.93 (td, J = 7.5, 1.4 Hz, 2H, H16), 6.24 (dd, J = 7.6, 1.1 Hz, 2H, H17). 13C{1H} NMR (101 MHz, CD2Cl2) δ 168.12 (C13), 154.34 (C6), 154.20 (C4), 151.49 (C8), 147.48 (C18), 144.39 (C12), 139.95 (C2), 139.14 (C10), 132.03 (C17), 130.91 (C16), 129.57 (C1), 126.26 (C3), 125.32 (C14), 123.51 (C15), 123.09 (C9), 120.50 (C11). HR-ESI MS: m/z calcd. for C32H24IrN4S: 689.1351; Found: 689.1316 [M−PF6]+.   IrNNNSNPF612345 6789101112131415161718  52 [Ir(ppy)2(DPSO)][PF6] (Ir-DPSO) [Ir(ppy)2Cl]2 (0.111g, 0.104 mmol, 1.0 eq.) and AgPF6 (0.110 g, 0.435 mmol, 4.0 eq.) were suspended in MeOH (20 mL) and the reaction mixture stirred for 12 h at room temperature. AgCl was removed by filtration over Celite and then the filtrate dried under reduced pressure. The resulting solid was re-dissolved in EtOH (20 mL) and DPSO (0.042 g, 0.207 mmol, 2.0 eq.) added, following which the solution was stirred at room temperature for 12 h. EtOH was removed in vacuo, the crude solid purified using column chromatography (silica, 99:1 CH2Cl2:MeOH) and the eluted product precipitated with hexanes giving a yellow crystalline solid (0.037 g, 0.044 mmol, 42%). 1H NMR (850 MHz, CD2Cl2) δ 8.83 (ddd, J = 5.9, 1.5, 0.8 Hz, 1H, H8), 8.59 (ddd, J = 5.6, 1.6, 0.7 Hz, 1H, H4), 8.38 (ddd, J = 8.0, 1.5, 0.7 Hz, 1H, H31), 8.27 (ddd, J = 8.0, 1.5, 0.7 Hz, 1H, H1), 8.22 – 8.20 (m, 3H, H2,,28,32), 7.96 (ddd, J = 8.3, 7.4, 1.5 Hz, 1H, H27), 7.92 – 7.85 (m, 3H, H10,11,23), 7.68 (ddd, J = 5.5, 1.6, 0.7 Hz, 1H, H34), 7.56 (ddt, J = 7.7, 1.4, 0.5 Hz, 1H, H14), 7.52 (ddd, J = 7.5, 5.6, 1.5 Hz, 1H, H3), 7.28 (ddd, J = 7.4, 5.9, 1.6 Hz, 1H, H9), 7.27 – 7.22 (m, 2H, H33), 7.11 (ddd, J = 7.8, 7.2, 1.1 Hz, 1H, H22), 7.04 (ddd, J = 7.7, 7.3, 1.2 Hz, 1H, H15), 7.03 – 6.98 (m, 1H, H16), 6.93 (ddd, J = 5.9, 1.5, 0.8 Hz, 1H, H29), 6.91 (ddd, J = 7.8, 7.2, 1.4 Hz, 1H, H21), 6.27 (ddd, J = 7.8, 1.2, 0.5 Hz, 1H, H20), 6.19 (ddd, J = 7.6, 1.2, 0.5 Hz, 1H, H17). 13C{1H} NMR (101 MHz, CD2Cl2) δ 169.10 (C12), 167.69 (C25), 164.38 (C6), 161.96 (C36), 154.60 (C4), 154.35 (C8), 152.87 (C34), 148.75 (C29), 145.52 (C18), 145.09 (C13), 144.81 (C19), 144.11 (C24), 141.55 (C32), 141.36 (C28), 139.90 (C27), 139.73 (C10), 132.59 (C17), 131.73 (C20), 131.68 (C16), 131.06 (C21), 128.93 (C3), 128.51 (C33), 126.23 (C23), 125.03 (C14), 124.51 (C15), 124.23 (C22), 124.17 (C2), 123.76 (C31), 123.33 (C9), 123.24 (C1), IrNNNSNPF612345 678910111213141516171819O2021222324 2526272829303132333435 36  53 121.58 (C26), 120.56 (C11). HR-ESI MS: m/z calcd. for C32H24IrN4OS: 705.1300; Found: 705.1277 [M−PF6]+  [Ir(ppy)2(DPSO2)][PF6] (Ir-DPSO2)  [Ir(ppy)2Cl]2 (0.090 g, 0.084 mmol, 1.0 eq.) and AgPF6 (0.104 g, 0.415 mmol, 4.0 eq.) were suspended in MeOH (20 mL). The reaction mixture was stirred for 12 h at room temperature, during which time AgCl formed as a precipitate. This was removed by filtration over Celite, followed by evaporation of the solvent. The resulting solid was re-dissolved in EtOH (20 mL) and DPSO2 (0.037 g, 0.168 mmol, 2.0 eq.) was added. The reaction mixture was left to stir overnight at room temperature, following which the EtOH was removed under reduced pressure and the crude solid dissolved in CH2Cl2. This mixture was passed over Celite three times and the product precipitated with hexanes, yielding a yellow powder (0.020 g, 0.023 mmol, 28%). 1H NMR (400 MHz, CD2Cl2) δ 8.72 (d, J = 7.8 Hz, 2H, H1), 8.36 – 8.31 (m, 2H, H4), 8.29 – 8.24 (m, 2H, H2), 7.99 – 7.94 (m, 2H, H8), 7.87 (td, J = 7.8, 1.5 Hz, 2H, H10), 7.75 – 7.69 (m, 2H, H14), 7.61 (d, J = 5.8 Hz, 2H, H11), 7.52 (m, 2H, H3), 7.13 – 7.05 (m, 4H, H9,15), 6.94 (td, J = 7.5, 1.4 Hz, 2H, H16), 6.13 (dd, J = 7.7, 1.2 Hz, 2H, H17). 13C{1H} NMR (101 MHz, CD2Cl2) δ 167.11 (C12), 155.41 (C4), 155.05 (C6), 152.09 (C11), 144.98 (C13), 142.52 (C18), 141.89 (C2), 139.51 (C10), 132.22 (C17), 131.71 (C3), 131.25 (C16), 127.54 (C1), 125.59 (C14), 124.52 (C15), 123.62 (C9), 120.36 (C8). HR-ESI MS: m/z calcd. for C32H24IrN4O2S: 721.1249; Found: 721.1232 [M−PF6]+.   IrNNNSNPF612345 6789101112131415161718OO  54 Synthesis of [Ir(ppy)2(4,4'-Me-DPS)][PF6] (Ir-4,4'-Me-DPS)  [Ir(ppy)2Cl]2 (0.107 g, 0.099 mmol, 1.0 eq.) and AgPF6 (0.101 g, 0.400 mmol, 4.0 eq.) were suspended in MeOH (20 mL) and stirred for 12 h at room temperature. Precipitated AgCl was removed by filtration over Celite and the dried filtrate redissolved in EtOH (20 mL) and 4,4'-Me-DPS (0.050 g, 0.199 mmol, 2.0 eq.) added. The reaction mixture was stirred at room temperature for 12 h. The filtrate was purified using column chromatography (silica, 100% CH3CN, followed by 97:3:1, CH3CN:H2O:KNO3 (sat.)) to yield a yellow crystalline solid (0.080 g, 0.093 mmol, 94%). 1H NMR (400 MHz, CD2Cl2) δ 8.13 (ddd, J = 5.9, 1.5, 0.7 Hz, 2H, H14), 8.00 – 7.95 (m, 2H, H17), 7.92 – 7.86 (m, 4H, H4,16), 7.72 (t, J = 1.5 Hz, 2H, H1), 7.67 (dd, J = 7.8, 1.3 Hz, 2H, H11), 7.13 (ddd, J = 7.4, 5.9, 1.5 Hz, 2H, H15), 7.01 (td, J = 7.5, 1.2 Hz, 2H, H10), 6.92 (ddd, J = 5.8, 2.0, 0.8 Hz, 2H, H3), 6.88 (td, J = 7.5, 1.4 Hz, 2H, H9), 6.20 (dd, J = 7.7, 1.2 Hz, 2H, H8), 2.35 (s, 6H, HMe). 13C{1H} NMR (101 MHz, CD2Cl2) δ 168.35 (C18), 153.86 (C6), 153.48 (C4), 152.81 (C2), 151.59 (C14), 148.17 (C12), 144.58 (C13), 139.14 (C16), 132.21 (C8), 131.00 (C9), 130.09 (C1), 127.32 (C3), 125.41 (C11), 123.50 (C10), 123.12 (C15), 120.53 (C17), 21.21 (CMe). HR-ESI MS: m/z calcd. for C34H28IrN4S: 717.1664; Found: 717.1727 [M−PF6]+.       IrNNNSNPF612345 6789101112131415161718CH3CH3Me  55 [Ir(ppy)2(4,4'-Me-DPSO)][PF6] (Ir-4,4'-Me-DPSO) [Ir(ppy)2Cl]2 (0.076 g, 0.073 mmol, 1.0 eq.) was suspended in MeOH (25 mL) along with AgPF6 (0.056 g, 0.220 mmol, 3.0 eq.) and stirred at room temperature for 12 h. The reaction mixture was passed over Celite to remove precipitated AgCl, and 4,4'-Me-DPSO (0.034 g, 0.146 mmol, 2.0 eq.) was added to the filtrate, after which the solution was left to stir for 12 h at room temperature. The product precipitated out of solution and was filtered, giving a yellow powder that required no further purification (0.041 g, 0.047 mmol, 61%). 1H NMR (850 MHz, CD2Cl2) δ 8.82 (ddd, J = 5.9, 1.5, 0.8 Hz, 1H, H8), 8.37 (d, J = 5.8 Hz, 1H, H4), 8.18 – 8.15 (m, 1H, H26), 8.14 (dq, J = 2.0, 0.6 Hz, 1H, H31), 8.04 (dt, J = 2.0, 0.6 Hz, 1H, H1), 7.95 (ddd, J = 8.3, 7.4, 1.5 Hz, 1H, H27), 7.89 – 7.84 (m, 3H, H10,11,23), 7.55 (ddt, J = 7.7, 1.3, 0.5 Hz, 1H, H14), 7.48 – 7.46 (m, 1H, H34), 7.29 (ddt, J = 5.8, 2.0, 0.7 Hz, 1H, H3), 7.26 (ddd, J = 7.4, 5.9, 1.7 Hz, 1H, H9), 7.08 (ddd, J = 7.9, 7.2, 1.2 Hz, 1H, H22), 7.05 – 7.03 (m, 1H, H15), 7.02 – 7.00 (m, 2H, H28,33), 6.98 – 6.96 (m, 2H, H16,29), 6.89 (ddd, J = 7.8, 7.2, 1.4 Hz, 1H, H1), 6.26 (ddd, J = 7.8, 1.2, 0.5 Hz, 1H, H20), 6.19 (ddd, J = 7.6, 1.2, 0.5 Hz, 1H, H17), 2.52 (d, J = 0.7 Hz, 3H, HMe-1), 2.47 (d, J = 0.7 Hz, 3H, HMe-2). 13C{1H} NMR (101 MHz, CD2Cl2) δ 169.09 (C12), 167.67 (C25), 163.55 (C6), 161.13 (C36), 154.86 (C2), 154.61 (C32), 154.24 (C8), 153.78 (C4), 152.06 (C34), 148.82 (C29), 146.07 (C18), 145.34 (C19), 145.07 (C13), 144.16 (C24), 139.72 (C27), 139.52 (C10), 132.61 (C17), 131.67 (C20), 131.54 (C16), 130.90 (C21), 129.65 (C3), 129.14 (C28), 126.11 (C23), 124.90 (C14), 124.25 (C15), 124.11 (C31,33), 123.91 (C22), 123.57 (C1), 123.19 (C9), 121.38 (C26), 120.39 (C11), 21.97 (CMe-1,Me-2). HR-ESI MS: m/z calcd. for C34H28IrN4OS: 733.1613; Found: 733.1542 [M−PF6]+. IrNNNSNPF612345 678910111213141516171819O2021222324 2526272829303132333435 36CH3CH3Me-1Me-2  56  [Ir(ppy)2(4,4'-Me-DPSO2)][PF6] (Ir-4,4'-Me-DPSO2) [Ir(ppy)2Cl]2 (0.092 g, 0.086 mmol, 1.0 eq.) and AgPF6 (0.090 g, 0.356 mmol, 4.1 eq.) were suspended in MeOH (15 mL) and the reaction mixture stirred for 12 h at room temperature. AgCl precipitated out of the solution, the mixture was filtered over Celite and the filtrate dried under reduced pressure. The resulting solid was re-dissolved in EtOH (20 mL) and 4,4'-Me-DPSO2 (0.043 g, 0.172 mmol, 2.0 eq.) added, following which the solution was stirred at room temperature for 12 h. The EtOH was then removed in vacuo and the crude mixture re-dissolved in CH2Cl2, following which the solution was filtered over Celite two times to remove a brown impurity. The resulting product was then precipitated out of solution with hexanes, yielding a yellow solid (0.025 g, 0.028 mmol, 33%). 1H NMR (400 MHz, CD2Cl2) δ 8.51 (d, J = 2.0 Hz, 2H, H1), 8.12 (d, J = 5.7 Hz, 2H, H4), 7.97 – 7.94 (m, 2H, H8), 7.89 – 7.83 (m, 2H, H10), 7.70 (dd, J = 7.8, 1.4 Hz, 2H, H14), 7.61 (d, J = 5.9 Hz, 2H, H11), 7.27 (ddd, J = 5.7, 2.1, 0.9 Hz, 2H, H3), 7.11 – 7.04 (m, 4H, H9,15), 6.92 (td, J = 7.5, 1.4 Hz, 2H, H16), 6.13 (dd, J = 7.7, 1.1 Hz, 2H, H17), 2.53 (s, 6H, HMe). 13C{1H} NMR (101 MHz, CD2Cl2) δ 167.16 (C12), 155.22 (C2), 154.65 (C4), 154.41 (C6), 152.09 (C11), 145.01 (C13), 143.14 (C18), 139.36 (C10), 132.21 (C3), 132.12 (C17), 131.14 (C16), 128.11 (C1), 125.49 (C14), 124.30 (C15), 123.53 (C9), 120.24 (C8), 21.91 (CMe). HR-ESI MS: m/z calcd. for C34H28IrN4O2S: 749.1562; Found: [M−PF6]+.   IrNNNSNPF612345 6789101112131415161718CH3CH3MeOO  57 CHAPTER 3: Substituent and Oxidation State Influence on Sulfur-Bridged Luminescent Copper(I) Complexes Showing Delayed Fluorescence  3.1 Introduction Luminescent materials based on transition metal complexes have long afforded substantial research efforts due to their ability to be implemented in a range of technologies. These include solar applications,6 photocatalysis18 and, in particular, lighting technologies such as organic light-emitting diodes (OLEDs)10 or light-emitting electrochemical cells (LEECs).11 Often these applications are dominated by the use of second- and third-row transition metals such as Ir,13,14 Ru131 and Pt132 due to their high degrees of colour tuneability, efficiencies and stabilities. These metals, however, are expensive and limited by low abundancies in the Earth’s crust, and as such efforts are being made to find more cost-effective alternatives. In particular, coinage metals with a d10 configuration such as Cu(I), Ag(I) and Au(I) are promising candidates, with copper(I) especially attracting considerable interest from the research community in the past decade.71,133,134 Emissive copper(I) complexes have been shown to adopt a diverse range of structures. These range from homo- or heteroleptic mononuclear species containing diimine (N^N), diphosphine (P^P) or imine-phosphine (N^P) ligands,57,58,135 to polynuclear species that can be bridged by the aforementioned ligands136 or halide atoms,137 to tetranuclear cubane structures.54 This ability for copper(I) to accept a variety of structures allows for an array of different photophysical properties.138 The photophysics of copper(I) complexes have been extensively studied in solution, in the solid state and in polymer environments.139 In the ground-state, the Cu(I) center is of a d10   58 configuration adopting a pseudotetrahedral (D2d) geometry, which on photoexcitation undergoes a metal-to-ligand charge transfer (MLCT) formally oxidizing the metal center to a d9 Cu(II) configuration. Once in this excited-state, the complex undergoes a fast (<1 ps) Jahn-Teller distortion to a more flattened geometry, opening up nonradiative decay pathways such as exciplex formation with coordinating species. This distorted geometry is preserved until relaxation to the ground-state. Structural modification of the ligands has been shown to inhibit this distortion through rational design aimed at introducing constraints to maintain a tetrahedral geometry, thus increasing the lifetime of the excited-state and the quantum efficiency of the complex.65 Furthermore, Cu(I) species can undergo an additional radiative decay pathway via thermally activated delayed fluorescence (TADF), whereby otherwise non-emissive triplet states can be converted to singlet states, resulting in higher emission quantum yields.71 In order to satisfy the demands of TADF a small S1–T1 energy gap (<0.42 eV) must exist for reverse intersystem crossing (RISC) to be viable.9,21  In Chapter 2: it was shown that on increasing the oxidation state at the sulfur from sulfide (S), to sulfoxide (SO), to sulfone (SO2) it was possible to switch from a blue-green emissive state of mainly 3LC character (in the case of S and SO) to one of 3MLCT/3LLCT character giving yellow emission (for SO2). The sulfur-bridged dipyridyl ligands were also found to adopt a larger bite angle than their bipyridine counterparts, which in conjunction with additional possible oxygen binding sites for the sulfoxide and sulfone analogues can lead to interesting binding modes. In this chapter the synthesis and characterization of a series of four new heteroleptic Cu(I) species utilizing sulfur-bridged dipyridyl ligands in either sulfide or sulfone oxidation state as the N^N component, and bis[2-(diphenylphosphino)-phenyl]ether (POP) as the P^P component are   59 discussed. It is shown that subtle steric changes at the 6- and 6'-positions of the N^N ligand can influence the formation of bridged dinuclear complexes.  3.2 Ligand Synthesis and Characterization Scheme 3-1 Synthesis of Cu(I) Complexes    The sulfide-bridged compounds DPS and Me-DPS were synthesized via nucleophilic aromatic substitution of thiourea with 2-bromopyridine and 2-bromo-6-methylpyridine, respectively (Scheme 3-1). Desired sulfone species DPSO2 and Me-DPSO2 were obtained through NBr Thiourea (0.5 eq.)EtOH, 78 °C, N2, 36–60 h30% H2O2 (15.0 eq.)NbC (0.25 eq)EtOH, 60 °C, 12–21 hDPS: R = HMe-DPS: R = MeR = H, Me, PhRDPSO2: R = HMe-DPSO2: R = MeNSNRRNSNRROOPPh2 PPh2CH2Cl2, 25 °C, 2 hNSNO n1.0 eq.NSN2.0 eq.Me MeCH2Cl2, 25 °C, 1 hCH2Cl2, 25 °C, 1 hNSN0.5 eq.Me MeCH2Cl2, 25 °C, 1 hO OCuPOPPh PhPhPhNNSBF4Cu-DPS: n = 0Cu-DPSO2: n = 2PPhPhCuOP CuNNSMeMePhPh[BF4]2NNSMeMeCu-Me-DPSO2POPPh PhPhPhCuNSOMeNMeOPOPPhPhPh PhCuO nCu-Me-DPS[Cu(MeCN)4]BF4[BF4]2O  60 oxidation of the appropriate diaryl sulfide pro-ligand with 30% H2O2 using a niobium carbide (NbC) catalyst. High yields were attained, and the four pro-ligands were characterized using NMR spectroscopy, mass spectrometry and infrared spectroscopy. 1H and 13C signals were assigned using COSY, HSQC and HMBC experiments.  3.3 Synthesis and Characterization of Cu(I) Complexes The four copper(I) complexes were prepared using the synthetic methods described by Linfoot et al.140 (Scheme 3-1). To a stirring solution of [Cu(MeCN)4][BF4] in CH2Cl2 at 25 ˚C was added the desired molar ratio of POP followed by the desired molar ratio of the N^N pro-ligand (N^N = DPS, DPSO2, Me-DPS or Me-DPSO2). Following precipitation from solution using Et2O, the species were all isolated in high yields as off-white to yellow powders that required no further purification. 1H, 13C and 31P NMR spectra were collected and assigned using COSY, HSQC and HMBC experiments, with the broad 31P signals that are observed attributed to quadrupole coupling with the copper atom. Complexes Cu-DPS and Cu-DPSO2 were synthesized using equimolar ratios of [Cu(MeCN)4][BF4], POP and the appropriate DPS or DPSO2 ligand and isolated as monometallic species. HR-ESI mass spectra show a product peak corresponding to [M–BF4]+ and exhibit a characteristic copper isotope pattern. Elemental analyses agree with the expected values for the assigned complexes.  3.3.1 Nuclear Magnetic Resonance (NMR) Spectroscopy Interestingly, on increasing steric bulk through addition of methyl groups in the 6- and 6'-positions on the dipyridyl ligand (Me-DPS and Me-DPSO2) bimetallic complexes Cu-Me-DPS   61 and Cu-Me-DPSO2 were isolated (Scheme 3-1). While analogous complex [Cu(POP)(tmbpy)][BF4] (tmbpy = 4,4',6,6'-tetramethyl-2,2´-bipyridine) is monometallic,76 the larger bite angle afforded by the addition of a sulfur atom between the two pyridyl rings coupled with the added bulk of the methyl substituents presumably causes the steric constraints to be too great to form the typical P,P and N,N-bound four-coordinate species.  NMR experiments for Cu-Me-DPS show a bimetallic complex with each discrete Cu(I) complexed with the Me-DPS ligand, and the two metal centres bridged by the POP ligand. This bridged binding mode for the POP ligand has not been reported before. The 1H NMR spectrum of Cu-Me-DPS shows sharp signals for the POP ligand, but broad signals attributed to the Me-DPS ligand, likely caused by fluxional behavior on the NMR timescale. The use of low-temperature NMR experiments was unable to elucidate a defined structure. Cu-Me-DPSO2 was found to also exist as a bimetallic species; however, instead of POP acting as a bridging ligand, Me-DPSO2 connects the two copper(I) atoms with each discrete metal center bound in a N,O fashion through one pyridyl ring and one of the sulfone oxygen atoms. Four-coordinate geometry about each Cu(I) is completed by a POP ligand. Syntheses of Cu-Me-DPS and Cu-Me-DPSO2 were also conducted using equimolar equivalents of [Cu(MeCN)4][BF4], diphosphine and diimine ligands but the bridged, bimetallic structures still preferentially formed.  3.3.2 Single Crystal X-ray Diffraction Single crystals of Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 were grown by vapor diffusion of Et2O into CH2Cl2 solutions of the complexes. Steric factors strongly affect the coordination chemistry of copper(I) when complexed with POP ligands. POP is a bulky bidentate   62 component and it has been shown that when two POP ligands coordinate to copper(I), the metal adopts a trigonal geometry to reduce steric hindrance. In [Cu(POP)2][BF4] it has been shown that only one POP is able to behave as a chelating unit, while the second ligand can only bind through one phosphorus atom forming a trigonal species.141 In contrast, copper(I) forms tetrahedral complexes in less sterically hindered systems, for example when one POP is present together with smaller ligands.142-144     63  Figure 3-1 ORTEP representations of single crystal structures of (a) Cu-DPS; (b) Cu-Me-DPS; (c) Cu-DPSO2; (d) Cu-Me-DPSO2. Ellipsoids are plotted at the 50% probability level. Hydrogen atoms and solvent molecules are removed for clarity. The four aforementioned copper(I) complexes were structurally characterized (Figure 3-1), with the data corroborating the NMR spectral observations. The four dipyridyl ligands each differ in both steric bulk and binding modes. DPS and Me-DPS are both bidentate ligands which coordinate to copper through the pyridyl nitrogen atoms, with Me-DPS possessing a larger steric bulk due to the presence of methyl groups in 6- and 6'-positions. DPSO2 and Me-DPSO2 each   64 have two different binding modes, either through the pyridyl nitrogen atoms (N,N), or through one nitrogen and one sulfone oxygen (N,O). These ligands can therefore adopt a preferential coordination mode in order to reduce steric tension within the complex. The result is that in the presence of methyl groups the complexes crystalize as dimers where either POP or Me-DPSO2 bridge two copper atoms.  