Halide Photoredox Chemistry - Chemical Reviews (ACS Publications)

Review Cite This: Chem. Rev. XXXX, XXX, XXX−XXX

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Halide Photoredox Chemistry Ludovic Troian-Gautier, Michael D. Turlington, Sara A. M. Wehlin, Andrew B. Maurer, Matthew D. Brady, Wesley B. Swords, and Gerald J. Meyer*

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Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599, United States ABSTRACT: Halide photoredox chemistry is of both practical and fundamental interest. Practical applications have largely focused on solar energy conversion with hydrogen gas, through HX splitting, and electrical power generation, in regenerative photoelectrochemical and photovoltaic cells. On a more fundamental level, halide photoredox chemistry provides a unique means to generate and characterize one electron transfer chemistry that is intimately coupled with X−X bond-breaking and -forming reactivity. This review aims to deliver a background on the solution chemistry of I, Br, and Cl that enables readers to understand and utilize the most recent advances in halide photoredox chemistry research. These include reactions initiated through outer-sphere, halide-to-metal, and metal-toligand charge-transfer excited states. Kosower’s salt, 1-methylpyridinium iodide, provides an early outer-sphere charge-transfer excited state that reports on solvent polarity. A plethora of new inner-sphere complexes based on transition and main group metal halide complexes that show promise for HX splitting are described. Long-lived charge-transfer excited states that undergo redox reactions with one or more halogen species are detailed. The review concludes with some key goals for future research that promise to direct the field of halide photoredox chemistry to even greater heights.

CONTENTS 1. Introduction, Background, and Motivation 1.1. Introduction 1.2. Background 1.2.1. Formal Reduction Potentials of Halogen Species 1.2.2. Halide Spectroscopy 1.2.3. Halogen Bonding in Polyhalide Species 1.3. Motivation and Applications 1.3.1. HX Splitting 1.3.2. Regenerative Solar Cells 2. Halide Photoredox Chemistry with Metal-toLigand Charge-Transfer (MLCT) Excited States 2.1. Excited-States and Quenching Mechanisms 2.1.1. Metal-to-Ligand Charge-Transfer Excited States 2.1.2. Stern−Volmer Analysis 2.1.3. Kinetic Analysis 2.1.4. Electron-Transfer Rate Constants 2.2. Dynamic Quenching 2.2.1. Iodide Photo-Oxidation 2.2.2. Triiodide Reduction 2.2.3. Diffusional Bromide and Chloride Photo-Oxidation 2.2.4. Photocatalytic Bromide Oxidation Using Diazonium Sacrificial Agents 2.3. Static and Dynamic Halide Photoredox Chemistry 2.3.1. Ion Pairing in Low Dielectric Constant Solvents 2.3.2. Supramolecular Assemblies with Halide Ions

© XXXX American Chemical Society

2.3.3. Ruthenium Complexes with Polycationic Ligands 2.3.4. Excited-State Ion Pairs That Undergo Halide Photorelease 2.3.5. Bromide and Chloride Photo-Oxidation 3. Halide-to-Acceptor Charge Transfer 3.1. Outer-Sphere Charge Transfer 3.1.1. Ion-Paired Charge Transfer 3.1.2. Halide-to-Viologen Charge Transfer 3.1.3. Charge Transfer in Supramolecular Assemblies 3.2. Halide-to-Metal Charge Transfer 3.3. Halide Oxidation Through X2 Elimination 3.3.1. Dirhodium HX Splitting Catalysts 3.3.2. Bimetallic Catalysts for X2 Elimination 3.3.3. Main Group HX Splitting Catalysts 3.3.4. Monomeric and Bimetallic Gold and Platinum HX Splitting Catalysts 3.3.5. Main Group LMCT Halide Oxidation 3.3.6. Halogen Photoelimination from First Row Transition Metals 4. Conclusions and Future Directions 4.1. Supramolecular Complexes That Recognize and Activate Specific Halide Ions 4.2. Identification of Photoredox Chemistry That Enables X−X Bond Formation and Electron Transfer to Occur in One Concerted Step 4.3. Organic Photoredox Catalysis with Halogen Species 4.4. Stabilization of Halogen Radicals

B B B B D F J J K N N N O Q R T T U U V V V W

X Z AA AC AC AC AE AF AG AG AG AI AJ AL AM AO AO AO

AO AP AP

Received: December 4, 2018

A

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Chemical Reviews 4.5. Quantification of Halide Outer-Sphere and Charge-Transfer Absorption Bands with Mulliken−Hush Theory 4.6. Synthesis of Radical Traps Selective for X• and X2 4.7. Systematic Characterization of Halogen Species in Organic Solvents 4.8. Solid-State Characterization of Halide Photoredox Chemistry 4.9. The Other Halides Author Information Corresponding Author ORCID Notes Biographies Acknowledgments References

Review

increase intersystem crossing yields through the heavy atom effect.33−38 This review focuses on fundamental developments in the photoredox chemistry of chlorine, bromine, and iodine. The instability of astatine and the toxicity of fluorine have precluded widespread study of these halogens’ redox chemistry, which will not be detailed herein. Furthermore, this review does not detail halogen oxides, such as XO−, XO2−, XO3−, and XO4−. The remainder of this section will describe the electrochemical and spectroscopic behavior of chlorine, bromine, and iodine and their applications in solar energy conversion that provide motivation for many of the studies that are reviewed. Section 2 focuses on long-lived excited states that are capable of diffusional electron-transfer reactions with halogen species. This research was inspired by solar applications yet also provides formal reduction potentials for halogen species and kinetic rate constants in a wide variety of coordination environments. Section 3 focuses on organic and metal complexes where the halide itself is involved in light absorption and subsequent reactivity. The organic complexes are of historic interest and provide early examples of outer-sphere charge transfer within ion pairs. Studies of the metal complexes were often motivated by photoreductive elimination of halogens for catalyst regeneration in HX splitting applications. A note on the nomenclature used throughout this review. Chlorine, bromine, and iodine refer to these compounds in their elemental states, i.e., Cl2, Br2, and I2. These are distinguished from the single odd electron atoms that are referred to as the halogen atom or X•, with X being the generic symbol for any halogen. The halide ions, particularly iodide, form polyhalides such as I2•−, I3−, and I4•−, that will be referred to as diiodide, triiodide, tetraiodide, and so on. Note that these names do not indicate whether unpaired electrons are present on the ion while the chemical symbols do. The term “halogen species” is all encompassing and includes atoms, halogens, and ions. In the tables and figures, a green color is used for chlorine, red for bromine, and purple for iodine species.

AP AP AP AP AP AP AP AP AQ AQ AQ AQ

1. INTRODUCTION, BACKGROUND, AND MOTIVATION 1.1. Introduction

Turmoil in the Middle East after the Arab−Israeli war in 1973 caused a dramatic drop in the global oil supply and an energy crisis in many parts of the world. The realization that the earth’s limited petroleum resources might not be widely available led to an expansion of research in renewable energy sources. During this time, investigations into hydrohalic acid (HX) splitting indicated that halide redox chemistry could be utilized for energy conversion and storage. This promoted the development of photoelectrochemical cells (PEC)1−4 for HX splitting and regenerative cells5 that utilized halides as redox mediators for electrical power generation. Research was initially conducted using bulk semiconductor photoelectrodes,6,7 with over 1000 journal articles published in the field of HX splitting from 1975−1983.8 In one notable advance, Texas Instruments developed photoelectrochemical cells for solar HBr splitting that were integrated with H2 and Br2 storage capabilities, enabling electric power generation when the sun was down. 9,10 However, economic and environmental concerns precluded early prototype devices from being brought to market. Nevertheless, the halide photoredox chemistry advanced during this time has recently been revitalized as the need for the conversion and storage of solar energy has grown. This is the subject of this review. Along with a heightened interest in solar HX splitting, halides have been explored in a variety of renewable energy applications.11−13 For example, the most effective dyesensitized solar cells (DSSCs) utilize iodide ions to shuttle electrons between a dye-sensitized electrode and a counter electrode.14−17 Bromide mediators have also been studied but seem to suffer from corrosion issues like those encountered by Texas Instruments in their HBr splitting cells.10 The redox chemistry of halides has also been implicated in the remarkable efficiency of lead halide perovskite solar cells, which have garnered considerable attention.18−27 In these perovskite materials, the halide is intimately involved in light absorption and charge transport. Halides also passivate semiconductor surfaces, most notably in quantum dot solar cells.28−32 Finally, iodine has historically been utilized to increase the electrical conductivity of light-absorbing organic polymers and to

1.2. Background

1.2.1. Formal Reduction Potentials of Halogen Species. Formal reduction potentials among the halogens follow the periodic trends emphasized in introductory chemistry courses. A reminder of these trends and values is given in Figures 1 and 2 and Tables 1 and 2. Among the halide series, iodide has the lowest electron affinity and is hence the most easily oxidized in aqueous solutions with formal E°(X •/X −) reduction potentials following the trend I < Br ≪ Cl. These reduction potentials are not quantified through standard electrochemical measurements, as only two-electron transfer processes are observed at metal electrodes.39 This occurs because single-electron transfer to a halogen species is intimately coupled with bond-breaking or bond-formation chemistry that yields products that undergo a second electron transfer at the applied potential required to initiate the first single-electron transfer. For this reason, the one-electron reduction potentials have been estimated through free-energy cycles or extracted from kinetic data with application of Marcus theory. The details of how the kinetic approach is implemented are given in Section 2 with a description of the assumptions made and the possible sources of error. B

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Figure 1. Radii, ionic radii, and interatomic distances for the indicated halogen species.

Figure 2. (left) Latimer-type diagram of the formal reduction potentials and equilibrium constants for halogen species in aqueous solution.40,49,50 Color code is green (Cl species), red (Br species), and purple (I species). (right) Latimer-type diagram for iodine species in CH3CN.51

Table 1. Rate Constants for Common Halide Reactions in Aqueous Solutiona

a

Color code is green (Cl species), red (Br species), and purple (I species).

X• + X− F X 2•−

A Latimer-type diagram for Cl2, Br2, and I2 in aqueous solution is given in Figure 2.40 The E°(X•/X−) potentials indicate that strong photo-oxidants are required for the oxidation of each halide to the halogen atom. A thermodynamically less demanding pathway exists that requires direct oxidation of two halides to yield X2•−. For example, E°(I•/I−) = 1.33 V vs NHE, whereas E°(I2•−/2I−) = 1.03 V vs NHE. Because the equilibrium constant for eq 1 is large for iodide (Keq = 1.1 × 105 M−1), both redox pathways yield the same I2•− product in concentrated iodide solutions, yet the initial oxidation of iodide to yield a single atom requires an additional 300 meV of free energy.

(1)

Kinetic evidence of a concerted mechanism in which two iodides are oxidized by one electron to yield the I2•− product in a single step exists in the stopped-flow literature.41−43 This concerted pathway is proposed to have a cationic species in the transition state that helps overcome the Coulombic repulsion of the two iodides. Access to the concerted mechanism represents an obvious goal for solar energy conversion as weaker photo-oxidants that absorb a greater fraction of the solar spectrum could be employed. Note that the kinetic rate constants for formation of X2•− from the X• atom and X− anion are large for all these halides, C

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Table 2. Absorption Maxima, Molar Absorption Coefficient As Well As Ionization Potential (IP) or Electron Affinity (EA) for the Indicated Halogen Species in Water unless Otherwise Noted

a

CH3CN. bCrystal lattice. cCH2Cl2. dMeasured with a lanthanide trihalide salt.

The reduction of halides to X2− likely occurs outside the window available for common electrolytes and has not been reported to our knowledge. The one electron reduction potential of I3− has been reported in aqueous and CH3CN electrolytes. The reductions are dissociative, I3− + e− → I− + I2•−, and an [I32−] intermediate has been proposed but, to our knowledge, has not been directly observed.56,57 The corresponding oneelectron reduction potentials for Br3− or Cl3− were not found in the literature. The formal reduction potentials of the halogens in organic solutions are far less developed than in water, with only a few isolated values available in the literature. The most wellestablished values are for iodine species in CH3CN electrolytes that are commonly used in dye-sensitized solar cells. These values are given on the right-hand side of Figure 2. Note that in acetonitrile the iodine formal reduction potential is 1.23 V vs NHE, which is 100 mV less positive than the value in water. Theoretical calculations indicate that halide charge transfer to water is significant, which could account for the more favorable reduction potentials in CH3CN.58−60 However, it should be emphasized that halide solvation remains poorly understood and the answers to fundamental questions such as “how many solvent molecules are coordinated to a halide species?” or “how quickly do ligated solvent molecules exchange with the bulk solvent?” remain unknown. 1.2.2. Halide Spectroscopy. The absorption maximum and molar absorption coefficients of key halogen species measured in water are gathered in Table 2. Note that the absorption maximum of the halogen atoms, X•, increases in energy as one goes down the periodic table while the halides, X−, show the opposite trend. This occurs because the halide absorption is a halide-to-solvent charge-transfer transition while that for the atoms is a solvent-to-halogen-atom charge transfer.61−64 Hence iodide as the best electron donor absorbs

Table 1. Hence the use of photoredox chemistry to generate halogen atoms is expected to quantitatively yield X2•− in concentrated aqueous halide solutions. This reaction may compete kinetically with reductive elimination of X2 in transition metal-catalyzed HX splitting. In organic solvents or in the presence of other reactants, the halogen atoms may undergo competitive radical reactions, such as H atom abstraction, that are well documented, with I• being the most persistent radical and Cl• the least.44−46 This rapid reactivity has precluded direct determination of the X•/X− selfexchange rate constants where the barrier is expected to be solely determined by the outer-sphere reorganization energy. The reduction of the elemental halogens, X2, to yield X2•− follows period trends, with Cl2 reduction being the most facile. The electron is added to an antibonding orbital and lowers the bond order to 1/2, resulting in a longer σ2σ*1 bond in X2•−. It has long been known that I2 and Br2 can be oxidized in highly acidic oxidizing solutions to yield the one electron products, I2+ and Br2+.52−55 This redox chemistry involves removal of an electron from a πg* antibonding orbital, and hence the intranuclear distance decreases. Higher nuclearity cations, such as I3+ are also known, however, attempts to identify formal reduction potentials of any cationic halogen species in the literature have been unsuccessful.53−55 The Latimer diagrams show that the dihalide species, X2•−, where X = Br or I, are unstable with respect to disproportionation, eq 2. The free-energy change associated with this reaction is large and represents a significant loss mechanism in halide photoredox chemistry. Table 1 shows that this reaction is very rapid, occurring near the diffusion limit. A potential advantage of this chemistry is that the products formed are not easily reduced by a single electron, particularly in CH3CN electrolytes. X 2•− + X 2•− F X3− + X−

(2) D

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light at lower energy than does the other halides. Likewise, the iodine atom as the poorest electron acceptor absorbs light at the shortest wavelength and highest energy. A large body of literature has described the formation of halogen atoms, such as Cl•,64,65,80−120 Br • , 64,65,81,85,87,88,90,100−102,106,112,114,115,118,121−133 and I•,64,65,106,114,121,124,128,131,134−159 in water or organic solvents. In pulse radiolysis, Cl• is formed by reaction with SO4• or HO• that are generated by electron beam irradiation of Na2S2O8 or water, respectively.50,99,105 In laser flash photolysis, ultraviolet irradiation of halide solutions leads to the formation of X• and a solvated electron.65,160 Chlorine atoms were shown to form charge-transfer complexes with arenes in a variety of solvents and such complexes were also invoked to rationalize the increased selectivity of photochlorination reactions.91 A linear relationship between the absorption maxima of Cl• and the ionization potential of the solvent provided strong evidence for the solvent-to-atom assignment that was further interpreted with Mulliken’s intermolecular charge-transfer resonance theory, Figure 3 and Table 3.64,82,161 Similar charge-transfer complexes have been observed with Br• and acetonitrile (ε = 1470 M−1cm−1)122 as well as from I• in water, methanol, and 2propanol.148,151,162

The absorption spectra of the halides in water is dominated by two absorption bands in the UV region. These have been assigned as charge-transfer-to-water transitions that are sometimes referred to simply as outer-sphere charge-transfer transitions.163,164 The appearance of two transitions is understood by the doublet X • products, 2 P 1/2 and 2 P3/2.165−167 In keeping with periodic trends, one would anticipate that iodide would absorb light at the longest wavelength (lowest energy) and chloride at the shortest value, a trend that is approximately followed given the spread of literature values, Table 2. Direct confirmation of the chargetransfer assignment comes from studies where the solvated electron was directly observed after pulsed laser excitation into these charge-transfer absorption bands. Lever and co-workers’ analysis of these absorption bands provided reorganization energies of ∼1.4 eV for Cl− (1.42 eV), Br− (1.39 eV), and I− (1.38). It was noted that these values were preliminary and should be used with caution.164 The absorption spectrum of X2 is well-known in the gas and condensed phases. The spectra are understood with the molecular orbital diagram shown in Figure 4. The figure shows that halogen mixing of valence p orbitals forms bonding and antibonding orbitals. In the case of these diatomic halogen molecules, following the Laporte rule, there are only two allowed transitions, the σg → σu* and πg* → σu*, the latter of which gives rise to the purple color of iodine. This lower energy transition has displayed a pronounced solvent dependence. For example, the absorption maximum for I2 is 655 nm in the gas phase,168 523 nm in cyclohexane,70 and 460 nm in water.75 This solvatochromism is attributed to bond length changes that induce partial donation from the solvent to the σu*.68 An increased X−X bond length substantially decreases π interactions while only slightly affecting σ interactions, resulting in changes to the lower energy transition without major deviations to the higher energy transition. As discussed further below, I2 forms adducts with Lewis bases through halogen bonding. Seminal work by Grossweiner and Matheson have shown that laser flash photolysis of aqueous solutions containing iodide, bromide, or chloride led to the formation of the corresponding dihalide radical anions I2•−, Br2•−, and Cl2•−,65 and absorption features in the near IR were identified for I2•− and Br2•−.121,127,144 These absorption features were not thoroughly investigated as the UV flash photoionization generated a large amount of solvated electrons that absorbed light intensely in the near IR region and obscured the low energy absorption features of I2•− and Br2•−. Similar absorption features should be present for Cl2•−, and have been reported in irradiated potassium chloride at liquid nitrogen temperature.169 More refined analysis has shown that these species give rise to three absorption bands as is predicted from their orbital mixing shown in Figure 5. Because of the charged nature of the X3− anions, the symmetry of the ion is effectively broken in polar solvent systems such that the transition from πu to σu* is no longer Laporte forbidden and the orbitals no longer have ungerade symmetry.68 The absorption spectra of X3− is understood with the molecular orbital diagram shown in Figure 4. Mixing of p orbitals results in similar behavior to X2 with normally symmetry forbidden transitions and only two allowed transitions. In the case of triiodide, additional absorption bands have been observed and attributed to antisymmetric geometries, for which linear symmetry no longer exists and

Figure 3. UV−visible absorption spectra measured 25 ns after pulsed 355 nm laser excitation of Cl2 in the indicated solvents. The absorption bands were assigned to charge-transfer bands from the solvent to Cl•. Adapted with permission from ref 82. Copyright 1989 Elsevier.

