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Hydrogen Peroxide Disproportionation and Organic Compound Oxidation by Peroxycarbonate Catalyzed by Manganese(II): Kinet...


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HYDROGEN PEROXIDE DISPROPORTI ONATION AND ORGANIC COMPOUND OXIDATION BY PEROXYCARBONATE CATALYZED BY MANGANESE(II): KINETICS AND MECHANISM By ANDREW P. BURKE A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL OF THE UNIVERSITY OF FLOR IDA IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY UNIVERSITY OF FLORIDA 2005

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Copyright 2005 by ANDREW P. BURKE

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to my wife and parents

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iv ACKNOWLEDGMENTS The work presented here would not have been possible without the help and support of a number of people. I would like to acknowledge th ese people individually for their contributions in making the following document possible. First, I would like to thank my advisor, Dr. David Richardson, for all of his help and support during these past five years. Dr. Richardson has helped to make me a better scientist. My presentation and writing skill s have vastly improved under his advisement, and they will prove useful in all my future endeavors. I have also had the opportunity to learn from Dr. Richardson the proper me thod for performing chemical kinetics. I would also like to thank the member s of the Richardson group, both past and present, for all of their help during the year s. Without the fun environment they created, working in the lab would have been much less enjoyable. I would like to thank my partner in crime, Dan Denevan. It was always nice having Dan to make jokes with and to have around to complain to about everything going wrong in the lab. I would also like to thank Dr. Ana Ison for all of her support duri ng this process. Ana was always around to discuss ideas about projects. I would also like thank her for the help she provided in acquiring the GC data. I would especially li ke to thank Dr. Celeste Regino. Celeste taught me the intricacies of HPLC and that in order for it to do what you want you must coddle it at all times. I would also like to thank Pat Butler for a ll of her hard work during my elementary education. As my SLD teacher Mrs. Butler worked with me constantly for many years

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v to help me cope with a disability I never t hought I would be able to overcome. Now, as I finish my dissertation to achieve my Ph.D., I appreciate even more all of the techniques she taught me to help me achieve my goals. I could not have accomplished this goal without the support of my family, especially my parents. My parents have al ways supported me in all of the decisions I have made and attending graduate school was no exception. Without their support, I would never have had the courage to face new challenges and persevere in the face of opposition. I would also like to thank my wife, Erin. I never expect ed to meet my wife in graduate school, nor did I expe ct her to be a chemistry gra duate student. She has been a constant support these past 5 years, and I would have given up this dream long ago were it not for her constant vigilance in driving me toward my goal. Most of all, I would like to thank God for all of His love and support through not only my graduate career, but my entire li fe. Through Him all things are possible.

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vi TABLE OF CONTENTS page ACKNOWLEDGMENTS.................................................................................................iv LIST OF TABLES.............................................................................................................ix LIST OF FIGURES.............................................................................................................x ABSTRACT....................................................................................................................... xx CHAPTER 1 INTRODUCTION........................................................................................................1 General Oxidation.........................................................................................................1 Reactive Oxygen Species.............................................................................................3 Hydrogen Peroxide.......................................................................................................4 Activation of Hydrogen Peroxide.................................................................................6 UV Activation.......................................................................................................6 Strong Base Activation..........................................................................................7 Strong Acid Activation..........................................................................................8 Acyl Hydroperoxides.............................................................................................8 Iron(II) Activation.................................................................................................9 Transition-metal Organo metallic Complexes......................................................10 Methyltrioxorhenium...........................................................................................11 Asymmetric Oxidation................................................................................................12 Sharpless Oxidation of Allylic Alcohols.............................................................12 Mn(III)-salen Epoxidation Catalysts...................................................................13 Chiral Ketone Epoxidation Catalysts..................................................................15 Peroxycarbonate.........................................................................................................16 Transition-metal Peroxycarbonate Complexes....................................................19 Transition-metal Activation of Peroxycarbonate in Solution..............................22 Scope of the Dissertation............................................................................................23 2 OXIDATION OF NUCLEOPHILIC ALKENES IN AQUEOUS MICELLAR MEDIA.......................................................................................................................25 Introduction.................................................................................................................25 Results and Discussion...............................................................................................28 Styrene Oxidation in Micellar Media in the Absence of Mn(II).........................28

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vii Large Scale Styrene Oxidation............................................................................28 Styrene Oxidation in Micellar Media in the Presence of Mn(II).........................31 Reaction Kinetics.................................................................................................34 Dependence of Styrene Oxidation on Surfactant Identity...................................35 Dependence of Styrene Oxidation on the Manganese(II) Source.......................36 Bicarbonate Dependence.....................................................................................37 Experimental...............................................................................................................38 Materials and Instrumentation.............................................................................38 Standardization of Sodium Bicarbonate Solutions..............................................39 Styrene Oxidation Reactions...............................................................................40 Large Scale Styrene Oxidations..........................................................................40 Synthesis of Mn(DS)2..........................................................................................41 Styrene Oxidation in SDS with Mn(II) and Mn(DS)2.........................................41 3 KINETIC INVESTIGATIONS OF THE MANGANESE(II) CATALYZED DISPROPORTIONATION OF HYDROGEN PEROXIDE IN THE PRESENCE OF BICARBONATE AND THE CO MPARISON TO NUCLEOPHILIC ALKENE EPOXIDATION........................................................................................42 Introduction.................................................................................................................42 Results and Discussion...............................................................................................44 Kinetics of Hydrogen Peroxide Decomposition..................................................44 Manganese(II) Dependence..........................................................................47 Bicarbonate Dependence..............................................................................48 Comparison of Hydrogen Peroxide Reac tion Kinetics to Nucleophilic Alkene Epoxidation Kinetics........................................................................................50 Manganese dependence on nucle ophilic alkene epoxidation.......................50 Bicarbonate dependence on nucle ophilic alkene epoxidation......................51 Hydrogen peroxide dependence on nuc leophilic alkene epoxidation..........52 Catalyst Lifetime Studies....................................................................................53 Examining the loss of activity......................................................................54 Multiple additions of distilled hydrogen peroxide and solid sodium bicarbonate................................................................................................55 Studies of the Manganese Source........................................................................56 Potassium permanganate..............................................................................56 [MnIV(Me3TACN)(OMe)3]PF6.....................................................................57 Mn(IV) catalyst stability..............................................................................57 Cis-trans Isomerization in the Mangane se(II) Catalyzed Alkene Epoxidation...61 Cis/Trans isomerization reactions with cis-2-butene-1,4-diol.....................65 Cis/Trans isomerization of maleic and fumaric acids..................................65 Examination of Sychev’s Radical Trap Experiments..........................................68 Proposed Mechanism of N -dealkylation.............................................................81 Support for the Single Electron Transfer Pathway..............................................85 Solvent Isotope Effect.........................................................................................86 Proposed Mechanism...........................................................................................90 Numerical Simulation of the Proposed Mechanism............................................93 Conclusions...............................................................................................................109

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viii Experimental.............................................................................................................113 Materials and Instrumentation...........................................................................113 Standardization of sodi um bicarbonate solutions.......................................114 Hydrogen peroxide de composition studies................................................115 Synthesis of [MnIV(Me3TACN)(OMe)3](PF6)...........................................115 Oxidation of N N -dimethyl-4-nitrosoanil ine (DMNA) by Oxone.............116 Oxidation of N N -Dimethyl-4-nitrosoaniline (DMNA) by H2O2/HCO3 -/Mn2+.......................................................................................................116 Oxidation of N N -diethyl-4-nitrosoaniline (DENA) by H2O2/HCO3 -/Mn2+117 4 ELECTROPHILIC ALKENE EPOXIDATION BY THE PEROXYCARBONATE DIANION........................................................................118 Introduction...............................................................................................................118 Results and Discussion.............................................................................................119 Effect of Mn(II) on Electr ophilic Alkene Epoxidation.....................................121 Effect of pH on the Oxidati on of Electrophilic Alkenes...................................125 Effect of Buffer Choice on Elect rophilic Alkene Epoxidation.........................130 Electrophilic Alkene Oxidation Kinetics...........................................................131 Discussion of the SecondOrder Rate Constants...............................................133 Conclusions...............................................................................................................139 Materials and Instrumentation..................................................................................142 Experimental.............................................................................................................143 Electrophilic Alkene Epoxidation.....................................................................143 Dibenzoylethylene Kinetics...............................................................................143 5 GENERAL CONCLUSIONS...................................................................................144 APPENDIX: VARIATIONS IN NUCLEO PHILIC ALKENE EPOXIDATION AND HYDROGEN PEROXIDE RATE CONSTANTS...................................................152 LIST OF REFERENCES.................................................................................................155 BIOGRAPHICAL SKETCH...........................................................................................163

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ix LIST OF TABLES Table page 1-1 Some common reactive oxygen species.....................................................................3 2-1 Comparison of Styrene Oxidation in CTACl and SDS for the Mn(II) catalyzed epoxidation. Reaction conditions: 0.05 M Styrene, 0.100 M CTACl or SDS, 0.25 M NH4HCO3, 1.00 M H2O2, and 10 M Mn(II). Errors are reported to the 95% confidence........................................................................................................36 2-2 Comparison of observed rate consta nts for differing manganese sources for micellar styrene oxidation. Reaction conditions: 0.05 M Styrene, 0.100 M SDS, 0.25 M NH4HCO3, 1.00 M H2O2, and 10 M Mn(II) or Mn(DS)2. Errors are reported to the 95% confidence................................................................................37 3-1 Comparison of observed rate cons tants for the decomposition of hydrogen peroxide, 0.100 M, in 0.20 M s odium bicarbonate with 3 and 4 M manganese(II) and permanganate. Errors reported are to the 95% confidence.......56 3-2 Comparison of observed rate cons tants for the decomposition of hydrogen peroxide (0.100 M, final concentration) in 0.20 M sodium bicarbonate with 3 and 4 M manganese(II) and [MnIV(Me3TACN)(OMe)3](PF6). Errors reported are to the 95% confidence........................................................................................61 3-3 Comparison of first-order rate cons tants for the epoxidation of sulfonated styrene in H2O and D2O. Reaction conditions: 0.001 M SS, 1.0 M Sodium Bicarbonate, 0.50 M Mn(II)...................................................................................86 3-4 Comparison of solvent isotope effect for hydrogen peroxide decomposition. Reaction Conditions: 0.40 M HCO3 -, 0.10 M H2O2.................................................90 4-1 The percent conversion of 1 in varying buffer at differing pH..............................131 4-2 pKa values in water and CTACl for several different dyes.151...............................139

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x LIST OF FIGURES Figure page 1-1 The sulfate dianion.....................................................................................................1 1-2 The oxidation of organic molecules is defined as formation of bonds to carbon with atoms that are more electronegative than carbon. Reduction is the loss of bonds to more electronegative atoms and bond formation with hydrogen................3 1-3 Superoxide dismutase enzymatically oxidizes the superoxide anion and two protons to hydrogen peroxide, anothe r reactive oxygen sp ecies. Hydrogen peroxide is the disproportionated by catalase to yield water and molecular oxygen........................................................................................................................4 1-4 The AO-process for the industrial production of hydrogen peroxide........................5 1-5 Illustration of a nucleophilic attack on hydrogen peroxide. The use of a general acid facilitates the proton transfer to yield the oxidized nucle ophile and water........6 1-6 The reactivity of olef ins with hydroxyl radicals.11.....................................................7 1-7 Polymerization of olefins by hydroxyl radical.11.......................................................7 1-8 Reactivity of electroph ilic olefins with nucleophilic oxidants, such as hydroperoxide, react to produce the epoxi de plus the oxidants’ corresponding leaving group, in this case hydroxide.........................................................................7 1-9 The reaction of an alkene with OH+ generates an interm ediate carbocation. A general base can then depr otonate the oxygen of the intermediate which results in ring closure to form the epoxide............................................................................8 1-10 Alkene oxidation by m -CPBA....................................................................................9 1-11 Activation of iron(III) tetrakis (pentafluorophenyl) porphyrin by hydrogen peroxide to produce a high oxi dation state iron complex.20.....................................11 1-12 The two dominant forms in the MTO/H2O2 system under acid conditions. The diperoxorhenium adduct reacts slightly slower than the monoperoxorhenium complex.23................................................................................................................12