Cu-DPS (Figure 3-1(a)) and Cu-DPSO2 (Figure 3-1(c)) both crystalize as monomers with the copper center presenting a distorted tetrahedral geometry, and in both cases the dipyridyl ligand binds in an N,N fashion. Cu-Me-DPS (Figure 3-1(b)) crystalizes as a dimer in which the POP bridges two copper atoms with a distorted trigonal geometry. Cu-Me-DPSO2 (Figure 3-1(d)) also crystalizes as a dimer, but differently from Cu-Me-DPS. In this case the Cu(I) atoms have tetrahedral geometry with the bridging ligand Me-DPSO2 ligand binding with N,O coordination modes to the two metal centers. Selected bond lengths and angles for the four complexes are reported in Table 3-1. All the complexes exhibit a distorted geometry. This is not surprising because DPS and DPSO2 ligands are more rigid compared to POP and as a result the N–Cu–N bite angles are expected to be smaller than the P–Cu–P bite angles. In addition, the P–Cu–P angles are anticipated to be quite large due to the presence of the bulky PPh3 groups. Structural studies show that in all the tetrahedral complexes the P–Cu–P bite angles are in the range of 113–117˚, and the N–Cu–N bite angles are between 93–97˚. The N–Cu–O angle in Cu-Me-DPSO2 is even smaller at 79.35˚. The trigonal Cu-Me-DPS complex also has a distorted structure, with a N–Cu–N angle of 96.47˚ and a P–Cu–N angle of 129.53˚. The Cu–P and Cu–N bond lengths are all similar within the series reported here and to other mononuclear copper(I) complexes in the literature.136     65 Table 3-1 Selected bond lengths (Å) and angles (˚) of Cu-DPS, Cu-DPSO2, Cu-Me-DPS, and Cu-Me-DPSO2.  Cu-DPS Cu-DPSO2 Cu-Me-DPS Cu-Me-DPSO2 Cu1–N1 2.0804(14) 2.1186(10) 2.038(3) 2.0480(16) Cu1–N2 2.0691(13) 2.1163(10) 2.040(3) – Cu1–O1 – – – 2.4134(13) Cu1–P1 2.2463(4) 2.3052(3) 2.2131(10) 2.2719(5) Cu1–P2 2.2659(4) 2.2508(3) – 2.2292(5) N1–Cu1–N2 93.89(6) 97.03(4) 96.47(12) – N1–Cu1–O1 – – – 79.35(5) P1–Cu1–P2 113.268(16) 114.029(12) – 117.12(2) P1–Cu1–N1 117.94(4) 99.95(3) 129.52(8) 108.44(4)   3.4 Electrochemical Properties The electrochemical behavior of the Cu(I) complexes was investigated using cyclic voltammetry (CV) of CH2Cl2 solutions vs. Fc/Fc+ and the data summarized in Table 3-2. Two oxidative features are seen for each complex, with the first attributed to the Cu+/Cu2+ oxidation.142 In comparing the mononuclear species, on oxidizing Cu-DPS to Cu-DPSO2 an increase in oxidation potential is seen (+1.03 V to +1.17 V) owing to the electron-withdrawing effect of the oxygen atoms of the sulfone-bridged ligand on the coordinating nitrogen atoms. These potentials are higher than those seen for the model complex [Cu(POP)(bpy)][PF6] which has an unsubstituted bpy ligand (+0.72 V).145     66 Table 3-2 Electrochemical data of Cu(I) complexes in CH2Cl2 solutions.  E(ox) (V) E(red) (V) EHOMO (eV) [a] ELUMO (eV) [a] ∆Eredox (V) Cu-DPS 1.03 [b] 1.64 [b] –0.93 [b] –6.13 –4.17 1.96 Cu-DPSO2 1.17 [b] 1.74 [b] –1.01 [b] –6.27 –4.09 2.18 Cu-Me-DPS 1.08 [b] 1.67 [b] –1.01  [b] –6.18 –4.09 2.09 Cu-Me-DPSO2 1.13 [b] 1.73 [b] –0.99 [b] –6.23 –4.11 2.12 Measurements were carried in CH2Cl2 at a scan rate of 50 mVs–1 with Fc/Fc+ employed as an internal standard, with data reported vs. SCE (Fc/Fc+ = 0.46 V in CH2Cl2).146 [a] HOMO and LUMO energies were calculated following EHOMO = –[Eox vs. Fc/Fc+ + 5.1] eV, ELUMO = –[Ered vs. Fc/Fc+ + 5.1] eV.147 [b] Irreversible redox wave.  In comparing the two bimetallic species an increase in oxidative potential is again observed for the sulfone vs sulfide complex (+1.08 V for Cu-Me-DPS and +1.13 V for Cu-Me-DPSO2) which can be attributed to electronic differences (as above), but possibly also to due to steric effects. It is known that by increasing the stabilization of the tetrahedral geometry about the copper center the Cu+/Cu2+ oxidation potential is raised,53 and so by adding substituents in the 6,6′-positions at the diimine ligand the flattening distortion is reduced. Additionally, in going from three-coordinate Cu-Me-DPS to four-coordinate Cu-Me-DPSO2 the stabilization of the tetrahedral geometry is further increased. Irreversible redox waves are seen for all four complexes, which is typical for transition metal complexes.145    67 3.5 Solution-state Photophysical Properties 3.5.1 Absorption Spectroscopy  Figure 3-2 Absorption spectra of (a) pro-ligands and (b) Cu(I) complexes in CH2Cl2 solutions. All spectra were recorded in ~ 2 ´ 10–5 M CH2Cl2 solutions at room temperature.  The  absorption spectra of sulfide and sulfone pro-ligands DPS, DPSO2, Me-DPS and Me-DPSO2 in CH2Cl2 solutions are shown in Figure 3-2(a). Features at energies higher than 380 nm are attributed to π–π* transitions, and absorptions in Me-DPS and Me-DPSO2 are red-shifted relative to those of their unsubstituted counterparts due to the electron-donating nature of the methyl substituents. Absorption spectra of the four Cu(I) complexes in CH2Cl2 solutions are shown in Figure 3-2(b). All show intense high energy bands assigned to ligand-based transitions and broad, lower energy bands in the range 325–450 nm due to metal-to-ligand charge transfer (MLCT) transitions, typical of [Cu(P^P)(N^N)]+ complexes,148 with the sulfone derivatives red-shifted compared to those of the sulfide.    68 3.5.2 Photoluminescence Spectroscopy  Figure 3-3 Emission spectra of Cu(I) complexes in CH2Cl2 solutions at room temperature. Concentrations ~ 2 ´ 10–5 M. λex = 370 nm.  The photoluminescence spectra of Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 were collected in deaerated CH2Cl2 solutions (Figure 3-3).  All complexes were found to be poorly emissive in solution presumably due to Jahn-Teller distortions in the excited-state leading to a quenching of the emission. The relative emission energies can be explained using steric and electronic arguments. As discussed in Section 1.3.5, Cu(I) complexes undergo a significant structural relaxation in the excited-state leaving a stabilized T1 state. While the absorption occurs at the optimal equilibrium geometry of S0, the emission occurs at the relaxed T1 geometry and therefore determines to a large extent the energy of the radiative decay.  Unsubstituted complex Cu-DPS is expected to show the largest geometrical change and therefore the greatest stabilization of the T1 energy and thus has the lowest energy emission (640 nm) of the four species. On oxidizing the ligand to Cu-DPSO2 the emission undergoes a large hypsochromic shift to 566 nm. Of the methyl-substituted species, Cu-Me-DPS exhibits the lowest   69 emission energy (632 nm) due to the distorted trigonal geometry about the copper centre, while the excited-state of Cu-Me-DPSO2 can be stabilized by the methyl groups in the 6,6′-positions of the bridging diimine ligand giving emission at 596 nm.  3.6 Solid state Photophysical Properties 3.6.1 Thin Film Photoluminescence  Figure 3-4 Solid state emission spectra of Cu(I) complexes as thin films drop-cast from MeOH on quartz. λex = 370 nm.  In the solid state at room temperature, Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 show much higher emission intensities in addition to a blue-shift when compared to their dissolved counterparts. This is characteristic of [Cu(P^P)(N^N)]+ complexes and is attributed to the more rigid environment severely restricting the degree of excited-state Jahn-Teller distortions. The photoluminescence spectra of the four Cu(I) complexes as thin films are shown in Figure 3-4 and their photophysical properties summarized in Table 3-3. In both the monometallic (Cu-DPS, λmax = 525 nm; Cu-DPSO2, λmax = 560 nm) and bimetallic (Cu-Me-DPS, λmax = 518 nm; Cu-Me-  70 DPSO2, λmax = 557 nm) species, oxidation to the sulfone results in a bathochromic shift of the emission wavelength. Functionalizing the diimine ligand with electron-withdrawing oxygen atoms at the sulfur reduces the electron density at the nitrogen atoms, which in turn reduces the ligands σ-donating and increases their π-accepting abilities. This destabilizes the LUMO, reducing the HOMO–LUMO gap and resulting in a lower energy emission.  Table 3-3 Solid state photophysical properties of Cu(I) complexes as neat thin films at room temperature in air. Complex [a] λmax [nm] [b] PLQY [b] τem [ns] [b][c] Cu-DPS 525 0.04 66.3 (8%), 451 (23%), 3720 (69%) Cu-DPSO2 560 <0.01 31.2 (11%), 218 (34%), 1930 (55%) Cu-Me-DPS 518 0.06 83.4 (9%), 467 (24%), 2810 (67%) Cu-Me-DPSO2 557 0.14 90.0 (3%), 546 (27%), 2280 (70%) [a] All samples drop-cast on to quartz from MeOH. [b] λex = 370 nm. [c] Lifetimes collected over both a 50 ns and 5 µs time regime in order to retain accuracy for the short and long components. Percentage contributions relate to 5 µs regime.  3.6.2 Thin Film Photoluminescence Quantum Yields (PLQY) Photoluminescence quantum yields (Table 3-3) follow a general trend of increasing as the environment about the copper becomes more constrained, likely due to improved stabilization of the tetrahedral geometry. Cu-Me-DPSO2 has the highest PLQY of 0.14, however this is lower than other POP-containing complexes that feature methyl groups in the 6,6′-positions of the diimine ligand (e.g. PLQY is 0.74 for [Cu(POP)(tmbpy)][BF4]).76 This could be due to the flexibility about the sulfur center providing additional non-radiative decay pathways when compared to bpy-based N^N ligands. Cu-DPSO2 is an outlier in this trend and is practically non-  71 emissive. In addition to the lack of steric bulk at the diimine, the electron-withdrawing oxygen atoms reduce the Cu–N bond strength enabling a greater stabilization of the T1 excited-state. Single crystal structures (Section 3.3.2) show that the Cu–N bond distances are longer for Cu-DPSO2 (2.116 and 2.119 Å) than for Cu-DPS (2.069 and 2.080 Å).  3.6.3 Low Temperature Thin Film Photoluminescence  Figure 3-5 Solid state emission spectra of (a) Cu-DPS and Cu-DPSO2; (b) Cu-Me-DPS and Cu-Me-DPSO2 at room temperature (solid line) and 77 K (dashed line). Thin films drop-cast from MeOH on quartz. λex = 370 nm.  As discussed in Section 1.3.6, TADF occurs when the S1–T1 energy gap is <0.43 eV and that for practical applications this energy gap should be no more than 0.12 eV. This energy gap can be determined experimentally for Cu(I) complexes using variable-temperature steady-state emission spectroscopy. At ambient temperature, the emission originates from the S1 state, but on cooling to 77 K the emission is from the T1, as observed by a red-shift in the emission spectrum due to the triplet state residing at a lower energy. An estimation of ΔE(S1–T1) can be made by taking the difference between the λmax emission at room temperature and 77 K. The emission   72 spectra of the four copper(I) complexes were therefore collected at room temperature and 77 K (Figure 3-5) and the ΔE(S1–T1) calculated (Table 3-4).  Table 3-4 ΔE(S1–T1) of Cu(I) complexes calculated from variable temperature solid state emission data.  Complex [a] λmax [nm] [b]  ΔE(S1–T1) [eV] Room Temperature 77 K Cu-DPS 525 536 0.04 Cu-DPSO2 560 584 0.05 Cu-Me-DPS 518 528 0.05 Cu-Me-DPSO2 557 578 0.08 [a] All samples drop-cast on to quartz from MeOH. [b] λex = 370 nm.   In all instances for Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 the S1–T1 energy gap was found to be less than the 0.12 eV required for the RISC process, indicating that the four complexes are candidates as TADF emitters.    73 3.6.4 Thin Film Photoluminescence Lifetimes  Figure 3-6 Emission lifetime measurements displaying the long lifetime component of thin films (a) Cu-DPS; (b) Cu-DPSO2; (c) Cu-Me-DPS and (d) Cu-Me-DPSO2 are plotted. NanoLED excitation source (λex = 370 nm).  Excited-state lifetimes of thin films of Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 were measured in air (Table 3-3) by time correlated single photon counting (TCSPC). All four complexes exhibited multi-exponential decay profiles which can be typical of Cu(I) complexes that display TADF, however the decay could also contain a mix of fluorescence and phosphorescence, or a combination of all three.10,71 In order to discern the radiative mechanism,   74 variable temperature emission lifetimes were collected over a temperature range of –196 to 28 °C (Figure 3-6 and Appendix Chapter 6:Listing B.5). As discussed in Section 1.3.6, at low temperatures a prompt fluorescence is not detected, instead a very bright, long-lived phosphorescence is seen. Upon increasing the temperature, the S1 state can be populated, resulting in an activation of the thermally delayed radiative pathway coupled with a drastic decrease in the emission lifetime and a blue-shift of the emission spectrum. The decay time of a system with thermally equilibrated excited state is determined by a Boltzmann-type relationship, and thus variable temperature excited-state lifetime experiments of copper(I) species displaying TADF show a sigmoidal curve.71 All species exhibit multi-exponential decay profiles, and when plotting the long lifetime components (which correspond to the thermalized emission of the triplet state), complexes Cu-DPS, Cu-DPSO2 and Cu-Me-DPSO2 exhibit a sigmoidal profile, indicating the TADF nature of their radiative mechanism. Cu-Me-DPS does not show this same pattern, however –196 °C is not always sufficient to isolate only the T1 emission (for that liquid helium temperatures are required) and so emission at low temperature in this instance is likely due to a mix of phosphorescence and TADF.  3.7 Time-resolved Photophysical Properties As discussed in Sections 1.3.7 and 3.6.3, luminescence from the T1 state can occur at 77 K, and ΔE(S1–T1) can be estimated from the emission spectra at low temperature and room temperature. The λmax value, however, is the most probable transition and therefore a more accurate determination for ΔE(S1–T1) can be performed if the zero–zero energy (E0–0) for each excited-state is calculated. E0–0 is the energy difference between the ground and excited-state, each taken at their zero vibrational levels.149 For copper(I) species, time-resolved photoluminescence spectroscopy   75 can be employed. Through monitoring emission at 77 K using different time regimes, the emission profiles of the fluorescence (ns, from S1 state) and phosphorescence (µs, from T1 state) can be collected. By using the onset values for each excited-state at 77 K, ΔE(S1–T1) can be more accurately determined.150  In conjunction with the Zysman-Colman and Samuel groups at the University of St. Andrews, Scotland, time-resolved photoluminescence spectroscopy was used to determine the ΔE(S1–T1) of Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 (Figure 3-7). Samples were prepared as neat powders and cooled to 77 K, following which the emission profiles at ns and µs regimes were collected. Time-resolved spectra in the time window of 100–200 ns (Cu-DPS, Cu-DPSO2 and Cu-Me-DPS) and 100–600 ns (Cu-Me-DPSO2) give the photoluminescence spectra for transitions emanating from the S1 state (grey traces), while spectra in the time windows 100–200 µs (Cu-Me-DPS) and 400–500 µs (Cu-DPS, Cu-DPSO2 and Cu-Me-DPSO2) give the photoluminescence spectra for T1 à S0 transitions. In all instances a shift to lower energy is seen when monitoring the longer time component, indicative of emission from a lower-energy state.   76  Figure 3-7 Time-resolved solid state photoluminescence spectra of (a) Cu-DPS; (b) Cu-DPSO2; (c) Cu-Me-DPS; (d) Cu-Me-DPSO2 at 77 K. Grey traces represent emission from the S1 state, and pink traces emission from the T1 state. Black lines indicate the determination of E0–0. Samples prepared as neat powders.   By determining E0–0 of each excited-state transition (represented by the black lines in Figure 3-7), ΔE(S1–T1) was calculated for each complex. Cu-DPS (0.03 eV), Cu-DPSO2 (0.01 eV), Cu-Me-DPS (0.07 eV) and Cu-Me-DPSO2 (0.12 eV) all have S1–T1 energy gaps far less than the 0.43 eV limit for RISC, and so at least one of the long-lived time components in the room   77 temperature emission lifetime data (Error! Reference source not found.) is likely due to a TADF mechanism. Comparisons of the time-resolved data to the steady-state data are presented in Table 3-5. All ΔE(S1–T1) values are close and, with the exception of Cu-DPSO, the trend remains the same whereby the energy gap becomes larger when functionalizing the slightly electron-donating methyl groups to the pyridyl rings. This shows that using λmax emission values from variable temperature steady-state emission data can give a good estimation as the S1–T1 energy gap.  Table 3-5 S1–T1 energy gaps of Cu(I) complexes as calculated using low temperature steady-state emission and low temperature time-resolved emission spectra.  Complex ΔE(S1–T1) [eV] Steady-State [a] Time-Resolved [b] Cu-DPS 0.04 0.03 Cu-DPSO2 0.05 0.01 Cu-Me-DPS 0.05 0.07 Cu-Me-DPSO2 0.08 0.12 [a] Calculated using λmax emission at room temperature and 77 K. All samples drop-cast on to quartz from MeOH. [b] E0–0 at room temperature and 77 K determined from emission onset, with ΔE(S1–T1) determined from E0–0 values. Samples prepared as neat powders.  3.8 Theoretical Calculations DFT and TD-DFT calculations were performed in order to elucidate the ground- and excited-states of the Cu(I) complexes and the character of the electronic transitions. The crystal structures discussed in Section 3.3.2 were used as starting points in the ground-state geometry   78 optimization. Calculated energy levels, oscillator strength (f) and orbital transitions are summarized in Table 3-6.  Table 3-6 TD-DFT calculated energy levels, oscillator strength and orbital transitions for Cu(I) complexes.  S1 T1  ΔE(S1–T1) [eV]   f E [eV] Contribution [H–L] Assignment E [eV] Contribution [H–L] Assignment Cu-DPS 0.0125 2.15 97.2% MLCT/LLCT 2.02 94.8% MLCT/LLCT 0.13 Cu-DPSO2 0.0012 1.06 97.6% MLCT/LLCT 0.98 98.4% MLCT/LLCT 0.08 Cu-Me-DPS 0.0160 2.72 93.3% MLCT/LLCT 2.59 88.8% MLCT/LLCT 0.13 Cu-Me-DPSO2 0.0071 1.33 99.1% MLCT/LLCT 1.31 96.7% MLCT/LLCT 0.02  For monometallic species Cu-DPS and Cu-DPSO2, the HOMO is largely located on the 3d orbital of the copper(I) center and the coordinating phosphorus ligand, while the LUMO is mainly distributed over the sulfur-bridged dipyridyl ligand. In the bimetallic complex Cu-Me-DPS, the HOMO is distributed over the 3d orbital of the copper(I) atom and the bridging POP ligand, with the LUMO located on the two dipyrdyl ligands, separated by the copper(I) atoms and the POP ligand. In the Cu-Me-DPSO2 dyad, the HOMO is again located across the copper(I) centre and POP ligand, with the LUMO distributed mainly over the bridging Me-DPSO2 ligand. The HOMO–LUMO transition can be assigned as metal-to-ligand charge transfer (MLCT).     79  Figure 3-8 Frontier orbital plots illustrating the HOMO and LUMO distributions for Cu(I) complexes.    