Table 3. Chlorine Atom Charge-Transfer Absorption Maxima in Halogenated Solvents of Known Ionization Potentials82 solvent

Ip (eV)

λmax (nm)

hνCT (eV)

C7F14 CFCl2CF2Cl CFCl3 CHCl3 CCl4 CH2Cl2 CHCl2CHCl2 CHBrCl2

13.20 11.99 11.77 11.42 11.47 11.35 11.10 10.88

235 280 300 330 330 335 340 365

5.28 4.43 4.13 3.76 3.76 3.70 3.65 3.40

E

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Figure 4. Molecular orbital diagram for X2 (left) and X3− (right).

1.2.3. Halogen Bonding in Polyhalide Species. The Latimer-type diagram reveals favorable equilibrium constants for Lewis acid−base chemistry between halogen species that would not be predicted based on their Lewis structures alone, eq 3. This chemistry results from what is now referred to as “halogen bonding,” wherein a sigma hole (σ-hole) present on the pole of a terminal halogen atom in polyhalide, halogen, or organohalide species accepts charge density from a Lewis base.170−176 X− + X 2 F X3−

(3)

A celebrated early example of halogen bonding is found in the equilibrium between I 2 and Lewis bases in organic solvents177−179 as well as the interaction between X2 and 1,4-dioxane.180−182 Many other examples exist today.170−175 The magnitude of the sigma hole increases with the principal quantum number Cl ≪ Br < I, Figure 6. Hence the heavier halogens offer great versatility as they are able to act as either electron acceptors in halogen bonds or as electron donors in oxidative addition reactions. This expectation of greater Lewis acidity with increasing quantum number is born out of the acid−base equilibrium shown in eq 3 with equilibrium constants of 750 (I), 16 (Br), and 0.2 (Cl) M−1 in water.183,184 The less stable Cl3− species has proven the most elusive out of the three, and the I3− equilibrium in water is the most studied and well established.39,185 In light of the limited availability of data, the equilibria have been analyzed with thermochemical cycles, and therefore the relative bond strengths of the X3− species were considered. In theoretical studies, the calculated bond energies of the X3− ions were similar and spanned only 7 kJ mol−1 between Cl3− and I3−.186 Using thermochemical cycles, Crawford et al. investigated the influence of solvation enthalpies as the dominant term in determining the equilibrium constants.187 There were clear differences in solvation enthalpies for X− (g) and X3− (g), but they were found to be small in comparison with the enthalpy of solvation of the elements, X2. As a result, the authors concluded that the trend in equilibrium constants was ultimately determined by the halogen X2 solvation enthalpy from their standard state,

Figure 5. Electronic transitions in dichloride, dibromide, diiodide, and tetraiodide.65−67,77,137,148

weak transitions in the UV−vis region can occur.68 Ultrafast measurements after pulsed laser excitation of I3− dissolved in ethanol revealed that the formation of I2•− occurred within 300 fs, consistent with a dissociative I3−* excited state.138 Similar dissociative excited states were invoked in the photochemistry of Cl3− and Br3−.64 The formation of I• and I2•− from I3−* appears to be very general and has been reported to occur in a plethora of solvents such as water, acetonitrile, propionitrile, butyronitrile, dimethoxyethane, methanol, ethanol, 2-propanol, and acetone.65,134,135,144,148,151 It is curious that this photochemistry, as well as that of the higher nuclearity polyiodides discussed below, does not yield stable diamagnetic products such as I2 and I− from I3−. F

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Figure 6. Electrostatic potential surfaces mapped onto total electron density for the X2 species. The surfaces were calculated using M06-L functional with a LANL2DZdp ECP basis set in the Gaussian09 package.

iodide ions. Water, for example, is not a suitable solvent, as the equilibrium constant associated with I− + I2 ⇌ I3− is only 750 M−1 at 25 °C. Organic and less protic or completely aprotic solvents have much larger equilibrium constants. In an early example by Fornier de Violet et al., flash photolysis was used to produce I2•− that eventually lead to the formation of I4•− in tBuOH.135 Two key absorption features appeared on the microsecond time scale with different rate constants. An absorption band centered at 590 nm (ε = 15000 ± 5000 M−1 cm−1) was assigned to I4•− and a band at 740 nm was assigned to I2•−. The proposed reaction mechanism indicated that I4•− was formed through the reactions of I• and I3−, eq 4.

which were 23.3 kJ/mol, 32.2 and 80.2 kJ/mol for Cl2(g), Br2(l), and I2(s), respectively.187 Hence, the unfavorable solvation of X2 by water was deemed responsible for the periodic trend in Keq for eq 3, i.e., I2 > Br2 > Cl2. In addition, a more electropositive sigma hole should facilitate a stronger interaction between the halide and halogen, resulting in a higher Keq,174 Figure 6. This follows the periodic trend in equilibrium constants with increasing principal quantum number in water. The Keq values for iodine in eq 3 are on the order of 102 M−1 in water and 107 M−1 in CH3CN. The 5 orders of magnitude increase in an organic solvent is difficult to rationalize or to predict theoretically. Furthermore, the periodic trend in Keq values have been reported to reverse in aprotic solvents such as acetonitrile, nitromethane, and acetone, i.e., Cl2 > Br2 > I2.188,189 Attempts to relate the reversed order of equilibrium constants in acetonitrile versus water to the free energy of transfer of X− from an aqueous to a nonaqueous solvent have been made. Unfortunately, the limited availability of solvation enthalpies of these halogen species in organic solvents make development of a robust predictive model difficult. The chlorine equilibrium in organic solvents has been particularly difficult to measure, and the value Keq = 1010 M−1 in acetonitrile is based on electrochemical measurements with platinum microelectrodes.190 However, other studies using platinum disk electrodes have shown that the trend for the halides was the same in water as in both acetonitrile and nitromethane.189,191−194 Such considerations underscore the importance of experimental determination of reduction potentials and equilibrium constants in nonaqueous conditions that could one day enable predictive models to be developed. However, the preponderance of data shows that the forward reaction in eq 3 is favored in nonaqueous solvents. 1.2.3.1. The Special Case of Polyiodides. As the most polarizable halogen with the greatest tendency to act as both a Lewis acid and a Lewis base, iodine displays diverse reactivity with high nuclearity species that warrants its own discussion.195 The formation of polyhalides has been reported both in fluid solution134,135 as well as in solid matrices.137,156,157 To prepare high nuclearity polyiodide species, conditions must be found where triiodide ions are in large excess with respect to

I• + I3− F I4•−

(4)

I5− F I4•− + I•

(5)

•

I +

I5−

F I6

•− •

(6)

I3−

t

The reaction of I and would be favored in BuOH because nonaqueous solvents increase the equilibrium concentration of polyiodide species. A similar approach was used for the formation of I6•−. The equilibrium I3− + I2 ⇌ I5− was achieved in tBuOH or acetonitrile solution at I2 concentrations greater than 10−2 M. The pentameric species, I5−, exhibited absorption features at 315 and 395 nm. Pulsed-light excitation of I5− lead to the formation of I4•− and I•, and the latter reacted with I5− to form I6•−. The absorption maximum of I6•− was centered at 540 nm. Both the absorption maximum of I4•− and I6•− were the same in CH3CN and tBuOH solvents, indicating that these transitions were not outer-sphere charge-transfer bands as was reported for I•.148 The tetranuclear I4•− anion was also synthesized by γirradiation of molecular iodine in frozen ethereal solvent at 77K.137 In this case, the proposed mechanism for I4•− formation involved the reaction of I2•− and I2. In a more recent study, Miller et al. characterized the dissociative photochemistry of I3− and have shown I4•− formation using a single crystal of tetra-n-butylammonium triiodide (TBAI3) at 80K, Figure 7.156,157 Femtosecond irradiation of I3− allowed population of an excited state that subsequently underwent G

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ultrafast, radiationless decay, creating two ground-state photoproducts: I2•− and I•. These two products were formed in a confined space that was vicinal to I3− molecules. Within a few picoseconds, the secondary reaction between I• and I3− was quantified, giving rise to spectral features associated with I4•−. Other polyiodide species ranging from I7− to I264− have been reported in the literature.195 While the specific synthetic methods vary from case to case, the general synthetic procedure of iodine addition to iodide or an already existing polyiodide, such as I5−, still holds, eqs 4−6. To stabilize higher nuclearity polyiodides species, the polyiodides are often isolated in the solid state with guest cations. Such examples include (Et4N)I7,196 [Fe(phen)3]I18,197 (MePh3P)I22,198 and (Me10Fc)4I26.199 Because of their solid-state nature, structural complexity, and lack of photoredox literature, this review will not focus on any of these larger polyiodide systems. 1.2.3.2. Mixed Halogen Species. There are numerous mixed halogen species commercially available. Iodine monochloride (ICl) and iodine monobromide (IBr) in particular are available and are used most commonly as sources of electrophilic I+ for organic synthesis.200,201 Iodine trichloride (ICl3) is also commercially available and is used as an analytical reagent for organic synthesis202 and as a topical antiseptic.203 Bromine monochloride (BrCl) is also commercially available. The strong oxidizing nature of BrCl allows for its use in analytical chemistry to determine low concentrations of mercury.204,205 In their early report, Grossweiner and Matheson attempted to form mixed halide radical anions such as IBr•− by flash photolysis of aqueous solutions containing alkali iodide and bromide.65 In all cases, the sole observable product was I2•−, even when the molar ratio of Br− to I− was 500:1. It was concluded that either I2•− was considerably more stable than IBr•− or that their absorption features coincided. Mixed-halide radical species have nonetheless been observed in rigid matrices. In a series of papers, Ershow et al. investigated the formation of mixed halide radical anions (ClBr•−, BrI•−) in

Figure 7. (top) Possible reaction mechanism of I3− in solid-state tetra-n-butylammonium triiodide (TBAI3). (a,b) The photoexcitation of I3− (a) initiates a reaction chain in both liquid phase (top) and within a crystal (bottom) that leads to its dissociation to I2− + I• as shown in (b). (c) In liquid, the solvent cage protects the nascent photoproducts and they recombine to reform I3−. In the crystal, a secondary reaction occurs and I• reacts with an adjacent I3− ion to give I4•−. Adapted with permission from refs 156 and 157. Copyright 2017 Springer Nature.

Table 4. Rate and Equilibrium Constants for the Chlorine Species reaction HO• + Cl− ⇌ HOCl•−

equilibrium

rate constant forward

rate constant reverse

solvent

ref

0.70 ± 0.13 M−1

4.3 ± 0.4 × 109 M−1 s−1

6.1 ± 0.8 × 109 s−1

H2O

104

H2O

50

SO4•− + Cl− ⇌ SO42− + Cl• OH•− + H+ ⇌ Cl• + H2O

Cl• + Cl− ⇌ Cl2•−

2.7 × 108 M−1 s−1 1.6 × 107 0.9−4.4 × 107

2.1 ± 0.7 × 1010 M−1 s−1

1.3 × 103 M−1 s−1 0.3−3.0 × 103 M−1 s−1

H2O H2O

104 113

1.9 × 105 M−1 1.4−2.8 × 105 M−1 1.4 ± 0.1 × 105 M−1

8.5 × 109 M−1 s−1 2.1 × 1010 M−1 s−1 8.5 ± 0.7 × 109 M−1 s−1 8.5 ± 0.7 × 109 M−1 s−1 8 × 109 M−1 s−1

1.1 ± 0.5 × 105 s−1 1.1 ± 0.4 × 105 s−1 6.0 ± 0.5 × 104 M−1 s−1

H2O H2O H2O H2O H2O H2O H2O

104,113 104 50 50 106 50 104,106

H2O H2O H2O

99 209

H2O

111

H2O

206

4.2 × 10 s 4

Cl2•− + Cl2•− ⇌ 2Cl− + Cl2

Cl• + Cl− ⇌ Cl2•− Cl2•− + Cl2•− ⇌ Cl3− + Cl−

−1

3.5 × 109 M−1 s−1 1.4 ± 0.3 × 1010 M−1 s−1 1.38 ± 0.13 × 1010 M−1 s−1 1.4 ± 0.2 × 105 M−1 2.6 × 109 M−1 s−1 H

6.0 ± 0.5 × 104 s−1

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Table 5. Rate and Equilibrium Constants for the Bromine Species reaction HO• + Br− ⇌ BrOH•−

Br• + HO− ⇌ BrOH•−

equilibrium

rate constant forward

321 M−1 321 M−1 333 M−1

1.06 ± 0.08 × 1010 M−1 s−1 1.06 × 1010 M−1 s−1 1.06 ± 0.08 × 1010 M−1 s−1 1.1 × 1010 M−1 s−1

3.1 × 103 M−1

1.3 × 1010 M−1 s−1

Br• + Br− ⇌ Br2•−

rate constant reverse

solvent

ref

3.3 × 107 s−1 3.3 ± 0.4 × 107 s−1 3.3 × 107 s−1

H2O H2O H2O H2O

127 85 127 132

4.2 × 106 s−1 4.2 ± 0.6 × 106 s−1

H2O H2O

132 127

H2O H2O H2O H2O H2O acetone acetone CH2Cl2

127 123 85 132 64 210 211 212

1010 M−1 s−1 3.3 2.2 6.3 1.6

× × × ×

103 105 105 104

M−1 M−1 M−1 M−1

1.0 × 1010 M−1 s−1 1.2 × 1010 M−1 s−1 1.1 ± 0.1 × 1010 M−1 s−1 1.1 ± 0.2 × 1010 M−1 s−1 1.1 ± 0.4 × 1010 M−1 s−1 5.4 ± 1 × 108 M−1 s−1

4.6 × 104 s−1 1.9 × 104 M−1 s−1 7 ± 2 × 105 s−1

Br2•− + Br• ⇌ Br3−

1.37 × 109 M−1 s−1

H2O

123

Br• + Br• ⇌ Br2

1.2 × 108 M−1 s−1

H2O

123

H2O H2O H2O

85 132 213−215

Br2 + Br− ⇌ Br3−

17.4 M−1 16.1 M−1 16.0 M−1

9.6 × 108 M−1 s−1

5.5 × 107 s−1

Br2•− + Br2•− ⇌ Br3− + Br−

1.6 × 109 M−1 s−1 1.9 × 109 M−1 s−1 3.4 × 109 M−1 s−1

H2O H2O H2O

123 132 206

Br2•− + Br2•− ⇌ Br2 + 2 Br−

3.0 × 109 M−1 s−1 2.2 × 109 M−1 s−1

H2O H2O

85 216

Table 6. Rate and Equilibrium Constants for the Iodine Species reaction

equilibrium

rate constant forward

1.1 ± 0.2 × 104 M−1

8.8 × 109 s−1 2.3 × 1010 M−1 s−1 1.3 × 1010 M−1 s−1 2.1 × 1010 M−1 s−1 1.4 × 1010 M−1 s−1 1.2 × 1010 M−1 s−1 9.8 ± 0.3 × 109 M−1 s−1 2.4 ± 0.2 × 1010 M−1 s−1 3.1 ± 0.3 × 1010 M−1 s−1 1.7 ± 0.1 × 1010 M−1 s−1 1.7 ± 0.1 × 1010 M−1 s−1 6.2 ± 0.1 × 109 M−1 s−1 1.5 ± 0.1 × 1010 M−1 s−1

I• + I− ⇌ I2•−

I2 + I− ⇌ I3−

107 M−1 720 M−1

rate constant reverse

solvent

ref

9 × 105 s−1

H2O CH3CN H2O C2H5CN CH3(CH2)2CN CH3O(CH2)2CN H2O CH3CN acetone acetone CH2Cl2 CH2Cl2 CH2Cl2

149 217 217 217 217 217 148 218,219 220 220 221 221 221

CH3CN H2O

222 148

I3− + e− → I2•− + I−

3.5 × 1010 M−1 s−1

H2O

223

I2 + e− → I2•−

5.2 × 1010 M−1 s−1

H2O

223

7.7 ± 1.5 × 109 M−1 s−1 4.5 × 109 M−1 s−1 1.5 × 109 M−1 s−1 3 × 109 M−1 s−1

H2O H2O H2O CH3CN

148 140 65,148 57

I2•− + I2•− ⇌ I3− + I− 12000

I

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Table 7. Rate and Equilibrium Constants for the Mixed-Halide Species reaction

equilibrium

rate constant forward

rate constant reverse

solvent

ref

Cl2•− + Br− ⇌ ClBr•− + Cl− Br• + Cl− ⇌ ClBr•− ClBr•− + Br− ⇌ Br2•− + Cl− Cl2•− + I− ⇌ 2Cl− + I• Br2•− + I− ⇌ BrI•− + Br− Br− + I• ⇌ BrI•− BrI•− + I− ⇌ I2•− + Br− BrI•− + BrI•− ⇌ BrI2•− + I−

3.6 × 106 1.3 × 102 1.9 × 103

4 × 109 M−1 s−1 1.1 × 1010 M−1 s−1 8.0 × 109 M−1 s−1 4.5 × 109 M−1 s−1 4.3 × 109 M−1 s−1 1.0 × 1010 M−1 s−1 5.8 × 109 M−1 s−1 3.0 × 109 M−1 s−1

