Mechanisms and Kinetics of Proton-Coupled Electron-Transfer Oxidation of Phenols by Na Song A disertation submited to the Graduate Faculty of Auburn University in partial fulfilment of the requirements for the Degree of Doctor of Philosophy Auburn, Alabama December 12, 2011 Keywords: Proton-coupled electron-transfer, phenol oxidation, outer-sphere, kinetics and mechanisms, transition metal complexes, stopped-flow. Copyright 2011 by Na Song Approved by David Stanbury, Chair, Profesor of Chemistry and Biochemistry Holly Elis, Asociate Profesor of Chemistry and Biochemistry Christian Goldsmith, Asistant Profesor of Chemistry and Biochemistry German Mils, Asociate Profesor of Chemistry and Biochemistry ii Abstract The kinetics of the aqueous oxidation of phenol by a deficiency of [IrCl 6 ] 2? has been investigated. The reaction initialy produces [IrCl 6 ] 3? and phenoxyl radicals. The inhibition caused by [IrCl 6 ] 3? can be prevented by use of dibromonitroso- benzenesulfonate (DBNBS) as a phenoxyl radical scavenger. The phenoxyl radicals primarily couple to form 4,4'-biphenol, 2,2'-biphenol, 2,4'-biphenol, and 4- phenoxyphenol. Further oxidation of these coupling products leads to a rather complex mixture of final products. The rate laws for the oxidation of the four coupling products by [IrCl 6 ] 2? have the same form as those for the oxidation of phenol itself: d[Ir V t = (kArOH+r-K/[ + ]) 1a ArOH]to[I V . Values for k ArOH and k ArO? have been determined at 25 ?C and are asigned to H 2 O-CPET (water asociated concerted proton coupled electron transfer) and electron-transfer mechanisms respectively. Kinetic simulations of a combined mechanism that includes the oxidation of phenol as wel as the subsequent reactions show that the degree of overoxidation is rather limited at high pH but quite extensive at low pH. This pH-dependent overoxidation leads to a pH-dependent stoichiometric factor in the rate law for oxidation of phenol, and causes some minor deviations in the rate law for oxidation of phenol. Empiricaly, these minor deviations can be acommodated by introduction of a third term in the rate law that includes a "pH- dependent rate constant", but this approach masks the mechanistic origins of the efect. iii One-electron oxidation of alkyl- and alkoxy-substituted phenols (2-methylphenol, 2,6-dimethylphenol, 2,4,6-trimethylphenol, 4-tert-butylphenol and 4-methoxyphenol) has been studied. pH-dependent stoichiometric factors corresponding to overoxidation are found with al substituted phenols except for 2,4,6-trimethylphenol and 4- methoxyphenol. In the 2,4,6-trimethylphenol reaction, the identification of product, 4- hydroxymethyl-2,6-dimethylphenol, rules out the overoxidation steps and there is no need to include an ?overoxidation pH-dependent? rate constant. The solvent H/D KIE?s for the phenols pathway provide further evidence for a H 2 O-CPET mechanism of oxidation of phenols by [IrCl 6 ] 2? . Overoxidation is also observed in the reaction betwen N-acetyl-L-tyrosinamide (a protected tyrosine derivative) and [IrCl 6 ] 2? . Fiting the data to a two-term rate law yields second-order rate constants of k ArOH = 5.4 ? 0.6 M ?1 s ?1 and k ArO ? = (4.5 ? 0.3) ! 10 7 M ?1 s ?1 . Analysis of the kinetic data of the oxidation of phenol by [Os(phen) 3 ] 3+ yields the rate law: ? d[O I ] t =2kdim [s I ] 2 ArOHto 2 I ( KArOH a+[] ) 2 where the reaction rate is second-order in both [Os(phen) 3 ] 3+ and phenol. K ArOH = 1.1 ! 10 ?10 M and K ArO ? = 7.0 are obtained from thermodynamics. k ArO ? is calculated to be 2.1 ! 10 9 M ?1 s ?1 acording to Marcus theory and this value is also supported by kinetic simulations. iv Acknowledgments The first person I would like to thank is my advisor, Dr. David Stanbury, for his kindnes and patience to guide my PhD journey, for his generosity to teach me not only the knowledge but also the atitude to be a scientist, and for al his help to navigate my academic way. I am very grateful to my commite members, Dr. German Mils, Dr. Christian Goldsmith and Dr. Holly Elis, for their advice with respect to my research, career and writing. I also want to thank Dr. Thomas Albrecht-Schmit who helped me at the beginning of this work. Many thanks to Dr. John Gorden for helping me to synthesize a complex, Dr. Douglas Goodwin for alowing me to use his stopped-flow instrument, Dr. Yonnie Wu for collecting mas spectroscopic data and Dr. Michael Meadows for helping me acquire NMR data. I also would like to thank my previous and current lab colleagues: Dr. Xiaoguang Wang who helped me a lot when I started this research, Dr. Thanasekaran Pounraj who gave me some useful advice and Mr. Nootan Bhatarai who discusses the research with me. I thank al my friends and colleagues at Auburn for their supports and encouragement. Finaly, I would like to thank my family, without their love and support I never could have gotten this far. I especialy thank my husband, Jie Wu, for always being on my side and helping me keep my faith. v Table of Contents Abstract ............................................................................................................................... ii! Acknowledgments .............................................................................................................. iv! List of Tables ..................................................................................................................... ix! List of Schemes ................................................................................................................ xii! List of Figures .................................................................................................................. xiv! Chapter 1 Literature Review ............................................................................................... 1! 1.1 Outer-Sphere Electron Transfer ................................................................... 1! 1.2 Proton-Coupled Electron-Transfer (PCET) ................................................. 7! 1.3 Oxidation of Phenols .................................................................................... 9! Chapter 2 Proton-Coupled Electron-Transfer Oxidation of Phenol by Hexachloroiridate(IV) ...................................................................................... 14! 2.1 Introduction ................................................................................................ 14! 2.2 Experimental Section ................................................................................. 15! 2.2.1 Reagents and Solutions .......................................................... 15! 2.2.2 Methods .................................................................................. 18! 2.3 Results ........................................................................................................ 22! 2.3.1 Oxidation of Phenol ............................................................... 22! 2.3.1.1 Hexachloroiridate(III) Efects ............................ 22! 2.3.1.2 General Base Catalytic Test. .............................. 25! vi 2.3.1.3 Spin Trapping Efect .......................................... 26! 2.3.1.4 Phenol Dependence ............................................ 29! 2.3.1.5 Dependence on p[H + ] ......................................... 30! 2.3.1.6 Inclusion of the k? Term ..................................... 32! 2.3.1.7 Kinetic Isotope Efect (KIE) .............................. 35! 2.3.2 Overoxidation of Phenol ........................................................ 36! 2.3.2.1 Coupling Products .............................................. 36! 2.3.2.2 pK a values of 4,4?-Biphenol and 2,4?-Biphenol . 39! 2.3.2.3 Overoxidation of Phenol ..................................... 41! 2.4 Discussion and Conclusion ........................................................................ 56! Chapter 3 Oxidation of alkyl- and alkoxy-substituted phenols by Hexachloroiridate(IV)75! 3.1 Introduction ................................................................................................ 75! 3.2 Experimental Section ................................................................................. 76! 3.2.1 Reagents and Solutions .......................................................... 76! 3.2.2 Methods .................................................................................. 78! 3.3 Results ........................................................................................................ 78! 3.3.1 The Oxidation of 2-Methylphenol .......................................... 78! 3.3.2 The Oxidation of 2,6-Dimethylphenol ................................... 88! 3.3.3 The Oxidation of 2,4,6-Trimethylphenol (TMP) ................... 95! 3.3.4 The Oxidation of 4-Methoxyphenol ..................................... 102! 3.3.5 The oxidation of 4-tert-Butylphenol .................................... 107! 3.4 Discussion and Conclusion ...................................................................... 111! 3.4.1 Mechanism ........................................................................... 111! vii 3.4.2 Mechanism of k ArO ? Term .................................................... 113! 3.4.3 Mechanism of k ArOH Term .................................................... 118! Chapter 4 Oxidation of Ac-Y-NH2 by Hexachloroiridate(IV) ....................................... 132! 4.1 Introduction .............................................................................................. 132! 4.2 Experimental Section ............................................................................... 132! 4.2.1 Reagents and Solutions ........................................................ 132! 4.2.2 Methods ................................................................................ 133! 4.3 Results ...................................................................................................... 134! 4.3.1 Hexachloroiridate(III) Inhibition .......................................... 134! 4.3.2 Spin Trapping Efect ............................................................ 137! 4.3.3 Ac-Y-NH 2 Dependence ........................................................ 138! 4.3.4 Dependence on p[H + ] ........................................................... 138! 4.3.5 Stoichiometry and Overoxidation ........................................ 142! 4.3.6 Control Experiments ............................................................ 145! 4.4 Discussion and Conclusion ...................................................................... 145! Chapter 5 Oxidation of phenol by tris-(1,10-phenanthroline)osmium(III) ..................... 149! 5.1 Introduction .............................................................................................. 149! 5.2 Experimental Section ............................................................................... 149! 5.2.1 Reagents and Solutions ........................................................ 149! 5.2.2 Methods ................................................................................ 151! 5.3 Results ...................................................................................................... 152! 5.3.1 Characterization of the Osmium Complexes ....................... 152! 5.3.2 Kinetics ................................................................................. 153! vii 5.4 Discussion and Conclusion ...................................................................... 167! References ....................................................................................................................... 173! Appendix A Experimental Details in Chapters 2?5 ........................................................ 180! ix List of Tables Table 2-1. Kinetic Dependence on Concentration of Buffer ............................................ 25! Table 2-2. Kinetic Data for the Reaction of Phenol with Ir IV in the Presence of PBN, DMPO, POBN, MNP, DMNBS and DBNBS ................................................. 28! Table 2-3. Kinetic Data for the Reaction of Phenol with Ir IV in the Presence of DBNBS 29! Table 2-4. Kinetic Dependence of Phenol Oxidation on DBNBS at High p[H + ] ............. 30! Table 2-5. Kinetics Data for the Reactions of Ir IV with the Phenol Coupling Products ... 43! Table 2-6. Comparison of the Experimental Data of 4,4?- and 2,4?-Biphenol Oxidation at Diferent p[H + ] ................................................................................................ 47! Table 2-7. The Mechanism of Phenol Reaction and the Simulation Model ..................... 64! Table 2-8. The Overoxidation Rate Constants at Diferent p[H + ] .................................... 65! Table 2-9. Comparison of the Experimental Data with Overoxidation and No Overoxidation Simulation Results of the Phenol Reaction at Diferent p[H + ] in Absence of DBNBS ......................................................................................... 67! Table 2-10. Simulation Product Concentrations of the Phenol Reaction at Diferent p[H + ] Without DBNBS ............................................................................................ 69! Table 2-11. Comparison of the Experimental Data with Overoxidation Simulation Results of the Phenol Reaction in the Presence of Ir III or Various DBNBS Concentration ................................................................................................. 71! Table 3-1. Kinetic Data for the Reaction Betwen Cresol and Ir IV with Added Ir III in the Presence of Spin Trapping Agents PBN, DMPO, POBN, MNP, DMNBS and DBNBS ............................................................................................................ 82! Table 3-2. Kinetic Efect of Cu 2+ and Dipic for Xylenol Oxidation ................................ 91! Table 3-3. Kinetic Data for the Reaction of Phenols with Ir IV ....................................... 114! x Table 3-4. Calculated Data for Phenols and the Reactions with Ir IV .............................. 115! Table 3-5. Thermodynamic Parameters for CPET and ET/PT Mechanisms .................. 122! Table 3-6. Kinetic data for Phenols Oxidation by ClO 2 ................................................. 124! Table 3-7. Thermodynamic and Kinetic Parameters for Phenol Oxidation by Diferent Oxidants ......................................................................................................... 127! Table 3-8. Activation Parameters for Cresol and MOP Oxidation by Ir IV ..................... 131! Table 4-1. Kinetic Data for the Reaction of Ac-Y-NH 2 with Ir IV in the Presence of DBNBS .......................................................................................................... 137! Table 5-1. UV-Visible Absorbance of Os Complexes in 0.01 M HCl Solution. ............ 156! Table 5-2. Kinetic Data for the Reaction of Phenol with Os III in the Presence of Dipic or DBNBS .......................................................................................................... 165! Table 5-3. The Mechanism of Phenol Reaction and the Simulation Model ................... 170! Table 5-4. Comparison of the Experimental Data with the Simulation Results at diferent p[H + ] .............................................................................................................. 171! Table 5-5. Comparison of the Experimental Data to the Simulation Results in the Presence of Various Concentration of DBNBS ............................................ 171! Table A-1. Kinetic Dependence of Phenol Oxidation on [Phenol] tot .............................. 180! Table A-2. Kinetic Dependence of Phenol Oxidation on p[H + ] Without DBNBS ......... 181! Table A-3. Kinetic Dependence of Phenol Oxidation on p[H + ] with DBNBS ............... 182! Table A-4. Nonlinear-Least-Squares Regresion Results of the k obs /[phenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Absence of DBNBS ...................................................................................................................... 183! Table A-5. Nonlinear-Least-Squares Regresion Results of the k obs /[phenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS ...................................................................................................................... 184! Table A-6. Comparison of k obs in H 2 O and in D 2 O ......................................................... 185! Table A-7. Kinetic Dependence on [Cresol] tot ................................................................ 185! xi Table A-8. Kinetic Dependence of Biphenols and 4-Phenoxyphenol Oxidation on p[H + ] ...................................................................................................................... 186! Table A-9. Kinetic Dependence of Cresol Oxidation on p[H + ] with DBNBS ............... 187! Table A-10. Nonlinear-Least-Squares Regresion Results of the k obs /[cresol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS .................................................................................................................... 188! Table A-11. Kinetic Dependence of Cresol Oxidation on Temperature ........................ 189! Table A-12. Kinetic Isotope Efect for Oxidation of Cresol .......................................... 189! Table A-13. Kinetic Dependence on [Xylenol] tot ........................................................... 190! Table A-14. Kinetic Dependence of Xylenol Oxidation on p[H + ] ................................. 191! Table A-15. Nonlinear-Least-Squares Regresion Results of the k obs /[xylenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models ....................................... 192! Table A-16. Kinetic Isotope Efect for Oxidation of Xylenol ........................................ 193! Table A-17. Kinetic Dependence on [TMP] tot ................................................................ 193! Table A-18. p[H + ] Dependence of TMP Oxidation ........................................................ 194! Table A-19. Kinetic Isotope Efect for Oxidation of TMP ............................................. 195! Table A-20. Kinetic Isotope Efect for Oxidation of MOP ............................................ 195! Table A-21. Kinetic Dependence of MOP Oxidation on p[H + ] ...................................... 196! Table A-22. Kinetic Dependence of MOP Oxidation on Temperature .......................... 197! Table A-23. Kinetic Dependence of TBP Oxidation on p[H + ] ....................................... 198! Table A-24. Nonlinear-Least-Squares Regresion Results of the k obs /[TBP] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS .................................................................................................................... 199! Table A-25. Kinetic Dependence on [Ac-Y-NH 2 ] tot ....................................................... 200! Table A-26. Kinetic Dependence of Ac-Y-NH 2 Oxidation on p[H + ] ............................. 201! xii Table A-27. Nonlinear-Least-Squares Regresion Results of the k obs /[Ac-Y-NH 2 ] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS ....................................................................................................... 202! Table A-28. Kinetic Dependence on [Os III ] 0 ................................................................... 203! Table A-29. Kinetic data for the oxidation with various [phenol] tot ............................... 203! Table A-30. Kinetic Dependence on [Os II ] 0 ................................................................... 204! Table A-31. Kinetic Dependence on p[H + ] ..................................................................... 204! xii List of Schemes Scheme 1-1. Electron Transfer Proces .............................................................................. 2! Scheme 1-2. PCET pathways .............................................................................................. 8! Scheme 1-3. The electron and proton flow in Photosystem II .......................................... 10! Scheme 1-4. The electron and proton flow in Clas I RNRs. ........................................... 11! Scheme 2-1. Structure of spin trapping agents tested ....................................................... 26! Scheme 2-2. Expected dimerization products of phenoxy radical and their composition acording to literature .................................................................................. 37! Scheme 3-1. Structures of reductants ................................................................................ 77! Scheme 4-1. The structures of Ac-Y-NH 2 and Ac-Phe-NH 2 .......................................... 133! Scheme 5-1. Structure of [Os(phen) 3 ] 2+ .......................................................................... 150! xiv List of Figures Figure 1-1. Potential energy surfaces for an outer-sphere electron transfer reaction ......... 5! Figure 2-1. 1 H NMR spectrum of 2,4?-biphenol in CDCl 3 ............................................... 19! Figure 2-2. UV-vis spectra of (NH 4 ) 2 [IrCl 6 ] and (NH 4 ) 3 [IrCl 6 ] in aqueous solution ....... 21! Figure 2-3. Kinetic traces of oxidation of phenol by Ir IV at diferent p[H + ] ..................... 23! Figure 2-4. Comparative traces of the phenol reaction with added Ir III and no added Ir III 24! Figure 2-5. Trace of the phenol reaction with added DBNBS .......................................... 27! Figure 2-6. Plot of k obs vs [phenol] tot ................................................................................. 31! Figure 2-7. Plot of k obs /[phenol] tot vs p[H + ] in the absence of DBNBS ............................ 33! Figure 2-8. Plot of k obs /[phenol] tot vs p[H + ] with DBNBS ................................................ 34! Figure 2-9. The UV-vis spectra of four coupling isomers ................................................ 38! Figure 2-10. Titration of 3.0 ! 10 ?5 M 4,4?-biphenol ....................................................... 40! Figure 2-11. Titration of 5.0 ! 10 ?5 M 2,4?-biphenol ....................................................... 42! Figure 2-12. UV-Vis spectra of products of 4,4?-/2,2?-/2,4?-biphenol oxidation ............. 45! Figure 2-13. The kinetic trace of 4,4?-biphenol reacting with Ir IV at 398 nm .................. 46! Figure 2-14. Titration of 1.5 mL 4,4?-biphenol by Ir IV ..................................................... 48! Figure 2-15. Kinetic trace of 4,4?-biphenol oxidation at 488 nm ..................................... 50! Figure 2-16. Plot of k obs /[substrate] tot vs p[H + ] ................................................................. 51! Figure 2-17. The kinetic trace of 2,2?-biphenol reacting with Ir IV at 406 nm .................. 52! Figure 2-18. UV-Vis spectra during the reaction betwen 2,4?-biphenol and Ir IV ........... 54! xv Figure 2-19. The kinetic trace of 2,4?-biphenol reacting with Ir IV at 398 nm .................. 55! Figure 2-20. UV-Vis spectra during the reaction betwen phenol and Ir IV ...................... 57! Figure 2-21. The kinetic trace and curve fit of the reaction betwen phenol and Ir IV ...... 58! Figure 2-22. Comparative pH dependence of the phenol reaction experimental data and simulation results ......................................................................................... 66! Figure 3-1. Kinetic trace of the Ir IV consumption in the cresol oxidation ........................ 80! Figure 3-2. Comparative traces of the cresol reaction with Ir IV and in the presence of added Ir III in H 2 O ........................................................................................... 81! Figure 3-3. Kinetic traces of the cresol reaction with added DBNBS and first order fit .. 83! Figure 3-4. Plot of k obs vs [cresol] tot .................................................................................. 85! Figure 3-5. Plot of k obs /[cresol] tot vs p[H + ] with DBNBS ................................................. 86! Figure 3-6. Plot of log (k obs /T) vs 1/T ............................................................................... 87! Figure 3-7. Trace of the xylenol reaction and first-order fit ............................................. 89! Figure 3-8. Comparative traces of the xylenol reaction with added 5 ! 10 ?4 M Ir III and 2.5 ! 10 ?3 M Ir III .................................................................................................. 90! Figure 3-9. Plot of k obs vs [xylenol] tot ............................................................................... 93! Figure 3-10. Plot of k obs /[xylenol] tot vs p[H + ] ................................................................... 94! Figure 3-11. Kinetic trace of the Ir IV consumption in the TMP oxidation ....................... 96! Figure 3-12. Comparative traces of the TMP reaction with added Ir III ............................ 97! Figure 3-13. Plot of k obs vs [TMP] tot ................................................................................. 99! Figure 3-14. Plot of k obs /[TMP] tot vs p[H + ] ..................................................................... 100! Figure 3-15. 1 H NMR identification of TMP oxidation Products .................................. 101! Figure 3-16. Kinetic trace of the Ir IV consumption in the MOP oxidation ..................... 104! Figure 3-17. Plot of k obs /[MOP] tot vs p[H + ] ..................................................................... 105! Figure 3-18. Plots of ln (k obs /T) vs 1/T ............................................................................ 106! xvi Figure 3-19. Kinetic trace of the Ir IV consumption in the TBP oxidation ...................... 108! Figure 3-20. Kinetic trace of the Ir IV consumption in the TBP oxidation with DBNBS 109! Figure 3-21. Plot of k obs /[TBP] tot vs p[H + ] ...................................................................... 110! Figure 3-22. Plot of ("G ? ? w 12 ) versus "G?? ................................................................. 117! Figure 3-23. Correlations betwen the second-order rate constants for reaction of phenols (k ArOH ) and phenoxide anions (k ArO ?) with Ir IV versus substituent constants ! + ................................................................................................................. 119! Figure 3-24. Plot of pK a versus substituent constants ! + ................................................ 120! Figure 3-25. Plot of kinetic isotopic efect versus substituent constants ! + ................... 123! Figure 3-26. Plot of ln k (ClO 2 ) versus ln k (Ir IV ) ........................................................... 125! Figure 3-27. Plot of ln k ArOH versus "G ? CPET for phenol oxidation by diferent oxidants .................................................................................................................... 129! Figure 3-28. Plot of ln k ArOH versus "G ? CPET for the Ir IV reduction by diferent phenols130! Figure 4-1. Trace of the Ac-Y-NH 2 reaction and first-order fit ...................................... 135! Figure 4-2. Comparative traces of the Ac-Y-NH 2 reaction with added Ir III ................... 136! Figure 4-3. Trace of the Ac-Y-NH 2 reaction with DBNBS and first-order fit ............... 139! Figure 4-4. Plot of k obs vs [Ac-Y-NH 2 ] tot ........................................................................ 140! Figure 4-5. Plot of k obs /[Ac-Y-NH 2 ] tot vs p[H + ] .............................................................. 141! Figure 4-6. Titration of Ac-Y-NH 2 with Ir IV at p[H + ] = 5.5 and 25 ?C ........................... 143! Figure 4-7. The titration curve of Ac-Y-NH 2 solution by Ir IV solution at p[H + ] = 5.5 ... 144! Figure 5-1. 1 H NMR spectrum of [Os(phen) 3 ]Cl 2 in D 2 O (aromatic region) ................. 154! Figure 5-2. UV-vis Spectra of [Os(phen) 3 ]Cl 2 and [Os(phen) 3 ]Cl 3 ................................ 155! Figure 5-3. Kinetic traces of the phenol oxidation by Os III ............................................ 157! Figure 5-4. Plot of t 1/2 vs 1/[Os III ] 0 .................................................................................. 158! Figure 5-5. Kinetic trace of phenol oxidation by Os III at 550 nm ................................... 160! xvii Figure 5-6. Plot of k obs vs [phenol] tot 2 ............................................................................. 161! Figure 5-7. Plot of k obs /[phenol] tot 2 vs 1/[Os II ] 0 2 ............................................................. 162! Figure 5-8. Plot of [Os II ] 0 2 k obs /[phenol] tot 2 versus p[H + ] ................................................ 164! Figure 5-9. Kinetic trace of phenol oxidation by Os III with DBNBS at 480 nm ............ 166! 1 Chapter 1 LITERATURE REVIEW 1.1 Outer-Sphere Electron Transfer Electron transfer is a fundamental step of many chemical and biological proceses, such as photosynthesis, nitrogen fixation and phosphorylation. 1-3 The study of electron transfer has been facilitated by technological advances in stopped-flow kinetics, electrochemistry and pulse radiolysis. Electron transfer can be clasified as either inner- sphere or outer-sphere. In an inner-sphere electron transfer reaction, the oxidant and reductant connect with each other through a covalent chemical bond, whereas, in an outer-sphere electron transfer reaction, the reactants asociate through more indirect interactions. One motivation for the study of outer-sphere electron transfer is that most biological redox is done through this mechanism. Furthermore, the combination of theoretical works of Collision theory and Marcus theory is applied succesfully to describe the outer-sphere electron transfer reaction, through which the prediction of the actual rate constants can be achieved. A detailed electron transfer proces betwen D (donor) and A (aceptor) is described in Scheme 1-1. Collision Theory. The first step in Scheme 1-1 is the difusion-controlled collision betwen D and A. In Debye?s colliding sphere model, D and A are both treated as spheres with radii r 1 and r 2 , and charges of z 1 and z 2 , respectively. The energy needed to bring the 2 Scheme 1-1. Electron Transfer Proces. two separated reactants (spheres from an infinite distance) to the closest approach distance (the center-to-center separation distance), r 12 = r 1 + r 2 , is known as the electrostatic energy or Coulombic work (w 12 ) as shown in eq 1-1. 