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xi 1-13 Nucleophilic attack of an olefin on the electrophilic oxygen of the hydrogen peroxide activated methyltrioxorhenium yields the oxidized nucleophile and regenerates MTO. Attack of a nucle ophile on the diperoxo complex generates the oxidized nucleophile and the monoperoxorhenium complex.24.........................12 1-14 Illustration of the asymmetric epoxida tion using the Sharpless method. Use of the (+) or (-)-tartrate allows for the oxyge n atom to be added to only one face of the allylic alcohol.25..................................................................................................13 1-15 A salen ligand...........................................................................................................1 3 1-16 spiro[2H-1-benzopyran-2,1'-cyclohexane]...............................................................14 1-17 Asymmetric epoxidation of al kenes can be easily achieved using peroxymonosulfate to generate a dioxirane in situ .30...............................................15 1-18 Structure of 1,2:4,5-diO -isopropylidene-D-erythro2,3-hexodiuro-2,6-pyranose used by Shi30 for the asymmetric epoxidation of alkenes using peroxymonosulfate to generate a dioxirane in situ ...................................................16 1-19 The equilibrium formation of bicar bonate and peroxycarbonate proceeds through CO2 as an intermediate.34............................................................................17 1-20 Fe(qn)2(O2C(O)O]Ph4P1.5MeOH0.5 (CH3)2NCHO.38..........................................18 1-21 Nucleophilic attack on the peroxycar bonate anion. An intramolecular proton transfer in the transition state allows for release of bicarbonate instead of hydroxide as in the case of hydrogen peroxide........................................................19 1-22 Generation of a metal peroxycarbonate (LnM(CO4)Xm) from its parent O2 complex, LnM(O2)Xm, by passing CO2 through a dry solution of the parent complex.42................................................................................................................19 1-23 Structure of the (Ph3P)2Pt(CO4) complex of Nyman.45............................................20 1-24 Routes for the oxidation of PR3 by (PEt2Ph)3RhCl(CO4).43 Route A shows the solution chemistry where a solvent molecu le displaces a phosphine before it is oxidized. Route B shows the solid state chemistry where coordination of ethylene occurs first with the displace ment of a phosphine ligand followed by oxidation of the ligand..............................................................................................21 1-25 Structure of products of styrene oxidation by [(PEt2Ph)3RhCl(CO4)] under a CO2/O2 atmosphere that indicate a radical mechanism.43........................................22 1-26 The structure of diethylenetri aminepentaacetic acid (DTPA)..................................22

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xii 2-1 The structures of three common surfact ants. Cetyltrimethylammonium chloride is a cationic surfactant, wh ile sodium dodecylsulfate is anionic. Triton X-100 is a non-ionic surfactant...............................................................................................26 2-2 The structure of a micelle with a concentration greate r then the cmc......................27 2-3 The graphical representation of an alkene dissolved in a micelle............................27 2-4 The reaction scheme for the oxidation of styrene by hydrogen peroxide in the presence of bicarbonate a nd cetyltrimethylammonium chloride (CTACl) without the presence of Mn(II). Hydrolysis of the product epoxide forms the corresponding diol. Reaction conditi ons: 0.05 M Styrene, 0.10 M CTACl, 2.00 M H2O2, 1.00 M NH4HCO3, 3 days.........................................................................28 2-5 A picture of a lighter-thanwater liquid-liquid extractor..........................................30 2-6 Reaction scheme used by Burgess40 in the mixed solvent epoxidation of styrene..32 2-7 Schematic representation for the oxi dation of styrene in surfactant with hydrogen peroxide and bicarbonate catalyzed by manganese(II). Reaction conditions: 50 mM styrene, 0.10 M CTACl, 2.00 M H2O2, and 1.00 M NH4HCO3, 30 minutes.............................................................................................32 2-8 HPLC chromatograms for the initial r eaction (top panel) and after 30 minutes (bottom panel) for the oxid ation of styrene with H2O2, HCO3 -, and Mn(II) in the presence of surfactant (CTACl). HPLC performed using a C18 reverse phase column using a non-linear gradient for 12 minutes. Mobile Phase: 25%:75% (v:v) CH3CN:H2O – 95%:5% CH3CN:H2O.............................................................33 2-9 Styrene area disappearance versus time from the HPLC analysis of styrene oxidation by hydrogen peroxide in micellar media in th e presence of bicarbonate and Mn(II). Reaction conditions: 0.05 M Styren e, 0.100 M CTACl, 0.25 M NH4HCO3, 1.00 M H2O2, 10 M Mn(II).................................................................34 2-10 ln(styrene area) versus time to find the first-order rate consta nt. The line is the linear regression to the data at the 95% confidence. The kobs is the negative slope of the line........................................................................................................35 2-11 Structure of manganese(II) bisdodecylsulfate..........................................................36 2-12 Graph of kobs versus [NH4HCO3] for the styrene oxidation in the presence of 0.100 M CTACl. Reaction conditions: 0.05 M Styrene, 0.100 M CTACl, 1.00 M H2O2, and 10 M Mn(II).....................................................................................38 3-1 Hydrogen peroxide decomposition in the presence of manganese(II) and bicarbonate...............................................................................................................45

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xiii 3-2 Plot of the ln([H2O2]) versus time for varying bicarbonate concentration. The 0.20 and 0.30 M bicarbonate reactions are typical of the accelerations noticed for these reactions. Reaction conditions: 0.10 M H2O2, 4 M Mn(II)....................46 3-3 The dependence of kobs on the [Mn(II)]. R eaction conditions: 0.10 M H2O2, 0.4 M HCO3 -, varying [Mn(II)]. y = ((7.98 0.62) x10-4)x, error reported to the 95% confidence........................................................................................................47 3-4 Plot of kobs versus [NaHCO3]2. Reaction conditions: 0.10 M H2O2 and 4 M Mn(II). y = ((2.08 0.25) x10-3)x, error reported to the 95% confidence..............49 3-5 Plot of kobs versus [Mn(II)] observed for nucleophilic alkene epoxidation (Bennett, 2002)46 y = ((2.09 0.25) x10-3)x, error reported to the 95% confidence................................................................................................................50 3-6 Plot of kobs versus [HCO3 -]2 which shows a second-order dependence. Reaction conditions: 0.001 M p -vinyl benzene sulfonate, 1.00 M Mn2+ ( ) 0.10 M H2O2 y = ((2.62 0.17) x10-3)x ( ) 0.50 M H2O2 y = ((1.19 0.23) x10-3)x ( ) 0.75 M H2O2 y = ((8.33 0.76) x10-4)x, errors reported to the 95% confidence. (Bennett, 2002)46......................................................................................................51 3-7 Plot of kobs on the [H2O2]. Reaction conditions: 0.001 M p-vinyl benzene sulfonate ( ) 1.00 M NaHCO3, 0.50 M Mn2+ ( ) 0.75 M NaHCO3, 0.50 M Mn2+ ( ) 1.00 M NaHCO3, trace metal catalysis (Bennett, 2002).46........................52 3-8 Plot of kobs versus # of additions of hydrogen peroxide to a spent solution in the catalyst lifetime study over multiple days. There is a 16 hr delay before addition 16 and 24..................................................................................................................53 3-9 Plot of kobs versus # of hydrogen peroxide addi tions for the catalyst lifetime study using distilled hydroge n peroxide and adding solid sodium bicarbonate. The loss of activity is now due only to dilution and the inability to maintain the bicarbonate concentration at a constant value..........................................................55 3-10 Molecular structure of [MnIV(Me3TACN)(OMe)3](PF6).63......................................57 3-11 Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of 1.00 sodium bicarbonate..........................................................................................58 3-12 Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of 0.50 M hydrogen peroxide.......................................................................................58 3-13 Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of 25 L (0.100 M, final concentration) hydrogen peroxide was done at 350 seconds. The absorbance first decays to 0 and within a matter of minutes, the solution is bright yellow...........................................................................................59

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xiv 3-14 UV-vis specta of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6). The solid line is the spectrum in the presence of 1.00 M sodi um bicarbonate. The dotted line is the spectrum of the solution after 1 eq of hydrogen peroxide was added......................60 3-15 Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of 50 L (0.200 M, final concentration) hydrogen peroxide was done at 312 seconds. Even after 6 minutes, th e yellow color does not develop.........................60 3-16 The concerted mechanism for the m -CPBA oxidation of nucleophilic alkenes resulting in the retention of stereochemistry............................................................62 3-17 The stepwise oxidation of an alkene by Mn(salen) and hydrogen peroxide is shown. Cis/trans isomerization occurs in the transition state, where the C-C sigma bond is able to rotate into th e more stable trans configuration......................62 3-18 The cis/trans isomerization noted Burg ess in his epoxidation of stilbene using the Mn(II), H2O2, bicarbonate system using a mixed solvent system of DMF/H2O. (Burgess, 2002)40...................................................................................63 3-19 Synthetic scheme for synthesis of 4,4’-sulfonated stilbene. (van Es, 1964)64.........63 3-20 13C NMR of cis and trans-2,3epoxybutane-1,4-diol in D2O using methanol as an internal standard..................................................................................................64 3-21 Epoxidation of cis-2-butene1,4-diol (0.60 M) with 1.00 M HCO3 -, Mn(II) (10 M), and H2O2 (6.00 M) after 30 minutes (l eft) and 18 hrs (right).........................65 3-22 The structures of maleic and fuma ric acids at the operating pH of 8.4....................66 3-23 1H NMR of maleic and fumaric acid oxides in D2O using methanol as an internal standard.......................................................................................................66 3-24 Epoxidation of maleic acid by hydrog en peroxide and manganese(II) in the presence of bicarbonate after 15 min. Reaction conditions: 0.10 M maleic acid, 1.00 M H2O2, 0.80 M NaHCO3, and 10 M Mn(II)................................................67 3-25 Fenton’s reagent can be used to oxidize benzene to phenol and biphenyl...............69 3-26 The influence of inhibitors on th e catalase process in the Mn(II)/HCO3/H2O2 system. [Mn(II)] = 4 x 10-6 M, [H2O2] = 0.10 M, pH 7.0, [HCO3 -] = 0.4 M, and T = 25 C: 0) kinetic curve with no inhibitors; 1), 2), 3), and 6) in the presence of DMNA as the inhibitor(at concentrations of 1 x 10-5,1.5 x 10-5, 2 x 10-5,and 4 x 10-5 M respectively; 4)in the presen ce of tetranitromethane(4 x 10-5 M); 5) in the presence of hydroquinone (1.5 x 10-5 M); 7) decomposition of H2O2 without Mn(II) ion (blank experiment). (Sychev, 1977)47....................................................70 3-27 The reaction of N N -dimethyl-4-nitrosoaniline with peroxymonosulfate to yield N N -dimethyl-4-nitroaniline.....................................................................................72

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xv 3-28 1H NMR of the crude reaction mi xture after an oxidation of N N -dimethyl-4nitrosoaniline by hydrogen peroxide in th e presence of bicarbonate and Mn(II). Reaction conditions: N N -dimehtyl-4-nitrosoaniline (1 g, 6.66 mmol), 0.400 M sodium bicarbonate, 10 M Mn(II), 6.64 M H2O2, 1 hr..........................................74 3-29 GC trace for a standard of N N -dimethyl-4-nitrsoaniline. Non-linear gradient for 30 minutes, detection by FID...................................................................................75 3-30 GC trace for the crude reaction material from the oxidation of N N -dimethyl-4nitrsoaniline from Figure 3-25. Lack of a peak near 14.637 min proves that no starting material remains. GC conditi ons: non-linear gradient for 30 minutes, Detection: FID..........................................................................................................76 3-31 GC trace (left figure) and 1H NMR (right figure) for Fr action 4 of the silica column. Identification of the product as N N -dimethyl-4-nitroaniline was confirmed by comparison with a GC trace and 1H NMR of an authentic sample...77 3-32 GC trace (left figure) and 1H NMR (right figure) for Fr action 9 of the silica column. Identification of the product as 4-nitroaniline was confirmed by comparison with a GC trace and 1H NMR of an authentic sample..........................78 3-33 1H NMR of an authentic sample of N -methyl-4-nitroaniline. Comparison with the crude reaction mixture confirms its presence as a product................................79 3-34 The solution collected from the reaction of N N -diethyl-4-nitrosoaniline was analyzed by Gas Chromatography (lower figure), which was compared to an authentic sample of acetaldehyde. 1H NMR (top figure) of the solution also confirmed that the product was acetaldehyde..........................................................80 3-35 The proposed mechanism for the oxidative N -dealkylation of amines by hydrogen peroxide in the presence of bicarbonate as catalyzed by manganese(II). The secondary amine pr oduced can cycle again as long as it contains a hydrogen on the carbon to the nitrogen. A second molecule of aldehyde or ketone will also be produced................................................................83 3-36 The structure of N N -dimethyl-2-amino-2-methyl-3phenylpropane, the substrate used by Miwa et al.73,74 for use in experiments w ith cytochrome P450 on the oxidative N -dealkylation mechanism.......................................................................85 3-37 Mechanism of epoxide hydrolysis in acidic media..................................................87 3-38 Possible epoxidation routes through a manganese(IV) oxo complex. None of the envisioned reactions ha s a proton transfer...............................................................88 3-39 Mechanism of oxygen transfer by attack of the alkene on a manganese(II) bound peroxycarbonate. The proton transfer in the transition state may account for the inverse isotope effect observed................................................................................89