1.1 × 102 M−1 s−1 8.5 × 107 s−1 4.3 × 106 M−1 s−1

H2O H2O H2O H2O H2O H2O H2O H2O

206 206 206 140 124 124 124 124

17.5 1.34 × 103

aqueous solutions.101,102,124,140,206 Pulse radiolysis of aqueous solutions of NaCl and NaBr yielded, depending on the ratio of halide ions present, Cl2•−, ClBr•−, and Br2•−. Furthermore, the formation of Cl3−, Cl2Br−, and ClBr2− ions were also quantified as reaction products, confirming therefore that mixed halide intermediates could be formed. The rate and equilibrium constants for the halide and mixed-halide species in aqueous and nonaqueous solvent are gathered in Tables 4−7. The maximum absorption of ClBr•− was centered at 350 nm with a molar absorption coefficient of 9300 M−1 cm−1. The formation of ClBr•− is proposed to occur through the reaction between Cl • and Br − , in a similar fashion as ClSCN •− and BrSCN•−.207,208 The formal reduction potential E°(ClBr•−/ Cl•,Br−) = 1.85 V was also determined. Attempts to prepare ICl•− have thus far been unsuccessful. It was shown that Cl2•− oxidized iodide to form I•, which subsequently reacted with I− to form I2•−. It was hence postulated that the inability to form ICl•− was a direct consequence of the reduction potentials: E°(Cl2•−/2Cl−) = 2.09 V, E°(Br2•−/2Br−) = 1.62 V, E°(Br•/Br−) = 1.93 V, and E°(I•/I−) = 1.33 V.49 In the case of ClBr•−, the E°(Cl2•−/ 2Cl−) and E°(Br•/Br−) are separated by 0.16 V, which would favor bond formation and generation of a mixed orbital in ClBr•−, whereas the large difference (0.76 V) between E°(Cl2•−/2Cl−) and E°(I•/I−) would only favor electron transfer. Support of this hypothesis was garnered by the fact that the oxidation of I− by Br2•− yielded both BrI•− and I2•−. In this case, the difference in formal reduction potentials between E°(Br2•−/2Br−) and E°(I•/I−) was 0.29 that could therefore enable formation of the mixed BrI•− species. The absorption maximum of BrI•− was centered at 370 nm with a molar absorption coefficient of 9650 M−1 cm−1 and a formal reduction potential E°(BrI•−/Br−,I−) = 1.25 V was also determined.124

Br > Cl > F. The sigma hole was present on the halogen carbon bond and could therefore accommodate the expected 180° R−X···I− bond angle. Kinetic studies found indeed that iodide oxidation was most rapid for the iodine containing sensitizer, Dye−I, and followed the trend, I > Br > Cl > F. This behavior could not be rationalized through thermodynamic considerations and was instead attributed to halogen bonding. For example, a distribution analysis revealed a second-order regeneration constant of 4.7 × 108 M−1 s−1 for Dye−F and 13.5 × 108 M−1 s−1 for Dye−I. The larger rate constant for Dye−I was also manifest in a larger Voc value and a higher light-to-electrical energy conversion efficiency. M

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unreasonable second-order rate constants, indicating that I3− was indeed the acceptor. Through systematic studies of the initial ratio of I− and I2 in a DSSC electrolyte, O’Regan and co-workers have provided compelling evidence that I2 is an important electron acceptor.282,283 A standard electrolyte is prepared with 0.5 M LiI and 0.05 M I2 in CH3CN, often with additional additives. The equilibrium concentration of I2 is expected to be micromolar, and increasing this value led to a lower DSSC efficiency. This low I2 concentration explains why recombination with TiO2(e−) is a rare event. There is some evidence that the unwanted recombination reaction occurs by an inner-sphere mechanism wherein I2 (or I3−) forms an adduct with the ground-state dye.282,286 These putative adducts have not been resolved spectroscopically, but kinetic evidence for their presence exists and is in line with their ability to form halogen bonds. O’Regan and co-workers found that sensitizers with N or S atoms can promote recombination as evidenced by a smaller Voc and short lifetime of the injected electron relative to sensitizers that do not have these atoms.287 Later studies showed that the location of the heteroatom was also important and that sensitizers with S atoms, for example, could be highly efficient in DSSCs.288 1.3.2.3. Aqueous DSSCs. The most efficient DSSCs have been realized in CH3CN electrolytes. Trace water was initially viewed as a detriment and was attributed to poor stability, sensitizer desorption, and poor reproducibility. However, more recent experiments have led to highly stable aqueous DSSCs with efficiencies as high as 5.64%.289−291 This efficiency is about 1/2 the value certified in CH3CN, and the Latimer diagram of Figure 2 provides a possible explanation for this. The 4 orders of magnitude lower Keq value for I− + I2 ⇌ I3− results in higher equilibrium concentrations of I2, which O’Regan has shown to be an important acceptor for injected electrons.282,283 In addition, the E°(I3−/I2•−,I−) reduction potential has an unusual solvent dependency and shifts positive by ∼300 mV in water, conditions that favor reduction by TiO2(e−)s. While the reduction potential of the TiO2(e−)s is also certainly solvent dependent and high mobility solvents help ensure rapid mass transport of iodide, those solvents that have large Keq values and very negative one-electron reduction potentials for I3− are expected to be good candidates for high efficiency DSSCs.

2.1. Excited-States and Quenching Mechanisms

2.1.1. Metal-to-Ligand Charge-Transfer Excited States. Visible light absorption by the parent complex [Ru(bpy)3]2+, where bpy is 2,2′-bipyridine, is a metal-to-ligand charge-transfer transition that formally promotes an electron from the metal to a bipyridine ligand, as shown in eq 10.294−301 hυ

[Ru II(bpy)3 ]2 + → [Ru III(bpy •−)(bpy)2 ]2 +*

(10)

After light absorption, rapid relaxation to a thermally equilibrated MLCT excited state is accompanied by significant charge transfer from the ligand back to the metal.302 The Lattimer-type diagram shows how the [Ru(bpy)3]2+ excitedstate is both a stronger oxidant and reductant than the [RuII(bpy)3]2+ ground state, Figure 11.

Figure 11. Formal reduction potential for [Ru(bpy)3]2+* in CH3CN at a fixed ionic strength of 0.1. Potentials are referenced vs SCE (+0.244 vs NHE).299

In the absence of an external quencher, the MLCT excited state lifetime relaxes to the ground state by three processes: (i) radiative decay through photoluminescence, (ii) nonradiative decay, and (iii) thermal population of ligand-field (LF) excited state(s). The latter process often leads to irreversible ligandloss chemistry.303,304 Equation 11 is used to characterize the excited-state lifetime, where ΔE represents the activation energy between the MLCT state and the LF state while kr and knr are the radiative and nonradiative rate constants. Decay through the LF states is often included in knr, such that 1 τ = (k + k ) and the quantum yield for photoluminescence, r

ϕ=

nr

kr . (k r + k nr)

Therefore, measurement of τ and ϕ enables

experimentalists to determine kr and knr. 1 τ= −ΔE / RT (k r + k nr + A e )

2. HALIDE PHOTOREDOX CHEMISTRY WITH METAL-TO-LIGAND CHARGE-TRANSFER (MLCT) EXCITED STATES Halides absorb light most effectively in the ultraviolet region. A strategy to achieve halide oxidation as described here is to use dyes to sensitize halide redox chemistry to visible light. As dyes in electronic excited states are stronger oxidants and reductants than the ground state, dye sensitization can initiate both reduction and oxidation of halogen species. Indeed, the oxidation of monohalide anions, X−, and the reduction of I3− have been realized with a particular class of transition metal “dyes” with a (dπ)6 ground-state configuration and low-lying metal-to-ligand charge-transfer (MLCT) excited states.57,211,219,292,293 In this section, we briefly discuss MLCT excited states and the common mechanisms for quenching these excited states. A brief description of Marcus theory and its employment in the determination of formal reduction potentials is also presented, followed by a review of halide redox chemistry initiated by visible light absorption.

(11)

In fluid solutions at room temperature, the MLCT excited state is luminescent with about a one microsecond lifetime that enables diffusional interactions with halide species.299,305,306 The main contributor to the excited-state lifetime is the nonradiative rate constant which is approximately 106 s−1, whereas the radiative rate constant is much smaller, 104 s−1. The solvent has been shown to influence these rate constants.307 Time-resolved resonance Raman and Stark spectroscopy have revealed that the excited electron is localized on a single ligand. For heteroleptic complexes this corresponds to the bipyridine ligand that is most easily reduced.308−316 The high stability of the Ru(II) complexes in the adjacent one-electron reduced and oxidized forms coupled with the ability to finetune formal reduction potentials at the molecular level make them ideal for fundamental study of halide photoredox chemistry. The ruthenium(II) polypyridyl and related N

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Figure 12. Ru(II), Cr(III), and Ir(III) complexes discussed in this section. Note that the respective charges are indicated with each complex’s name.

quenching rate constant (kq) through the excited-state lifetime in the absence of quencher, τ0, eq 12. Note that large kq values do not necessarily indicate an underlying electron-transfer reaction. Alternative quenching reactions, such as energy transfer or, for organic dyes, enhanced intersystem crossing to nonemissive triplet states may be operative.306 Complementary transient absorption spectroscopic measurements that provide direct identification of the reaction products with rate constants that can be compared to those quantified through quenching studies are hence of great value.

complexes that will be discussed in this section are gathered in Figure 12. We note that excellent reviews of the details of MLCT excited states are available in the literature and will not be detailed further here.294−301 Halide photoredox chemistry promoted by MLCT excited states has been shown to occur by both dynamic (diffusional) and static mechanisms. 218,219,273,317−320 It is therefore worthwhile to briefly describe these mechanisms and review diffusion as it relates to excited-state electron transfer and the subsequent reactivity of the electron-transfer products. The remainder of this section focuses on recent experimental results with specific complexes. 2.1.2. Stern−Volmer Analysis. In 1919, Otto Stern and Max Volmer provided a robust analysis that has withstood the test of time and enables one to quantify quenching of luminescent excited states by redox active species in fluid solution.321 As the name implies, “quenchers” are species that quench the excited state as measured by a decrease in the steady-state photoluminescence intensity, PLI0, and/or the lifetime, τ0. Stern−Volmer analysis requires that conditions be identified where the excited-state concentration is fixed and the quencher, Q, concentration is systematically varied. Plots of PLI0/PLI and τ0/τ versus the quencher concentration provide mechanistic insights. When such data is linear and coincident, the slope provides the Stern−Volmer constant, KSV, and a dynamic quenching mechanism that involves diffusion is operative. Such dynamic quenching is also equivalently referred to as “diffusional” or “collisional” quenching. The magnitude of (KSV)−1 provides the quencher concentration necessary to quench 1/2 of the excited states. In a dynamic reaction, KSV is also directly related to the bimolecular

τ (PLI 0) = 0 = 1 + KSV[Q ] = 1 + kqτ0[Q ] τ (PLI)

(12)

In an alternative static mechanism, a quencher, Q, decreases the PL intensity without influencing the excited-state lifetime τ0. Under these conditions, the quencher forms a nonemissive ground-state adduct. Diffusion of the quencher or the excited state is not operative in a static electron-transfer mechanism. Static quenching appears as a loss in the initial amplitude, α0, of time-resolved PL decays. Plots of α0/α and PLI0/PLI versus the quencher concentration provide the equilibrium constant for the quencher−chromophore adduct, KS, eq 13. The formation of adducts often results in a significant change in the ground-state absorption, so care must be taken to identify excitation wavelengths where the fraction of light absorbed by the chromophore is constant as demanded by the Stern− Volmer analysis. α (PLI 0) = 0 = 1 + KS[Q ] α (PLI) O

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Figure 13. Excited-state quenching of [Ru(bpy)2(deeb)]2+* by iodide in CH3CN (a,b,c) and in CH2Cl2 (d,e,f). Steady-state (a,d) and timeresolved (b,e) photoluminescence quenching with increasing concentration of iodide. The decrease of signal amplitude in (e) is associated with static quenching, whereas a decrease in excited-state lifetime is correlated to dynamic quenching. (c,f) Stern−Volmer plot using either the integrated photoluminescence (PLI), the excited-state lifetime (τ), or the initial amplitude of the time-resolved photoluminescence (α). The equilibrium constant is extracted from the amplitudes (α0/α), whereas the quenching rate constants, kq, are abstracted from excited-state lifetime (τ0/τ).

Figure 14. Proposed mechanism for dynamic excited-state quenching by iodide.

Such steady-state PLI quenching data can be quantitatively modeled when KD and KS are known. This combined Stern− Volmer analysis displays a quadratic dependence on the quencher concentration which underlies the upward curvature, eq 14.

Note that eq 13 assumes the ground-state adduct is nonluminescent, yet cases have been identified where the electron-transfer rate constant for halide oxidation within the adduct is slow and competitive, with excited-state relaxation leading to incomplete quenching of the excited state.211,212,292 This can lead to very different behavior in Stern−Volmer plots as is described below for “ter-ionic” complexes.220 Under many experimental conditions with ionic excitedstates and halogen species, both static and dynamic quenching mechanisms are operative. This is often manifest as an upward curvature in a Stern−Volmer plot of PLI0/PLI versus [Q].

(PLI 0) α τ = 0 × 0 = 1 + (KD + KS)[Q ] + KDKS[Q ]2 (PLI) α τ (14)

Iodide quenching of [Ru(bpy)2(deeb)] *, where deeb is 4,4′-(CO2CH2CH3)2-2,2′-bipyridine, provides an illustrative 2+

P

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Figure 15. Gibbs free energy surfaces that represent the electron donor and acceptor reactants (blue) and products (red) as the electronic coupling matrix element (HDA) is increased from 0 (nonadiabatic) to over 1000 cm−1 (adiabatic). The decrease in the Gibbs free energy expected for an adiabatic electron-transfer reaction is emphasized, |ΔG°ad| < |ΔG°|.

referred to as the activated rate constant (M−1 s−1) as described further by Sutin.323

example of Stern−Volmer analysis. 317−319 In CH3 CN solutions, the quenching is purely dynamic, as demonstrated by the Stern−Volmer plot in Figure 13, where the excited-state lifetime decreased with added iodide and the initial amplitude remained unperturbed, Figure 13b. Hence, plots of τ0/τ and PLI0/PLI were coincident. However, in dichloromethane solution, not only did the lifetime decrease, but the initial amplitude also decreased as a function of iodide concentration, Figure 13e. The Stern− Volmer plot (PLI0/PLI) showed the characteristic upward curvature, consistent with a combination of both dynamic and static quenching mechanisms. It was shown through Benesi− Hildebrand and Stern−Volmer analysis that iodide ion pairs were formed with a large equilibrium constant, Keq = 59700 M−1 in dichloromethane solution.317,319 This iodide ion pairing process was reversible and could be tuned when an inert salt such as tetrabutylammonium hexafluorophosphate (TBAPF6) was used to control the solution’s ionic strength. 2.1.3. Kinetic Analysis. It is of interest to relate the quenching rate constant kq extracted from a Stern−Volmer plot to the true electron-transfer rate constant. Complementary spectroscopic data that identifies and kinetically resolves the appearance of the electron-transfer products with a rate constant equal to kq constitutes direct evidence of an excited-state electron-transfer quenching process. A composite mechanism similar to all bimolecular electron-transfer reactions in fluid solution is shown in Figure 14. Light absorption creates an MLCT excited state that undergoes diffusion and forms an adduct with iodide with a rate constant kdiff. This adduct is often referred to as an “encounter complex” that is shown in the brackets of Figure 14. Electron transfer within the encounter complex yields an iodine atom and the reduced Ru complex. These electrontransfer products may undergo back-electron transfer to generate ground-state products or may diffuse apart with a “cage escape” yield, ϕce, greater than zero.322 The electrontransfer products that escape the cage eventually recombine to form the ground-state reactants with a rate constant, kbet. A steady-state approximation applied to the encounter complex provides eq 15 that relates kq to kdiff and the product of the equilibrium constant for encounter complex formation, KA = kdiff/k−diff, and ket. Within this encounter complex, electron transfer occurs by a first-order reaction with units of s−1. As the units of KA are M−1, the product KA and ket are

1 1 of an excited state diffuses with the lifetime1 = + kq kdiff KAket (15)

Diffusion in fluid solution is well understood. Diffusion limited rate constants kdiff = 106−1011 M−1 s−1 are typical for two neutral molecules in organic solutions at room temperature.324,325 The kdiff values have been determined with eq 16, where NA is Avogadro’s number, and D is the diffusion coefficient for each reactant.325 Values of D have been measured by electrochemical methods for redox active species, while diffusion of the excited states is discussed further below. Alternatively, D may be calculated through the Stokes− Einstein relation, where η is the solvent viscosity and r is the radius of the assumed spherical reactant, eq 17. The parameter β is the effective reaction radius defined by eq 18, where R is the sum of the reactant radii, r, Rc is the Onsager radius, eq 19, κ is the Debye length, eq 20, z is the ionic charge, and I is the ionic strength of the solution. 324

kdiff = 4πNA(DRu + DX )β D=

β=

(16)

kBT 6πηr

(17)

R ce R cκ [e(R c / R) − 1]

(18)

z Ruz Xe 2 4πεrε0kBT

(19)

Rc =

κ −1 =

εrε0kBT 2000e 2NAI

(20)

The association equilibrium constant has been estimated with eq 21.326 i4y KA = 1000jjj zzzπR3NAe(−R c / R)e R cκ /(1 + κR) k3{

(21)

These equations yield theoretical estimates of the diffusion rate constant, kdiff and the association constant, KA, which together Q

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[CrII(H2O)6]2+ quantitatively yielded [CoII(NH3)5(H2O)]2+ and [CrIIICl(H2O)5]2+ through a proposed inner-sphere mechanism involving an elusive chloro-bridge [Co - - Cl - Cr]4+ intermediate. Related halogen bridged intermediates may be present in the bimolecular electron transfer reactions described in section 2 and are of clear relevance to the halideto-metal charge-transfer excited states discussed in section 3. The magnitude of the electronic coupling is expected to impact ΔG°, and hence formal reduction potentials for halogen species extracted from kinetic measurements.334 Strong electronic coupling lowers the Gibbs free energy change for electron transfer, i.e. |ΔG°ad| < |ΔG°|. Hence, if electronic coupling is sufficiently strong in the encounter complex, the free energy change may be significantly different than that expected. This in turn would lead to incorrect estimations of formal reduction potentials of important halogen species. In principle, such errors could be avoided by making certain that the halogen reduction potentials measured are insensitive to the nature of the excited state utilized, as is discussed further below. In practice, experimental details such as solubility, light absorption, and the need for long-lived excited states that undergo diffusional reactions often preclude robust analysis with different photosensitizers. The reorganization energy, λ, is defined as the free energy necessary for electron transfer from the donor to the acceptor without nuclear motion when ΔG° = 0.339,340 The total reorganization energy is usually separated into a sum of innersphere (bond lengths and angles) and outer-sphere (solvent and electrolyte) contributions.332,341 For the oxidation of halide anions to the halogen atom, there are no contributions from the inner-sphere and the total reorganization energy change is determined by the solvent or electrolyte. A small inner-sphere contribution to the total reorganization energy is expected for MLCT excited states.342 The outer-sphere reorganization energy in organic solvents can be crudely estimated with dielectric continuum theory and values of 1−1.5 eV are typical in organic solvents. The reduction of polynuclear halides may involve a significant inner-sphere reorganization energy, particularly when a dissociative reaction is operative. For example, if I3− reduction results in dissociation and directly yields I2•− and I•, and hence does not involve an I32− intermediate, then the I−I bond dissociation energy is included within the total reorganization energy.343 The Gibbs free energy change for excited-state electron transfer, ΔG°, was considered by Rehm and Weller.344,345 IUPAC defines this quantity for a mole of generic donors (D) and acceptors (A) that can be photoexcited to yield D+ and A−, eq 24.345

with kq can be used to estimate the electron-transfer rate constant for X− oxidation, ket. The descriptions of dynamic or collisional quenching provide context to the role excited-state lifetimes play in the quenching efficiency. As metastable species, excited states do not diffuse for indefinite periods of time as they eventually decay by radiative or nonradiative relaxation to the ground state. The Einstein equation relates the average distance (Δx2)1/2 an excited state diffuses with the lifetime τ, eq 22. Δx 2 = 2Dτ