4 This term depends on dielectric medium (dielectric constant D) and the total ionic strength (?) of the solution. w 12 = z12 e Dr(+!?) (1-1) Here, the constant " is the reciprocal Debye radius and e is the electron charge (4.803 ! 10 ?10 esu). Since al of the reactions in this thesis are aqueous, D is set to 78.4. If the distance r is in angstroms and ? in molar, " = 0.328 ? 1/2 /mol 1/2 and w 12 can be simplified to eq 1-2 in kilojoules per mole. D z 1 +1 ! A z 2 ?1 ! D z 1 ! A z 2 ! Precursor D z 1 +1 A z 2 ?1 Successor !G? !G?? D z 1 ! A z 2 ! r 1 ! r 2 ! w 12 w 21 3 w 12 = 17. z2 r2(+038r1?) (1-2) The collision results in the transient formation of ?precursor? complex, [D z1 , A z2 ], subsequent electron transfer forms the short-lived ?succesor? complex, [D z1+1 , A z2?1 ]. The electron transfer details wil be explained by Marcus theory in the next part. We define the corrected Gibbs free energy, "G??, as the free energy diference betwen the succesor and precursor. This difers from the standard Gibbs free energy, "G?, which is the energy diference betwen separated reactants and separated products. This is important because Marcus theory correlates rate constants to "G?? rather than "G?. The succesor complex disociates to give the final products of the electron transfer, D z1+1 and A z2?1 . The energy asociated with this proces (?w 21 ) is electrostatic energy w 21 = 7. (z1+)2! r2038r? (1-3) The corrected Gibbs free energy "G?? is related to "G? by eq 1-4. "G?? = "G? ? w 12 + w 21 (1-4) The Gibbs free energy, "G?, in eq 1-4 can be calculated acording to eq 1-5 using the standard reduction potentials of the oxidant and reductant. 4 "G? = ?Z (E oxidant ? E reductant ) F (1-5) Marcus Theory for Electron Transfer. Rudolph A. Marcus was recognized with a Nobel Prize in Chemistry in 1992 for his theoretical description of the outer-sphere electron transfer mechanism. Over the past few decades, Marcus theory has been increasingly applied in the study of chemistry and biology. 5-11 In Marcus theory, the free energy of activation, "G ? , is related to the corrected Gibbs free energy of the reaction, "G??, by the quadratic equation 1-6: 2 "G ? = w 12 + ! 4 (1 + ?' ) 2 (1-6) The w 12 term is the same as that described by eqs 1-1 and 1-2. # is the energy asociated with the outer- and inner-sphere reorganization that acompanies the transition from the precursor to the succesor states. For an outer-sphere reaction, the energy barrier is determined by the changes in bond lengths and angles of the donor and aceptor molecules, which is quite smal in most cases. However, the energy barrier for reorientation of solvent molecules is a major component due to the diferent electronic properties and charge distribution betwen the succesor and precursor. "G ? , # and "G?? are represented schematicaly in Figure 1-1. The potential energy surfaces of electron transfer betwen the reactants and products are depicted in Figure 1-1. The horizontal axis is the reaction coordinate, which corresponds to the motions of al atomic nuclei. The left parabolic curve R represents the 5 Figure 1-1. Potential energy surfaces for an outer-sphere electron transfer reaction. potential surface of reactants plus the surrounding medium with the minimum point indicating the nuclear coordinate for an equilibrium configuration of precursor complexes. Whereas, the curve P represents the potential surface of products plus the surrounding medium with the minima denoting the equilibrium nuclear coordinates of succesor complexes. The intersection of the two parabolas, the transition state, is the only place where electron transfer can occur while complying with both Franck-Condon principle (that is, electron transfer occurs so rapidly that no change in nuclear coordinates can occur during the transfer) and the first law of thermodynamics (conservation of energy). However, weak electronic interaction betwen D and A splits the potential energy surfaces, and the orbital mixing generates the resonance energy, which produces the electronic coupling. This can be described by the electronic matrix element, H AB . The energy diference betwen two minima is the corrected Gibbs free energy of the reaction, AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G " =w12+ ! 4 ( G#? ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G " =w12+ ! 4 ( #? ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P ! P 2H AB k et = ! Z exp(?"G ? /RT) ! =w12+ 4 ( "? ) 2 "G ? AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " 12 4 ( #? ! ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G " =w12+ ! 4 ( G#? ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P ! P 2H AB k et = ! Z exp(?"G ? /RT) ! =12+ 4 ( "? ) 2 "G ? AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " 12 4 ( #? ! ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G " =w12+ ! 4 ( G#? ) 2 !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P ! P 2H AB k et = ! Z exp(?"G ? /RT) ! =w12+ 4 ( "? ) 2 "G ? AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " 1( #? ) !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! 2 4 !G ? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " =w1+ 4 ( #? ) !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P ! P 2H AB k et = ! Z exp(?"G ? /RT) ! =12+ 4 ( "? )"G ? AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " 12( #? ) !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! !G ? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P AUBURN UNIVERSITY! Chemistry and Biochemistry! College of Sciences and Mathematics! " =w12+ 4 ( #? ) !G ? !G?? ! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P !G?? P ! P 2H AB k et = ! Z exp(?"G ? /RT) ! =w12+ 4 ( ? )"G ? 6 "G??. The energy diference betwen the precursor and the intersection is the free energy of activation, "G ? , which is related to electron transfer rate constant, k et , through the Eyring equation 1-7. k et = $ Z exp(?"G ? /RT) (1-7) Here, Z is the collision frequency and reported to be 10 11 M ?1 s ?1 . 2 $ is the transmision coeficient, which represents the probability of electron transfer per pasage. $ is large (= 1) for an adiabatic reaction where the change of nuclear coordinates is sufficiently slow when pasing the intersection and the system stays at equilibrium. However, when the jumping occurs with a high pasing velocity through the intersection, or when the spliting of the potential surface is smal due to weak electronic interaction, the probability $ of going from the precursor to the succesor surface is very smal ($ << 1), and this proces is defined as nonadiabatic. Marcus Cross Relation (MCR). The reorganization energy, #, can be obtained from the two relative self-exchange reactions 1-8 and 1-9. * D z1 + D z1+1 ! * D z1+1 + D z1 k 11 , # 11 (1-8) * A z2 + A z2?1 ! * A z2?1 + A z2 k 22 , # 22 (1-9) [D z1 , A z2 ] (Precursor) ! [D z1+1 , A z2?1 ] (Succesor) k 12 , # 12 , K 12 (1-10) k 11 and # 11 correspond to the reaction betwen oxidized and reduced forms of the electron donor D and k 22 and # 22 belong to the reaction betwen oxidized and reduced forms of 7 electron aceptor A. For both reactions 1-9 and 1-10, "G?? = 0 and acording to eq 1-6, "G ? = w + #/4. An asumption (eq 1-11) 12 used to derive the MCR (eqs 1-12 to 1-15) 2 is that the reorganization energy for the cross-reaction (eq 1-10), # 12 , is equal to the average of # 11 and # 22 . # 12 = !1+ 2 (1-11) k 12 = (k 11 k 22 K 12 f 12 ) 1/2 W 12 (1-12) lnf12= [l1+(w!21)/RT] 2 42/Z (1-13) W 12 = exp[(#w 12 # w 21 + w 11 + w 22 )/2RT] (1-14) ij = 17.zi j rij(+0328r?) (1-15) As noted above, Z is the collision frequency (10 11 M ?1 s ?1 ), w 11 , w 22 , w 12 and w 21 correspond to Coulombic works. K 12 is the equilibrium constant. W 12 and f 12 are adjustable factors that depend on the ionic strength, reaction media, radii and charges of reactants. 1.2 Proton-Coupled Electron-Transfer (PCET) In 1959 Halpern proposed that the oxidation of formate by permanganate in acidic solutions might undergo electron transfer coupled with proton transfer. 13 The coupling of electron and proton through in phosphorylation was recognized two years later by 8 Mitchel. 14 PCET has subsequently been studied extensively through theoretical and experimental means. 15-24 In redox reactions involving weak acids and bases, deprotonation can occur with electron transfer. When one or more protons and electrons are transferred, two fundamental mechanisms are most frequently proposed: sequential and concerted PCET. Three possible pathways are considered for transferring one electron and one proton, as shown in Scheme 1-2: Electron transfer followed by proton transfer (ET/PT), proton transfer prior to electron transfer (PT/ET) and concerted PCET (CPET) mechanism (protons and electrons are transferred simultaneously). Scheme 1-2. PCET pathways. Here, DH is the electron and proton donor and A is the electron aceptor. CPET may occur in acidic media where high H + concentrations limit the deprotonation of DH, and consequently PT/ET pathway. Marcus Theory for H 2 O-CPET. In aqueous solution, the solvent water serves as the proton aceptor in CPET. 25-27 The self-exchange CPET reaction is described in eq 1- 16, with # CPET as the self-exchange reorganization energy. A z 2 + DH A z 2 ?1 + D ? + H + A z 2 ?1 + DH +? A z 2 + D ? + H + ET PT PT ET CPET !G ? !G?? !!! 2H AB Successor P ot e nt i a l E ne r gy Reaction Coordinate Precursor R R P P 9 ( * DH???H 2 O) + (D ? , + H 3 O) ! ( * D ? , + H 3 O) + (DH???H 2 O) k CPET , # CPET (1-16) Here, ????? represents hydrogen bonds betwen DH (for example, phenol) and H 2 O. 28 Combined with the self-exchange reaction of the oxidant (1-9), the total reorganization energy can be calculated by # = (# CPET + # 22 )/2. 26 1.3 Oxidation of Phenols Since they were first isolated from coal tar by Runge in 1834, 29 phenol and its derivatives have been widely found in nature, for example, serving as components in many antibiotics, colorants, flavonoids, neurotransmiters and hormones. 30-34 Other phenols are medicinaly and industrialy vital. 35-37 Tyrosine, one of the 22 proteinogenic amino acids, contains a phenol side chain which is esential for its biological functions; In a phosphorylation proces, for example, tyrosine residues are tagged with phosphate groups and act as receptors in signal transduction. 38 Tyrosine sulfation is another proces where a sulfate group is added to tyrosine residues, through which the protein-protein interactions are strengthened. 39 One important function of tyrosine is that it serves as an electron donor in photosystem II, reducing the oxidized chlorophyll. 40,41 In photosynthesis, the photons of sunlight are absorbed by chlorophyll A in photosystem II and excite P 680 to the excited state, P 680 + . During the excitation proces, electrons are released and pased through a redox reaction to pheophytin, then subsequently to plastoquinone, where protons are transferred to the corresponding quinol. Eventualy electrons are transferred to photosystem I where NADP is reduced to NADPH. The excited P 680 + is a strong 10 oxidizing agent with redox potential of 1.26 V 42 and has to be reduced to its ground state in order to absorb more photons. Water is a good source of electrons and protons. P 680 + can oxidize water releasing electrons and protons with dioxygen as a by-product. This reaction is stepwise: a tyrosine residue Yz is directly oxidized by P 680 + through a proton- coupled electron-transfer proces, sequentialy, the tyrosine radical formed during this step is reduced by the oxygen evolving complex (OEC), a manganese cluster, where water is split into oxygen. 43 Scheme 1-3 shows the flow of the electron and proton from water to plastoquinone in photosystem II. Scheme 1-3. The electron and proton flow in Photosystem II. Another esential function of tyrosine is found in DNA replication and repair, where long-range PCET is believed to occur. During this proces, an electron and a proton are transferred betwen cysteine (C439) in clas I ribonucleotide reductases (RNRs) to the P680 + P680 Pheophytin Plastoquinone e e, H + Tyrosine OEC H 2 O O 2 + H + e, H + e e 11 orthogonal PCET through which the electron and proton are transferred to diferent aceptors. 18,44-46 Scheme 1-4 shows the PCET in Clas I RNRs. Scheme 1-4. The electron and proton flow in Clas I RNRs. Obtained from on September 2011. One-electron oxidation of phenols in aqueous media is currently of intense interest due to the abovementioned biochemical systems. The structural constraints of these systems enforce long-distance electron transfer and outer-sphere mechanisms. The current consensus is that these reactions are subject to general base catalysis. 47, 48 However, the pH dependence when water is the proton aceptor remains enigmatic: there is general agreement that the reactions are pH-independent at high acidity, increase in rate with increasing pH, and reach a limiting rate at high pH, but some reports indicate an approximate inverse half-order dependence on [H + ] in the intermediate pH region, 49, 50 while others support an inverse first-order dependence. 51, 52 The half-order dependence 8/28/11 6:47 P MJoAnne Stubbe Research Group - MIT Pa ge 2 of 6http ://web.mit.edu/biochemistry/research.htm Proposed biosynthesis, activation, and regulation of RNR in E. coli Cotruvo, J.A.; Stubbe, J. An Active Dimanganese(III)-Tyrosyl Radical Cofactor in Escherichia coli Class Ib Ribonucleotide Reductase. Biochemistry Articles ASAP. Lohman, G.J.S.; Gerfen, G.J.; Stubbe, J. Inactivation of Lactobacillus leichmannii ribonucleotide reductase by F 2 CTP: adenosylcobalamin destruction and formation of a nucleotide based radical. Biochemistry Just Accepted. Lohman, G.J.S.; Stubbe, J. Inactivation of Lactobacillus leichmannii ribonucleotide reductase by F 2 CTP: covalent modification. Biochemistry Just Accepted. Shanmugam, M.; Doan, P.E.; Lees, N.S.; Stubbe, J.; Hoffman, B.M. Identification of Protonated Oxygenic Ligands of Ribonucleotide Reductase Intermediate X. J. Am. Chem. Soc. 2009, 131, 3370-3376. Artin, E,; Wang, J.; Lohman, G.J.S.; Yokoyama, K.; Yu, G.; Griffin, R.G.; Bar, G.; Stubbe, J. Insight into the Mechanism of Inactivation of Ribonucleotide Reductase by Gemcitabine 5'-Diphosphate in the Presence or Absence of 12 has led to the controversial proposal of a pH-dependent driving force. 49, 50 The question of whether the reactions in acidic media (pH 1#3) operate through a sequential or concerted PCET mechanism is unresolved. One approach to gaining insight into PCET is to study reactions in which one of the reactants is a typical inorganic outer-sphere electron-transfer reagent with no acid/base properties: this constraint confines much of the electron-proton coupling considerations to the other reaction partner. The transition metal complexes, such as [IrCl 6 ] 2? and [Os(phen) 3 ] 3+ , have proven to be valuable outer-sphere electron-transfer oxidants. 53-60 In Chapter 2, we present kinetic data on the oxidation of phenol and the four coupling products (4,4?-/2,2?-/2,4?-biphenol and 4-phenoxyphenol) by [IrCl 6 ] 2? , obtained under conditions where bufer catalysis is insignificant. These data support an inverse first-order dependence on [H + ] near neutral pH, a concerted PCET mechanism at low pH, and a degree of overoxidation that is pH-dependent. The oxidation of phenol by Ir IV was the subject of a clasic study by Cecil and Litler 40 years ago. 51 That paper showed that the reaction yields Ir III and a variety of phenolic oxidation products and that it has a simple two-term rate law (described in the following chapter). Limitations of the study by Cecil and Litler included the inability to exclude completely the efects of Ir III , the instrumental restriction to measuring relatively slow rates, the unknown driving forces for the rate-limiting steps, the unknown rate of self-reaction of the phenoxyl radicals, the unknown pK a of ArOH ?+ , and uncertainty as to whether the acid pathway (k ArOH ) was a sequential proces of electron transfer to form ArOH ?+ followed by its deprotonation or whether it was a concerted PCET proces. Here we use DBNBS as a phenoxyl radical scavenger to eliminate the kinetic efects of Ir III , we use stopped-flow methods to obtain 13 kinetic data under conditions where the rates are much faster, we use numerical modeling and the now-known properties of the phenoxyl radicals to confirm the basic mechanism, and we use kinetic isotope efects (KIEs) to probe the concertednes of the ArOH oxidation pathway. Chapter 3 extends the studies of oxidation by Ir IV to alkyl- and alkoxy-substituted phenols (2-methylphenol, 2,6-dimethylphenol, 2,4,6-trimethylphenol, 4-tert-butylphenol and 4-methoxyphenol). The reaction of 2,4,6-trimethylphenol yields 4-hydroxymethyl- 2,6-dimethylphenol rather than the coupling products obtained from phenol. Apparently the 2,4,6-trimethylphenoxyl radical disproportionates with oxidation at the 4-methyl position, which eliminates the overoxidation proces and gives an excelent inverse first- order dependence on [H + ] near neutral pH. Marcus theory is applied to explain both electron transfer and H 2 O-CPET proceses. Chapter 4 describes the oxidation of one of the tyrosine derivatives, N-acetyl- tyrosine amide, by Ir IV . Overoxidation is observed and the values of k ArOH and k ArO ? are obtained by fiting the data to a two-term rate law. The oxidation of phenol by tris-(1,10-phenanthroline)osmium(III) is presented in Chapter 5. Analysis of the kinetic data yields a second-order dependence on [Os(phen) 3 ] 3+ and phenol, and an inverse second-order dependence on [Os(phen) 3 ] 2+ and H + . 14 Chapter 2 PROTON-COUPLED ELECTRON-TRANSFER OXIDATION OF PHENOL BY HEXACHLOROIRIDATE(IV) This Chapter is based on the following papers and reprints were made with permision from America Chemical Society. (a) Song, N.; Stanbury, D. M. Inorg. Chem. 2008, 47, 11458. (b) Song, N.; Stanbury, D. . Inorg. Chem., DOI: 10.1021/ic201897m. 2.1 Introduction Redox reactions in aqueous solution are often acompanied by changes in protonation. When the reactions proced via one-electron steps the phenomenon of proton-coupled electron transfer (PCET) is often involved. The breadth of importance of PCET is imense, and PCET is of great significance because it is often an absolute criterion of reactivity. 23 One approach to gaining insight into PCET is to study reactions where one of the reactants is a typical inorganic outer-sphere electron-transfer reagent with no acid/base properties: this constraint confines much of the electron-proton coupling considerations to the other reaction partner. Oxidation of phenols has become central in developing the concepts of PCET, in part because of its importance in revealing the function of tyrosine in redox proteins, but also because these reactions are quite amenable to study. Acordingly, the oxidation of phenol by [IrCl 6 ] 2? has been the focus of a clasic publication, 51 and it is now of interest as a model for reactions where electron 15 transfer occurs in concert with proton transfer to the solvent: H 2 O-CPET. 25-27 Despite its importance, the oxidation of phenol by [IrCl 6 ] 2? presents certain dificulties. One of these is that the phenolic products are a complex mixture, 51 and it is unclear how the post-rate-limiting steps lead to this mixture. It is also unclear whether these later stages in the reaction have any influence on the measurements of the putative rate-limiting steps. Of further concern is the pH dependence of the rates, which displays deviations from the clasic two-term rate law for such reactions 27 and might provide additional evidence in support of a pH-dependent rate constant as has been reported previously for the oxidation of phenol by [Ru(bpy) 3 ] 3+ . 50 Herein is reported a study on the [IrCl 6 ] 2? oxidations of the four major products derived from coupling of the phenoxyl radical. These rates are then incorporated into an overal mechanism for the oxidation of phenol, which shows that overoxidation is responsible for many of the reaction products. Moreover, it is shown that the degree of overoxidation is pH-dependent and that this can acount for the observed deviations from the clasical two-term rate law. 2.2 Experimental Section 2.2.1 Reagents and Solutions Al commercial chemical reagents were used as received except as noted. Amonium hexachloroiridate(III) monohydrate (Ir III ), 2,6-dimethylacetanilide, chlorosulfonic acid, 3-chloroperoxybenzoic acid, disodium, N-tert-butyl-phenyl-nitrone (PBN), %-(4-pyridyl-N-oxide)-N-tert-butylnitrone (POBN), 2-methyl-2-nitrosopropane (MNP), 5,5-dimethyl-1-pyrroline N-oxide (DMPO), deuterium oxide, 3,5- dibromosulfanilic acid sodium salt, 2-iodophenol, 4-hydroxyphenylboronic acid, 16 paladium(II) acetate, 1,1?-bis(diphenylphosphino)-ferrocence (dppf), potasium carbonate, sodium acetate anhydrous, cacodylic acid and sodium hydroxide were purchased from Sigma-Aldrich Chemicals Co. Perchloric acid, amonium perchlorate, sodium chloride, amonium chloride, copper(II) nitrate trihydrate, acetic acid, monochloroacetic acid, ethanol, diethyl ether, 1,4-dioxane, petroleum ether, ethyl acetate, acetonitrile, hydrochloric acid and hydrogen peroxide were from Fisher Scientific Co. 2,2?-Biphenol, 4,4?-biphenol and 4-phenoxyphenol are commercialy available from Acros Organics. Amonium hexachloroiridate(IV) (Ir IV ) was purchased from Alfa or prepared acording to the literature 61 by the addition of amonium chloride (Fisher) to a solution of sodium hexachloroiridate(IV) hexahydrate. Phenol (Fluka) was recrystalized from a 75% weight/weight water solution as described in the literature. 62 Al solutions were freshly prepared with deionized water provided by a Barnstead NANO Pure Infinity ultrapure water system, and purged with argon gas prior to the reactions to prevent potential complications caused by O 2 . In order to increase the concentration of the solution, 4,4?-biphenol or 2,4?-biphenol was first disolved in CH 3 CN or ethanol and then diluted with water to make a reaction solution where les than 1% w/w organic solvent was present. The ionic strength was adjusted by lithium perchlorate trihydrate (GFS) and was approximately equal in both oxidants and reductants solutions to prevent Schlieren efects (or refractive index efect 63 ). Selected buffer solutions (acetate, monochloroacetate, and cacodylate buffers) were applied to control the pH if necesary. Preparation of Sodium 3,5-Dibromo-4-nitrosobenzenesulfonate (DBNBS). This compound was synthesized from 3,5-dibromosulfanilic acid sodium salt acording to the 17 method of Kaur et al. 64 and further purified using procedure type C of Hamilton et al.. 65 A mixture of 3,5-dibromosulfanilic acid (10 mmol), anhydrous sodium acetate (10 mol), 7.9 mL of a 30% aqueous hydrogen peroxide solution and 30 mL of glacial acetic acid was warmed gently to disolve the solid. The solution was stored at room temperature for 14 days. Then the crude yelow solid product was collected and washed with glacial acetic acid, cold ethanol, diethyl ether/1,4-dioxane (1:1), and cold ethanol again. A pale-yelow powder was obtained after drying. Yield: 30%. Mp: > 300 ?C. 1 H NMR (D 2 O): $ 8.30 (s, 2H). 13 C NMR (D 2 O): $ 119.10 (aryl-CBr), 131.03 (aryl-CH), 141.00 (aryl-CSO 3 Na), 147.61 (aryl-CNO). Preparation of Sodium 2,4-Dimethyl-3-nitrosobenzene Sulfonate (DMNBS). This compound was prepared acording to the literature. 66,67 The 2,6-dimethylacetanilide (18 mmol) was added into 10 mL of cold ClSO 3 H with stirring. After reacting for 15 min at 5 ?C, 1h at 25 ?C and 10 min at 40 ?C, the mixture was poured onto cracked ice to form a white solid of sulfonyl chloride. Then this compound was treated with concentrated HCl under reflux for 1 h before neutralization by NaOH to produce sodium 2,6- dimethylaniline-3-sulfonate. Sodium sulfonate (10 mmol) was disolved in 30 mL of methanol and oxidized by m-chloroperoxybenzoic acid (10 mmol) with stirring for 1 h at room temperature. Then ether was added to precipitate the final product. Yield: 80%. 1 H NMR (D 2 O): $ 2.56 (s, 3H), 2.75 (s, 3H), 7.54 (d, 1H), 8.13 (d, 1H). Preparation of 2,4?-Biphenol. 2,4?-Biphenol was prepared through a previously described Pd-catalyzed Suzuki-Miyaura coupling reaction. 68 Under dry nitrogen, 4- hydroxyphenylboronic acid (1.4 mmol), Pd(OAc) 2 (0.14 mmol), bis(diphenylphosphino)- ferrocence (dppf) (0.14 mmol), and K 2 CO 3 (3 mmol) were added to a solution of 2- 18 iodophenol (1 mmol) in 10 mL THF. The reaction was stirred under reflux for 1 day and monitored by thin-layer chromatography. After the reaction was completed, the crude product was purified by column chromatography on silica gel (eluent PE/EtOAc = 3.5/1). Then, sublimation was performed under vacuum at 160 ?C to remove the non-volatile metal residue impurities, and resublimation was carried out at 110 ?C to remove the volatile impurities. Yield: 50%. Mp: 161.2?162.8 ?C. 1 H NMR (CDCl 3 ) (Figure 2-1): $ 4.89 (s, 1H), 5.14 (s, 1H), 6.94?6.99 (m, 4H), 7.20?7.24 (m, 2H), 7.34?7.36 (m, 2H). 2.2.2 Methods A Corning 450 pH/ion meter was used with a Metler Toledo InLab 421 or InLab Semi-Micro-L combination pH electrode. The reference electrode electrolyte was replaced with 3 M NaCl to prevent the precipitate of KClO 4 . Electrode calibrations at ? = 0.1 M (LiClO 4 ) were carried out with 0.01?0.1 M perchloric acid. With the known H + concentration and pH reading, the activity coeficient ! (= 0.839 ? 0.04) was obtained from equation p[H + ] = pH + log !, where p[H + ] is equal to ?log [H + ]. When the pK a of 4,4?/2,4?-biphenol is measured, an alkaline error may occur with a pH value higher than 11. Thus, we calibrated the electrode with 1 ! 10 ?3 M NaOH at ? = 0.1 M (LiClO 4 ) and found that the true pH is equal to the apparent pH plus 0.3. Al measurements were performed at 25.0 ? 0.1 ?C. The kinetics experiments were carried out on a Hi-Tech SF-51 stopped-flow spectrophotometer with OLIS 4300 data acquisition and analysis software. UV-vis spectra were monitored on a HP-8453 diode- array spectrophotometer equipped with a Brinkman Lauda RM6 thermostatted water bath to maintain the temperature at 25 ?C, and the solutions were prepared in a 1.0 cm quartz cells. Because (NH 4 ) 2 [IrCl 6 ] has strong absorbance around 488 nm while the 19 Figure 2-1. 1 H NMR spectrum of 2,4?-biphenol in CDCl 3 . This spectrum matches the previously reported results. 69 5.05.25.45.65.86.06.26.46.66.87.07.27.4 ppm 4.8855.1366.9436.9646.9927.2017.2217.2387.2407.2437.3397.361 1.011.004.151.882.09 PC 1.00 GB 0 LB 0.00 Hz SSB 0 WDW no SF 400.1800074 MHz SI 131072 SFO1 400.1824713 MHz PL1 1.60 dB P1 11.50 usec NUC1 1H ======== CHANNEL f1 ======== TD0 1 D1 1.00000000 sec TE 295.1 K DE 6.50 usec DW 60.400 usec RG 181 AQ 3.9584243 sec FIDRES 0.126314 Hz SWH 8278.146 Hz DS 2 NS 16 SOLVENT CDCl3 TD 65536 PULPROG zg30 PROBHD 5 mm BBO BB?1H INSTRUM Avance400 Time 10.03 Date_ 20110118 PROCNO 1 EXPNO 1 NAME syn2,4??diphenol 20 corresponding product (NH 4 ) 3 [IrCl 6 ] does not, as shown in Figure 2-2, al kinetics data were obtained by monitoring the absorbance of Ir IV at # max (488 nm) with " 488 = (3.9 ? 0.1) ! 10 3 M ?1 cm ?1 . 55 In order to detect the DBNBS influence on the reactions of phenol, we also observed the absorbance change over the wavelength 400?525 nm at intervals of 25 nm. For the phenol and 4-phenoxyphenol reactions, the observed pseudo-first-order rate constants were obtained from the fiting of kinetic traces over 5 half-lives to first-order exponential functions. For 4,4?-biphenol and 2,2?-biphenol the observed initial rate constants were determined from the slope of the linear regresion of logarithm of absorbance at 488 nm within the first half-life, while for 2,4?-biphenol the rate constants were obtained from double-exponential fits. The Specfit/32 version 3.0.15 global analysis system was applied to simulate the reaction traces, and the GraphPad Prism 4 or 5 software was used to analyze the rate law with 1/Y 2 weighting. Each reported observed rate constant k obs is the average of at least five individual shots. The average k obs values are fit to rate law in Prism, and therefore, the reported uncertainties reflect the scater of k obs . 1 H and 13 C NMR spectra were acquired on a Bruker AV 400 MHz spectrometer; chemical shifts in CDCl 3 are relative to TMS and in D 2 O are relative to DS. The melting points were obtained using an Electrothermal IA 9100 digital melting point apparatus. Cyclic voltamograms (CV) and Osteryoung square-wave voltamograms (OSWV) were performed on a BAS 100B electrochemical analyzer equipped with a BAS C3 cel stand and a purging and stirring system; a glasy carbon electrode acts as the working electrode, Ag/AgCl (3.0 M NaCl) is the reference electrode (E? = 0.205 V vs NHE), 55 and a Pt wire is used as the auxiliary electrode. 21 Figure 2-2. UV-vis spectra of (NH 4 ) 2 [IrCl 6 ] (solid line) and (NH 4 ) 3 [IrCl 6 ] (dashed line) in aqueous solution. T = 25 ?C. 200 320 440 560 680 800 0 1000 2000 3000 4000 5000 Wavelength, nm ! , M - 1 c m - 1 Ir IV Ir III 22 Quantum calculations of electronic spectra were performed with the Spartan 08 software package. 70 2.3 Results The kinetics traces for the consumption of 1 ! 10 ?4 M Ir IV in its reaction with a large exces of phenol were obtained at various p[H + ] at 488 nm. Figure 2-3a exhibits a typical kinetic trace for a reaction with 0.44 M phenol in 0.05 M HClO 4 (p[H + ] = 1.3). Such kinetic traces do not give good fits with either first- or second-order rate laws. As the p[H + ] increases, good-quality pseudo-first-order fits are obtained at both p[H + ] = 2.9 and 4.8 with 0.044 M phenol (Figure 2-3b and Figure 2-3c). At high p[H + ] (= 7.1), the reaction with 1.8 ! 