80 The spatial separation of HOMO and LUMO is distinct, so a relatively small singlet–triplet energy gap ΔE(S1–T1) can be predicted, indicative of TADF. Indeed, TD-DFT calculations using the B3LYP method in singlet and triplet state optimized geometries give relatively small values of ΔE(S1–T1) (Table 3-6), much smaller than the critical value for TADF of 0.43 eV. The major transitions of singlet and triplet excited states can be assigned as HOMO–LUMO transitions.  Pronounced MLCT electronic transitions result in significant changes of nuclear coordinates in excited states.151,152 From the oscillator strength data in Table 3-6, Cu-DPSO2 has the smallest value (f = 0.0012), suggesting that this complex undergoes the most significant geometry distortion and therefore contributes to the low luminescence quantum yield observed. Additionally, calculations were performed using the wb97xd functional with effective core potential SDD for copper and 6-31G(d,p) for all other atoms. This method is regarded as being more sophisticated than those used above, however the results were not satisfying. For complexes Cu-DPS and Cu-Me-DPS ΔE(S1–T1) were both around 0.30 eV, with the emission energies of Cu-DPS under-estimated (S1: 2.75 ev; T1: 2.38 eV) and the emission energies of Cu-Me-DPS over-estimated (S1: 2.75 ev; T1: 2.38 eV) compared to the experimental data. The optimization of the singlet excited-state for Cu-DPSO2 and both the singlet and triplet excited-states for Cu-Me-DPSO2 did not converge. For these reasons the B3LYP calculations were used.  3.9 Conclusions Four heteroleptic Cu(I) complexes of the class [Cu(POP)(N^N)][BF4] are reported where N^N = sulfur-bridged dipyridyl ligands functionalized either with or without methyl substituents in the 6,6′-positions of the pyridyl rings and sulfide or sulfone oxidation states at the sulfur center. Single crystal structures of Cu-DPS and Cu-DPSO2 confirm that the complexes are coordinated   81 in a distorted tetrahedral geometry. However upon introducing additional bulk with methyl groups the steric constraints about the copper center become too great and dimers are formed. Cu-Me-DPS is a bimetallic species which adopts a trigonal geometry with a bridging POP ligand, whereas Cu-Me-DPSO2 (with additional binding sites at the sulfone oxygen atoms) forms a bimetallic species with the diimine moiety acting as a bridging ligand linking the two copper atoms. These structures are retained when the complexes are dissolved in solution. The complexes were all weakly emissive in solution, but luminance improved in the solid state, with PLQYs up to 0.14 for the most sterically constrained complex, Cu-Me-DPSO2. The sulfone complexes (Cu-DPSO2 and Cu-Me-DPSO2) both show a red-shifted emission profile when compared to their sulfide counterparts (Cu-DPS and Cu-Me-DPS) due to the electron-withdrawing nature of the oxygen atoms at the sulfur of the diimine ligand. All four complexes showed multiexponential decay profiles consisting of a prompt emission in the nanosecond regime, followed by longer components up to microseconds in length. Through the use of low-temperature time-resolved photoluminescence spectroscopy, in addition to computational methods, the S1–T1 energy gaps for each complex was determined and it was found that the decay profile is likely due to a TADF mechanism. Here we have shown that with changes in steric constraints and sulfur oxidation state a variety of different complex geometries can be obtained. Cu-Me-DPS represents the first examples of a dimetallic Cu(I) species bridged by a POP ligand, and Cu-Me-DPSO2 represents the first example of a sulfone-based N^O-bound diimine ligand. Cu-Me-DPSO2 could show promise for application in LEECs, and the novel binding modes seen in the complex could open avenues towards interesting, luminescent coordination polymers.    82 3.10 Experimental Details 3.10.1 General All experiments were conducted in atmospheric conditions unless otherwise stated. Solvents used for synthesis were of reagent grade and were used without any further purification. HPLC grade solvents were used for spectroscopic studies. 2-bromo-6-methylpyridine and niobium carbide were purcashed from Alfa Aesar, [Cu(MeCN)4][BF4] and 30% H2O2 solutions were purchased from Sigma-Aldrich, and thiourea from TCI. All reagents were used as received. Bis(pyridin-2-yl)sulfane (DPS) and 2,2′-sulfonyldipyridine (DPSO2) were synthesized using the methods described in Chapter 1:.  3.10.2 Spectroscopy 1H, 13C{1H}, 31P{1H}, COSY, NOESY, HSQC and HMBC NMR spectra were collected using a Bruker AV-400 spectrometer and referenced first to TMS and then to the residual protonated solvent peak. NMR solvents (Aldrich or Cambridge Isotope Laboratories) were used as received. Electrospray ionization mass spectrometry data was obtained using a Bruker Esquire LC ion trap mass spectrometer and elemental analysis determined using a Thermo Flash 2000 Elemental Analyzer. Infrared spectroscopy was performed on an attenuated total reflection (ATR) crystal using a Perkin-Elmer Frontier FT-IR spectrometer. UV-vis absorption spectra were recorded on a Varian-Cary 5000 UV-Vis-near-IR spectrophotometer. Fluorescence data were collected on a Photon Technology International (PTI) QuantaMaster 50 fluorimeter fitted with an integrating sphere, double excitation monochrometer and utilizing a 75 W Xe arc lamp as the source. Emission lifetime data were collected using a Horiba Yvon Fluorocube TCSPC apparatus. A 370 nm NanoLED was used as the excitation source, pulsing at a repetition rate of 100 kHz and   83 with a lifetime of 1.2 ns. Broadband emission was monitored by a CCD detector at the desired wavelengths using appropriate low pass filters. The lifetime data were fitted using the DAS6 Data Analysis software package. Solution spectra were collected using 1 cm2 quartz cells (Starna Cells) and spectra of the neat solids were collected by drop-casting from MeOH on to quartz slides (Ted Pella, Inc.).  3.10.3 X-Ray Crystallography SCXRD raw data were collected on a Bruker APEX DUO diffractometer using graphite monochromated MoKα radiation (λ = 0.71073) at 90 K and to a resolution of 0.77 Å. The data elaboration was performed with the software APEX2.153 Structures were solved using SUPERFLIP154 and refined using full-matrix least-squares on F2 within the CRYSTALS155 suite. Hydrogen atoms were initially refined with restraints on bond lengths and angles, after which the positions were used as the basis for a riding model.156 All non-hydrogen atoms were refined anisotropically.  3.10.4 Computational Details Density functional theory (DFT) calculations were carried out using the Gaussian 16 Rev.A03 suite of programs. The B3LYP functional123,124,157 with 6-31G** basis set (for C, H, S N P and O atoms) and the LANL2DZ effective-core pseudopotential (for Cu)125 was employed to optimize the geometry of the ground-states (S0) of all complexes with input structures extracted from the X-ray crystallographic data. Time-dependent DFT (TD-DFT) calculations were also performed to understand the nature of first singlet (S1) and triplet excited states (T1).128,129,158    84 3.10.5 Synthesis bis(6-methylpyridin-2-yl)sulfane (Me-DPS) 2-Bromo-6-methylpyridine (0.66 mL, 5.81 mmol, 1.0 eq.) was added to a solution of thiourea (0.21 g, 2.76 mmol, 0.48 eq.) in 100 mL of EtOH under an N2 atmosphere and heated to reflux for 36 h. The reaction mixture was cooled to room temperature, the solvent removed under reduced pressure and the resultant residue dissolved in CH2Cl2 (20 mL) and washed with water (3 × 20 mL) and brine (1 × 50 mL). The organic layer was dried with MgSO4, filtered and the CH2Cl2 removed in vacuo. The crude product was purified over silica gel using a solvent gradient from 1:0 CH2Cl2:MeOH to 9:1 CH2Cl2:MeOH, yielding a brown oil (0.558 g, 2.72 mmol, 99%). 1H NMR (400 MHz, CD2Cl2) δ 7.50 (t, J = 7.7 Hz, 2H, H5), 7.19 (dt, J = 7.9, 0.8 Hz, 2H, H6), 7.02 (ddd, J = 7.6, 0.9, 0.5 Hz, 2H, H4), 2.50 (s, 6H, HMe). 13C{1H} NMR (CD2Cl2, 100 MHz) δ 159.82 (C1), 156.80 (C3), 137.69 (C5), 123.05 (C6), 121.67 (C4), 24.62 (CMe). ESI-HR MS: Calcd. for C12H13N2S: 217.0799; Found: 217.0798 [M+H]+  6,6'-sulfonylbis(2-methylpyridine (Me-DPSO2) A round-bottomed flask was charged with Me-DPS (0.167 g, 0.772 mmol, 1.0 eq.) and NbC (0.020 g, 0.193 mmol, 0.25 eq.). 10 mL EtOH and 30% aqueous solution of H2O2 (1.09 mL, 11.6 mmol, 15.0 eq.) were added and the reaction mixture was heated to 60 ºC for 12 h. The mixture was cooled to room temperature, diluted with saturated Na2S2O3 (10 mL) and extracted with CH2Cl2 (3 × 10 mL). The combined organics were dried over MgSO4, filtered and the solvent removed in vacuo yielding a white powder that required no further purification (0.190 g, 0.765 mmol, 99%). 1H NMR (400 MHz, CD2Cl2) δ 8.07 (d, J = 7.7 Hz, 2H, H6), 7.85 (t, J = 7.8 Hz, 2H, H5), 7.36 (d, J = 7.8 Hz, 2H, H4), 2.51 (s, 6H, HMe). NSN123456CH3 CH3MeO ONSN123456CH3 CH3Me  85 13C{1H} NMR (101 MHz, CD2Cl2) δ 160.67 (C1), 156.90 (C3), 138.43 (C5), 127.71 (C4), 121.54 (C6), 24.53 (CMe). IR (solid): ṽ (σ SO2) 1450, 1313 cm–1. ESI-HR MS: Calcd. for C12H12N2O2SNa: 271.0517; Found: 271.0520 [M+Na]+.   [Cu(POP)(DPS)][BF4] (Cu-DPS) [Cu(MeCN)4]BF4 (0.092 g, 0.292 mmol, 1.0 eq.) and POP (0.158 g, 0.292 mmol, 1.0 eq.) were dissolved in 12 mL CH2Cl2 and stirred at room temperature for 2 h, after which DPS (0.055 g, 0.292 mmol, 1.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated in vacuo, the product precipitated with Et2O, filtered and washed with Et2O, yielding an off-white powder that required no further purification (0.214 g, 0.244 mmol, 84%). 1H NMR (400 MHz, CD3OD) δ 8.16 – 8.10 (m, 2H, H6), 7.90 – 7.81 (m, 4H, H3,5), 7.39 – 7.19 (m, 22H, H8,9,10,14), 7.14 – 7.08 (m, 2H, H4), 7.05 (td, J = 7.6, 1.1 Hz, 2H, H13), 6.99 (dtd, J = 8.2, 2.5, 1.0 Hz, 2H, H15), 6.82 (dtd, J = 7.8, 4.0, 1.6 Hz, 2H, H12). 13C{1H} NMR (101 MHz, CD3OD) δ 159.11 (C16), 154.90 (C1), 152.31 (C6), 140.64 (C5), 135.02 (C12), 134.74 (t, J = 8.1 Hz, C8), 133.41 (C14), 132.22 (t, J = 16.5 Hz, C7), 131.47 (C10), 130.29 (C3), 129.98 (C9), 126.30 (C13), 125.92 (t, J = 14.4 Hz, C11), 125.41 (C4), 121.69 (C15). 31P{1H} NMR (162 MHz, CD3OD) δ –14.20. ESI-HR MS Calcd. for C46H36N2OP2S63Cu: 789.1320; Found: 789.1323 [M–BF4]+. Elemental Analysis: Calcd. C46H36BF4N2OP2SCu: C, 62.99; H, 4.14; N, 3.19; S, 3.65. Found: C, 62.71; H, 4.26; N, 3.15; S, 3.64.   CuPOPNNSBF412314 11107 4 5613891215 16  86 [Cu(POP)(DPSO2)][BF4] (Cu-DPSO2)  A solution of [Cu(MeCN)4]BF4 (0.079 g, 0.250 mmol, 1.0 eq.) and POP (0.134 g, 0.250 mmol, 1.0 eq.) in 12 mL CH2Cl2 was stirred at room temperature for 2 h, after which DPSO2 (0.055 g, 0.250 mmol, 1.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated under reduced pressure, the product precipitated with Et2O, filtered and washed with Et2O, giving a yellow powder (0.210 g, 0.231 mmol, 92%). 1H NMR (400 MHz, CD3OD) δ 8.52 (d, J = 4.8 Hz, 2H, H6), 8.36 (d, J = 7.9 Hz, 2H, H3), 8.14 (td, J = 7.8, 1.6 Hz, 2H, H5), 7.54 (ddd, J = 7.7, 4.9, 1.2 Hz, 2H, H4), 7.46 – 7.37 (m, 4H, H10), 7.31 (ddt, J = 10.2, 7.3, 2.7 Hz, 18H, H8,9,14), 7.02 (td, J = 7.6, 1.1 Hz, 2H, H13), 6.97 (dtd, J = 8.2, 2.5, 1.0 Hz, 2H, H15), 6.75 (dtd, J = 7.9, 4.1, 1.6 Hz, 2H, H12). 13C{1H} NMR (101 MHz, CD3OD) δ 159.13 (t, J = 5.9 Hz, C16), 157.31 (C1), 152.08 (C6), 140.67 (C5), 135.28 (C12), 134.81 (t, J = 8.3 Hz, C8), 133.31 (C14), 132.00 (t, J = 17.2 Hz, C7), 131.53 (C10), 130.02 (t, J = 4.9 Hz, C9), 129.64 (C4), 126.13 (C13), 125.54 (C3), 125.31 (t, C11), 121.44 (C15). 31P{1H} NMR (162 MHz, CD3OD) δ –17.56.  ESI-HR MS Calcd. for C46H36N2O3P2S63Cu: 821.1225; Found: 821.12183 [M–BF4]+. Elemental Analysis: Calcd. C46H36BF4N2O3P2SCu: C, 60.77; H, 3.99; N, 3.08; S, 3.53. Found: C, 60.44; H, 4.24; N, 2.76; S, 3.41.      CuPOPNNSBF412314 11107 4 5613891215 16 OO  87 [(Me-DPS)Cu(POP)Cu(Me-DPS)][BF4]2 (Cu-Me-DPS) [Cu(MeCN)4]BF4 (0.176 g, 0.559 mmol, 2.0 eq.) and POP (0.140 g, 0.278 mmol, 1.0 eq.) were dissolved in 15 mL CH2Cl2 and stirred at room temperature for 2 h, after which Me-DPS (0.121 g, 0.559 mmol, 2.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated in vacuo, the product precipitated with Et2O, filtered and washed with Et2O, yielding an off-white powder that required no further purification (0.290 g, 0.228 mmol, 82%). 1H NMR (400 MHz, CD3OD) δ 7.91 (br s, 8H, H4,5), 7.51 (br m, 2H, H6) 7.39 (m, 8H, H10,14), 7.28 (t, J = 7.5 Hz, 8H, H9), 7.14 (m, 10H, H8,15), 7.02 (t, J = 7.5 Hz, 2H, H13), 6.64 (dtd, J = 5.9, 4.1, 1.5 Hz, 2H, H12), 2.21 (br s, 12H, HMe). 13C{1H} NMR (101 MHz, CD3OD) δ 159.33 (C16), 135.54 (C12), 134.44 (t, C8), 133.63 (C14), 131.74 (C10), 131.30 (t, J = 19.6 Hz, C7), 130.14 (t, J = 5.0 Hz, C9), 126.31 (C13), 124.36 (C11), 121.40 (C5). 31P{1H} NMR (162 MHz, CD3OD) δ –15.83. Elemental Analysis: Calcd. C60H52B2F8N4OP2S2Cu2: C, 56.66; H, 4.12; N, 4.41; S, 5.04. Found: C, 56.64; H, 4.12; N, 4.40; S, 5.10.        PPhCuNNSMeMeOP CuNNSMeMePhPh[BF4]2123456107891112 13141516Me  88 [(POP)Cu(Me-DPSO2)Cu(POP)][BF4]2 (Cu-Me-DPSO2) [Cu(MeCN)4]BF4 (0.064 g, 0.202 mmol, 2.0 eq.) and POP (0.109 g, 0.202 mmol, 1.0 eq.) were dissolved in 10 mL CH2Cl2 and stirred at room temperature for 2 h, after which Me-DPSO2 (0.025 g, 0.101 mmol, 1.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated in vacuo, the product precipitated with Et2O, filtered and washed with Et2O, yielding a yellow powder that required no further purification (0.157 g, 0.097 mmol, 96%). 1H NMR (400 MHz, CD3OD) δ 8.00 (d, J = 7.5 Hz, 4H, H4,5), 7.53 (d, J = 7.2 Hz, 2H, H6), 7.47 – 7.40 (m, 8H, H10), 7.39 – 7.27 (m, 36H, H8,9,14), 7.00 (t, J = 7.4 Hz, 8H, H13,15), 6.69 (dtd, J = 8.0, 4.0, 1.6 Hz, 4H, H12), 2.26 (s, 6H, HMe). 13C{1H} NMR (101 MHz, CD3OD) δ 13C{1H} NMR (101 MHz, CD3OD) δ 161.65 (C1), 159.34 (t, C16), 140.50 (C5), 135.41 (C12), 134.77 (t, J = 8.3 Hz, C8), 133.25 (C14), 132.20 (t, J = 17.6 Hz, C7), 131.58 (C10), 130.02 (t, J = 4.9 Hz, C3,6,9), 126.02 (C13), 124.96 (t, J = 15.5 Hz, C11), 122.78 (C4), 121.45 (C15), 24.28 (CMe). 31P{1H} NMR (162 MHz, CD3OD) δ –18.44. Elemental Analysis: Calcd. C84H68B2F8N2O4P4SCu: C, 62.04; H, 4.22; N, 1.72; S, 1.97. Found: C, 62.02; H, 4.55; N, 1.60; S, 1.93.    POPPhPhPh PhCu[BF4]2NSOMeNMePOPCu O1234561078914 1113 1215 16Me  89 CHAPTER 4: Thermochromic Solid state Emission of Dipyridyl Sulfoxide Copper(I) Complexes  4.1 Introduction Luminescent materials have seen a surge of interest within the last two decades for their use in a broad range of applications such as biological labelling,159 sensing,160,161 lighting and displays.26,162 These applications often involve emission in the solid state, with organic light-emitting diodes (OLEDs) and light-emitting electrochemical cells (LECs) being two of the more common implementations. It has been shown that the emissive properties (such as colour, emission lifetime etc.) in these devices can be manipulated, and for metal complexes this can be achieved through ligand design27,88,89,163 or via intermolecular interactions.164-168 Copper(I) complexes have been shown to be excellent candidates for low-cost, solid state emitters due to their high Earth-abundance and low cost, coupled with their ability to luminesce brightly at room temperature,71 resulting in complexes with interesting photophysical characteristics. Nishihara et al. have reported that stimuli-responsive ring rotation in copper(I) complexes can be used to switch physical properties, including a dual luminescence.169,170 In addition, copper(I) complexes have shown emissive thermochromic properties as clusters,54,171-175 in coordination polymers176 and in trinuclear species177 with a high sensitivity to temperature, their environment and the rigidity of their medium.171 The photophysics of [Cu(disphosphine)(diimine)]+ complexes have received considerable attention. These complexes exhibit an absorption in the visible light region due to a metal-to-ligand charge transfer (MLCT) transition resulting in an intense luminescence.148,178-183 In particular,   90 Cu(I) complexes with an (oxydi-2,1-phenylene)bis(diphenylphosphine) (POP) diphosphine ligand have been extensively researched, and the introduction of a bulky diimine ligand into the coordination sphere has been shown to increase the excited state lifetime of the complex by preventing structural relaxation from tetrahedral to square-planar geometry and inhibiting solvent coordination. The Wolf group has shown that the degree of sulfur oxidation can have significant effects on the electronic properties of a system, both in conjugated organic molecules77-80 and inorganic complexes (Chapter 1:. In this work, sulfur-bridged diimine ligands based on di(pyridine-2-yl)sulfanes are used to probe how ligand binding mode and intermolecular interactions affect emission behaviour in Cu(I) complexes. Similar to 2,2'-bipyridine, sulfur-bridged diimine ligands can bear differing substituents at the pyridyl rings. However, the addition of the sulfur bridge in the 2,2' positions both increases the bite angle of the chelator and gives an additional site for further electronic tuning. On oxidizing the sulfur to sulfoxide, a third binding site becomes available leading to a variety of binding modes, which, when coupled with the tetrahedral geometry of the bridge and the electron-withdrawing nature of the S=O bond, gives flexibility to the ligand allowing coordination changes at copper. Herein, we report two heteroleptic Cu(I) complexes Cu-DPSO and Cu-Me-DPSO containing sulfoxide ligands based on di(pyridine-2-yl)sulfanes (DPSO and Me-DPSO). The structural properties of the complexes were characterized using NMR experiments, X-ray crystallography, mass spectrometry and elemental analysis, and the photophysical behavior was studied as a function of temperature. Cu-Me-DPSO was found to luminesce either yellow or orange depending on morphology and temperature.  This behavior is attributed to a change in   91 excited state coordination geometry, fundamentally differing from thermochromism in other copper complexes involving emission from Cu–Cu states.  4.2 Ligand Synthesis and Characterization Scheme 4-1 Synthesis of Cu(I) species Cu-DPSO and Cu-Me-DPSO.   The diimine sulfide precursors DPS and Me-DPS were prepared via a nucleophilic aromatic substitution of 2-bromopyridine or 2-bromo-6-methylpyridine using thiourea (Scheme 4-1).96 The desired sulfoxide proligands (DPSO and Me-DPSO) were then synthesized by oxidizing the appropriate sulfide compound with 30% H2O2 in the presence of glacial CH3COOH and purified using column chromatography.97 Structural characterization was performed using 1H and 13C NMR spectroscopy (including COSY, HSQC and HMBC experiments), infrared spectroscopy and high resolution mass spectrometry.    