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xvi 3-40 Oxidation of a nucleophilic alkene by two sequential r eactions with the carbonate radical. The carbocation interm ediate formed explains the loss of retention observed for cis-alkenes............................................................................90 3-41 Proposed mechanism for hydrogen per oxide decomposition and nucleophilic alkene epoxidation in the presence of bicarbonate catal yzed by Mn(II)..................91 3-42 Proposed generation of the activ e manganese catalyst by from the [MnII(HCO3)2(HCO4)]complex by a 2 electron oxidation of manganese to form a high valent [Mn-O2-]2+ complex............................................................................92 3-43 Simulation of the dependence on the con centration of the act ive catalyst with varying bicarbonate. Simulation conditions: 1.00 M hydrogen peroxide and 4 M Mn(II). y = ((6.55 0.37)x10-7)x, error reported to the 95% confidence.........96 3-44 Plot of simulation results for [Mn(HCO3)]+ versus [HCO3 -]. The [Mn(HCO3)]+ quickly saturates due to the large equilibrium constant of 19.05. Simulation conditions: 1.00 M H2O2, 4 M Mn(II)...................................................................97 3-45 Simulation of the dependence on the con centration of the act ive catalyst with varying [HCO3]. Simulation Conditions: 1.00 M hydrogen peroxide and 4 M Mn(II). y = ((4.60 0.22)x10-7)x, error reported to the 95% confidence................98 3-46 The generated curve for the hydroge n peroxide dependence on nucleophilic alkene oxidation. The poin ts represent the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each H2O2 concentration. Reaction a nd simulation conditions: 0.5 M Mn(II), 1.00 M bicarbonate, 0.001 M Sulfonated Styrene (SS)........................................................99 3-47 The generated curve for the bicar bonate dependence on nucleophilic alkene oxidation. The points re present the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each [HCO3] concentration. Reaction and Simulation Conditions: 0.5 M Mn(II), 0.10 M hydrogen peroxide, 0.001 M Sulfonated Styrene (SS).........................................................................................100 3-48 The generated curve for the mangane se dependence on nucleophilic alkene oxidation. The points re present the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each [Mn(II)] concentration. Reaction and simulation conditions: 1.00 M bicarbonate, 0.55 M hydrogen peroxide, 0.001 M Sulfonated Styrene...................................................................100 3-49 A typical numerical simulation plot attempting to model the hydrogen peroxide decay curves. Points represent observed ln[H2O2] versus time, while the line is the simulated ln([H2O2]) versus time. Reaction and simulation conditions: 0.10 M H2O2, 0.30 M HCO3 -, 4.0 M Mn(II)................................................................102

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xvii 3-50 Simulation of hydrogen peroxide decay at lower bicarbonat e concentration. Points represent data, while the line is the simulated decay. Reaction and simulation conditions: 0.10 M H2O2, 0.10 M HCO3, 4 M Mn(II).......................106 3-51 A plot of [HCO3]2 versus “ kobs” for the hydrogen peroxide decomposition. Reaction and simulation conditions: 0.10 M H2O2, 4 M Mn(II)..........................107 3-51 A plot of “ kobs” versus [Mn(II)] for the hydrogen peroxide decomposition. Reaction and simulation conditions: 0.10 M H2O2, 0.40 M HCO3........................107 3-53 Simulated hydrogen peroxide de pendence for the hydrogen peroxide decomposition reactions. Simula tion conditions: 0.90 M HCO3, 0.5 M Mn(II).108 3-54 Simulated hydrogen peroxide de pendence for the hydrogen peroxide decomposition reactions. Simula tion conditions: 0.40 M HCO3, 3.0 M Mn(II).108 3-55 Plot of ln([H2O2]) versus time. The points represen t observed data and the line is the simulation. Reaction and simulation conditions:0.10 M H2O2, 0.40 M HCO3, 3.0 M Mn(II). As the plot indicate s the reaction accelerates as the decomposition occurs.............................................................................................109 4-1 The resonance structure of an electrophilic alkene, an -unsaturated ketone, explains the reactivity with nucleophilic oxidants. The -carbon of the alkene, as seen in the resonance structure, is mo re electropositive and will be the site of attack by a nucleophilic oxidant.............................................................................120 4-2 The mechanism of electrophilic alkene oxidation by the hydroperoxide anion is illustrated. The nucleophilic oxidant adds at the electrophilic carbon, the carbon. Reformation of the ketone moiety causes either the displacement of the hydroperoxide anion, regenerating the star ting alkene and hydr operoxide, or ring closure to form the epoxide and the hydroxide anion............................................120 4-3 The mechanism of electrophilic al kene epoxidation by the peroxycarbonate dianion is illustrated. The mechanism is identical to that of hydroperoxide oxidation, except that the nucleofuge of the peroxycarbonate dianion is the carbonate dianion...................................................................................................120 4-4 Electrophilic alkenes used in this study.................................................................121 4-5 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 at 60 min, in the presence of 1.00 M sodium bicarbonate (50% conversion)..............................................................121 4-6 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 with 4 M Mn(II) in the presence of 1.00 M sodium bicarbonate at 60 min (44% conversion)..................................122 4-7 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 with 5 mM DTPA in the presence of 1.00 M sodium bicarbonate at 60 min (50% conversion)..................................123

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xviii 4-8 1H-NMR of 1 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 8.6 after 15 min (50% conversion)...............................................................124 4-9 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 7.8 in 24 hrs (75% conversion).....................................................................125 4-10 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 8.6 at 24 hrs (50% conversion).....................................................................126 4-11 Speciation of methacrylic acid (M AA) at 0.05 M and hydrogen peroxide (0.30 M) as a function of pH. The maximum concentration of MAA and hydroperoxide happens to occur at pH 7.8, the buffering pH of sodium bicarbonate.............................................................................................................127 4-12 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 7.8 with 4 M Mn(II) in 15 min (>90% conversion)...................................128 4-13 1H-NMR of 3 epoxidized by H2O2 in 1.00 M sodium bicarbonate at pH 7.8 in 24 hrs (44% conversion).............................................................................................129 4-14 1H-NMR of 3 epoxidized by H2O2 in 1.00 M sodium bicarbonate at pH 8.6 in 24 hrs (66% conversion).............................................................................................130 4-15 Plot of log( kobs) vs pH for the oxidation of 4 by H2O2...........................................132 4-16 Plot of log( kobs) vs pH for the oxidation of 4 by -OCl............................................133 4-17 Reaction mechanism determined by Ro senblatt for the nucleophilic oxidation of electrophilic alkenes.131..........................................................................................133 4-18 Structure of the substrate used by Rosenblatt and Broome,131 o chlorobenzylidenemalononitrile.............................................................................134 4-19 A reaction coordinate diagram for the addition of a nucleophilic oxidant to an electrophilic alkene. If th e ring closure of the epoxide is the rate determining step, as seen in the diagram by a larger energy barrier for the formation of the epoxide, the identity of the nucleofuge (Z ) will determine the reactivity of the oxidant. The more stable the nucleofuge, the more reactive the oxidant will be..135 4-20 A reaction coordinate diagram for the addition of a nucleophilic oxidant to an electrophilic alkene. If the addition of the oxidant is the rate determining step, as seen in the diagram by a larger energy barrier to the formation of the intermediate, the identity of the nucle ofuge (Z) will no longer determine the reactivity of the oxidant. The more basic oxidant will react faster.......................136 4-21 Substrates used by Bunton and Minko ff to study the oxidatio n of electrophilic alkenes by the hydroperoxide anion.......................................................................137

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xix A-1 Variation in the S + A Products rate constant...................................................152 A-2 Variation in the equilib rium constant for A + H2O2 B...................................153 A-3 Variation in the equilibrium c onstant for the formation of “A”.............................153 A-4 Variation in the rate constant for A radicals.....................................................154

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xx Abstract of Dissertation Pres ented to the Graduate School of the University of Florida in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy HYDROGEN PEROXIDE DISPROPORTI ONATION AND ORGANIC COMPOUND OXIDATION BY PEROXYCARBONATE CATALYZED BY MANGANESE(II): KINETICS AND MECHANISM By Andrew P. Burke August 2005 Chair: David E. Richardson Major Department: Chemistry The investigation of the mechanism of hydrogen peroxide di sproportionation and alkene epoxidation in aqueous solutions of bi carbonate at near neutral pH as catalyzed by manganese(II) is described. Current literat ure proposes that a free hydroxyl radical pathway based on Fenton chemistry is responsible for the hydrogen peroxide decay. This proposed mechanism does not adequately expl ain the unique requirement of bicarbonate in these reactions. Also, the proposed free hydroxyl radical mechanism does not explain why no radically coupled products in the oxida tion of nucleophilic alkenes are detected. We suggest that manganese(II) is activated by peroxycarbonate, a hydr ogen peroxide and bicarbonate adduct, to form a high oxidation st ate manganese(IV) complex. In addition, it is proposed that the carbonate radical an ion is also a product of the reaction of peroxycarbonate in the presence of metal cations. Both th e carbonate radi cal anion and the high oxidation state manganese(IV) comple x are believed to be the main reactive

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xxi oxygen donors in this system responsible for the observed reactivity. Numerical simulations of the hydrogen peroxide and nucleophilic alkene epoxidation by hydrogen peroxide in solutions of bicarbonate a nd manganese(II) have also been conducted. While the epoxidation of nucleophilic alke nes by hydrogen peroxide in bicarbonate solutions is catalyzed by manganese(II), the same is not true for the epoxidation of electrophilic alkenes. Inves tigations have been conducted using several water soluble alkenes. For these reactions, the addition of manganese(II) has been shown to inhibit the oxidation by decomposing the active oxida nt, the hydroperoxide anion. Kinetic investigations of the oxidation of dibenz oylethylene in micellar media by hydroperoxide and hypochlorite will also be presented. The obs erved second-order rate constant for the oxidation by hydroperoxide is 660 40 M-1s-1 and that for hypochlorite is 118 2 M-1s-1.

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1 CHAPTER 1 INTRODUCTION General Oxidation Oxidation-reduction (redox) reactions are some of the most important chemical reactions. These reactions are responsible for the formation of compounds from their elements, the generation of electricity, and combustion reactions, some of which produce energy at the cellular level. Redox reactions are always coupled, and the number of electrons transferred must be equal in numb er between the oxidation and reduction halfreactions. Redox reactions can be easily determined by identify ing the oxidation states of the atoms in the ions and molecules involved in the reaction. Lewis st ructures provide an easy convention by which oxidation states may be assigned to atoms. Typically, all bonds must be assumed to be completely ioni c, and the more electronegative atom of the bonded pair is allocated the pair of electrons. For example, consider the sulfate dianion, SO4 2-. S O OO O 2Figure 1-1. The sulfate dianion The sulfur-oxygen bonds are polar covalent polarized toward the oxygen atoms. Each oxygen atom is given an oxidation number of -2, eight valence electrons versus six

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2 for the free atom. The sulfur atom is give n an oxidation number of +6, zero valence electrons versus six for the free atom. The char ge of the ion is given by the sum of all the formal oxidation charges, in this case 6-8 = -2. The use of oxidation states can now be used to easily identify redox reactions. As an example, consider the reacti on of sulfite and permanganate anions in acidic solution to yield the sulfate anion and manganese(II). 5SO3 2+ 2MnO4 + 6H+ 5SO4 2+ 2Mn2+ + 3H2O (1-1) In this example, the sulfur at om of sulfite begins in the +4 oxidation state and is in the +6 oxidation state on the product side, as is seen in the Equation 1-1, a loss of two electrons. By the definition of oxidation a nd reduction, this is an oxidation. By definition, oxidation reactions must be coupl ed to a reduction reaction, the permanganate must be gaining electrons. In the example above this is seen to be true since the reactant manganese of permanganate is in the +7 oxida tion state and a gain of 5 electrons yields manganese(II), as seen on the product side of the reaction. In organic chemistry, the assignment of oxi dation states is not as simple as the above example with the sulfate dianion.1,2 It has been the trad itional method, therefore, to define oxidation in organic chemistry as the “loss” of electrons by forming bonds with elements that are more electronegative th an carbon, such as oxygen or nitrogen. Reduction then is the “gain” of electrons by breaking bonds with more electronegative atoms and forming bonds with hydrogen.1,2 For example, ethanol can be transformed to form acetaldehyde. In this process, ethanol looses a bond to hydrogen and gains a bond to oxygen, an oxidation. Similarl y, acetic acid can be converted to acetaldehyde. In this process, acetic acid looses a bond to o xygen and gains a bond to hydrogen, a reduction.