10 ?3 M phenol is much faster and becomes non-pseudo-first-order again (Figure 2-3d). These results show a strong p[H + ] efect on phenol oxidation by Ir IV . As shown below, the deviations at low pH arise from inhibition by Ir III , while the deviations at high pH are due to the absorbance of phenolic products. 2.3.1 Oxidation of Phenol 2.3.1.1 Hexachloroiridate(II) Effects Kinetic Inhibition by Ir III . At p[H + ] = 1.3, Figure 2-4a shows the efects of adding 5-fold exces of Ir III under conditions otherwise identical with those in Figure 2-3a. The retarding efect of Ir III at this pH is sen to be quite strong and clearly can be expected to cause deviations from pseudo-first-order kinetics of the type shown in Figure 2-3a. On the other hand, at p[H + ] = 7.1, no Ir III inhibition is observed with the addition of 5- and 10-fold exces of Ir III (Figure 2-4b, c), and so the deviations from pseudo-first order 23 Figure 2-3. Kinetic traces of oxidation of phenol by 1 ! 10 ?4 M of Ir IV at diferent p[H + ]. Lower boxes show the experimental traces (solid lines) and the pseudo-first-order fits (dashed lines). Upper boxes show the residuals in the fits. ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) [Phenol] tot = 0.44 M; [HClO 4 ] = 0.05 M. (b) [Phenol] tot = 0.044 M; p[H + ] = 2.9 (0.02 M monochloroacetate buffer). (c) [Phenol] tot = 0.044 M; p[H + ] = 4.8 (0.02 M acetate buffer). (d) [Phenol] tot = 1.8 ! 10 ?3 M; p[H + ] = 7.1 (0.02 M cacodylate buffer). -0.05 0.00 0.05 a R e s i d u a l s 0 20 40 60 80 0.0 0.1 0.2 0.3 p[H + ] = 1.3 Time, s A 488 -0.05 0.00 0.05 b R e s i d u a l s 0 20 40 60 80 100 0.0 0.1 0.2 0.3 p[H + ] = 2.9 Time, s A 488 -0.02 0.00 0.02 d R e s i d u a l s 0.0 0.5 1.0 1.5 0.0 0.1 0.2 0.3 p[H + ] = 7.1 Time, s A 488 -0.02 0.00 0.02 c R e s i d u a l s 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.1 0.2 0.3 p[H + ] = 4.8 Time, s A 488 24 Figure 2-4. Comparative traces of the phenol reaction with added Ir III (dashed line) and no added Ir III (solid line) in H 2 O. [Ir IV ] 0 = 1 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) [Phenol] tot = 0.44 M; [Ir III ] = 5 ! 10 ?4 M; [HClO 4 ] = 0.05 M. (b) [Phenol] tot = 1.8 ! 10 ?3 M; [Ir III ] = 5 ! 10 ?4 M; p[H + ] = 7.1 (0.02 M cacodylate buffer). (c) [Phenol] tot = 1.8 ! 10 ? 3 M; [Ir III ] = 1 ! 10 ?3 M; p[H + ] = 7.1 (0.02 M cacodylate buffer). 0 120 240 360 480 0.00 0.12 0.24 0.36 0.48 0.60 no added Ir III with added 5-fold Ir III Time, s A 488 a 0.0 0.5 1.0 1.5 2.0 no added Ir III with added 5-fold Ir III b Time, s 0.0 0.5 1.0 1.5 2.0 2.5 0.00 0.12 0.24 0.36 0.48 0.60 no added Ir III with added 10-fold Ir III c Time, s A 488 25 kinetics at this pH (Figure 2-3d) are due to factors other than inhibition by Ir III . Basicity of Ir III . The basicity of Ir III was probed by measuring the pH dependency of the cyclic voltametry of 1 ! 10 ?3 M of Ir IV in the p[H + ] range of 0?2 at 1 M ionic strength (LiClO 4 ). p[H + ] values were adjusted by HClO 4 acording to p[H + ] = ? log [HClO 4 ]. Reversible CVs (E 1/2 = 0.717 ? 0.007 V vs Ag/AgCl) were obtained over this pH range. No pH dependence of E 1/2 was observed under al of these conditions, which implies that Ir III is not significantly basic even in 1 M H + . Bruhn et al. observed a smal but significant dependence of E 1/2 over this same pH range in H + /Na + media; 71 we atribute this efect to the difering specific interaction coeficients of these two ions. 2.3.1.2 General Base Catalytic Test. Tests for general base catalysis were performed as follows: First, for the phenol reaction at p[H + ] = 6.5 (shown in Table 2-1), the cacodylate buffer concentration was varied from 2 to 20 mM, and after correction for pH fluctuations, no significant corresponding rate variation was detected. Table 2-1. Kinetic Dependence on Concentration of Buffer. a [buffer] ! 10 ?3 , M (k obs /[phenol] tot ) !10 ?3 , M ?1 s ?1 p[H + ] 2.0 2.31 6.44 4.0 2.72 6.49 20 3.12 6.57 a Al the reactions were run under pseudo-first-order conditions. [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] tot = 4.4 ! 10 ?3 M; ? = 0.1 M (LiClO 4 ); The p[H + ] values were maintained using cacodylate buffers; T = 25 ?C. 26 Second, tests were also performed on the reaction betwen 1.8 ! 10 ?3 M of phenol and 1 ! 10 ?4 M of (NH 4 ) 2 IrCl 6 at p[H + ] = 7 (0.02 M cacodylate buffer) at 0.1 M ionic strength (LiClO 4 ). The concentration of amonium was varied from 0.2 to 0.7 mM because the oxidant used was (NH 4 ) 2 [IrCl 6 ]; no kinetic efect was detected. Third, if phenol itself were a catalyst, then a second-order dependence on [phenol] would be expected, contrary to our observations as described below. In summary, these results show that the kinetic data presented herein on phenol oxidation by Ir IV are free of complications arising from general base catalysis. This conclusion is in agreement with the recent report from Irebo et al. and Bonin et al. that general base catalysis in outer-sphere phenol oxidation can be insignificant under certain conditions. 26,48,72 2.3.1.3 Spin Trapping Effect Several conventional spin traps 67,73-79 (ilustrated in Scheme 2-1) were investigated for their efects on the kinetics of the phenol/ Ir IV reaction in 0.05 M H + . Scheme 2-1. Structure of spin trapping agents tested. N O PBN N O DMPO N O N O PBN NO MNP H 3 CCH 3 N O DMNBS SO 3 - Br Br N O DBNS - O 3 S 27 Figure 2-5. Trace of the phenol reaction with added DBNBS. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] tot = 0.44 M; [DBNBS] = 10 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -5!10 -03 0 5!10 -03 R e s i d u a l s 0 10 20 30 40 0.0 0.1 0.2 0.3 0.4 Curve fit Reaction trace Time, s A 488 28 We found DBNBS is quite efective in this regard. Figure 2-5 ilustrates the dramaticaly improved fit to a first-order rate law achieved with only 10 mM DBNBS. The other spin traps PBN, DMPO, POBN, and MNP barely afect the reaction, and DMNBS is only partialy efective (Table 2-2). Table 2-2. Kinetic Data for the Reaction of Phenol with Ir IV in the Presence of PBN, DMPO, POBN, MNP, DMNBS and DBNBS. a Spin Trapping Agent t 1/2 , s SD b No Spin Trapping Agent 8.0 0.016 c 1 mM PBN 8.0 0.011 c 1 mM DMPO 8.0 0.016 c 2 mM POBN 8.0 0.012 c 4 mM MNP 8.0 0.006 c 2 mM DMNBS 6.0 0.011 c 1 mM DBNBS 4.4 0.004 a [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] tot = 0.44 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b SD = standard deviation of curve fit residuals over the first 4 half lives. c Succesive half lives increase progresively. These results are not unexpected, given the known specificity of DBNBS for phenoxyl radicals. 79 Tests in 0.05 M H + with concentrations of DBNBS ranging from 0.1 to 15 mM (Table 2-3) show that the efects saturate at about 5 mM DBNBS. A standard DBNBS concentration of 10 mM is used in the experiments described below. The approach to 29 Table 2-3. Kinetic Data for the Reaction of Phenol with Ir IV in the Presence of DBNBS. a [DBNBS] ! 10 3 , M t 1/2 , s Residuals b 0 8.0 0.05 0.10 6.7 0.03 1.0 4.4 0.02 5.0 2.7 0.008 10 2.7 0.006 15 2.5 0.006 a [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] tot = 0.443 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b Maximum residuals of the first-order fits over the first five half-lives. saturation is expected to be dependent on both the eficiency of phenoxyl radical scavenging and the DBNBS dimerization equilibrium (K = 1.3 ! 10 ?3 M). 80 The efects of DBNBS are also significant at p[H + ] 7. The results (Table 2-4) show that DBNBS removes the deviations from pseudo-first-order kinetics that are otherwise sen at this pH. The efect is atributed to DBNBS interception of phenoxyl radicals, the dimerization of which would lead to absorbing intermediates (biphenoquinones), as shown below. 2.3.1.4 Phenol Dependence In the presence of 10 mM DBNBS, the oxidations of phenol by Ir IV in 0.05 M HClO 4 were carried out with phenol concentration (22.2?443) ! 10 ?3 M under pseudo- first-order conditions. Another set of experiments was also performed at p[H + ] = 5.1 in 30 Table 2-4. Kinetic Dependence of Phenol Oxidation on DBNBS at High p[H + ]. a p[H + ] [DBNBS] ! 10 3 , M k obs , s ?1 STDV b 7.11 0 22.4 6.73 ! 10 ?3 7.09 5.0 23.2 1.23 ! 10 ?3 6.94 10 17.3 5.49 ! 10 ?4 a [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] tot = 1.8 ! 10 ?3 M; ? = 0.1 M (LiClO 4 ); The p[H + ] values were maintained using 0.02 M cacodylate buffers; T = 25 ?C. b Standard deviation of first-order curve fit residuals over 1 s. the absence of DBNBS with phenol concentration (1.77?44.3) ! 10 ?3 M. The kinetic data are collected in Table A-1 (in appendix A). The linear plots of k obs versus [phenol] shown in Figure 2-6a-b show that the conditional rate constants are 0.612 ? 0.001 M ?1 s ?1 at p[H + ] = 1.3 and 106 ? 4 M ?1 s ?1 at p[H + ] = 5.1. These results demonstrate that the rate law is first-order with respect to [phenol] tot and the kinetics are sensitive to p[H + ]. 2.3.1.5 Dependence on p[H + ] A simple two-term rate law arises from the asumption that both phenol and the phenolate anion can react with Ir IV , as shown by eqs 2-1 and 2-2: ? d[Ir V ] t = (kArOH+r-Ka/[ + ]) 1 ArOH]to[I V (2-1) obs [r]t (rr-a/[ + ) ] (2-2) 31 Figure 2-6. Plot of k obs vs [phenol] tot . [Ir IV ] 0 = 1 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) [HClO 4 ] = 0.05 M; [phenol] tot = (22.2?443) ! 10 ?3 M; [DBNBS] = 10 mM. (b) p[H + ] = 5.1 (0.02 M acetate buffer); [phenol] tot = (1.77?44.3) ! 10 ?3 M. Solid lines are linear fits. Data from Table A-1. 0.0 0.1 0.2 0.3 0.4 0.5 0.0 0.1 0.2 0.3 p[H + ] = 1.3 a [Phenol] tot , M k obs , s -1 0.00 0.01 0.02 0.03 0.04 0.05 0 1 2 3 4 5 p[H + ] = 5.1 b [Phenol] tot , M k obs , s -1 32 Here, K a is the acid disociation constant of phenol, pK a, ArOH = 9.79 at ? = 0.1 M 81 . k ArOH and k ArO ? represent the reactivities of phenol and phenolate anion. Betwen p[H + ] = 2.4 and 6.8, the oxidation of phenol by Ir IV in the absence of DBNBS gives good-quality pseudo-first-order fits, as shown in Figure 2-3b and Figure 2-3c. This feature enables us to study the p[H + ] efect on the rate constants from p[H + ] = 2.46 to 6.74 with 1 ! 10 ?4 M Ir IV and (4.43?44.3) ! 10 ?3 M phenol. Selected buffers were employed to maintain the p[H + ] values. The data are summarized in Table A-2 and the plot of k obs /[phenol] tot versus p[H + ] is shown in Figure 2-7. A nonlinear least-squares fit of the data to eq 2-2 shows that the rates conform to this two-term rate law with k ArOH = 0.54 ? 0.02 M ?1 s ?1 and k ArO ? = (5.0 ? 0.1) ! 10 6 M ?1 s ?1 . As described above, in the presence of 10 mM DBNBS, the kinetic traces of reactions betwen 1 ! 10 ?4 M of Ir IV and (1.77?44.3) ! 10 ?3 M of phenol obey pseudo- first-order kinetics over the wider p[H + ] range of 1?7. Experiments above pH = 7 were not performed because of the known instability of Ir IV at high pH. 82 The details of these experiments are summarized in Table A-3, and they show that the reaction obeys eq 2-2 (Figure 2-8) with the rate constants about 50% greater than those without the addition of DBNBS: k ArOH = 0.77 ? 0.03 M ?1 s ?1 and k ArO ? = (8.0 ? 0.2) ! 10 6 M ?1 s ?1 . As discussed below, this rate increase induced by DBNBS is atributed to difering stoichiometric factors. 2.3.1.6 Inclusion of the k? Term It has been reported that eq 2-2 is inadequate to describe the reaction of phenol with [Ru(bpy) 3 ] 3+ , and a ?pH-dependent rate constant? (the k? term) was introduced as in eq 2- 3 to fit the data. 50 33 Figure 2-7. Plot of k obs /[phenol] tot vs p[H + ] in the absence of DBNBS. [Ir IV ] 0 = 1 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.6 < p[H + ] < 7.0. Solid line is the fit to eq 2-2 and the dashed line is the fit to eq 2-3. Data from Table A-2. 0 2 4 6 8 10 -1 10 0 10 1 10 2 10 3 10 4 10 5 3-term rate law fit 2-term rate law fit p[H + ] k o b s / [ p h e n o l ] t o t , M ? 1 s ? 1 34 Figure 2-8. Plot of k obs /[phenol] tot vs p[H + ] with DBNBS. [Ir IV ] 0 = 1 ! 10 ?4 M; [DBNBS] = 10 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.4 < p[H + ] < 7.0. Solid line is the fit to eq 2-2 and the dashed line is the fit to eq 2-3. Data from Table A-3. 0 2 4 6 8 10 -1 10 0 10 1 10 2 10 3 10 4 10 5 3-term rate law fit 2-term rate law fit p[H + ] k o b s / [ p h e n o l ] t o t , M ? 1 s ? 1 35 kobs [ArOH]t = ArO+k?10 .5p[H + ] kArO-10 (p[H + ]?Ka) (?Ka) (2-3) When the pH-dependent data for the oxidation by Ir IV with no added DBNBS are fit with eq 2-3, the optimized values are k ArOH = 0.40 ? 0.06 M ?1 s ?1 , k ArO ? = (4.9 ? 0.1) ! 10 6 M ?1 , and k? = (6.5 ? 3) ! 10 ?3 M ?1 s ?1 (Table A-4). This fit yields slightly improved residuals, but the values of k ArOH and k ArO ? are virtualy unchanged. Over the p[H + ] range of 1?7, the maximum contribution of the k? term to the total rate is 16% and occurs around p[H + ] = 2.7. Although this k? term has only marginal statistical significance, we show below that it is a consequence of pH-dependent stoichiometric factors. In the presence of DBNBS, fiting the data to eq 2-3 yields slightly improved residuals relative to eq 2-2 (Table A-5). The fited values for k ArOH and k ArO ? are only slightly changed: k ArOH = 0.59 ? 0.03 M ?1 s ?1 , k ArO ? = (7.3 ? 0.1) ! 10 6 M ?1 and k? = (2.0 ? 0.3) ! 10 ?2 M ?1 s ?1 . The maximum contribution from the k? term is 28% and appears at the same p[H + ] as in the result without DBNBS. The origin of the k? term in the presence of DBNBS is currently unknown but is speculated to arise from a pH-dependent overoxidation of the DBNBS/phenoxyl adduct. 2.3.1.7 Kinetic Isotope Effect (KIE) The deuterium KIE in the reaction betwen phenol and Ir IV was measured by comparing the rates in a D 2 O solution to those in normal H 2 O. The experiments were performed at [H + ] = 0.09 M in the presence of 10 mM DBNBS, and thus the KIE refers to k ArOH . The rates are independent of the pH under these conditions, so the equilibrium 36 isotope efect on K a is not an isue. Al data are shown in Table A-6. The KIE value is 3.5 ? 0.3, which clearly indicates a primary KIE and implies cleavage of the O-H bond in the rate-limiting step. Although this result is taken as evidence of a concerted PCET mechanism (se below), the measured KIE is not extremely large, so it is conceivable that the measured k ArOH value includes a smal contribution from a paralel sequential PCET mechanism. Focusing on the (major) concerted PCET component, the solvent (water) must be acting as the proton aceptor because Ir III is not appreciably basic and no other bases appear in the rate law. A similar KIE and mechanism were recently reported for the oxidation of phenol by three Ru III (bpy) 3 -type complexes 25 and for the oxidation of hydroxylamine by Ir IV . 60 2.3.2 Overoxidation of Phenol 2.3.2.1 Coupling Products As we have described elsewhere, 27 the second step of the reaction betwen phenol and Ir IV is the self-reaction of the phenoxyl radicals. This self-reaction is currently believed to occur through C-C and C-O coupling to generate the four major expected product isomers in Scheme 2-2. Acording to the pulse irradiation experiments, 83 the yields of these products are 25% for 4,4?-biphenol, 18% for 2,2?-biphenol, 44% for 2,4?- biphenol and 9% for 4-phenoxyphenol. The remaining 4% corresponds to unidentified products. The UV-vis spectra of the four coupling isomers in H 2 O or a 0.2% ethanol/water mixture are shown in Figure 2-9. 4,4?-Biphenol has strong absorption at 263 nm with molar absorptivity (& 263 ) of 2.1 ! 10 4 M ?1 cm ?1 . The maximum absorption wavelength is red-shifted to 279 nm when the structure changes to 2,2-biphenol and its & 279 value is 5.9 37 Scheme 2-2. Expected dimerization products of phenoxy radical and their composition acording to literature. 83 C-C ! Coupling C-O Coupling 4,4?-biphenol (25%)  2,2?-biphenol (18%)  2,4?-biphenol (44%) 4-phenoxyphenol (9%) Other products (4%) 2! 38 Figure 2-9. The UV-vis spectra of four coupling isomers. The spectra of 2,2?-biphenol and 4-phenoxyphenol were recorded in water, and those of 4,4?- and 2,4?-biphenol were obtained in 0.2% ethanol-water solution. T = 25 ?C. 200 250 300 350 400 0.0 5.0!10 03 1.0!10 04 1.5!10 04 2.0!10 04 2.5!10 04 4-phenoxyphenol 2,2'-biphenol 4,4'-biphenol 2,4'-biphenol Wavelength, nm " , M - 1 c m - 1 39 ! 10 3 M ?1 cm ?1 . Two absorption bands are observed for 2,4?-biphenol: 250 nm and 283 nm (& 250 = 9.0 ! 10 3 M ?1 cm ?1 and & 283 = 5.2 ! 10 3 M ?1 cm ?1 ). The spectrum of 4- phenoxyphenol exhibits a weaker and broader absorption at 278 nm with & 278 = 2.2 ! 10 3 M ?1 cm ?1 . 2.3.2.2 pK a values of 4,4?-Biphenol and 2,4?-Biphenol Spectrometric titration with NaOH was used to determine the two pK a values for 4,4?-biphenol because controversial results were found previously. 84-86 At 0.1 M ionic strength (LiClO 4 ), 3 ! 10 ?5 M 4,4?-biphenol was titrated from pH = 6.5 to 12.6 with various concentrations of NaOH. The titration was repeated three times and the results are reproducible. UV-vis spectra of one of the titrations are shown in Figure 2-10a. These spectra exhibit a loss of absorbance at 260 nm due to consumption of the biphenol and a gain of absorbance at 288 nm. The titration curve at 288 nm, after volume corrections, is sigmoidal, as shown in Figure 2-10b. Satisfactory fits of this curve to a model involving a single ionizable proton could not be obtained. On the other hand, an excelent fit was obtained with a model that included two ionizable protons, as in eq 2-4: A= !HOAr-c 1+0 (p-Ka2) 1 (pa-H) + !-OAr-c 10 (pKa2H)(pKa1+2-pH) !HOArc 0 (p-Ka1) + (2p-Ka12) (2-4) Equation 2-4 expreses for the total absorption (A) as a function of the pH, pK a1 , and pK a2 , the total biphenol concentration (c), and the molar absorptivities (&) for al three protonated and deprotonated forms. The result of the fit to eq 2-4 is shown in Figure 2-10b, and the derived values are pK a1 = 9.66 ? 0.05 and pK a2 = 10.92 ? 0.4, & HOArArO ? = 40 Figure 2-10. Titration of 3.0 ! 10 ?5 M 4,4?-biphenol. ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) Absorbance spectra of 4,4?-biphenol during titration with various concentrations of NaOH in aqueous solution. pH betwen 6.5 and 12.6. (b) Titration curve at 288 nm after volume correction. Solid line is the fit to eq 2-4. 220 260 300 340 380 0.00 0.16 0.32 0.48 0.64 0.80 pH = 6.5 pH = 12.6 a Wavelength (nm) Abs 6.0 7.4 8.8 10.2 11.6 13.0 0.20 0.34 0.48 0.62 0.76 0.90 pK a1 = 9.66 ? 0.05 pK a2 = 10.92 ? 0.37 b pH A 288nm, corr 41 2.4 ! 10 4 M ?1 cm ?1 , &? OArArO ? = 2.7 ! 10 4 M ?1 cm ?1 and & HOArArOH = 8.0 ! 10 3 M ?1 cm ?1 . This pK a1 value is very close to that reported by Jonsson et al., 85 and the pK a2 value is as expected for succesive pK a s. A similar titration of 2,4?-biphenol was performed three times and the results are reproducible. Figure 2-11a shows a complex series of UV-vis spectra of one of the titrations. The curve at 310 nm (Figure 2-11b) is sigmoidal, but the curve at 285 clearly shows that three species are involved. A fit of the absorbance data at 285 nm to eq 2-4 yields the following values: pK a1 = 9.65 ? 0.04, pK a2 = 10.96 ? 0.06, & HOArArO ? = 8.5 ! 10 3 M ?1 cm ?1 , &? OArArO ? = 6.1 ! 10 3 M ?1 cm ?1 and & HOArArOH = 5.2 ! 10 3 M ?1 cm ?1 . A fit at 310 nm yields pK a1 = 9.68 ? 0.04, pK a2 = 10.89 ? 0.05, & HOArArO ? = 4.8 ! 10 3 M ?1 cm ?1 , &? OArArO ? = 8.8 ! 10 3 M ?1 cm ?1 and & HOArArOH = 4.9 ! 10 2 M ?1 cm ?1 . The agreement betwen the pK a values at these two wavelengths is excelent. The two pK a values for 2,2?-biphenol are quite diferent from those of the 4,4? and 2,4? isomers. 85 This diference is atributed to hydrogen bonding betwen the two O atoms, which is only possible for the 2,2?-isomer. The pK a1 values for al four coupling isomers are summarized in Table 2-5. 2.3.2.3 Overoxidation of Phenol As shown below, Ir IV oxidizes al four of these coupling isomers rapidly, which leads to overoxidation in the phenol reaction. A full acount of the phenol reaction is thus dependent on the details of each of these overoxidation pathways. The Oxidation of 4,4?-Biphenol. Qualitatively, the reaction of Ir IV with 4,4?- biphenol is signaled by the rapid loss of absorbance at 488 nm, which is characteristic of Ir IV . A concurrent absorbance increase occurs at 398 nm, and this absorbance decays at 42 Figure 2-11. Titration of 5.0 ! 10 ?5 M 2,4?-biphenol. ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) Absorbance spectra of 2,4?-biphenol during titration with various concentrations of NaOH in aqueous solution. pH betwen 6.7 and 12.7 (b) Titration curves at 285 nm and 310 nm after volume correction. Solid lines are the fits to eq 2-4. 220 260 300 340 380 0.00 0.13 0.26 0.39 0.52 0.65 pH = 6.7 pH = 12.7 a Wavelength (nm) Abs 6.0 7.6 9.2 10.8 12.4 14.0 0.0 0.1 0.2 0.3 0.4 0.5 285 nm pK a1 = 9.65 ? 0.04 pK a2 = 10.96 ? 0.06 310 nm pK a1 = 9.68 ? 0.04 pK a2 = 10.89 ? 0.05 b pH A corr 43 Table 2-5. Kinetics Data for the Reactions of Ir IV with the Phenol Coupling Products. a Substrate k ArOH , M ?1 s ?1 k ArO ?, M ?1 s ?1 pK a1 E f , V ArO ? /ArO ? E f , V ArO ? , H + /ArOH phenol 0.77 ? 0.03 (8.0 ? 0.2) ! 10 6 9.79 b 0.80 f 1.38 4,4?-biphenol (4.6 ? 0.4) ! 10 4 (6.5 ? 0.7) ! 10 8 9.66 c 0.64 d 1.21 2,2?-biphenol 6.6 ? 2.6 (4.0 ? 0.7) ! 10 6 7.60 d 1.00 d 1.45 2,4?-biphenol (4.0 ? 0.6) ! 10 3 (1.3 ? 0.3) ! 10 8 9.67 c 4-phenoxy phenol (1.2 ? 0.1) ! 10 3 (1.1 ? 0.2) ! 10 8 9.90 e a Ar = 4,4?-/2,2?-/2,4?-HOC 6 H 4 C 6 H 4 , 4-C 6 H 5 OC 6 H 4 . Rate constants obtained by fiting the k obs results in Table A-8 to eq 2-2 and with the K a ?s constrained to the values given above. b Reference 81. c This work. d Reference 85. e The reported pK a1 = 9.81 at ? = 0.25 M in reference 87 is converted to pK a1 = 9.90 at ? = 0.1 M using Davies equation. f E f corrected in this work from E? in reference 88. 44 longer time scales. A UV-vis spectrum of the 398 nm intermediate generated from the reaction of 2.5 ! 10 ?5 M 4,4?-biphenol with 2.5 ! 10 ?5 M Ir IV at p[H + ] = 5.5 is shown in Figure 2-12. This absorbance decays with an initial rate of 1.7 ! 10 ?4 s ?1 as shown in Figure 2-13. When the same experiment is performed at p[H + ] = 2.5, the maximum absorbance at 398 nm and its initial decay rate are close to those obtained at p[H + ] = 5.5, as shown in Table 2-6. The species absorbing at 398 nm is asigned as 4,4?- biphenoquinone, which was claimed to be observed at 400 nm in the oxidation of phenol by other metal complexes. 89-90 Our results imply that the decomposition of 4,4?- biphenoquinone is almost pH-independent. The initial (maximum) absorbance of 4,4?- biphenoquinone at 398 nm yields a molar absorptivity 5.2 ! 10 4 M ?1 cm ?1 in aqueous solution, based on the asumption that 1 mol is produced per 2 mol of Ir IV ; this value for & 398 is consistent with the reported value (% 5 ! 10 4 M ?1 cm ?1 ). 91 In a spectrophotometric titration at p[H + ] = 5.5, 4.8 ! 10 ?5 M Ir IV was added to 1.5 mL of 4.8 ! 10 ?5 M 4,4?-biphenol, generating the titration curves in Figure 2-14. At 398 nm, the absorbance (after correction for dilution) rises linearly to an abrupt end point, and after the end point, the absorbance is stable. The end point corresponds to a consumption ratio of n IrIV /n 4,4?-biphenol of 2.5:1. Apparently, the 4,4?-biphenoquinone product undergoes further oxidation in the presence of 4,4?-biphenol. The absorbance at 488 nm (due to Ir IV ) remains esentialy zero (Figure 2-14b) prior to the same abrupt end point observed at 398 nm, implying that Ir IV is fully consumed up to the end point. The principal reaction is thus 2[IrCl 6 ] 2? + 4,4?-biphenol ! 2[IrCl 6 ] 3? + 4,4?-biphenoquinone + 2H + (2-5) 45 Figure 2-12. UV-Vis spectra of products of 4,4?-/2,2?-/2,4?-biphenol oxidation. [Ir IV ] 0 = 2.5 ! 10 ?5 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. The solid line is the 4,4?-biphenol oxidation: [4,4?-biphenol] 0 = 2.5 ! 10 ?5 M; p[H + ] = 5.5 (0.02 M acetate buffer). The short dashed line was obtained from the 2,2?-biphenol oxidation: [2,2?-biphenol] 0 = 1 ! 10 ?3 M; p[H + ] = 4.0 (0.02 M acetate buffer). The long dashed line results from the 2,4?-biphenol reaction: [2,4?-biphenol] 0 = 2 ! 10 ?4 M; [HClO 4 ] = 0.05 M. The absorption with a maximum peak at approximately 400 nm is asumed to be the spectrum from a product of the oxidation reaction. The molar absorptivity, &, was calculated based on the stoichiometry that one equivalent of biphenoquinone product is obtained from two equivalents of Ir IV . 300 370 440 510 580 650 0 14000 28000 42000 56000 70000 4,4'-biphenol oxidation 2,2'-biphenol oxidation 2,4'-biphenol oxidation Wavelength (nm) ! , M - 1 c m - 1 46 Figure 2-13. The kinetic trace of 2.5 ! 10 ?5 M 4,4?-biphenol reacting with 2.5 ! 10 -5 M Ir IV at 398 nm (solid line). p[H + ] = 5.5 (0.02 M acetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. Slope = ?1.7 ! 10 ?4 s ?1 . Solid line is the experimental trace, and dashed line shows the zero-order fit. 0 100 200 300 400 500 0.50 0.55 0.60 0.65 0.70 Reaction trace Linear regression fit Time, s A 398 47 Table 2-6. Comparison of the Experimental Data of 4,4?- and 2,4?-Biphenol Oxidation at Diferent p[H + ]. a p[H + ] b [biphenol] tot ! 10 4 , M A 398 , max Decay rate, c s ?1 4,4?-biphenol 2.52 0.25 0.640 (1.70 ? 0.02) ! 10 ?4 5.49 0.25 0.653 (1.57 ? 0.01) ! 10 ?4 2,4?-biphenol 1.30 2.0 (7.00 ? 0.12) ! 10 ?3 2.52 0.25 (3.90 ? 0.27) ! 10 ?3 2.51 2.5 (3.55 ? 0.60) ! 10 ?3 5.47 0.25 0.146 (7.35 ? 0.12) ! 10 ?3 5.81 2.5 0.146 (1.19 ? 0.02) ! 10 ?2 6.85 2.0 0.204 d (1.74 ? 0.25) ! 10 ?2d 7.00 0.50 0.113 (1.22 ? 0.04) ! 10 ?2 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; ? = 0.1 M (LiClO 4 ). b p[H + ] = 1.3 is obtained with 0.05 M of HClO 4 . The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for p[H + ] = 2.5, acetate buffer for p[H + ] = 5.5 and 5.8, and cacodylate buffer for p[H + ] = 6.8?7.0. c Initial rates were calculated for 4,4?-biphenol, while first-order rate constants were obtained for 2,4?-biphenol. d Obtained from stopped- flow measurements. 48 Figure 2-14. Titration of 1.5 mL 4,4?-biphenol by Ir IV . [4,4?-biphenol] 0 = 4.8 ! 10 ?5 M; [Ir IV ] 0 = 4.8 ! 10 ?5 M; p[H + ] = 5.5 (0.02 M acetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) Monitored at 488 nm. (b) Absorbance at 488 nm corrected using A cor = A obs *V total /V inital . (c) Monitored at 398 nm. (d) Absorbance at 398 nm corrected using A cor = A obs *V total /V inital . 0 2!10 -07 4!10 -07 6!10 -07 8!10 -07 0.0 0.2 0.4 0.6 0.8 1.0 n Ir(IV) , mol A 398 c 0 2!10 -07 4!10 -07 6!10 -07 8!10 -07 0.00 0.03 0.06 0.09 0.12 0.15 a n Ir(IV) , mol A 488 0 2!10 -07 4!10 -07 6!10 -07 8!10 -07 0 1 2 3 4 5 n Ir(IV) , mol A 398, corr d 0 2!10 -07 4!10 -07 6!10 -07 8!10 -07 0.0 0.3 0.6 0.9 1.2 1.5 n Ir(IV) , mol A 488, corr b 49 Kinetic studies were performed with an exces of 4,4?-biphenol over Ir IV , but because of its low solubility, the initial concentrations of 4,4?-biphenol were not high enough to ensure strictly pseudo-first-order conditions. Thus, values of k obs were obtained from the slope of the linear regresion of the logarithm of absorbance at 488 nm within the first half-life (Figure 2-15). Rates were determined at p[H + ] 1?7, and the data are shown in Table A-8. These pH-dependent data conform to eq 2-2 (Figure 2-16), and they yield the following rate constants: k ArOH = (4.6 ? 0.4) !10 4 M ?1 s ?1 and k ArO ? = (6.5 ? 0.7) ! 10 8 M -1 s ?1 (Table 2-5). The Oxidation of 2,2?-Biphenol. UV-vis spectroscopy was used to detect the products of the reaction betwen 1 ! 10 ?3 M 2,2?-biphenol and 2.5 ! 10 ?5 M Ir IV at p[H + ] = 4.0. As shown in Figure 2-12, a weak absorbance peak appears at 406 nm with a broad shoulder observed in the range of 450?600 nm. This spectrum is atributed to the absorption of 2,2?-biphenoquinone, with & 406 = 1.4 ! 10 4 M ?1 cm ?1 in an aqueous solution. It decomposes slowly as is shown in Figure 2-17. The kinetic experiments on the oxidation of 2,2?-biphenol by Ir IV were performed by using stopped-flow spectroscopy to monitor the loss of Ir IV at 488 nm at p[H + ] 1?7 with at least a 10-fold molar exces of biphenol over Ir IV . However, because of the absorbance of products at 488 nm and their subsequent decay, deviations from pseudo-first-order kinetics were observed at some p[H + ]. Therefore, the initial rate constants were calculated from the slopes of semilog plots of (A ? A & ) vs t. The reaction rates increase dramaticaly as the p[H + ] increases, although the oxidation is extremely slow at low p[H + ], as shown in Figure 2-16. Values for k ArOH and k ArO ? obtained by fiting the data in Table A-8 to eq 2-2 are 6.6 ? 3 M ?1 s ?1 and (4.0 ? 0.7) ! 10 6 M ?1 s ?1 , respectively. 50 Figure 2-15. Kinetic trace of 4,4?-biphenol oxidation at 488 nm. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [4,4?-biphenol] = 1 ! 10 ?4 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. The insert is the linear regresion of log A in the first half time. Slope = ?1.764 ? 0.021 s ?1 (R = 0.9978) and the observed initial rate constant (k obs = ?ln(10) ! slope) is calculated to be 4.06 s ?1 . 0.0 0.6 1.2 1.8 2.4 3.0 0.00 0.03 0.06 0.09 0.12 Time, s A 488 0.00 0.04 0.08 0.12 0.16 -1.3 -1.2 -1.1 -1.0 -0.9 Time, s log A 488 51 Figure 2-16. Plot of k obs /[substrate] tot vs p[H + ]. [Ir IV ] 0 = 2.5 ! 10 ?5 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.6, acetate buffer for 3.7 < p[H + ] < 5.4, and cacodylate buffer for 5.5 < p[H + ] < 7.1. Solid lines are fits to eq 2-2. Data from Table A-8. 0 2 4 6 8 0 1 2 3 4 5 6 7 OHHO OH O p[H + ] l o g ( k o b s / [ A r O H ] t o t , M - 1 s - 1 ) 52 Figure 2-17. The kinetic trace of 1 ! 10 ?3 M 2,2?-biphenol reacting with 2.5 ! 10 ?5 M Ir IV at 406 nm. p[H + ] = 4.0 (0.02 M acetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is the first-order fit. 0 120 240 360 480 600 0.10 0.12 0.14 0.16 0.18 0.20 Time, s A 406 53 The Oxidation of 2,4?-Biphenol. The products of the reaction betwen 2 ! 10 ?4 M 2,4?-biphenol and 2.5 ! 