92 4.3 Synthesis and Characterization of [Cu(P^P)(N^N)[BF4] Complexes Cu(I) complexes (Cu-DPSO and Cu-Me-DPSO) were prepared via reaction of [Cu(MeCN)4][BF4] with POP, followed by the addition of the appropriate dipyridyl sulfoxide proligand.76,141 The products both precipitated with diethyl ether yielding yellow powders that required no further purification (Scheme 4-1). All complexes were characterized using 1H, 13C{1H} and 31P{1H} NMR spectroscopy, high resolution mass spectrometry, X-ray crystallography, powder XRD and elemental analysis. HR-ESI spectra show peaks corresponding to [M–BF4]+ and exhibit a characteristic copper isotope splitting pattern. 1H and 13C signals were assigned using COSY, NOESY, HSQC and HMBC NMR experiments.  4.3.1 Nuclear Magnetic Resonance (NMR)  Figure 4-1 1H NMR spectra of Cu-DPSO and Cu-Me-DPSO in CD3CN at 25 ºC (400 MHz).  The room-temperature NMR spectra of both Cu-DPSO and Cu-Me-DPSO show clear, defined signals with no evidence for exchange or dissociation occurring in CD3CN (Figure 4-1). Sulfoxide ligands DPSO and Me-DPSO can bind to metals in three possible binding modes: in a   93 bidentate fashion via either both pyridyl rings (N,N), through one pyridyl ring and one oxygen (N,O) or through a pyridyl ring and the sulfur lone pair (N,S). For both Cu-DPSO and Cu-Me-DPSO, the NMR spectra are consistent with a symmetric N,N bound species with equivalent pyridyl rings. The POP ligand is bound in the usual bidentate fashion chelating via the phosphorus groups in both complexes, with a single peak in each case in the 31P{1H} NMR spectrum, with quadrupolar broadening attributed to the copper atoms. On cooling a CD3CN solution of Cu-Me-DPSO to –40 ºC no change in the 1H NMR spectrum was observed (Figure 4-2), indicating that over this temperature range the complex is either in fast equilibrium or present only as a single structure.  Figure 4-2 Low temperature 1H NMR spectra of Cu-Me-DPSO in CD3CN from 25 ºC to –40 ºC (400 MHz).    94 4.3.2 Single Crystal X-Ray Diffraction  Figure 4-3 Solid state structures of (a) Cu-DPSO (N,N and N,O coordination modes, same unit cell). (b) Cu-Me-DPSO (N,O coordination mode). Ellipsoids are plotted at the 50% probability level. Hydrogen atoms and solvent molecules are removed for clarity.  X-ray quality crystals of each complex were obtained through the slow diffusion of Et2O into a CH2Cl2 solution (Figure 4-3). Structural studies show the binding mode in these mononuclear copper(I) complexes in the solid state in the presence of the bulky POP chelating ligand. Cu-DPSO comprises the less bulky DPSO ligand and crystallizes in the space group P-1 with the asymmetric unit containing two distinct Cu(I) complexes, each with a different coordination mode for the DPSO ligand (Figure 4-3(a); N,N and N,O binding modes). The packing is stabilized by intermolecular hydrogen bonding between the sulfoxide oxygen of Cu-DPSO (N,N) and the α-proton of the non-bound pyridyl ring of Cu-DPSO (N,O), with a bond distance of 2.442 Å.  (a) (b)Cu-DPSO Cu-Me-DPSON,N-bound N,O-bound N,O-bound  95 Table 4-1  Selected bond lengths (Å) and angles (º) of Cu-DPSO (N,N-bound), Cu-DPSO (N,O-bound) and Cu-Me-DPSO (N,O-bound).  Cu-DPSO (N,N-bound) Cu-DPSO (N,O-bound) Cu-Me-DPSO (N,O-bound) Cu1–N1 2.092(2) [a] 2.066(2) 2.0687(18) Cu1–N12 2.092(2) – – Cu1–O1 – 2.1777(18) 2.1975(17) Cu1–P1 2.2643(7) [b] 2.2478(7) 2.2380(6) Cu1–P2 2.2753(7) [c] 2.2851(7) 2.2809(5) N11–Cu1–N12 91.82(8) – – N1–Cu1–O1 – 81.16(8) 83.81(7) P1–Cu1–P2 112.42(2) [d] 116.14(3) 114.73(2) [a] Cu11–N11. [b] Cu11–P11. [c] Cu11–P12. [d] P11–Cu11–P12.  Only the N,O binding mode is found in the crystal structure of Cu-Me-DPSO (Figure 4-3(b)). In this conformation one of the pyridyl rings of Me-DPSO is rotated, orienting the methyl group away from the copper center. Cu-Me-DPSO contains a diimine ligand with methyl groups in the 6,6'-positions, resulting in greater steric strain near the metal center; in the N,N binding mode the two methyl groups will point towards the phenyl groups of the POP ligand. Binding through both pyridyl nitrogen atoms therefore appears to be disfavored in the crystal. This N,O-bound structure is not consistent with the solution structure determined by NMR spectroscopy; the presence of two different structures in the solution and solid state suggests that the energy difference between N,N and N,O binding modes is relatively small. The N,S binding mode was not observed with either DPSO or Me-DPSO ligands in the solid state or in solution, presumably due to the greater strain that would result from an acute N–Cu–S bond angle.   96 Selected bond lengths and angles are reported in Table 4-1. Structural analysis shows that the copper(I) complexes have a d10 tetrahedral structure and avoid mutual repulsion of the POP and diamine ligands by adopting a distorted tetrahedral geometry.184 The P–Cu–P bite angle was anticipated to be quite large due to the presence of the bulky PPh3 groups, whereas the N–Cu–N and N–Cu–O angles were expected to be smaller due to the higher steric constraints of the diimine ligands. Structural studies reveal that in all the complexes the P–Cu–P angles are in the range 112–116º while the N–Cu–N angle is 91.82º in N,N-bound Cu-DPSO and N–Cu–O angles are 81.16º and 83.81º for N,O-bound Cu-DPSO and Cu-Me-DPSO respectively. The three structures observed have similar Cu–P, Cu–N and Cu–O bond lengths, which are comparable to those measured in other related mononuclear copper(I) complexes.75,76,136,185 The shortest inter-complex Cu(I)–Cu(I) bond distances were found to be 9.733 Å and 10.872 Å for Cu-DPSO and Cu-Me-DPSO, respectively. These values exceed the sum of the van der Waals radius of copper(I) (2.80 Å)186 and thus metal–metal interactions are not seen.    97 4.3.3 Powder X-Ray Diffraction (pXRD)  Figure 4-4 (a) pXRD of Cu-DPSO compared to the simulated powder pattern generated from the single crystal structure (N,N and N,O bound). (b) pXRD spectra of amorphous a-Cu-Me-DPSO (grey) heated to 180 ºC (teal) and then cooled to room temperature, retaining crystalline species c-Cu-Me-DPSO (pink)  Powder X-ray diffraction (pXRD) scans were recorded for both complexes at room temperature, with all powders isolated via rapid precipitation through addition of Et2O to a concentrated CH2Cl2 solution of the complex. Cu-DPSO shows crystallinity with diffraction peaks matching those calculated from the diffraction data of the single crystal X-ray structure (Figure 4-4(a)), confirming that the isolated powder contains a mixture of complexes bound both N,N and N,O through the diimine ligand. The isolated methyl-substituted complex, however, is amorphous as isolated at room temperature (a-Cu-Me-DPSO, Figure 4-4(b)).    98  Figure 4-5 (a) pXRD of Cu-DPSO during heating to 180 ºC. (b) pXRD of Cu-DPSO during cooling to –180 ºC. (c) pXRD comparing c-Cu-Me-DPSO (pink) to the diffraction pattern generated from the single crystal data of N,O-bound Cu-Me-DPSO (black). a-Cu-Me-DPSO: amorphous state. c-Cu-Me-DPSO: crystalline state. (d) pXRD cycle of c-Cu-Me-DPSO cooled from room temperature to –180 ºC and back to room temperature.  The yellow powdered sample of Cu-DPSO was first heated to 180 ºC, then cooled to –180 ºC, and heated back to room temperature with no appreciable changes in the diffractogram (Figure 4-5(a) and Figure 4-5(b) respectively), indicating the structure was retained throughout this treatment. On heating a-Cu-Me-DPSO to 180 ºC however, the sample crystallized and the   99 crystallinity was maintained on cooling back to room temperature (c-Cu-Me-DPSO, Figure 4-4(b)) giving a yellow-orange powder. The diffraction peaks observed for this sample do not match those predicted from the single crystal X-ray structure data, indicating that this structure does not contain ligand Me-DPSO bound in the N,O form (Figure 4-5(c)). On cooling c-Cu-Me-DPSO to –180 ºC a small shift in the diffraction pattern is observed indicating a change in the crystalline lattice (Figure 4-5(d)).  The change could be a contraction in the unit cell as on warming to room temperature the diffraction pattern returned to that originally seen. NMR studies showed that both a-Cu-Me-DPSO and c-Cu-Me-DPSO exhibit identical 1H NMR spectra when dissolved in CD3CN or non-coordinating CD2Cl2, indicating that in solution only one state is preferred and that no reactions occur during the heating process.    100 4.3.4 Differential Scanning Calorimetry (DSC)  Figure 4-6 DSC data of (a) Cu-DPSO and (b) Cu-Me-DPSO. Both samples were scanned over a temperature range of 25–180 ºC with three cycles.  Differential scanning calorimetry (DSC) of Cu-DPSO over a temperature range of 25–180 ºC (Figure 4-6(a)) shows no marked change over each of three scans, however a-Cu-Me-DPSO over the same temperature range shows a strong endothermic feature in the first heating scan with an onset at 110 ºC which is attributed to the crystallization of a-Cu-Me-DPSO to c-Cu-Me-DPSO (Figure 4-6(b)). Once crystallized, c-Cu-Me-DPSO exhibits a glass transition (Tg ~100 ºC) on further heating cycles. The crystallization may be due to a flattening distortion of the complex towards a pseudo-tetrahedral geometry, as it has been noted that a slight red-shift in the absorption   101 of ground-state Cu(I) complexes accompanies changes of this type,187,188 consistent with the slight change in colour of Cu-Me-DPSO from yellow to yellow-orange.  4.3.5 Fourier Transform Infrared Spectroscopy (FT-IR)  Figure 4-7 FT-IR spectra of a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) as neat solids at 25 ºC.  Infrared spectroscopy provides a useful handle for elucidating the structural transformation as the sulfoxide functional group has a strong stretching band. If the binding mode were changed from N,N to N,O upon crystallization, the SO stretch would be expected to shift to a lower frequency upon coordination to the metal center. However in comparing the spectra of a-Cu-Me-DPSO to c-Cu-Me-DPSO (Figure 4-7) the sulfoxide stretching band (1050 cm–1) does not shift, therefore it is concluded that the complex maintains the N,N binding motif on crystallization.    102 4.4 Solution-state Photophysical Properties 4.4.1 Pro-ligand Photophysical Properties  Figure 4-8 (a) UV-vis absorption spectra of DPSO (blue trace) and Me-DPSO (pink trace). (b) Normalized excitation (dashed line) and emission (solid line) spectra of DPSO (blue trace) and Me-DPSO (pink trace). λex = 310 nm. Solutions of CH2Cl2 at 25 ºC, concentrations ~1 × 10–5 M.   The absorption spectra of sulfoxide pro-ligands DPSO and Me-DPSO were recorded in ~1 × 10–5 M solutions of CH2Cl2 (Figure 4-8(a) and Table 4-2) with features at energies higher than 300 nm corresponding to π–π* transitions. Both compounds exhibit blue fluorescence with λmax = 426 and 422 nm for DPSO and Me-DPSO, respectively, with each compound also displaying a higher energy shoulder in the near-UV (Figure 4-8(b)). The excitation spectra of both pro-ligands shows that the emission stems from a π–π* transition with λmax = 311 nm.    103 4.4.2 Cu(I) Complex Photophysical Properties  Figure 4-9 (a) Absorption and emission spectra of Cu-DPSO (blue) and Cu-Me-DPSO (pink) in sparged (Ar) CH2Cl2 solutions (concentrations ~3 × 10–5 M). Emission spectra are corrected for absorbance at the excitation wavelength (lex = 380 nm). (b) Absorption spectra of a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) in CH2Cl2 at 25 ºC (concentrations ~5 × 10–5 M).  Absorption spectra of Cu-DPSO and Cu-Me-DPSO in CH2Cl2 are shown in Figure 4-9(a) and the photophysical data for the copper(I) complexes are summarized in Table 4-2. In CH2Cl2, Cu-DPSO and Cu-Me-DPSO both show intense absorption bands below 300 nm which are assigned to π–π* ligand centered (LC) transitions of the diimine ligands and the POP ligand. Additional shoulders between 300–400 nm not observed in the free ligands are assigned to d(Cu)–π*(N,N) metal-to-ligand charge-transfer (MLCT) transitions mixed with π(PP)–π*(N,N) ligand-to-ligand charge-transfer (LLCT) transitions,148 and the lack of features near 450 nm indicates that the bis-diimine complex is not present in solution. The absorptions of  a-Cu-Me-DPSO and c-Cu-Me-DPSO in CH2Cl2 were compared and found to be identical (Figure 4-9(b)) indicating that the ground state structure of both complexes reverts to the same geometry in solution.   104  Table 4-2 Photophysical data for compounds in CH2Cl2 (~1 × 10–5 M).  Absorption Emission Compound λmax (nm) ε (M−1 cm−1) λem (nm) PLQY τem [c] DPSO 266 1.3 × 104 426 [a] – N/A [d] Me-DPSO 272 1.4 × 104 422 [a] – N/A [d] Cu-DPSO 250, 355 2.3 × 104, 3.0 × 103 532 [b] <0.01 3 ns (63%), 11 ns (37%) Cu-Me-DPSO 271, 365 2.7 × 104, 1.9 × 103 422 [b] 0.09 4 ns  [a] lex = 310 nm. [b] lex = 380 nm. [c] lex = 370 nm. [d] Reliable emission lifetimes of the pro-ligands could not be collected due to lack of absorption at 370 nm.  Emission spectra of both complexes were recorded in CH2Cl2 after sparging with Ar for 30 min. Cu-DPSO displays a very weak, broad yellow emission when excited at 380 nm (λmax = 532 nm, PLQY < 0.01) indicative of a MLCT-based transition (Figure 4-9(a)). The low intensity in solution is attributed to a flattening distortion in the vibrationally relaxed state, typical of copper(I) complexes of this type.180 Cu-Me-DPSO shows a stronger blue emission upon excitation at 380 nm (λmax = 422 nm, PLQY = 0.09), significantly blue-shifted from that of Cu-DPSO. The more intense emission is due to the steric bulk afforded by the methyl groups on the diimine ligand hindering the Jahn-Teller distortion and reducing the level of quenching. This emission profile is similar to that of pro-ligand Me-DPSO in solution, and may therefore be due to an inter-ligand charge transfer (ILCT) state. Both complexes have emission lifetimes consistent with fluorescence, and no change was observed in the photophysical properties before and after sparging with Ar.   105  Figure 4-10 Excitation spectra of Cu-DPSO and Cu-Me-DPSO in CH2Cl2 at 25 ºC (concentrations ~3 × 10–5 M).  [Cu(PP)(NN)]+ complexes are known to be labile in solution141 so it is possible that the emission seen in Cu-Me-DPSO could be from the formation of [Cu(POP)2]+, however this homoleptic complex emits at ~480 nm in CH2Cl2,144 60 nm red-shifted from the emission of Cu-Me-DPSO. [Cu(POP)2]+ also does not display absorption at energies lower than 350 nm in CH2Cl2, while Cu-Me-DPSO shows a strong MLCT band centered at 365 nm. The excitation spectrum of Cu-DPSO contains both ligand-based (λmax = 292 nm) and MLCT-based (λmax = 350 nm) contributions, while the spectrum of Cu-Me-DPSO contains primarily one MLCT transition, with λmax = 380 nm and a small shoulder at ~335 nm (Figure 4-10).    106 4.5 Cu(I) Complex Solid state Photophysical Properties 4.5.1 Photoluminescence Spectroscopy  Figure 4-11 (a) Room temperature solid state emission of Cu-DPSO (blue), a-Cu-Me-DPSO (grey) and c-Cu-Me-DPSO (pink) drop-cast from MeOH on quartz. (lex = 380 nm); (b) Photograph of the three complexes under UV hand lamp irradiation. (lex = 365 nm)  Thin films of the neat powders were prepared by drop-casting onto quartz from MeOH. When irradiated with UV light (λex = 380 nm) Cu-DPSO gives green-yellow luminescence (λem = 524) and a-Cu-Me-DPSO gives yellow luminescence (λem = 542 nm at 25 ºC) with broad, unstructured emission bands attributed to a MLCT (Figure 4-11(a) and Table 4-3).71,182,189 However, when a thin film of the crystallized sample c-Cu-Me-DPSO was irradiated, it showed a pronounced bathochromic shift in emission with λem = 658 nm upon excitation. A large N-Cu-N bite angle (in comparison to that in [Cu(POP)(bpy)]+ systems76) coupled with the additional steric constraints of methyl groups in the 6- and 6'-positions on the diimine ligand of Cu-Me-DPSO creates a bulky environment in conjunction with the POP ligand, which when rapidly precipitated from solution results in amorphous a-Cu-Me-DPSO being isolated. Due to the similar emission   107 energy of this compound when compared to Cu-DPSO, it is likely that the Cu(I) center in a-Cu-Me-DPSO possesses a tetrahedral environment aided by the bulky methyl groups on the diimine ligand. However, once heated above ~100 ºC a structural change occurs which triggers a red-shift in the emission wavelength. c-Cu-Me-DPSO likely acquires a more flattened distortion about the metal center as it is known that a bathochromic shift in the emission of Cu(I) result from pseudo-Jahn-Teller distortions.190  Table 4-3 Solid state emission data for thin films of Cu-DPSO, a-Cu-Me-DPSO and c-Cu-Me-DPSO.  Solid state Emission [a] λem at 25 ºC (nm) λem at –196 ºC (nm) Cu-DPSO 524 534 a-Cu-Me-DPSO 542 558 c-Cu-Me-DPSO 658 556 [a] Thin films drop-cast from MeOH on to quartz. (lex = 380 nm)  4.5.2 Low-Temperature Photophysical Properties  Figure 4-12 Variable-temperature solid state emission spectra of (a) Cu-DPSO and (b) a-Cu-Me-DPSO from –196 to 27 ºC. Thin films dropcast on quartz from MeOH. (λex = 380 nm)   108  Low-temperature solid state emission data helps elucidate the cause of the large red-shift in emission from a-Cu-Me-DPSO to c-Cu-Me-DPSO. The solid state emission spectra of a thin film of Cu-DPSO drop-cast from MeOH on a quartz slide were collected from –196 to 25 ºC (Figure 4-12(a)). The emission intensity decreased and a blue-shift in λmax from 534 to 524 nm was observed on warming from –196 ºC to room temperature. Cu(I) complexes are known to often have a small ΔE(S1–T1) and the T1 state can be populated even at liquid nitrogen temperatures.10 The observed blue-shift is therefore a consequence of thermal activation of the energetically higher-lying S1 state and is typical of the behavior of [Cu(PP)(NN)]+ complexes, with a ΔE(S1–T1) of 0.04 eV seen for Cu-DPSO.   