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3 O H OH O OH O H Oxidized Reduced Figure 1-2. The oxidation of organic molecules is define d as formation of bonds to carbon with atoms that are more electr onegative than carbon. Reduction is the loss of bonds to more electronegative atoms and bond formation with hydrogen. Reactive Oxygen Species Molecular oxygen and reactive oxygen species (ROS) are the main oxygen sources for oxidation processes and are highly reac tive oxygen donor molecule s with the ability to react with a wide variety of substrates.3 Table 1-1 lists some of the most commonly encountered radical and non-ra dical reactive oxygen species. Table 1-1. Some common reactive oxygen species Radical Non-radical Superoxide, O2 - Hydrogen Peroxide, H2O2 Hydroxyl, OH Hypochlorous acid, HOCl Peroxyl, ROO Alkyl Hydroperoxide, ROOH Alkoxyl, RO Hydroperoxyl, HOO In the human body, for example, the effects of reactive oxygen species have been measured by examining the oxidative stress on cells.4 Oxidative stress is defined as the imbalance between the cellular production of reactive oxygen species and the antioxidant mechanisms in existence to remove them.4 The effects of these reactive oxygen species have been linked to chronic disease and aging.5,6 The human body has several mechanisms, including enzymes and radical scavengers, that can interven e with reactive oxygen species.4 For example, superoxide

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4 dismutase converts the superoxi de radical anion plus two pr otons to hydrogen peroxide. The product hydrogen peroxide, which is yet another reactive oxygen species, can then be enzymatically decomposed by catalase to water and oxygen. The human body also uses a number of radical scavengers, such as as corbate, urate, and to copherol, to rid cells of high concentrations of reactive oxygen species. O2Superoxide Dismutase2 H+H2O22 H2O2Catalase 2 H2O + O2 Figure 1-3. Superoxide dismutase enzymati cally oxidizes the superoxide anion and two protons to hydrogen peroxide, anothe r reactive oxygen sp ecies. Hydrogen peroxide is the disproportionated by catalase to yield water and molecular oxygen. Hydrogen Peroxide Hydrogen peroxide (H2O2) is a common reactive oxygen species which is environmentally friendly due to its decompositi on to molecular oxygen and water. It is a weak, nonspecific, electroph ilic oxidant with an E = 1.77 V vs. NHE7 that has been used as a bleaching agent for over a century.8 Hydrogen peroxide is only a weak oxidant under mild conditions. Recent interest in the use of H2O2 as a terminal oxidant has come from increasing pressure in th e industrial sector to find mo re environmentally friendly oxidation reagents.8 Many industries are beginning to use hydrogen peroxide in the treatment of wastewater. Recently, hydrogen peroxide was shown to remove cyanide from thermoelectric power station wastewater.9 Hydrogen peroxide is also an important co mmercial chemical in the production of epoxides. In this case, hydrogen peroxide is used to generate peracids that are then used in the epoxidation of numerous alkenes.10 The activation of hydr ogen peroxide to form peracids will be presented shortly in the introduction.

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5 Hydrogen peroxide is produced co mmercially by the AO-Process,10 which involves the hydrogenation of a 2-alkyl-9,10-anthra quinone to the corresponding hydroquinone. The hydroquinone produced is then oxidized w ith oxygen, or air, to regenerate the anthraquinone and produce hydrogen peroxide (Figure 1-4). Th e hydrogen peroxide produced is extracted with water, while th e organic components can be recycled back through the hydrogenation step. O O R OH OH R Catalyst H2O2H2O2 Figure 1-4. The AO-process for the indus trial production of hydrogen peroxide. Hydrogen peroxide is an excel lent environmental choice fo r two reasons. First, the decomposition products are molecular oxygen and water. Second, due to its relative inactivity, specific methods of activation must be used which can tune the reactivity to the particular oxidative process required. Figur e 1-5 illustrates the nucleophilic attack of a substrate on hydrogen peroxide. A general acid can act as a proton transfer agent to assist in the cleaving of th e peroxide bond to form the oxi dized nucleophile and water.

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6 Nu: H O O H Nu O H O H NuOH+ + OH-NuO + H2O +HA Nu O H O H A H -HA Figure 1-5. Illustration of a nucleophilic attack on hydrogen peroxide. The use of a general acid facilitates th e proton transfer to yiel d the oxidized nucleophile and water. Activation of Hydrogen Peroxide UV Activation While hydrogen peroxide may be used for some types of oxidations, activation is required for use in a wider variety of reac tions. For instance, solutions of hydrogen peroxide can be irradiated using UV radiation to homolytic ally cleave the peroxide bond to form two hydroxyl radi cals, Equation 1-2. H2O2 2 HO (1-2) h The hydroxyl radical is a poten t, nonspecific, one-electr on oxidant that can readily react with alkenes (Figure 16) by addition to the double bond.11 The resulting organic radical can then react with anot her hydroxyl radical to form the diol, or in the presence of iron(II) and acid, the alcohol. Depending on the nature of the double bond, radical polymerization can also occur, as seen in Figure 1-7.

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7 R' R'' H H HO + R' R'' H H HO HO R' R'' H H HO HO Fe2+H+R' R'' H H HO H Figure 1-6. The reactivity of olefins with hydroxyl radicals.11 R HO R HO R HO R HO R R Polymer Figure 1-7. Polymerization of olefins by hydroxyl radical.11 Strong Base Activation Other methods for the activation of hydrogen peroxide are known. The reaction of hydrogen peroxide with a str ong base generates the hydroper oxide anion, as seen in Equation 1-3, which is an effective nucle ophilic oxidant. The hydr operoxide anion can epoxidize an electrophilic alkene as seen in Figure 1-8. H2O2 + -OH -OOH+H2O (1-3) R'' R R' OHOO-R'' R R' O OOH O R '' R' O R+ OHFigure 1-8. Reactivity of elect rophilic olefins with nucleophilic oxidants, such as hydroperoxide, react to produce the epoxi de plus the oxidants’ corresponding leaving group, in this case hydroxide.

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8 Strong Acid Activation Hydrogen peroxide is also activated by st rong acids. Protonation of one of the oxygens in hydrogen peroxide results in pol arization of the O-O bond to generate OH+, a strong electrophilic oxidant that can react with nucleophiles, su ch as alkenes. Water is the other product of the react ion (Equation 1-4). H2O2 HO+ + H2O (1-4) H+ R OH+R O H AR O + HA Figure 1-9. The reaction of an alkene with OH+ generates an intermediate carbocation. A general base can then deprotonate the oxygen of the intermediate which results in ring closur e to form the epoxide. Acyl Hydroperoxides Acyl hydroperoxides, a broad category of oxidants includ ing the organic peracids, are electrophilic oxidants often used for the he terolytic oxidation of organic substrates. Peracids are synthesized from the acid-catal yzed equilibrium between hydrogen peroxide and the acid form of the peracid, as seen in Equation 1-5.12 In the absence of a catalyst, the equilibrium is slow. In order to isolate the peracid fr om the equilibrium mixture, continuous distillation or an ex traction step must be used. RCO2H + H2O2 RCO3H + H2O (1-5) H+ A common example of a peracid us ed in organic oxidations is mchloroperoxybenzoic acid (m-CPBA). The mechanism of the reaction of an organic nucleophile, an alkene, proceeds as shown in Figure 1-10. This example uses m-CPBA as the oxidant. The rate of peracid epoxidation of alkenes is influenced by three main factors. First, the reac tion is dependent on the type of double bond. Second, the

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9 substituents of the peracid affect its ability to oxidize an alkene. Third, the rate of reaction is reduced in the presence of coordi nating solvents, such as ethers, which form intermolecular H-bonds.13 The kinetic aspects of peracid oxidations are as follows: 1) the reaction is second order, 2) th e reaction is stereospecific, m eaning that cis-alkenes will react to give cis-epoxides and trans-alkenes will react to give trans-epoxides, and 3) the rate of reaction is increased with in creasing strength of the formed acid.13 R R' O O Cl HO R R' O O H RR' O + HO O Cl O R'' Figure 1-10. Alkene oxidation by m-CPBA Iron(II) Activation One of the best known and most studied processes for the ac tivation of hydrogen peroxide is by iron(II) salts, the combination of which is known as Fenton’s reagent. The most accepted mechanism for the activation, introduced by Haber and Weiss14,15 and studied extensively by Barb et al.,16-19 involves the redox cycl e of iron(II). The mechanism is shown in Equations (1-6)-(1-10). Fe2+ + H2O2 Fe3+ + •OH + OH(1-6) Fe2+ + •OH Fe3+ + OH(1-7) •OH + H2O2 HOO• + H2O (1-8) Fe2+ + HOO• Fe3+ + HOO(1-9) Fe3+ + HOO• Fe2+ + O2 + H+ (1-10)

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10 Equation 1-6 represents th e initiation reaction by which hydrogen peroxide is activated to form a free hydroxyl radical. Further chain pr opagation reactions involving the hydroxyl radical ultimately produce peroxyl radicals, as shown in Equation 1-8 with a hydrogen abstraction. The hydroxyl and per oxyl radicals are reactive oxygen species known to be powerful oxidizing agents and have been implicated in aging and chronic disease. In the chain termination step, a pe roxyl radical reacts w ith iron(III) to yield a proton, molecular oxygen, and regeneration of iron(II) to complete the redox cycle. Transition-metal Organometallic Complexes Another method for hydrogen peroxide activ ation is through the use of transitionmetal cation complexes with organo ligands, including porphyrins. There are a wide variety of metal complexes described in the literature with any numb er of differing metal cations including Cu(I), Cu(II), Ni(II), Co(II ), Co(III), Fe(II), Fe(III), Mn(II), and Mn(III).8 The mechanisms by which these metal porphyrin complexes activate hydrogen peroxide vary and include production of fr ee hydroxyl radicals as well as formation of high valent metal-oxo species.8 In the case of free hydroxyl radical formation, mechanisms based on Fenton type chemistr y are proposed, as seen in the [Cu(phen)2]+ example in Equation 1-11. In the case of hi gh valent metal-oxo formation reactions, such [Cu(phen)2]+ + H2O2 [Cu(phen)2]2+ + HO + HO(1-11) as seen by Traylor et al20 in the activation of hyd rogen peroxide by iron(III) tetrakis(pentafluorophenyl) porphyrin, the mech anism of oxidation is believed to occur via an oxygen transfer from a high valent iron complex.

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11 N N N N F F F FF F FF F F F F F F F F F F F F Fe H2O2"Fe O" Figure 1-11. Activation of iron(III) tetrakis(pentafluor ophenyl) porphyrin by hydrogen peroxide to produce a high oxi dation state iron complex.20 Methyltrioxorhenium Methyltrioxorhenium (MTO), first intr oduced by Beattie and Jones in 1979,21 in combination with hydrogen peroxide provide s a useful system for the oxidation of alkenes and other substrates such as al kynes and ketones, as reported by Herrmann.22 Considerable research on the kinetics of the reaction of MT O with hydrogen peroxide has been studied by Espenson, who has shown th at two predominant species exist in the MTO/H2O2 system, shown in Figure 1-12, that ar e more stable under acidic conditions.23 Both are efficient oxygen donors to substrates such as phosphines a nd sulfides. In the reaction of alkenes with MTO/H2O2, both of the peroxide adducts react to form epoxide, but the monoperoxide tends to react s lightly faster than the diperoxide.24 There are three major drawbacks to the use of MTO as an al kene epoxidation catalys t: 1) MTO has a low stability in the presence of pe roxide 2) MTO is both expens ive and difficult to synthesize and 3) because reactions are conducted unde r acidic conditions, ring-opening of acid sensitive epoxides is possible.23

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12 ReO O O +H2O2, -H2O -H2O2, +H2O Re O O O O Re O O O O O MTO diperoxorhenium(VII) monoperoxorhenium(VII) +H2O2, -H2O -H2O2, +H2O Figure 1-12. The two dominant forms in the MTO/H2O2 system under acid conditions. The diperoxorhenium adduct react s slightly slower than the monoperoxorhenium complex.23 Re OO O ReO O O O Re O O O O O +H2O2Nu:Nu: +H2O2-H2O2-H2O2NuONuO Figure 1-13. Nucleophilic attack of an ol efin on the electrophili c oxygen of the hydrogen peroxide activated methyltrioxorhenium yields the oxidized nucleophile and regenerates MTO. Attack of a nucle ophile on the diperoxo complex generates the oxidized nucleophile and the monoperoxorhenium complex.24 Asymmetric Oxidation In addition to activating hydrogen peroxide fo r use as an oxidizing agent, it is also advantageous to control the addition of th e oxygen atom to substrates, for instance, alkenes, to generate only one epoxide enantio mer. A number of different methods have been proposed in the literature and have used a number of differen t techniques to assure that the oxygen atom only adds to one face of the alkene. Three of these methods are discussed in detail below. Sharpless Oxidation of Allylic Alcohols A simple and relatively inexpensive met hod for asymmetric epoxidation of allylic alcohols using titanium(IV) tetraisopropoxide, tetrabutyl hydroperoxi de (TBHP), and (+) or (-)-diethyl tartrate was introduced by Sharpless in 1980.25 The use of the (+) or (-)diethyl tartrate facilitates th e addition of the oxygen to one face of the alkene as shown in Figure 1-14. One of the most interesting asp ects of this epoxidation system is the ability

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13 to add the oxygen to a particular face, depe nding on which tartrate enantiomer is used, regardless of the substituti on pattern of the alkene. R' R''' R'' HO “O” “O” Figure 1-14. Illustration of the asymmetric epoxidation using the Sharpless method. Use of the (+) or (-)-tartrate allows for the oxygen atom to be added to only one face of the allylic alcohol.25 Mn(III)-salen Epoxidation Catalysts Chiral manganese(III)-salen complexes (the salen ligand is illustrated in Figure 115) are another example of an asymmetric alkene epoxidation cat alysts. Although the Sharpless method used tartrate as an additive for achie ving asymmetric epoxidation, manganese(III)-salen catalysts rely on the ch irality of the complex to provide for the asymmetric addition of the oxygen atom. The popularity of the epoxidation system comes from the ease in synthesis of the catal yst and the use of cheap, readily available oxidants, such as iodosyl benzene and hypochlorite. N N H H R R R R OH HO Figure 1-15. A salen ligand.