10 ?5 M Ir IV at [HClO 4 ] = 0.05 M were investigated by UV-vis spectroscopy. As is shown in Figure 2-18, an intermediate that has an absorbance maximum at 398 nm is produced rapidly, and then it undergoes decomposition with a first-order rate constant of 7.2 ! 10 ?3 s ?1 (3.5 !10 ?2 s ?1 was obtained by the stopped-flow instrument at 488 nm, as shown in Table A-8). The first UV-vis spectrum obtained after mixing (~ 10 s) is shown in Figure 2-12; it has a main peak at 398 nm and a minor peak at 488 nm (overlapping with that of Ir IV ). Asignment of this spectrum to 2,4?- biphenoquinone is supported by density functional theory calculations at B3LYP/6-31G* level that produce an electronic spectrum with a major peak at ~320 nm and a secondary peak at 460 nm.With the asumption of a 2:1 Ir IV /2,4?-biphenoquinone stoichiometry, the experimental spectrum yields & 398 = 3.0 ! 10 4 M ?1 cm ?1 and & 488 = 4.0 ! 10 3 M ?1 cm ?1 . At p[H + ] = 6.9 and 398 nm, the kinetic traces are triphasic, showing a rapid rise in absorbance, a slower smal-amplitude fal, and then an even slower large-amplitude fal (Figure 2-19). The first phase occurs on the same time frame as the initial absorbance loss at 488 nm (se below). This triphasic behavior is thus atributed to the rapid production of 2,4?-biphenoquinone in the first phase, followed by its biphasic decay. Because of the relatively rapid decay in the second phase, it was not possible to obtain an acurate value for & 398 at this pH, a value of only 1.6 ! 10 4 M ?1 cm ?1 corresponding to the maximum absorbance. At 488 nm the kinetic traces exhibit biphasic decay. Double-exponential fits were used to obtain the fast pseudo-first-order rate constants for the consumption of Ir IV (k obs ) as wel as the slower first-order rate constant for the decay of 2,4?-biphenoquinone (k obs,2 ) 54 Figure 2-18. UV-Vis spectra during the reaction betwen 2 ! 10 ?4 M of 2,4?-biphenol and 2.5 ! 10 ?5 M of Ir IV . [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C; 30 s interval betwen spectra. The inset shows the kinetic trace at 398 nm. k obs = 7.18 ! 10 ?3 s ?1 . 300 360 420 480 540 600 0.0 0.1 0.2 0.3 0.4 0.5 Wavelength, nm Abs 0 120 240 360 480 600 0.0 0.1 0.2 0.3 0.4 Reaction trace Curve fit Time, s A 398 55 Figure 2-19. The kinetic trace of 2 ! 10 ?4 M 2,4?-biphenol reacting with 2.5 ! 10 ?5 M Ir IV at 398 nm. p[H + ] = 6.9 (0.02 M cacodylate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is the experimental trace the dashed line shows the pseudo-first-order fit for the decomposition part. The insert shows the enlarged reaction trace at the beginning 0.5 s. 0 20 40 60 80 100 0.00 0.05 0.10 0.15 0.20 0.25 Time, s A 398 0.0 0.1 0.2 0.3 0.4 0.5 0.00 0.05 0.10 0.15 0.20 0.25 Time, s A 398 56 (shown in Table A-8). The p[H + ] dependence of the Ir IV reduction was studied at p[H + ] 1?7. A fit of the values of k obs to eq 2-2 generated k ArOH = (4.0 ? 0.6) !10 3 M ?1 s ?1 and k ArO ? = (1.3 ? 0.3) ! 10 8 M ?1 s ?1 (Table 2-5). The Oxidation of 4-Phenoxyphenol. Kinetic studies of the reaction of 0.3 mM 4- phenoxyphenol with 2.5 ! 10 ?5 M Ir IV were carried out in the p[H + ] range of 1?7. The decay traces of Ir IV at 488 nm display excelent pseudo-first-order behavior under these conditions (Table A-8). A fit of the data to rate law (2-2) led to the following parameters: k ArOH = (1.2 ? 0.1) ! 10 3 M ?1 s ?1 and k ArO ? = (1.1 ? 0.2) ! 10 8 M ?1 s ?1 (Table 2-5). Evidence of Overoxidation of Phenol. As is shown in Figure 2-20, the reaction betwen 0.0443 M phenol and 1 ! 10 ?4 M Ir IV at p[H + ] = 2.5 displays a loss of absorbance at 488 nm, signaling the consumption of Ir IV , and there is a concurrent rise in absorbance at 398 nm that is asigned to the formation of biphenoquinones arising from oxidation of the biphenol coupling products. Figure 2-21 displays the absorbance at 398 nm on a longer time scale to demonstrate the rise-fal character of the signal. A double- exponential nonlinear regresion fit generates two observed rate constants: the one related to the fast step is equal to 3.6 ! 10 ?2 s ?1 and it corresponds to the consumption of Ir IV at 488 nm, while the one related to the slow step is equal to 3.7 ! 10 ?3 s ?1 which corresponds to the decay of 2,4?-biphenoquinone (Table 2-6). 2.4 Discusion and Conclusion The initial steps in the oxidation of phenol by Ir IV are now well established. 27,51 First, depending on the pH either phenol itself or its conjugate base undergoes reversible one- electron oxidation to produce the phenoxyl radical and [IrCl 6 ] 3? . Second, the phenoxyl 57 Figure 2-20. UV-Vis spectra during the reaction betwen 4.43 ! 10 ?2 M of phenol and 1 ! 10 ?4 M of Ir IV . p[H + ] = 2.5 (0.02 M monochloroacetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C; 5 s interval betwen spectra. The inset shows the kinetic traces at 398 nm (solid line) and 488 nm (dashed line). 300 360 420 480 540 600 0.0 0.2 0.4 0.6 0.8 Wavelength, nm Abs 0 20 40 60 80 0.0 0.1 0.2 0.3 0.4 0.5 398 nm 488 nm Time, s Abs 58 Figure 2-21. The kinetic trace and curve fit of the reaction betwen 4.43 ! 10 ?2 M phenol and 1 ! 10 ?4 M Ir IV at 398 nm. p[H + ] = 2.5 (0.02 M monochloroacetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. k obs, fast = 3.55 ! 10 ?2 s ?1 and k obs, slow = 3.73 ! 10 ?3 s ?1 . Solid line is the experimental trace; dashed line shows the two-exponential fit. 0 60 120 180 240 300 0.00 0.15 0.30 0.45 0.60 Reaction trace Curve fit Time, s A 398 59 radicals undergo bimolecular C?C coupling to produce biphenols and C?O coupling to produce phenoxyphenol. Crucial evidence for this mechanism includes the pH dependence of the rates, the kinetic inhibition by Ir III and its removal by the spin-trap DBNBS, and the simulation of these efects with a kinetic model employing realistic values for the four redox rate constants and the radical coupling overal rate constant. Early on, Cecil and Litler recognized that this model is incomplete because it fails to acount for the observed production of diphenoquinone; 51 they proposed that the initial diphenol coupling products were oxidized further, generating diphenoquinone. Our in- situ UV-vis data provide further evidence for this overoxidation mechanism. Our current studies of the direct oxidations of the four principal phenoxyl coupling products alow quantitative tests of the overoxidation mechanism and provide an explanation of the apparent pH-dependent rate constant (k? in eq 2-3). As shown in Figure 2-16, al four coupling products are oxidized by Ir IV with rates that are highly sensitive to the pH. Moreover, they al (and phenol itself) exhibit two-term rate laws, as given by eq 2-2. A notable diference betwen the oxidations of the biphenols and phenol is that the biphenols undergo net two-electron oxidation to yield quinones, while phenol undergoes a net one-electron oxidation to yield radical coupling products. Evidently, the biphenols achieve their two-electron oxidations through one- electron oxidation and their conjugate bases to form semiquinones, followed by the rapid one-electron conversion of the semiquinones to the quinones. Radical coupling occurs in the case of phenol because the phenoxyl radical is not easily oxidized to its cation. The overal oxidation of 4,4?-biphenol by Ir IV as in eq 2-5 is close to thermoneutral at pH = 1 becuase E?(4,4?-biphenoquinone, 2H + /4,4?-biphenol) = 0.94 V vs NHE 92 while 60 E?(Ir IV /Ir III ) = 0.893 V). 93 This reaction becomes more favorable as the pH increases. In view of the overal stoichiometry (eq 2-5) and the two-term rate law (eq 2-2), a reasonable reaction mechanism is 4,4?-HOC 6 H 4 C 6 H 4 OH ! 4,4?-HOC 6 H 4 C 6 H 4 O ? + H + K a (2-6) Ir IV + 4,4?-HOC 6 H 4 C 6 H 4 OH ! Ir III + 4,4?-HOC 6 H 4 C 6 H 4 O ? + H + K ArOH , k ArOH , k ?ArOH (2-7) Ir IV + 4,4?-HOC 6 H 4 C 6 H 4 O ? ! Ir III + 4,4?-HOC 6 H 4 C 6 H 4 O ? K ArO ?, k ArO ?, k ?ArO ? (2-8) 2[4,4?-HOC 6 H 4 C 6 H 4 O ? ] ! 4,4?-biphenol + 4,4?-biphenoquinone k disp, 4,4 (2-9) Ir IV + 4,4?-HOC 6 H 4 C 6 H 4 O ? ! Ir III + 4,4?-biphenoquinone + H + k semi (2-10) Here, direct oxidation of the biphenol generates the biphenosemiquinone through a concerted electron-proton-transfer mechanism. This step is uphil and, hence, is shown as reversible because E?(4,4?-HOC 6 H 4 C 6 H 4 O ? ,H + /4,4?-HOC 6 H 4 C 6 H 4 OH) = 1.21 V [calculated from E?(ArO ? , H + /ArOH) = E?(ArO ? /ArO ? ) + 0.059 pK a with E?(ArO ? /ArO ? ) = 0.64 V vs NHE 85 ]. Oxidation of the conjugate base of biphenol occurs through simple electron transfer; both of these are in direct analogy with the oxidation of phenol. The fate of the semiquinone could be either disproportionation as in eq 2-9, or further oxidation, as in eq 2-10. Disproportionation was reported previously for the semiquinone in the absence of good oxidants; 89 although the disproportionation rate constant was not determined, it can be expected to be large because the reaction has a large driving force: E?(4,4?-biphenoquinone,2H + /4,4?-biphenol) = 0.94 V vs NHE 92 and E?(4,4?- HOC 6 H 4 C 6 H 4 O ? ,H + /4,4?-HOC 6 H 4 C 6 H 4 OH) = 1.21 V vs NHE. On the other hand, 61 semiquinones are easily oxidized, 94 and the oxidation of 4,4?-biphenosemiquinone is quite favorable: E?(4,4?-biphenoquinone,H + /4,4?-biphenosemiquinone) = 0.67 V (as calculated from the above data). Moreover, the semiquinone is relatively acidic with a pK a of 6.3 for 4,4?-HOC 6 H 4 C 6 H 4 O ? . 85 Clearly, a large rate constant can be anticipated for k semi . The low steady-state concentration of the semiquinone ensures that its reaction with Ir IV wil be dominant. 2,2?-biphenol is more acidic than 4,4?-biphenol by about 2 pK units. This diference is atributed to the formation of an internal hydrogen bond betwen the two O atoms oxygens in the 2,2?-biphenolate monoanion. The value of k ArO? is about a factor of 100 les for 2,2?-biphenol than for 4,4?-biphenol (for reasons discussed below). The result of these two efects is that the rates of oxidation of the two biphenol isomers are quite similar at pH = 7. On the other hand, the value for k ArOH is 10000-fold les for 2,2?- biphenol, so the pH-independent region of the rate law is limited to only quite acidic conditions (Figure 2-16). Despite these quantitative diferences, we infer the same qualitative mechanism for oxidation of the two substrates. Unlike 2,2?-biphenol, 2,4?-biphenol is incapable of internal hydrogen bonding and thus resembles 4,4?-biphenol quite closely in its oxidation by Ir IV . 4-Phenoxyphenol has a pK a value and values for k ArOH and k ArO? that are similar to those of 4,4?-biphenol, and hence its rate-pH profile is similar. Upon oxidation, it yields the 4-phenoxyphenoxyl radical rather than a semiquinone, and hence the ultimate product is due to radical coupling. The above considerations lead to the following mechanism for the overal oxidation of phenol: 62 Ir IV + C 6 H 5 OH ! Ir III + C 6 H 5 O ? + H + K ArOH , k ArOH , k ?ArOH (2-11) Ir IV + C 6 H 5 O ? ! Ir III + C 6 H 5 O ? K ArO ?, k ArO ?, k ?ArO ? (2-12) C 6 H 5 OH ! C 6 H 5 O ? + H + K a (2-13) 2C 6 H 5 O ? ! 4,4?-biphenol k dim,4,4 (2-14-a) 2C 6 H 5 O ? ! 2,2?-biphenol k dim,2,2 (2-14-b) 2C 6 H 5 O ? ! 2,4?-biphenol k dim,2,4 (2-14-c) 2C 6 H 5 O ? ! 4-phenoxyphenol k dim,pop (2-14-d) Ir IV + 4,4?-biphenol ! Ir III + 4,4?-HOC 6 H 4 C 6 H 4 O ? + H + k overox1,4,4 (2-15-a) Ir IV + 2,2?-biphenol ! Ir III + 2,2?-HOC 6 H 4 C 6 H 4 O ? + H + k overox1,2,2 (2-15-b) Ir IV + 2,4?-biphenol ! Ir III + 2,4?-HOC 6 H 4 C 6 H 4 O ? + H + k overox1,2,4 (2-15-c) Ir IV + 4-phenoxyphenol! Ir III + 4-C 6 H 5 OC 6 H 4 O ? + H + k overox1,pop (2-15-d) Ir IV + 4,4?-HOC 6 H 4 C 6 H 4 O ? ! Ir III + 4,4?-biphenoquinone + H + k overox2,4,4 (2-16-a) Ir IV + 2,2?-HOC 6 H 4 C 6 H 4 O ? ! Ir III + 2,2?-biphenoquinone + H + k overox2,2,2 (2-16-b) Ir IV + 2,4?-HOC 6 H 4 C 6 H 4 O ? ! Ir III + 2,4?-biphenoquinone + H + k overox2,2,4 (2-16-c) 2[4-C 6 H 5 OC 6 H 4 O ? ] ! 4-products k dec,pop (2-16-d) The first step in this mechanism corresponds to the reversible concerted proton- electron-transfer oxidation of neutral phenol, and the second step is the reversible outer- sphere electron-transfer from the phenolate anion. The acid/base reaction relating phenol and phenolate (eq 2-13) is asumed to be at equilibrium because of rapid proton transfer. Values for k ArOH and k ArO? are directly measured as described above, while values for K ArOH and K ArO? are calculated from the forward rate constants and the E f values at ? = 63 0.1 M: E f (Ir IV /Ir III ) = 0.893 V, 93 E f (C 6 H 5 O ? /C 6 H 5 O ? ) = 0.80 V, corrected from 0.79 V at ? = 0.0 M 88 by log ' = ?Az i 2 ? 1/2 /(1 + ? 1/2 ), and E f (C 6 H 5 O ? ,H + /C 6 H 5 OH) = 1.38 V [calculated from E f (ArO ? , H + /ArOH) = E f (ArO ? /ArO ? ) + 0.059pK a ]. Dimerization of the phenoxyl radical, which can be partialy rate-limiting under certain conditions, forms the four major coupling isomers: 4,4?-, 2,2?-, and 2,4?-biphenol and 4-phenoxyphenol. Under the assumption that the dimerization rate constant producing each species is proportional to its yield, k dim,4,4 , k dim,2,2 , k dim,2,4 and k dim,pop in eqs 2-14 a-d are 2.88 ! 10 8 M ?1 s ?1 , 2.07 ! 10 8 M ?1 s ?1 , 5.06 ! 10 8 M ?1 s ?1 , and 1.04 ! 10 8 M ?1 s ?1 , respectively, acording to the reported overal 2k dim value 95 and yield of each isomer. 83 Further one-electron oxidations of the phenol coupling products are shown in eqs 2- 15 a-d without reference to the pH dependence evident in Figure 2-16. In our simulations described below, the rate constants for these steps, k overox1,4,4 , k overox1,2,2 , k overox1,2,4 and k overox1,pop , are conditional on the pH and are derived from the parameters in Table 2-5. Eqs 2-16 a-c depict the semiquinones as undergoing oxidation by Ir IV ; in the simulations, we asign rapid rate constants for these oxidations, but the results are insensitive to the exact values used. Eq 2-16-d depicts a rapid second-order decay through dimerization for the phenoxyphenoxyl radical based on analogy with other phenoxyl radicals. Kinetic simulations of the oxidation of phenol based on the above mechanism were performed by the use of the Specfit/32 computer program, 96 with the exact model specified in Table 2-7. These simulations yielded decays of [Ir IV ], the half-lives of which were then used to generate simulated pseudo-first-order rate constants, k obs,sim . Figure 2-22a (data from Table 2-9) shows that these results give an excelent fit to the experimental pH dependence of the reaction. Figure 2-22b compares the simulated results 64 Table 2-7. The Mechanism of Phenol Reaction and the Simulation Model. Equations Kinetic Parameter a Reactions in the model Species eq 2-11 k ArOH = 0.31 k ?ArOH = 5.0 ! 10 7 A + B ! C + D + E k 1 C + D + E ! A + B k ?1 A = Ir IV B = C 6 H 5 OH eq 2-12 k ArO ? = 4.0 ! 10 6 k ?ArO ? = 1.0 ! 10 5 A + F ! C + D k 2 C + D ! A + F k ?2 C = Ir III D = C 6 H 5 O ? eq 2-13 K a = 1.6 ! 10 ?10 B ! E + F k 3 E + F ! B k ?3 E = H + F = C 6 H 5 O ? K = K a /K w B + G ! F + H k 4 F + H ! B + G k ?4 G = OH ? H = H 2 O Buffer K a, buffer I ! J + E k 5 J + E ! I k ?5 I = Buffer J = Conjugate Base K = K a, buffer /K w I + G ! J + H k 6 J + H ! I + G k ?6 eq 2-14-a eq 2-14-b eq 2-14-c eq 2-14-d k dim, 4,4 = 2.9 ! 10 8 k dim, 2,3 = 2.1 ! 10 8 k dim, 2,4 = 5.1 ! 10 8 k dim, pop = 1.0 ! 10 8 2 * D ! K k 7 2 * D ! L k 8 2 * D ! M k 9 2 * D ! N k 10 K = 4,4?-biphenol L = 2,2?-biphenol M = 2,4?-biphenol N = 4-phenoxyphenol eq 2-15-a eq 2-15-b eq 2-15-c eq 2-15-d k overox1, 4,4 : known b k overox1, 2,2 : known b k overox1, 2,4 : known b k overox1, pop : known b A + K ! C + O + E k 11 A + L ! C + P + E k 12 A + M ! C + Q + E k 13 A + N ! C + R + E k 14 O = 4,4?-HOArArO ? P = 2,2?-HOArArO ? Q = 2,4?-HOArArO ? R = 4-C 6 H 5 OC 6 H 4 O ? eq 2-16-a eq 2-16-b eq 2-16-c eq 2-16-d k overox2, 4,4 = 1.0 ! 10 9 k overox2, 2,2 = 1.0 ! 10 9 k overox2, 2,4 = 1.0 ! 10 9 k dec, pop = 1.0 ! 10 9 A + O ! C + S + E k 15 A + P ! C + T + E k 16 A + Q ! C + U + E k 17 2 * R ! V k 18 S = 4,4?-BPQ T = 2,2?-BPQ U = 2,4?-BPQ V = 4-P eq 2-18 k DBNBS = 2.0 ! 10 5 D + W ! X k 19 W = DBNBS eq 2-19 k adduct = 1.0 ! 10 7 A + X ! C + Y k 20 X = Adduct ? eq 2-17 k dim, DBS = 1.3 ! 10 6 k ?dim, DBS = 1.0 ! 10 9 Z ! 2 * W k 21 2 * W ! Z k ?21 Y = Adduct + Z = DBNBS dimer a Rate constants k (M ?1 s ?1 ), acid disociation constant K a nd water disociation constant K w = 1.0 ! 10 ?14 . b See Table 2-8. 65 Table 2-8. The Overoxidation Rate Constants at Diferent p[H + ] Obtained from the Curve Fit of Figure 2-16. p[H + ] k overox1, 4,4 , M ?1 s ?1 k overox1, 2,2 , M ?1 s ?1 k overox1, 2,4 , M ?1 s ?1 k overox1, pop , M ?1 s ?1 1.30 4.6 ! 10 4 8.6 4.0 ! 10 3 1.2 ! 10 3 2.46 4.6 ! 10 4 36 4.0 ! 10 3 1.2 ! 10 3 2.86 4.6 ! 10 4 80 4.0 ! 10 3 1.2 ! 10 3 3.39 4.6 ! 10 4 2.4 ! 10 1 4.0 ! 10 3 1.2 ! 10 3 3.99 4.8 ! 10 4 1.0 ! 10 3 4.2 ! 10 3 1.4 ! 10 3 4.59 5.2 ! 10 4 4.0 ! 10 3 5.1 ! 10 3 1.7 ! 10 3 5.36 7.9 ! 10 4 2.4 ! 10 4 1.0 ! 10 4 4.3 ! 10 3 6.18 2.6 ! 10 5 1.4 ! 10 5 4.5 ! 10 4 2.1 ! 10 4 6.74 8.2 ! 10 5 4.9 ! 10 5 1.5 ! 10 5 7.5 ! 10 4 66 Figure 2-22. Comparative pH dependence of the phenol reaction experimental data and simulation results. (a) The experimental data and simulated results with overoxidation. (b) The simulated results with and without overoxidation. (c) The simulated result with overoxidation and its 3-term rate law curve fit. 2 3 4 5 6 10 -1 10 0 10 1 10 2 10 3 10 4 k exp k sim, overox (a) p[H + ] k o b s / [ p h e n o l ] t o t , M - 1 s - 1 2 3 4 5 6 k sim, overox k sim, no overox (b) p[H + ] 0 2 4 6 8 k sim, overox 3-term rate law fit (c) p[H + ] 67 Table 2-9. Comparison of the Experimental Data with Overoxidation and No Overoxidation Simulation Results of the Phenol Reaction at Diferent p[H + ] in Absence of DBNBS. a p[H + ] [phenol] tot ! 10 3 , M k exp b M ?1 s ?1 k sim, overox c M ?1 s ?1 k sim, no overox d M ?1 s ?1 Stoich. factor e 2.46 44.3 7.26 ! 10 ?1 7.30 ! 10 ?1 4.18 ! 10 ?1 1.75 2.86 44.3 1.15 1.21 7.26 ! 10 ?1 1.67 3.39 44.3 2.71 2.82 1.84 1.54 3.99 44.3 8.87 8.92 6.55 1.36 4.59 44.3 30.4 30.7 25.2 1.22 5.36 44.3 170 164 149 1.10 6.18 4.43 1.24 ! 10 3 1.32 ! 10 3 978 1.35 6.74 4.43 4.67 ! 10 3 4.67 ! 10 3 3.56 ! 10 3 1.31 a [Ir IV ] 0 = 1 ! 10 ?4 M; The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.9 < p[H + ] < 5.4, and cacodylate buffer for 6.1 < p[H + ] < 6.8. b From Table A-2. c Calculated from the half live of overoxidation simulation model with k overox1 and k overox2 ; The series of rate constants k overox1 at diferent p[H + ] are generated from each curve in Figure 2-16. d Calculated from the half live of no overoxidation simulation model (k overox = 0 and k disp = 0). e pH-dependent stoichiometric factor = k sim, overox /k sim, no overox . 68 from the full mechanism with those obtained when overoxidation (eqs 2-15 a-d and 2-16 a-d) is excluded from the mechanism, and it shows that overoxidation increases the net rates at low pH but les so at higher pH. As a result, overoxidation leads to a pH dependence that deviates systematicaly from the simple two-term rate law in eq 2-2. Figure 2-22c shows that the simulations with overoxidation included give an excelent fit to the three-term rate law (eq 2-3), with fited values of k ArOH = 0.38 ? 0.14 M ?1 s ?1 , k ArO ? = (4.9 ? 0.2) ! 10 6 M ?1 s ?1 and k? = (8.2 ? 7) ! 10 ?3 M ?1 s ?1 . Although the statistical uncertainty in k? is large, the value of k? is in good agreement with the value derived from the experimental data. Evidently, the origin of the k? term in the fit of empirical rate law (2-3) to the experimental results is the pH-dependent influence of overoxidation on the rate of consumption of Ir IV . If overoxidation were not occurring, the rate constants in eq 2-2 would agree exactly with those in the mechanism, corresponding to a stoichiometric factor of unity in the rate law. An efect of overoxidation is that the rate constants derived from eq 2-2 are somewhat larger than the rate constants for the elementary steps, which means that overoxidation introduces stoichiometric factors greater than unity. These stoichiometric factors, calculated as the ratio k sim (with overoxidation)/k sim (without overoxidation), decrease systematicaly from 1.75 at pH 2.46 to 1.10 at pH 5.36 (Table 2-9). They also depend somewhat on the phenol concentration. It is the pH dependence of the stoichiometric factor that leads to the k? term in rate law (2-3). Another consequence of the pH-dependent stoichiometric factor is that the yields of the reaction products are also pH-dependent. As shown in Table 2-10, at pH 2.46 the simulated yield of the initial coupling products per 1 mol of Ir IV (corrected for the 1:2 69 Table 2-10. Simulation Product Concentrations of the Phenol Reaction at Diferent p[H + ] Without DBNBS. a Species p[H + ] = 2.46 p[H + ] = 5.36 [Product] ! 10 6 , M % of Ir IVb [Product] ! 10 6 , M % of Ir IVb 4,4?-BP 1.15 ! 10 ?3 9.6% 6.98 72.9% 2,2?-BP 4.67 6.98 2,4?-BP 6.46 ! 10 ?2 18.6 4-POP 0.138 3.95 4,4?-BPQ 7.28 90.4% 4.27 27.1% 2,2?-BPQ 0.582 1.16 2,4?-BPQ 12.8 1.25 4-P 1.26 0.059 a [Ir IV ] 0 = 1 ! 10 ?4 M; p[H + ] = 2.46 was maintained by means of a 0.02 M monochloroacetate buffer; p[H + ] = 5.36 was maintained using a 0.02 M acetate buffer. b Percentage of Ir IV was converted for four products. One equivalent of biphenol product was formed from the reaction of phenol with two equivalents of Ir IV , and one equivalent of biphenoquinone product was obtained from four equivalents of Ir IV . 70 stoichiometric ratio) is 9.6% and the yield of the overoxidation products (corrected for the 1:4 stoichiometric ratio) is 90.4%. However, at pH 5.36 these yields cross over to 72.9% and 27.1%, respectively. In the presence of the spin-trap DBNBS, three more steps are added to the above mechanism: (DBNBS) 2 ! 2DBNBS K dim,DBNBS , k dim,DBNBS , k ?dim,DBNBS (2-17) C 6 H 5 O ? + DBNBS ' adduct ? k DBNBS (2-18) Ir IV + adduct ? ' Ir III + adduct + k adduct (2-19) As mentioned above, only the monomer of DBNBS can scavenge the phenoxyl radical. Therefore, the dimerization step of DBNBS with equilibrium constant K dim,DBNBS of 1.3 ! 10 ?3 M should be included in our mechanism. 80 The phenoxyl radical is scavenged by DBNBS to form an adduct which undergoes a rapid oxidation by Ir IV , as in eqs 2-18 and 2-19. Competition betwen the scavenging of the phenoxyl radical by DBNBS and its dimerization is evident in the kinetic saturation dependence on [DBNBS], as shown in Table 2-3. Simulations of this saturation efect at p[H + ] = 1.3, performed with the mechanism in Table 2-7, are sensitive to the value of k DBNBS . As shown in Table 2-11, a good fit is obtained when the rate constant of C 6 H 5 O ? /DBNBS adduct formation, k DBNBS , is 2.0 ! 10 5 M ?1 s ?1 . This rate constant is large enough to enable 10 mM concentrations of DBNBS to scavenge the phenoxyl radicals completely, preventing their backreaction with Ir III and ensuring that the rate-limiting step in the oxidation by Ir IV is the initial oxidation of phenol. 71 Table 2-11. Comparison of the Experimental Data with Overoxidation Simulation Results of the Phenol Reaction in the Presence of Ir III or Various DBNBS Concentration. a [Ir III ] 0 , M [DBNBS], mM t 1/2, exp , b s t 1/2, sim , c s 0 0 8.0 5.4 0 0.1 6.7 5.0 0 1.0 4.4 3.8 0 5.0 2.7 3.0 0 10 2.7 2.8 0 15 2.5 2.8 5 ! 10 ?4 0 255 271 a [Ir IV ] 0 = 1 ! 10 ?4 M; [phenol] = 0.443 M; [HClO 4 ] = 0.05 M. b From Table 2-2. c From overoxidation simulation model with k overox1 , k overox2 , k DBNBS , k adduct , k dim,DBS and k ?dim,DBS ; the series of rate constants k overox1 at p[H + ] = 1.3 is from Table 2-8. 72 A remaining question is the origin of the pH dependence of the [Ru(bpy) 3 ] 3+ /phenol reaction reported by Sjodin et al., which required the three-term eq 2-3. 50 First, we note that Bonin et al. were unable to reproduce the efect and found the two-term eq 2-2 to be adequate. 25 Second, the degree of overoxidation to be expected depends on the phenol concentration, and this concentration was not disclosed in the original report. Third, the degree of overoxidation should also depend on the rate constants for [Ru(bpy) 3 ] 3+ oxidation of the phenolic coupling products, and these rate constants are unknown. However, it can be shown that a significant degree of overoxidation should occur in acidic media if favorable choices are made for the reactant concentrations and rate constants. Rate Constant Trends. Table 2-5 shows that the values for k ArO? range from 4.0 ! 10 6 to 6.5 ! 10 8 M ?1 s ?1 for the five phenolate ions considered in this study. These rate constants correspond to electron-transfer reactions, and hence it is reasonable to use the cross-relationship of Marcus theory, as given in eqs 1-12 to 1-15 to rationalize their variations. 97 In these equations, k 12 is the cross-electron-transfer rate constant (k ArO? = 8.0 ! 10 6 M ?1 s ?1 for phenoxide), and k 11 and k 22 are the self-exchange rate constants of the C 6 H 5 O ? /C 6 H 5 O ? and Ir IV /Ir III redox couples, respectively. A value for k 22 of 2 ! 10 5 M ?1 s ?1 is used in the calculation, 98 and 1 ! 10 11 M ?1 s ?1 is used for Z, the collision frequency. 99 Z i , Z j are ionic charges of the reactants, R is the ideal gas constant, and r is the center to center distance betwen two reactants when they are approaching to each other. The radii of [IrCl 6 ] 2 and C 6 H 5 O ? are 4.1 ? 58 and 2.5 ?, respectively, estimated from Corey? Pauling?Koltun atomic models. ? is the ionic strength. w ij is the electrostatic energy 73 betwen reactants i and j. If the distance r is in angstroms and ? in molar, then w 12 can be calculated acording to eq 1-15 in kilojoules per mole. With al of these parameters and the experimental values of k 12 and K 12 , k 11 is calculated from the above equations as 2.3 ! 10 6 M ?1 s ?1 . The phenoxide anion and the phenoxyl radical are predicted to have quite similar structures, with the largest diference being a 0.015 ? change in the C?O bond length. 100 Such a smal structural change should not contribute significantly to the self- exchange barrier. Estimates of the solvation contribution to the self-exchange barrier ("G os ? ) are dificult to make because the phenolate charge is highly localized on the oxygen end of the anion; however, this charge localization should cause "G os ? to be larger than that for a spherical anion of comparable size. As a result, the significant overal self-exchange barrier implied by the k 11 above value is atributed principaly to the solvent barrier. A significantly greater value for k 11 of 1.9 ! 10 8 M ?1 s ?1 was previously measured directly by electron spin resonace line broadening. 101 Apparently, the solvent barrier is reduced in the actual self-exchange proces, possibly through a weak asociation betwen the radical and the phenoxide anion. In the case of the 4,4'-biphenoxide anion [E?(ArO ? /ArO ? ) = 0.64 V], 85 k ArO? is considerably larger than that for phenoxide itself. Most of this increase can be ascribed to the greater driving force for the reaction, but there is also some contribution from a greater self-exchange rate constant (k 11 = 4 ! 10 7 M ?1 s ?1 ). A reduced self-exchange barrier can be atributed to the delocalized electronic structure of the semiquinone radical. The 2,2'-biphenoxide anion is considerably more dificult to oxidize [E?(ArO ? /ArO ? ) = 1.00 V] 85 than phenoxide, but the two substrates have quite similar k ArO? values. A large self-exchange rate constant of 3 ! 10 8 M ?1 s ?1 is required by these data. It is 74 conceivable that the internal hydrogen bonding in the 2,2' isomer reduces the degree of hydrogen bonding with the solvent and thus leads to a greater k 11 value. Corresponding discussions of the 2,4'-biphenoxide and 4-phenoxyphenoxide rates wil require determination of the relevant E? values. We have previously argued that the direct oxidation of phenol by Ir IV (eq 2-11, k ArOH ) involves proton transfer to the solvent in concert with electron transfer (H 2 O-CPET). 27 This mechanistic asignment was based largely on the significant solvent deuterium KIE, the high acidity of the ArOH ?+ radical cation, and the low basicity of Ir III . This conclusion is strengthened by the CV measurements described above, which show that Ir III is not significantly protonated in 1 M H + . Further support for this H 2 O-CPET mechanism is provided by a linear free-energy relationship that relates the rates of oxidation of phenol to the E? values for Ir IV and a set of three Ru III oxidants. 25 It sems reasonable to asign a H 2 O-CPET mechanism to al of the phenol reactions in Table 2-5. Recently, Bonin et al. have developed a theoretical treatment of CPET reactions with water (and other bases) as the proton aceptor; 26 this theory predicts, in the absence of other efects, that these reactions should display a typicaly Marcusian dependence of the rates on the driving forces. Qualitatively, Table 2-5 shows that this expectation is met in comparing phenol with 4,4'-biphenol and in comparing 4,4'-biphenol with 2,2'-biphenol. However, 2,2'- biphenol reacts 9 times faster than phenol, even though it is 0.07 V more dificult to oxidize. This apparent contradiction may signal the importance of several other variables in the theory of H 2 O-CPET. 75 Chapter 3 OXIDATION OF ALKYL- AND ALKOXY-SUBSTITUTED PHENOLS BY HEXACHLOROIRIDATE(IV) 3.1 Introduction Substituted phenols have been widely found in nature and have important pharmaceutical, medicinal and chemical applications. 30-37 The mono- or poly-substituents alter the properties of phenol such as electronic efects, steric hindrance, hydrophobicity, hydrophilicity and hydrogen-bonding efect. 102 Among those substituents, the methyl, t- butyl and methoxy groups are the simplest ones that have been involved in research over one century and their properties are wel investigated. The diversities in pK a values and reduction potentials of these phenols generate a wide range of driving forces for their oxidation, and in conjunction with the kinetic data obtained from experiments can elucidate the mechanism of the oxidations. The oxidation of four alkyl-substituted phenols (2-methylphenol, 2,6- dimethylphenol, 2,4,6-trimethylphenol and 4-tert-butylphenol) and one alkoxyl- substituted phenol (4-methoxyphenol) by [IrCl 6 ] 2? are investigated in this chapter. 