(22)

Suppose the excited state has a diffusion coefficient of 2.5 × 10−5 cm2/s. If the excited state has a lifetime of 10 ns, then the average distance it can diffuse (Δx2)1/2 is about 70 Å. However, if the lifetime were 1 μs, then the excited state could diffuse for 700 Å. This underscores the importance of longlived excited states when diffusional quenching is operative. Indeed, for practical applications of luminescent excited states such as in O2 sensors, excited states with lifetimes less than 50 ns are not influenced by O2 dissolved in water. The excitedstate lifetime is simply not long enough to encounter O2 molecules through diffusion. However, long-lived luminescent excited states are very sensitive to O2 and have been successfully employed for sensing and singlet oxygen applications.327−331 2.1.4. Electron-Transfer Rate Constants. It is often of interest to contrast an electron-transfer rate constant, ket, determined experimentally with those predicted theoretically. Furthermore, with well-defined excited states, theory can provide valuable insights into the electron-transfer parameters associated with the redox chemistry of halogen species. An excellent starting point for such discussions is the seminal work of Marcus, who reduced the many-fold potential surfaces for electron transfer to two parabolic Gibbs free energy surfaces that represent the electron-transfer reactants and products as a function of a single reaction coordinate with fixed force constants, Figure 15.332 The semiclassical expression for ket is given by eq 23, where HAB is the electronic coupling between the donor and the acceptor, λ is the reorganization energy, and ΔG° is the Gibbs free energy change for electron transfer. Below, these three terms are described, and the validity of this familiar Marcus equation is discussed within the context of halide photoredox chemistry. ket =

2π |HAB|2 ℏ

2 1 ji (λ + ΔG°) zyz expjjj− z j 4λkBT zz{ 4πλkBT k

(23)

The mixing of the donor and acceptor wave functions at the instance of electron transfer is contained in the matrix element HAB.333,334 In outer-sphere charge-transfer complexes like Kosower’s salt described in section 3, the molar absorption coefficient and energy of the charge-transfer absorption provide direct information on the magnitude of HAB.335,336 In bimolecular diffusional electron-transfer reactions, the coupling within the encounter complex is generally unknown. With a highly polarizable donor like iodide, the coupling may be sufficiently strong that the semiclassical relation given in eq 23 is no longer valid. Indeed, halides are ambidentate ligands that are known to bridge between discrete redox centers and mediate the electronic coupling.337 A classic example of this was reported in a seminal study from Henry Taube in 1953.338 In this example, the reaction between [CoIIICl(NH3)5]2+ and

ΔG° = NA{e[E°(D+•/D) − E°(A /A−•)] + ω(D+•A−•) − ω(DA)} − ΔE0,0

(24)

Here, E°(D+•/D) − E°(A/A−•) represents the difference in reduction potentials. It is also common to include the free energy stored in the excited state (ΔGES) within a sensitizer reduction potential as an excited-state reduction potential.346−348 The magnitude of ΔGES is estimated either through the photoluminescence onset or through a Franck−Condon line shape analysis of the photoluminescence recorded at 77K.349−351 The term ΔE0,0 is the energy separation between the lowest vibrational levels of the ground and excited state. The term ω represents the electrostatic work required to bring R

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Figure 16. Description of the effects of changing the difference in charge, Δz, between D and A (A), the solvent dielectric constant, εr (B), and the intermolecular distance (C) have on the work term. Data were calculated from eq 29, with (A) the given Δz values, the εr of the solvents indicated, and an intermolecular distance of 6 Å, (B) the given intermolecular distances, Δz values indicated and solvents shown, and (C) the given change in εr for the Δz values shown and the intermolecular distances indicated.

Work terms estimated through a Coulomb’s law approximation typically assume point charges separated at a van der Waals distance.340,354 This assumption does not account for the case where ions have multiple charges in specific locations. In theory, if the individual partial charge of every atom on the donor and acceptor was known along with the electron-transfer distance, the calculated work terms could be more accurately estimated than by the simple sphere model. Recently, density functional theory along with natural population analysis has been proposed as a means to provide reasonable estimates of the individual atomic charges in complex systems.212,221 However, in these cases, a transition metal acceptor was combined with a simple halide donor to simplify the reaction chemistry of interest that is discussed further toward the end of this section. Three variables in Coulomb’s law analysis affect the work term: the intermolecular distance, a, dielectric constant, ε, and the difference in charge between products and reactant (Δz). As seen in eq 29, the work term is directly proportional to Δz, which determines both its magnitude and sign. Figure 16A shows the effect of modulating Δz between 6+ and 6-. The xintercept at Δz = +1 is the only case in which the work term is equal to zero. The intermolecular distance, a, dependence decreases in magnitude with increasing a, Figure 16B. Finally, the solvent dielectric constant also controls the magnitude of the work term. As the dielectric constant decreases, the work term becomes larger, Figure 16C. These three factors control the magnitude and sign of the work term, and while each has the potential to be individually tuned, one factor may influence the others making systematic studies difficult. As an example of how the work term may play a role in halide electron-transfer chemistry, consider photoinduced electron transfer from iodide (I−) to the photoexcited ruthenium polypyridyl complex [Ru(deeb)3]2+*. The difference in formal reduction potentials is −0.22 V in CH3CN.273 The ruthenium acceptor (2+ charge) and iodide donor (1− charge) gives a Δz of 3+ and therefore a positive work term value of +0.13 V in CH3CN. In water, this value is more than halved to +0.06 V, whereas in CH2Cl2, the work term more than doubles to +0.54 V. Therefore, while electron transfer would be expected to proceed readily in both CH3CN and H2O, in CH2Cl2 the electron transfer would have a 300 mV uphill driving force. This highlights how under special circumstances the magnitude of the work term may control the redox chemistry.

the two reactants together and to separate them after electron transfer. The value of ω is typically estimated through a Coulomb’s law analysis, eq 25 and 26, where z is the ionic charge of the reactants or products, εo is the vacuum permittivity, εr is the dielectric constant of the solvent or relative medium static permittivity, and a is the donor− acceptor intermolecular electron-transfer distance.345 ω(D+•A−•) =

ω(DA) =

z(D+•)z(A−•)e 2 4πε0εra

z(D)z(A)e 2 4πε0εra

(25)

(26)

Note that the equations are shown for the interactions of a neutral donor and acceptor pair. As only one electron is transferred in a typical photoinduced electron-transfer process, the work term equations can be simplified. Equation 25 can be replaced by eq 27 to match the one electron lost and gained by the donor and acceptor, respectively. Substitution of eq 26 and 27 into eq 24 gives eq 28, in which the overall work term is now defined by eq 29. This defines the work term by the difference in the charge of the reactant donor and acceptor, removing confusion, and simplifying analysis. ω(D+•A−•) =

[z(D) + 1][z(A) − 1]e 2 4πε0εra

l o ΔG° = NA o me[E°(D+•/D) − E°(A /A−•)] o o n o [z(A) − z(D) − 1]e 2 | o + − ΔE0,0 } o o 4πε0εra ~ ω = Gw =

[z(A) − z(D) − 1]e 2 [Δz − 1]e 2 = 4πε0εra 4πε0εra

(27)

(28)

(29)

Focusing on photoinduced electron-transfer chemistry, many studies utilized neutral donor−acceptor pairs,344,352,353 while others have probed the work term in cases where at least one or both of the reactants are charged.352,354−356 In many cases, the work terms can be safely ignored as their magnitude is less than 2kT (∼50 mV).354,357,358 In other cases where the solvent dielectric constant is small and/or the charge, z, is high, ω should be considered.359−361 This is often the case for halide redox chemistry in organic solvents. S

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Values of ΔG° for halide photoredox chemistry are of particular interest for excited states with known reduction potentials as they provide valuable estimates of halide formal reduction potentials that are not easily garnered by other experimental means.40 Hence, when the overall driving force for the reaction and the excited-state reduction potential of the sensitizer are known, the reduction potential of the halide can be estimated. Indeed, the one-electron reduction potential of the halides (I−, Br−, Cl−) have been estimated in acetonitrile and in acetone in different studies211,292 as well as the one electron reduction potential of triiodide.57 Ideally, ket values are determined for a series of excited-state complexes whose reduction potentials have been tuned without significant change to the geometry and electronic coupling within the encounter complex. With such data, the pre-exponential terms may be lumped together and a consistent value of Gibbs’ free energy for the reaction can be found, eq 30. ket = Ae−(ΔG

o

+ λ)2 /4λRT

(30)

2.2. Dynamic Quenching

2.2.1. Iodide Photo-Oxidation. In early iodide quenching studies of [Ru(bpy)3]2+*, Demas and co-workers reported Ksv < 1 M−1, implying that a stronger photo-oxidant was necessary to realize efficient quenching.362−365 Toward this end, a set of Cr(III) complexes with long-lived ligand-field excited states were found to be quenched by iodide with kq = 1 × 109 M−1 s−1.366,367 Lever and co-workers reported that ruthenium complexes bearing 2,2′-bipyrazine type ligands were potent photo-oxidants.362−364 Indeed, the MLCT excited state of [Ru(bpz)2(deeb)]2+ was efficiently quenched by iodide in CH3CN solutions. Stern−Volmer analysis (vide infra) of the PLI and lifetime yielded a kq of 6.6 × 1010 M−1 s−1 that was fully consistent with dynamic quenching. Direct evidence of an electron-transfer quenching mechanism was reported by Gardner et al. in 2008.219 Pulsed light excitation of [Ru(bpz)2(deeb)]2+ in iodide CH3CN solution resulted in the appearance of the oneelectron reduced ruthenium complex with a rate constant, k1, that was equal to kq. This indicated that the reduced Ru complex was a primary photochemical product of the excitedstate reaction. The oxidized iodide species, I2•−, was also observed spectroscopically in the expected 1:1 stoichiometry with the reduced Ru complex. The delayed formation of I2•− was consistent with a two-step mechanism, eqs 31 and 32, where I• was formed in the first step through excited-state electron transfer, and I• reacted with iodide to form I2•− in the second step.

Figure 17. (top) Rate constants for the formation of I2•− and the monoreduced [Ru(bpz)2(deeb)]+ complex as well as excited-state decay as a function of iodide concentration. (bottom) Absorption difference spectra measured in acetonitrile at the indicated time delay after pulsed 532 nm light excitation of a solution containing [Ru(bpz)2(deeb)]2+ and 500 mM iodide. Adapted with permission from ref 218. Copyright 2009 American Chemical Society.

photoproduct that subsequently reacted with iodide to yield I2•−.317,319 Justification for this proposal came from parallel studies, where iodine atoms were generated with ultraviolet light and the rate constant for I2•− formation was found to be within experimental error the same as that for the dyesensitized reaction, k2 = 2.4 × 1010 M−1 s−1.217,218 Later, Farnum et al. utilized a series of ruthenium(II) polypyridyl complexes to tune the driving force for iodide oxidation in acetonitrile, Figure 12.273 The same two-step mechanism reported for [Ru(bpz)2(deeb)]2+ was found to be operative (eqs 31 and 32). Surprisingly, large rate constants of 109 M−1 s−1 were quantified even when ΔG° was near zero. Such rapid reactivity without a free energy incentive suggests strong electronic coupling within the encounter complex and the possibility of adiabatic electron-transfer pathways.338,368 The second-order rate constants for I2•− formation were identical for all ruthenium complexes except [Ru(dmb)2(bpz)]2+ and [Ru(bpy)2(deeb)]2+ that were rate limited by the formation of iodine atoms (Table 8). The series of ruthenium complexes and the corresponding kinetic data presented in Table 8 represent a textbook example of how Marcus theory can be used to estimate formal

[Ru III(bpz−)(bpz)(deeb)]2 +* + I− k1

→ [Ru II(bpz−)(bpz)(deeb)]+ + I• k2

I• + I− → I•− 2

(31) (32)

Figure 17 shows transient absorption spectra at indicated delay times with overlaid simulations based on the 1:1 stoichiometry of monoreduced ruthenium complex and I2•−. The monoreduced ruthenium complex absorbs light at 500 nm, while I2•− has spectral features below 400 nm and above 600 nm. The rate constant for the appearance of I2•− (2.4 × 1010 M−1 s−1) was found to be smaller than kq (6.6 × 1010 M−1 s−1), indicating that I2•− was not a primary photoproduct. Instead, the iodine atom was proposed to be the primary T

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Table 8. Summary of Thermodynamic Parameters, Rate Constants, and Formal Reduction Potentials in Acetonitrile from ref 273 compound

ΔGes

ΔG°*

[Ru(dmb)2(bpz)]2+ [Ru(bpy)2(deeb)]2+ [Ru(deeb)2(dmb)]2+ [Ru(deeb)3]2+ [Ru(dmb)2(deeb)]2+ [Ru(deeb)2(bpz)]2+ [Ru(bpz)2(deeb)]2+

1.87 1.99 2.02 2.13 2.15 2.07 2.16

0.04 0.00 −0.07 −0.22 −0.25 −0.27 −0.40

k1 (× 1010) (0.57 (0.79 (3.5 (4.9 (5.6 (5.4 (6.6

± ± ± ± ± ± ±

0.06) 0.11) 0.7) 0.2) 0.4) 0.7) 0.2)

k2 (× 1010)

kact

E°(I•/−) calculatedb

± ± ± ± ± ± ±

(6.2 ± 0.6) × 109 (9.0 ± 0.6) × 109 (7.9 ± 1.7) × 1010 (2.1 ± 0.3) × 1011 (4.5 ± 0.1) × 1011 (3.3 ± 0.4) × 1011

1.20 1.22 1.18 1.27 1.25 1.29

(0.54 (0.71 (2.6 (2.2 (2.2 (2.3 (2.4

a

0.20) 0.20) 0.1) 0.2) 0.3) 0.2) 0.2)

a

b

Correction could not be applied because k1 was within experimental error equal to kd. Calculated using eqs 19, 20, 21, and 23.

reduction potentials. The value of kact, i.e., the activated rate constant (eq 15), in the case of [Ru(bpy)2(deeb)]2+ was 9.0 × 109 M−1 s−1. The Onsager radius Rc was 2.99 nm, eq 19. A Debye length, κ = 4.25 × 108 m−1, at an ionic strength of 8 mM that with ionic radii of 7 and 2.06 Å for the ruthenium complex and iodide, respectively, provided an association constant, KA = 127 M−1, eq 21. The electron-transfer rate constant, ket = 7.10 × 107 s−1, was then determined by dividing kact by KA. The electron-transfer driving force, ΔG° = −9 meV was determined using eq 23, with assumption that λ = 1 eV. The preexponential factor that encompasses the matrix element HAB, λ and several constants was estimated as 1012 s−1 and was used as A in eq 30 for all the ruthenium complexes. Values of preexponential factors, as reported by Sutin, usually range between 1012 and 1013 s−1 for ruthenium polypyridyl complexes.323 The reduction potential E(I•/−) = 0.981 V vs SCE (1.225 V vs NHE) was then calculated from ΔG° = −9 meV and E(Ru2+*/+) = 0.99 V vs SCE. A value of E(I•/−) = 1.23 V vs NHE was determined by averaging the values obtained for all six complexes in Table 8 that coincides with the accepted literature value. 2.2.2. Triiodide Reduction. In an early report, Clark et al. reported the quenching of ruthenium excited state by triiodide in organic solutions and when anchored to TiO2 interfaces.293 In dichloromethane solutions, both static and dynamic excitedstate quenching mechanisms were operative, indicative of ground-state ion pairing between I3− and [Ru(bpy)2(deeb)]2+. Only dynamic quenching was observed in acetonitrile solution.

Figure 18. Transient absorption difference spectra recorded after 532 nm laser excitation of an argon-purged acetonitrile solution that contained Ru2+, I−, and I3−. The flash-quench cycle where Ru2+ is the chromophore, I− is the electron donor, and I3− is the electron acceptor is indicated. Reproduced with permission from ref 369. Copyright 2013 American Chemical Society.