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14 Hydrogen peroxide can also be used as the terminal oxidant, although decomposition of the oxidant by the cat alyst has been observed. Katsuki26 in 1994 reported that hydrogen peroxide could be used with Mn(I II)-salen cata lysts for the epoxidation of chromene. The yields were low (17-53 %), but with good ee (93-96 %). It was noted by Katsuki that in order for the reaction of Mn(III)-salen catalysts to epoxidize alkenes with hydrogen peroxide, an axial ligand was required. For the epoxidations of chromene, N-methylimidazole was used as the axial ligand. It has also been noted that the use of carboxylate salts as additives are useful in the epoxidation of alkenes by Mn(III)-salen catalysts.27 In 1998, Pietikainen found that the use of manganese(III)-salen with 30 % hydrogen pe roxide, along with ammonium acetate, oxidized spiro[2H-1-benzopyran-2,1'-cyclohex ane] in 90 % yield and an ee of 91 %.28 O Figure 1-16. spiro[2H-1-ben zopyran-2,1'-cyclohexane] Enantiomeric excess (ee) provides a me thod for reporting the yield of one enantiomer in comparison to the other. In the case above for the oxidation by Pietikainen, a 91% ee was repor ted. This means that 9 % (100% 91%) of the product is racemic, implying that the remaining mixture is 4.5 % of each enantiomer. The total yield of the predominant enantiomer is th en 95.5 % (91 % + 4.5 %), while the other enantiomer is 4.5 %.

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15 Chiral Ketone Epoxidation Catalysts The use of chiral dioxiranes for the e poxidation of alkenes was first reported by Curci et al. in 1984.29 Shi et al.30,31 has observed that chiral ketones are effective catalysts for the asymmetric epoxidation of alkenes by the in situ generation of dioxiranes using potassium peroxymonosulfate (Oxone), as seen in Figure 1-17. R 1R2O HSO5 -HSO4 -R1O O R2R1R2R1R2O Figure 1-17. Asymmetric e poxidation of alkenes can be easily achieved using peroxymonosulfate to generate a dioxirane in situ.30 These reactions are performed in mi xed solvent systems, usually 1,2dimethoxymethane and water. The alkene is soluble in th e organic solvent, while the Oxone is soluble in the aqueous layer. The ketone catalyst used is soluble in both water and the organic solvent, thus allowing the keto ne to act as a phase transfer catalyst. The ketone is oxidized by peroxymonosulfate in the aqueous layer to form the dioxirane, which then transfers to the organic layer wher e it can oxidize the alkene and regenerate the starting ketone. The asymme tric epoxidation is facilitate d by the chiral nature of the ketone. As the generated dioxirane nears th e alkene, the oxygen is transferred to only one face of the alkene. The ee’s that have been reported using 1,2:4,5-di-Oisopropylidene-D-erythro-2,3-hexodiuro-2,6pyranose, whose structure is shown in Figure 1-18, range from 12 – 98%. The lowest ee’s are for the epoxi dation of cis-alkenes

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16 (12 – 56.2%), while an ee of 76.4 – 98% has been observed for the trans-alkenes. The difference in the ee between cis and trans-alkene s has been attributed to the approach the catalyst can make to the alkene in th e transition states for the two alkenes.30 O O O O O O Figure 1-18. Structure of 1,2:4,5-di-O-isopropylidene-D-erythr o-2,3-hexodiuro-2,6pyranose used by Shi30 for the asymmetric epoxidation of alkenes using peroxymonosulfate to generate a dioxirane in situ. Peroxycarbonate Recent work in our group has found that bi carbonate is an effective activator of H2O2,32,33 known as BAP, b icarbonate a ctivated p eroxide. Equilibrium between bicarbonate and H2O2 produces the peroxycarbonate anion (HCO4 -), as seen in Equation 1-12. HCO3 + H2O2 HCO4 -+H2O (1-12) The mechanism by which this equilibrium occurs has been determined by Yao34 and has been found to proceed through carbon di oxide as an intermediate. The presence of carbonic anhydrase or the carboni c anhydrase model complex 1,4,7,10tetraazacyclododecanezinc(II) accel erates the equilibrium reaction through catalysis of the dehydration of bicarbonate and possibl e catalysis of the perhydration pathway.34 The complete equilibrium processes is shown in Figure 1-19.

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17 CO2HO2 + H+ H2O2HOO OH O -OO OH O Ka (H2CO4)-O2H H2O -OH + H+HO O OH -O O OH-OH SSO Ka (H2CO3) Figure 1-19. The equilibrium formation of bicarbonate and peroxycarbonate proceeds through CO2 as an intermediate.34 Peroxycarbonate is a strong oxidant with an E (HCO4 -/HCO3 -) of 1.8 0.1 V vs. NHE.35 Inorganic salts and metal complexes of peroxycarbonate have been isolated and analyzed by X-ray crystallogra phy and vibrationa l spectroscopy.36,37 The analysis indicates that peroxycarbonate is a tr ue peroxide with a structure of HOOCO2 -. Recently, an iron(III) complex has been isolated and characterized by X-ray crystallography, as seen in Figure 1-20.38 Synthesis of metal comple xes of peroxycarbonate will be presented in the next section of this in troduction. Peroxycarbo nate should not be confused with sodium percarbonate, which is simply the cocrystallite of sodium carbonate and H2O2 (Na2CO3•1.5 H2O2). Peroxycarbonate is a moderately active oxidant for organic substrates, including sulfides and alkenes.39-42 The increase in reactivity over hydrogen peroxide can be

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18 attributed to the nature of the leaving group during a nucleophilic attack of the electrophilic oxygen of peroxide and peroxycarbonate. In the case of hydrogen peroxide, a general acid is required as a proton transfer agent, as seen in Figure 1-5. In the case of peroxycarbonate, however, an intramolecular proton transfer can re lease bicarbonate as the leaving group, as seen in Figure 1-21. Because bicarbona te is a weaker base than hydroxide, peroxycarbonate is a stronger elect rophile over hydrogen peroxide by a factor of about 300 based on studies with sulfides.41 Figure 1-20. Fe(qn)2(O2C(O)O]Ph4P1.5MeOH0.5 (CH3)2NCHO.38 Typical reactions of substrates with per oxycarbonate are slow, but still much faster than background reactions with hydrogen peroxi de alone. For instance, epoxidations of water soluble alkenes in the absence of bicarbonate yields negligible products in 24 hours. With the addition of bicarbonate, however, NMR analysis shows 90% conversion to the corresponding epoxide in 15 hours. A similar trend is also seen in sulfide oxidation.43,44

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19 O O O OH Nu: O O O O Nu H NuO + HCO3 Figure 1-21. Nucleophilic attack on the pe roxycarbonate anion. An intramolecular proton transfer in the tran sition state allows for rel ease of bicarbonate instead of hydroxide as in the ca se of hydrogen peroxide. Transition-metal Peroxycarbonate Complexes Transition-metal complexes containing the peroxycarbonate dianion ligand, CO4 2-, are known. The general formula for these peracids are LnM(CO4)Xm, where L = an ancillary ligands, n = 2 or 3, M = Pd, Pt, Rh or Ir, X = a halogen, and m = 0 or 1.42 The peroxycarbonate complexes are generally synthesized by passing carbon dioxide gas through a solution of the LnM(O2)Xm parent complex dissolved in a dry solvent. Two possible mechanisms for the formation of the peroxycarbonate complexes are shown in Figure 1-22. Oxygen label studies43 indicate that the carbon dioxide does not insert into the M-O bond (pathway 2), but instead inserts into the O-O bond (pathway 1). These complexes are classified as heterolytic oxidants and are good electrophilic oxida nts which react with nucleophiles such as alkenes and phosphines.43 M O O * 1 2 M O O * or M O O * CO2M O O O O * M OO * CO2M O O O O * Figure 1-22. Generation of a metal peroxycarbonate (LnM(CO4)Xm) from its parent O2 complex, LnM(O2)Xm, by passing CO2 through a dry solution of the parent complex.42

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20 For example, Nyman et al.44,45 observed that when carbon dioxide was passed through a dry benzene solution of (Ph3P)2PtO2, the platinum peroxycarbonate complex, whose structure is shown in Figure 1-23, wa s obtained. The complex was identified based on its infrared spectrum, chemical properties, and elemental analysis. Pt Ph3P Ph3P OO O O Figure 1-23. Structure of the (Ph3P)2Pt(CO4) complex of Nyman.45 Estimation of the peroxide content or oxidation power of the complex was attempted, but it was stated that no completely satisfactory method was found.44 In general, the (Ph3P)2Pt(CO4) complex in the presence of acidified iodide solutions did produce an immediate color change attributed to the release of iodine, but the color faded. The loss of color was thought to occur via th e oxidation of the tr iphenylphosphine, but no data were presented indicating that th e oxidation products were identified. Unfortunately, attempts to oxidize or ganic species were not undertaken. In 2001, Aresta et al.43 examined the reactivity of (PEt2Ph)3RhCl(CO4) which was synthesized by passing CO2 through the parent O2 complex. She describes both the solution and solid state oxidation of one of the phosphine ligands. In solution, a solvent molecule displaces a phosphine ligand from the coordination sphere of the metal. The phosphine is proposed to act as a nucleophile and attack the electrophilic oxygen of the peroxycarbonate ligand. This reaction th en yields the corresponding carbonato Rh complex and the oxidized phosphine (Figure 1-24, Route A). In the solid state reaction, the presence of ethylene and the Rh complex does not yield ethylene oxide. The ethylene displaces a phosphine ligand from the coordination

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21 sphere (Figure 1-24, Route B). The displaced phosphine then attacks the peroxycarbonate ligand and is oxidized. Fr om these experiments it was concluded that the mechanism of oxidation does not occur by an intramolecular oxygen transfer from the peroxycarbonate ligand directly to the phosphine. For this particular complex, the phosphine must be displaced firs t before it can be oxidized. Rh O O O O Cl R3P R3P PR3A B Solution Solid State -PR3 -PR3 + Rh O O O Cl R3P Solv R3P + R3P=O Rh O O O Cl PR3R3P + R3P=O Figure 1-24. Routes for the oxidation of PR3 by (PEt2Ph)3RhCl(CO4).43 Route A shows the solution chemistry where a solvent mo lecule displaces a phosphine before it is oxidized. Route B shows the solid state chemistry where coordination of ethylene occurs first with the displace ment of a phosphine ligand followed by oxidation of the ligand. The (PEt2Ph)3RhCl(CO4) complex has also been observed to oxidize more reactive olefins, such as styrene.43 When styrene (0.1 mL, 0.873 mmol) in 2 mL THF was allowed to react with [(PEt2Ph)3RhCl(CO4)] (0.100g, 0.16 mmol) under a CO2/O2 atmosphere (10:1 v:v), benzaldehyde, phenyl acetaldehyde, phenyl methyl ketone, and styrene oxide were observed by GC/MS in a ratio of 1:3:3:5, respectively. The presence of benzaldehyde, phenylacetaldehyde, and phe nyl methyl ketone suggest that the mechanism of oxidation occurs via radical chemistry as opposed to a simple oxygen transfer in which case styrene oxide would be the only product.

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22 O H Benzaldehyde O H Phenylacetaldehyde O Phenyl methyl ketone Figure 1-25. Structure of products of styrene oxidation by [(PEt2Ph)3RhCl(CO4)] under a CO2/O2 atmosphere that indicate a radical mechanism.43 Transition-metal Activation of Peroxycarbonate in Solution Recent work by Burgess et al.39,40 has shown that the addition of certain transitionmetal cations increases the catalytic rate of alkene epoxidation in solutions of hydrogen peroxide and bicarbonate in mixed solvent syst ems. Of the inorganic metal salts tested, manganese(II) sulfate produced th e greatest increase in the epoxidation reaction. Along with an increase in the rate of epoxidation, the addition of Mn(II ) to solutions of H2O2 and bicarbonate also enhances the rate of H2O2 disproportionation. The rate of disproportionation is enhanced to such a degr ee that methods must be employed to deal with the excessive amount of heat ev olved. Recent studies by Bennett46 have shown that the addition of the chelating agent diethylen etriaminepentaacetic acid (DTPA) inhibits the oxidation of alkenes, but onl y in some cases. This has been attributed to the removal of extraneous metal cations from the bicarbonate salts.46 N N N O OH O HO O OOH OH O OH Figure 1-26. The structure of diethylen etriaminepentaacetic acid (DTPA)

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23 In 1977, Sychev et al.47 reported that hydrogen peroxide is rapidly disproportionated in the presence of bicarbonate and free Mn(II). In a series of papers from 1977 to 1984,47-56 an investigation on the reaction mechanism of H2O2 decomposition with manganese(II) and bicar bonate was conducted. His proposed mechanism assumes that Mn(II) follows the Fenton type chemistry of Fe(II) with its reaction with H2O2 and therefore, proceeds via a free hydroxyl radical pathway. His work, however, does not provide adequate deta il into the necessity of the bicarbonate ion in this reaction. Addition of similar anions, such as acetate, phosphate, oxalate, or borate, does not result in H2O2 disproportionation when Mn(II ) is introduced. Also, the explanation provided by Sychev does not enlighten us in the observation that alkenes are cleanly oxidized to epoxide wit hout detection of usual radica l coupled products, as seen in the work by Aresta.43 Scope of the Dissertation The goal of this current st udy is to further understand the reactive nature of the peroxycarbonate anion and dianion. Chapte r 2 will discuss the reaction of Mn(II) and bicarbonate in the oxidation of styrene in mi cellar media. Both small and large scale oxidations of styrene were attempted and the results of these experiments will be presented. Questions arising from these e xperiments led us to i nvestigate the hydrogen peroxide disproportiona tion reaction further. The importance of peroxycarbonate in the Mn(II) catalyzed disproportionation of hydrogen peroxide will be the focus of Chapter 3. Kinetic investiga tions of the reaction have been conducted, and the re sults of these experiments will be presented. The lifetime of the catalyst was also investigated. The similarities between hydrogen peroxide disproportionation and nucleophilic alkene oxidation using Mn (II) was also of interest

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24 during this study. A proposed mechanism for the hydrogen peroxide decomposition and alkene epoxidation will be introduced and numerical simulations of various proposed models for the disproporti onation and alkene epoxidat ion will be presented. Chapter 4 will discuss the use of the pe roxycarbonate dianion as a nucleophilic oxidant for epoxidation of elect rophilic alkenes. The resu lts of experiments with the peroxycarbonate dianion will be compared with kinetic measurements using other nucleophilic oxidants. A summary and discussion of possible future work will comprise Chapter 5.