4- Hydroxymethyl-2,6-dimethylphenol is identified to be one of the products of 2,4,6- trimethylphenol oxidation, which rules out the overoxidation proces. Overoxidation could not occur in the reaction betwen Ir IV and 4-methoxyphenol due to the rapid 76 reaction rate. Two-term and three-term rate laws (applied if overoxidation are observed) are used to obtain the rate constants of oxidation of phenols (k ArOH ), phenolates (k ArO ?), and overoxidations (k?). Mechanisms are proposed for these reactions, and Marcus theory is applied to explain both electron transfer and H 2 O-CPET proceses. 3.2 Experimental Section 3.2.1 Reagents and Solutions Al commercial chemical reagents were used as received except as noted. 2- Methylphenol (cresol), 2,6-dimethylphenol (xylenol), 2,4,6-trimethylphenol (abbreviated as TMP), 4-methoxyphenol (abbreviated as MOP), methoxybenzene (anisole), amonium hexachloroiridate(III) monohydrate (abbreviated as Ir III ), disodium 2,6- pyridinedicarboxylate (abbreviated as dipic), deuterium oxide, sodium acetate anhydrous, cacodylic acid and sodium hydroxide were purchased from Sigma?Aldrich Chemicals Co. 4-tert-Butylphenol (abbreviated as TBP) was received from Fluka. Perchloric acid, amonium perchlorate, sodium chloride, amonium chloride, copper(II) nitrate trihydrate, acetic acid, monochloroacetic acid, formaldehyde solution, toluene, and ethyl acetate were obtained from Fisher Scientific Co. Amonium hexachloroiridate(IV) (abbreviated as Ir IV ) was purchased from Alfa or prepared acording to the literature 61 by addition of amonium chloride (Fisher) to a solution of sodium hexachloroiridate(IV) hexahydrate. Al spin trapping agents utilized in this study were as described in Chapter 2. The structures of reductants used in this work are shown in Scheme 3-1. Eforts to study the oxidation of 4-methylphenol were stymied by our inability to obtain this reagent in sufficient purity. Al atempts yielded a significant contamination 77 Scheme 3-1. Structures of reductants. by 2-methylphenol. Al solutions were freshly prepared with deionized water provided by a Barnstead NANO Pure Infinity ultrapure water system, and purged with argon gas prior to the reactions to prevent potential complications caused by O 2 . The ionic strength was adjusted by lithium perchlorate trihydrate (GFS) and was approximately equal in both oxidants and resultants solutions to prevent Schlieren efects (or refractive index efect 63 ). Selected buffer solutions (acetate, monochloroacetate, and cacodylate buffers) were applied to control the pH if necesary. Preparation of 4-Hydroxymethyl-2,6-dimethylphenol. This compound was prepared as described previously. 103 A 37% formaldehyde solution (0.818 g) was added to a 1 mL toluene solution of 2,6-dimethylphenol (10 mmol). Then, A 48% NaOH solution (0.833 g) was slowly added into the above solution at 10 ?C with stirring. After reacting 20 h at 25 ?C, the solution was poured into 15 mL water and neutralized by acetic acid. The precipitate was collected and recrystalized from ethyl acetate to yield the final product. Yield: 60%. Mp: 104?105 ?C. 1 H NMR (D 2 O): $ 2.22 (s, 6H), 4.48 (s, 2H), 7.04 (s, 2H). OHOH OH OCH 3 OH t-Bu OH 2-Methylpenol (Crsl) 2,6-Dimethylpenol (Xlnol) 2,46-Trimethylpenol (MP) 4-Methoxyphenol (P) 4-ter-tylphenol (TP) 78 3.2.2 Methods A Corning 450 pH/ion meter was used with a Metler Toledo InLab 421 or InLab Semi-Micro-L combination pH electrode. The reference electrode electrolyte was replaced with 3 M NaCl to prevent the formation of KClO 4 precipitate. With the known H + concentration and pH reading, the activity coeficient ! (= 0.839 ? 0.04) was obtained from equation p[H + ] = pH + log !, where p[H + ] is equal to ?log [H + ]. Al measurements were performed at 25.0 ? 0.1 ?C. The kinetics experiments were carried out on a Hi-Tech SF-51 stopped-flow spectrophotometer with OLIS 4300 data acquisition and analysis software. UV-vis spectra were monitored on a HP-8453 diode array spectrophotometer equipped with a Brinkman Lauda RM6 thermostated water bath to maintain the temperature at 25 ?C. Al kinetics data were obtained by monitoring the absorbance of Ir IV at 488 nm. The observed pseudo-first-order rate constants were obtained from fiting kinetic traces over five half lives to first-order exponential functions and each reported observed rate constant is the average of at least 5 shots. The Specfit/32 version 3.0.15 global analysis system was applied to simulate the reaction traces, and the GraphPad Prism 4 or 5 software was used to analyze the rate law with 1/Y 2 weighting. 1 H NMR spectra were acquired on a Bruker AV 400 MHz spectrometer; chemical shifts in D 2 O are relative to DS. The melting points were obtained using an Electrothermal IA 9100 digital melting point apparatus. 3.3 Results 3.3.1 The Oxidation of 2-Methylphenol 79 A kinetic trace for the consumption of 1 ! 10 ?4 M Ir IV in its reaction with 0.01 M cresol under acidic conditions ([HClO 4 ] = 0.05 M), as shown in Figure 3-1, does not yield a good fit to a first-order rate law. Hexachloroiridate(II) Inhibition. We tested the Ir III inhibition efect on the cresol oxidation at diferent p[H + ] (= 1.3, 2.52 and 4.24) with a 5-fold exces of Ir III . A strong inhibiting efect was observed when the reaction was run in 0.05 M HClO 4 with 0.02 M cresol (Figure 3-2a). This inhibition implies an outer-sphere electron transfer mechanism for the oxidation of cresol by Ir IV . As the p[H + ] increased to 2.52, the reaction became faster and the additional Ir III had les of an impact (Figure 3-2b). Subsequently, we detected a smal perturbation with added Ir III at p[H + ] = 4.24 (Figure 3-2c). Spin Trapping Effect. In order to get beter fits to pseudo-first order kinetics and minimize the inhibition by Ir III , the cresol reactions were run in the presence of spin trapping agents. The conventional spin traps ilustrated in Scheme 2-1 were investigated for their efects on the kinetics of the cresol/ Ir IV reaction in 0.05 M H + . The reaction details are summarized in Table 3-1. The oxidation of 0.02 M cresol by 1 ! 10 ?4 M Ir IV with a 5-fold exces of Ir III in presence of 2 mM DBNBS was found to efectively decrease the half-life of the reaction, whereas the other scavengers have litle impact. The reaction betwen 0.01 M cresol and 1 ! 10 ?4 M Ir IV with 1 mM DBNBS yields an excelent pseudo-first-order fit as shown in Figure 3-3. This demonstrates that DBNBS can act as an eficient phenoxyl radical scavenger under these conditions. 2-Methylphenol Dependence. In presence of 1 mM DBNBS, the reactions betwen various concentrations of cresol (0.001?0.02 M) and 1 ! 10 ?4 M of Ir IV were carried out under acidic conditions (0.05 M HClO 4 ). The kinetic data are shown in Table A-7. When 80 Figure 3-1. Kinetic trace of the Ir IV consumption in the cresol oxidation. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 1 ! 10 ?4 M; [cresol] tot = 0.01 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -0.02 0.00 0.02 R e s i d u a l s 0 50 100 150 200 0.0 0.1 0.2 0.3 Reaction trace Curve fit Time, s A 488 81 Figure 3-2. Comparative traces of the cresol reaction with Ir IV (solid line) and in the presence of added Ir III (dashed line) in H 2 O. [Ir IV ] 0 = 1 ! 10 ?4 M; [Ir III ] 0 = 5 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. (a) [Cresol] tot = 0.02 M; [HClO 4 ] = 0.05 M. (b) [Cresol] tot = 0.01 M; p[H + ] = 2.52 (0.02 M monochloroacetate buffer). (c) [Cresol] tot = 0.01 M; p[H + ] = 4.24 (0.02 M acetate buffer). 0 8 16 24 32 Time, s (b) p[H + ] = 2.52 0 2 4 6 8 10 0.0 0.1 0.2 0.3 0.4 0.5 (c) p[H + ] = 4.24 Time, s A 488 0 60 120 180 240 0.0 0.1 0.2 0.3 0.4 0.5 (a) p[H + ] = 1.30 Time, s A 488 82 Table 3-1. Kinetic Data for the Reaction Betwen Cresol and Ir IV with Added Ir III in the Presence of Spin Trapping Agents PBN, DMPO, POBN, MNP, DMNBS and DBNBS. a Spin Trapping Agent t 1/2 , s No Spin Trapping Agent 61.1 2 mM PBN 59.7 2 mM DMPO 62.6 2 mM POBN 61.5 2 mM MNP 59.4 2 mM DMNBS 55.8 2 mM DBNBS 7.20 a [Ir IV ] 0 = 1 ! 10 ?4 M; [cresol] tot = 0.02 M; [Ir III ] 0 = 5 ! 10 ?4 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 83 Figure 3-3. Kinetic traces of the cresol reaction with added DBNBS (solid line) and first order fit (dashed line). [Ir IV ] 0 = 1 ! 10 ?4 M; [cresol] tot = 0.01 M; [DBNBS] = 1 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -6!10 -03 0 6!10 -03 R e s i d u a l s 0 20 40 60 80 100 0.0 0.1 0.2 0.3 Reaction trace Curve fit Time, s A 488 84 plotting k obs versus [cresol] (Figure 3-4), a good straight line is obtained with a slope equal to 14.06 ? 0.04 M ?1 s ?1 and y-intercept equal to (1.02 ? 0.5) ! 10 ?3 s ?1 , which is equal, within error, to zero. The results indicate that the rate of the reaction exhibits a first order dependence on cresol concentration. Dependence on p[H + ]. The influence of p[H + ] on the rate of the reaction was examined in the p[H + ] range of 1?7 with 1 mM DBNBS. Al values of k obs are given in Table A-9. A plot of k obs /[cresol] tot versus p[H + ] reveals that the rates increase regularly with increasing p[H + ]. Generaly, the protonated and deprotonated forms of cresol can react with Ir IV through kineticaly distinguishable terms, k ArOH and k ArO ?, as shown in the two-term rate law (eq 2-2). A nonlinear least-squares fit of k obs as a function of p[H + ] (Figure 3-5) gives k ArOH = 16 ? 1 M ?1 s ?1 and k ArO ? = (5.2 ? 0.2) ! 10 7 M ?1 s ?1 when holding pK a = 10.09 at ? = 0.1 M and 25 ?C. 81 A three-term rate law, eq 2-3, was also applied to the data in Table A-9. The values for k ArOH , k ArO ? and k? are: k ArOH = 13 ? 1 M ?1 s ?1 , k ArO ? = (4.5 ? 0.2) ! 10 7 M ?1 and k? = (0.17 ? 0.04) M ?1 s ?1 , respectively, as displayed in Table A-10. The contribution of the additional k? term was calculated over the p[H + ] range of 1?7. The maximum value is 28% and occurs at around p[H + ] = 3.6. Temperature Dependence. The oxidation of 0.02 M cresol by 2.5 ! 10 ?5 M of Ir IV in 0.01 M HClO 4 was studied at diferent temperatures: 8, 15, 25, 35 and 45 ?C. The values of k obs are summarized in Table A-11. Under these conditions, the k ArOH term dominates the reaction kinetics and k ArO ? is negligible. It is reasonable to asume that k obs represents k ArOH . When plotting ln (k ArOH /T) versus 1/T (Figure 3-6), a linear relationship is obtained with the slope of ?(5.1 ? 0.1) ! 10 3 K and intercept equal to 13.6 ? 0.2. 85 Figure 3-4. Plot of k obs vs [cresol] tot . [Ir IV ] 0 = 1 ! 10 ?4 M; [cresol] tot = (1.0?20) ! 10 ?3 M; [DBNBS] = 1 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is linear fit. Data from Table A-7. 0.0000 0.0044 0.0088 0.0132 0.0176 0.0220 0.00 0.06 0.12 0.18 0.24 0.30 [Cresol] tot , M k obs , s -1 86 Figure 3-5. Plot of k obs /[cresol] tot vs p[H + ] with DBNBS. [Ir IV ] 0 = (2.5?10) ! 10 ?5 M; [cresol] tot = (1.0?20) ! 10 ?3 M; [DBNBS] = 1 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.3 < p[H + ] < 3.5, acetate buffer for 3.7 < p[H + ] < 5.4, and cacodylate buffer for 5.7 < p[H + ] < 7.1. Solid line is the fit to eq 2-2, and the dashed line is the fit to eq 2-3. Data from Table A-9. 0.0 1.6 3.2 4.8 6.4 8.0 10 0 10 1 10 2 10 3 10 4 10 5 10 6 2-term rate law fit 3-term rate law fit p[H + ] k o b s / [ c r e s o l ] t o t , M - 1 s - 1 87 Figure 3-6. Plot of log (k obs /T) vs 1/T. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [cresol] tot = 0.02 M; [HClO 4 ] = 0.01 M; ? = 0.1 M (LiClO 4 ). Solid line is the fit to eq 3-1. Data from Table A-11. 3.0!10 -03 3.1!10 -03 3.3!10 -03 3.4!10 -03 3.6!10 -03 3.7!10 -03 -5.0 -4.3 -3.6 -2.9 -2.2 -1.5 1/T, K -1 ln (k ArOH /T) 88 Acording to Eyring?s transition state theory eq 3-1, ln k ArOH T =ln k B h + !S ? R ? !H ? RT ! (3-1) the activation parameters can be calculated as "H ? = 42.0 ? 1 kJ mol ?1 and "S ? = ?84.0 ? 2 J mol ?1 K ?1 after alowance for ln (k B /h) = 23.8. Here k B is Boltzmann?s constant, h is Planck?s constant, and R is the gas constant. Kinetic Isotope Effect. The deuterium kinetic isotope efect of the reaction betwen 0.02 M of cresol and 1 ! 10 ?4 M of Ir IV was studied under two conditions, p[H + ] = 1.30 and 2.34. Acording to the p[H + ] dependence plot in Figure 3-5, both p[H + ]s are on the plateau and the equilibrium isotope efects on K a is not an isue. The observed rate constants measured in presence of 1 mM DBNBS in D 2 O and normal H 2 O are given in Table A-12. Both results yield the KIE of 2.9 ? 0.2. Products. Although no reported # max values are found for two methyl groups substituted 4,4?-biphenoquinones, they are expected to appear in the wavelength range of 398 nm (# max of 4,4?-biphenoquinone) to 421 nm (# max of 3,3?,5,5?-tetramethyldiphenoqu- inone 104 ). Acording to our UV-vis spectra, a weak absorption at 412 nm was detected at the end of the cresol and Ir IV reaction, which is indicative of the corresponding biphenoquinones and overoxidation takes place in cresol reaction. 3.3.2 The Oxidation of 2,6-Dimethylphenol A kinetic experiment for the reaction of 1 ! 10 ?4 M Ir IV and 0.025 M xylenol was carried out under acidic conditions ([HClO 4 ] = 0.05 M). The trace of Ir IV consumption 89 Figure 3-7. Trace of the xylenol reaction (solid line) and first-order fit (dashed line). [Ir IV ] 0 = 1 ! 10 ?4 M; [xylenol] tot = 0.025 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -3!10 -03 0 3!10 -03 R e s i d u a l s 0 2 4 6 8 0.0 0.1 0.2 0.3 0.4 Curve fit Reaction trace Time, s A 488 90 Figure 3-8. Comparative traces of the xylenol reaction (solid line) with added 5 ! 10 ?4 M Ir III (long dashed line) and 2.5 ! 10 ?3 M Ir III (long dashed line). [Ir IV ] 0 = 1 ! 10 ?4 M; [xylenol] tot = 0.025 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 0 4 8 12 16 20 0.0 0.1 0.2 0.3 0.4 0.5 0.6 with added 5-fold Ir III no added Ir III with added 25-fold Ir III Time, s A 488 91 can be wel fit to pseudo-first-order kinetics (Figure 3-7) without the need for adding spin trapping agent. Hexachloroiridate(II) Inhibition. The inhibiting effect of Ir III was examined with 1 ! 10 ?4 M Ir IV and 0.025 M xylenol in 0.05 M HClO 4 in the presence of 5 and 25-fold exces of Ir III . A weak inhibition was observed from Figure 3-8 with 5 ! 10 ?4 M Ir III , which was enhanced by increasing the concentration of the added Ir III to 2.5 ! 10 ?3 M. For our general kinetics studies no additional Ir III is present and the concentration of Ir IV is always below 1 ! 10 ?4 M. Under these conditions Ir III inhibition can be ignored. Copper Catalysis. Our previous studies on outer-sphere electron-transfer oxidation show that trace copper ions can catalyze the reaction. 60,73,76,105 The efect of copper catalysis was tested for the oxidation of xylenol in acidic aqueous solution, and the results are shown in Table 3-2. Table 3-2. Kinetic Efect of Cu 2+ and Dipic for Xylenol Oxidation. a Cu 2+ and Dipic k obs /[xylenol] tot , M ?1 s ?1 35.8 5 ! 10 ?6 M Cu 2+ 35.9 1 ! 10 ?3 M dipic 35.7 a Cu 2+ added as Cu(NO 3 ) 2 , [Ir IV ] 0 = 1 ! 10 ?4 M; [xylenol] tot = 0.025 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. In presence of 5 ! 10 ?6 M Cu(NO 3 ) 2 , the observed rate constant for the oxidation of 0.025 M xylenol by 1 ! 10 ?4 M Ir IV remains the same within error. Moreover, acording to our 92 previous reported results on the oxidation of thioglycolic acid and cysteine, 77,73 2,6- pyridinedicarboxylic acid (dipic) efectively inhibits copper catalysis. Consequently, 1 ! 10 ?3 dipic was added to the reaction betwen 0.025 M of xylenol and 1 ! 10 ?4 M of Ir IV in 0.05 M HClO 4 . No effect observed with the addition of dipic, which confirms that there is no copper catalysis in the oxidation of xylenol. 2,6-Dimethylphenol Dependence. Under pseudo-first-order condition, 1 ! 10 ?4 M Ir IV oxidized various concentration of xylenol (0.0025 ? 0.025 M) in 0.05 M HClO 4 . The observed rate constants of the reaction are summarized in Table A-13. The plot of k obs versus [xylenol] tot shown in Figure 3-9 exhibits a straight line with slope = 35.5 ? 1.4 M ?1 s ?1 and a negligible intercept. The linear relationship indicates that the rate law is first- order with respect to [xylenol] tot . Dependence on p[H + ]. A study of the p[H + ] dependence of the kinetics was carried out with various xylenol and Ir IV concentrations ([xylenol] tot = (0.25?5) ! 10 ?3 M; [Ir IV ] = (0.25?1) ! 10 ?4 M) as ilustrated in Table A-14 under pseudo-first-order condition over the p[H + ] range of 1?7. The data set follows eq 2-2, and the nonlinear least-squares fit of k obs /[xylenol] tot versus p[H + ] plot shown in Figure 3-10 yields k ArOH = 47 ? 2 M ?1 s ?1 and k ArO ? = (2.0 ? 0.1) ! 10 8 M ?1 s ?1 when holding pK a = 10.38 at ? = 0.1 M (LiClO 4 ) and 25 ?C. This pK a is obtained by using Davies equation (log ' = ?Az i 2 ? 1/2 /(1 + ? 1/2 )) to convert the reported value, 10.62 at ? = 0 and 25 ?C 81 , into pK a at ? = 0.1 M. Figure 3-10 also displays the nonlinear fit of the data to the 3-term rate law with k ArOH , k ArO ? and k? as 38 ? 1 M ?1 s ?1 , k ArO ? = (1.71 ? 0.03) ! 10 8 M ?1 and k? = (0.40 ? 0.04) M ?1 s ?1 , respectively. These results are summarized in Table A-15. The maximum contribution of the additional k? term is 28% (p[H + ] = 3.7). 93 Figure 3-9. Plot of k obs vs [xylenol] tot . [Ir IV ] 0 = 1 ! 10 ?4 M; [xylenol] tot = (2.5?25) ! 10 ?3 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is linear fit. Data from Table A-13. 0.000 0.006 0.012 0.018 0.024 0.030 0.0 0.2 0.4 0.6 0.8 1.0 [Xylenol] tot , M k obs , s -1 94 Figure 3-10. Plot of k obs /[xylenol] tot vs p[H + ]. [Ir IV ] 0 = (2.5?10) ! 10 ?5 M; [xylenol] tot = (2.5?50) ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.5 < p[H + ] < 5.4, and cacodylate buffer for 5.5 < p[H + ] < 7.0. Solid line is the fit to eq 2-2 and the dashed line is the fit to eq 2-3. Data from Table A-14. 0.0 1.6 3.2 4.8 6.4 8.0 10 1 10 2 10 3 10 4 10 5 2-term rate law fit 3-term rate law fit p[H + ] k o b s / [ x y l e n o l ] t o t , M - 1 s - 1 95 Kinetic Isotope Effect. The kinetic isotope efect was measured with 0.025 M xylenol and 1 ! 10 ?4 M Ir IV at high acidity ([HClO 4 ] = 0.05 M) where the rate constants are p[H + ] independent. The experimental details are shown in Table A-16. The ratio of observed rate constants in H 2 O and D 2 O are obtained as 2.6 ? 0.1. Products. The UV-vis spectrum of the xylenol and Ir IV mixture at the end of the reaction exhibited an extremely strong absorption at 421 nm which is consistent with the formation of 3,3?,5,5?-tetramethyldiphenoquinone as a product acording to literature. 104 3.3.3 The Oxidation of 2,4,6-Trimethylphenol (TMP) A typical kinetic decay of 1 ! 10 ?4 M Ir IV in the reaction of TMP is shown in Figure 3-11. In 0.05 M HClO 4 solution and with only 1.2 ! 10 ?3 M TMP present, the reaction was stil fast compared to the oxidation of cresol and xylenol at the same p[H + ] but with higher concentration of reductants. A good-quality pseudo-first order trace was observed and is exhibited in Figure 3-11. Therefore, we expected that Ir III inhibition would not be a problem for the TMP reaction. Hexachloroiridate(II) Inhibition. In order to check for a possible Ir III efect, the oxidation of 1.2 ! 10 ?3 M TMP by 1 ! 10 ?4 M Ir IV in presence of Ir III was examined under acidic conditions ([HClO 4 ] = 0.05 M), as shown in Figure 3-12. With high amount of exces Ir III (= 2.5 ! 10 ?3 M) we detected a relatively smal inhibition phenomenon. This proved that the degree of kinetic inhibition by Ir III is considerably weak and that the reaction could be studied without a spin trapping agent. 2,4,6-Trimethylphenol Dependence. Under acidic conditions (0.01 M HClO 4 ), various concentrations of TMP were oxidized by 1 ! 10 ?4 M Ir IV , and the results are presented in Table A-17. A plot of k obs versus [TMP] tot in the range of (0.3?3) ! 10 ?3 M 96 Figure 3-11. Kinetic trace of the Ir IV consumption in the TMP oxidation. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 1 ! 10 ?4 M; [TMP] tot = 1.2 ! 10 ?3 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -4!10 -03 0 4!10 -03 R e s i d u a l s 0 2 4 6 8 10 0.0 0.1 0.2 0.3 0.4 Reaction trace Curve fit Time, s A 488 97 Figure 3-12. Comparative traces of the TMP reaction (solid line) with added Ir III (dashed line). [Ir IV ] 0 = 1 ! 10 ?4 M; [TMP] tot = 1.2 ! 10 ?3 M; [Ir III ] 0 = 2.5 ! 10 ?3 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 0 5 10 15 20 25 0.00 0.12 0.24 0.36 0.48 0.60 with added 25-fold Ir III no added Ir III Time, s A 488 98 TMP (Figure 3-13) is linear with a slope = (1.09 ? 0.02) ! 10 3 M ?1 s ?1 and a negligible intercept, indicating a first-order dependence of the rate on the TMP concentration. Dependence on p[H + ]. The influence of p[H + ] on the reaction rate was investigated with 2.5 ! 10 ?5 M of Ir IV and 3 ! 10 ?4 M of TMP over the p[H + ] range of 1?7. The experimental data are summarized in Table A-18, and a plot of k obs /[TMP] tot versus p[H + ] is shown in Figure 3-14. The nonlinear-least-square fit to the rate law eq 2-2 yields k ArOH = (1.09 ? 0.01) ! 10 3 M ?1 s ?1 and k ArO ? = (2.58 ? 0.03) ! 10 8 M ?1 s ?1 with pK a = 10.65 at ? = 0.1 M and 25 ?C (=10.89 at ? = 0 and 25 ?C 81 ). Since an excelent 2-term rate law fit is obtained for TMP reaction with R square of 0.9998, data need not be fited to the 3- term rate law. Kinetic Isotope Effect. The oxidations of 1.5 ! 10 ?3 M and 3 ! 10 ?3 M of TMP by 1 ! 10 ?4 M of Ir IV were detected in H 2 O and D 2 O at high acidity ([HClO 4 ] = 0.01 M). Under this condition, the reaction is pseudo-first order and p[H + ] independent. Al the data are shown in Table A-19, and the ratio of observed rate constants in H 2 O and D 2 O is obtained as 2.04 ? 0.01. Products Identification. The UV-vis spectrum of TMP and Ir IV mixture at the end of the reaction did not show a detectable absorption around 400 nm, which is quite diferent from the other phenols. We studied the reaction products by using 1 H NMR spectroscopy. The spectrum of TMP exhibits three distinct peaks as shown in Figure 3-15a: two peaks in the high-field region ($ = 2.18 and 2.19 ppm) correspond to the protons on methyl groups and one peak in the low-field region ($ = 6.90 ppm) represents the two equivalent aromatic protons. The 1 H NMR spectrum of the reaction solution was detected by mixing 1 ! 10 ?3 M of Ir IV and 3 ! 10 ?3 M of TMP in D 2 O, as shown in 99 Figure 3-13. Plot of k obs vs [TMP] tot . [Ir IV ] 0 = 2.5 ! 10 ?5 M; [TMP] tot = (0.3?3.0) ! 10 ?3 M; [HClO 4 ] = 0.01 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is linear fit. Data from Table A-17. 0.0000 0.0008 0.0016 0.0024 0.0032 0.0040 0.0 0.8 1.6 2.4 3.2 4.0 [TMP] tot , M k obs , s ?1 100 Figure 3-14. Plot of k obs /[TMP] tot vs p[H + ]. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [TMP] tot = 3.0 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.5, acetate buffer for 3.6 < p[H + ] < 5.5, and cacodylate buffer for 5.6 < p[H + ] < 7.0. Solid line is the fit to 2-term rate law eq 2-2. Data from Table A-18. 0.0 1.6 3.2 4.8 6.4 8.0 10 2 10 3 10 4 10 5 10 6 p[H + ] k o b s / [ T M P ] t o t , M - 1 s - 1 101 Figure 3-15. 1 H NMR identification of TMP oxidation Products. (a) The 1 H NMR spectrum of TMP in D 2 O. (b) The 1 H NMR spectrum of 4-hydroxymethyl-2,6- dimethylphenol in D 2 O. (c) A sample of the reaction solution with 3 mM TMP oxidized by 1 mM Ir IV in D 2 O. (d) Spiked reaction solution in c with 4-hydroxymethyl-2,6- dimethylphenol. !" #" $" %" OH C 3 CH 3 H 3 C OH C 3 H 2 C H 3 C OH [IrCl 6] 2+ : TMP = 1 : 3! !"#$%& OH CH 3 CH 3 H 3 C OH C 3 H 2 C H 3 C OH Ir IV :TMP=1:3 Spike (a) (b) (c) (d) !" #" !" & # " & $ " $"$" %" '" %" ("(" & ! " & ' " & % " & ( " & # "& ! " & % " & ' " & $ " & ( " & # "& ! " & % " & ' " & $ " & ( " 102 Figure 3-15c. Comparing to the spectrum of TMP, three additional peaks ($ = 2.22, 4.48 and 7.04 ppm) were observed for the reaction solution, which must be the product of the oxidation of TMP. After analyzing the chemical shifts and integrals, we predicted that the product is 4-hydroxymethyl-2,6-dimethylphenol, which was confirmed via comparison to an independently prepared standard (its 1 H NMR spectrum is shown in Figure 3-15b). Spiking the reaction solution with 4-hydroxymethyl-2,6-dimethylphenol increased the intensity of the three resonances, as shown in Figure 3-15d, further bolstering the identification. Additional information was also obtained from the integrals. The diference in the integrals for the rest of TMP and the product 4-hydroxymethyl-2,6- dimethylphenol (2.22 ppm) reveals that the stoichiometry of Ir IV : TMP is equal to 2:1 if al Ir IV was reduced. Overoxidation. In order to get insight into the overoxidation of TMP, the oxidation of 4-hydroxymethyl-2,6-dimethylphenol was examined with 1 ! 10 ?4 M Ir IV and 1 ! 10 ?3 M 4-hydroxymethyl-2,6-dimethylphenol at p[H + ] = 1.3 and 4.4. The values of k obs /[4- hydroxymethyl-2,6-dimethylphenol] tot are 193 and 505 M ?1 s ?1 , respectively. In comparison with the values of k obs /[TMP] tot , 1.06 ! 10 3 M ?1 s ?1 at p[H + ] = 1.3, and 1.31!10 3 M ?1 s ?1 at p[H + ] = 4.5 (shown in Table A-18), we found that the oxidation of TMP is 2?5 times faster than that of 4-hydroxymethyl-2,6-dimethylphenol in the p[H + ] range of 1.3?4.5. Besides, the high concentration of TMP enables it to be more rapidly oxidized by Ir IV than 4-hydroxymethyl-2,6-dimethylphenol, thereby limiting further oxidation. 3.3.4 The Oxidation of 4-Methoxyphenol A typical kinetic trace of Ir IV reduction in the MOP oxidation is shown in 103 Figure 3-16. The reaction was carried out with 3 ! 10 ?4 M of MOP and 2.5 ! 10 ?5 M of Ir IV and can be wel fited to a pseudo-first order curve with smal deviations, which means that under these systems the Ir III inhibition is negligible and no trapping agent is required for the kinetic studies. Dependence on p[H + ]. The p[H + ] dependence was studied with 2.5 ! 10 ?5 M of Ir IV and 3 ! 10 ?4 M of MOP over the p[H + ] range of 1?7. The experimental data are summarized in Table A-21, and a plot of k obs /[MOP] tot versus p[H + ] is shown in Figure 3-17. When fiting the data to the rate law eq 2-2, it yields k ArOH = (6.1 ? 0.1) ! 10 4 M ?1 s ?1 and k ArO ? = (3.3 ? 0.1) ! 10 8 M ?1 s ?1 with pK a = 9.96 at ? = 0.1 M and 25 ?C. 81 Kinetic Isotope Effect. The deuterium kinetic isotope efect of reaction betwen 2.5 ! 10 ?5 M of Ir IV and 3 ! 10 ?4 M of MOP was determined in H 2 O and D 2 O at high acidity ([HClO 4 ] = 0.05 M). Under this condition, the reaction is pseudo-first order and pH-independent. Al the data are shown in Table A-20 and the ratio of observed rate constants in H 2 O and D 2 O are obtained as 1.9 ? 0.1. Temperature Dependence. The reaction betwen 3 ! 10 ?4 M of MOP and 2.5 ! 10 ? 5 M of Ir IV in 0.1 M HClO 4 was carried out at diferent temperatures: 8, 15, 25, 35 and 45 ?C. The values of k obs are summarized in Table A-22. Under these acidic conditions, it is reasonable to asume that k obs represents k ArOH which dominates the reaction kinetics. A plot of ln (k ArOH /T) versus 1/T in Figure 3-18 shows a linear relationship with the slope of ?(2.7 ? 0.2) ! 10 3 K and intercept equal to 14.4 ? 1. Acording to the Eyring equation (3-1), the activation parameters can be calculated as "H ? = 22.8 ? 1 kJ mol ?1 and "S ? = ? 78.0 ? 5 J mol ?1 K ?1 using ln (k B /h) = 23.8. 104 Figure 3-16. Kinetic trace of the Ir IV consumption in the MOP oxidation. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -4!10 -03 0 4!10 -03 R e s i d u a l s 0.0 0.2 0.4 0.6 0.8 1.0 0.00 0.03 0.06 0.09 0.12 Reaction trace Curve fit Time, s A 488 105 Figure 3-17. Plot of k obs /[MOP] tot vs p[H + ]. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.5, and cacodylate buffer for 5.6 < p[H + ] < 7.0. Solid line is the fit to 2-term rate law eq 2-2. Table A-21. 0.0 1.6 3.2 4.8 6.4 8.0 10 4 10 4.5 10 5 10 5.5 10 6 p[H + ] k o b s / [ M O P ] t o t , M - 1 s - 1 106 Figure 3-18. Plots of ln (k obs /T) vs 1/T. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; [HClO 4 ] = 0.1 M; T = 25 ?C. Solid line is the fit to eq 3-1. Data from Table A-22. 0.00300 0.00317 0.00334 0.00351 0.00368 4.2 4.6 5.0 5.4 5.8 6.2 1/T, K ?1 ln (k ArOH /T) 107 Control Experiments. In order to gain insight into the reactivity of the methoxy group of MOP, the reactions betwen 0.01 M of anisole and 1.0 ! 10 ?4 M of Ir IV were tested at 0.1 M ionic strength in 0.05 M HClO 4 and at p[H + ] = 7.0. No oxidation was observed under both conditions suggesting that the methoxy group is unreactive, which confirms that only the hydroxyl group on MOP participates in the reaction with Ir IV . 3.3.5 The oxidation of 4-tert-Butylphenol The oxidation of 2 ! 10 ?3 M of TBP by 2.5 ! 10 ?5 M of Ir IV was studied in 0.05 M HClO 4 , and the decay of Ir IV at 488 nm is presented in Figure 3-19. When fiting the kinetic trace to a pseudo-first order rate law, we did not obtain a satisfactory result. In order to improve the kinetic behavior, we added 1 mM of DBNBS into the above reaction and obtained a good-quality first-order fit, as shown in Figure 3-20. Dependence on p[H + ]. The p[H + ] influence on the reaction rate was studied with 2.5 ! 10 ?5 M of Ir IV and (3.0?9.0) ! 10 ?4 M of TBP over the p[H + ] range of 1?7. The experiments are carried out in the presence of 1 mM DBNBS and the data are ilustrated in Table A-23. The plot of k obs /[TBP] tot versus p[H + ] in Figure 3-21 is fited to the rate law eq 2-2, and yields k ArOH = 29 ? 3 M ?1 s ?1 and k ArO ? = (3.6 ? 0.3) ! 10 7 M ?1 s ?1 with pK a = 10.07 at ? = 0.1 M and 25 ?C (pK a = 10.31 at ? = 0 and 25 ?C 81 ). When fiting the data in Table A-23 to the 3-term rate law eq 2-3, as shown in Figure 3-21, the following values for k ArOH , k ArO ? and k? are obtained: k ArOH = 18 ? 1 M ?1 s ?1 , k ArO ? = (2.4 ? 0.1) ! 10 7 M ?1 and k? = (0.48 ? 0.03) M ?1 s ?1 , as displayed in Table A-24. The contributions of the additional k? term were calculated over the p[H + ] range of 1?7. The maximum value is 56% and occurs around p[H + ] = 3.9. 