The rate constant for I3− reduction was used with Marcus theory to estimate the formal reduction potential E°(I3−/ (I2•−,I−)) = −0.33 V vs NHE. This value was considerably more negative than the accepted value in water and helped explain how I3− avoids recombination with electrons injected into TiO2. These TiO2 electrons are weak reductants with a TiIV/III reduction potential more positive than −0.33 V vs NHE, which therefore makes I3− reduction a thermodynamically uphill reaction. 2.2.3. Diffusional Bromide and Chloride PhotoOxidation. Rabani et al. described photodriven bromide oxidation by visible light excitation of [Ir(bpy)2(C3,N′-bpy)]2+ in aqueous solutions.370 The C3,N′-bpy is a 2,2′-bipyridine ligand that is bound to the metal center through one nitrogen and one cyclometalating C−Ir bond.371−374 In steady-state experiments, where the iridium complex was excited in the presence of bromide and oxygen, both tribromide (Br3−) and hydrogen peroxide (H2O2) were produced. The quantum yields of products were found to depend on the initial concentration of reactants and the solution pH. However, the UV−vis absorption spectra showed no indications of groundstate adduct formation between the iridium complex and bromide. Photoluminescence from the iridium complex was quenched by bromide, and the PL maximum shifted to higher energy with increased bromide concentration, behavior that

[Ru III(bpy)2 (deeb−)]2 +* + I3− → [Ru III(bpy)2 (deeb)]3 + + I32 −

(33)

Although the electron-transfer products were not directly detected, cathodic photocurrents in dye-sensitized solar cells indicated an oxidative quenching mechanism was operative when the complex was immobilized on the oxide surface. If this mechanism also occurred in fluid solution (eq 33), then quantitative back-electron transfer from I32− to the oxidized Ru complex must occur within the encounter complex. To better understand the one electron reduction of I3−, flash-quench studies were performed (Figure 18).57,369 In these experiments, the MLCT excited state was reduced by iodide in solutions that contained a large excess of I3−. In this manner, the reduced Ru product was most likely to undergo collisional reactions with I3−. Indeed, the diiodide radical was identified as the reaction product. Spectral modeling of the transient absorption spectra with the reduced complex and I2•− absorption clearly indicated that I2•− was a direct product of I3− reduction. U

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Bromide was then oxidized by RuIII, resulting in a recovery of the ground-state ruthenium complex and a bromine radical (eq 35). Control experiments showed that the oxidative quenching of the excited state by ArN2+ was 2 orders of magnitude faster than reductive quenching by bromide. The bromine radical formed in solution was found to react with bromide to form Br2•− within the diffusion limit. Spectral evidence for Br2•− was obtained, and the rate constant of Br2•− formation was determined to be k = 5 × 109 M−1 s−1. Overall, the turnover number (TON) for bromine formation by ruthenium photocatalysis was estimated to be 230. In a follow-up study of these consecutive bimolecular reactions, a series of six different ruthenium complexes [Ru(bpy)3]2+, [Ru(bpy)2(deeb)]2+, [Ru(deeb)2(dmb)]2+, [Ru(deeb)2(bpy)]2+, [Ru(deeb)3]2+, and [Ru(deeb)2(bpz)]2+ was investigated.382 These complexes were evaluated for their photophysical and redox properties and were found to have RuIII/II potentials that ranged from 1.51 to 1.91 V vs NHE. Oxidative quenching with ArN 2 + was shown to be thermodynamically favorable for all complexes, with ΔG° ranging from 700 to 170 mV. The [Ru(deeb)2(bpz)]2+* excited-state was a weak photoreductant that was not quenched significantly by ArN2+. For the other five complexes, the bromide oxidation by the RuIII complexes was quantified spectroscopically. The formation of bromine was also confirmed by postreaction studies with 1hexene.383 Indeed, the steady-state illumination of a solution containing [Ru(bpy)2(deeb)]2+ in the presence of ArN2+ and bromide resulted in the formation of Br2 that reacted with 1hexene to yield 1,2-bromohexane. Excited-state quenching and bromide oxidation for the six complexes were evaluated using Marcus theory. The work terms for the reactions were evaluated and found to be very similar for all the complexes. Reorganization energies were also estimated from eq 23 by plotting the rate constants versus the Gibbs free energy. The reorganization energy for ArN2+/0 was estimated to be 3.07 eV, while the Br•/− was estimated to 2.33 eV. The large reorganization energy for ArN2+ was explained by the bond breaking involved in the reorganization.

was attributed to an excited-state complex, or exciplex, between bromide and [Ir(bpy)2(C3,N′-bpy)]2+*. In time-resolved PL experiments, bromide was reported to dynamically quench the excited state, but at higher concentrations, the kinetics were better described by a biexponential kinetic model. The shorter lifetime was ascribed to the iridium complex quenched by bromide, and the longer lifetime was assigned to the exciplex, which in turn was dependent on the bromide concentration. Pulsed laser excitation resulted in an absorption change at 360 nm, and this transient species was concluded to be Br2•−. Very recently, an iridium complex, [Ir(dFCF3ppy)2(dtb)]+,375 where dF-CF3ppy is 3,5-difluoro-2-[5(trifluoromethyl)-2-pyridinyl-N]phenyl-C, has been proposed to generate Cl• upon light-induced oxidation of chloride.376 This proposal was based on prolonged photolysis studies where hydrogen abstraction products were quantified. Excitedstate quenching experiments in acetonitrile with TBACl yielded Stern−Volmer constants, KSV = 2.5 and 1.9 M−1, which would correspond to a small quenching rate constant of 7.6 × 105−1.0 × 106 M−1 s−1, assuming an excited-state lifetime of 2.5 μs. HCl oxidation to Cl• has also been proposed using acridinium derivatives.377 MacMillan recently reported evidence for the photooxidation of bromide using [Ir(dFCF3ppy)2(dtb)]+*.378−380 Despite the absence of spectroscopic evidence toward electron-transfer products, the large Stern−Volmer rate constant (KSV = 18000 M−1) led to the conclusion that bromide was oxidized to Br•. 2.2.4. Photocatalytic Bromide Oxidation Using Diazonium Sacrificial Agents. Yun-Da Tsai et al. have reported the photocatalytic oxidation of bromide by ruthenium polypyridyl species in acetonitrile solution.381,382 In their early study, photooxidation of bromide was achieved with the [Ru(deeb)2(dmb)]2+* excited state. However, the secondorder rate constant for bromide oxidation k = 2 × 106 M−1 s−1 was small, resulting in inefficient excited-state quenching attributed to a small thermodynamic driving force. In an alternative approach, the excited state was first oxidized with a sacrificial electron acceptor (bromobenzenediazonium tetrafluoroborate ArN2+BF4−), yielding the RuIII complex that subsequently oxidized bromide to the bromine atom, Br•, Figure 19.

2.3. Static and Dynamic Halide Photoredox Chemistry

2.3.1. Ion Pairing in Low Dielectric Constant Solvents. Solvent polarity can influence reaction mechanisms. As was mentioned previously, dynamic quenching of [Ru(bpy)2(deeb)]2+* by iodide was exclusively observed in acetonitrile (ε = 38) while static electron transfer was dominant in dichloromethane (ε = 9), Figure 13. One can imagine that a solvent with an intermediate dielectric constant, such as acetone (ε = 21) would show intermediate behavior. In 2005, Clark et al., followed by others, provided demonstrations of a static electron-transfer quenching mechanism for iodide photooxidation by [Ru(bpy)2(deeb)]2+* (Figure 13).317−319 It was shown through Benesi−Hildebrand and Stern−Volmer analysis that iodide ion-pairs were formed with a large equilibrium constant, Keq = 59700 M−1 in dichloromethane solution. This ion-pairing process was reversible and could be inhibited by increasing the ionic strength of the solution with an inert salt such as tetrabutylammonium hexafluorophosphate (TBAPF6). Nanosecond transient absorption experiments with the ion pairs revealed that the monoreduced complexes were formed within the instrument’s response time, confirming therefore that the electron transfer occurred within the ion-paired excited states.

Figure 19. Catalytic cycle describing the mechanism of bromide oxidation.

The first reaction was the irreversible electron transfer from the excited state to the ArN2+ to yield the RuIII complex (eq 34). [Ru III]2 +* + ArN+2 → [Ru III]3 + + Ar• + N2

(34)

[Ru III]3 + + Br − → [Ru II]2 + + Br •

(35) V

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The reaction yielded the monoreduced compound and an iodine atom as primary photoproducts. Diiodide was formed with a smaller rate constant, showing that it was again a secondary photochemical reaction product. Ion pairs are often divided into two classes: “contact” or “solvent-separated” ion pairs. A contact ion pair, also referred to as a tight or intimate ion pair, is produced when no solvent molecules are located between the ions, as opposed to a solvent-separated ion pair where solvent molecules lie between the ions. The electronic coupling between the ions is expected to be smaller for the solvent-separated ion pairs than for the contact ion pair. Common techniques to investigate the nature of ion pairs, which are encompassed in the broader context of host−guest chemistry, are NMR,220,221,292,384−395 UV−visible,220,221,292,384,389,390,394,395 and photoluminescence spectroscopy220,221,292,306,384,389,394,395 as well as isothermal titration calorimetry (ITC).342,396−398 To quantify ion pair structures without excited-state reactivity, Ward et al. utilized chloride as an “innocent” ion, as its large reduction potential precluded excited-state electron transfer.384 A series of ruthenium complexes, i.e., [Ru(bpy)3−x(deeb)x]2+ where x = 0, 1, 2 or 3, and chloride was studied by 1H NMR, UV−vis, and photoluminescence spectroscopy in CH2Cl2. Large downfield shifts of the 3,3′-H atoms of the ancillary 2,2′-bipyridine was evidenced by 1H NMR spectroscopy of [Ru(bpy)2(deeb)]2+ with increasing amounts of chloride. The downfield shifts were rationalized by hydrogen bonding between chloride and the 3,3′-hydrogen atoms that resulted in an elongation of the C−H bond and hence a decreased electron density. The 6,6′-H atom resonances also exhibited downfield shifted upon chloride addition, albeit to a lesser extent. Very interestingly, only 0.1 equiv of Cl− were necessary to induce significant changes in the 1H NMR spectra, with no evidence for signal broadening or multiple sets of resonances that would correspond to the singly ion-paired species and the nonchloride paired complexes. At the concentration used, it was concluded that the chloride ion undergoes rapid intermolecular exchange on the NMR time scale. The higher affinity for ion pairing with the 2,2′-bipyridine ligand over the deeb ligand suggested that synthetic engineering of ligands could be used to promote ion pairing in solution at specific locations.220,221,394 The photophysical study of the homoleptic compounds was unfortunately complicated by ligand-loss photochemistry that is often observed for RuII polypyridyl complexes.303 This ligand loss occurs through population of the ligand-field state and is often observed with potent photo-oxidants or when the reduction potentials of the ancillary ligands are very similar. Crystal structures of [Ru(bpy)2(deeb)]2+ crystallized in the presence of chloride, Figure 20, were obtained where the closest chlorides were proximate to the 3,3′-H atoms of the 2,2′-bipyridine (at 2.5 and 2.8 Å) as well as next to the 6,6′-H atoms of the deeb ligand (2.7 Å). 2.3.2. Supramolecular Assemblies with Halide Ions. Recently, [Ru(dtb)2(dea)]2+, where dtb is 4,4′-di-tert-butyl2,2′-bipyridine and dea is 4,4′-diethanolamide-2,2′-bipyridine, was reported.221 The dea ligand was specially designed to favor interactions with halides and was indeed shown to form 1:1 adducts with chloride, bromide, and iodide in acetonitrile and dichloromethane. The stoichiometry and the position of ionpairing was identified through 1H NMR spectroscopy and Job plots. 1H NMR titration with these halides resulted in significant shifts of the 3,3′-H atoms of the dea ligand as well as with the resonances associated with the amide and the

Figure 20. ORTEP diagram for [Ru(bpy)2(deeb)]2+ showing the closest four chloride ions. The interactions of the chlorides (lightgreen dashed lines) with the 3,3′-H atoms of the 2,2′-bipyridine are more clearly seen in the viewing angle on the left, and the interactions with the 6,6′-H atoms of the deeb ligand are more clearly seen in the view on the right. Ellipsoids drawn at the 50% probability level. Reproduced with permission from ref 384. Copyright 2013 American Chemical Society.

alcohol groups. This indicated that the dea ligand wrapped around the different halides, with the 3,3′, NH, and OH protons directing the interaction, Figure 21. Evidence for interaction with the 3,3′-H atoms of the 4,4′-di-tert-butyl-2,2′bipyridine was also found at high chloride concentration in CD2Cl2. The addition of one equivalent of chloride to [Ru(dtb)2(dea)]2+ in dichloromethane solution led to a 100% increase in the photoluminescence quantum yield. This

Figure 21. Plausible dynamic quenching of [Ru(dtb)2(dea)]2+* by iodide. The presence of the electron on the dea ligand is emphasized by the orange-colored bipyridine. Purple spheres represent iodine species, whereas green spheres represent chloride. Ancillary dtb ligands are omitted for clarity. Reproduced with permission from ref 221. Copyright 2016 American Chemical Society. W

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increase in quantum yield was also accompanied by a blue-shift in the PL spectra and an increase in excited-state lifetime in agreement with the energy gap law.399 The addition of one equivalent of iodide also enhanced the excited state, but to a lesser extent.221 More interestingly, the further addition of iodide caused drastic excited-state quenching. Steady-state and time-resolved experiments showed that this quenching process was solely dynamic in nature, as evidenced by the corresponding Stern−Volmer plots. The excited-state quenching rate constant by iodide was 2.7 × 1010 M−1 s−1. A blue-shift accompanied the increased PLI observed in the 1:1 halide adducts that corresponded to about a 60 meV increase in the free energy of the excited state. Yet, this increase alone could not account for the tremendous turn-on in excited-state quenching in the 1:1 ion pair. Several excitedstate quenching mechanisms were considered to understand how iodide first enhanced the excited-state photoluminescence and lifetime and then quenched it at higher iodide concentration. These mechanisms are shown in Figure 21: (A) a concerted mechanism in which electron transfer and I−I bond formation occurred in one concerted step, (B) a standard excited-state dynamic quenching by iodide, and (C,D) the dynamic quenching of the halide−ion paired. Mechanism A was disproved as the formation of I2•−, as measured by transient absorption, was delayed compared to the formation of the monoreduced complex. The second mechanism was also disregarded, as there was no evidence for excited-state quenching in the absence of the iodide ion pair. Indeed, experiment performed in the presence of TBAClO4 showed no evidence for excited-state quenching. The excited-state quenching by iodide occurred with the same quenching rate constant when {[Ru(dtb)2(dea)]2+, Cl−}+ was initially formed by the addition of one equivalent of TBACl, kq = 2.7 × 1010 M−1 s−1. A Debye−Hückel analysis was very informative and showed that the excited-state reaction occurred with a monocationic excited-state, i.e., {[Ru(dtb)2(dea)]2+, X−}+. These experiments were hence more in line with mechanisms C and D in Figure 21, where the excited-state ion-paired complex is dynamically quenched by iodide through the twostep mechanism that was previously described. It was concluded that coordination of I− (or Cl−) by the dea ligand was kinetically fast relative to halide photo-oxidation. Once the iodide was coordinated to dea, it was inert as excited-state oxidation of the H-bonded iodide was thermodynamically uphill. Work term calculations were used to better understand the thermodynamics for electron transfer. Natural population analysis and density functional theory allowed assignment of a partial charge to each atom of [Ru(dtb)2(dea)]2+. With this, the Coulombic charge interaction between the two species could be calculated at any distance. Contour plots of the calculated stabilization energy at van der Waals distances showed a 500 meV Coulombic incentive for the halide to form an adduct with the ruthenium complex in the dea ligand, Figure 22. This 500 meV stabilization represented a considerable decrease in the reaction driving force that would inhibit static excited-state electron transfer, leading instead to dynamic excited-state quenching from iodide in solution. 2.3.3. Ruthenium Complexes with Polycationic Ligands. Most mononuclear ruthenium compounds used for excited-state electron transfer were dicationic (2+) in charge. Swords et al. developed a series of highly charged ruthenium

Figure 22. Contour plots of the calculated Gibbs free energy stabilization, Gw in eV, over the plane containing the dea ligand in the absence (A), and presence (B), of the chloride ion pair. All atoms within 1 Å of this plane are shown as small colored dots. The dea ligand is superimposed in white. Reproduced with permission from ref 221. Copyright 2016 American Chemical Society.

polypyridyl complexes bearing the dicationic tmam ligand, where tmam is 4,4′-bis(trimethylaminomethyl)-2,2′-bipyridine, which allowed for the synthesis of ruthenium(II) complexes bearing a total charge of 4+ ([Ru(deeb)2(tmam)]4+), 6+ ([Ru(tmam)2(deeb)]6+), and 8+ ([Ru(tmam)3]8+), the latter being the highest charged mononuclear ruthenium complex to date, Figure 23.400 Equilibrium constants for iodide ion-pairing in acetonitrile increased with the cationic charge, Keq = 4000 (4+), 4400 (6+), and 7000 (8+) M−1, respectively. 1H NMR experiments showed that iodide interacted with the tmam ligands, localizing in a “binding pocket”, with the key shifts being displayed by the 3,3′-H atoms of the tmam ligand. Furthermore, the ion-pairing of iodide led to resolution of the −CH 2 -N + (CH 3 ) 3 enantiotopic protons. Excited-state quenching was shown to occur efficiently through a combination of static and dynamic excited-state quenching, with a bimolecular quenching rate constant of 6.3 × 1010 M−1 s−1 for [Ru(deeb)2(tmam)]4+. This value was very close to the reported diffusion-limited electron transfer of 6.4 × 1010 M−1 s−1 in acetonitrile. Quenching rate X

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was investigated by 1H NMR, UV−vis, as well as time-resolved and steady-state photoluminescence. The atomic resolution provided by 1 H NMR spectroscopy was particularly informative. As expected, the 3,3′ protons on the tmam ligands displayed large downfield shifts as well as the alkyl protons on the tmam ligand resolved in a roofed doublet with increased iodide concentration in deuterated acetone. Interestingly, additional shifts associated with the 5,5′-H atoms of the tmam ligand as well as with the 6-H atoms of the ancillary 4,4′di-tert-butyl-2,2′-bipyridine ligands suggested that two iodides were ion-paired with [Ru(dtb)2(tmam)]4+ and formed a socalled “ter-ionic complex”, Figure 24. The first ion-paired iodide was thought to locate within the tmam binding pocket, whereas the second was suggested to be located between the tmam ligand and the ancillary ligands around the 5,5′ position. UV−vis spectroscopy provided two equilibrium binding constants, Keq1 = 1.7 ± 0.5 × 106 M−1 and Keq2 = 1.5 ± 0.5 × 105 M−1. The formation of a ter-ionic complex was also supported by natural bond order analysis and density functional theory that located the two iodides 5.8 Å apart of each other. The formation of this ter-ionic complex and the influence of each additional iodide can be understood by Figure 25, where the work term is represented for the tmam and for a plane located 2−3 Å above and below this tmam plane. When the structure was optimized without any iodide, Figure 25a, the work term map clearly indicated strong Coulombic incentive to bind a halide in the area between the substituents on the tmam ligand where the strongest incentive (450 mV) locates an iodide in the designated tmam binding pocket. This can be seen in the purple region of horizontal contour surfaces located on the tmam ligand, Figure 25a. DFT also showed that there was an approximately 250 mV incentive to bind a second iodide once the first iodide was located in the tmam binding pocket, Figure 25b. Indeed, the work term map indicated that the positions above or below the tmam ligand, closer to the ruthenium center, was the preferred location for the second ion-pairing event. Note that these positions are equivalent in the singly ion-paired species and that the shifts observed by 1H NMR spectroscopy agreed with such an ion pair. Once the {[Ru(dtb)2(tmam)]4+,2I−}2+ ter-ionic complex formed, there was no specific location for additional iodide ions to associate with the complex, Figure 25c. The excited-state behavior of [Ru(dtb)2(tmam)]4+* in the presence of iodide was unique. Indeed, excited-state quenching resulted in Stern−Volmer plots that saturated at high iodide concentrations, both in the PL intensity and in the excitedstate lifetime. As stated in section 2.1.2, Stern−Volmer plots of τ0/τ and I0/I are coincident and linear with a slope that provides the quenching rate constant when the quenching is purely dynamic. When the excited-state quenching is solely static, as for example in the case of a nonemissive ground state adduct, then τ0 is independent of the quencher concentration and plots of I0/I versus the quencher concentration give the equilibrium constant for the ground-state adduct. In many cases, the excited-state quenching is a combination of both dynamic and static quenching and results therefore in an upward curvature. However, in this study, the Stern−Volmer plots displayed saturation at high iodide concentration, Figure 26. The saturation and nonlinear behavior in Stern−Volmer plots was rationalized by the contributions from three photoluminescent species that were present in solution