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25 CHAPTER 2 OXIDATION OF NUCLEOPHILIC ALKENES IN AQUEOUS MICELLAR MEDIA Introduction Prior work by Bennett46 and Yao33 has shown that alkenes can be oxidized to the corresponding epoxides using bicarbonate-activated pero xide. Studies of hydrophobic alkenes were conducted in mi xed solvent systems to all ow for solubility of the alkene, while wa ter-soluble alkenes were chosen for pure water studies. In general, epoxidations were found to be slow using peroxycarbonate solutions, but faster than background oxidations by hydrogen peroxide al one. It was also found that reactions proceed fa ster in pure water than in mixed solvent systems. The faster reaction in water has b een attributed to a proton tran sfer, which proceeds faster in pure water than in mixed solvent systems.32 It was also noted by Burgess et al.40 that the addition of transition metal salts in the pr esence of hydrogen peroxi de and bicarbonate in H2O/DMF solutions accelerated th e oxidation of alkenes. Of the transition-metals tested, manganese(II) sulfate produced the most dramatic increase in oxidation. In order to take full advantage of the use of transition-metal sa lts for alkene oxidation, surfactants were used to allow for the oxidation of hydrophobic alkenes in aqueous solution in the presence and absence of manganese(II) salts. The use of surfactants is also advantageous since the work by Bennett46 indicates that the oxidation of alkenes tends to proceed faster in pure water. Surfactants are long chain alkanes with hydrophobic tails and polar head groups, which allow for solubility in water. Three examples of common surfactants,

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26 cetyltrimethylammonium chlori de (CTACl), sodium dodecyls ulfate (SDS), and Triton X100, are found in Figure 2-1. N 15 Cl-cetyltrimethylammonium chloride O 11 S O O O Na+sodium dodecylsulfate (H3C)3CCH2(H3C)2COCH2CH2OH x x = 10 T r itonX-100 Figure 2-1. The structures of three co mmon surfactants. Cetyltrimethylammonium chloride is a cationic surfactant, while sodium dodecylsulfate is anionic. Triton X-100 is a nonionic surfactant. When dissolved in water, surfactants will begin to organize themselves into micelles after the concentration reaches a cr ucial level known as the critical micelle concentration (cmc),57 which is a unique value for each surfactant and depends on the ionic strength of the solution. The det ection of micellization can be accomplished by observation of the surface te nsion, refractive index, or conductivity (for ionic surfactants).57 At the cmc, the hydrophobic tails will begi n to congregate, expelling water from the forming micelle’s core, while the polar head groups will arrange to allow for the maximum interaction with water, as seen in Fi gure 2-2. The Stern la yer is defined as the area around the micelle where the polar head groups are located, as are their counter ions.57

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27 Normally, hydrophobic molecules are unable to dissolve is aqueous solution. However, if micelles are present, a hydrophobic molecule, such as an alkene, is able to penetrate into the hydrophobic core of the mice lle and dissolve, as seen in Figure 2-3. Hydrophobic Core Pola r Head Group Counterion Stern Layer + + + + + + + + + + + +-+ Figure 2-2. The structure of a micelle with a concentration greater then the cmc. + + + + + + + + + + + + R'R'' R' R'' R' R'' R'R'' R' R'' Figure 2-3. The graphical representation of an alkene dissolved in a micelle.

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28 Results and Discussion Styrene Oxidation in Micellar Media in the Absence of Mn(II) Initially, styrene was oxidized in micella r media using the BAP method without the introduction of manganese(II) salts (Figure 2-4). Styren e (50 mM), CTACl (100 mM), H2O2 (2.00 M), and ammonium bicarbonate (1. 00 M) in a volume of 250 mL were allowed to react in water for 3 da ys in the dark with stirring. Analysis of the products by HPLC showed that approximately 90% of the starting styrene had reacted to yield the corres ponding epoxide (~90 %), although significant hydrolysis to the corresponding diol has also been detected (~10%). H2O2, HCO3 -CTACl O H2O OH OH Figure 2-4. The reaction scheme for the oxida tion of styrene by hydroge n peroxide in the presence of bicarbonate and cetyltr imethylammonium chloride (CTACl) without the presence of Mn (II). Hydrolysis of th e product epoxide forms the corresponding diol. Reaction conditi ons: 0.05 M Styrene, 0.10 M CTACl, 2.00 M H2O2, 1.00 M NH4HCO3, 3 days Large Scale Styrene Oxidation In addition to small scale epoxidations of styrene in micellar media, large scale epoxidations were attempted. For a typical large scale epoxidation, 5 mL of styrene (175 mM), 230 mL CTACl (350 mM), 2.00 M H2O2, 39.6 g NH4HCO3 (1.00 M), and enough water (to bring the volume to 500 mL) were allo wed to react in a to tal volume of 500 mL with stirring for 3 days in the dark at room temperature. For the small scale epoxidations, purification of the unreacted styrene and the styrene oxide product was unnecessary due to the use of HPLC for the analysis of the reaction pr oducts. For the large scale epoxidations, however, a purific ation method was required.

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29 First, extraction with methylene chloride was attempted for the isolation of the styrene and styrene oxide. Since the surf actant would probably pr oduce an emulsion, it was thought that if the reacti on solution were diluted the em ulsion would dissipate within a short amount of time. Unfortunately, this was not the case. The emulsion formed, even when the reaction was diluted to 2.5 L. On occasion, allowing the emulsion to stand overnight would allow for small amounts of methylene chloride to be isolat ed from the extraction. Upon drying and removal of solvent, styrene oxide, the surfactant, and water were observed by 1H NMR analysis. While the surfactant, CTACl in this case, is a cationic species and should not dissolve in organic solvents, the surfactant c ould form a reverse micelle, where the polar head groups now surround a small amount of water and the hydrophobi c tails extend into the organic solvent.57 This would explain why both su rfactant and water are observed by 1H NMR. Given the unsatisfactory results from extr action, another method for purification of the organic products was required. The sec ond method used for the purification of large scale styrene oxidations was li quid-liquid extraction. This purification method uses the same principal as extraction, but without the tendency to form emulsions. At first, for the large scale epoxidations, ether was used fo r the liquid-liquid extr actions since only a lighter-than-water liquid-liqui d extraction apparatus was immediately available. The setup of the apparatus is shown in Figure 2-5. For the process of liquid-liquid extraction, a constant volume of organic solvent can be used to extract the organic product fr om the aqueous layer. This provides a convenient method for extracting slightly soluble organic products with a minimum

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30 amount of organic solvent, as opposed to nor mal extraction procedures which typically require larger volumes of organic solvent. Condenser Distilling Organic Solvent Reflux Arm Return Arm Organic Layer Aqueous Layer Funnel Condenser Distilling Organic Solvent Reflux Arm Return Arm Organic Layer Aqueous Layer Funnel Figure 2-5. A picture of a lighter-than-water liquid-liquid extractor. Initially, the organic solvent is layered over the aqueous soluti on. In addition, a small amount of organic solvent is placed in a round bottom flask connected to the reflux arm. Once the organic solvent in the round bottom has begun refluxing, it will be

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31 liquefied in the condenser and collects in th e funnel. As the orga nic solvent collects in the funnel, small amounts of the solvent will be pushed out the end of the funnel into the aqueous layer. As the organic solvent rises through the aqueous layer, a small amount of organic product will diffuse into the droplet. After a few minutes, enough organic solvent will have added to the solvent layered over the aqueous solution to allow the organic solvent to drip down the return arm b ack to the refluxing solvent. In this way, the organic product will slowly accumulate in the round bottom flask. After the extraction is complete, the solvent in the round bottom can be dried and the solvent removed to give the desired product, as woul d be done for a normal extraction process. The time of completion for the extraction must be determined experimentally for the unique conditions in which the extractor is be ing used. Heavier th an water liquid-liquid extractions also exist and extract the organic product in a similar way. For the large scale epoxidations, it was found that the highest yield of epoxide was observed after 3 days of liqui d-liquid extracting. In addi tion to the styrene oxide, the corresponding diol was also present (~15% ), when analyzed by HPLC. This is reasonable since after 3 days of reacting a 10% conversion to the diol is observed. The additional 3 days of extraction accounts for the additional 5% conversion of the epoxide to the diol product. On occasion, surfactant and water were still observed in the purified product, even when care was taken to assu re that no emulsions were formed when layering the organic solvent over th e aqueous layer in the extractor. Styrene Oxidation in Micellar Med ia in the Presence of Mn(II) Recent work by Burgess et al.39,40 has shown that the in troduction of transitionmetal salts to oxidations of hydrophobic alkenes by H2O2 and bicarbonate in mixed solvent systems of dimethyl formamide (D MF) and water show a significant rate

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32 enhancement, Figure 2-6.39,40 Upon addition of a transitionmetal salt, epoxidations with reaction times greater than 48 hours were decreased to only 16 hours. The main drawback, specifically the long reaction times, to Burgess’ method has been the slow addition of the H2O2 and bicarbonate solutions to DMF to minimize precipitation.40 This rate enhancement is greatest when the i norganic salt added is manganese(II) sulfate. H2O2, HCO3 -, Mn2+DMF/H2O O Figure 2-6. Reaction scheme used by Burgess40 in the mixed solvent epoxidation of styrene. Our current method of alkene epoxidati on using micellar media offers a better alternative to the mixed solvent system em ployed by Burgess, Figure 2-7. The main drawback seen by Burgess, namely the slow addition of the H2O2/bicarbonate solution, is not an issue in micellar media. The organic substrate is dissolved by the micelle, and all of the remaining reactants are freely soluble in water, so precipitation is no longer of concern. When a test reaction was perfor med using styrene (50 mM), CTACl (100 mM), H2O2 (2.00 M), NH4HCO3 (1.00 M), and only 10 M MnSO4, the epoxidation was complete in less than 30 minutes as seen by the HPLC chromatograms in Figure 2-8, as opposed to 3 days in the absence of manganese(II). H2O2, HCO3 -, Mn2+Surfactant, H2O O Figure 2-7. Schematic representation for th e oxidation of styrene in surfactant with hydrogen peroxide and bicarbonate catalyzed by manganese(II). Reaction conditions: 50 mM styrene, 0.10 M CTACl, 2.00 M H2O2, and 1.00 M NH4HCO3, 30 minutes.

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33 Figure 2-8. HPLC chromatograms for the initia l reaction (top panel) and after 30 minutes (bottom panel) for the oxid ation of styrene with H2O2, HCO3 -, and Mn(II) in the presence of surfactant (CTACl). HPLC performed using a C18 reverse phase column using a non-linear grad ient for 12 minutes. Mobile Phase: 25%:75% (v:v) CH3CN:H2O – 95%:5% CH3CN:H2O

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34 Reaction Kinetics Kinetic experiments were conducted to determine the dependence of various conditions, including the identity of the surfact ant, the source of the manganese, and the bicarbonate concentration, on the manganese( II) catalyzed oxidations of styrene in micellar media. For these reactions, aliquo ts of reaction solutions were removed over time and added to a solution of bovine catalase to destroy any remaining H2O2 and therefore, quench the reaction. The aliquot s were then diluted with acetonitrile and analyzed by HPLC. Figure 2-9 shows a representative graph demonstrating the disappearance of styrene versus time. From a pl ot of the ln(styrene area) versus time, the first-order rate constant can be determined from the slope of line, as shown in Figure 210. 0 500000 1000000 1500000 2000000 2500000 3000000 3500000 4000000 07001400210028003500 Time, secondsStyrene Area Figure 2-9. Styrene area disa ppearance versus time from th e HPLC analysis of styrene oxidation by hydrogen per oxide in micellar media in the presence of bicarbonate and Mn(II). Reaction co nditions: 0.05 M Styrene, 0.100 M CTACl, 0.25 M NH4HCO3, 1.00 M H2O2, 10 M Mn(II).