108 Figure 3-19. Kinetic trace of the Ir IV consumption in the TBP oxidation. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [TBP] tot = 2 ! 10 ?3 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -0.06 0.00 0.06 R e s i d u a l s 0 120 240 360 480 600 0.00 0.03 0.06 0.09 0.12 Reaction trace Curve fit Time, s A 488 109 Figure 3-20. Kinetic trace of the Ir IV consumption in the TBP oxidation with DBNBS. Lower box shows the experimental trace (solid line) and the pseudo-first-order fit (dashed line). Upper box shows the residuals in the fit. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [TBP] tot = 2 ! 10 ?3 M; [DBNBS] = 1 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -2!10 -03 0 2!10 -03 R e s i d u a l s 0 60 120 180 240 300 0.00 0.03 0.06 0.09 0.12 Curve fit Reaction trace Time, s A 488 110 Figure 3-21. Plot of k obs /[TBP] tot vs p[H + ]. [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 1 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.3 < p[H + ] < 3.4, acetate buffer for 3.5 < p[H + ] < 5.3, and cacodylate buffer for 5.5 < p[H + ] < 7.1. Solid line is the fit to eq 2-2 and the dashed line is the fit to eq 2-3. Data from Table A-23. 0.0 1.6 3.2 4.8 6.4 8.0 10 0 10 1 10 2 10 3 10 4 10 5 10 6 p[H + ] k o b s / [ T B P ] t o t , M - 1 s - 1 2-term rate law fit 3-term rate law fit 111 3.4 Discusion and Conclusion 3.4.1 Mechanism As we reported in the previous work, 106 a similar mechanisms are proposed for ortho-substituted phenols, cresol and xylenol: first, the reactions betwen Ir IV and neutral phenols undergo a concerted proton-coupled electron-transfer oxidation to produce the corresponding phenoxyl radical, Ir III and a proton, as shown in eq 3-2. K ArOH , k ArOH and k ? ArOH are the equilibrium constant and rate constants for this step. The oxidation of the conjugate base phenolate anions is a reversible outer-sphere one-electron transfer proces (eq 3-3), forming a phenoxyl radical. K ArO ?, k ArO ? and k ?ArO ? represent the equilibrium constant and second-order rate constants of this pathway. Second, the phenoxyl radicals undergo bimolecular C?C coupling to produce biphenols and C?O coupling to produce phenoxyphenol (eq 3-5), which is subsequently oxidized by Ir IV to form biphenosemiquinones or phenoxyphenoxyl radicals (eq 3-6). k dim corresponds to the coupling rate constant and k overox,1 is the rate constant of further overoxidation. Lastly, biphenosemiquinones are easily oxidized by Ir IV to form biphenoquinones with rate constant of k overox, 2 , as ilustrated in eq 3-7. Phenoxyphenoxyl radicals can dimerize with rate constant k dec acording to eq 3-8. Eqs 3-2 to 3-5 are also proposed for the oxidation of para-substituted phenol, TBP, which only forms the corresponding 2,2? product. Further oxidations by Ir IV are shown in eqs 3-6 and 3-7. Because no overoxidation was detected from our experiments, the oxidation of another para-substituted phenol, MOP, only undergoes steps 3-2 to 3-5. For the TMP reaction, the ortho- and para- positions of TMP are al occupied with methyl groups, therefore, dimerizations through C-C and C-O couplings of the 112 corresponding phenoxyl radical generated from eqs 3-2 and 3-3 are impossible. The disproportionation pathway undergoes, as depicted in eq 3-9, based on the observation of 4-hydroxymethyl-2,6-dimethylphenol as the oxidation product, and k disp is given for the rate constant of this step. Compared with the TMP oxidation, the reaction betwen Ir IV and 4-hydroxymethyl-2,6-dimethylphenol is 2?5 times slower. Therefore, the overoxidation proces does not occur with TMP. K ArOH , k ArOH , k ?ArOH (3-2) K ArO ?, k ArO ?, k ?ArO ? (3-3) K a (3-4) k dim (3-5) Ir IV + Coupling Products ' Ir III + biphenosemiquinone or phenoxyphenoxyl radical + H + k overox, 1 (3-6) Ir IV + biphenosemiquinone ' Ir III + biphenoquinones + H + k overox, 2 (3-7) 2 phenoxyphenoxyl radical ' polymer k dec (3-8) k disp (3-9) OH Ir V + Ir I ++ HR O R O Ir V + Ir I +R O R + H OH R O R 2 O RCoupling Products O OH H 2 O + 2 + OH OH 113 Here, R represents 2-CH 3 , 2,6-(CH 3 ) 2 , 4-OCH 3 and 4-C(CH 3 ) 3 . The forward second-order rate constants in eqs 3-2 and 3-3 were obtained experimentaly and are summarized in Table 3-3. 3.4.2 Mechanism of k ArO ? Term Marcus Theory. At high pH, most of the reductants are phenoxide anions instead of phenols and their oxidation by Ir IV occurs through a simple outer-sphere reversible electron transfer mechanism. With the size and unbalanced charge correction term W 12 , we could use the cross-relationship of the Marcus theory (eqs 1-12 to 1-15) to obtain more acurate self-exchange rate constants for phenoxide anions. The cross electron- transfer rate constant k 12 , k ArO ?, was obtained experimentaly and ilustrated in Table 3-3. The radii of phenoxides are estimated using CPK Atomic Models. Other parameters are the same as those described for the phenol oxidation in Chapter 2. The calculated self- exchange rate constants of the phenoxide radical/anion couple, k 11 , are shown in Table 3-4. The driving force (?"G?) for the electron transfer step betwen Ir IV and phenoxide anion can be calculated from the diference betwen the potential for Ir IV /Ir III couple (= 0.893 V vs NHE at ? = 0.1 M) 93 and that for ArO ? / ArO ? couple shown in Table 3-4. The corrected Gibbs free energy, "G??, is related to "G? through the relation of eq 1-4 where w 12 (summarized in Table 3-4) is the coulombic work betwen [IrCl 6 ] 2? and phenoxide anion, and w 21 is that betwen [IrCl 6 ] 3? and the phenoxyl radical (equal to zero). From the experimental rate constants, k ArO ?, we can calculate the values for activation barrier ("G ? ) acording to eq 1-7 using Z = 10 11 M ?1 s ?1 . The reorganization energy (#) can be estimated using the adiabatic Marcus relationship in which the activation barrier is 114 Table 3-3. Kinetic Data for the Reaction of Phenols with Ir IV . substrate k ArOH , a M ?1 s ?1 k ArO ?, a M ?1 s ?1 pK a KIE ! +a Phenol 0.77 ? 0.03 (8.0 ? 0.2) ! 10 6 9.79 b 3.5 ? 0.3 0 Cresol 16 ? 1 (5.2 ? 0.2) ! 10 7 10.09 b 2.9 ? 0.2 ?0.20 c Xylenol 47 ? 2 (2.0 ? 0.1) ! 10 8 10.38 d 2.6 ? 0.1 ?0.40 e TMP (1.09 ? 0.01) ! 10 3 (2.58 ? 0.03) ! 10 8 10.65 d 2.04 ? 0.01 ?0.71 e MOP (6.1 ? 0.1) ! 10 4 (3.3 ? 0.10) ! 10 8 9.96 1.9 ? 0.1 ?0.78 TBP 29 ? 3 (3.6 ? 0.3) ! 10 7 10.07 d ?0.26 POP (1.2 ? 0.1) !10 3 (1.1 ? 0.2) !10 8 9.90 f ?0.50 a Brown and Hamet substituent constant. The value for para-substituted phenols are from reference 107 except where specificaly noted. b Obtained from reference 81 at ? = 0.1 M and 25?C. c ! + = 0.66 ! ! + (para-cresol) with ! + (para-cresol) = ?0.31 acording to reference 107. d In reference 81, pK a values of xylenol, TMP and TBP are reported at zero ionic strength, 10.62, 10.89 and 10.31, respectively. They are converted in to pK a s at ? = 0.1 M by using Davies equation. e Sum of each substituent constant. f The reported pK a1 = 9.81 at ? = 0.25 M in reference 87 is converted to pK a1 = 9.90 at ? = 0.1 M by using Davies equation. 115 Table 3-4. Calculated Data for Phenols and the Reactions with Ir IV . substrate r, a ? E?(ArO ? / ArO ? ),V k 11 , b M ?1 s ?1 f 12 c w 12 d "G?? e ArO ? ! ArO? "G ?f ArO ? ! ArO? phenol 2.5 0.80 g 2.3 ! 10 6 0.74 3.2 ?12.3 23.4 cresol 2.9 0.77 h 3.4 ! 10 7 0.61 2.9 ?15.0 18.7 xylenol 3.1 0.56 i 8.2 ! 10 5 0.11 2.8 ?34.8 15.4 TMP 3.2 0.50 j 2.4 ! 10 5 0.05 2.8 ?40.8 14.8 MOP 2.9 0.55 g 2.0 ! 10 6 0.08 2.9 ?36.2 14.1 TBP 3.5 0.77 j 1.4 ! 10 7 0.63 2.6 ?14.6 19.6 4,4?-BP 3.8 0.64 k 4.0 ! 10 7 0.18 2.5 ?26.8 12.5 2,2?-BP 4.4 1.00 k 3.0 ! 10 8 0.83 2.2 8.21 25.1 a Radii are estimated from CPK Atomic Models. b Self-exchange rate constant calculated from eqs 1-12 to 1-15 for phenolates/phenoxide radicals couples. c Constant in eq 1-12. d Coulombic work betwen [IrCl 6 ] 2? and phenoxide anion with unit in kJ mol ?1 . e Corrected Gibbs free energy in kJ mol ?1 and calculated from eqs 1-4 and 1-5, with w 21 = 0. f In kJ mol ?1 and calculated acording to eq 1-7 with k ArO ? as k et . g Reference 88. h Reference 108. i This work. j Reference 109. k Reference 85. 116 dependent of the driving force as shown in eq 1-6. A nonlinear fit of "G?? and ("G ? ? w 12 ) for al phenols including unsubstituted phenol, 4,4?-biphenol and 2,2?-biphenol to eq 1-6 gives the solid curve in Figure 3-22 with the reorganization energy # 12 = 103 ? 4 kJ mol ?1 . The reorganization energy # 12 could also be estimated acording to the additivity postulate (eq 1-11), in which asumes that the reorganization energy for the overal reaction, # 12 , is the average of the reorganization energies of the two individual self- exchange reactions, # 11 and # 22 . From the self-exchange rate constant of the Ir IV /Ir III redox couple (k 22 = 2 ! 10 5 M ?1 s ?1 ), # 22 is calculated to be 130 kJ mol ?1 when holding "G? = 0. # 11 = 106 kJ mol ?1 for the phenoxide anion/phenoxyl radical couple was calculated from its self-exchange rate constant, k 11 = 2.3 ! 10 M ?1 s ?1 , and the method used to calculate the self-exchange rate constant is described in chapter 2. Therefore, a # 12 value of 118 kJ mol ?1 for phenol reaction is calculated acording to eq 1-11. This result is larger than the average # 12 value (103 ? 4 kJ mol ?1 ) obtained from the above-mentioned "G???("G ? ? w 12 ) relation because of the diferent self-exchange barriers for each substrate. The Hammett Correlation. Because of the strong resonance interaction betwen electron donating substituents and the electron-deficient center of phenol, the Brown and Hamet substituent constant ! +107 is applied to the Hamet equation 3-10, k = k 0 10 ((!+) (3-10) where k are the second-order rate constants for the reaction of Ir IV with the substituted phenoxide anions and k 0 is that for the unsubstituted phenoxide anions. ! is the reaction constant. The ! + values for para-substituted phenols are taken from the literature, 107 117 Figure 3-22. Plot of ("G ? ? w 12 ) versus "G??. The solid line is the fit to eq 1-6. -100 -70 -40 -10 20 50 -5 7 19 31 43 55 Cresol Xylenol TMP MOP TBP Phenol 4,4'-biphenol 2.2'-biphenol !G?' ArO -, kJ mol ?1 ! G ! ? w 1 2 , k J m o l ? 1 118 while that for ortho-cresol (! + = ?0.20) is calculated acording to the relationship provided by Jonsson et al., 110 which is 0.66 times that of the ! + for para-cresol (= ?0.31). For the reactions of phenols containing more than one substituent, bi-substituted phenol (xylenol) and tri-substituted phenol (TMP), ! + is normaly taken as the sum of each substituent constant. Al ! + are summarized in Table 3-4 including that of 4- phenoxyphenol. When plotting the second-order rate constants of phenoxide anions (k ArO ?) versus ! + as shown in Figure 3-23, a correlation is obtained with ! equal to ?2.5 ? 0.3. The decreasing trend is obeyed by al reactions and implies that the oxidations are aided by electron donating groups. This, in turn, suggests that the reaction center (here, the phenoxyl oxygen) is more positive in the transition state. Figure 3-24 shows a plot of pK a of al phenols at 0.1 M ionic strength versus ! + . A linear correlation is achieved with the slope of ?1.2 ? 0.1 excluding the reaction of MOP and POP. This result indicates that the oxygen atom atached on the benzene ring affects the correlation betwen ! + and acid-base behavior of phenols. 3.4.3 Mechanism of k ArOH Term In the two-term rate law, k ArOH refers to the oxidation of phenols in acidic media (in the p[H + ] range of 1?3). Under these conditions two possible pathways are considered: a sequential mechanism, including electron transfer prior to the proton transfer (ET/PT), and a concerted PCET (CPET) mechanism. An atempt to distinguish these two proceses is performed as described below. Thermodynamics. The driving force for the concerted proton-coupled electron transfer ("G? CPET ) and the stepwise ET/PT proces ("G? ET/PT ) can be calculated from the diference betwen the redox potential of Ir IV /Ir III couple and that of ArO ? , H + /ArOH 119 Figure 3-23. Correlations betwen the second-order rate constants for reaction of phenols (k ArOH ) and phenoxide anions (k ArO ?) with Ir IV versus substituent constants ! + . Values of k ArOH are labeled with the closed circle and k ArO ? with the open circle. The line is the nonlinear regresion to eq 3-10. -1.00 -0.76 -0.52 -0.28 -0.04 0.20 10 -2 10 3 10 8 10 13 k ArO ? k ArOH ! + k, M ?1 s ?1 120 Figure 3-24. Plot of pK a versus substituent constants ! + . The open circle point represents data for the oxidation of MOP and the open square point represents the results for the reaction with 4-phenoxyphenol. The closed circle points resulted from the oxidation of the phenol, cresol, xylenol, TMP, and TBP. Solid line is the linear regresion of closed circle points. -1.00 -0.76 -0.52 -0.28 -0.04 0.20 9.0 9.5 10.0 10.5 11.0 11.5 POP MOP ! + pK a 121 couple or ArOH +? /ArOH couple. Eqs 3-11 and 3-12 show the calculation of E?(ArO ? , H + /ArOH) and E?(ArOH +? /ArOH), respectively. E?(ArO ? , H + /ArOH) = E?(ArO ? /ArO ? ) + 0.059 ! pK a, ArOH (3-11) E?(ArOH +? /ArOH) = E?(ArO ? /ArO ? ) + 0.059 ! (pK a, ArOH ? pK a, ArOH+? ) (3-12) The reported pK a, ArOH+? values for phenol, cresol and MOP are ?2.75 111 , ?1.99 112 and ? 1.41 112 , respectively. No data are available for the other phenols. Table 3-5 lists the reduction potentials and "G?. As expected, the driving force for ET/PT pathway in phenol, cresol and MOP oxidation is larger than that for CPET pathway. Acording to the Marcus theory, in this region where "G? is greater than ?# the activation barrier "G ? is increased with an increase of "G?. Therefore, the ET/PT proces has a larger barrier to overcome, which disfavors the stepwise pathway. The Hammett Correlation. A plot of the second-order rate constants of phenols (k ArOH ) versus the Brown and Hamet substituent constant ! + is also shown in Figure 3-23 together with the plot of k ArO ? versus ! + . A decreasing trend is observed for both k ArOH and k ArO ?, which implies that the oxidations are aided by electron donating groups and the reaction center (here the phenoxyl oxygen) becomes more positive in the transition state. The reaction constant ! is obtained as ?6.0 ? 0.2 when the data are fit to the Hamet equation 3-10 where k refers to k ArOH for substituted phenol and k 0 refers to that for phenol (= 0.77 M ?1 s ?1 ). Compared to the ! value for k ArO ? (?2.5 ? 0.3), the more negative value obtained for k ArOH reveals a moderately higher sensitivity to substituent efects for the rates of reaction with phenols than those with phenoxide anions. 122 Table 3-5. Thermodynamic Parameters for CPET and ET/PT Mechanisms substrate CPET ET/PT E?(ArO ? , H + /ArOH), a V "G? CPET , b kJ mol ?1 "G ? CPET , c kJ mol ?1 pK a, ArOH+? E?(ArOH +? /ArOH), d V "G? ET/PT , e kJ mol ?1 phenol 1.38 46.7 46.8 ?2.75 f 1.54 62.4 cresol 1.36 45.6 45.6 ?1.99 g 1.48 56.9 xylenol 1.17 27.2 28.7 TMP 1.13 22.7 25.2 MOP 1.14 23.6 25.8 ?1.41 g 1.22 31.6 TBP 1.36 45.4 45.5 4,4?-BP 1.21 30.8 31.7 2,2?-BP 1.45 53.8 54.4 a Calculated acording eq 3-11 with E?(ArO ? /ArO ? ) listed in Table 3-4 and pK a, s in Table 3-3. b Equal to ?ZF[E?(Ir IV /Ir III ) ? E?(ArO ? , H + /ArOH)]. c Calculated by "G ? CPET = [# (1 + "G? CPET /#) 2 ]/4 with # = (# ox + # CPET )/2 (# CPET = 43.4 kJ mol ?1 and # ox, IrIV = 130 kJ mol ?1 ). d Calculated acording eq 3-12. e Equal to ?ZF[E?(Ir IV /Ir III ) ? E?(ArOH +? /ArOH)]. f Reference 111. g Reference 112. 123 Figure 3-25. Plot of kinetic isotopic efect versus substituent constants ! + . The closed circle points are the oxidation of the phenol, cresol, xylenol, TMP and TBP and solid line is the linear regresion. Data from Table 3-3. -1.00 -0.76 -0.52 -0.28 -0.04 0.20 1.0 1.7 2.4 3.1 3.8 4.5 ! + KIE 124 KIE?Hammett Relationship. The deuterium kinetic isotope efects of the reactions betwen Ir IV and al phenols, except for TBP and POP, versus the Hamet substituent constant ! + are plotted in Figure 3-25. A good linear relationship is observed with the slope of 2.0 ? 0.1. Comparison with Phenols Oxidation by ClO 2 . The oxidations of phenol, cresol, TMP, MOP and TBP by ClO 2 were studied by Hoigne and Bader. 113 They used a two- term rate law to obtain the values for k ArOH and k ArO ?, as shown in Table 3-6. Table 3-6. Kinetic data for Phenols Oxidation by ClO 2 . a Reductant k ArOH , M ?1 s ?1 k ArO ?, M ?1 s ?1 phenol 0.4 ? 0.1 (4.9 ? 0.5) ! 10 7 cresol 16 ? 4 (4.4 ? 0.4) ! 10 8 TMP (3.9 ? 0.2) ! 10 3 (4.0 ? 1.0) ! 10 9 MOP (2.5 ? 0.2) ! 10 4 (1.7 ? 0.6) ! 10 9 TBP 11 ? 3 (1.5 ? 0.1) ! 10 8 a Reference 114. When plotting ln k of phenol oxidation by ClO 2 versus ln k of Ir IV reactions, a linear relationship, as shown in Figure 3-26, is obtained both for k ArOH and k ArO ? with slopes of 1.0 ? 0.1 and 1.1 ? 0.2, respectively. This result suggests that the oxidation of phenols by Ir IV is a series of common phenol reactions and that the mechanism of the Ir IV reactions is similar to those of other oxidants, like ClO 2 . 125 Figure 3-26. Plot of ln k (ClO 2 ) versus ln k (Ir IV ). The closed circle points are the oxidation of the phenol, cresol, TMP, MOP and TBP by ClO 2 and the open circle points represent these oxidation by Ir IV . Solid lines are the linear regresion. -5 1 7 13 19 25 -5 1 7 13 19 25 k ArO ? k ArOH ln k, Ir IV ln k, ClO 2 126 Marcus Theory. Acording to Bonin, 115 the reorganization energy (#) of a concerted proton-coupled electron-transfer reaction with water acting as the proton aceptor could be estimated acording to eq 3-13. # = !ox+ CPET 2 (3-13) # CPET is the self-exchange reorganization energy of the reaction described by eq 3-14. ( * ArOH??H 2 O) + (ArO ? , + H 3 O) ! ( * ArO ? , + H 3 O) + (ArOH??H 2 O) (3-14) Here, ????? represents hydrogen bonds betwen phenols and H 2 O. # CPET is mainly atributed to solvent organization. The value for phenol reaction was reported as 43.4 kJ mol ?125 and derived from the oxidations by three diferent oxidants including [IrCl 6 ] 2 . The extension of their results to the ClO 2 oxidation is discussed here by considering the diference of total reorganization energy caused by the diferent oxidants. As mentioned in Chapter 2, the self-exchange rate constant of the Ir IV /Ir III redox couple is equal to 2 ! 10 5 M ?1 s ?1 , therefore, # ox is calculated to be 130 kJ mol ?1 . The self- exchange rate constant for the ClO 2 /ClO 2 redox couple is reported as 160 M ?1 s ?1 , 54 which means that # ox is equal to 201 kJ mol ?1 . A kinetic study of phenol oxidations by [Ru(bpy) 3 ] 3+ , [Ru(methyl-bpy) 3 ] 3+ and [Ru(bpy)(ester-bpy) 2 ] 3+ is reported by Bonin, 25 and the data are shown in Table 3-7. # ox for the [Ru(bpy) 3 ] 3+ /[Ru(bpy) 3 ] 2+ couple is 45.6 kJ mol ?1 , 116 and al three Ru III /Ru II couples have the same # ox value. 115 With al these # ox values, the total reorganization energy # could be calculated acording to eq 3-13 with 127 Table 3-7. Thermodynamic and Kinetic Parameters for Phenol Oxidation by Diferent Oxidants. Oxidant k ArOH , M ?1 s ?1 # ox , kJ mol ?1 E?(Ox), V "G? CPET , a kJ mol ?1 "G ? CPET , b kJ mol ?1 [Ru(bpy) 3 ] 3+c 3.0 ! 10 5 45.6 d 1.27 10.3 16.9 [Ru(methyl-bpy) 3 ] 3+c 4.0 ! 10 3 45.6 1.09 27.6 29.2 [Ru(bpy)(ester-bpy) 2 ] 3+c 1.0 ! 10 8 45.6 1.47 ?9.03 7.07 ClO 2 e 0.4 201 f 0.95 40.8 54.3 [IrCl 6 ] 2? 0.77 130 0.89 46.7 51.3 a Equal to ?ZF[E?(ox) ? E?(ArO ? , H + /ArOH)]. b Calculated from "G ? CPET =[# (1+ "G? CPET /#) 2 ]/4 with # = (# ox + # CPET )/2 and # CPET = 43.4 kJ mol ?1 . c Reference 25. d Reference 116. e Reference 114. f Reference 54. 128 # CPET = 43.4 kJ mol ?1 . The diferent reduction potentials of these oxidants correspond to diferent driving forces ("G? CPET ). Acording to the Marcus theory, the activation barrier ("G ? CPET ) could be calculated with the known # and "G? CPET values by holding the coulombic work terms w 12 and w 21 to be zero. A plot of ln k ArOH versus "G ? CPET , shown in Figure 3-27, yielded a good linear relationship with slope of ?0.40 ? 0.02 and an intercept of 20.3 ? 0.7. This slope confirms the theoretical value 0.40 (=1000/RT) which suggests that Marcus theory can be applied to H 2 O-CPET reaction. The intercept value gives the collision frequency, Z, equal to 6.6 ! 10 8 M ?1 s ?1 which is smaler than the value in the simple electron transfer proces (= 10 ! 10 11 M ?1 s ?1 ). A test about the validity of # CPET is performed using a two-fold increase (# CPET = 86.8 kJ mol ?1 ). The same calculation was made for "G ? CPET , and the plot of ln k ArOH versus "G ? CPET still yielded a good straight line and reproduced the theoretical slope. However, the value of Z increased by almost ten times to 5.3 ! 10 9 M ?1 s ?1 compared to that derived with # CPET = 43.4 kJ mol ?1 . Although the acurate # CPET value cannot be obtained, the results demonstrate that Marcus theory could be succesfully applied for H 2 O-CPET phenol oxidation and the # CPET values are the same for all above-mentioned reactions with diferent oxidants. The "G ? CPET values for the oxidation of the phenols by Ir IV are calculated acording to Marcus theory (eqs 1-4 and 1-6) and summarized in Table 3-5. The reorganization energy (#) for each reaction was obtained from eq 3-13 with # CPET = 43.4 kJ mol ?1 and # ox = 130 kJ mol ?1 . "G? CPET values are also listed in Table 3-5 and the employed coulombic work terms w 12 and w 21 were equal to zero. A plot of ln k ArOH versus "G ? CPET is shown in Figure 3-28 and no specific correlation is observed. This result reveals that the # CPET 129 Figure 3-27. Plot of ln k ArOH versus "G ? CPET for phenol oxidation by diferent oxidants. Solid line is the linear regresion. Data from Table 3-7. 0 12 24 36 48 60 -5 1 7 13 19 25 !G ? CPET , kJ mol ?1 ln k ArOH [Ru(bpy)(ester-bpy) 2 ] 3+ [Ru(bpy) 3 ] 3+ [Ru(methyl-bpy) 3 ] 3+ [IrCl 6 ] 2? ClO 2 130 Figure 3-28. Plot of ln k ArOH versus "G ? CPET for the Ir IV reduction by diferent phenols. 10 22 34 46 58 70 -5 -1 3 7 11 15 phenol cresol xylenol TMP MOP TBP 4,4'-BP 2,2'-BP !G ? CPET , kJ mol ?1 ln k ArOH 131 values are influenced by the substituents of phenol, which alter the physical properties of phenol and the hydrogen bonding betwen phenols and H 2 O is expected to change. Temperature Dependence. The temperature dependence of cresol and MOP are summarized in Table 3-8. Table 3-8. Activation Parameters for Cresol and MOP Oxidation by Ir IV . a substrate "G ? , b kJ mol ?1 (H ? , kJ mol ?1 (S ? , J mol ?1 K ?1 E a , kJ mol ?1 "G ? CPET , c kJ mol ?1 Cresol 67.1 42.0 ?84.1 44.5 45.6 MOP 46.0 22.8 ?78.0 25.3 25.8 a At 298 K. b "G ? = (H ? ? 298 ! (S ? . c Obtained with # CPET = 43.4 kJ mol ?1 (derived from phenol reaction). These reactions were performed at near room temperature, where the Arrhenius activation energy E a is roughly equal to (H ? plus 2.5 kJ mol ?1 . The E a values are consistent with the activation barrier "G ? CPET obtained with # CPET = 43.4 kJ mol ?1 , which is derived from phenol reaction. However, the "G ? CPET values derived from Marcus theory are les than the "G ? values by % 20 kJ mol ?1 , obtained from the temperature dependence experiments for k ArOH . 132 Chapter 4 OXIDATION OF Ac-Y-NH 2 BY HEXACHLOROIRIDATE(IV) 4.1 Introduction As one of the most important alpha-amino acids in biological chemistry, tyrosine is of interest to many researchers. The phenol functional group facilitates the electron transfer betwen tyrosine and other cofactors in photosystem II 40,41 and RNR. 18,117 The C- and N-termini provide tyrosine three pK a ?s and redox-activity which complicates kinetic studies. In this chapter, a N-acetyl and C-amide protected tyrosine, N-acetyl-tyrosine amide, is used as a reductant for Ir IV . The kinetic study of this reaction wil provide insight into the mechanisms of biological electron transfer reactions. Overoxidation is observed, and two-term and three-term rate laws are applied to obtain the rate constants: k ArOH , k ArO ? and k?. 4.2 Experimental Section 4.2.1 Reagents and Solutions Al commercial chemical reagents were used as received except as noted. Amonium hexachloroiridate(III) monohydrate (abbreviated as Ir III ), deuterium oxide, 133 sodium acetate anhydrous, cacodylic acid and sodium hydroxide were purchased from Sigma?Aldrich Chemicals Company. Perchloric acid, acetic acid, monochloroacetic acid were obtained from Fisher Scientific Co. Amonium hexachloroiridate(IV) (abbreviated as Ir IV ) was purchased from Alfa. The spin-trapping agent DBNBS was synthesized as described in Chapter 2. N-Acetyl-tyrosine amide (Ac-Y-NH 2 ) and N-acetyl-phenylalanine amide (Ac-Phe-NH 2 ) were purchased from Bachem, and their structures are shown in Scheme 4-1. Scheme 4-1. The structures of Ac-Y-NH 2 and Ac-Phe-NH 2 . Al solutions were freshly prepared with deionized water provided by a Barnstead NANO Pure Infinity ultrapure water system, and purged with argon gas prior to the reactions to prevent potential complications caused by O 2 . The ionic strength was adjusted with lithium perchlorate trihydrate (GFS) and is approximately equal in both oxidants and reductants solutions to prevent Schlieren efects (or refractive index effect 63 ). Selected buffer solutions (acetate, monochloroacetate, and cacodylate buffers) were applied to control the pH if necesary. 4.2.2 Methods A Corning 450 pH/ion meter was used with a Metler Toledo InLab 421 or InLab HO HN H 2 O O N-Acetyl-rosine amide (NH 2 ) HN H 2 O O N-Acetyl-phenylaine amide (cP-NH 2 ) 134 Semi-Micro-L combination pH electrode. The reference electrode electrolyte was replaced with 3 M NaCl to prevent the formation of KClO 4 precipitate. With the known H + concentration and pH reading, the activity coeficient ! (= 0.839 ? 0.04) was obtained from equation p[H + ] = pH + log !, where p[H + ] is equal to ?log [H + ]. Al measurements were performed at 25.0 ? 0.1 ?C. The kinetics experiments were carried out on a Hi-Tech SF-51 stopped-flow spectrophotometer with OLIS 4300 data acquisition and analysis software. UV-vis spectra were monitored on a HP-8453 diode array spectrophotometer equipped with a Brinkman Lauda RM6 thermostated water bath to maintain the temperature at 25 ?C. Al kinetics data were obtained by monitoring the absorbance of Ir IV at 488 nm. The observed pseudo-first-order rate constants were obtained from fiting kinetic traces over five half lives to first-order exponential functions and each reported observed rate constant is the average of at least 5 shots. The Specfit/32 version 3.0.15 global analysis system was applied to simulate the reaction traces, and the GraphPad Prism 4 or 5 software was used to analyze the rate law with 1/Y 2 weighting. 4.3 Results The kinetic trace of a typical reaction betwen Ac-Y-NH 2 and Ir IV is shown in Figure 4-1. 0.004 M of Ac-Y-NH 2 is oxidized by 2.5 ! 10 ?5 M of Ir IV in 0.05 M HClO 4 at 0.1 M ionic strength. Although we atempted to use a smal amount of Ir IV to obtain an improved first-order kinetics, the consumption curve of Ir IV does not fit a first-order exponential function wel. 4.3.1 Hexachloroiridate(II) Inhibition 135 Figure 4-1. Trace of the Ac-Y-NH 2 reaction (solid line) and first-order fit (dashed line). [Ac-Y-NH 2 ] tot = 0.004 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M, [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -0.02 0.00 0.02 R e s i d u a l s 0 400 800 1200 1600 2000 0.00 0.03 0.06 0.09 Times, s A 488 136 Figure 4-2. Comparative traces of the Ac-Y-NH 2 reaction (solid line) with added Ir III (dashed line). [Ac-Y-NH 2 ] tot = 8.0 ! 10 ?4 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [Ir III ] = 12.5 ! 10 ?5 M; p[H + ] = 3.5 (0.02 M acetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. 0 160 320 480 640 800 0.00 0.03 0.06 0.09 0.12 with added 5-fold Ir III no added Ir III Time, s A 488 137 The solid line in Figure 4-2 displays the reaction traces of 8.0 ! 10 ?4 M Ac-Y-NH 2 and 2.5 ! 10 ?5 M Ir IV at p[H + ] = 3.5, while the dashed line shows the kinetic trace of the same reaction with the addition of a 5-fold exces of Ir III . A significant diference is observed when comparing these two reactions, which implies that the oxidation of Ac-Y- NH 2 by Ir IV is afected by Ir III . 4.3.2 Spin Trapping Effect In order to remove the Ir III inhibition, the phenoxyl radical spin trapping agent, DBNBS, was applied to the Ir IV reaction. Various concentrations of DBNBS (0.1?10 mM) were added into the solutions containing 0.004 M of Ac-Y-NH 2 and 2.5 ! 10 ?5 M of Ir IV in 0.05 M HClO 4 . The first half-lives of al these reactions are summarized in Table 4-1. Table 4-1. Kinetic Data for the Reaction of Ac-Y-NH 2 with Ir IV in the Presence of DBNBS. a [DBNBS], mM t 1/2 , s 0 165 0.1 133 0.5 74.7 1.0 58.3 2.0 50.0 5.0 36.7 10 36.7 a [Ac-Y-NH 2 ] tot = 0.004 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 138 The reaction rate increases with increasing concentrations of DBNBS and approaches a maximum value when 5 mM DBNBS were added. Subsequently, the rate does not change even after adding more DBNBS. In the presence of 5 mM DBNBS, a good first- order fit was obtained as shown in Figure 4-3 at 488 nm. This result suggests that 5 mM DBNBS is sufficient to prevent Ir III inhibition. 4.3.3 Ac-Y-NH 2 Dependence At p[H + ] = 2.8, the oxidations of (8.0?72) ! 10 ?4 M of Ac-Y-NH 2 by 2.5 ! 10 ?5 M of Ir IV were examined in presence of 5 mM DBNBS under pseudo-first-order conditions. Al kinetic data are collected in Table A-25. When plotting the observed rate constants versus the total concentrations of Ac-Y-NH 2 , a linear relationship is obtained as shown in Figure 4-4 with the slope of 13.0 ? 0.1 M ?1 s ?1 . This result indicates that the rate law is first-order with respect to Ac-Y-NH 2 as shown in eq 4-1: k obs = k [Ac-Y-NH 2 ] tot (4-1) 4.3.4 Dependence on p[H + ] The p[H + ] dependence was studied using 2.5 ! 10 ?5 M of Ir IV and (3.0?64) ! 10 ?4 M of Ac-Y-NH 2 over the p[H + ] range of 1?7 in the presence of 5 mM DBNBS. The experimental data are summarized in Table A-26, and a plot of k obs /[Ac-Y-NH 2 ] tot versus p[H + ] is shown in Figure 4-5. The simple two-term rate law, eq 2-2, is applied to analyze the data. K a is the acid disociation constant of Ac-Y-NH 2 , pK a = 9.90. 