Figure 4-13 (a) Variable temperature solid state emission spectra of c-Cu-Me-DPSO from –196 to 27 ºC (λex = 380 nm). (b) Variable temperature, solid state emission lifetimes of c-Cu-Me-DPSO at –196 and 25 ºC. To maintain accuracy, a 50 µs time range was used for –196 ºC and 20 µs for 25 ºC. Longest time component for each complex reported (λex = 370 nm). Thin films dropcast on quartz from MeOH.     109 Emission spectra of a thin film of a-Cu-Me-DPSO were collected from –196 ºC to 25 ºC and a blue-shift of 558 to 542 nm observed, similar to the results for Cu-DPSO (Figure 4-12(b)). In contrast, c-Cu-Me-DPSO in the solid state exhibited significant emission thermochromism on warming from –196 ºC to 25 ºC (Figure 4-13(a)), with a red-shift of 102 nm from a high energy (HE) band at 556 nm to a low energy (LE) band at 658 nm. This thermochromic emission behavior was found to be reversible and rapid. As a thin film cooled in liquid nitrogen, c-Cu-Me-DPSO gives a bright yellow emission when irradiated with a handheld UV lamp, and then when heated (either rapidly with a Varitemp heat gun or left to slowly warm to room temperature) gives orange emission. This process can be repeatedly cycled with no noticeable change to the emission colour thus demonstrating reversible thermochromism. Similar behavior has been observed in transition metal complexes where a bathochromic shift accompanies occupancy of an energetically lower-lying state (such as a T1 state) that reduces the HOMO–LUMO energy gap. However, variable-temperature emission lifetime data for c-Cu-Me-DPSO exhibited an increase in lifetime on cooling, with a long-lived component ranging from 6.7 µs at 25 ºC, to 37.0 µs at –196 ºC (Figure 4-13(b)). While the emission lifetime increases at low temperature, the higher energy emission indicated that yellow emission band is not due to a of population of a lower-energy phosphorescent state. Based on these results, we propose that the observed thermochromism is the result of a structural change about the copper center. The dynamics of [Cu(PP)(NN)]+ complexes after photoexcitation have been well documented.181,191 The ground state d10 electronic configuration of these complexes favors a pseudo-tetrahedral geometry,  which alters to a tetragonally distorted environment on excitation to the MLCT excited state as the Cu(I) is formally oxidized to a d9 Cu(II) center.59 Following this, the complex undergoes a pseudo-Jahn-Teller distortion on a sub-  110 picosecond timescale, resulting in a flattening distortion of the molecular geometry.60 These classical Cu(I) dynamics are consistent with the observations in Cu-DPSO.   Figure 4-14 (a) Photographs showing the thermochromic behaviour of thin films of c-Cu-Me-DPSO in the ground-state and under UV lamp excitation (lex = 365 nm) at different temperatures. (b) The proposed excited-state conformational change for the thermochromic emission of c-Cu-Me-DPSO.  It has been known for some time, however, that Cu(II) bis-diimine complexes can undergo a thermochromic distortion in the ground-state, whereby a blue-shift in colour is seen on cooling which results from a square-planar to tetrahedral distortion.192-194 We believe that this phenomenon is occurring in the excited, Cu(II)-like state of c-Cu-Me-DPSO (Figure 4-14). When the complex is in the ground-state, c-Cu-Me-DPSO resides in a pseudo-tetrahedral geometry. On irradiating at room temperature the complex undergoes a Jahn-Teller distortion and an orange emission is observed. However when the complex is excited in a cooled state, a distortion towards a  geometry closer to tetrahedral about the now-Cu(II) metal center is preferred, resulting in a diminishing of the orange emissive LE band and the growth of a yellow emissive HE band. The methyl groups in the 6- and 6'-positions on the diimine ligand help stabilize this transition, which explains why the   111 thermochromism is not observed in Cu-DPSO. It is postulated that the yellow emission in a-Cu-Me-DPSO is due to the less rigid environment of the amorphous powder allowing a more tetrahedral geometry to be achieved.   Figure 4-15 (a) Variable temperature UV-vis absorption spectra of c-Cu-Me-DPSO from –196 to 27 ºC. (b) Variable temperature excitation spectra of c-Cu-Me-DPSO measuring the HE emission at 556 nm. (c) Variable temperature excitation spectra of c-Cu-Me-DPSO measuring the LE emission at 658 nm. (d) Variable temperature excitation spectra of Cu-DPSO measuring the emission at 524 nm. Thin films dropcast on quartz from MeOH.   112 Further evidence for the thermochromism being related to the excited-state is found in the variable temperature UV-vis absorption and excitation spectra. On warming a thin film of c-Cu-Me-DPSO from –196 ºC to 25 ºC no change is observed in the lower energy MLCT region of the spectrum (Figure 4-15(a)), however a pronounced change in seen in the excitation spectra. On monitoring the HE emission band, a transition is seen in the excitation spectrum at λmax ≈ 325 nm (Figure 4-15(b)). This transition is strongest at –196 ºC and diminishes on warming to room temperature. Conversely, when monitoring the LE emission band minimal transitions are seen at –196 ºC, but on warming an excitation feature at λmax ≈ 390 nm grows in (Figure 4-15(c)). The excitation wavelength (λex = 380 nm) for the emission experiments traverses the overlap of these two excitation bands giving “full” yellow character at liquid nitrogen temperatures and “full” orange character at room temperature, with a mix of the two states present at intermediate temperatures. There is no evidence of an excitation growing in at energies lower than 450 nm, indicating that the orange emission is not caused by the formation of a [Cu(Me-DPSO)2][BF4] complex. The spectrum of Cu-DPSO shows no shift in the excitation band, merely a diminishing in intensity on warming (Figure 4-15(d)). The emissive behavior during the crystallization transition from a-Cu-Me-DPSO to c-Cu-Me-DPSO was also monitored to help elucidate the origin of the thermochromic emission. a-Cu-Me-DPSO was drop-cast from MeOH onto a quartz slide, heated in a cryostat and the emission spectra monitored (Figure 4-16). At temperatures up to 87 ºC no change in the HE emission band are seen, however at 107 ºC the LE orange emission begins to grow in, and monitoring at this temperature over a period of 30 minutes shows the HE band diminishing and the LE band growing in. If the thermochromism were due to temperature-dependent population of a different electronic state, it would be anticipated that growth of the LE orange emission band would occur in seconds   113 or quicker, and not on the timescales observed. The temperature range in which the colour-change occurs agrees with the temperature required for onset of crystallization as observed by DSC. These data give further evidence towards a thermochromic emission caused by a structural change.  Figure 4-16 Solid state emission spectra of the crystallization process from a-Cu-Me-DPSO to c-Cu-Me-DPSO. At 87 ºC (black trace) the thin film shows HE emission, and at 107 ºC (initial) the LE band has already started to grow in. Over the next 30 min the HE band diminishes and the LE band increases in intensity (pink trace).  4.6 Conclusions The ability to control or tune photoluminescence properties is advantageous in the design of photofunctional materials, particularly those which react to an external stimulus. While there have been other examples of thermochromic emission behavior in copper(I) systems, those were   114 in halide clusters only where Cu–Cu interactions play a role. In this work we have developed two new [Cu(disphosphine)(diimine)]+ complexes, Cu-DPSO and Cu-Me-DPSO, where sulfoxide-bridged diimine ligands are present. Cu-DPSO displayed photophysical properties typical of other Cu(I) complexes. However, on addition of methyl substituents in the 6- and 6´-positions of the diimine ligand to give complex Cu-Me-DPSO, we were able to access a variety of different photophysical properties. Cu-Me-DPSO is initially isolated as an amorphous solid (a-Cu-Me-DPSO) which under UV irradiation gives a yellow emission. Upon heating to 180 ºC, a-Cu-Me-DPSO forms a crystalline species which is retained on cooling to room temperature (c-Cu-Me-DPSO). This crystalline species gives markedly bathochromically shifted (>100 nm) orange emission, which is attributed to a flattening conformational change during the crystallization process due to the increased bulk at the diimine ligand. Low-temperature solid state emission experiments revealed a small red-shift in emission colour for Cu-DPSO, typical for Cu(I) complexes. On the other hand, reversible thermochromic emission was observed in c-Cu-Me-DPSO, whereby the complex undergoes a large colour change from orange at room temperature to yellow at –196 ºC. It is inferred that this hypsochromic shift is not due to the population of a lower energy triplet state at low temperatures (as is often the case in copper(I) complexes), but is instead attributed to a geometric distortion in the excited-state. At room temperature the excited c-Cu-Me-DPSO complex exists in a more flattened state giving an orange emission, however at –196 ºC the complex prefers to adopt a pseudo-tetrahedral geometry giving rise to a yellow emission. We show that subtle changes in the properties of the diimine ligand in [Cu(disphosphine)(diimine)]+ complexes can have a dramatic impact on the properties of the complex. In utilizing a sulfur-bridged dipyridyl ligand it is not only possible to alter the steric and   115 electronic properties through substitution at the pyridyl rings, but also alter the degree of oxidation at the sulfur atom with addition of electron-withdrawing oxygen atoms. This gives an additional level of control over the electronics, allowing for finely-tuned electronic properties in these ligands.  In showing that a copper(I) complex can exhibit reversible thermochromism, avenues that were inaccessible due to the limitations of thermochromic copper halide clusters can now be realized. For example, inexpensive copper(I) metallopolymers that offer a response to thermal stimuli could be fabricated. The simple synthetic procedures to the complexes reported, coupled with the absence of toxic heavy metals, give avenues towards practical applications in optoelectronics, including light-emitting diodes, display technologies and chemical and biological sensors.  4.7 Experimental Details 4.7.1 General All experiments were conducted under an inert nitrogen atmosphere using standard Schlenk-line techniques, unless otherwise stated. 2-Bromopyridine, 2-bromo-6-methylpyridine and tetrakis(acetonitrile)copper(I) tetrafluoroborate were purchased from Aldrich. POP was purchased from Strem and thiourea obtained from TCI. All compounds were used as received. All other solvents were obtained from commercial sources and used without any further purification.  4.7.2 Spectroscopy 1H NMR, 13C NMR and 31P NMR spectra were collected using a Bruker AV-400 spectrometer and referenced to the residual protonated solvent peak. NMR solvents (from either   116 Cambridge Isotope Laboratories or Aldrich) were used as received. Electrospray ionization mass spectrometry data were obtained using a Bruker Esquire LC ion trap mass spectrometer and elemental analysis determined using a Thermo Flash 2000 Elemental Analyzer. Infrared spectra were collected using a Perkin Elmer Frontier FT-IR spectrometer and DSC data collected using a Netzsch DSC 216 Polyma. Electronic absorption spectra were recorded on a Varian-Cary 5000 UV−vis−near-IR spectrophotometer. Fluorescence data were collected on a Photon Technology International (PTI) QuantaMaster 50 fluorimeter fitted with an integrating sphere, double excitation monochrometer and utilizing a 75 W Xe arc lamp as the source. Emission lifetime data were collected using a Horiba Yvon Fluorocube TCSPC apparatus. A 370 nm NanoLED source pulsing at a repetition rate of 100−1000 kHz was used for excitation. Broad-band emission was monitored by a CCD detector at wavelengths >450 nm using a low-pass filter. Data were fitted using the DAS6 Data Analysis software package. Samples for low temperature emission were cooled using an Oxford Instruments Optistat DN. Solution spectra were collected using 1 cm2 quartz cells (Starna Cells) and spectra of the neat solids were collected by drop-casting from MeOH on to quartz slides (Ted Pella, Inc.).  4.7.3 X-Ray Crystallography Single crystal X-ray data were collected using a Bruker APEX DUO diffractometer with graphite monochromated Mo Kα radiation (λ = 0.71073 Å) at 90 K. Raw frame data were processed using APEX2.195 Structure solution was performed using SUPERFLIP and refinement was carried out using full-matrix least-squares on F within the CRYSTALS suite.154,155 All non-hydrogen atoms were refined with anisotropic displacement parameters. The positions of the hydrogen atoms were calculated using the Fourier difference map and were initially refined with   117 restraints on bond lengths and angles, after which the positions were used as the basis for the riding model.156 Powder X-ray diffraction patterns were collected using a Bruker D8 Advance Powder Diffractometer and data analyzed using Bruker DiffracPlus EVA software.  4.7.4 Synthesis di(pyridin-2-yl)sulfane (DPS) Using a procedure adapted from Khan et al.96 a two-necked round-bottomed flask was charged with  thiourea (5.86 g, 77.0 mmol, 0.5 eq.), placed under an N2 atmosphere and dissolved in 500 mL EtOH. To this, 2-bromopyridine (15.0 mL, 154 mmol, 1.0 eq.) was added and the reaction mixture heated to reflux for 36 h, after which the flask was cooled to room temperature and the solvent removed in vacuo. The crude product was dissolved in CH2Cl2 (50 mL), washed with water (3 × 50 mL) then brine (1 × 100 mL) and the organic layer dried with MgSO4, filtered and the solvent evaporated under reduced pressure. Following purification using a silica gel column and a solvent gradient from 1:0 CH2Cl2:MeOH to 9:1 CH2Cl2:MeOH, a yellow-brown oil was isolated (13.5 g, 92%). The spectroscopic data matched that of the literature.196  bis(6-methylpyridin-2-yl)sulfane (Me-DPS) 2-bromo-6-methylpyridine (0.66 mL, 5.81 mmol, 1.0 eq.) was added to a solution of thiourea (0.210 g, 2.76 mmol, 0.48 eq.) in 100 mL of EtOH under an N2 atmosphere and heated to reflux for 36 h. The reaction mixture was cooled to room temperature the solvent removed and the resultant residue dissolved in CH2Cl2 (20 mL) and washed with water (3 × 20 mL) and brine (1 × 50 mL). The organic layer was dried with MgSO4, filtered and the CH2Cl2 removed in vacuo. The crude product was purified by column NSN123456Me Me  118 chromatography on silica gel using a solvent gradient from 1:0 CH2Cl2:MeOH to 9:1 CH2Cl2:MeOH, yielding a brown oil (0.558 g, 2.58 mmol, 94%). 1H NMR (400 MHz, CD2Cl2) δ 7.50 (t, J = 7.7 Hz, 2H, H4), 7.19 (dt, J = 7.9, 0.8 Hz, 2H, H3), 7.02 (ddd, J = 7.6, 0.9, 0.5 Hz, 2H, H5), 2.50 (s, 6H, HMe). 13C{1H} NMR (CD2Cl2, 101 MHz) δ 159.82 (C2), 156.80 (C6), 137.69 (C4), 123.05 (C3), 121.67 (C5), 24.62 (CMe). HR-ESI MS: Calcd for C10H8N2OSNa: 227.0255; Found: 227.0255 [M+Na]+.  2,2'-sulfinyldipyridine (DPSO) DPS (0.173 g, 0.919 mmol, 1.0 eq.) was dissolved in 2 mL glacial acetic acid, to which a 30% solution of H2O2 (0.35 mL, 3.68 mmol, 4.0 eq.) was added dropwise. The reaction mixture was stirred at 25 ºC for 42.5 h, following which 6 M NaOH was added until basic, forming a yellow suspension which turned white after stirring for 30 min. Brine (50 mL) was added and the crude product extracted with CH2Cl2 (3 × 10 mL) followed by drying of the combined organics over MgSO4 and removal of the CH2Cl2 in vacuo. Purification over a silica gel column using a 95:5 CH2Cl2:MeOH solvent system resulted in an off-white powder (0.107 g, 0.524 mmol, 57%). 1H NMR (400 MHz, CD2Cl2) δ 8.57 (d, J = 4.6 Hz, 2H, H4), 7.97 (d, J = 7.9 Hz, 2H, H1), 7.88 (td, J = 7.7, 1.6 Hz, 2H, H2), 7.39 – 7.31 (m, 2H, H3). 13C{1H} NMR (CD2Cl2, 100 MHz) δ 164.84 (C6), 150.23 (C4), 138.40 (C2), 125.47 (C3), 119.96 (C1). HR-ESI MS: m/z calcd. for C10H8N2ONaS: 227.0255; Found: 227.0255 [M+Na]+. IR (neat): ṽ (σ SO) 1037 cm−1.    NSN123456O  119  6,6'-sulfinylbis(2-methylpyridine) (Me-DPSO)  Me-DPS (0.199 g, 0.920 mmol, 1.0 eq.) was dissolved in 12 mL glacial acetic acid and heated to 25 ºC. To this, 30% H2O2 (0.35 mL, 2.68 mmol, 4.0 eq.) was added dropwise, and the solution left to stir for 23 h. The reaction mixture was cooled, basified using 6 M NaOH and the crude product extracted using CH2Cl2 (3 × 15 mL). The combined organics were dried over MgSO4, filtered and the solvent removed in vacuo. Column chromatography on silica gel utilizing a 95:5 CH2Cl2:MeOH solvent system gave an off-white powder (0.100 g, 0.431 mmol, 47%). 1H NMR (CD2Cl2, 400 MHz): δ 7.77 – 7.70 (m, 4H, H3,4), 7.21 – 7.15 (m, 2H, H5), 2.53 (s, 6H, HMe). 13C{1H} NMR (CD2Cl2, 101 MHz) δ 164.40 (C2), 159.91 (C6), 138.47 (C4), 125.29 (C5), 117.21 (C3), 24.48 (CMe). IR (solid): (σ SO) 1037 cm-1. HR-ESI MS: Calcd for C12H12N2OSNa: 255.0568; Found: 255.0569 [M+Na]+   [Cu(POP)(DPSO)][BF4] (Cu-DPSO)  [Cu(MeCN)4]BF4 (0.077 g, 0.245 mmol, 1.0 eq.) and POP (0.132 g, 0.245 mmol, 1.0 eq.) were dissolved in CH2Cl2 and stirred at room temperature for 2 h, after which DPSO (0.050 g, 0.245 mmol, 1.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated in vacuo, the product precipitated with Et2O, filtered and washed with Et2O, giving a yellow-white powder (0.194 g, 0.217 mmol, 89%). 1H NMR (CD3CN, 400 MHz): δ 8.48 (d, J = 4.8 Hz, 2H, H6), 8.01 – 7.94 (m, 4H, H3,5), 7.47 – 7.34 (m, 14H, H4,9,10), 7.34 – 7.25 (m, 10H, H8,14), 7.00 (td, J = NSN123456Me MeOCuPOPNNS OBF412314 11107 4 5613891215 16  120 7.6, 1.1 Hz, 2H, H13), 6.98 – 6.92 (m, 2H, H15), 6.73 (ddt, J = 6.2, 4.0, 2.4 Hz, 2H, H12). 13C{1H} NMR (CD3CN, 101 MHz) δ 164.85 (C1), 158.81 (C16), 151.06 (C6), 139.76 (C5), 135.02 (C12), 134.47 (t, C8), 132.99 (C14), 131.99 (t, C7), 131.22 (C10), 129.84 (t, C9), 126.60 (C4), 125.77 (C13), 124.47 (t, C11), 121.12 (C15), 120.87 (C3). 31P{1H} NMR (CD3CN, 161 MHz): δ –14.1. HR-ESI MS: Calcd for C46H36N2O2P2SCu: 805.1269; Found: 805.1276 [M–BF]+. Anal. Calcd for C46H36BF4N2O2P2SCu: C, 61.86; H, 4.06; N, 3.14; S, 3.59. Found: C, 61.82; H, 4.23; N, 2.98; S, 3.47.  [Cu(POP)(Me-DPSO)][BF4] (Cu-Me-DPSO) [Cu(MeCN)4]BF4 (0.068 g, 0.215 mmol, 1.0 eq.) and POP (0.116 g, 0.215 mmol, 1.0 eq.) were dissolved in CH2Cl2 and stirred at room temperature for 2 h, after which Me-DPSO (0.050 g, 0.215 mmol, 1.0 eq.) was added and the solution left to stir for a further 1 h. The reaction mixture was then concentrated in vacuo, the product precipitated with Et2O, filtered and washed with Et2O, giving a yellow-white powder (0.151 g, 0.164 mmol, 76%). 1H NMR (CD3CN, 400 MHz): δ 7.82 (t, J = 7.7 Hz, 2H, H5), 7.70 (d, J = 7.7 Hz, 2H, H4), 7.41 (dt, J = 13.6, 6.8 Hz, 12H, H8,9), 7.35 – 7.23 (m, 12H, H8,14,16), 6.99 (t, J = 7.5 Hz, 2H, H13), 6.93 (d, J = 8.0 Hz, 2H, H15), 6.71 (dt, J = 7.2, 3.3 Hz, 2H, H12), 2.48 (s, 6H, HMe). 13C{1H} NMR (CD3CN, 101 MHz) δ 164.69 (C1), 160.36 (C3), 158.91 (C16), 139.40 (C5), 135.11 (C12), 134.49 (t, C8), 132.93 (C14), 132.21 (t, C7), 131.18 (C10), 129.82 (t, C9), 126.07 (C6), 125.71 (C13), 124.44 (C11), 121.05 (C15), 117.87 (C4), 24.19 (CMe). 31P{1H} NMR (CD3CN, 161 MHz): δ –14.6. HR-ESI MS: Calcd for CuPOPNNS OMeMeBF412314 11107 4 5613891215 16Me  121 C48H40N2O2P2SCu: 833.1582; Found: 833.1600 [M–BF4]+. Anal. Calcd for C48H40BF4N2O2P2SCu: C, 62.85; H, 4.38; N, 3.04; S, 3.48. Found: C, 62.41; H, 4.41; N, 2.85; S, 3.40.      