Figure 23. Ruthenium(II) complexes bearing dicationic 4,4′-bis(trimethylaminomethyl)-2,2′-bipyridine (tmam) ligand.220,292,394,400 Note that every tmam ligand contributes an additional +2 charge to the dicationic ruthenium polypyridyl complex.

constants for [Ru(tmam)2(deeb)]6+ and ([Ru(tmam)3]8+ could not be obtained as the PLI and lifetime ratios were nonlinear with iodide concentration. The use of complexes bearing tmam ligands for ion-pairing were recently translated to the dye-sensitized TiO 2 interface, where a static regeneration mechanism was operative in the case of [Ru(tmam)2(dcb)]6+ and an anionic cobalt redox mediator.276 Ground-state ion pairing was clearly evident at the sensitized TiO2 interface that nearly doubled the light-to-electrical energy conversion efficiency compared to a sensitizer in which no ion pairing occurred. With the promising results obtained for ion pairing of iodide in acetonitrile, the rational development was to aim to favor higher stoichiometry ion pairs by decreasing the solvent dielectric constant. Hence, Wehlin et al. recently reported the ion-pairing properties of a tmam complex with halides in acetone, Figures 23 and 24.220,394 The formation of ion pairs

Figure 24. Proposed structure of {[Ru(dtb)2(tmam)]4+,2I−}2+ with the physical location of the two iodides determined by 1H NMR in deuterated acetone. Reproduced with permission from ref 220. Copyright 2018 American Chemical Society. Y

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Figure 25. Representations of ground-state structure of (a) [Ru(dtb)2(tmam)]4+, (b) {[Ru(dtb)2(tmam)]4+,I−}3+, and (c) {[Ru(dtb)2(tmam)]4+,2I−}2+. Overlaid on each structure are three representative contour plots of the calculated stabilization energy, Gw in eV, calculated as a work term over each plane with zero, one or two iodides, respectively.

dynamically quenched by iodide and was the predominant luminescent species at high iodide concentration. At iodide concentrations where the ter-ionic species was the predominant species, the formation of the monoreduced complex and I2•− reached a plateau and the two rate constants differed by only 60 ± 10 ns. This offset between the growth of the reduced compound, and the growth of I2•− was ascribed to the rate constant for I−I bond formation. The insensitivity of these reactions to the iodide concentration indicated that the formation of I2•− occurred within the ter-ionic complex. The following mechanism for the excited-state quenching within the ter-ionic complex was proposed, Figure 27: The highly cationic [Ru(dtb)2(tmam)]4+ complex promoted ground-state supramolecular assembly of the two iodides in close proximity to each other, a crucial component to overcome the statistical barrier of a three-species encounter complex. Visible light excitation led to electron transfer to generate the monoreduced Ru and I• within the ter-ionic complex. Within 60 ns, the iodine atom diffused toward the second iodide to form I2•−. It was speculated that the Coulombic repulsion between the reduced tmam ligand and the iodide increased the distance between the two iodine species precluded observation of the concerted mechanism. This mechanism underscores the importance of excited-state localization and suggests that more rapid bond formation could be achieved if the excited state were localized away from the ion-pairing ligand. 2.3.4. Excited-State Ion Pairs That Undergo Halide Photorelease. The demonstration that excited-state localization could induce Coulombic repulsion of an ion-paired halide was clearly evident in a study by Turlington et al. that showed chloride photorelease within a supramolecular assembly.395 Two complexes were synthesized in order to examine the influence of the excited-state dipole on the excited-state equilibrium with the possibility for photorelease of chloride. The two complexes shared a common daea ligand, where daea is 4,4′-(aminoethylamino)ethanolamine-2,2′-bipyridine, that was specifically designed to promote ion-pairing with halides. The ancillary ligands were either 4,4′-di-tert-butyl2,2′-bipyridine (dtb) or 4,4′-bis(trifluoromethyl)-2,2′-bipyridine (btfmb) that were used to synthesize [Ru(dtb)2(daea)]2+ and [Ru(btfmb)2(daea)]2+. Both complexes were shown to form ion-pairs with chloride in dichloromethane with large equilibrium constant, Keq ∼ 4 × 10 6 M −1. In [Ru(dtb)2(daea)]2+, the excited-state dipole was localized on the daea ligand, i.e., [RuIII(dtb)2(daea−)]2+*, and hence toward the associated halide while in [Ru(btfmb)2(daea)]2+, the excited state was localized on the btfmb ligand, i.e., [RuIII(btfmb)(btfmb−)(daea)]2+*, and hence away from the associated halide.

Figure 26. (top) Time-resolved PL decays of [Ru(dtb)2(tmam)]4+* with the addition of up to 15 equiv of TBAI. Inset shows the corresponding steady-state PL spectra. (bottom) Corresponding Stern−Volmer analysis with fit overlaid according to a modified Stern−Volmer equation. Reproduced with permission from ref 220. Copyright 2018 American Chemical Society.

throughout the titration. The first species was [Ru(dtb)2(tmam)]4+·4PF6− with a known excited-state lifetime. Control experiments at high ionic strength with inert salts showed that this excited state was not dynamically quenched by iodide. The second species was the mono-ion-paired {Ru4+:I−}3+* species which was dynamically quenched by iodide. The final species taken into consideration was the terionic complex {Ru4+:2I−}2+*. This ter-ionic complex was not Z

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Figure 28. (a) Square scheme for ground- and excited-state equilibria of [Ru(dtb)2(daea)]2+ with chloride. The blue shading indicates the ligand on which the excited state was localized in the MLCT excited state, and the green sphere represents chloride. (b) Transient photoluminescence spectra obtained 45 ns (purple square) and at longer (purple to red) time delays after pulsed 500 nm laser excitation of [Ru(dtb)2(daea)]2+, abbreviated Ru-daea, in the presence of one equivalent of chloride. The photoluminescence spectra in the absence of chloride is also given for reference (black triangles, dashed line). The bold black arrow indicates the spectral shift observed for chloride. The colored arrow indicates the time dependent spectral shift measured. Reproduced with permission from ref 395. Copyright 2018 American Chemical Society.

Figure 27. (top) Observed rate constants for the formation of the monoreduced Ru3+ complex and I2•− as a function of the iodide equivalents. The region highlighted in blue corresponds to diffusional excited-state quenching, leading to the formation of I2•− through the reaction of I• and I−. The region highlighted in green corresponds predominantly to the static excited-state quenching and I−I bond formation within the ter-ionic complex. (bottom) Proposed mechanism of iodide photo-oxidation within the ter-ionic assembly. Light excitation forms an excited state localized on the tmam ligand, which induces a Coloumbic repulsion between the ligand and the iodide in the pocket. After electron transfer from the iodide localized closer to the metal center, the iodine ion has to move across to the remaining iodide in order to form the covalent diiodide bond. Reproduced with permission from ref 220. Copyright 2018 American Chemical Society.

compared to the ground state for {[Ru(dtb)2(daea)],Cl−}+, whereas a 45-fold increase in the excited-state equilibrium constant was determined for {[Ru(btfmb)2(daea)],Cl−}+.395 2.3.5. Bromide and Chloride Photo-Oxidation. Bromide and chloride photo-oxidation, while more thermodynamically challenging than iodide oxidation, offer advantages such as the energy stored in the halogen bonds and the relative abundance of both elements in seawater. Because the one electron E°(X•/−) reduction potentials are more positive than that for iodide, the need for more potent photo-oxidants was anticipated. An early example of bromide oxidation by a ruthenium polypyridyl type complex was performed with [Ru(bpz)2(deeb)]2+ in acetone.211 Two distinct mechanisms of bromide oxidation were revealed. The first was diffusional quenching of the excited state by bromide, which occurred with a rate constant of 8.1 × 1010 M−1 s−1. The second mechanism, operative at high bromide concentrations, was

Visible light excitation of these two chloride-ion pairs resulted in drastically different behaviors. In the case of {[Ru(dtb)2(daea)],Cl−}+, the photoluminescence spectra measured after pulsed laser excitation were shown to redshift with time, Figure 28. This indicated that the photoluminescence shifted in the direction that corresponded to the nonchloride-paired complex. This behavior was therefore consistent with the gradual photorelease of chloride in the excited state. Such a shift in the PL spectra was not observed for {[Ru(btfmb)2(daea)],Cl−}+. Kinetic analysis, as reported for photoacids and photobases,401 revealed that the excitedstate equilibrium constant decreased by a factor of 20 AA

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Figure 29. Reaction mechanism for the light-induced reaction between [Ru(bpz)2(deeb)]2+ and bromide (orange sphere) in acetone. The orange shading on the bpz ligand represents the additional electron density present in the excited state or in the reduced complex.

Figure 30. Three-dimensional contour diagrams developed from the application of Coulomb’s Law and natural population analysis. Contours are plotted at the potential energy for electron transfer from a −1 charge to the ruthenium based partial charges. Therefore, the contour at 540 mV for the non-ion-paired species indicates the work term for the singly ion-paired bromide, and that at 390 mV is for the doubly ion-paired species. Reproduced with permission from ref 212. Copyright 2017 American Chemical Society.

extent (0.09 and 0.12 ppm, respectively) compared to very small shifts of the 3,3′-H atoms of the 2,2′-bipyrazine ligand. Moving from acetone to CH2Cl2 drastically influenced the ground- and excited-state chemistry.212 The titration of bromide into [Ru(bpz)2(deeb)]2+ CH2Cl2 solution resulted in the formation of two consecutive bromide ion pairs, {[Ru(bpz)2(deeb)]2+, Br−}+, and {[Ru(bpz)2(deeb)]2+, 2Br−} with unique isosbestic points, such that two equilibrium constants were determined, Keq1 > 106 M−1 and Keq2 = 2.4 ± 0.4 × 105 M−1. Such a stoichiometry was also revealed by 1H NMR spectroscopy, where the addition of bromide induced large downfield shifts in the 5,5′-H atoms of the 2,2′-bipyrazine ligands as well as in the 6,6′-H atoms of the deeb ligand. This indicated that the bromide ions were located close to the ruthenium center. Resonances due to the 3,3′-H atoms of the 2,2′-bipyrazine and the deeb ligands shifted to a smaller extend and in the opposite direction, i.e., upfield. Very interestingly, resonances associated with the 6,6′-H atoms of the 2,2′bipyrazine ligand shifted downfield upon the first bromide ionpairing event and reverted back upfield upon ion-pairing with the second bromide. All in all, it was concluded that both bromide ions were located next to the ruthenium center, one above the deeb ligand and one below. Density functional theory agreed with the proposed structure, Figure 30.

termed an inner-sphere, or static, quenching mechanism. However, unlike standard static quenching where the adduct is nonphotoluminescent, the ion-pair was photoluminescent in this case. Biexponential kinetics were observed with one lifetime that was independent of Br− concentration and the other that was strongly Br− dependent and corresponded to dynamic excited-state quenching. The electron-transfer rate was measured to be ket = 2.5 × 107 s−1. These are described in Figure 29. The primary electron-transfer products were determined to be the reduced ruthenium complex and a bromine atom (Br•). The bromine atom subsequently reacted with Br− to generate Br2•− with a rate constant of 1.1 × 1010 M−1 s−1. The formal one-electron reduction potential of bromine in acetone was estimated as E°(Br•/−) = 1.44 ± 0.02 V vs NHE using Marcus theory.210,211 Prolonged irradiation led to the formation of both the cis and trans isomers of [Ru(bpz)(deeb)Br2] that were isolated by column chromatography.210 In acetone, a single equilibrium to yield {[Ru(bpz)2(deeb)]2+,Br−} was determined Keq = 8400 M−1.211 1H NMR titrations indicated that the ion-paired bromide was preferentially located next to the ruthenium center.210 Indeed, resonances associated with 5-H atoms of the 2,2′-bipyrazine ligand and 6-H atoms of the deeb ligand shifted to a larger AB

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Figure 31. Time-resolved PL decays measured after pulsed 445 nm excitation of a 50 μM [Ru(bpz)2(deeb)]2+ dichloromethane solution with 0 to 1 (a) and 1 to 2 (b) equivalents of bromide. Reproduced with permission from ref 212. Copyright 2017 American Chemical Society.

In dichloromethane, [Ru(bpz)2(deeb)]2+* was efficiently quenched by bromide through a static quenching mechanism, evidenced by the formation of the nonluminescent ion-pair [Ru2+, Br−]+.400 Indeed, with less than one equivalent of bromide, only the photoluminescence intensity was quenched, while the excited-state lifetime remained unchanged (τ = 900 ns). Between one and two equivalents of bromide, the photoluminescence intensity increased and the excited-state lifetime became biexponential (τ1 = 900 ns and τ2 = 65 ns). When greater than two equivalents of bromide were present, excited-state decay was monoexponential, with a lifetime of 65 ns, Figure 31. This excited-state lifetime corresponded to a static electron-transfer rate constant of 1.5 ± 0.2 × 107 s−1 within the {[Ru(bpz)2(deeb)]2+, 2Br−} ion pair. Transient absorption spectroscopy was performed under conditions where {[Ru(bpz)2(deeb)]2+, Br−}+ was the predominant species. In that case, only the monoreduced complex was observed spectroscopically, with no evidence for Br2•−. It was rationalized that the equilibrium concentration of bromide was so small that the generated Br• was unable to encounter bromide prior to undergoing back-electron transfer with the monoreduced complex. In contrast, Br2•− was observed transiently after photoexcitation of {[Ru(bpz)2(deeb)]2+, 2Br−} in the presence of excess solvated bromide. The cage-escape yields were determined to be 0.05 and 0.2 for the singly and doubly bromide ion-paired species, respectively. The driving force for electron transfer was estimated based on reduction potentials. It was shown that bromide oxidation was favored by 240 and 250 meV for the singly and doubly bromide ion-paired species with corresponding calculated work term contributions of 540 and 660 meV, respectively. Taken together, this data indicated that the reaction driving force for the excited-state electron transfer was uphill by 290 meV and 420 meV for the singly and doubly bromide ion-paired species. The 130 meV decrease in the driving force for electron transfer upon the second ion-pairing was most likely responsible for the enhanced excited-state lifetime and smaller electron-transfer rate constant. The fact that this uphill reaction occurred so efficiently suggested strong electronic coupling between the bromide and the ruthenium center. Wehlin et al. reported evidence for chloride photooxidation in acetone and acetone/water mixtures using a series of 2,2′bipyrazine substituted ruthenium complexes, namely [Ru(bpz)3]2+, [Ru(bpz)2(tmam)]4+, and [Ru(bpz)2(deeb)]2+, Figures 12 and 23.292 Of these three, [Ru(bpz)3]2+ was the most potent photooxidant (Ru2+*/+ = 1.8 V vs NHE), while the two other complexes shared about the same excited-state

reduction potential (Run+*/(n−1)+ = 1.7 V vs NHE). The barrier between the MLCT excited state and the dissociative ligandfield (LF) states was determined. The activation energy required to cross to the dissociative ligand-field state was Ea = 3100 cm−1 for [Ru(bpz)2(tmam)]4+, while it was significantly smaller for [Ru(bpz)2(deeb)]2+ and [Ru(bpz)3]2+, i.e., Ea = 2600 and 2700 cm−1, respectively. This trend was supported by steady-state photolysis expe riments whe re [Ru(bpz)2(tmam)]4+ was much more stable toward ligand loss photochemistry, whereas the two other complexes were prone to it. The design of [Ru(bpz)2(tmam)]4+ took advantage of a strong photo-oxidation conferred by 2,2′-bipyrazine ligands with the ion-pairing ability of the dicationic tmam ligand, allowing therefore greater stability in the ground and excited states. While dynamic quenching was observed for all three complexes (kq = 4.9, 1.2, and 2.5 × 1010 M−1 s−1), [Ru(bpz)2(tmam)]4+* was also efficiently quenched through a static process. The appearance of the monoreduced [Ru(bpz)2(tmam)]3+ was indicative of electron-transfer quenching. Oxidized chloride species were however not observed spectroscopically. Evidence for the presence of Cl2 or Cl• was garnered from photolysis studies in the presence of a halogen trap, i.e., cis-dimethylbutadiene (DMBD) in deuterated acetone. 1H NMR analysis revealed that the resonances associated with DMBD decreased by a factor of 2, which was accompanied by the appearance of new resonances. The kq values, corrected for diffusion and excited-state encounter complex formation, allowed the free energy change for the excited-state electron-transfer reaction, ΔG° to be estimated through Marcus theory. Using the ΔG° and the [Ru(bpz)2(tmam)]4+*/3+ reduction potential, the formal oneelectron reduction potential of chloride in acetone was estimated to be E°(Cl•/−) = 1.87 V vs NHE.