45 k ArOH and k ArO ? represent the reactivity of Ac-Y-NH 2 and the corresponding phenolate species. A 139 Figure 4-3. Trace of the Ac-Y-NH 2 reaction in the presence of DBNBS (solid line) and first-order fit (dashed line). [Ac-Y-NH 2 ] tot = 0.004 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M, [DBNBS] = 5 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. -0.002 0.000 0.002 R e s i d u a l s 0 100 200 300 400 500 0.00 0.03 0.06 0.09 Times, s A 488 140 Figure 4-4. Plot of k obs vs [Ac-Y-NH 2 ] tot . [Ac-Y-NH 2 ] tot = (8.0?72) ! 10 ?4 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 5 mM; p[H + ] = 2.8 (0.02 M monochloroacetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. Solid line is linear fit. Data from Table A-25. 0.000 0.002 0.004 0.006 0.008 0.00 0.02 0.04 0.06 0.08 0.10 [Ac-Y-NH 2 ] tot , M k obs , s -1 141 Figure 4-5. Plot of k obs /[Ac-Y-NH 2 ] tot vs p[H + ]. [Ac-Y-NH 2 ] tot = (3.0?64) ! 10 ?4 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 5 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.4 < p[H + ] < 7.4. Solid line is the fit to eq 2-2 and the dashed line is the fit to eq 2-3. Data from Table A-26. 0.0 1.6 3.2 4.8 6.4 8.0 10 -1 10 0 10 1 10 2 10 3 10 4 10 5 10 6 3-term rate law fit 2-term rate law fit p[H + ] k o b s / [ A c - Y - N H 2 ] t o t , M - 1 s - 1 142 nonlinear least-squares fit of the data to eq 2-2 yielded k ArOH = 5.4 ? 0.6 M ?1 s ?1 and k ArO ? = (4.5 ? 0.3) ! 10 7 M ?1 s ?1 . Figure 4-5 shows a fit of the data in Table A-26 to the 3-term rate law (eq 2-3), values for k ArOH , k ArO ? and k? are obtained, with k ArOH = 2.9 ? 0.2 M ?1 s ?1 , k ArO ? = (3.6 ? 0.1) ! 10 7 M ?1 s ?1 and k? = 0.22 ? 0.02 M ?1 s ?1 , as displayed in Table A-27. The contributions of the additional k? term were calculated over the p[H + ] range of 1?7. The maximum value is 49% and occurs around p[H + ] = 2.8. 4.3.5 Stoichiometry and Overoxidation A spectrophotometric titration of Ac-Y-NH 2 with Ir IV was performed to determine the consumption ratio. The experiments were investigated in 0.02 M acetate buffer at p[H + ] = 5.5. Figure 4-6a shows the absorption spectra changes resulting from titrating 2.4 mL of 1.0 ! 10 ?4 M Ac-Y-NH 2 with 2.3 ! 10 ?3 M of Ir IV solution. Al spectra were obtained after the completion of oxidation. The absorption spectrum after adding 2.3 ! 10 ?7 mol of Ir IV is displayed in Figure 4-6b. Under this condition, the peak at 275 nm (maximum absorbance of Ac-Y-NH 2 with " 275 = 1.5 ! 10 3 M ?1 cm ?1 ) reveals that some Ac-Y-NH 2 is stil present in the solution and al added Ir IV is consumed, therefore, the absorption of Ir IV is not an isue. A new peak at 416 nm and a shoulder around 340 nm are observed, which indicates the absorption of products. Same maximum absorption (275 nm) has been found for Ac-Y-NH 2 and its analogous species N-acetyltyrosine, 118 so the coupling product of Ac-Y-NH 2 oxidation is expected to have maximum absorption at around 286 nm in analogy to the N-acetyltyrosine dimer. However, no peak is observed at this wavelength. The peak at 416 nm may originate from the overoxidation product quinone species. The titration curves at 340 nm (Figure 4-7a) and 488 nm (Figure 4-7b) 143 Figure 4-6. Titration of Ac-Y-NH 2 (2.4 mL, 1.0 ! 10 ?4 M) with 2.3 ! 10 ?3 M of Ir IV at p[H + ] = 5.5 (0.02 M acetate buffer) and 25 ?C. Al spectra were obtained after the completion of the oxidation. Dashed lines in (a) and (b) are the spectra obtained after adding 2.3 ! 10 ?7 mol of Ir IV . 200 300 400 500 600 700 0.00 0.24 0.48 0.72 0.96 1.20 aFinal Titration Point Before Titration Wavelength, nm Abs 250 300 350 400 450 500 0.00 0.06 0.12 0.18 0.24 0.30 Ac-Y-NH 2 275 nm 416 nm 340 nm b Wavelength, nm Abs 144 Figure 4-7. The titration curve of Ac-Y-NH 2 (2.4 mL, 0.1 mmol) solution by 2.5 ! 10 ?5 M of Ir IV solution at p[H + ] = 5.5 (0.02 M acetate buffer) and 25 ?C. Closed circles represent the data corrected for volume changes. Dashed lines are the linear regresion fit. (a) 340 nm. (b) 488 nm. The insert shows an oxidation trace (solid line) and a bi-phase exponential fit (dashed line) after titrating with 9.2 ! 10 ?7 mol of Ir IV . 0.0 5.0!10 -07 1.0!10 -06 1.5!10 -06 2.0!10 -06 0.00 0.28 0.56 0.84 1.12 1.40 End Point b n Ir(IV) , mol A 488, corr 0 7201440216028803600 0.20 0.24 0.28 0.32 0.36 0.40 Oxidation trace Biphasic exponential fit Time, s A 488 0.0 5.0!10 -07 1.0!10 -06 1.5!10 -06 2.0!10 -06 0.00 0.05 0.10 0.15 0.20 0.25 End Point a n Ir(IV) , mol A 340, corr 145 are corrected for volume changes using A cor = A obs *V total /V inital . The increase of absorption after the end point at 340 nm in Figure 4-7a is due to the absorption of Ir IV . The consumption ratio of n IrIV /n Ac-Y-NH2 amounted to 2.9, which indicates that overoxidation occurs under these conditions. Significant oxidation is detected even after the end point as shown in the insert in Figure 4-7b when titrating with 9.2 ! 10 ?7 mol of Ir IV . The kinetic decay trace fits a biphasic exponential curve wel with a fast rate constant of (4.7 ? 0.1) ! 10 ?3 s ?1 and a slow rate constant of (3.7 ? 0.1) ! 10 ?4 s ?1 . Al titration curves in Figure 4-6a were collected at the end of the show rate. The biphasic phenomenon may be atributed to the further oxidation of the decomposition products of the Ac-Y-NH 2 overoxidation, and the addition of Ir IV could acelerate this proces. No oxidation is observed after adding 1.38 ! 10 ?6 mol of Ir IV . 4.3.6 Control Experiments In order to gain insight about the reactivity of the caboxamide groups of Ac-Y-NH 2 , the reactions betwen 3.0 ! 10 ?3 M of Ac-Phe-NH 2 and 1.0 ! 10 ?4 M of Ir IV were tested at 0.1 M ionic strength in 0.05 M HClO 4 and at p[H + ] = 7.0. No oxidation was observed under both conditions, which confirms that only the hydroxyl group of Ac-Y-NH 2 participates in the reaction with Ir IV . 4.4 Discusion and Conclusion A mechanism similar to the one established for phenols in chapter 2 and 3 is proposed for the reaction betwen Ir IV and Ac-Y-NH 2 . Concerted proton-coupled electron-transfer takes place when the reductant is Ac-Y-NH 2 acording to eq 4-2. K ArOH , 146 k ArOH and k ?ArOH are the equilibrium constant and rate constants for this reaction. Reversible outer-sphere one-electron transfer occurs when the conjugate base, the phenolate anion, of Ac-Y-NH 2 is the major reactant (eq 4-3). K ArO ?, k ArO ? and k ?ArO ? represent the equilibrium constant and second-order rate constants of this pathway. The relevant acid/base equilibrium betwen Ac-Y-NH 2 and its phenolate anion is described by eq 4-4 with K a s the acid disociation constant. The phenoxyl radicals formed in proceses 4-2 and 4-3 undergo bimolecular C?C coupling to produce 2,2?-biphenol derivative (eq 4-5), which undergoes further oxidation by Ir IV to form the corresponding biphenosemiquinone (eq 4-6). k dim represents the coupling rate constant and k overox, 1 is the rate constant of further overoxidation. As we mentioned before, biphenosemiquinone is readily oxidized by Ir IV with formation of biphenoquinone (rate constant = k overox, 2 ) is described in reaction 4-7. K ArOH , k ArOH , k ?ArOH (4-2) K ArO ?, k ArO ?, k ?ArO ? (4-3) K a (4-4) k dim (4-5) k overox, 1 (4-6) k overox, 2 (4-7) OHIr V + Ir ++ H R O R OIr V + Ir + R O R OH R + HO R 2OR OH R H R Ir V + Ir ++ H OH R OH R OH R O R Ir V + Ir ++ H OH R O R O R O R 147 Here, R represent ?CH 2 CH(CONH 2 )(NHCOCH 3 ). In the presence of the spin trap DBNBS, the dimerization step (eq 4-8) of DBNBS with an equilibrium constant K dim, DBNBS , of 1.3 ! 10 ?3 M is included in the mechanism. The Ac-Y-NH 2 phenoxyl radical is scavenged by DBNBS to form an adduct which undergoes a rapid oxidation by Ir IV as described in eqs 4-9 and 4-10. DBNBS (dimer) ! 2 ! DBNBS K dim, DBNBS , k dim, DBNBS , k ?dim, DBNBS (4-8) Ac-Y-NH 2 ? + DBNBS ' adduct ? k DBNBS (4-9) Ir IV + adduct ? ' Ir III + adduct + k adduct (4-10) Marcus Theory. The electron-transfer rate constant, k ArO ?, is analyzed using the Marcus theory equations 1-12 to 1-15 where k 12 is the cross electron-transfer rate constant (k ArO ? = 4.5 ! 10 7 M ?1 s ?1 ), k 11 and k 22 are the self-exchange rate constants of the Ac-Y-NH 2 ? /Ac-Y-NH 2 ? and Ir IV /Ir III redox couples. A value for k 22 of 2 ! 10 5 M ?1 s ?198 is used in the calculation, 1 ! 10 11 M ?1 s ?1 is used for Z, the collision frequency. 99 Z i , Z j are the ionic charges of the reactants, R is the ideal gas constant, and r is the center to center distance betwen two reactants when they are approaching each other. The radii of [IrCl 6 ] 2? and Ac-Y-NH 2 ? are 4.1 ? 58 and 4.2 ?, respectively, estimated from CPK Atomic Models. ? is the ionic strength. The value of K 12 was calculated to be 1.8 ! 10 4 from the redox potential of the Ac-Y-NH 2 ? /Ac-Y-NH 2 ? couple (= 0.64 V vs NHE 45 ) and the Ir IV /Ir III couple (= 0.893 V vs NHE at ? = 0.1 M 93 ). With al these parameters and the experimental value of k 12 and K 12 , k 11 is calculated from above equations as 2.9 ! 10 5 M ?1 s ?1 . 148 The rate constant of the H 2 O-CPET proces, k ArOH , is also analyzed with the Marcus theory, seting # CPET equal to 43.4 kJ mol ?1 . The driving force, "G? CPET , is calculated to be 32.8 kJ mol ?1 with the reduction potential of the Ac-Y-NH 2 ? /Ac-Y-NH 2 couple as 1.23 V vs NHE. With the known # CPET and "G? CPET values, the activation barrier, "G ? CPET , is estimated to be 33.4 kJ mol ?1 from equation 1-6, where the coulombic work term w 12 is equal to zero. Overoxidation. The large stoichiometric ratio (= 2.9) of n IrIV /n Ac-Y-NH2 suggests that the oxidation of Ac-Y-NH 2 by Ir IV is highly afected by extensive overoxidation. Although in the presence of DBNBS, the reaction may also be influenced by overoxidation, which results in a large maximum contribution of the additional k? term (49%) in a three-term rate law over the p[H + ] range of 1?7. A significant maximum contribution of the additional k? (56%) is also found in the oxidation of TBP where a large substituted group is located on the para position. This result indicates that the large substituent on para position may enhance the overoxidation efect, which leads to a significant pH-dependent rate constant. 149 Chapter 5 OXIDATION OF PHENOL BY TRIS-(1,10-PHENANTHROLINE)OSMIUM(II) 5.1 Introduction Tris-(1,10-phenanthroline)osmium(III) is another substitution-inert oxidant which has a reduction potential close to that of [IrCl 6 ] 2? , but difers in that it is a cation. As shown in Marcus theory, the electrostatic energy is dependent on the charges of the oxidant, which influences the activation barrier ("G ? ) and thereby the rate of electron transfer. The oxidation of phenol by tris-(1,10-phenanthroline)osmium(III) is reported in this chapter. The Os III complex was obtained by conversion of the corresponding Os II complex with bromine. Al kinetic data were obtained in the presence of Os II in order to avoid complications arising from exces bromine. A second order rate law on [Os III ] and [phenol] tot is established, and a mechanism is proposed. 5.2 Experimental Section 5.2.1 Reagents and Solutions 150 Al commercial chemical reagents were used as received except as noted. Sodium acetate anhydrous, cacodylic acid and sodium hydroxide were purchased from Sigma? Aldrich Chemicals Company. Perchloric acid, sodium chloride, amonium chloride, acetic acid, monochloroacetic acid, ethanol, diethyl ether, 1,4-dioxane, petroleum ether, ethyl acetate, bromine, acetonitrile, hydrochloric acid and hydrogen peroxide were from Fisher Scientific Co. Phenol (Fluka) was recrystalized from a 75% w/w water solution as described in the literature. 62 Pure [Os(phen) 3 ]Cl 2 (Os II ) was prepared following a procedure of a prior study 59 and its structure is shown in Scheme 5-1. Scheme 5-1. Structure of [Os(phen) 3 ] 2+ . Al aqueous solutions were freshly prepared with deionized water provided by a Barnstead NANO Pure Infinity ultrapure water system, and purged with argon gas prior to the reactions to prevent potential complications caused by O 2 . A solution of % 0.1 M bromine in acetonitrile was prepared under argon gas and kept in the dark. The bromine stock solution was then diluted and the concentration was quantified spectraly acording to the absorbance of bromine at ) max (392 nm) with " 392 = 183 ? 4 M ?1 cm ?1 . 119 The ionic strength was adjusted by NaCl and was approximately equal in both oxidants and 2+ N N N Os 151 reductants solutions to prevent Schlieren efects (or refractive index efect 63 ). Selected buffer solutions (acetate, monochloroacetate, and cacodylate buffers) were applied to control the pH if necesary. Preparation of [Os(phen) 3 ] 3+ Solutions. The Os III aqueous solution (concentration of CH 3 CN is les than 1% v/v) was prepared in situ under argon gas by adding a 0.01 M bromine-acetonitrile solution (saturated with Ar) into the Os II solution in 0.01 M HCl right before the reaction. 56 In order to prevent complications caused by bromine, an exces of Os II was maintained. The solutions were protected from light at al times. 5.2.2 Methods A Corning 450 pH/ion meter was used with a Metler Toledo InLab Semi-Micro-L combination pH electrode. Electrode calibrations at ? = 0.1 M (LiClO 4 ) were carried out with 0.01?0.1 M perchloric acid. With the known H + concentration and pH reading, the activity coeficient ! (= 0.839 ? 0.04) was obtained from equation p[H + ] = pH + log !, where p[H + ] is equal to ?log [H + ]. Al measurements were performed at 25.0 ? 0.1 ?C. The kinetics experiments were carried out on a Hi-Tech SF-51 stopped-flow spectrophotometer with OLIS 4300 data acquisition and analysis software. UV-vis spectra were monitored on a HP-8453 diode array spectrophotometer equipped with a Brinkman Lauda RM6 thermostated water bath to maintain the temperature at 25 ?C. Al kinetics data were obtained by monitoring the absorbance of at 480 nm or 550 nm with a 375-nm optical UV cut-off filter to prevent photoreactions. k obs is the observed second-order rate constant based on concentration acording to eq 5-1. 152 1 [Os I ] ! I 0 =kobs t (5-1) The observed second-order rate constant based on absorbance, k obs ?, was obtained from fiting kinetic traces over 4 half-lives to the second-order function acording to eq 5-2. Each reported experimental rate constant is the average of at least 4 shots. 1 A!"0 =kobs' t (5-2) k obs is calculated from k obs ? by multiplying of the molar absorptivity diferences of Os II and Os III (" OsI ? " OsII ) at the particular wavelength, as shown in eq 5-3. k obs = k obs ? (" OsI ? " OsII ) (5-3) The Specfit/32 version 3.0.15 global analysis system was applied to simulate the reaction traces, and the GraphPad Prism 5 software was used to analyze the rate law with 1/Y 2 weighting. 1 H NMR spectra were acquired on a Bruker AV 400 MHz spectrometer; chemical shifts in D 2 O are relative to DS. 5.3 Results 5.3.1 Characterization of the Osmium Complexes 153 Figure 5-1 shows the 1 H NMR spectrum of [Os(phen) 3 ]Cl 2 in the low-field region. There are four signals at chemical shifts of 7.58, 8.08, 8.29 and 8.42 ppm, which are asigned to the four non-equivalent protons in the phenanthroline rings. The UV-vis spectrum of [Os(phen) 3 ]Cl 2 in acidic aqueous solution, as shown in Figure 5-2, displays a maximum absorption peak at 430 nm with " 430 = (1.80 ? 0.15) ! 10 4 M ?1 cm ?1 , which is consistent with the reported value 1.85 ! 10 4 M ?1 cm ?1 . 59 In order to avoid complications arising from the potential biphenoquinone products which exhibit an absorption at around 400 nm, we performed our kinetic measurements at 480 nm where " 480 = (1.67 ? 0.13) ! 10 4 M ?1 cm ?1 . Under pseudo-second-order conditions, high amounts of Os II were required to be present in the solution, which produced an extremely strong absorbance at 480 nm that exceded the detection limit. Therefore, we followed the reactions at 550 nm, where Os II absorbs relatively weakly with " 550 = (5.15 ? 0.34) ! 10 3 M ?1 cm ?1 . We obtained the UV-vis spectrum of [Os(phen) 3 ]Cl 3 (Figure 5-2) via bromine oxidation of Os II in 0.01 M HCl. The maximum absorption diference betwen Os III and Os II appears around 480 nm. All molar absorptivities of both Os III and Os II at 430, 480 and 550 nm are summarized in Table 5-1. The molar absorptivity diferences of Os II and Os III (" OsI ? " OsII ) at particular wavelengths are also displayed in Table 5-1. 5.3.2 Kinetics The oxidation of 0.02 M phenol by 0.67 ! 10 ?5 M Os III was carried out in 0.02 M acetate buffer at p[H + ] = 4.1 in the presence of 1.56 ! 10 ?5 M Os II . The reaction trace at 480 nm is shown in Figure 5-3a and can be fit to a pseudo-second-order model. The kinetic trace of the reaction betwen 0.73 ! 10 ?5 M of Os III and 0.002 M phenol with 1.28 154 Figure 5-1. 1 H NMR spectrum of [Os(phen) 3 ]Cl 2 in D 2 O (aromatic region). 7.67.77.87.98.08.18.28.38.4 ppm 7.5687.5817.5887.6028.0728.0758.0858.0888.2888.4168.4198.4378.440 1.101.041.011.00 PC 1.00 GB 0 LB 0.00 Hz SSB 0 WDW no SF 400.1799189 MHz SI 32768 SFO1 400.1824713 MHz PL1 2.80 dB P1 10.75 usec NUC1 1H ======== CHANNEL f1 ======== TD0 1 D1 1.00000000 sec TE 293.5 K DE 6.50 usec DW 60.400 usec RG 362 AQ 3.9584243 sec FIDRES 0.126314 Hz SWH 8278.146 Hz DS 2 NS 16 SOLVENT D2O TD 65536 PULPROG zg30 PROBHD 5 mm BBO BB?1H INSTRUM Avance400 Time 12.46 Date_ 20091211 PROCNO 1 EXPNO 1 NAME Os(phen)3Cl2 !" #" $" %" & !" & #" & $" & %" 155 Figure 5-2. UV-vis Spectra of [Os(phen) 3 ]Cl 2 (solid line) and [Os(phen) 3 ]Cl 3 (dashed line) in 0.01 M HCl solution. 380 450 520 590 660 730 0 4000 8000 12000 16000 20000 Os III Os II Wavelength, nm ! , M - 1 c m - 1 156 Table 5-1. UV-Visible Absorbance of Os Complexes in 0.01 M HCl Solution. " 430 , M ?1 cm ?1 " 480 , M ?1 cm ?1 " 550 , M ?1 cm ?1 Os II (1.80 ? 0.1) ! 10 4 1.67 ! 10 4 5.15 ! 10 3 Os III 1.61 ! 10 3 6.30 ! 10 2 9.83 ! 10 2 (" OsI ? " OsII ) 430 , M ?1 cm ?1 (" OsI ? " OsII ) 480 , M ?1 cm ?1 (" OsI ? " OsII ) 550 , M ?1 cm ?1 1.64 ! 10 4 1.61 ! 10 4 4.16 ! 10 3 ! 10 ?5 M of Os II at 550 nm at p[H + ] = 5.1 is shown in Figure 5-3b. A good-quality pseudo-second-order fit was also obtained. Os III Dependence. In the presence of more than 10-fold exces of Os II over Os III , the oxidations of 0.002 M phenol by various concentrations of Os III were studied at p[H + ] = 5.1. Kinetic traces of reactions were acquired at 550 nm. The observed first half-lives are summarized in Table A-28. A plot of half-life versus 1/[Os III ] 0 yielded a straight line, as shown in Figure 5-4, with a slope of (1.56 ? 0.03) ! 10 ?5 M s. This result implies that the rate law is second-order with respect to the initial concentration of Os III acording to eq 5-4. ? d[Os I ] t =kobs [ I ] 2 (5-4) Phenol Dependence. Under pseudo-second-order conditions, a series of reactions 157 Figure 5-3. Kinetic traces of the phenol oxidation by Os III . ? = 0.1 M (NaCl); T = 25?C. Lower box shows the experimental trace (solid line) and the pseudo-second-order fit (dashed line). Upper box shows the residuals of the fit. (a) Monitored at 480 nm; [Os III ] 0 = 0.67 ! 10 ?5 M; [Os II ] 0 = 1.56 ! 10 ?5 M; [phenol] = 0.02 M; p[H + ] = 4.1 (0.02 M acetate buffer). (b) Monitored at 550 nm; [Os III ] 0 = 0.79 ! 10 ?5 M; [Os II ] 0 = 1.27 ! 10 ?4 M; [phenol] = 0.002 M; p[H + ] = 5.1 (0.02 M acetate buffer). -2!10 -03 0 2!10 -03 R e s i d u a l s 0 1 2 3 4 5 0.25 0.28 0.31 0.34 0.37 Reaction trace Curve fit a Time, s A 480 -1!10 -03 0 1!10 -03 R e s i d u a l s 0 10 20 30 0.65 0.66 0.67 0.68 0.69 Reaction trace Curve fit b Time, s A 550 158 Figure 5-4. Plot of t 1/2 vs 1/[Os III ] 0 . Al the reactions were run under pseudo-second-order conditions at 550 nm. [Os III ] 0 = (0.49?1.15) ! 10 ?5 M; [Os II ] 0 = (12.2?12.7) ! 10 ?5 M; [phenol] = 0.002 M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 5.1 (0.02 M acetate buffer). Solid line is the linear fit. Data from Table A-28. 0.5 1.2 1.9 2.6 3.3 4.0 5.0!10 04 1.0!10 05 1.5!10 05 2.0!10 05 2.5!10 05 1/[Os III ] 0 , M -1 t 1/2 , s 159 betwen 1.21 ! 10 ?5 M of Os III and various concentrations of phenol (0.002?0.02 M) were followed at 550 nm at a p[H + ] of around 5.5 (acetate buffer) in the presence of 12.4 !10 ?5 M Os II . Figure 5-5 shows the kinetic trace of a reaction with 0.002 M phenol and its pseudo-second-order fit. Kinetic data of al the four reactions are summarized in Table A-29. The linear plot of the observed second-order rate constant versus [phenol] tot 2 , shown in Figure 5-6, yielded a slope of (9.20 ? 0.13) ! 10 9 M ?3 s ?1 . Thus, the rate law is second-order with respect to the total concentration of phenol as described in eq 5-5. k obs = k [phenol] tot 2 (5-5) Os II Dependence. The oxidations of 0.0025 or 0.005 M phenol by (0.47?1.38) ! 10 ? 5 M of Os III in the presence of (0.68?1.51) ! 10 ?5 M of Os II were studied in the presence of 0.02 M acetate buffer at p[H + ] = 4.7 and were followed at 480 nm. The first half-life and observed second-order rate constants are summarized in Table A-30. A linear relationship betwen k obs /[phenol] tot 2 and 1/[Os II ] 0 2 was observed and is shown in Figure 5-7 with a slope of (3.39 ? 0.11) M ?1 s ?1 and an intercept of (2.58 ? 0.16) ! 10 10 M ?3 s ?1 . It should be noted that under those conditions, for example, the oxidation by 1.38 ! 10 ?5 M of Os III in presence of 0.68 ! 10 ?5 M of Os II , the concentration of Os II during the reaction can not be treated as a constant value. And the rate constants obtained from pseudo-second-order kinetics are much slower than those without the influence of Os II . This would increase the slope and decrease the intercept in Figure 5-7. However, the observed linear relationship suggests that the rate law follows an inverse second-order dependence on the initial concentration of Os II acording to eq 5-6. 160 Figure 5-5. Kinetic trace of phenol oxidation by Os III at 550 nm. ? = 0.1 M (NaCl); T = 25?C. Lower box shows the experimental trace (solid line) and the pseudo-second-order fit (dashed line). Upper box shows the residuals in the fit. [Os III ] 0 = 1.21 ! 10 ?5 M; [Os II ] 0 = 12.4 ! 10 ?5 M; [phenol] = 0.002 M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 5.4 (0.02 M acetate buffer). -0.001 0.000 0.001 R e s i d u a l s 0 4 8 12 16 20 0.63 0.65 0.67 0.69 0.71 Reaction trace Curve fit Time, s A 550 161 Figure 5-6. Plot of k obs vs [phenol] tot 2 . Al the reactions were studied at 550 nm using pseudo-second-order conditions. [Os III ] 0 = 1.21 ! 10 ?5 M; [Os II ] 0 = 12.4 ! 10 ?5 M; [phenol] = (0.002?0.02) M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] values of 5.44?5.68 were maintained using a 0.02 M acetate buffer. Solid line is the linear fit. Data from Table A-29. 0.0000 0.0001 0.0002 0.0003 0.0004 0.0005 0 1!10 06 2!10 06 3!10 06 4!10 06 5!10 06 [Phenol] tot 2 , M 2 k o b s , M - 1 S - 1 162 Figure 5-7. Plot of k obs /[phenol] tot 2 vs 1/[Os II ] 0 2 . Al the reactions were run under ?pseudo-second-order? conditions at 480 nm. [Os III ] 0 = (0.47?1.38) ! 10 ?5 M; [Os II ] 0 = (0.68?1.51) ! 10 ?5 M; [phenol] tot = (2.5?5.0) ! 10 ?3 M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 4.7 (0.02 M acetate buffer). Solid line is the linear fit. Data from Table A-30. 0.0 7.0!10 09 1.4!10 10 2.1!10 10 2.8!10 10 2.5!10 10 4.5!10 10 6.5!10 10 8.5!10 10 1.1!10 11 1.2!10 11 1/[Os II ] 0 2 , M -2 k o b s / [ p h e n o l ] t o t 2 , M - 3 s - 1 163 kobs =' 1 [O I ]0 2 (5-6) p[H + ] Dependence. The phenol oxidations by (0.39?0.67) ! 10 ?5 M of Os III were studied over the p[H + ] range of 4.0?6.4 at 480 nm. The detailed reaction conditions and kinetic data are collected in Table A-31. Acording to the pre-equilibrium approximation and the mechanism proposed in the next section, a rearrangement of the k obs expresion (eq 5-18) as shown below yields [Os I ]0 2 kobs ArHt = dim( KArOH a+10 !p[] ) 2 (5-7) Here, K a is the acid disociation constant of phenol, pK a, ArOH = 9.79 at ? = 0.1 M 81 . k dim represents the dimerization rate constant of phenoxyl radicals. The value of K ArOH (= 1.1 ! 10 -9 M) was calculated from the redox potential of the C 6 H 5 O ? , H + /C 6 H 5 OH and Os III /Os II couples. E f (C 6 H 5 O ? /C 6 H 5 O ? ) = 0.797 V is corrected from 0.79 V at ? = 0.0 M 88 by means of the equation log ' = ?Az i 2 ? 1/2 /(1 + ? 1/2 ) and E f (C 6 H 5 O ? , H + /C 6 H 5 OH) = 1.376 V is calculated from E f (ArO ? , H + /ArOH) = E f (ArO ? /ArO ? ) + 0.0592pK a . The redox potential of the Os III /Os II couple was obtained from the supporting data of reference 56 as 0.846 V vs NHE at ? = 0.1 M. When plotting [Os II ] 0 2 k obs /[phenol] tot 2 versus p[H + ] and fiting the data to eq 5-7, a good quality fit was achieved with log 2k dim of 9.49, as shown in Figure 5-8. This value is close to the reported log 2k dim value of 9.36. 95 If we consider a 5 mV uncertainty in the redox potential for each couple the resulting uncertainty of log 164 Figure 5-8. Plot of [Os II ] 0 2 k obs /[phenol] tot 2 versus p[H + ]. Al the reactions were run under ?pseudo-second-order? conditions at 480 nm. [Os III ] 0 = (0.39?0.67) ! 10 ?5 M; [Os II ] 0 = (1.51?1.81) ! 10 ?5 M; [phenol] tot = (0.2?20) ! 10 ?3 M; ? = 0.1 M (NaCl); T = 25 ?C. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: acetate buffer for 4.0 < p[H + ] < 5.4, and cacodylate buffer for p[H + ] = 6.34. Solid line is fit to eq 5-7. Data from Table A-31. 3.5 4.2 4.9 5.6 6.3 7.0 10 -1 10 0 10 1 10 2 10 3 10 4 10 5 p[H + ] [ O s I I ] 0 2 k o b s / [ p h e n o l ] t o t 2 , M - 1 s - 1 165 2k dim is 0.30 . Spin Trapping Effect. The influence of spin trapping agent DBNBS was tested for the phenol oxidation by Os III . 1 or 10 mM of DBNBS was added into the reaction mixture containing 1.26 ! 10 ?5 M of Os III and 0.0025 M of phenol in the presence of 0.75 ! 10 ?5 M of Os II at p[H + ] = 4.71. Al data are summarized in Table 5-2. Table 5-2. Kinetic Data for the Reaction of Phenol with Os III in the Presence of Dipic or DBNBS. a [DBNBS], mM b t 1/2, s 0 0.11 1.0 0.08 10 0.03 [Dipic], mM c t 1/2, s 0 0.18 1.0 0.18 a Al the reactions were monitored at 480 nm. ? = 0.1 M (NaCl); T = 25?C. b [Os III ] 0 = (1.23?1.38) ! 10 ?5 M; [Os II ] 0 = (0.68?0.79) ! 10 ?5 M; [phenol] tot = 0.0025 M; p[H + ] = 4.71 (0.02 M acetate buffer). c [Os III ] 0 = 0.67 ! 10 ?5 M; [Os II ] 0 = 1.56 ! 10 ?5 M; [phenol] tot = 0.02 M; p[H + ] = 4.05 (0.02 M acetate buffer). A decrease of the first half-life is observed with increasing DBNBS concentration. A kinetic trace of the reaction with 10 mM DBNBS is shown in Figure 5-9. A good quality pseudo first-order fit was obtained at 480 nm, which implies that the phenoxyl radical is 166 Figure 5-9. Kinetic trace of phenol oxidation by Os III with DBNBS at 480 nm. ? = 0.1 M (NaCl); T = 25?C. Lower box shows the experimental trace (solid line) and the pseudo- first-order fit (dashed line). Upper box shows the residuals in the fit. [Os III ] 0 = 1.29 ! 10 ?5 M; [Os II ] 0 = 0.71 ! 10 ?5 M; [phenol] = 0.0025 M; [DBNBS] = 10 mM; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 4.71 (0.02 M acetate buffer). -0.005 0.000 0.005 R e s i d u a l s 0.00 0.12 0.24 0.36 0.48 0.60 0.0 0.1 0.2 0.3 Reaction trace First-order fit Time, s A 480 167 scavenged by the spin tapping agent DBNBS. Metal Catalysis. Acording to previous reported results on the oxidation of thioglycolic acid and cysteine, 73, 77 2,6-pyridinedicarboxylic acid (dipic) is an efective inhibitor of the copper catalysis. 1 ! 10 ?3 M of dipic was added to the reaction betwen 0.02 M of phenol and 0.67 ! 10 ?5 M of Os III at p[H + ] = 4.05 with 15.6 ! 10 ?5 M of Os II . The experiments were monitored at 480 nm and the data are collected in Table 5-2. No efect was observed, which suggests that no copper catalysis is involved in the oxidation of phenol by Os III , and that catalytic efects of other metals could also be excluded. 5.4 Discusion and Conclusion Acording to the above results, the following mechanism, eqs 5-8 to 5-11, is proposed to acount for the oxidation of phenol by Os III in the absence of spin traps: Os III + C 6 H 5 OH ! Os II + C 6 H 5 O ? + H + K ArOH , k ArOH , k ?ArOH (5-8) Os III + C 6 H 5 O ? ! Os II + C 6 H 5 O ? K ArO ?, k ArO ?, k ?ArO ? (5-9) C 6 H 5 OH ! C 6 H 5 O ? + H + K a (5-10) 2 ! C 6 H 5 O ? ! Coupling Products k dim (5-11) This mechanism looks identical to the outer-sphere proton-coupled electron-transfer oxidation of phenol by Ir IV as we described in Chapter 2. However, the kinetics is quite diferent: the first two steps (eqs 5-8 and 5-9) rapidly approach equilibrium instead of being the rate-limiting steps as shown in Ir IV reaction. Therefore, we propose a second- order rate law and this rate law can be derived as follows: 168 ? d[Os I ] t =2kdim [Ar ? ] 2 (5-12) Acording to the pre-equilibrium approximation, the expresion of concentration of the phenoxyl radical (eq 5-16) is obtained from eqs 5-13 to 5-15. KArOH = [s I ]r ? [ + ] (5-13) r- [ I ]r ? s ? (5-14) [ArOH] tot = [ArOH] + [ArO ? ] (5-15) [ArO ? = [s I rto I ] ( KArOH a+[] ) (5-16) Substitution of eq 5-16 into eq 5-12 gives the rate law (eq 5-17) and the observed rate constant can be expresed by eq 5-18. ? d[Os I ] t =2kdim [s I ] 2 ArOHto 2 I ( KArOH a+[] ) 2 (5-17) kobs di [r]to s I 0 2 ( r a[ + ] ) 2 (5-18) Thus, the rate law obtained for phenol oxidation by Os III follows an inverse second-order with respect to [Os II ] and [H + ], and second-order with respect to [Os III ] and [phenol] tot . 169 Kinetic simulations based on the above mechanism were performed with the model listed in Table 5-3. The experimental half-lives obtained in the absence of DBNBS at diferent p[H + ] values are consistent with those from simulations (Table 5-4). In the presence of spin trap agent DBNBS, three more steps were added to the mechanism: C 6 H 5 O ? + DBNBS ' adduct ? k DBNBS (5-19) Os III + adduct ? ' Os II + adduct + k adduct (5-20) DBNBS (dimer) ! 