122 CHAPTER 5: Sulfur-Bridged Bimetallic Ru(II)-Re(I) Systems – Modulating CO2 Reduction with Sulfur Oxidation State  5.1 Introduction  Global energy consumption is rapidly increasing, driven by a population boom and industrialization. In 2001, global energy consumption was 4.3 × 1020 J197 with this usage projected to triple by 2050.198 The majority of this energy is obtained from fossil resources: oil, coal and natural gas. When burned, these fuels generate a significant amount of CO2 which is then released into the atmosphere, and this increase in atmospheric CO2 concentration results in global warming. Additionally, fossil fuels are a finite resource and are not just an energy source, but also a feedstock for many chemical starting materials. It is therefore of great importance to reduce the volume of CO2 released in to the atmosphere, but also to convert this otherwise wasted by-product into something more industrially useful.  5.1.1 Photocatalytic CO2 Reduction With more efficient methods required to convert solar energy to useable chemical energy, researchers have taken inspiration from Nature in a bid to mimic photosynthesis and ultimately harness this solar radiation through storage of energy in chemical bonds and reactive chemical species.199,200 Photocatalytic reduction is a technique in which CO2 is converted into useful higher energy components such as CO or formic acid, and various metal complexes have been reported which are catalytically active for this process.201 Photochemical CO2 reduction involves two   123 components: a redox photosensitizer (PS) that absorbs light in the visible region; and a catalytic center at which the reduction occurs. The scope of this work will focus on Ru(II) as the PS unit, and Re(I) as the catalyst. Since the first reported example of luminescence from tris-2,2´-bipyridine ruthenium(II) by Paris and Brandt nearly 60 years ago,202 a significant number of members of the coordination chemistry community have focused their research on ruthenium(II) complexes. While expensive, ruthenium has proven itself a useful tool through modification of its transition metal complexes, giving rise to a wealth of tunable spectroscopic and redox properties, stemming from a low-lying, intense Ru(dπ) to ligand(π*) metal-to-ligand charge transfer (MLCT) transition. These properties have helped establish ruthenium(II) complexes as viable candidates for solar energy conversion, as molecular devices, in the medical field towards new therapeutics and bio-probes, and for use in photocatalysis.203 Ruthenium(II) complexes with polypyridyl ligands are ideal for use as photosensitizers204 due to their long excited-state lifetimes, redox properties, chemical stability, and excited-state reactivity.205 The catalytic center in light-driven CO2 reduction systems has predominantly been based on fac-Re(bpy)(CO)3X (X = Cl– and Br–) due to the stability, high product selectivity and efficiencies of these complexes.201 The use of Re(I) complexes as monometallic units for CO2 reduction was first demonstrated by Lehn, Ziessel and Hawecker in 1983.206 It has been shown, however, that supramolecular photocatalytic systems – whereby the PS and catalytic units are conjoined – can greatly increase the reduction efficiency.207    124  Figure 5-1 Depiction of a supramolecular photocatalyst assembly for CO2 reduction, where (a) illustrates the generation of the one electron reduced species (OERS) of the photocatalyst, and (b) illustrates the catalytic cycle for CO2 reduction.  A supramolecular photocatalyst consists of a visible-light absorbing PS unit and a catalytic unit bound together via a bridging ligand (Figure 5-1). The electron transfer between the PS and the catalyst is accelerated when compared to relying on intermolecular interactions for a solution containing free PS and free catalyst. This enhances the durability of the system as the unstable redox states can be consumed more rapidly, helping improve the efficiency of the system.208 The reaction mechanism can be broken into two sections, the first being the generation of the one   125 electron reduced species (OERS) of the photocatalyst (Figure 5-1(a)), and the second being the catalytic cycle for the production of CO (Figure 5-1(b)). The mechanism proceeds as follows: 1. Irradiation of the metallic PS gives the 3MLCT excited-state. 2. The excited PS unit is quenched by a sacrificial electron donor (D), generating the one electron reduced species (OERS) of the PS. 3. A rapid intramolecular electron transfer from the OERS of the PS to the catalytic center (Cat). 4. CO2 binds to the photocatalyst. 5. Additional OERS of the PS–Cat remaining in solution functions as a redox photosensitizer and supplies an electron, forming an intermediate. 6. The electron for the second reduction process is supplied again by remaining OERS of PS-Cat, resulting in the formation of CO and the regeneration of PS–Cat.  5.1.2 Ru(II)-Re(I) Supramolecular Photocatalytic Systems In 2005, Ishitani and co-workers demonstrated the first successful Ru(II)-Re(I) supramolecular photocatalyst for CO2 reduction (24, Figure 5-2) consisting of a [Ru(N^N)]2+ PS unit and a fac-Re(N^N)(CO)3Cl catalyst bridged by two 4-methyl-bpy ligands connected with a -CH2CH(OH)CH2- chain.209 When irradiating with visible light (> 500 nm) and using 1-benzyl-1,4-dihydronicotinamide (BNAH) as a sacrificial electron donor, CO was found to be produced with a high efficiency and selectivity (ΦCO = 0.12, TONCO = 170). These Ru(II)-Re(I) systems have since garnered considerable research and it has been shown that the bridging ligand plays a major role in the efficiency of the systems, a selection of which are found in Figure 5-2.    126   Figure 5-2 Selected Ru(II)-Re(I) systems.209-211  It was found that the length of the alkyl chain linking the Ru(II) and Re(I) centers impacted the efficiency of the supramolecular catalyst.210 When the alkyl linker bridging the species increased in length the efficiency decreased. For 25 containing an ethyl linked bridge, ΦCO = 0.13 and TONCO = 180. However on increasing the length to a butylene (26) or hexylene (27) chain, the catalytic activity diminished (ΦCO = 0.11 and TONCO = 120) with similar properties seen between the two species with longer bridges. An electronic interaction between the two metal centers in the excited state was only observed in 25 and not in the 26 or 27 dyads. This causes a higher reductive quenching efficiency of the 3MLCT state by BNAH in 25, leading to increased reduction of CO2.  The electronics at the bridge also affect the performance of the photocatalysts.208 28 is structured with the two bipyridyl moieties of the bridge joined by a carbon-carbon double bond,211   127 which is comparable to 25 that is connected via the saturated analogue. 25 displays a significant enhancement in efficiency over 28 (TON = 180 and 50, respectively). Extended conjugation at the bridging ligand lowers the reducing power of the OERS of rhenium which can inhibit its reductive catalytic activity. It has been reported that a threshold for the photocatalytic activity of mononuclear Re complexes exists at E1/2red = –1.4 V, and it is this occurrence that explains the difference in activity: E1/2red of 28 was –1.34 V, while 25 was –1.77 V.  In this work we investigate how the oxidation state at sulfur affects the CO2 reductive properties of two new Ru(II)-Re(I) supramolecular catalysts. A sulfur-bridged bisphenanthroline is utilized as the connector between the two metal units, allowing us to maintain almost-identical geometric structures between the sulfide and sulfone oxidation states.  5.2 Ligand Synthesis and Characterization The series of ligands described in this work was prepared using modifications of previously reported procedures (Scheme 1). Ligand PSP was prepared in a one-pot synthesis through reaction of Li2S with commercially available 5,6-epoxy-5,6-dihydro-1,10-phenanthroline in refluxing tert-butanol.212 This pro-ligand was then oxidized to the sulfone (PSO2P) analogue with 30% H2O2 in the presence of a NbC catalyst.99 Both bridging pro-ligands were isolated in moderate to high yields and structural characterization performed using 1H, 13C{1H}, COSY, HSQC and HMBC NMR experiments, in addition to mass spectrometry and infrared spectroscopy.    128 Scheme 5-1 Synthesis of Ru(II)-Re(I) complexes Ru-PSP-Re and Ru-PSO2P-Re.  5.3 Synthesis and Characterization of Ru(II) Complexes cis-Ru(phen)2Cl2 was synthesized as per the literature.213 Monometallic Ru(II) complexes Ru-PSP and Ru-PSO2P were prepared using an adapted procedure from Johansson et al. utilizing a microwave synthesizer.214 cis-Ru(phen)2Cl2 was reacted with an excess of the desired pro-ligand (Scheme 1) and the crude products purified via column chromatography over silica gel (F60) using   129 CH3CN:H2O:KNO3 (aq) as the eluting solvent. The pure products were isolated as hexafluorophosphate salts through the addition of a saturated aqueous NH4PF6 solution that resulted in the formation of precipitates that were orange in colour in all cases. Monometallic rhenium complexes Re-PSP and Re-PSO2P were synthesized by reaction of the desired pro-ligand with Re(CO)5Br in MeOH heated to reflux in the dark. Purification occurred by purification with column chromatography using silica gel (F60) using first CH3CN then 98:2 CH2Cl2:MeOH as the eluting solvents. Bridged Ru(II)-Re(I) species were prepared through reaction of Ru-PSP or Ru-PSO2P with Re(CO)5Br under a nitrogen atmosphere in the dark followed by purification over silica gel in a similar manner to the monometallic Ru(II) species. 1H and 13C NMR spectra were assigned using 1H, 13C{1H}, COSY, ROESY, HSQC and HMBC experiments. The HR-MS spectra show product peaks corresponding to [M−PF6]+ exhibiting characteristic ruthenium and rhenium isotope patterns.  Both Ru-PSP-Re and Ru-PSO2P-Re exist as diastereomers due to chirality at both the ruthenium and rhenium centers. This unfortunately made full assignment of the complexes difficult, even with a field strength of 20 T in the NMR spectrometer; however through analysis of the bridge protons in Ru-PSO2P-Re it was possible to observe two different species (Figure 5-3(a)). The protons in the 6- and 6´-positions on the PSO2P bridge are singlets in 1H NMR with signals at δ = 9.50 and δ = 8.23 for the phenanthroline moieties bound at Ru and Re, respectively. ROESY experiments observing through-space correlations between the bridge and the phen ligands bound to the Ru center were used to determine the binding of PSO2P. The singlet seen at each position is in fact two overlapping signals (for the two diastereomers) as determined from HSQC experiments (Figure 5-3(b)) indicating that the proton signal corresponds to two separate   130 carbon environments. From COSY, ROESY, HSQC, HMBC and 13C{1H} experiments it was possible to distinguish the bridging ligands of the two distinct isomers.   Figure 5-3 (a) 1H NMR shifts for the bridging ligand in Ru-PSO2P-Re. Black text designates signals from one diastereotopic species, red text for the other. Chemical shifts with only one number indicates overlapping, indistinguishable species. (b) HSQC NMR spectrum (CD3CN, 850 MHz, 25 ºC) showing two diastereotopic carbon signals correlating to one broad singlet.  5.4 Electrochemical Properties The electrochemical behavior of each monometallic and bimetallic species was determined using cyclic voltammetry (CV) and square wave voltammetry (SWV) referenced to a Fc/Fc+ standard. CVs were collected at a scan rate of 0.1 Vs–1 and a 25 Hz wave form was applied in SWV. From CV data (Figure 5-4 and Table 5-1), both ruthenium monometallic complexes (Ru-PSP and Ru-PSO2P) display a quasi-reversible anodic event at ~1.0 V which is consistent with a Ru(II)/Ru(III) couple. Oxidation of the sulfur bridge to Ru-PSO2P was found to be accompanied by a +40 mV shift in anodic peak potential with no significant change in ΔEp and is taken to be a   131 result of the greater electron-withdrawing character of the sulfone group. The rhenium monometallic species (Re-PSP and Re-PSO2P) are oxidized irreversibly at potentials between 50 and 100 mV positive of the analogous ruthenium complexes. Sulfur oxidation state appears to impose a greater effect on the Re(I)/Re(II) couple, where the wave was found to shift positively by 100 mV with a change from a sulfide to sulfone bridging group. The corresponding oxidative wave for the bimetallic species (Ru-PSP-Re and Ru-PSO2P-Re) occurs at an intermediate potential between the corresponding Ru and Re analogues, and the current response appears as a superposition of the two monometallic complexes. However, the change in peak potential as function of sulfur oxidation state is nearly equivalent to that observed for the Ru subunit.     132  Figure 5-4 Cyclic voltammograms of (a) Ru(II) monometallic complexes; (b) Re(I) monometallic complexes; (c) Ru(II)-Re(I) bimetallic species. All spectra taken in dry, N2-sparged CH3CN solutions at room temperature.  Due to overlapping peaks in the CV, the cathodic response from each complex was determined with SWV (Table 5-1). The most positive reduction potential of the monometallic species is generally assigned to a ligand-based reduction and was found to occur at potentials between –1.41 and –1.65 V for these complexes. The separation in peak potential between analogous Re and Ru subunits was found to be much smaller for this wave than the metal-based couples, where Epc1 (peak cathodic potential) was found to be less than 12 mV negative for the Ru   133 complexes. In contrast, oxidation of the sulfur bridge from sulfide to sulfone resulted in a 220 mV negative shift in Epc1 for both cases. This result supports the assignment of Epc1 to a ligand-based reduction and indicates that the oxidation state of the linking sulfur bears a larger influence on redox properties of the bridging ligand than on the metal center. Additionally, the slightly higher reduction potentials measured for Re-PSP (–1.63 V) and Re-PSO2P (–1.41 V) suggests that downhill electron transfer can occur from Ru to Re in the dyads, with a limited loss of energy.  Table 5-1 Electrochemical data of Ru and Re mono- and bimetallic species.  E(ox) [V] [a] E(red) [V] [b] Ru-PSP 0.96 [c] –1.64 Ru-PSO2P 1.00 [c] –1.42 Re-PSP 1.00 [d] –1.63 Re-PSO2P 1.11 [d] –1.41 Ru-PSP-Re 1.00 [c] –1.56 Ru-PSO2P-Re 1.04 [c] –1.30 [a] Cyclic voltammetry. [b] Square-wave voltammetry. [c] Quasi-reversible wave. [d] Irreversible wave. All experiments conducted in dry CH3CN solutions sparged with N2.  The first reduction event for the S and SO2 linked bimetallic complexes (Ru-PSP-Re and Ru-PSO2P-Re) occurred at significantly less negative potentials (–1.56 and –1.30 V, respectively) than that of either of their monometallic analogues, with Ru-PSO2P-Re affected to a greater extent. The lower reduction potential for the sulfone species in comparison to the sulfide is due to the additional conjugation in the bridging ligand reducing the π*-orbital energy of the phenanthroline unit at the rhenium(I).215 The positive change in Ep1 between the mono- and   134 bimetallic bridging ligand may be explained by added electron withdrawing character brought on by coordination to the electropositive Re center.  5.5 Photophysical Properties 5.5.1 Absorption Spectroscopy  Figure 5-5 Absorption spectra of PSP and PSO2P in CH2Cl2 solutions at room temperature.  Figure 5-5 shows the UV-vis absorption spectra of pro-ligands PSP and PSO2P measured in CH2Cl2 at room temperature. Both compounds contain strong features at energies higher than 300 nm are attributed to π–π* transition on the phenanthroline rings, analogous to those observed in 1,10-phenanthroline.216 Both compounds also exhibit lower energy features tailing to 380 nm and 345 nm for PSP and PSO2P, respectively, that are attributed to the introduction of the sulfur bridge. This low energy band is broader for the sulfide relative to the sulfone species. In the sulfide-bridged compound it is possible to achieve multiple conformations about the sulfur leading to differing degrees of interaction between the phenanthroline rings, contributing to a broadness in the band. In the sulfone, however, the absence of lone pairs on the sulfur prevents these interactions   135 and results in a more structured absorption. This has been seen in calculations of poly-p-phenyl sulfide217 and in oligo- and polythienyl sulfur-bridged compounds by our group.80   Figure 5-6 Absorption spectra of (a) Ru-PSP-Re (blue) and its monometallic model complexes Ru-PSP (grey) and Re-PSP (purple). (b) Ru-PSO2P-Re (pink) and model complexes Ru-PSO2P (grey) and Re-PSO2P (teal). CH3CN solutions of concentration ~3 × 10–6 M at room temperature.  Figure 5-6(a) shows the UV-vis absorption spectrum of Ru-PSP-Re and its model mononuclear complexes Ru-PSP and Re-PSP, and Figure 5-6(b) shows Ru-PSO2P-Re and the model species Ru-PSO2P and Re-PSO2P. In both dyads, broad absorptions between 380–500 nm are due to a 1MLCT at the ruthenium, and those between 350–440 nm contain contributions from 1MLCT absorption of both the ruthenium and rhenium metals, however as with other Ru(II)-Re(I) systems the 1MLCT of the Re is often obscured.209 Strong absorptions from 200–300 nm are due to π–π* transitions at the diimine ligands.    136 5.5.2 Photoluminescence Spectroscopy  Figure 5-7 (a) Emission spectra of Ru-PSP (light grey), Ru-PSO2P (dark grey), Ru-PSP-Re (blue) and Ru-PSO2P-Re (pink). Spectra were normalized to the absorbance at the excitation wavelength (λex = 450 nm). (b) TCSPC data for Ru-containing monometallic and bimetallic species (λex = 453 nm). CH3CN solutions of concentration ~3 × 10–6 M at room temperature.  The photoluminescence spectra of Ru-PSP-Re, Ru-PSO2P-Re and their Ru monometallic counterparts were measured in CH3CN solutions and are shown in Figure 5-7 and presented in Table 5-2(a). Emission spectra were normalized for absorbance at the excitation wavelength (450 nm) Following irradiation, all ruthenium-containing complexes exhibited a broad, featureless emission profile indicative of radiative decay from a Ru-based 3MLCT excited-state.40 The dyads were each slightly (5–8 nm) red-shifted of their monometallic Ru counterparts indicating that there is a weak electronic interaction between the two metal centers.210 However, when comparing the two bimetallic species, the emission maximum of Ru-PSO2P-Re was red-shifted 24 nm compared to Ru-PSP-Re, which in addition to the quenching of emission intensity, suggests that increasing the oxidation at the sulfur bridge in turn increases the electronic communication between the Ru   137 and Re centers. While a decrease in excited-state lifetime would be expected in concert with a reduction in emission intensity, in Figure 5-7(b) and Table 5-2 it is shown that τem increases for Ru-PSO2P-Re (1.46 µs) over Ru-PSP-Re (1.27 µs). This could be explained by increased delocalization of the excited electron at the sulfone unit compared to the sulfide, as seen in the monometallic Ru species (Ru-PSP = 0.93 µs; Ru-PSO2P = 2.20 µs).  