3. HALIDE-TO-ACCEPTOR CHARGE TRANSFER 3.1. Outer-Sphere Charge Transfer

3.1.1. Ion-Paired Charge Transfer. Pyridinium salts with iodide provide a landmark example of ion-pair charge-transfer (IPCT) excited states, Figure 32. Pioneering work by Kosower investigated this charge transfer and utilized it to develop a sensitive method for determining solvent polarity.402−404 In initial reports of 1-methylpyridinium iodide, now known as Kosower’s salt, a new visible absorption band was observed that was not present in the separate absorption spectra of the parent ions.405 Prior studies of pyridinium reactivity led the AC

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shifted to high energy. As previously described in section 1.2.2, iodide also displays two absorption maxima in water that have been assigned as charge-transfer to H2O (2P1/2 and 2P3/2), and the energy difference between these transitions is also 21.7 kcal mol−1. The authors invoked these energy states of iodide to explain the two IPCT bands. Subsequent studies on pyridinium cations with aromatic donors refuted this explanation, as two charge-transfer bands were also observed in the absence of iodide.409 Additionally, Poziomek et al. found that functionalization of the pyridinium ring caused ΔTE to vary from 10 to 28 kcal mol−1.410 As substituents on the acceptor influenced ΔTE, the two band structure was not due to the iodide donor. The study instead showed that the absorption profile was successfully modeled by the presence of two energetically similar acceptor orbitals on the pyridinium cation. Work by Kosower on pyridinium radicals generated through pulse radiolysis also supported this conclusion.411 These radicals displayed a low energy absorption which correlated with the energy difference between the chargetransfer bands in the ground state. In IPCT, the ion-pair structure (contact, solvent separated, etc.) is expected to have a large impact on the electron-transfer properties. To probe the ion pairing in these complexes, Kosower et al. quantified how the substituents on pyridinium rings influenced the equilibrium constants.412 Two classes of interactions were proposed. In the so-called localized case, the iodide would associate with particular atoms on the pyridinium ring. Conversely, in the generalized case, the iodide would interact equally with all atoms in the aromatic ring. To distinguish between the two, the authors determined the equilibrium constants for a series of pyridinium complexes with bulky functional groups at different positions on the ring. The position of the sterically demanding groups did not impact the equilibrium constant, which favored the assignment of a “general” ion-pairing interaction. Substituents on the pyridinium ring, which only had a minor influence on the ion-pairing equilibrium, had a much more pronounced effect on the charge-transfer energy.413 Electrondonating hydrocarbon groups increased the electron density on the pyridinium ring, which blue-shifted the charge-transfer transition. Increasing the number of alkyl substituents on the ring was generally found to additively increase the energy of the charge-transfer transition. A series of 1-benzylpyridinium rings shown in Figure 32 corroborated that the energy of the charge-transfer was influenced by substituents further away from the electron acceptor.414 Electron donating and withdrawing functional groups on the benzyl substituent had the expected effect of raising or lowering the charge-transfer energy, respectively. A plot of the charge-transfer energies versus the Hammett constants resulted in a linear correlation with a ρ value of −1.70. The IPCT absorption of pyridinium iodides has been probed by flash photolysis of 1-ethyl-4-carbomethoxypyridinium iodide in benzene.415 Light excitation of the ion pair resulted in a new transient absorption spectrum that was consistent with the presence of a pyridinium radical and an iodine atom. The decay of these photoproducts was found to proceed through two separate pathways, as evidenced by biexponential decay kinetics with both a fast (109−1010 M−1 s−1) and a slow (107−108 M−1 s−1) second-order rate constant. In the fast pathway, the photogenerated iodine atom and pyridinium radical were proposed to react within the solvent cage. In the slow pathway, photogenerated iodine atoms that

Figure 32. Neutral and cationic acceptors for iodide charge transfer.

authors to hypothesize that a covalent bond was formed between the iodide and the aromatic ring. Substituents on the pyridinium were therefore expected to hinder this reactivity by significantly impacting the formation of this putative covalent species. However, upon preparation of a series of methyl substituted pyridinium iodides, the equilibrium constants were found to be very similar to that of the unsubstituted pyridinium, which excluded covalent bond formation as the cause of the visible absorption.406 In light of this result, the absorption feature was then assigned to a charge-transfer transition from the iodide ion to the pyridinium ring, a transition that is now often referred to as an outer-sphere charge transfer. Studies on the sensitivity of pyridinium iodide absorption bands to the solvation environment provided evidence for assignment as an IPCT transition.407 In these studies, a variety of solvents (methanol, ethanol, and acetone) with varying water content gave a gradient of polarities in which the absorption energy of 1-methyl-4-carbomethoxy pyridinium iodide was determined. Two absorption bands were observed that were found to be highly solvent dependent and displayed a pronounced negative solvatochromism, i.e., a blue-shift of the absorption band in polar solvents. This agreed with a stabilization of the charged, ground-state adduct in polar media that increased the energy required to excite the ion pair. This negative solvatochromism was observed both in neat solvents and as water was added to the organic solutions. In nonpolar solvents, iodide-to-pyridinium charge-transfer bands shift to a low enough energy (380 nm for 1-methylpyridinium iodide in CH2Cl2) that a second, high energy absorption feature can often be resolved.402,408 This high energy feature has also been shown to be sensitive to the solvent. In the case of 1-methylpyridinium iodide, the energy difference (ΔTE) between the two transitions remained constant (21.8 kcal mol−1) as EtOH was added to the chloroform solution, even though both absorption maxima AD

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modeled by the addition of authentic spectra of the methyl viologen radical and dichloride, Cl2•−. These products underwent back-electron transfer with a rate constant of 4 × 108 s−1.427 Further analysis on the picosecond time scale showed that MV•+,Cl• was generated when the singly ionpaired species (MV2+,Cl−) was present. However, a much faster back-electron transfer rate constant, 109 s−1, was determined in this case.431 The authors concluded that the energy barrier for back-electron transfer to the chlorine atom was less than that for the dichloride. When the photolysis was performed in alcohol, the solvent was found to react with the photogenerated radicals. In these photolysis studies, the authors noted that rate constants for back-electron transfer were much larger for bromide and iodide than for chloride. Because of the short lifetimes and the low quantum yields, they were not able to characterize the bromide and iodide salts by photolysis.431 Subsequent work detailing the flash photolysis of MV2+(I−)2 in water has characterized the ultrafast processes that occur upon charge transfer with iodide.430 Upon excitation of MV2+(I−)2, Figure 33, a transient absorption feature was

escaped the solvent cage were proposed to react and form molecular iodine. The iodine subsequently reacted with the pyridinium radical to generate ground-state products. In light of more recent work on iodide photoredox chemistry, it seems more likely that the acceptor was I2•− and not I2. The extensive work on pyridinium iodides has provided a comprehensive example of IPCT between aromatic cations and iodide. Other aromatic cations such as pyrylium,416 transstyrylpyridinium,417 and tropylium418,419 have also been shown to undergo IPCT with halides. For example, the chloride, bromide, and iodide salts of tropylium have been synthesized and studied. In CH2Cl2, the charge-transfer band shifted to higher energy as the halide became harder to oxidize.418 Subsequent work utilized tropylium and iodide as ligands for a molybdenum complexes, as shown in Figure 32.420 Significantly, the authors observed a low energy absorption (735 nm) that was assigned to a ligand-to-ligand charge transfer. Additionally, an iodide to tropylium charge-transfer was observed at 640 nm in a related molybdenum complex (Figure 32) when the tropylium was coordinated to the metal center and iodide was the counterion. The bis(arene)iron dication in Figure 32 was also reported to undergo IPCT transitions.421 The 2+ charge of the iron complex effectively facilitated ion-pairing with halides and other anions in solution. When the anion was an electron donor such as iodide, a new visible absorption was present. The energy of this absorption was dependent on the electron donating ability of the anion. The addition of inert salts to solutions of the charge-transfer complexes decreased the absorption intensity, consistent with competitive equilibria. 3.1.2. Halide-to-Viologen Charge Transfer. A large body of work has studied the electron acceptor properties of viologens, so they are reviewed separately in spite of their close relationship to the pyridinium salts discussed above. Viologens have proven to be robust electron acceptors, and are therefore used in a wide variety of energy storage applications.422−426 The halide salts of methyl viologen display a slight yellow color in solution even though the constituent cations and anions are themselves colorless. This visible absorption has been attributed to a halide-to-viologen charge transfer,427−430 and, consistent with Mulliken’s theory of charge transfer, the energy of the transition is linearly dependent on the ionization energy of the anions.429 Like the pyridinium complexes, the chargetransfer bands of methyl viologens are also sensitive to the polarity of the solvent and display a negative solvatochromism.428 The ion-pairing thermodynamics of methyl viologen with the halides have been investigated. The 2+ charge of this MV2+ salt allows for two equilibria to be established, as described by equations 36 and 37.429 MV2 + + X− F (MV2 +X−)

(36)

(MV2 +X−) + X− F (MV2 +(X−)2 )

(37)

In water, the first equilibrium was typically on the order 101 M−1 and was more favored than the second, 10−1 M−1, by at least an order of magnitude. These equilibria were suggested to have a large effect on the observed photoproducts after the charge transfer from the halide to the viologen. The flash photolysis of MV2+(Cl−)2 in water and alcohol solutions has been reported by Ebbesen and Ferraudi.427,431 The photolysis of this complex in water generated an absorption spectrum on the nanosecond time scale that was

Figure 33. (top) Time-resolved transient absorption spectra following 400 nm light excitation of a methylviologen−iodide complex. (bottom) Single wavelength absorption changes monitored at 610 nm in aqueous solution containing 0.03 M of methylviologen and different concentrations of iodide. Solid lines represent the best biexponential fits. Reproduced with permission from ref 430. Copyright 2001 Elsevier. AE

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observed at 610 nm which was indicative of the viologen radical, MV•+. On longer time scales, additional absorption features at 450 and 700 nm appeared that corresponded with the spectrum of diiodide (I2•−). The decay of the methyl viologen radical was well described with biexponential kinetics, yielding two lifetimes of approximately 1 and 20 ps. During iodide titration experiments, these two lifetimes remained constant, but their associated amplitudes changed. More specifically, the long-lived component’s contribution to the spectrum increased with the iodide concentration. The two transient decays were proposed to arise from electron transfer in the two ion-paired states given in eqs 36 and 37. The backelectron transfer reactions were characterized at a fixed MV2+ concentration with different concentrations of iodide. At lower iodide concentrations, the single ion pair MV2+,I− was expected to be the predominate species in solution. The fast 1 ps lifetime, which contributed most significantly to the transient spectrum at low iodide concentration, was therefore attributed to back-electron transfer from the MV•+ radical to the iodine atom. At higher iodide concentrations, the amount of the double ion pair MV2+,(I−)2 increased. At these iodide concentrations, the 20 ps lifetime contributed more to the transient spectrum and was therefore assigned to back-electron transfer from MV•+ to diiodide (I2•−). These electron-transfer reactions were treated theoretically with Marcus theory, and the calculated electron-transfer rates for the iodide species, as well as the chloride species studied by Ebbesen, were found to be in good agreement with the experimental values. This analysis led the authors to conclude that back electron transfer in both MV2+,I and MV2+,(I−)2 adducts was in the Marcus inverted kinetic region (−ΔG° > λ). 3.1.3. Charge Transfer in Supramolecular Assemblies. Methods for promoting outer-sphere charge transfer that do not involve ion pairing have also been investigated. As early as 1965, iodide was shown to form a charge-transfer complex with carbon tetrachloride through an outer-sphere charge-transfer interaction.432 In solution, these adducts formed with a one-toone stoichiometry, and IR spectroscopy suggested that the halide associated with a face of the carbon tetrahalide tetrahedron.433,434 However, in the solid state, crystal packing positioned the halides closest to the apexes of the carbon tetrachloride.435 The equilibrium constants for adduct formation decreased in magnitude when the halide was changed from iodide to bromide. The charge-transfer with iodide was also found to occur at a lower energy than bromide.434,436 More recent work by Kochi has exploited these charge-transfer interactions of carbon tetrabromide with halides in crystal engineering.437,438 The absorption bands for outer-sphere halide charge transfer contain information on fundamental electron-transfer parameters. Kochi et al. utilized this to quantify charge-transfer transitions from Cl−, Br−, and I− to carbon tetrabromide as the electron acceptor in CH2Cl2 solutions.438 The measured absorption energy Eop is directly related to the sum of the Gibbs free energy change and the reorganization energy, eq 38. Eop = ΔG° + λ

Figure 34. Spectral changes upon incremental addition of Pr4N+Br− into a 7.5 mM solution of CBr4 in dichloromethane. Concentration of Pr4N+Br−: 0 mM (a), 18 mM (b), 38 mM (c), 56 mM (d), 75 mM (e), 94 mM (f). The spectrum of Pr4N+Br− is shown for comparison purpose (g). Inset represents the Mulliken correlation between the energy of the charge-transfer band and the oxidation potential of the donor in the complexes of CBr4 acceptor with the indicated halide. Reproduced with permission from ref 438. Copyright 2003 American Chemical Society.

Hush theory, eq 39. Here the parameters of the halide charge transition, namely the full-width at half-max Δν1/2 (in eV), molar absorption coefficient at the absorbance maximum, εmax (M−1cm−1), and Eop (eV), enable direct calculation of HDA.439 This analysis also requires an estimation of the charge-transfer distance, r (Å). The shortest geometric distance between the halide and the bromine atoms in CBr4 obtained from X-ray crystallography measurements of the charge-transfer crystals was used. This likely represents an upper limit to the true charge-transfer distance. HDA =

0.0206 Eopεmax Δν1/2 r

(39)

From this analysis, HDA values in the 0.5−0.6 eV range were reported that represent remarkably strong coupling for a neutral electron acceptor. Surprisingly, they did not report the reorganization energies for the optical charge-transfer transitions. Also surprising is the lack of similar analysis in the vast halide literature. As rare examples, Farnum et al. reported Eop values consistent with outer-sphere charge transfer from iodide to a Ru bipyridine complex. Jarzeba et al. used a similar type analysis for Cl− and I− charge transfer to methyl viologen.430 Nevertheless, Mulliken−Hush analysis of outer-sphere halide charge-transfer absorption bands appears to be an area ripe for future research. Outer-sphere charge transfer has also been facilitated through halide interactions with electron poor aromatic rings and olefins.404,440−442 Davis et al. showed that a new visible absorption was present in solutions of 1,3,5-trinitrobenzene with iodide and in solutions of chloranil with bromide.440 When varying the solvent, the energy of the charge-transfer transition for both adducts increased with increasing solvent polarity, analogous to Kosower’s salt. In polar solvents, plots of the charge-transfer energy versus measures of the solvent polarity gave the expected slope of near unity. However, in nonpolar solvents, the slope deviated from unity. To explain this, the authors invoked ion pair formation between the

(38)

Figure 34 shows that the optical charge transfer shifted toward higher energy with the reduction potential of the halogen as would be expected based on formal reduction potentials, i.e., Cl− > Br− > I−. Note that Eop is equivalent to hvCT. The magnitude of the electronic coupling matrix element, HDA, was calculated using semiclassical Mulliken− AF

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of key importance.448 Figure 36 illustrates possible electronic transitions within an octahedral symmetry, MX6n− complex.

anionic donor and its accompanying cation. The presence of an ion pair was proposed to stabilize the ground state, which would increase the energy of the charge-transfer transition. To test this hypothesis, the authors investigated the relationship between cation size and the energy of the charge-transfer transition.442 The use of alkali metal and various tetraalkylammonium iodides was found to have a large influence on the energy of the charge-transfer transition in nonpolar solvents like t-butanol, but no effect was observed in ethanol. In nonpolar solvents, increasing the steric bulk of the ammonium cations positioned the positive charge farther from the supramolecular assembly, which caused the chargetransfer energy to decrease. In polar solvents, ion pair formation was disfavored, and so the size of the cation did not have an effect. Kochi et al. recognized that the use of electron deficient aromatics and olefins provided a useful means of recognizing halides in solution and therefore developed a number of electron poor organic compounds that displayed anion−π interactions.443 These assemblies all facilitated charge-transfer from an associated halide to the electron-poor π-system, as observed through the growth of visible absorption features upon halide addition in CH3CN. The charge-transfer transition energy was found to increase from iodide to bromide to chloride, as predicted by Mulliken based on the formal halide reduction potential.

Figure 36. Partial MO diagram in an octahedral symmetry showing possible LMCT transitions in MX6n− complex.

The transitions from σ to t2g (green) and π to t2g (red) are of less significance in an octahedral complex as these transitions are weak due to low orbital overlap.448 The other two transitions are strong and ultimately of interest for LMCT purposes as long as electron transfer results in localization of the oxidizing equivalent on the halide.16 Unlike the outer-sphere charge transfer discussed in the previous section, there is significantly more mixing of the donor and acceptor wave functions in the ground and excited states of metal halide complexes. As such, the use of formal oxidation states is of less use as indicated by the δ+ and δ− in Figure 35. LMCT excitation is better described as chargetransfer chemistry as opposed to electron transfer. Nevertheless, oxidized halide species are poor ligands themselves, and as is described below, LMCT excitation often results in reductive elimination, albeit with low quantum yields.

3.2. Halide-to-Metal Charge Transfer

One approach for driving excited-state halide oxidation is through the use of ligand-to-metal charge-transfer (LMCT) excited states. In transition metal halide complexes where this often represents the sole and lowest lying excited state, every absorbed photon results in direct oxidation of the halide species and substantially reduces energy losses. With a given metal, the expected LMCT transition energy will increase in the sequence I− < Br− < Cl− < F− as the halides become more difficult to oxidize.444,445 Along the same line of reasoning, higher formal oxidation state metals are better donors and will results in a decrease in the LMCT energy.446 In practice, halide-to-metal charge-transfer transitions sometimes occur in the ultraviolet region with other overlapping electronic transition. As such, they are often difficult to access for meaningful photochemistry. In addition, the yield of oxidized halide is often low, behavior that is attributed to rapid back-electron transfer. For some metal halide complexes, LMCT excitation results in reductive photoelimination of halogens. Oxidative addition and reductive elimination represent important mechanistic steps in many catalytic cycles and study of LMCT excited states can provide insights into the detailed reaction chemistry. The schematic LMCT transitions shown in Figure 35 implies the localization of charge to generate “oxidized and reduced species”.447 The intricacies of the frontier orbitals are

3.3. Halide Oxidation Through X2 Elimination

3.3.1. Dirhodium HX Splitting Catalysts. Halide oxidation has been accomplished by the Nocera group in their pioneering work to develop hydrohalic acid (HX) splitting catalysts for solar fuel generation. Excellent reviews are recommended for a detailed description of their approach.12,449 Bimetallic mixed-valence complexes with a metal−metal bond, Figure 37, were characterized as competent photocatalysts for the two-electron process required to reduce protons to H2 and oxidize halides to X2 or X•, which then reacted with a halogen trap.450 Early catalyst (denoted Mn− Mn) behavior was consistent with the catalytic cycle shown in Figure 38.451 Oxidative addition of HX to both metal centers in A gave the dihydride−dihalide Mn+2−Mn+2. Photoinduced elimination of H2 yielded the catalyst as Mn+1−Mn+1, which then underwent halide transfer to give the disproportionation product Mn−Mn+2. Photoexcitation yielded LMCT excited states that facilitated halogen elimination in the presence of sacrificial olefin reagents, or halogen traps, which regenerated the catalyst (Mn−Mn), and closed the catalytic cycle. The first example of the homogeneous catalytic conversion of HX (X = Cl or Br) to H2 and X2 was achieved with the dirhodium catalyst A, Figure 37 and 38.450 The success of this

Figure 35. A general depiction of a LMCT transition, yielding an “oxidized” ligand and a “reduced” metal center, in an octahedral metal complex upon photoexcitation. AG

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Figure 37. Dirhodium HX splitting catalysts.

Figure 39. Photodriven halogen elimination from a mixed-valence RhIII−RhI complex (B) and valence-symmetric RhII−RhII complex (C) result in similar quantum yields. Transient absorption experiments suggested that both complexes go through a common intermediate during catalysis.

transient absorption spectroscopy provided evidence for a common intermediate during catalysis with the two complexes, as was demonstrated by a common absorption maximum at 440 nm, Figure 40. This intermediate was proposed to be the halide-bridged species D, and the authors hypothesized that this intermediate facilitated halogen elimination in both cases, which resulted in similar quantum yields. Subsequently, the authors prepared two Cl-bridged dirhodium complexes through variation of the L-type ligand, Figure 37 (E and F).