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35 y = -0.0014x + 15.385 10 11 12 13 14 15 16 0500100015002000250030003500 Time, secondsln (Styrene Area)kobs= (1.43 0.10) x10-3, s-1 y = -0.0014x + 15.385 10 11 12 13 14 15 16 0500100015002000250030003500 Time, secondsln (Styrene Area)kobs= (1.43 0.10) x10-3, s-1 Figure 2-10. ln(styrene area) vers us time to find the first-order rate constant. The line is the linear regression to the data at the 95% confidence. The kobs is the negative slope of the line. Dependence of Styrene Oxidation on Surfactant Identity For all reactions previously desc ribed, the surfactant used was cetyltrimethylammonium chloride a cationic surfactant. In order to determine whether the active species is charge d, the oxidation of styrene was performed under the same conditions as described except for the substi tution of sodium dodecylsulfate (SDS) for CTACl. If the active catalyti c species is positively charge d, the reaction should proceed faster in the anionic micelle due to attract ive forces between the micelle and the active oxidant. If the active oxidant is negatively charged, the reaction s hould be slower in the anionic micelle due to the repulsive force be tween the micelle and the active oxidant. Conversely, if the active catalyst is uncharged, no difference in the rate of the reaction should be observed. From the data presented in Table 2-1, the obser ved rate constants for the oxidation of styrene in th e two surfactants give the same observed first-order rate constants (within error), therefore, the conc lusion must be that the active manganese oxidant is uncharged.

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36 Table 2-1. Comparison of Styrene Oxid ation in CTACl and SDS for the Mn(II) catalyzed epoxidation. Reaction cond itions: 0.05 M Styrene, 0.100 M CTACl or SDS, 0.25 M NH4HCO3, 1.00 M H2O2, and 10 M Mn(II). Errors are reported to the 95% confidence. Surfactant kobs, s-1 Cetyltrimethylammonium chloride (1.39 0.10) x10-3 Sodium dodecylsulfate (1.50 0.15) x10-3 Dependence of Styrene Oxidation on the Manganese(II) Source In addition to testing whether the surfact ant made an impact on the reaction, the addition of the metal was also examined. When bulk manganese(II) is added to the solution containing SDS, the manganese(II) ions must exchange with sodium ions at the surface of the micelle. Manganese(II) bisdodecylsulfate (Mn(DS)2) was synthesized to allow for the metal to be added already bound to the surfactant. Mn(DS)2 is precipitated by the addition of saturated sodium dodecylsulfate and satu rated manganese(II) chloride. Since the manganese(II) is bound to the surfactant, all of the manganese(II) should be bound to the micelle surface. S S O O O O O O OO Mn 11 11 Figure 2-11. Structure of manganese(II) bisdodecylsulfate. When the reaction was performed using 10 M Mn(DS)2, an identical observed first-order rate constant, within error, was observed in comparison with the reaction with the addition of bulk manganese(II). The results of this experiment indicate that with the addition of bulk manganese(II), the metal is in rapid equilibrium with the sodium ions at the micelle surface. So, in the case of Mn(DS)2, the metal is not remaining at the micellar surface, but is rapidly being released into the bulk soluti on by exchange with sodium

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37 ions. It is, therefore, unn ecessary to add the metal alrea dy bound to the su rfactant, since the bulk metal rapidly exchanges with sodium ions at the micellar surface. Table 2-2. Comparison of obs erved rate constants for di ffering manganese sources for micellar styrene oxidation. Reacti on conditions: 0.05 M Styrene, 0.100 M SDS, 0.25 M NH4HCO3, 1.00 M H2O2, and 10 M Mn(II) or Mn(DS)2. Errors are reported to the 95% confidence. Manganese Source kobs, s-1 Bulk Manganese (Mn2+) (1.50 0.15) x10-3 Mn(DS)2 (1.43 0.07) x10-3 Bicarbonate Dependence The bicarbonate dependence on the styrene oxidations in micellar media was also investigated. When a plot of kobs versus [HCO3 -] was made, as seen in Figure 2-12, the rate has saturated by 0.25 M bicarbonate. This finding was unexpected, since previous work on sulfide and alkene oxidations did not show saturation w ith bicarbonate at concentrations similar to those used here.32,41 It is apparent, therefore, that the addition of the surfactant is concentrati ng the active oxidant near the mi celle surface in order for the oxidation to be saturating. Since the active oxi dant had not been examined at this time, it was impossible to make any conclusive remarks about how the micelle was affecting the epoxidation of styrene by manganese(II) and bicarbonate with hydroge n peroxide in the presence of surfactant. In order to furthe r understand the use of the manganese(II) as a catalyst for the epoxidation of alkenes, the hydrogen peroxide disp roportionation reaction needed to be better understood. The hydr ogen peroxide disproportionation can be examined in the absence of surfactant, since all of the reactants are fully water soluble. If the active oxidant can be identified, further experiments can then be designed to probe the nature of the oxidant in micellar oxidati ons of alkenes. The results of experiments with the disproportionati on reaction are presented in the next chapter.

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38 0.0 0.5 1.0 1.5 2.0 2.5 00.20.40.60.81 [NH4HCO3], Mkobs 1000, s-1 Figure 2-12. Graph of kobs versus [NH4HCO3] for the styrene oxidati on in the presence of 0.100 M CTACl. Reaction conditions: 0.05 M Styrene, 0.100 M CTACl, 1.00 M H2O2, and 10 M Mn(II). Experimental Materials and Instrumentation Sodium bicarbonate, sodium acetate, styr ene, and manganese (II) sulfate were all analytical grades and obtained from Fisher (Atlanta, GA). Cetyltrimethylammonium chloride was obtained from Al drich (St. Louis, MO). H ydrogen peroxide (35 and 50%) was obtained from Fisher (Atlanta, GA) and st andardized often by i odometric titration. Water was purified using a Barnstead E-Pure 3-Module Deionization System. Extraneous metal ions from salt solutions were removed by passing the solutions through a Chelex 100 resin obtained from Aldrich (St. Louis, MO). Sodium bicarbonate solutions were standardized using the method below before use to assure concentration. UV-vis kinetic experiments were obt ained using a He wlett-Packer 8453 spectrophotometer using 1.0 cm quartz cells from Starna Cells, Inc. Temperature was maintained at 25 ( 0.1) C using a Fi sher Isotemp 1600S water bath circulator.

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39 Styrene oxidation reactio ns were analyzed by High Performance Liquid Chromatography (HPLC) using a Rainin HP LX solvent delivery system with a C-18 reverse phase column. The method consisted of a non-linear gradient of mobile phase A:H2O and mobile phase B:CH3CN from 75:25 A:B to 5:95 over a 12 minute period. Products were detected at 221 nm. Standardization of Sodium Bicarbonate Solutions Solutions of sodium bicarbonate were standardized before each kinetic experiment to determine the concentrati on eluting from the Chelex 100 column. All solutions were delivered using volumetric pi pets. A 10 mL aliquot of sodium bicarbonate solution exiting the Chelex 100 column was a placed in a clean, dry beaker. An excess amount of sodium hydroxide solution of a known con centration, by titration with potassium hydrogen phthalate, was added to the beaker The solution was stirred to allow for complete deprotonation of the bicarbonate to form the carbonate dianion. An excess barium chloride solution is then added to precipitate a ll of the carbonate dianion as barium carbonate. Phenolphthalein is then a dded to the mixture to give a pink color due to the residual hydroxide ion. The mixture is titrated us ing a known concentration of hydrochloric acid until th e solution just turns clear. The number of moles of hydrochloric acid added is equal to the excess moles of sodium hydroxide. The difference between the number of moles from the acid titration a nd the number of mo les of hydroxide ion initially added equals the number of moles of bicarbonate present in the initial 10 mL aliquot (Equation 2-1). The molarity of the solution can then be determined. #moles OHinit-#moles OHexcess = #moles bicarbonate (2-1)

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40 Styrene Oxidation Reactions Kinetic experiments of styr ene oxidations were performed in micellar solutions and analyzed by the decreasing area of the styren e peak by HPLC. Reactions were performed using 0.05 M styrene, varying ammonium bicarbonate, 1.00 M H2O2, and 0.10 M surfactant, where the surfactant was either ce tyltrimethylammonium chloride or sodium dodecylsulfate (CTACl and SDS, respectively). For reactions using manganese(II), as either manganese sulfate of Mn(DS)2, 10 M was added to the reac tions. Aliquots (100 L) of the reaction mixture were taken over ti me and quenched using a catalase solution, which converts any excess H2O2 to O2 and H2O, thus removing the terminal oxidant. Each of the aliquots is then diluted with CH3CN to an appropriate concentration for HPLC analysis. The kobs is calculated using pseudo-firs t order plots of the ln(styrene area) vs. time. First-order plots we re linear for the conditions studied. Large Scale Styrene Oxidations Large scale styrene epoxidations were c onducted in micellar solutions. Reactions were performed using 175 mM Styr ene (5 mL), 350 mM CTACl, 1 M NH4HCO3, and 2.0 M H2O2 in a total volume of 500 mL, where the re maining volume is water. The reaction was allowed to stir in the dark for 3 days. The reaction mixture was then diluted to 1 L and poured into a lighter than water liquid-liquid extractor. Ether was then layered on top of the aqueous reaction solution, and the extr actor was allowed to extract for 3 days. After cooling the receiving flas k after 3 days, magnesium sulfate was added to dry the solvent. After filtering off the magnesi um sulfate and removing the solvent under reduced pressure, the organic residue was analyzed by 1H NMR.

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41 Synthesis of Mn(DS)2 Mn(DS)2 was synthesized by mixing equal parts saturated manganese(II) chloride and saturated sodium dodecylsulfate in water to form a while solid. The solid was then filtered and washed with ice-cold water. Calculated for C24H50S2O8Mn4H2O Calc: C: 43.82% H:8.89% S:9.75% O:29.19%Mn :8.35% Found C: 43.65% H:9.10% Styrene Oxidation in SDS with Mn(II) and Mn(DS)2 Styrene oxidations were performed us ing 0.05 M styrene, 0.500 M ammonium bicarbonate, 1.00 M H2O2, and 0.10 M sodium dodecylsulfate. Manganese(II) was added as either 10 M bulk manganese(II) or 10 M Mn(DS)2. Aliquots of the reaction mixture were taken over time and quenched using a catalase solution, wh ich converts any excess H2O2 to O2 and H2O. Each aliquot is then diluted with CH3CN to an appropriate concentration for HPLC analysis. The kobs are calculated using ps eudo-first order plots of the ln(styrene area) vs. time. First order plots were lin ear for the conditions studied.

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42 CHAPTER 3 KINETIC INVESTIGATIONS OF THE MANGANESE(II) CATALYZED DISPROPORTIONATION OF HYDROGEN PEROXIDE IN THE PRESENCE OF BICARBONATE AND THE COMPARISON TO NUCLEOPHILIC ALKENE EPOXIDATION Introduction Investigations on the use of manganese(II) in the catalys is of alkene epoxidation presented in the preceding ch apter raise some interesti ng questions about the hydrogen peroxide disproportionation under the reaction conditions. Namely, what is the active manganese species? Why is bicarbonate necessary for the reaction to proceed? In 1977, Sychev et al.47 reported that hydrogen peroxi de disproportionates rapidly in the presence of bicarbonate and free manganese (II). In a series of papers from 1977 to 1984,47-56 Sychev studied the mechanism of pe roxide decomposition. His proposed mechanism assumes that manganese(II) follows the Fenton chemistry of iron(II) in its reaction with peroxide, and th erefore, proceeds via a free hy droxyl radical pathway. His work, however, does not provide adequate deta il into the unique catalytic ability of the bicarbonate ion in this reaction or the identity of the active manganese species. In 1990, Stadtman et al.58 also investigated the use of the manganese(II) catalysis as a method for oxidation of amino acids. Th e decomposition of hydrogen peroxide by manganese(II) with bicarbonate was also exam ined during their studies. As with the work of Sychev, Stadtman does not provide an explanation for the unique reactivity of bicarbonate or speculate about the active manganese species re sponsible for the catalysis.