2 ! DBNBS K dim, DBNBS , k dim, DBNBS , k ?dim, DBNBS (5-21) As discussed previously, only the monomer of DBNBS can scavenge the phenoxyl radical, and the equilibrium constant, K dim, DBNBS , of the dimerization of DBNBS is 1.3 ! 10 ?3 M. A first-order rate law of phenol oxidation by Os III with DBNBS is proposed: ? d[Os I ] t = k DBNBS [DBNBS][ArO ? ] (5-22) [ I ] t =DBNS[] [Os I ArHto I ] ( KArOH a+[] ) (5-23) The simulation results of scavenging of the phenoxyl radical by DBNBS with a k DBNBS value (= 2.0 ! 10 5 M ?1 s ?1 ) together with those obtained from the phenol oxidation by Ir IV are shown in Table 5-5. Good agreements are obtained from the comparisons of the simulation results and the experimental data. The simulations are influenced by the removal of dimerization step in the presence of 1 mM DBNBS, whereas, no change is 170 Table 5-3. The Mechanism of Phenol Reaction and the Simulation Model. Equations Kinetic Parameter a Reactions in the model b Species eq 5-8 k ArOH = 5.7 k ?ArOH = 5.2 ! 10 9 A + B ! C + D + E k 1 C + D + E ! A + B k ?1 A = Os III B = C 6 H 5 OH eq 5-9 k ArO ? = 2.0 ! 10 9 k ?ArO ? = 3.0 ! 10 8 A + F ! C + D k 2 C + D ! A + F k ?2 C = Os II D = C 6 H 5 O ? eq 5-10 K a = 1.6 ! 10 ?10 B ! E + F k 3 E + F ! B k ?3 E = H + F = C 6 H 5 O ? K = K a /K w B + G ! F + H k 4 F + H ! B + G k ?4 G = OH ? H = H 2 O Buffer K a, buffer I ! J + E k 5 J + E ! I k ?5 I = Buffer J = Conjugate Base K = K a, buffer /K w I + G ! J + H k 6 J + H ! I + G k ?6 eq 5-11 k dim = 2.3 ! 10 9 2 * D ! K k 7 K = Coupling Products eq 5-19 k DBNBS = 2.0 ! 10 5 D + L ! M k 19 L = DBNBS eq 5-20 k adduct = 1.0 ! 10 7 A + M ! C + N k 20 M = Adduct ? eq 5-21 k dim, DBS = 1.3 ! 10 6 k ?dim, DBS = 1.0 ! 10 9 O ! 2 * L k 21 2 * L ! O k ?21 N = Adduct + O = DBNBS dimer a Rate constants k (M ?1 s ?1 ), acid disociation constant K a nd water disociation constant K w = 1.0 ! 10 ?14 . b k 3 , k ?3 , k 4 , k ?4 , k 5 , k ?5 , k 6 and k ?6 are the difusion-controlled forward and reverse rate constants obtained acording to the equilibriums. 171 Table 5-4. Comparison of the Experimental Data with the Simulation Results at diferent p[H + ]. a p[H + ] b [phenol] tot ! 10 3 , M t 1/2, exp , s t 1/2, sim , s 4.05 20 0.18 0.19 4.71 5.0 0.18 0.18 5.39 1.0 0.27 0.28 6.34 0.2 0.11 0.09 a [Os III ] 0 = (0.39?0.67) ! 10 ?5 M; [Os II ] 0 = (1.51?1.81) ! 10 ?5 M; The following buffers (0.02 M) were employed to maintain constant p[H + ] values: acetate buffer for 4.0 < p[H + ] < 5.4, and cacodylate buffer for p[H + ] = 6.34 . Table 5-5. Comparison of the Experimental Data to the Simulation Results in the Presence of Various Concentration of DBNBS. a [DBNBS], mM t 1/2, exp , s t 1/2, sim , s 1.0 0.08 0.11 10 0.03 0.05 a [Os III ] 0 = (1.23?1.29) ! 10 ?5 M; [Os II ] 0 = (0.71?0.79) ! 10 ?5 M; [phenol] tot = 0.0025 M; p[H + ] = 4.71 (0.02 M acetate buffer). 172 observed with 10 mM DBNBS. These results confirm the competition betwen scavenging the phenoxyl radical by DBNBS and their dimerization. Marcus Theory. In equations 1-12 to 1-15, k 12 is the cross electron-transfer rate constants, k 11 and k 22 are the self-exchange rate constants of C 6 H 5 O ? /C 6 H 5 O and Os III /Os II redox couples. k 11 is 1.9 ! 10 8 M ?1 s ?1 acording to literature 101 and a value for k 22 of 3 ! 10 8 M ?1 s ?1 is used in the calculation. 1 ! 10 11 M ?1 s ?1 is used for Z, the collision frequency. 99 Z i , Z j are the ionic charges of the reactants, R is the ideal gas constant, and r is the center to center distance betwen two reactants when they are approaching to each other. The radii of Os III and C 6 H 5 O ? are 6.7 56 ? and 2.5 ?, respectively, estimated from CPK Atomic Models. ? is the ionic strength. w ij is the electrostatic energy betwen reactants i and j. If the distance r is in angstroms and ? in molar, then w 12 can be calculated acording to eq 1-15 in kilojoules per mole. K 12 (K ArO ? = 7.0) was calculated from K ArO ? = K ArOH /K a with K ArOH = 1.1 ! 10 ?9 M and pK a = 9.79. With al these parameters k 12 (= k ArO ?) is calculated to be 2.1 ! 10 9 M ?1 s ?1 and k ?ArO ? is equal to 3.0 ! 10 8 M ?1 s ?1 . 173 References (1) Chance, B.; Nishimura, M. P. Natl. Acad. Sci. USA 1960, 46, 19. (2) Marcus, R. A.; Sutin, N. Biochim. Biophys. Acta 1985, 811, 265. (3) Stanbury, D. M. Adv. Inorg. Chem. 2003, 54, 351. (4) Kozik, M.; Baker, L. C. W. J. Am. Chem. Soc. 1990, 112, 7604. (5) Marcus, R. A. J. Phys. Chem. 1968, 72, 891. (6) Fukuzumi, S.; Wong, C. L.; Kochi, J. K. J. Am. Chem. Soc. 1980, 102, 2928. (7) Hwang, J. K.; Warshel, A. J. Am. Chem. Soc. 1987, 109, 715. (8) Arnaut, L. G.; Formosinho, S. J. J. Photochem. Photobiol. A: Chem. 1993, 75, 1. (9) Weinstock, I. A. Chem. Rev. 1998, 98, 113. (10) Roth, J. P.; Yoder, J. C.; Won, T. J.; Mayer, J. M. Science 2001, 294, 2524. (11) Barzykin, A. V.; Frantsuzov, P. A.; Seki, K.; Tachiya, M. Adv. Chem. Phys. 2002, 123, 511. (12) Marcus, R. A. J. Chem. Phys. 1965, 43, 679. (13) Taylor, S. M.; Halpern, J. J. Am. Chem. Soc. 1959, 81, 2933. (14) Mitchel, P. Nature 1961, 191, 144. (15) Binstead, R. A.; Moyer, B. A.; Samuels, G. J.; Meyer, T. J. J. Am. Chem. Soc. 1981, 103, 2897. (16) Fang, J. Y.; Hames-Schifer, S. J. Chem. Phys. 1997, 106, 8442. (17) Cukier, R. I.; Nocera, D. G. Annu. Rev. Phys. Chem. 1998, 49, 337. 174 (18) Stubbe, J.; Nocera, D. G.; Ye, C. S.; Chang, M. C. Y. Chem. Rev. 2003, 103, 2167. (19) Mayer, J. M.; Hrovat, D. A.; Thomas, J. L.; Borden, W. T. J. Am. Chem. Soc. 2002, 124, 11142. (20) Costentin, C.; Evans, D. H.; Robert, M.; Saveant, J. M.; Singh, P. S. J. Am. Chem. Soc. 2005, 127, 12490. (21) Siegbahn, P. E. M.; Blomberg, M. R. A. Chem. Rev. 2010, 110, 7040. (22) Mayer, J. M. Annu. Rev. Phys. Chem. 2004, 55, 363. (23) Huynh, M. H. V.; Meyer, T. J. Chem. Rev. 2007, 107, 5004. (24) Hames-Schifer, S.; Stuchebrukhov, A. A. Chem. Rev. 2010, 110, 6939. (25) Bonin, J.; Costentin, C.; Louault, C.; Robert, M.; Routier, M.; Saveant, J. M. P. Natl. Acad. Sci. USA 2010, 107, 3367. (26) Bonin, J.; Costentin, C.; Louault, C.; Robert, M.; Saveant, J. M. J. Am. Chem. Soc. 2011, 133, 6668. (27) Song, N.; Stanbury, D. M. Inorg. Chem. 2008, 47, 11458. (28) Abraham, M. H.; Abraham, R. J.; Byrne, J.; Grifiths, L. J. Org. Chem. 2006, 71, 3389. (29) Tyman, J. H. P. Synthetic and natural phenols; Elsevier: Amsterdam [The Netherlands] ; New York, 1996. (30) Elestad, G. A.; Kunstmann, M. P.; Whaley, H. A.; Paterson, E. L. J. Am. Chem. Soc. 1968, 90, 1325. (31) Marrian, G. F.; Haslewood, G. A. D. Biochem. J. 1932, 26, 1227. (32) Robertson, A.; Robinson, R. J. Chem. Soc. 1928, 1460. (33) Fomum, Z. T.; Ayafor, J. F.; Wandji, J. J. Nat. Prod. 1987, 50, 921. (34) Feldberg, W.; Toh, C. C. J Physiol-London 1953, 119, 352. (35) Titeler, M.; Lyon, R. A.; Glennon, R. A. Psychopharmacology 1988, 94, 213. (36) Johnson, M. P.; Hoffman, A. J.; Nichols, D. E. Eur. J. Pharmacol. 1986, 132, 269. 175 (37) Guiton, J.; Faveta, P.; Degoute, C. S.; Perdrix, J. P.; Dufresne, C.; Boulieu, R. Brit. J. Anaesth. 2002, 88, 653. (38) Higashiyama, S.; Iwabuki, H.; Morimoto, C.; Hieda, M.; Inoue, H.; Matsushita, N. Cancer Sci. 2008, 99, 214. (39) Betelheim, F. R. J. Am. Chem. Soc. 1954, 76, 2838. (40) Blomberg, M. R. A.; Siegbahn, P. E. M. BBA-Bioenergetics 2004, 1655, 45. (41) Babcock, G. T.; Tommos, C. BBA-Bioenergetics 2000, 1458, 199. (42) Rappaport, F.; Guergova-Kuras, M.; Nixon, P. J.; Diner, B. A.; Lavergne, J. Biochemistry 2002, 41, 8518. (43) Haumann, M.; Mulkidjanian, A.; Junge, W. Biochemistry 1999, 38, 1258. (44) Yokoyama, K.; Uhlin, U.; Stubbe, J. J. Am. Chem. Soc. 2010, 132, 8385. (45) Seyedsayamdost, M. R.; Yee, C. S.; Reece, S. Y.; Nocera, D. G.; Stubbe, J. J. Am. Chem. Soc. 2006, 128, 1562. (46) Reece, S. Y.; Stubbe, J.; Nocera, D. G. BBA-Bioenergetics 2005, 1706, 232. (47) Fecenko, C. J.; Thorp, H. H.; Meyer, T. J. J. Am. Chem. Soc. 2007, 129, 15098. (48) Irebo, T.; Reece, S. Y.; Sjodin, M.; Nocera, D. G.; Hamarstrom, L. J. Am. Chem. Soc. 2007, 129, 15462. (49) Reece, S. Y.; Nocera, D. G. J. Am. Chem. Soc. 2005, 127, 9448. (50) Sjodin, M.; Irebo, T.; Utas, J. E.; Lind, J.; Merenyi, G.; Akermark, B.; Hamarstrom, L. J. Am. Chem. Soc. 2006, 128, 13076. (51) Cecil, R.; Litler, J. S. J. Chem. Soc. B 1968, 1420. (52) Costentin, C.; Robert, M.; Saveant, J. M. J. Am. Chem. Soc. 2007, 129, 5870. (53) Stanbury, D. M.; Wilmarth, W. K.; Khalaf, S.; Po, H. N.; Byrd, J. E. Inorg. Chem. 1980, 19, 2715. (54) Stanbury, D. M.; Lednicky, L. A. J. Am. Chem. Soc. 1984, 106, 2847. (55) Sarala, R.; Stanbury, D. M. Inorg. Chem. 1990, 29, 3456. (56) Sarala, R.; Rabin, S. B.; Stanbury, D. M. Inorg. Chem. 1991, 30, 3999. 176 (57) Hung, M. L.; Mckee, M. L.; Stanbury, D. M. Inorg. Chem. 1994, 33, 5108. (58) Sun, J. F.; Stanbury, D. M. J. Chem. Soc., Dalton Trans. 2002, 785. (59) Hung, M. L.; Stanbury, D. M. Inorg. Chem. 2005, 44, 9952. (60) Makarycheva-Mikhailova, A. V.; Stanbury, D. M.; Mckee, M. L. J. Phys. Chem. B 2007, 111, 6942. (61) Kauffman, G. B.; Teter, L. A. Inorg. Synth. 1966, 8, 223. (62) Perrin, D. D.; Armarego, W. L. F. Purification of Laboratory Chemicals, 3rd Edition; Pergamon Pres, 1988. (63) Braithwaite, A.; Henthorn, K.; Eliott, G. E. P.; Marshal, B. W.; Smith, A. C.; Emslie, J. J.; Kaseke, C. T.; Tyson, J. F.; Ham, G.; Grifiths, P. D. Anal. Proc. 1981, 18, 60. (64) Kaur, H.; Leung, K. H. W.; Perkins, M. J. J. Chem. Soc., Chem. Commun. 1981, 142. (65) Hamilton, L.; Nielsen, B. R.; Davies, C. A.; Symons, M. C. R.; Winyard, P. G. Fre Rad. Res. 2003, 37, 41. (66) Kanetani, F.; Yamaguchi, H. B. Chem. Soc. Jpn. 1981, 54, 3048. (67) Konaka, R.; Sakata, S. Chem. Let. 1982, 411. (68) Amblard, F.; Govindarajan, B.; Lefkove, B.; Rapp, K. L.; Detorio, M.; Arbiser, J. L.; Schinazi, R. F. Bioorg. Med. Chem. Let. 2007, 17, 4428. (69) Padwa, A.; Wisnief, T. J.; Walsh, E. J. J. Org. Chem. 1989, 54, 299. (70) Spartan '08; Wavefunction, Inc: Irvine, CA. (71) Bruhn, H.; Nigam, S.; Holzwarth, J. F. Faraday Discuss. 1982, 74, 129. (72) Bonin, J.; Costentin, C.; Robert, M.; Saveant, J. M. Org. Biomol. Chem. 2011, 9, 4064. (73) Hung, M. L.; Stanbury, D. M. Inorg. Chem. 2005, 44, 3541. (74) Ram, M. S.; Stanbury, D. M. Inorg. Chem. 1985, 24, 4233. (75) Huang, H.; Sommerfeld, D.; Dunn, B. C.; Eyring, E. M.; Lloyd, C. R. J. Phys. Chem. A 2001, 105, 3536. 177 (76) Wang, X. G.; Stanbury, D. M. Inorg. Chem. 2008, 47, 1224. (77) Saha, B.; Hung, M. L.; Stanbury, D. M. Inorg. Chem. 2002, 41, 5538. (78) O'Brien, P.; Salacinski, H. J. Arch. Toxicol. 1996, 70, 787. (79) Zalomaeva, O. V.; Trukhan, N. N.; Ivanchikova, I. D.; Panchenko, A. A.; Roduner, E.; Talsi, E. P.; Sorokin, A. B.; Rogov, V. A.; Kholdeeva, O. A. J. Mol. Catal. A: Chem. 2007, 277, 185. (80) Ide, H.; Hagi, A.; Ohsumi, S.; Murakami, A.; Makino, K. Biochem. Int. 1992, 27, 367. (81) Martel, A. E.; Smith, R. M.; Motekaitis, R. J. NIST Critically Selected Stability Constants of Metal Complexes Database, 7.0 U. S. Department of Commerce: Gaithersburg, MD, 2003. (82) Fine, D. A. Inorg. Chem. 1969, 8, 1014. (83) Ye, M.; Schuler, R. H. J. Phys. Chem. 1989, 93, 1898. (84) Das, T. N. J. Phys. Chem. A 2001, 105, 5954. (85) Jonsson, M.; Lind, J.; Merenyi, G. J. Phys. Chem. A 2002, 106, 4758. (86) Jonsson, M.; Lind, J.; Merenyi, G. J. Phys. Chem. A 2003, 107, 5878. (87) Stradins, J.; Hasanli, B. J. Electroanal. Chem. 1993, 353, 57. (88) Lind, J.; Shen, X.; Eriksen, T. E.; Merenyi, G. J. Am. Chem. Soc. 1990, 112, 479. (89) Papina, A. A.; Koppenol, W. H. Chem. Res. Toxicol. 2007, 20, 2021. (90) Al-Ajlouni, A.; Bakac, A.; Espenson, J. H. Inorg. Chem. 1993, 32, 5792. (91) Al-Ajlouni, A. M.; Shawakfeh, K. Q.; Rajal, R. Kinet. Catal. 2009, 50, 88. (92) Pelizeti, E.; Mentasti, E. J. Inorg. Nucl. Chem. 1977, 39, 2227. (93) Margerum, D. W.; Chelappa, K. L.; Bossu, F. P.; Burce, G. L. J. Am. Chem. Soc. 1975, 97, 6894. (94) Neta, P.; Grodkowski, J. J. Phys. Chem. Ref. Data 2005, 34, 109. (95) Tripathi, G. N. R.; Schuler, R. H. Chem. Phys. Let. 1982, 88, 253. 178 (96) Binstead, R. A.; Jung, B.; Zuberbuhler, A. D. Specfit/32 Global Analysis System, version 3.0; Spectrum Software Asociates: Marlborough, MA USA, 2000. (97) Zuckerman, J. J. Inorganic Reactions and Methods; VCH: Derfield Beach, FL, 1986; Vol. 15. (98) Hurwitz, P.; Kustin, K. Trans. Faraday Soc. 1966, 62, 427. (99) Ram, M. S.; Stanbury, D. M. J. Phys. Chem. 1986, 90, 3691. (100) McDonald, W. J.; Einarsdottir, O. J. Phys. Chem. A 2008, 112, 11400. (101) Schuler, R. H.; Neta, P.; Zemel, H.; Fesenden, R. W. J. Am. Chem. Soc. 1976, 98, 3825. (102) Hansch, C.; Leo, A.; Taft, R. W. Chem. Rev. 1991, 91, 165. (103) Kaji, M.; Ogami, K.; Endo, T. J. Appl. Polym. Sci. 1999, 72, 953. (104) Saito, K.; Tago, T.; Masuyama, T.; Nishide, H. Angew. Chem. Int. Ed. 2004, 43, 730. (105) Wang, X. G.; Stanbury, D. M. J. Phys. Chem. A 2004, 108, 7637. (106) Song, N.; Stanbury, D. M. Inorg. Chem., Acepted. (107) Brown, H. C.; Okamoto, Y. J. Am. Chem. Soc. 1958, 80, 4979. (108) Roder, M.; Foldiak, G.; Wojnarovits, L. Radiat. Phys. Chem. 1999, 55, 515. (109) Jonsson, M.; Lind, J.; Reitberger, T.; Eriksen, T. E.; Merenyi, G. J. Phys. Chem. 1993, 97, 8229. (110) Jonsson, M.; Lind, J.; Eriksen, T. E.; Merenyi, G. J. Chem. Soc. Perk. T. 2 1993, 1567. (111) Das, T. N. J. Phys. Chem. A 2005, 109, 3344. (112) Dixon, W. T.; Murphy, D. J. Chem. Soc. Farad. T. 2 1978, 74, 432. (113) Hoigne, J.; Bader, H. Water Res. 1994, 28, 45. (114) Tratnyek, P. G.; Hoigne, J. Water Res. 1994, 28, 57. (115) Saveant, J. M.; Bonin, J.; Costentin, C.; Louault, C.; Robert, M. J. Am. Chem. Soc. 2011, 133, 6668. 179 (116) Young, R. C.; Kene, F. R.; Meyer, T. J. J. Am. Chem. Soc. 1977, 99, 2468. (117) Stubbe, J.; Riggs-Gelasco, P. Trends Biochem. Sci. 1998, 23, 438. (118) Michon, T.; Chenu, M.; Kelershon, N.; Desmadril, M.; Gueguen, J. Biochemistry 1997, 36, 8504. (119) Calahan, R. W.; Brown, G. M.; Meyer, T. J. Inorg. Chem. 1975, 14, 1443. 180 Appendix A EXPERIMENTAL DETAILS IN CHAPTERS 2?5 Table A-1. Kinetic Dependence of Phenol Oxidation on [Phenol] tot . a p[H + ] = 1.3 b , with DBNBS p[H + ] = 5.1 c , no DBNBS [phenol] tot ! 10 3 , M k obs ! 10 3 , s ?1 [phenol] tot ! 10 3 , M k obs , s ?1 22.2 14.4 1.77 0.270 44.3 27.8 8.86 1.13 266 163 22.2 2.64 443 272 44.3 4.78 a Al the reactions were run under pseudo-first-order conditions. [Ir IV ] 0 = 1 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b [HClO 4 ] = 0.05 M; [phenol] tot = (22.2?443) ! 10 ?3 M; [DBNBS] = 10 mM. c p[H + ] = 5.1 (0.02 M acetate buffer); [phenol] tot = (1.77?44.3) ! 10 ? 3 M. 181 Table A-2. Kinetic Dependence of Phenol Oxidation on p[H + ] Without DBNBS. a p[H + ] b [phenol] tot ! 10 3 , M k obs /[phenol] tot , M ?1 s ?1 2.46 44.3 7.26 ! 10 ?1 2.57 44.3 8.16 ! 10 ?1 2.71 44.3 9.58 ! 10 ?1 2.86 44.3 1.15 2.97 44.3 1.30 3.01 44.3 1.43 3.16 44.3 1.80 3.39 44.3 2.71 3.62 44.3 4.46 3.80 44.3 6.12 3.99 44.3 8.87 4.19 44.3 13.1 4.37 44.3 19.8 4.59 44.3 30.4 4.78 44.3 47.9 4.98 44.3 70.7 5.18 44.3 108 5.36 44.3 170 5.64 4.43 353 5.76 4.43 498 5.96 4.43 771 6.18 4.43 1.24 ! 10 3 6.35 4.43 1.84 ! 10 3 6.57 4.43 3.12 ! 10 3 6.74 4.43 4.67 ! 10 3 a Al the reactions were run under pseudo-first-order conditions. [Ir IV ] 0 = 1 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.6 < p[H + ] < 7.0. 182 Table A-3. Kinetic Dependence of Phenol Oxidation on p[H + ] with DBNBS. a p[H + ] b [phenol] tot ! 10 3 , M k obs /[phenol] tot , M ?1 s ?1 1.00 44.3 6.80 ! 10 ?1 1.20 44.3 6.71 ! 10 ?1 1.40 44.3 7.51 ! 10 ?1 1.60 44.3 7.55 ! 10 ?1 1.80 44.3 8.45 ! 10 ?1 2.00 44.3 9.50 ! 10 ?1 2.20 44.3 1.10 2.61 44.3 1.39 2.68 44.3 1.49 2.74 44.3 1.60 2.96 44.3 2.18 3.26 44.3 3.72 3.38 44.3 4.00 3.65 44.3 7.53 3.80 44.3 10.4 3.92 44.3 13.3 4.09 44.3 19.5 4.36 44.3 31.3 4.49 44.3 41.8 4.67 44.3 61.9 4.83 44.3 85.9 4.99 44.3 121 5.16 44.3 161 5.41 4.43 352 5.63 4.43 574 5.85 4.43 922 6.08 4.43 1.40 ! 10 3 6.24 4.43 2.04 ! 10 3 6.41 4.43 2.97 ! 10 3 6.60 4.43 4.18 ! 10 3 6.73 4.43 5.66 ! 10 3 6.94 1.77 9.75 ! 10 3 a [Ir IV ] 0 = 1 ! 10 ?4 M; [DBNBS] = 10 mM; ? = 0.1 M (LiClO 4 ); T = 25 !C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.4 < p[H + ] < 7.0. 183 Table A-4. Nonlinear-Least-Squares Regresion Results of the k obs /[phenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Absence of DBNBS. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 0.540 0.395 k ArO ?, M ?1 s ?1 5.05 ! 10 6 4.90 ! 10 6 k?, M ?1 s ?1 6.54 ! 10 ?3 Std. Error k ArOH , M ?1 s ?1 2.59 ! 10 ?2 6.36 ! 10 ?2 k ArO ?, M ?1 s ?1 7.80 ! 10 4 9.51 ! 10 4 k?, M ?1 s ?1 2.74 ! 10 ?3 95 % Confidence Intervals k ArOH , M ?1 s ?1 0.486 to 0.594 0.264 to 0.527 k ArO ?, M ?1 s ?1 4.89 ! 10 6 to 5.21 ! 10 6 4.70 ! 10 6 to 5.09 ! 10 6 k?, M ?1 s ?1 8.68 ! 10 ?4 to 1.22 ! 10 ?2 Goodnes of Fit Degrees of Freedom 23 22 R squared 0.9992 0.9994 Weighted Sum of Squares (1/(Y*Y)) 8.93 ! 10 ?2 7.02 ! 10 ?2 Sy.x 6.23 ! 10 ?2 5.65 ! 10 ?2 184 Table A-5. Nonlinear-Least-Squares Regresion Results of the k obs /[phenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 0.769 0.594 k ArO ?, M ?1 s ?1 8.02 ! 10 6 7.32 ! 10 6 k?, M ?1 s ?1 2.01 ! 10 ?2 Std. Error k ArOH , M ?1 s ?1 3.316 ! 10 ?2 2.93 ! 10 ?2 k ArO ?, M ?1 s ?1 2.03 ! 10 5 1.47 ! 10 5 k?, M ?1 s ?1 2.80 ! 10 ?3 95 % Confidence Intervals k ArOH , M ?1 s ?1 0.701 to 0.836 0.534 to 0.654 k ArO ?, M ?1 s ?1 7.60 ! 10 6 to 8.44 ! 10 6 7.01 ! 10 6 to 7.61 ! 10 6 k?, M ?1 s ?1 1.44 ! 10 ?2 to 2.59 ! 10 ?2 Goodnes of Fit Degrees of Freedom 30 29 R squared 0.9978 0.9992 Weighted Sum of Squares (1/(Y*Y)) 0.382 0.138 Sy.x 0.113 6.90 ! 10 ?2 185 Table A-6. Comparison of k obs in H 2 O and in D 2 O. a [phenol] tot ! 10 3 , M k obs ! 10 3 , s ?1 , D 2 O k obs ! 10 3 , s ?1 , H 2 O KIE = k H /k D 44.3 8.35 27.9 3.34 133 24.8 96.1 3.87 222 43.9 142 3.22 a [Ir IV ] 0 = 1 ! 10 ?4 M; [HClO 4 ] = 0.09 M; ? = 0.1 M; [DBNBS] = 10 mM; T = 25 ?C. Table A-7. Kinetic Dependence on [Cresol] tot . a [cresol] tot ! 10 3 , M k obs ! 10 2 , s ?1 1.0 1.48 5.0 7.12 10 14.2 20 28.2 a [Ir IV ] 0 = 1 ! 10 ?4 M; [DBNBS] = 1 mM; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 186 Table A-8. Kinetic Dependence of Biphenols and 4-Phenoxyphenol Oxidation on p[H + ]. a Substrate p[H + ] b [Substrate] tot ! 10 4 , M k obs /[Substrate] tot , M ?1 s ?1 k obs , 2 , c s ?1 4,4?-biphenol 1.30 1.0 4.06 ! 10 4 4.37 1.0 5.54 ! 10 4 5.94 1.0 1.74 ! 10 5 6.91 0.50 1.16 ! 10 6 2,2?-biphenol 1.30 14 8.57 3.81 10 8.81 ! 10 2 5.34 10 2.05 ! 10 4 6.72 3.0 3.47 ! 10 5 2,4?-biphenol d 1.30 2.0 3.54 ! 10 3 3.51 !10 ?2 2.48 2.5 3.22 ! 10 3 5.57 !10 ?2 3.83 1.0 5.63 ! 10 3 2.52 !10 ?2 5.29 1.0 8.69 ! 10 3 3.89 !10 ?2 6.85 2.0 2.02 ! 10 5 1.20 4-phenoxy phenol 1.30 3.0 7.43 ! 10 2 2.46 3.0 1.31 ! 10 3 3.59 3.0 1.43 ! 10 3 4.44 3.0 1.90 ! 10 3 4.95 3.0 2.62 ! 10 3 6.41 3.0 3.45 ! 10 4 7.06 3.0 1.45 ! 10 5 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C; b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.6, acetate buffer for 3.7 < p[H + ] < 5.4, and cacodylate buffer for 5.5 < p[H + ] < 7.1. c Slow observed first-order rate constant from two-exponential fit. d Two-exponential fit was applied to biphasic trace. 187 Table A-9. Kinetic Dependence of Cresol Oxidation on p[H + ] with DBNBS. a p[H + ] [cresol] tot ! 10 3 , M [Ir IV ] 0 ! 10 5 , M k obs /[cresol] tot , M ?1 s ?1 1.30 20 2.5 14.2 1.60 20 2.5 15.1 2.00 20 2.5 16.3 2.39 20 2.5 17.1 2.68 10 10 17.0 3.04 10 10 20.8 3.47 10 10 30.9 3.73 10 5.0 41.4 4.12 10 5.0 79.0 4.33 1.0 5.0 134 4.53 1.0 2.5 186 4.93 1.0 5.0 432 5.33 1.0 5.0 965 5.78 1.0 5.0 2.42 ! 10 3 6.18 1.0 5.0 5.55 ! 10 3 6.63 1.0 5.0 1.42 ! 10 4 7.08 1.0 5.0 3.88 ! 10 4 a [DBNBS] = 1 mM; ? = 0.1 M (LiClO 4 ); T = 25 !C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.3 < p[H + ] < 3.5, acetate buffer for 3.7 < p[H + ] < 5.4, and cacodylate buffer for 5.7 < p[H + ] < 7.1. 188 Table A-10. Nonlinear-Least-Squares Regresion Results of the k obs /[cresol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 15.8 13.2 k ArO ?, M ?1 s ?1 5.24 ! 10 7 4.54 ! 10 7 k?, M ?1 s ?1 0.168 Std. Error k ArOH , M ?1 s ?1 0.915 0.864 k ArO ?, M ?1 s ?1 2.48 ! 10 6 2.16 ! 10 6 k?, M ?1 s ?1 4.02 ! 10 ?2 95 % Confidence Intervals k ArOH , M ?1 s ?1 13.8 to 17.7 11.3 to 15.0 k ArO ?, M ?1 s ?1 4.72 ! 10 7 to 5.77 ! 10 7 4.07 ! 10 7 to 5.00 ! 10 7 k?, M ?1 s ?1 8.14 ! 10 ?2 to 0.254 Goodnes of Fit Degrees of Freedom 15 14 R squared 0.9975 0.9988 Weighted Sum of Squares (1/(Y*Y)) 0.280 0.132 Sy.x 0.137 9.71 ! 10 ?2 189 Table A-11. Kinetic Dependence of Cresol Oxidation on Temperature. a T, ?C k obs , s ?1 8.0 0.072 15 0.114 25 0.221 35 0.393 45 0.653 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [cresol] tot = 0.02 M; [HClO 4 ] = 0.01 M; ? = 0.1 M (LiClO 4 ). Table A-12. Kinetic Isotope Efect for Oxidation of Cresol. a Expt. No. k obs , s ?1 , D 2 O k obs ! 10 3 , s ?1 , H 2 O KIE = k H /k D 1 b 0.077 0.240 3.12 2 b 0.077 0.239 3.10 3 c 0.087 0.238 2.74 4 c 0.086 0.239 2.78 a [Ir IV ] 0 = 1 ! 10 ?4 M; [cresol] = 0.02 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b [HClO 4 ] = 0.05 M; [DBNBS] = 1 mM. c p[H + ] = 2.34 (0.02 M monochloroacetate buffer). 190 Table A-13. Kinetic Dependence on [Xylenol] tot . a [xylenol] tot ! 10 3 , M k obs , s ?1 2.50 0.096 8.00 0.305 13.0 0.444 19.0 0.653 25.0 0.914 a [Ir IV ] 0 = 1 ! 10 ?4 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 191 Table A-14. Kinetic Dependence of Xylenol Oxidation on p[H + ]. a p[H + ] [Ir IV ] 0 ! 10 5 , M [Xylenol] tot ! 10 4 , M k obs /[xylenol] tot , M ?1 #s ?1 1.30 10 50 40.1 1.60 10 50 40.7 2.00 10 50 43.0 2.40 10 50 47.4 2.67 10 50 48.6 2.83 10 50 51.3 3.00 10 50 56.4 3.18 10 50 63.7 3.34 10 50 73.3 3.58 10 50 86.7 3.73 10 50 104 3.93 10 50 132 4.13 10 50 174 4.39 10 50 264 4.53 10 50 350 4.75 10 25 590 4.93 10 25 828 5.14 10 25 1.24 ! 10 3 5.36 10 25 1.86 ! 10 3 5.53 2.5 5.0 2.85 ! 10 3 5.73 2.5 5.0 4.17 ! 10 3 5.94 2.5 5.0 6.48 ! 10 3 6.16 2.5 5.0 1.01 ! 10 4 6.35 2.5 2.5 1.69 ! 10 4 6.54 2.5 2.5 2.41 ! 10 4 6.74 2.5 2.5 3.65 ! 10 4 a ? = 0.1 M (LiClO 4 ); T = 25 ?C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.5 < p[H + ] < 5.4, and cacodylate buffer for 5.5 < p[H + ] < 7.0. 192 Table A-15. Nonlinear-Least-Squares Regresion Results of the k obs /[xylenol] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 46.7 37.6 k ArO ?, M ?1 s ?1 1.97 ! 10 8 1.71 ! 10 8 k?, M ?1 s ?1 0.403 Std. Error k ArOH , M ?1 s ?1 1.83 1.06 k ArO ?, M ?1 s ?1 6.09 ! 10 6 3.25 ! 10 6 k?, M ?1 s ?1 3.76 ! 10 ?2 95 % Confidence Intervals k ArOH , M ?1 s ?1 43.0 to 50.5 35.4 to 39.8 k ArO ?, M ?1 s ?1 1.84 ! 10 8 to 2.09 ! 10 8 1.64 ! 10 8 to 1.77 ! 10 8 k?, M ?1 s ?1 0.325 to 0.481 Goodnes of Fit Degrees of Freedom 24 23 R squared 0.9974 0.9996 Weighted Sum of Squares (1/(Y*Y)) 0.293 5.00 ! 10 ?2 Sy.x 0.110 4.66 ! 10 ?2 193 Table A-16. Kinetic Isotope Efect for Oxidation of Xylenol. a k obs , s ?1 , D 2 O k obs ! 10 3 , s ?1 , H 2 O KIE = k H /k D 0.337 0.884 2.62 0.354 0.913 2.58 0.359 0.926 2.58 0.347 0.912 2.62 a [Ir IV ] 0 = 1 ! 10 ?4 M; [xylenol] = 0.025 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. Table A-17. Kinetic Dependence on [TMP] tot . a [TMP] tot , ! 10 3 M k obs , s ?1 0.3 0.32 1.2 1.27 2.1 2.23 3.0 3.27 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [HClO 4 ] = 0.01 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 194 Table A-18. p[H + ] Dependence of TMP Oxidation. a p[H + ] b k obs /[TMP] tot , M ?1 s ?1 1.30 1.06 ! 10 3 1.60 1.07 ! 10 3 2.00 1.07 ! 10 3 2.41 1.08 ! 10 3 2.64 1.09 ! 10 3 3.01 1.10 ! 10 3 3.36 1.11 ! 10 3 3.76 1.11 ! 10 3 4.13 1.17 ! 10 3 4.53 1.31 ! 10 3 4.94 1.69 ! 10 3 5.38 2.55 ! 10 3 5.75 4.42 ! 10 3 6.18 9.64 ! 10 3 6.35 1.45 ! 10 4 6.93 4.88 ! 10 4 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [TMP] tot = 3.0 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 !C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.5, acetate buffer for 3.6 < p[H + ] < 5.5, and cacodylate buffer for 5.6 < p[H + ] < 7.0. 195 Table A-19. Kinetic Isotope Efect for Oxidation of TMP. a [TMP] tot ! 10 3 , M k obs , s ?1 D 2 O k obs , s ?1 H 2 O KIE = k H /k D 3.0 1.74 3.55 2.05 3.0 1.73 3.56 2.06 1.5 0.82 1.67 2.02 a [Ir IV ] 0 = 1 ! 10 ?4 M; [HClO 4 ] = 0.01 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. Table A-20. Kinetic Isotope Efect for Oxidation of MOP. a k obs , s ?1 , D 2 O k obs , s ?1 , H 2 O KIE = k H /k D 7.44 15.1 2.03 7.63 14.7 1.93 7.71 14.4 1.87 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; [HClO 4 ] = 0.05 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C. 196 Table A-21. Kinetic Dependence of MOP Oxidation on p[H + ]. a p[H + ] b k obs /[MOP] tot , M ?1 #s ?1 1.30 5.68 ! 10 4 1.60 5.80 ! 10 4 2.00 5.88 ! 10 4 2.97 6.05 ! 10 4 3.36 6.01 ! 10 4 3.79 6.09 ! 10 4 4.18 6.25 ! 10 4 4.58 6.31 ! 10 4 4.98 6.62 ! 10 4 5.41 7.34 ! 10 4 5.75 8.26 ! 10 4 6.15 1.17 ! 10 5 6.54 1.91 ! 10 5 6.94 3.61 ! 10 5 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; ? = 0.1 M (LiClO 4 ); T = 25 ?C; b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.5, and cacodylate buffer for 5.6 < p[H + ] < 7.0. 197 Table A-22. Kinetic Dependence of MOP Oxidation on Temperature. a T, ?C k obs , s ?1 8.0 8.11 15 11.6 25 16.6 35 22.7 45 28.9 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [MOP] tot = 3 ! 10 ?4 M; [HClO 4 ] = 0.1 M; T = 25 ?C. 198 Table A-23. Kinetic Dependence of TBP Oxidation on p[H + ]. a p[H + ] [TBP] tot ! 10 3 , M k obs /[TBP] tot , M ?1 s ?1 1.30 0.9 19.7 1.60 0.9 22.0 2.00 0.9 24.1 2.30 0.7 26.0 2.52 0.7 27.5 2.80 0.7 30.0 3.05 0.7 34.9 3.12 0.5 37.7 3.52 0.5 53.1 3.87 0.5 75.0 4.26 0.5 123 4.65 0.5 213 4.93 0.5 331 5.26 0.5 572 5.50 0.3 987 5.91 0.3 2.12!10 3 6.23 0.3 4.80!10 3 6.66 0.3 9.99!10 3 7.02 0.3 2.04!10 4 a [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 1 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.5 < p[H + ] < 5.3, and cacodylate buffer for 5.5 < p[H + ] < 7.1. 199 Table A-24. Nonlinear-Least-Squares Regresion Results of the k obs /[TBP] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 28.8 18.2 k ArO ?, M ?1 s ?1 3.61 ! 10 7 2.40 ! 10 7 k?, M ?1 s ?1 0.483 Std. Error k ArOH , M ?1 s ?1 2.62 0.744 k ArO ?, M ?1 s ?1 3.34 ! 10 6 8.71 ! 10 5 k?, M ?1 s ?1 3.05 ! 10 ?2 95 % Confidence Intervals k ArOH , M ?1 s ?1 23.3 to 34.3 16.6 to 19.7 k ArO ?, M ?1 s ?1 2.91 ! 10 7 to 4.32 ! 10 7 2.21 ! 10 7 to 2.58 ! 10 7 k?, M ?1 s ?1 0.418 to 0.548 Goodnes of Fit Degrees of Freedom 17 16 R squared 0.9892 0.9994 Weighted Sum of Squares (1/(Y*Y)) 1.14 5.98 ! 10 ?2 Sy.x 0.258 6.11 ! 10 ?2 200 Table A-25. Kinetic Dependence on [Ac-Y-NH 2 ] tot . a [Ac-Y-NH 2 ] tot ! 10 3 , M k obs , s ?1 0.8 0.010 2.4 0.031 4.0 0.053 5.6 0.072 7.2 0.093 a Al the reactions were run under pseudo-first-order conditions. [Ac-Y-NH 2 ] tot = (8.0?72) ! 10 ?4 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 5 mM; p[H + ] = 2.8 (0.02 M monochloroacetate buffer); ? = 0.1 M (LiClO 4 ); T = 25 ?C. 201 Table A-26. Kinetic Dependence of Ac-Y-NH 2 Oxidation on p[H + ]. a p[H + ] [Ac-Y-NH 2 ] tot ! 10 3 , M k obs /[Ac-Y-NH 2 ] tot , M ?1 s ?1 1.30 6.40 3.98 1.60 6.40 4.72 2.00 6.40 5.47 2.42 2.25 7.44 2.87 4.80 11.3 3.17 4.80 18.0 3.22 2.10 19.7 3.61 2.10 34.0 3.77 2.10 46.5 4.15 2.10 99.4 4.54 1.20 230 4.96 1.20 518 5.33 1.20 1.10!10 3 5.92 0.3 3.97!10 3 6.33 0.3 1.01!10 4 6.71 0.3 2.26!10 4 6.89 0.3 3.56!10 4 7.19 0.3 5.99!10 4 a [Ac-Y-NH 2 ] tot = (3.0?64) ! 10 ?4 M; [Ir IV ] 0 = 2.5 ! 10 ?5 M; [DBNBS] = 5 mM; ? = 0.1 M (LiClO 4 ); T = 25 ?C. b p[H + ] = ?log [HClO 4 ] in the p[H + ] range of 1.0?2.2. The following buffers (0.02 M) were employed to maintain constant p[H + ] values: monochloroacetate buffer for 2.4 < p[H + ] < 3.4, acetate buffer for 3.6 < p[H + ] < 5.4, and cacodylate buffer for 5.4 < p[H + ] < 7.4. 202 Table A-27. Nonlinear-Least-Squares Regresion Results of the k obs /[Ac-Y-NH 2 ] tot vs p[H + ] Plot with 2-Term and 3-Term Rate Law Models in the Presence of DBNBS. 2-term rate law (eq 2-2) 3-term rate law (eq 2-3) Best-fit values k ArOH , M ?1 s ?1 5.380 2.885 k ArO ?, M ?1 s ?1 4.47 ! 10 7 3.59 ! 10 7 k?, M ?1 s ?1 0.222 Std. Error k ArOH , M ?1 s ?1 0.590 0.235 k ArO ?, M ?1 s ?1 2.92 ! 10 6 9.89 ! 10 5 k?, M ?1 s ?1 1.90 ! 10 ?2 95 % Confidence Intervals k ArOH , M ?1 s ?1 4.129 to 6.632 2.384 to 3.386 k ArO ?, M ?1 s ?1 3.85 ! 10 7 to 5.09 ! 10 7 3.38 ! 10 7 to 3.80 ! 10 7 k?, M ?1 s ?1 0.182 to 0.263 Goodnes of Fit Degrees of Freedom 16 15 R squared 0.9909 0.9992 Weighted Sum of Squares (1/(Y*Y)) 0.732 0.068 Sy.x 0.214 6.72 ! 10 ?2 203 Table A-28. Kinetic Dependence on [Os III ] 0 . a [Os III ] 0 ! 10 5 , M t 1/2 , s k obs , M ?1 s ?1 0.49 3.0 7.14 ! 10 4 0.64 2.3 5.95 ! 10 4 0.79 1.8 6.84 ! 10 4 1.15 1.2 7.11 ! 10 4 a Al the reactions were run under pseudo-second-order conditions at 550 nm. [Os II ] 0 = (12.2?12.7) ! 10 ?5 M; [phenol] = 0.002 M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 5.1 (0.02 M acetate buffer). Table A-29. Kinetic data for the oxidation with various [phenol] tot . a [phenol] tot , M p[H + ] b k obs , M ?1 s ?1 0.002 5.44 8.12 ! 10 4 0.008 5.63 7.08 ! 10 5 0.014 5.67 1.89 ! 10 6 0.020 5.68 3.75 ! 10 6 a Al the reactions were run under pseudo-second-order conditions at 550 nm. [Os III ] 0 = 1.21 ! 10 ?5 M; [Os II ] 0 = 12.4 ! 10 ?5 M; ? = 0.1 M (NaCl); T = 25?C. b The p[H + ] values were maintained using 0.02 M acetate buffers. 204 Table A-30. Kinetic Dependence on [Os II ] 0 . a [Os II ] 0 ! 10 5 , M [Os III ] 0 ! 10 5 ,M [phenol] tot ! 10 3 , M t 1/2 , s k obs , M ?1 s ?1 0.68 1.38 2.5 0.11 6.27 ! 10 5 0.76 1.23 5.0 0.04 2.09 ! 10 6 1.04 0.92 5.0 0.08 1.46 ! 10 6 1.51 0.47 5.0 0.18 1.01 ! 10 6 a Al the reactions were run under ?pseudo-second-order? conditions at 480 nm. [Os III ] 0 = (0.47?1.38) ! 10 ?5 M; [Os II ] 0 = (0.68?1.51) ! 10 ?5 M; [phenol] tot = (2.5?5.0) ! 10 ?3 M; ? = 0.1 M (NaCl); T = 25?C; p[H + ] = 4.7 (0.02 M acetate buffer). Table A-31. Kinetic Dependence on p[H + ]. a p[H + ] b [phenol] tot ! 10 3 , M [Os II ] 0 ! 10 5 , M k obs , M ?1 s ?1 4.05 20 1.56 7.46 ! 10 5 4.71 5.0 1.51 1.01 ! 10 6 5.39 1.0 1.71 9.07 ! 10 5 6.34 0.2 1.81 2.19 ! 10 6 a Al the reactions were run under ?pseudo-second-order? conditions at 480 nm. [Os III ] 0 = (0.39?0.67) ! 10 ?5 M; ? = 0.1 M (NaCl); T = 25 ?C. b The following buffers (0.02 M) were employed to maintain constant p[H + ] values: acetate buffer for 4.0 < p[H + ] < 5.4, and cacodylate buffer for p[H + ] = 6.34