Figure 38. Proposed mechanism for HX splitting with complex A.

catalysis was due in part to the ability of the catalyst to support mixed-valency in the metal centers. In these complexes, the three bridging dfpma ligands (dfmpa = bis[difluorophosphino]methylamine) stabilized the rhodium in the Rh0, RhI, and RhII oxidation states, as the π-acceptor phosphine orbitals accept charge density from the nitrogen or the metal. This aided in the disproportionation reaction which yielded an intermediate that underwent reductive elimination upon irradiation with UV light, Figure 38. This elimination reaction, which is of interest in the context of this review, has subsequently been studied in significant detail as the low monochromatic quantum yield for the halogen elimination step (0.6%) was found to limit the overall efficiency. Subsequent research enabled an increased reaction efficiency by introducing modifications to the catalyst to improve the photodriven halide oxidation. In recent work, catalyst design has sought to incorporate structural features which mimic intermediates in the cycle, thereby minimizing reorganizational energies.452−454 Studies of complexes B and C provided a useful step forward in this approach.453 The Nocera group recognized that photocatalysis with the complexes resulted in similar quantum yield for Cl2 elimination despite the initial difference between the mixed-valence and valence-symmetric oxidation states of the complexes, Figure 39. In this case,

Figure 40. Transient absorption spectra obtained by laser flash photolysis of complexes C (black) and B (red) in Figure 37. Spectra are normalized to the absorbance at 440 nm. Normalized absorption spectra of C (dashed black) and B (dashed red) are also shown. Reproduced with permission from ref 453. Copyright 2013 Royal Society of Chemistry. AH

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ligand) was prepared and photolyzed in the solid state, causing CO release and the formation of the halide bridged intermediate D. The spectral features of D were consistent with the intermediate observed in HX splitting with B and C, which suggested that the chloride-bridged structure was indeed on the catalytic cycle. Subsequently, a series of dirhodium catalysts were prepared that featured a single L-type ligand and a bridging chloride or bromide ligand into the ground state, Figure 37 (E−G).454 The synthesis was achieved through the use of aryl isocyanides as the single L-type ligand instead of the previously used alkyl isocyanides. These catalysts simplified the catalytic cycle by removing the photoinitiated ligand loss step. Hydrogen production with these complexes was found to be 2 to 4 times more efficient than that measured for the previous Rh2L2 examples. Unexpectedly, the photoresting state of the catalyst, as determined through UV−visible absorption spectroscopy, was found to be M instead of E−G. This showed that H2 evolution, and not halide elimination, was the rate-limiting step for these complexes. This was a significant departure from the previous generation of catalysts, where X2 elimination always limited the efficiency of catalysis. 3.3.2. Bimetallic Catalysts for X2 Elimination. The elimination of X2 also proceeded through the use of bimetallic complexes where one metal mediates H2 formation, and a second, more oxidizing metal such as platinum, gold, or iridium facilitates the difficult M−X bond activation required for halogen elimination.455−457 This strategy was successfully employed in a PtIII(d7)−AuII(d9) complex where the quantum yield for the photoelimination of chloride with visible light in the presence of a halogen trap was 5.7%.455 Subsequent work on this and other bimetallic complexes, reviewed in reference 12, Figure 42 (O−Q), resulted in even higher quantum yields for chlorine photoelimination with values ranging from 10% to 38%. Complexes containing Pt(III) typically gave the highest quantum yields. Catalysts based on this design, but with main group elements such as tellurium (Te) or antimony (Sb), Figure 42 (R−S), instead of a second transition metal center, were also shown to facilitate the two-electron reactivity at the transition metal center with comparable quantum yields (4− 14%) for X2 formation.458,459 The mechanism of chloride photoelimination in main group transition metal platinum complexes was recently investigated.460 Pulsed-light excitation revealed the presence of reaction intermediates in a family of Pt(III) complexes. These intermediates did not correspond to the reductive elimination products, and the authors attributed them to the singly reduced Pt complexes caused by the release of a halogen atom X•. Photocrystallography of Q supported this assignment, as a single Pt−Cl bond was elongated upon irradiation, which suggested stepwise cleavage of Pt−Cl bonds instead of a concerted reductive elimination reaction. The halogen elimination quantum yield was shown to decrease as the strength of the Pt−Cl bond increased. Although this correlation was expected, it showed that the development of both highly endothermic and efficient energy storage reactions are at odds. Control of the reaction mechanism loosened this correlation, as was shown with studies of the fac−N and mer−N isomers of the RhII−PtIII complexes. Both complexes displayed a similar thermodynamic barrier for halogen elimination, but the reaction quantum yield was four times greater for fac−N than mer−N, Figure 43. The higher efficiency for the fac−N

These complexes displayed a ground-state absorbance at 450 nm, which closely aligned with that observed transiently and assigned to intermediate D. This supported the structural assignment of a chloride-bridged intermediate that facilitated halogen photoelimination. To further characterize the proposed intermediate, photocrystallographic studies were performed.452 To ensure, however, that the solid-state and solution reactions proceeded through the same mechanism, transient absorption experiments on solid films of B and C were first completed. These solid-state experiments showed absorption features that matched the solution studies, indicative of a common reaction mechanism in both phases. Crystallographic studies in the dark and under light illumination provided structural information on the light driven structural changes as shown in Figure 41, with the atomic positions mapped onto each other.

Figure 41. Thermal ellipsoid plots of photocrystallography results with photoinduced structure (solid) superimposed on dark structures (faded). (top) Rh2[I,III] (B) plot: Rh1−Rh2−Cl3 91.05(5)° (dark); 83.2(2)° (photoinduced). (bottom) Rh2[II,II] (C) plot: Rh1−Rh2− Cl3 91.15(5)° (dark), 78(2)° (photoinduced). Reproduced with permission from ref 452. Copyright 2014 American Chemical Society.

Notably, light-excitation caused the Rh(1)−Rh(2)−Cl(3) bond angle to contract from 91.05 to 83.2° for complex B and from 91.15 to 78° for complex C, which allowed the bound Cl to approach a bridging geometry between the two Rh centers, Figure 41. The chloride-bridged species was not directly observed in these photocrystallographic experiments. The loss of a L-type ligand was expected to precede formation of this intermediate. Presumably, the bulky AdCN (AdCN = 1cyanoadamantane) ligand was unable to dissociate in the solid state. Therefore, a series of analogous complexes with volatile CO ligands was synthesized, Figure 37 (H−K). In this series, irradiation of I resulted in the loss of a CO ligand and formation of μ-Cl species J, while prolonged irradiation led to X2 elimination and formation of K. Using this strategy, intermediate D was successfully synthesized and characterized. To do so, dirhodium complex L (with a single CO and AdCN AI

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Figure 42. Bimetallic complexes for halogen elimination.

has launched investigations into complexes where a main group element is the sole mediator for halogen elimination. Many main group elements in fact seem ideal for this application, as they preferably perform two-electron redox chemistry, as the one-electron processes are more energetically demanding.461−464 Examples of halogen elimination remain uncommon, but examples of tellurium465−467 and antimony468,469 have been reported. An early example was T− X2, shown in Figure 44, that was reported by Seferos et al. and

Figure 43. Comparative studies of fac vs mer RhII−PtIII complexes provided control of the X2 elimination mechanism. In the fac−N complex, concerted X2 elimination enhanced the quantum yield by a factor of 4 despite the similar thermodynamic barriers for the two complexes.

isomer was attributed to the detailed elimination mechanism. While probing this mechanism, a transient intermediate with absorption features similar to that of the parent complex was formed within the duration of the laser pulse (8 ns). This intermediate then converted smoothly to the photoproduct with a rate constant of 3.6 × 104 s−1. The similarity of the absorption spectra of fac−N and the intermediate suggested that the two species were structurally related. The authors concluded that upon light excitation, the adjacent equatorial chloride ligands reacted to give the σ-Cl2 species, which then eliminated Cl2. The quantum yield of halogen elimination from fac−N (23.4%) as compared to mer−N (6.6%) showed that control of reaction mechanism could be used to mitigate thermodynamically uphill energy storage reactions. 3.3.3. Main Group HX Splitting Catalysts. The work on main group participation in transition metal based HX splitting

Figure 44. Tellurophenes as reported by Seferos et al.466

utilized a tellurophene functional group, where the tellurium center had accessible II and IV oxidation states.466 In this work, illumination of the dihalide−Te(IV) complex reductively eliminated a halogen atom, but variation of the halide (X = Cl, Br) and halogen trap concentration only gave a maximum efficiency of 0.19%. DFT calculation suggested that the nature AJ

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Figure 45. Photolysis of U−Br2 with 447.5 nm light over time monitored by (a) UV−visible and (b) 1H NMR spectroscopy. Adapted with permission from ref 467. Copyright 2015 Royal Society of Chemistry under CC BY 3.0.

Figure 46. Expected trap products based on a stepwise (left) or concerted (right) mechanism of halogen elimination.

Figure 47. Mechanism of bromine photoelimination from U−Br2 (top) and proposed photochemical and photophysical mechanisms for U−Br2 upon excitation at 450 nm and U upon excitation and 355 nm.

with a concomitant rise in those associated with U, therefore confirming the light-induced halogen elimination. Mechanistic studies on a series of tellurophenes shed light on the bromine elimination pathway.465 A concerted Br2 and stepwise elimination mechanisms were envisioned, and to distinguish between them, the halogen trap products were characterized. If the concerted reductive elimination were operative and free Br2 was released, the halogen trap, 2,3dimethyl-2-butene (DMB), was expected to undergo anti addition with the halogen to give 2,3-dibromo-2,3-dimethyl butane (DMB-Br2). In the stepwise mechanism, a bromine radical would be released that would extract a hydrogen atom from DMB to give HBr and an allylic radical that subsequently reacted with a second bromine atom to give 1-bromo-2,3dimethyl-2-butene (DMB-Br) (Figure 46). Characterization of the trap products showed that DMB-Br was predominately formed, consistent with the stepwise elimination producing a bromine atom intermediate. Further evidence for bromine

of the HOMO−LUMO transition was responsible for the low quantum yield, as the excited state did not have significant Te−X antibonding character but was instead delocalized across the entire molecule. In subsequent work, Seferos et al. designed U−X2 compounds, Figure 44, in which the lowest energy excited state was localized on the tellurophene and was comprised of significant Te−X antibonding character.467 Irradiation of the lowest energy absorption band resulted in rapid halogen elimination, as shown in Figure 45. A halogen elimination quantum yield of nearly 17% (when X = Br) at trap concentrations of 5 M was determined, which was 85 times greater than the previously reported tellurophene.466 This quantum yield compared favorably with many of the leading transition metal complexes.456,457 Similar observations were made by 1H NMR when U−Br2 was irradiated in the presence of 0.9 M of DMBD halogen trap, Figure 45. Over 20 s of illumination, the resonances associated with U−Br2 decreased AK

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functional theory calculations indicated that the lowest energy excited states of this complex were antibonding with respect to the Pd−Cl and Sb−Cl bonds, and irradiation of this transition indeed lead to the photoelimination of Cl2 with a quantum yield of 0.58% when halogen trap concentrations of 2.0 M were used. Nocera et al. have provided a subsequent example of halogen elimination mediated solely by antimony.469 The Sb III and V oxidation states appeared well suited for halogen oxidative addition and reductive elimination, and the authors were able to exploit this when the main group Sb metal was complexed by a corrole ligand. Crystallography, 1H NMR, and UV−vis absorption spectroscopy revealed that changes in the antimony oxidation state were accompanied by significant structural changes. In the III oxidation state, the antimony resided above the plane of the ligand, but oxidative addition of either PhICl2 or Br2 resulted in an octahedral complex where the SbV resided in the ligand plane. Qualitative conversion from SbV−Br and SbV−Cl to SbIII was achieved through steady-state photolysis as is shown in Figure 49. In the various excitation wavelengths

atom intermediates came from experiments with excess bromide (Br−) that efficiently reacted with bromine atoms to give Br2•−. Transient absorption studies of the brominated tellurium complex U−Br2 using pulsed 450 nm excitation revealed the presence of short-lived (20−70 ns) intermediates before the formation of dehalogenated products.465 However, with pulsed 355 nm excitation, U−Br2 displayed a different transient absorption spectrum that was equivalent to that of U, which has been attributed to the triplet excited state of that complex, Figure 47. In light of the identical transient absorption spectra, the authors concluded that excitation of U−Br2 at 355 nm caused halogen elimination and photoproduct excitation to the triplet excited state within the duration of the laser pulse (5−8 ns). Therefore, the reductive elimination of bromine atoms was so fast that it could not be quantified under these conditions. The transient observed with 450 nm excitation possessed a lifetime of >10 ns and was thus inconsistent with an intermediate on the halogen elimination pathway. Furthermore, the absence of luminescence suggested that the transient species was not the singlet excited state. Therefore, the authors concluded that the intermediate was the U−Br2 triplet excitedstate. This was supported by experiments wherein U−Br2 produced singlet oxygen (1O2) using pulsed 450 nm excitation. These results suggested that the excited tellurium bromide complexes decayed through a variety of competitive pathways. In one pathway, the singlet excited state led to the desired halogen elimination. In the second pathway, intersystem crossing to the triplet excited state occurred, which either produced singlet oxygen in aerobic conditions or relaxed to the ground state. A recent report by Gabbaı̈ et al. has shown that antimony can also participate in the reductive elimination of halogen atoms.468 In the antimony−palladium complex V shown in Figure 48, reaction with chlorine did not result in oxidative addition at the palladium center, as was expected from related work on platinum−antimony complexes.459 Instead, addition occurred across the palladium−antimony bond, formally oxidizing both atoms by one, which yielded W. Density

Figure 49. Steady-state photolysis (λexc > 305 nm) of SbV−Corrole in degassed THF. The initial spectrum (green) converts to that of SbIII− Corrole (brown) over the course of 2 min. Adapted with permission from ref 469. Copyright 2018 American Chemical Society.

tested, SbV−Br was found to give higher reductive elimination quantum yields than SbV−Cl. For SbV−Br, a maximum ϕ of 0.88% was achieved with 435 nm excitation, while SbV−Cl had a maximum ϕ of 0.17% with 315 nm excitation. Trap-free halogen elimination from SbV−Br and SbV−Cl could not be achieved, as the halogen atoms reacted with the corrole ligand. The reported quantum yields were similar to many of the tellurophene complexes studied by Seferos et al.465−467 3.3.4. Monomeric and Bimetallic Gold and Platinum HX Splitting Catalysts. To more extensively focus on the halide oxidation step of HX splitting catalysis, transition metal complexes have been developed that only perform the halide oxidation reaction. Recently, advances have been achieved with monomeric transition metal complexes that successfully catalyze this reaction. Nocera et al. have reported a series of mono and bimetallic AuIII−halide complexes, X−Z in Figure 50, in which the metal center supported the complete twoelectron chemistry required for the reductive elimination of X2.383 In this work, X2 formation was driven through photochemical elimination from the gold center. It was found that two metal centers were not necessary for this to occur, and in fact, monomeric gold(III) complexes did perform the photochemical reactions. Although the mechanism was not

Figure 48. Antimony−palladium complexes as reported by Gabbaı̈ et al.468 AL

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Chemical Reviews

Review

products (DMB-Br) from Br• and HBr addition were formed in higher yields when larger concentrations of the trap were present. The authors concluded that the reactive excited states could react directly with the olefin trap, thereby changing the mechanism of bromine photoelimination. In the analogous chloride complexes, no evidence for Cl2 elimination was obtained from characterization of the trap products or from photolysis in the solid state. Control experiments with only Cl2 and 2-hexene in solution revealed that anti addition is primarily observed under these conditions. However, the photolysis of the platinum complexes in the presence of this trap yielded a 50−50 mixture of the syn and anti addition products, which was inconsistent with Cl2 addition. Additionally, the steric size of the platinum complexes was found to dictate photochlorination products in cases where both 1-hexene and trans-3-hexene were present, which would not be expected if Cl2 was the reactive species. Cl2 is known to selectively chlorinate the electron rich but sterically encumbered trans-3-hexene as opposed to 1-hexene in a 12:1 ratio. For less sterically hindered complexes, the photolysis primarily resulted in chlorination of trans-3-hexene, although with a reduced selectivity of 3:1. However, platinum complexes were found to preferentially chlorinate the sterically accessible 1-hexene. In all these cases, the trap products were consistent with reaction of the platinum excited state with added alkenes, and not with the formation of Cl2. 3.3.5. Main Group LMCT Halide Oxidation. Utilizing LMCT excited states, the Vogler group was able to drive halide oxidation in several main group metal halide species including tin(II), lead(II), antimony(III), and bismuth(III) bromo complexes.472 Previous work from their group had characterized the absorption and emission spectra of the chloro complexes and reported no appreciable evidence of photochemistry.473−479 The study was extended to the bromo complexes to ascertain the full role of the ligand on the electronic spectra. Absorption spectra in the visible and UV regions were assigned to transitions between the s and p orbitals of the metal, and a LMCT. The qualitative MO diagram shows the two major orbital transitions of interest, Figure 52. The high-energy absorption was assigned to the LMCT transition involving charge transfer from the σ orbitals of the halide to the p orbitals of the metal. This transition was measured at wavelengths less than 250 nm for all s2 metals in their study and as such is of limited practical interest due to the required high energy excitation.472−479 Low energy bands were also observed that were assigned to metal-localized sp excited states. Figure 52 compares bromide and chloride LMCT transitions to the same metal, [MX6]n−. Note that the σ orbitals of bromide are closer in energy to the metal orbitals than are chloride’s, resulting in greater mixing. For the chloride compounds, two distinct absorption bands are expected, a high energy LMCT and a sp transition. For the corresponding bromo compounds, the increased mixing provides lower energy LMCT/(sp) and sp/(LMCT) transitions than the corresponding chloro complex. These assignments were reinforced by redshifts observed in the lowest energy absorption band as additional halide ions were coordinated to the metal complexes. Ultimately, the mixing of the LMCT and sp transitions leads to the photoredox chemistry of the metal halide species to oxidized halide products and reduced metals. Vogler et al. reported a 30% quantum yield for formation of oxidized bromide products with 366 nm (sp/

Figure 50. Mono- and bimetallic gold complexes as reported by Nocera et al.383

definitively established, halogen trapping by olefin experiments gave products that were predominately consistent with X2 and not halogen radical formation, suggesting a concerted X2 reductive elimination mechanism. The series of Pt(PEt3)2(R)(Br)3 complexes shown in Figure 51 was reported by Sharp et al. for the reductive elimination of

Figure 51. Platinum complexes as reported by Sharp et al.470,471

Br2.470,471 In the R = CF3Ph case, a Br2 elimination quantum yield of 82% was achieved, while the remainder of the series gave quantum yields ranging from