Table 5-2 Photophysical properties of Ru-PSP-Re, Ru-PSO2P-Re and their respective monometallic model complexes.[a] Complex lmax (nm) [b] τem (µs) [c] Ru-PSP 603 0.93 [d] Ru-PSO2P 625 2.20 [e] Ru-PSP-Re 608 1.27 [d] Ru-PSO2P-Re 632 1.46 [d] [a] Measured in CH3CN at 25 ºC. Samples were sparged with Ar for 30 min prior to irradiation and left under a blanket of Ar for the duration of the studies. [b] Excitation wavelength was 450 nm. [c] Excitation wavelength was 453 nm. [d] Collected over a 10 µs time window. [e] Collected over a 20 µs time window.  5.6 Preliminary Photocatalytic Studies For the photocatalytic reaction, a CH3CN-triethanolamine (TEOA) (5:1, v/v) solution was prepared containing either Ru-PSP-Re or Ru-PSO2P-Re (0.05 mM) and BNAH (0.1 M) as the sacrificial reductant. The sample was saturated with CO2 and irradiated with a 500 nm LED setup (2.7 × 10–8 einstein s–1, FWHM = 30 nm). This emission source was chosen as it selectively excites only the Ru unit, and not into the Re directly. Ru(II)-Re(I) supramolecular CO2 photocatalysts typically form CO as the major product, with H2 and HCOOH as by-products.208 After irradiation for 24 h, TONCO values were found to   138 be very low for Ru-PSP-Re and Ru-PSO2P-Re (0.8 and 0.9, respectively) which was surprising as for both complexes bubbling was observed in the solution during the photocatalysis, to the point of having pressurized the sealed cell after just 1 h. In both studies GC data shows that H2 is not evolved, and solvent suppressed 1H NMR experiments show that no HCOOH is observed. The GC traces for Ru-PSP-Re (Figure 5-8(a)) and Ru-PSO2P-Re (Figure 5-8(b)) each show a significant peak overlapping with that for CO2, however this peak is due to volatized CH3CN produced during the CO2 sparging.   Figure 5-8 Gas chromatograms of (a) Ru-PSP-Re and (b) Ru-PSO2P-Re.  Further experiments are required to elucidate the gas being evolved. While not seen for Re(I) photocatalyic CO2 reduction, possible volatile reduction products include methane, ethane and ethylene. Atmospheric mass spectrometry will be employed to give an online detection over a mass range of 0–100 m/z. While a number of gases have the same m/z (CO2, N2 and C2H6 with m/z = 28), the study will be crucial in monitoring if any sulfur-containing volatiles are produced   139 from decomposition of the Ru(II)-Re(I) catalyst. Additionally, GC-MS will be conducted for further analysis of the gas produced.  5.7 Conclusions and Future Work Two new Ru(II)-Re(I) dyads have been synthesized with their metal centers tethered by a sulfur-bridged bisphenanthroline bridging ligand with a view for use as supramolecular photocatalysts for CO2 reduction. In changing the oxidation state at the sulfur center from sulfide to sulfone it is hoped that the degree of electron transfer from the Ru(II) photosensitizer to the Re(I) catalytic center can be modulated. Electrochemical data indicates that Ru-PSP-Re should be the more efficient catalyst of the two due to its lower reduction potential; however, photophysical data suggests Ru-PSO2P-Re has a greater degree of electronic communication between the two metals due to a red-shift and quenching of emission observed in the photoluminescence spectra. Preliminary photocatalytic data show that both complexes show poor efficiency for generation of CO with TONCO < 1. GC and solvent suppression 1H NMR data show that neither H2 nor HCOOH are produced as by-products, but a substantial peak overlapping with the CO2 signal is seen in the GC trace, corresponding to an unidentified compound. Further studies involving atmospheric-MS and GC-MS must be performed in order to elucidate the unknown volatile species.    140 5.8 Experimental Details 5.8.1 General All experiments were conducted under air unless otherwise stated. Solvents used for synthesis were of reagent grade and were used without any further purification. HPLC grade solvents were used for spectroscopic studies. Microwave reactions were performed in a Biotage Initiator 2.5 microwave synthesizer. 5,6-epoxy-5,6-dihydro-[1,10]-phenanthroline, ruthenium(III) chloride and 30% H2O2 solutions were purchased from Sigma-Aldrich. Niobium carbide was purchased from Alfa Aesar. All reagents were used as received. cis-Ru(phen)2Cl2 was prepared as per the literature.213  5.8.2 Spectroscopy 1H, 13C{1H}, COSY, ROESY, HSQC and HMBC NMR spectra were collected using Bruker NMR spectrometers and referenced first to TMS and then to the residual protonated solvent peak. Pro-ligand spectra were collected on a Bruker AV-400 spectrometer, spectra of the monometallic species on a Bruker AV-600 spectrometer, and spectra of the mixed-metal systems on a Bruker AV-850 spectrometer. NMR solvents (from Aldrich) were used as received. Electrospray ionization mass spectrometry data was obtained using a Bruker Esquire LC ion trap mass spectrometer. Infrared spectroscopy was performed on an attenuated total reflection (ATR) crystal using a Perkin-Elmer Frontier FT-IR spectrometer. UV-vis absorption spectra were recorded on a Varian-Cary 5000 UV-Vis-near-IR spectrophotometer. Fluorescence data were collected on a Photon Technology International (PTI) QuantaMaster 50 fluorimeter fitted with an integrating sphere, double excitation monochrometer and utilizing a 75 W Xe arc lamp as the source. Emission lifetime data were collected using a Horiba Yvon Fluorocube TCSPC apparatus.   141 A 453 nm NanoLED was used as the excitation source, pulsing at a repetition rate of 100 kHz. Broadband emission was monitored by a CCD detector at wavelengths >550 nm using a low pass filter. Data were fitted using the DAS6 Data Analysis software package. Sample solutions were first sparged with Ar for 30 min and then maintained under a blanket of Ar for the duration of the measurements in 1 cm2 quartz cells (Starna Cells) fitted with a rubber septum.  5.8.3 Electrochemistry All electrochemical measurements were performed using a Metrohm Autolab PGSTAT12 in an air tight single compartment cell with the use of an Ag/AgNO3 pseudo reference electrode.  A 7 mm2 glassy carbon disc and a platinum mesh were used as the working and counter electrodes, respectively. Each experiment was referenced to a Fc/Fc+ couple and carried out in a N2 sparged CH3CN solution with analytically pure 0.1 M n-Bu4NPF6 as the supporting electrolyte. The working electrode was regenerated and cleaned between measurements by polishing with alumina paste and sonicating in H2O and acetone, successively. Measurements were performed at a 1 mM concentration of analyte except in the case of the Re monometallic species which were run at lower concentration due to limited solubilities in acetonitrile.    142 5.8.4 Photocatalytic Setup  Figure 5-9 (a) Photocatalytic setup under irradiation. (b) LED housing used for photocatalytic experiments. (c) Emission profile of the LED source (peak wavelength = 503 nm, FWHM = 30 nm).  Figure 5-9(a) shows the photocatalytic setup used, consisting of a custom-made housing (Figure 5-9(b)) utilizing two LUMEX “Emerald Green” LEDs (SSL-LX5095UEGC, peak wavelength = 503 nm, FWHM = 30 nm, Figure 5-9(c)) wired in series and driven by an AC adaptor (YHi 898-1015-U12S) giving a continuous DC supply of 12.31 V. The housing was attached to a borosilicate cuvette (0.8 cm pathlength) which itself is fused to a Schlenk flask giving a total volume of 74 mL, from which headspace analysis could be conducted. During photocatalysis a 12 cm fan (CoolerMaster A12025-12CB-3BN-F1) running with an input of 9 V was used to maintain   143 cuvette temperature. A discussion on the determination of the incident light intensity using an Aberchrome 670 actinometer can be found in Chapter 6:Listing D.1.  5.8.5 Synthesis di(1,10-phenanthrolin-5-yl)sulfane (PSP)  While under an N2 atmosphere, a solution of Li2S (0.591 g, 12.8 mmol, 1.2 eq.) in tert-butanol (250 mL) was added dropwise to a suspension of 5,6-epoxy-5,6-dihydro-[1,10]-phenanthroline (2.10 g, 10.7 mmol, 1.0 eq.) in tert-butanol (60 mL) at 60 ºC over a 40 minute period. The reaction was monitored by TLC (silica gel; EtOAc:MeOH:NH4OH 10:1:0.5) and subsequently stirred for 18 h at 60 ºC, following which the solvent was removed under reduced pressure and the crude product dissolved in CHCl3 (300 mL) and washed with sat. NH4Cl (40 mL). The aqueous layer was then further extracted with CHCl3 (3 × 100 mL) and the combined organics washed with brine (2 × 100 mL). The organic layer was dried with MgSO4, filtered and the solvent removed in vacuo. The resulting crude solid was recrystallized from MeOH:H2O (2:1) giving off-white needles (2.72 g, 6.95 mmol, 65%). The spectroscopic and spectrometric data matches that of the literature.212 1H NMR (400 MHz, CDCl3) δ 9.30 (dd, J = 4.4, 1.6 Hz, 2H), 9.23 (dd, J = 4.4, 1.7 Hz, 2H), 8.79 (dd, J = 8.4, 1.6 Hz, 2H), 8.05 (dd, J = 8.2, 1.7 Hz, 2H), 7.74 – 7.68 (m, 4H), 7.64 (dd, J = 8.1, 4.4 Hz, 2H). HR-ESI MS: m/z calcd. for C24H15N4S: 391.1017; Found: 391.1015 [M+H]+.     NNS NN  144 5,5´-sulfonylbis(1,10-pheananthroline) (PSO2P)  To EtOH (50 mL) was added PSP (0.496 g, 1.27 mmol, 1.0 eq.), NbC (0.033 g, 0.320 mmol, 0.25 eq.) and 30% H2O2 solution (1.90 mL, 19.1 mmol, 15.0 eq.). The reaction mixture was then heated to 60 ºC for 18 h. After cooling to room temperature, sat. Na2S2O3 solution (25 mL) was added, followed by H2O (75 mL), forming a white precipitate which was extracted with EtOAc (3 × 30 mL). The combined organics were washed with H2O (30 mL) and brine (30 mL), dried over MgSO4, filtered and the solvent removed under reduced pressure to give a white solid that required no further purification (0.510 g, 1.21 mmol, 95%). 1H NMR (400 MHz, CDCl3d) δ 9.37 (dd, J = 4.4, 1.7 Hz, 2H, H3), 9.21 (dd, J = 4.3, 1.6 Hz, 2H, H12), 9.05 (dd, J = 8.6, 1.6 Hz, 2H, H10), 8.90 (s, 2H, H8), 8.46 (dd, J = 8.1, 1.7 Hz, 2H, H5), 7.81 (dd, J = 8.1, 4.4 Hz, 2H, H4), 7.63 (dd, J = 8.6, 4.3 Hz, 2H, H11). 13C{1H} NMR (101 MHz, CDCl3) δ 153.86 (C3), 151.37 (C12), 148.14 (C1), 146.58 (C14), 138.00 (C5), 134.83 (C8), 133.06 (C10), 131.98 (C7), 125.94 (C6), 124.41 (C4), 123.80 (C11), 123.59 (C9). IR (neat): ṽ (σ SO2) 1288 and 1126 cm−1. HR-EI MS: m/z calcd. for C24H14N4O2S: 422.0837; Found: 422.0834 [M]+.   [Ru(phen)2(PSP)][PF6]2 (Ru-PSP)  To a microwave vial was added Ru(phen)2Cl2 (0.470 g, 0.883 mmol, 1.0 eq.), and PSP (0.383 g, 0.981 mmol, 1.11 eq.) and ethylene glycol (11 mL). The suspension was reacted in a microwave for 40 min at 200 ºC and cooled to room temperature, following which the solvent was diluted with H2O NNS NNOO1234561078914111312NNRuNNNNNN123456107891411131215 16201718192122232425262728293031323334353637 38394041424344454647484950515253545556S[PF6]2  145 (80 mL) and NH4PF6 added to precipitate the crude compound. The resulting orange precipitate was filtered and washed with H2O and Et2O and then purified over silica gel (F60) using CH3CN:H2O:KNO3 (aq.) (88:8:10) as the eluting solvent. The first major orange band was collected. After an ion exchange with NH4PF6 the pure product was isolated as an orange solid (0.439 g, 0.385 mmol, 44%). 1H NMR (600 MHz, CD3CN) δ 9.20 (dd, J = 4.3, 1.7 Hz, 1H, H12), 9.18 (dd, J = 4.4, 1.6 Hz, 1H, H3), 8.97 (dd, J = 8.5, 1.2 Hz, 1H, H47), 8.74 (dd, J = 8.2, 1.7 Hz, 1H, H5), 8.59 (tdd, J = 15.9, 8.3, 1.3 Hz, 4H, H19,24,33,38), 8.37 (dd, J = 8.0, 1.7 Hz, 1H, H10), 8.35 (s, 1H, H8), 8.26 (d, J = 0.7 Hz, 2H, H35,36), 8.23 (d, J = 2.1 Hz, 2H, H21,22), 8.12 (dd, J = 8.4, 1.2 Hz, 1H, H52), 8.09 (dd, J = 5.3, 1.2 Hz, 1H, H45), 8.07 (dd, J = 5.2, 1.3 Hz, 1H, H17), 8.04 (dd, J = 5.2, 1.2 Hz, 1H, H40), 8.01 (dd, J = 5.2, 1.3 Hz, 1H, H31), 7.98 (dd, J = 5.4, 1.3 Hz, 1H, H26), 7.90 (dd, J = 5.3, 1.3 Hz, 1H, H54), 7.78 (dd, J = 8.1, 4.3 Hz, 1H, H11), 7.72 – 7.64 (m, 5H, H4,18,39,46,50), 7.62 (dd, J = 8.2, 5.2 Hz, 1H, H32), 7.59 (dd, J = 8.2, 5.3 Hz, 1H, H25), 7.42 (dd, J = 8.4, 5.2 Hz, 1H, H53). 13C{1H} NMR (151 MHz, CD3CN) δ 154.47 (C45), 154.10 (C40), 154.04 (C17), 153.92 (C31), 153.86 (C26), 153.45 (C54), 152.41 (C12), 151.61 (C3), 149.38 (C43), 148.81 (C15,28,29,42), 148.11 (C1,56), 147.79 (C14), 137.78 (C19,24,33,38), 137.22 (C10), 136.56 (C8), 136.41 (C52), 134.82 (C5), 134.40 (C47), 131.95 (C20,23,34,37), 131.58 (C51), 130.48 (C49), 129.60 (C9), 129.18 (C6,48), 129.01 (C35,36), 128.98 (C21,22), 127.55 (C50), 127.43 (C7), 127.13 (C53), 126.95 (C46), 126.83 (C18,39), 126.80 (C25,32), 124.90 (C11), 124.76 (C4). MALDI-TOF MS: m/z calcd. for C48H30F6N8PRuS: 991.1028; Found: 991.1032 [M–PF6]+.       146 [Ru(phen)2(PSO2P)][PF6]2 (Ru-PSO2P)  Ru(phen)2Cl2 (0.284 g, 0.533 mmol, 1.0 eq.) and PSO2P (0.250 g, 0.591 mmol, 1.1 eq.) and ethylene glycol (15 mL) were heated in a microwave at 200 ºC for 40 min. The reaction was cooled to room temperature and the ethylene glycol diluted with H2O (80 mL). NH4PF6 was added to a stirred solution of the reaction mixture, following which the precipitate was filtered and washed with H2O and Et2O. The crude product was purified over a silica gel column (F60) using CH3CN:H2O:KNO3 (aq.) (88:8:10) as the eluting solvent where the first major orange band was collected, which after anion exchange with NH4PF6 yielded the pure complex as an orange powder (0.130 g, 0.111 mmol, 21%). 1H NMR (600 MHz, CD3CN) δ 9.39 (s, 1H, H8), 9.25 (d, J = 4.4 Hz, 2H, H50,54), 9.19 (dd, J = 8.8, 1.2 Hz, 1H, H47), 9.14 – 9.11 (m, 2H, H3,5), 8.81 (dd, J = 8.3, 1.3 Hz, 1H, H10), 8.72 (dd, J = 8.2, 1.8 Hz, 1H, H52), 8.60 – 8.52 (m, 4H, H19,24,32,38), 8.22 (s, 2H, H21,36), 8.20 (s, 2H, H22,35), 8.14 (dd, J = 5.4, 1.3 Hz, 1H, H12), 7.99 (dd, J = 5.4, 1.2 Hz, 1H, H45), 7.93 (ddd, J = 6.9, 5.2, 1.3 Hz, 3H, H17,31,40), 7.91 (dd, J = 5.3, 1.3 Hz, 1H, H26), 7.87 (dd, J = 8.1, 4.3 Hz, 1H), H53, 7.71 (m, 2H, H4,11), 7.59 (dd, J = 8.4, 5.4 Hz, 1H, H39), 7.57 – 7.53 (m, 3H, H25,32,46), 7.51 (dd, J = 8.3, 5.2 Hz, 1H, H18). 13C{1H} NMR (151 MHz, CD3CN) δ 156.95 (C12), 154.90 (C45), 154.70 (C54), 154.21 (C17), 154.19 (C40), 153.86 (C31), 153.77 (C26), 151.85 (C3), 151.14 (C14), 149.87 (C43), 148.92 (C56), 148.70 (C15,42), 148.64 (C28,29), 147.33 (C1), 139.70 (C52), 139.31 (C10), 137.91 (C19,38), 137.88 (C24), 137.85 (C33), 137.73 (C49), 135.00 (C50), 134.56 (C47), 134.40 (C8), 134.05 (C5), 131.96 (C30,37), 131.90 (C23,34), 129.61 (C9), 128.99 (C21,36), 128.97 (C22,35), 127.84 (C11), 127.39 (C46), 127.25 (C48), 127.19 (C51), 126.82 (C32), 126.80 (C25), 126.69 (C34), 126.66 (C18), NNRuNNNNNN123456107891411131215 16201718192122232425262728293031323334353637 38394041424344454647484950515253545556S[PF6]2O O  147 125.46 (C53), 124.83 (C4), 124.41 (C7). IR (neat): ṽ (σ SO2) 1312 and 1136 cm−1. MALDI-TOF MS: m/z calcd. for C48H30F6N8O2PRuS: 1023.0930; Found: 1023.0928 [M–PF6]+.  Re(PSP)(CO)3Br (Re-PSP) PSP (0.231 g, 0.599 mmol, 1.32 eq.) was dissolved in refluxing MeOH (250 mL), to which Re(CO)5Br (0.185 g, 0.455 mmol, 1.0 eq.) in 90 mL MeOH was added dropwise. The solution left to reflux for 18 h in the dark, following which the reaction mixture was cooled to room temperature and the solvent removed via rotary evaporation. The crude powder was purified using column chromatography (silica, F60) first using CH3CN, followed by 98:2 CH2Cl2:MeOH, giving a yellow solid. (0.149 g, 0.201 mmol, 44%). 1H NMR (600 MHz, CDCl3) δ 9.50 (ddd, J = 8.3, 5.1, 1.3 Hz, 1H, H3), 9.33 (dd, J = 4.4, 1.7 Hz, 1H, H12), 9.30 (dd, J = 4.3, 1.6 Hz, 1H, H31), 9.27 (ddd, J = 7.1, 5.0, 1.4 Hz, 1H, H40), 9.07 (ddd, J = 8.5, 5.6, 1.4 Hz, 1H, H5), 8.69 (ddd, J = 8.4, 6.7, 1.7 Hz, 1H, H33), 8.29 (dd, J = 8.1, 1.7 Hz, 1H, H10), 8.25 (s, 1H, H8), 8.05 (td, J = 8.4, 1.4 Hz, 1H, H38), 7.98 (ddd, J = 8.4, 5.1, 3.2 Hz, 1H, H4), 7.75 (dd, J = 8.0, 4.4 Hz, 1H, H11), 7.70 (ddd, J = 8.4, 6.0, 4.5 Hz, 2H, H32,39), 7.24 (s, 1H, H36). 13C{1H} NMR (151 MHz, CDCl3) δ 153.86 (C3), 152.83 (C40), 152.08 (C12), 151.54 (C31), 147.44 (C1), 146.82 (C14), 146.70 (C29), 146.06 (C38), 136.83 (C42), 136.66 (C10), 136.11 (C8), 134.90 (C3), 134.24 (C33), 130.42 (C37), 129.49 (C7), 128.44 (C35), 128.41 (C9), 126.36 (C39), 126.10 (C4), 125.28 (C3g), 124.44 (C31), 124.20 (C11). IR (neat): ṽ (σ CO) 2019, 1884 and 1866 cm–1. HR-EI MS: m/z calcd. for C27H15BrN4O3ReS: 738.9578; Found: 738.9575 [M]+.   1 2456107891411NNSNN2930313233343536373839404142 ReCOCOCOBr13123  148 Re(PSO2P)(CO)3Br (Re-PSO2P)  To a MeOH (250 mL) solution heated to reflux of PSO2P (0.239 g, 0.566 mmol, 1.48 eq.) was added dropwise a MeOH (90 mL) solution of Re(CO)5Br (0.155 g, 0.382 mmol, 1.0 eq.). The reaction was heated at reflux for 18 h in the absence of light, following which the solution was cooled to room temperature and the solvent removed in vacuo. The crude solid was then purified over silica gel (F60) using first CH3CN followed by 98:2 CH2Cl2:MeOH as the eluting solvents, yielding an orange powder. (0.145 g, 0.188 mmol, 49%). 1H NMR (600 MHz, CDCl3) δ 9.53 (ddd, J = 8.6, 5.2, 1.4 Hz, 1H, H12), 9.48 – 9.44 (m, 1H, H3), 9.43 – 9.39 (m, 2H, H5,40), 9.28 (dd, J = 4.3, 1.6 Hz, 1H, H31), 9.04 (ddd, J = 8.6, 3.4, 1.6 Hz, 1H, H33), 8.97 (d, J = 2.3 Hz, 1H, H36), 8.81 (d, J = 1.3 Hz, 1H, H8), 8.65 (ddd, J = 8.2, 6.3, 1.4 Hz, 1H, H10), 8.53 (dt, J = 8.2, 1.8 Hz, 1H, H38), 7.98 (ddd, J = 8.4, 5.1, 3.7 Hz, 1H, H11), 7.92 (ddd, J = 8.7, 5.1, 4.0 Hz, 1H, H4), 7.85 (dd, J = 8.1, 4.4 Hz, 1H, H39), 7.72 (ddd, J = 8.6, 4.3, 1.5 Hz, 1H, H32). 13C{1H} NMR (151 MHz, CDCl3) δ 156.07 (C12), 154.47 (C40), 154.43 (C3), 151.82 (C31), 148.83 (C14), 148.26 (C42), 147.73 (C1), 139.59 (C10), 138.43 (C38), 137.68 (C7), 135.71 (C5), 133.36 (C36), 133.29 (C33), 133.16 (C35), 131.22 (C8), 128.24 (C9), 127.11 (C11), 126.80 (C4), 126.37 (C6), 125.97 (C37), 124.81 (C39), 124.38 (C32), 123.64 (C34). IR (neat): ṽ (σ SO2) 1304 and 1126 cm−1; ṽ (σ CO) 2016, 1884 and 1865 cm–1. HR-EI MS: m/z calcd. for C27H15BrN4O5ReS: 770.9476; Found: 770.9484 [M]+.      1 2456107891411NNSNN2930313233343536373839404142OOReCOCOCOBr13123  149 [Ru(phen)2(PSP)Re(CO)3Br][PF6]2 (Ru-PSP-Re)  A round-bottomed flask was charged with Ru-PSP (0.426 g, 0.373 mmol, 1.0 eq.) and Re(CO)5Br (0.181 g, 0.448 mmol, 1.2 eq.) and placed under an N2 atmosphere. A 9:1 MeOH:EtOH (500 mL) solution was added and the reaction mixture refluxed for 18 h in the dark then cooled to room temperature, following which the solvent was removed in vacuo. The crude product was purified over silica gel (F60) with a CH3CN:H2O:KNO3 (aq.) (95:5:5) solvent system. The nitrate salt of the complex underwent a salt metathesis with NH4PF6 and was filtered, yielding an orange powder. (0.466 g, 0.299 mmol, 80%). 1H NMR (850 MHz, CD3CN) δ 9.49 (dddd, J = 16.1, 10.8, 5.0, 1.3 Hz, 2H), 9.43 (dddd, J = 11.6, 8.0, 5.0, 1.4 Hz, 2H), 9.19 – 9.06 (m, 2H), 8.96 – 8.89 (m, 2H), 8.70 – 8.52 (m, 11H), 8.37 – 8.25 (m, 10H), 8.21 (s, 1H), 8.19 (s, 1H), 8.14 – 8.07 (m, 7H), 8.05 – 7.96 (m, 10H), 7.72 – 7.68 (m, 4H), 7.68 – 7.62 (m, 6H), 7.57 (dddd, J = 14.2, 9.3, 5.6, 1.2 Hz, 2H). 13C{1H} NMR (214 MHz, CD3CN) δ 156.28, 156.27, 155.11, 155.10, 155.02, 154.95, 154.93, 154.87, 154.73, 154.68, 154.66, 149.81, 149.75, 149.08, 148.88, 148.87, 148.86, 148.85, 148.66, 148.63, 148.62, 148.27, 148.25, 147.91, 147.85, 147.83, 140.56, 140.53, 139.60, 139.57, 139.52, 139.47, 138.50, 138.48, 137.89, 137.88, 137.86, 137.85, 137.84, 137.52, 137.47, 137.44, 137.38, 136.98, 136.90, 136.88, 135.05, 134.96, 134.87, 133.83, 133.77, 133.75, 133.69, 133.64, 133.61, 133.36, 133.29, 133.28, 133.22, 133.16, 132.68, 132.59, 132.53, 132.45, 132.42, 132.38, 132.19, 132.16, 132.07, 132.06, 132.04, 132.02, 132.01, 131.69, 131.59, 131.56, 131.51, 131.49, 131.45, 131.43, 131.29, 131.21, 131.16, 131.08, 130.92, 129.10, 129.08, 129.05, 128.20, 128.20, 128.10, 128.09, 127.78, 127.75, 127.72, 127.66, 127.64, 127.63, 127.60, 127.41, 127.39, 127.37, 127.28, 127.26, 127.24, 126.91, NNS NNRuNNNN[PF6]2ReCOCOCOBr  150 126.90, 126.88, 126.85. IR (neat): ṽ (σ CO) 2022, 1913 and 1894 cm–1. MALDI-TOF MS: m/z calcd. for C51H30BrF6N8O3PRuReS: 1346.9588; Found: 1346.9583 [M–PF6]+.  [Ru(phen)2(PSO2P)Re(CO)3Br][PF6]2 (Ru-PSO2P-Re)  Under an N2 atmosphere, and in the absence of light, Ru-PSO2P (0.127 g, 0.108 mmol, 1.0 eq.) and Re(CO)5Br (0.053 g, 0.130 mmol, 1.2 eq.) were heated to reflux for 18 h in a 9:1 MeOH:EtOH (500 mL) solution. Following cooling, the solvents were removed under reduced pressure and the crude residue purified using column chromatography over silica gel (F60) with CH3CN:H2O:KNO3 (aq.) (95:5:5) as the eluting solvent. The product was converted to the hexafluorophosphate salt through addition of NH4PF6 (aq.) and then filtered, yielding an orange solid. (0.129 g, 0.105 mmol, 97%). 1H NMR (850 MHz, CD3CN) δ 9.56 (dt, J = 5.2, 0.9 Hz, 1H), 9.56 – 9.54 (m, 2H), 9.53 (s, 1H), 9.50 (s, 2H), 9.46 (dt, J = 5.1, 1.0 Hz, 2H), 9.40 (tt, J = 7.7, 1.1 Hz, 2H), 9.17 (dt, J = 8.9, 0.9 Hz, 2H), 9.06 (td, J = 8.7, 1.4 Hz, 2H), 8.88 (dt, J = 8.4, 1.6 Hz, 2H), 8.62 – 8.60 (m, 3H), 8.58 (dddd, J = 6.6, 4.7, 3.3, 1.1 Hz, 4H), 8.57 – 8.55 (m, 1H), 8.26 (d, J = 0.6 Hz, 2H), 8.25 (s, 2H), 8.23 (d, J = 0.8 Hz, 2H), 8.22 – 8.18 (m, 4H), 8.14 – 8.13 (m, 1H), 8.13 – 8.11 (m, 1H), 8.03 (ddd, J = 7.0, 5.2, 1.1 Hz, 2H), 8.01 – 7.98 (m, 3H), 7.97 (td, J = 5.3, 1.5 Hz, 3H), 7.95 – 7.93 (m, 4H), 7.77 (d, J = 5.5 Hz, 1H), 7.76 (d, J = 5.4 Hz, 1H), 7.64 – 7.59 (m, 7H), 7.59 – 7.57 (m, 1H), 7.57 – 7.55 (m, 1H), 7.55 – 7.53 (m, 1H). 13C{1H} NMR (214 MHz, CD3CN) δ 159.32, 158.88, 157.06, 157.05, 156.70, 156.68, 155.88, 155.85, 155.81, 155.77, 155.53, 155.49, 155.46, 155.45, 153.10, 153.07, 151.62, 151.43, 150.37, 150.35, 150.34, 150.32, 150.31, 150.29, 150.15, 143.84, 143.80, NNS NNRuNNNN[PF6]2ReCOCOCOBrO O  151 141.15, 139.65, 139.62, 139.59, 139.58, 139.55, 139.54, 138.57, 138.55, 138.51, 138.50, 138.44, 138.06, 138.05, 137.21, 137.10, 137.04, 137.02, 136.12, 136.04, 135.99, 133.65, 133.63, 133.62, 133.58, 131.31, 131.29, 131.06, 131.04, 130.67, 130.66, 130.63, 130.57, 129.98, 129.95, 129.61, 129.59, 129.56, 129.54, 129.26, 129.20, 128.85, 128.84, 128.71, 128.67, 128.50, 128.48, 128.47, 128.38, 128.37, 128.32, 128.28. IR (neat): ṽ (σ SO2) 1322 and 1140 cm−1; ṽ (σ CO) 2022, 1914 and 1889 cm–1. MALDI-TOF MS: m/z calcd. for C51H30BrF6N8O5PRuReS: 1373.9474; Found: 1373.9476 [M–PF6]+.   152 CHAPTER 6: Conclusions and Future Work  The oxidation of sulfur atoms can be achieved simply and safely by a variety of methods with the use of low-cost, commercially available reagents. While at face value this appears to be a subtle structural change, in actuality the effect on the electronic properties is far from small. It is hoped that through this thesis other scientists are inspired to probe and manipulate this versatile sulfur atom in order to tailor new materials for a variety of photophysical applications. In Chapter 2 six new cyclometalated iridium(III) complexes bearing sulfur-bridged dipyridyl ligands were synthesized and characterized. In changing the sulfur center from sulfide (S), to sulfoxide (SO) to sulfone (SO2) it was revealed that this oxidation can serve as a switch to alter the emission of the complex from that of blue/green to yellow. This leads to a new class of ligand capable of a “two-level” tuning of the photoluminescence colour, whereby a combination of substituents at the pyridyl rings and sulfur oxidation state can be used to finely modify the electronics. While the PLQY for the complexes was found to be low, thereby limiting use in emissive devices, the additional degree of control over the excited-state properties is expected to be advantageous for applications in photoredox catalysis. The colour-tuning ability could be a useful property for fluorescence microscopy applications probing oxidative pathways, whereby the sulfide (with blue-green emission) is oxidized to the sulfone in vitro to give a yellow emissive species allowing the determination of oxygen-rich and oxygen-poor cell environments. With a view to creating lower-cost emitters, in Chapter 3 heteroleptic copper(I) complexes bearing POP diphosphine ligands and sulfur-bridged diimine ligands were prepared.   153 The sulfur-bridged dipyridyl ligands were functionalized either with or without methyl substituents in the 6,6′-positions of the pyridyl rings and sulfide or sulfone oxidation states at the sulfur center. Ligands without methyl functionalization formed monometallic species, but once the bulky substituents were introduced bimetallic adducts were isolated due to steric constraints about the copper. For the sulfide (Cu-Me-DPS), a bimetallic species consisting of a bridging POP ligand is seen, but the sulfone (Cu-Me-DPSO2) forms a bimetallic species with the diimine moiety acting as a bridging ligand linking the two copper atoms due to the additional binding sites at the sulfone oxygen atoms. These arrangements were observed in solution and as single crystal structures. PLQYs were modest, up to 0.14 for the most sterically constrained complex, Cu-Me-DPSO2, Through the use of low-temperature time-resolved photoluminescence spectroscopy, in addition to computational methods, the emissive pathway was determined to be due to a TADF mechanism. Cu-Me-DPSO2 could show promise as an emissive material in a LEEC device, and further structural modifications at the 6- and 6′-positions of the diimine ligand could be of interest to see if PLQYs can be increased through additional steric constraints about the copper(I) atom. The ability to control or tune photoluminescence properties is advantageous in the design of photofunctional materials, particularly those which react to an external stimulus. Further illustrating subtle alterations can result in dramatic changes in properties, in Chapter 4, two new [Cu(diphosphine)(diimine)]+ complexes bearing sulfoxide-bridged diimine ligands are presented. Cu-DPSO displayed photophysical properties typical of other Cu(I) complexes. However, on addition of methyl substituents in the 6- and 6′-positions of the diimine ligand to give complex Cu-Me-DPSO, we were able to access a variety of different photophysical properties. Cu-Me-DPSO is initially isolated as an amorphous solid which under UV irradiation gives a yellow emission, but upon heating to 180 ºC, a crystalline species (c-Cu-Me-DPSO) is formed which   154 gives orange emission due to a flattening conformational change during the crystallization process. c-Cu-Me-DPSO was also shown to exhibit reversible thermochromic emission, whereby the complex undergoes a large colour change from orange at room temperature to yellow at –196 ºC. Based on data from solid state variable-temperature excitation and absorption spectra this phenomenon is attributed to a change in coordination geometry about the copper atom in the excited state. At low temperatures, a pseudo-tetrahedral geometry is preferred, giving higher energy emission, while at higher temperatures a flattened geometry dominates giving a lower energy emission under UV irradiation. In showing that a copper(I) complex can exhibit reversible thermochromism, avenues that were inaccessible due to the limitations of thermochromic copper halide clusters can now be realized. For example, inexpensive copper(I) metallopolymers that offer a response to thermal stimuli could be fabricated. The simple synthetic procedures to the complexes reported, coupled with the absence of toxic heavy metals required, lends itself to practical applications in optoelectronics, including light-emitting diodes, display technologies and chemical and biological sensors. In Chapter 5 the use of supramolecular ruthenium(II)-rhenium(I) species as photocatalysts for CO2 reduction is introduced. In these systems the Ru(II) moiety acts as a visible-light absorbing photosensitizer (PS) unit, and the Re(I) as the catalytic center. The reaction mechanism proceeds by the selective photoexcitation of a one-electron reduced species (OERS) at the Ru(II) center which then proceeds via an intramolecular electron transfer to the Re(I) catalytic center, where CO2 reduction then occurs. When the two metal centers are tethered together the electron transfer is accelerated. Previous research in the Wolf group has shown that sulfur oxidation state can vary the nature of charge transfer in sulfur-bridged organic chromophores, and   155 so this project was designed to probe the degree of electron transfer in a tethered Ru(II)-Re(I) photocatalytic system and thus the impact on CO2 reduction. In using a sulfur-bridged bisphenanthroline ligand to connect the two metal centers, two new Ru(II)-Re(I) supramolecular photocatalysts were prepared with sulfide (Ru-PSP-Re) and sulfone (Ru-PSO2P-Re) oxidation states at the sulfur, in addition to their monometallic Ru(II) and Re(I) analogues. Electrochemical data were used to indicate the suitability of each dyad towards CO2 reduction. It has been reported that a threshold for the photocatalytic activity of mononuclear Re complexes exists at E1/2red = –1.4 V, and this translates to Ru(II)-Re(I) supramolecular catalysts with far lower efficiencies for catalysts at potentials more positive than this value. Reduction potentials for Ru-PSP-Re (–1.56 V) and Ru-PSO2P-Re (–1.30 V) suggest that the sulfide species is a better candidate for the photocatalysis. However, on analysis of the photophysical data, a bathochromic shift of 25 nm and a decrease in emission intensity is observed on going from Ru-PSP-Re to Ru-PSO2P-Re indicating a greater degree of electronic communication between the two metal centers for the sulfone species, which could aid in the efficiency of the electron transfer of the OERS. PLQYs are required to quantify the degree of emission quenching, and transient absorption measurements could aid in elucidating the nature of the excited-state. While the TON of both complexes for formation of CO was found to be very low, significant bubbling was observed from the solution under irradiation to the point of pressurizing the sealed vessel. From GC data it is seen that this gas is not CO or H2, the two volatile products typically produced in these systems, and from solvent suppression 1H NMR studies no formic acid was produced either. In order to probe whether the gas formed is a decomposition product of the catalyst, atmospheric MS will be performed in real-time, scanning 0–100 m/z to see if any volatile sulfur-containing species are evolved. Atmospheric-MS can also help identify other possible low   156 mass products, such as ethane or ethylene. Additionally, GC-MS will be conducted for further analysis of the gas produced. Overall, the work discussed in this thesis illustrates that through altering the oxidation state at sulfur can change the photophysical properties of a variety of metals, with copper, iridium, rhenium and ruthenium highlighted in this body of work, and as a consequence the complexes reported have the potential to be used in a wide variety of applications, from light-emitting applications such as OLEDs and LEECs, to photocatalytic reduction of CO2, to photocatalysis and uses in bioimaging. 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Soc. 2015, 137, 2738–2747.   174 APPENDIX  Appendix A: Tunable Emission of Iridium(III) Complexes Bearing Sulfur-Bridged Dipyridyl Ligands  Listing A.1 Frontier orbital diagrams for Ir-DPS, Ir-DPSO and Ir-DPSO2.    175 Listing A.2 Frontier orbital diagrams for Ir-4,4'-Me-DPS, Ir-4,4'-Me-DPSO and Ir-4,4'-Me-DPSO2.    176 Listing A.3 Unpaired-electron spin density contours for Ir(III) complexes.      177 Listing A.4 Lowest triplet excited states calculated at the TD-DFT B3LYP/(6-31G**+LANL2DZ) level. Vertical excitation energies (E), dominant monoexcitations with contributions (within parentheses) greater than 15%, nature of the electronic transition and description of the excited state are summarized for Ir(III) complexes.  Monoexcitation Energy Nature Description Ir-DPS H→ L+2 (55 %) 2.76 eV d π(Ir) + πppy → π*ppy 3LC/ 3MLCT Ir-DPSO H→ L+2 (26 %) 2.77 eV d π(Ir) + πppy → π*ppy 3LC/ 3MLCT Ir-DPSO2 H→ L (90 %) 2.65 eV d π(Ir) + πppy → π*anc 3MLCT/ 3LLCT Ir-4,4'-Me-DPS H→ L (54 %) 2.76 eV d π(Ir) + πppy → π*ppy 3LC/ 3MLCT Ir-4,4'-Me-DPSO H→L +1 (46 %) 2.76 eV d π(Ir) + πppy → π*ppy 3LC/ 3MLCT Ir-4,4'-Me-DPSO2 H→ L (42 %) 2.74 eV d π(Ir) + πppy → π*anc 3MLCT/ 3LLCT      178 Listing A.5 Summary of crystallographic data for Ir-DPS. Compound  Ir-DPS Empirical formula  C16H12F3Ir0.50N2P0.50S0.50 Formula weight  416.89 Temperature  173(2) K Wavelength  0.71073 Å Crystal system  Triclinic Space group  P -1 Unit cell dimensions a = 14.7218(16) Å a= 67.320(5)º.  b = 15.0655(17) Å b= 76.875(5)º.  c = 16.9581(19) Å g = 73.592(5)º. Volume 3299.2(7) Å3 Z 8 Density (calculated) 1.679 Mg/m3 Absorption coefficient 4.221 mm-1 F(000) 1624 Crystal size 0.302 x 0.100 x 0.020 mm3 Theta range for data collection 1.313 to 26.482º. Index ranges -18<=h<=17, -18<=k<=18, -21<=l<=17 Reflections collected 46631 Independent reflections 13136 [R(int) = 0.0269] Completeness to theta = 25.242º 96.7 %  Refinement method Full-matrix least-squares on F2 Data / restraints / parameters 13136 / 0 / 835   179 Goodness-of-fit on F2 1.055 Final R indices [I>2sigma(I)] R1 = 0.0351, wR2 = 0.0887 R indices (all data) R1 = 0.0433, wR2 = 0.0945 Extinction coefficient n/a Largest diff. peak and hole 2.791 and -1.682 e.Å-3      180  Appendix B:   Listing B.1 Summary of crystallographic data for Cu-DPS Compound Cu-DPS Chemical formula C46H36CuN2OP2SBF4 Formula weight 877.13 a (Å) 17.6927(7) b (Å) 13.6703(5) c (Å) 17.7523(7) a (deg) 90 β (deg) 105.7929(10) γ (deg) 90 Unit cell volume (Å3) 4131.6(3) Temperature (K) 90 Crystal system Monoclinic Space group P 21/n Z 4 Absorption coefficient (μ/mm-1) 0.713 Total number of reflections 72233 Unique reflections 12659 Final R1 values (I > 2s(I) ) 0.0351 Final wR(F2) values (all data) 0.0941    181 Listing B.2 Summary of crystallographic data for Cu-DPSO2. Compound Cu-DPSO2 Chemical formula C46H36CuN2O3P2S, BF4 Formula weight 909.13 a (Å) 11.9139(4) b (Å) 13.1490(5) c (Å) 15.6199(6) a (deg) 96.3361(9) β (deg) 112.0557(11) γ (deg) 93.6302(9) Unit cell volume (Å3) 2239.40(14) Temperature (K) 90 Crystal system Triclinic Space group P -1 Z 2 Absorption coefficient (μ/mm-1) 0.664 Total number of reflections 33113 Unique reflections 13668 Final R1 values (I > 2s(I) ) 0.0295 Final wR(F2) values (all data) 0.0767      182 Listing B.3 Summary of crystallographic data for Cu-Me-DPS. Compound Cu-Me-DPS Chemical formula C60H52Cu2N4OP2S2, 2(BF4), 2(CH2Cl2) Formula weight 1441.69 a (Å) 16.8444(16) b (Å) 13.7286(13) c (Å) 27.057(2) a (deg) 90 β (deg) 90 γ (deg) 90 Unit cell volume (Å3) 6256.9(10) Temperature (K) 90 Crystal system Orthorhombic Space group P b c n Z 4 Absorption coefficient (μ/mm-1) 1.038 Total number of reflections 13008 Unique reflections 7149 Final R1 values (I > 2s(I) ) 0.0475 Final wR(F2) values (all data) 0.1141      183 Listing B.4 Summary of crystallographic data for Cu-Me-DPSO2. Compound Cu-Me-DPSO2 Chemical formula C84H68Cu2N2O4P4S, 2(BF4) Formula weight 1626.06 a (Å) 23.1483(9) b (Å) 27.8735(11) c (Å) 14.5524(6) a (deg) 90 β (deg) 91.3552(9) γ (deg) 90 Unit cell volume (Å3) 9386.9(6) Temperature (K) 90 Crystal system Monoclinic Space group C 2/c Z 4 Absorption coefficient (μ/mm-1) 0.602 Total number of reflections 56010 Unique reflections 10825 Final R1 values (I > 2s(I) ) 0.0401 Final wR(F2) values (all data) 0.1103      184 Listing B.5 Variable temperature emission lifetime plots for Cu-DPS, Cu-DPSO2, Cu-Me-DPS and Cu-Me-DPSO2 from –196 to 28 °C.   All samples drop-cast from MeOH on to quartz slides. Excitation with a NanoLED source (370 nm) with a pulse width of 1.2 ns repeating at a rate of 10 kHz.    185 Appendix C: Thermochromic Solid state Emission of Dipyridyl Sulfoxide Copper(I) Complexes  Listing C.1 Summary of crystallographic data for Cu-DPSO. Compound Cu-DPSO Chemical formula 2(C46H36CuN2O2P2S), 2(BF4) Formula weight 1786.34 a (Å) 10.2401(7) b (Å) 19.7604(14) c (Å) 21.4189(15) a (deg) 101.9908(16) β (deg) 99.5004(16) γ (deg) 99.8941(15) Unit cell volume (Å3) 4084.8(5) Temperature (K) 90 Crystal system Triclinic Space group P -1 Z 2 Absorption coefficient (μ/mm-1) 0.725 Total number of reflections 36823 Unique reflections 18738 Final R1 values (I > 2s(I) ) 0.0413 Final wR(F2) values (all data) 0.1013    186 Listing C.2 Summary of crystallographic data for Cu-Me-DPSO. Compound Cu-Me-DPSO Chemical formula C48H40CuN2O2P2S, BF4, CH2Cl2 Formula weight 1006.15 a (Å) 12.4822(11) b (Å) 14.5739(14) c (Å) 15.8204(14) a (deg) 74.485(2) β (deg) 88.104(2) γ (deg) 72.284(2) Unit cell volume (Å3) 2637.9(4) Temperature (K) 90 Crystal system Triclinic Space group P -1 Z 2 Absorption coefficient (μ/mm-1) 0.667 Total number of reflections 48576 Unique reflections 22839 Final R1 values (I > 2s(I) ) 0.0662 Final wR(F2) values (all data) 0.2034      187 Appendix D: Sulfur-Bridged Bimetallic Ru(II)-Re(I) Systems – Modulating CO2 Reduction with Sulfur Oxidation State  Listing D.1 Actinometry for the photocatalytic LED setup. Light intensities at 503 nm were determined using the commercial furyl fulgide Aberchrome 670218 (TCI) following the procedure of Herder et al. 219 A toluene solution of Aberchrome 670 (4 mL, 1.4 × 10–4 M) was irradiated in a quartz cell (Starna Cells, 1 cm path length) for 5 min using a 365 nm hand lamp. Irradiation at 503 nm using the LED housing was conducted in 6 steps with 4 s irradiation times. I0 is calculated from the depletion of Aberchrome 670 absorbance at 519 nm:  𝐼' =	– ∆𝐴-./	01∆𝑡 	× 	 𝑉1000	 ×	𝛷-'8	01 	×	𝜀-./	01 	× 	L	 × 	(1– 10–<′)  ∆A519 nm  =  change in absorbance at 519 nm ∆t  = change in time V  = volume of Aberchrome 670 solution (mL) Φ503 nm  = 0.30 ε519 nm  = 7760 M–1 cm–1 L  = Path length of cuvette (cm) A'  = Initial absorbance at the irradiation wavelength (503 nm)  

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