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43 In this chapter, kinetic investig ations on the manganese(II) catalyzed decomposition of hydrogen peroxide in the pres ence of bicarbonate will be presented. The bicarbonate, manganese(II), and hydrogen peroxide dependencies measured during this study are similar to those reported by Sychev,47 but differ slightly from those observed by Stadtman.58 The observed differences in th e dependence are probably due to the conditions under which the reaction was exam ined. For this work and for that of Sychev, the hydrogen peroxide and bicarbonat e concentrations were in the 100-500 mM range, with a 0-10 M Mn(II) concentration. For the studies performed by Stadtman, the hydrogen peroxide and bicarbonate concentratio ns were in the 30 mM range, with much larger Mn(II) concentrations of 0.10 mM. The dependencies meas ured during this study are also similar to those observed by Bennett46 for the epoxidation of nucleophilic alkenes by hydrogen peroxide in the presence of bicarbonate catalyzed by manganese(II). Investigations on the lifetime of the catalys t have also been conducted. While the catalyst is still active upon addition of hydrogen peroxide to spent solutions, the reactivity is about half of the original activity. Thes e results are consistent with those reported by Stadtman. Two factors have been identified to explain the loss of activity. Experiments have also been conducted to examine whether the manganese source has an impact on the reaction. Three different manganese sources were used in this study and include manganese(II) sulfate, potassium permanganate, and a Mn(IV)-TACN complex. The results of these experiments i ndicate that the source of the manganese has no effect on the observed rate of hydrogen peroxide decomposition. Experiments were also conducted using ci s-alkenes to discern information about the mechanism of oxygen transfer from the active catalyst. Results from these

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44 experiments indicate that the oxygen atom is not being transferred in a concerted fashion, since cis/trans isomerization is observed in the epoxidation of nuc leophilic alkenes. Investigations of the radical traps used in the work by Sychev led to the interesting result that amines can react with this system to yield N-dealkylated produ cts. Reports on oxidative N-dealkylation are not common in the literature. Curre nt knowledge of oxidative N-dealkylation comes from studies with cytochrome P-450. The most accepted mechanism for the oxidative N-dealkylation of amines begi ns with a single electron transfer from the amine to a metal in a high oxidation state. A mechanism using manganese species that will be proposed in this study will be presented. Solvent isotope effects have been measur ed for the nucleophilic alkene epoxidation and hydrogen peroxide decomposition. A large, inverse isotope effect was observed for the alkene oxidation, while a normal is otope effect was observed for the hydrogen peroxide decomposition. The results of th e nucleophilic alkene oxi dation are consistent with those observed by Bennett. An attempt to justify the difference in the observed solvent isotope effects will be presented. Finally, a mechanism based on the work from this study will be presented in an attempt to explain both the nucleophilic alkene epoxi dation and hydrogen peroxide decomposition. Numerical simulation has been employed as a tool in the attempt to understand the reaction kinetics. Several models will be presented and discussed. Results and Discussion Kinetics of Hydrogen Peroxide Decomposition Hydrogen peroxide di sproportionation, Figur e 3-1, was followed spectrophotometrically by mon itoring the decreasing peak intensity at 263 nm. The hydrogen peroxide concentration was held constant at 0.10 M, while the sodium

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45 bicarbonate was varied in concentrati on from 0.0-0.60 M. The manganese(II) concentration was varied in the range from 0-30 M. 2 H2O2 2 H2O + O2Mn(II) NaHCO3, buffer Figure 3-1. Hydrogen peroxi de decomposition in the pr esence of manganese(II) and bicarbonate. In order to maintain a cons tant pH of 8.4 and an ionic strength of 1.00 M, two buffering systems were tried. Initially, s odium phosphate buffer was used, but it was discovered that the phosphate anion causes the manganese to precipitate. Also, the phosphate buffer did not maintain the pH. Even with higher con centrations, the pH would continue to rise during the experiment. Eventually, sodium acetate was employed to control ionic strength and stabilize pH. While sodium acetate is not actually a buffe r at pH 8.4, the addition of the salt in the solutions allowed for the stabilization of the pH. To exclude acetate as participating in the mechanism, a set of experiments were condu cted in the absence of acetate. While the pH did change slightly during these reactions the observed rate constants were nearly identical to those in the presence of acetate. Solutions of hydrogen peroxide, bicarbonate and sodium acetate were allowed to equilibrate for at least 15 mi nutes before kinetic experiment s were performed. Addition of a solution of the manganese(II) was always pr eformed last to initiate the reaction. To ensure that the order of reagent addition does not have an effect on the reaction, experiments were conducted which were ini tiated by the addition of hydrogen peroxide. These experiments gave nearly identical obs erved rate constants as those initiated by manganese(II), therefore, the order of reagent addition makes no significant difference to the overall reaction rates.

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46 The disproportionation reaction was monito red for greater than two half-lives. Linear first-order plots were only obtained under certain conditions from the absorbance decay by plotting the equation ((At-A)/(A0A)) versus time, as seen in Figure 3-2. In most cases, however, linear first-order plots we re not obtained. For these instances, the first-order plots were used as an attempt to approximate an observed first-order rate constant. In all cases, the reaction acceler ates near the end of the reaction. The hypothesis is that hydrogen peroxi de is an inhibitor of the re action. As the concentration of hydrogen peroxide drops, an acceleration in the peroxide decomposition will occur. Further details concerning this aspect of the reaction will be presented in the section discussing the numerical simulations. -2.5 -2 -1.5 -1 -0.5 0 02000400060008000 Time, secondsln (Abs) 0.1M NaHCO3 0.2M NaHCO3 0.3M NaHCO3 Figure 3-2. Plot of the ln([H2O2]) versus time for varying bicarbonate concentration. The 0.20 and 0.30 M bicarbonate reactions ar e typical of the accelerations noticed for these reactions. Reaction conditions: 0.10 M H2O2, 4 M Mn(II). The slope of the line produced by linear regr ession, of those plots that are linear, is the negative of kobs. For reactions that we re not linear, pseudo “kobs” were attempted to be used to analyze the data for these reactions. Background disproportionation of peroxide in bicarbonate and acetat e buffered solutions was negligible for the time scale of the catalytic di sproportionations.

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47 Manganese(II) Dependence The plot of kobs versus Mn(II) concentration gives a straight line with a y-intercept of 0 for the range of Mn(II) from 0–6.0 M, as seen in Figure 3-3, indicating a first-order dependence of Mn(II) on the react ion in that range. Since the hypothesis is that the catalytic disproportionati on of peroxide is metal dependent, a y-intercept of 0 is expected. The straight line plot indicates that only one manganese ion is found in the transition state complex. These data indicate that a multiple metal center complex is not involved in the disproportionation reaction under aqueous conditions. This is significant since much of the current literature on metal catalyzed peroxide disproportionation focuses on complexes with multiple metal centers.59-62 Scatter in the Mn(II) data begins to appear at about 6.0 M. The turnover in the data is attr ibuted to precipitation of manganese(II) carbonate. 0.00E+00 2.00E-03 4.00E-03 6.00E-03 8.00E-03 051015202530 [Mn2+], Mkobs, s-1 Figure 3-3. The dependence of kobs on the [Mn(II)]. Reac tion conditions: 0.10 M H2O2, 0.4 M HCO3 -, varying [Mn(II)]. y = ((7.98 0.62) x10-4)x, error reported to the 95% confidence. The solubility-product constant (Ksp) is defined as the equi librium constant between the ions in solution and the precipitated solid. An example using manganese(II) carbonate is given below.

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48 MnCO3(s) Mn2+(aq) + CO3 2-(aq) (3-1) Ksp = [Mn2+]*[CO3 2-] (3-2) The Ksp is reached at ~6.50 M Mn(II), given a constant bi carbonate concentration of 0.40 M, a pH of 8.4, and a Ksp of 2.72 x10-7, the conditions under which the Mn(II) dependence was measured. This corresponds well with the turnove r in the manganese dependence, Figure 3-3. This tu rnover was also observed by Sychev47 for his experiments on peroxide disp roportionation in bicarbonate so lutions with manganese(II), but further data as to why th e turnover occurred were not presented. The explanation given in the text simply referred to po ssible precipitation of manganese salts. Bicarbonate Dependence Studies of the dependence of sodium bicar bonate on the rate la w indicate a secondorder dependence, as seen in Figure 3-4. The plot of kobs versus [HCO3 -]2 produces a straight line with a y-intercept of 0. The data begins to scatte r at higher bicarbonate concentration for two reasons. First, the react ions become very fast and getting accurate data becomes more difficult. Second, th e higher bicarbonate reactions have a more pronounced curve in the ln plots (as discussed earlier), therefore the “kobs” reported is not a true first-order rate consta nt. As with Mn(II), since th e disproportionation of peroxide is dependent on the bicarbonate concentration, a y-intercept of 0 is expected. For the sodium bicarbonate, however, a second-orde r dependence reveals that two bicarbonate ions are present in the transition state comple x of the rate determining step. Presumably, one of these bicarbonate ions is in the form of peroxy carbonate, while the other may simply be coordinated to the metal center. Similar results were found for the nucleophilic alkene epoxidations studied by Bennett46 and the hydrogen peroxide decomposition

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49 studies of Sychev.47 Stadtman,58 on the other hand, reports a third-order dependence on the bicarbonate concentration. The difference in the order of the reaction could be due to the reaction conditions that we re used, since the hydrogen pe roxide and bicarbonate were much lower than this study and the mangane se was much higher. Currently, with the proposed mechanisms that will be presented later, a third-order dependence has not be observed in the numerical simulations. 0.00E+00 1.00E-03 2.00E-03 3.00E-03 4.00E-03 5.00E-03 6.00E-03 7.00E-03 8.00E-03 00.050.10.150.20.250.30.35 [NaHCO3]2, M2kobs, s-1 Figure 3-4. Plot of kobs versus [NaHCO3]2. Reaction conditions: 0.10 M H2O2 and 4 M Mn(II). y = ((2.08 0.25) x10-3)x, error reported to the 95% confidence. The overall rate equation de fined by the results above is given in Equation 3-3. At this time, no reliable method has been f ound to study the hydrogen peroxide dependence, so, its dependence for the hydrogen peroxide decomposition has yet to be observed. Numerical simulations which will be presented later indicate that there is an inverse dependence of the hydrogen peroxide concentration. It is for this reason that the ln plots curve near the end of the reactions. As th e hydrogen peroxide decays, the reactions begin to accelerate.

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50 v = kobs[Mn(II)][HCO3 -]2[H2O2]x (3-3) Comparison of Hydrogen Peroxide Reac tion Kinetics to Nucleophilic Alkene Epoxidation Kinetics In order to better understand the nature of the active catalyst in the manganese dependent hydrogen peroxide di sproportionation, the results of the kinetic investigation of hydrogen peroxide decomposition need to be compared to the results observed by Bennett46 for the manganese(II) dependent nucle ophilic alkene epoxidation observed in pure water using sulfonated styrene and 4-vinyl benzoic acid. A summ ary of her results for the study of sulfonated styrene are presented here. Manganese dependence on nucleophilic alkene epoxidation Bennett46 examined the manganese depende nce on the oxidation of sulfonated styrene in bicarbonate solu tion (1.00 M) with 1.00 M hydrogen peroxide. The dependence on manganese was shown to be linear in the range of 0 – 5.0 M, Figure 3-5. 0 0.002 0.004 0.006 0.008 0.01 012345 [Mn2+], Mkobs, s-1 Figure 3-5. Plot of kobs versus [Mn(II)] observed for nucleophilic alkene epoxidation (Bennett, 2002)46 y = ((2.09 0.25) x10-3)x, error reported to the 95% confidence.

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51 These data are consistent with the manganese(II) dependence observed for the hydrogen peroxide decomposition presented ear lier and support the proposal that only one manganese ion is presen t in the active catalyst. Bicarbonate dependence on nucleophilic alkene epoxidation Bennett46 also examined the bicarbonate de pendence on the oxidation of sulfonated styrene with manganese(II). The dependence of HCO3 on the oxidation was shown to be second-order (Figure 3-6), which is also seen in the manganese dependent hydrogen peroxide decomposition. Once again, these da ta indicate that two bicarbonate ions are present in the active ca talyst. Presumably, one of these bicarbonate ions is present as a peroxycarbonate ion while the other bicarbonate may simply be coordinated to the metal center. 0.0000 0.0005 0.0010 0.0015 0.0020 0.0025 0.0030 0.000.200.400.600.801.001.20 [HCO3 -]2, M2kobs, s-1 Figure 3-6. Plot of kobs versus [HCO3 -]2 which shows a second-order dependence. Reaction conditions: 0.001 M p-vinyl benzene sulfonate, 1.00 M Mn2+ ( ) 0.10 M H2O2 y = ((2.62 0.17) x10-3)x ( ) 0.50 M H2O2 y = ((1.19 0.23) x10-3)x ( ) 0.75 M H2O2 y = ((8.33 0.76) x10-4)x, errors reported to the 95% confidence. (Bennett, 2002)46

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52 Hydrogen peroxide dependence on nucleophilic alkene epoxidation Bennett46 also observed the hydrogen peroxide dependence for the oxidation of sulfonated styrene with manganese(II). The dependence of H2O2 on the oxidation was shown to have an inverse relationship with increasing peroxide concentration (Figure 37). 0.00 2.00 4.00 6.00 8.00 10.00 12.00 14.00 16.00 0.000.200.400.600.801.001.20 [H2O2], M104*kobs,s-1 Figure 3-7. Plot of kobs on the [H2O2]. Reaction conditions: 0.001 M p-vinyl benzene sulfonate ( ) 1.00 M NaHCO3, 0.50 M Mn2+ ( ) 0.75 M NaHCO3, 0.50 M Mn2+ ( ) 1.00 M NaHCO3, trace metal catalysis (Bennett, 2002).46 Two possibilities exist for the downward trend observed in the hydrogen peroxide dependence. One explanation involves the r eaction of hydrogen peroxide to form a less reactive intermediate of the manganese catalyst. This explanation is plausible given the work by Espenson on the oxidation of nucle ophilic alkenes by me thyltrioxorhenium (MTO).24 In the case of MTO, the addition of a second hydrogen peroxide molecule generates a diperoxo complex which has a slig htly lower epoxidation rate constant than does the monoperoxide intermediate. The other explanation for the downturn in th e reaction has come from work in this study using numerical simulation to model th e decomposition and epoxidation kinetics. From this work, it appears that hydrogen peroxide may actually inhibit its own

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53 decomposition at higher concentr ations. Further details abou t this possibility will be presented with the numerical simulations. Catalyst Lifetime Studies In addition to examining the kinetics of the manganese(II) catalyzed decomposition of hydrogen peroxide, the lifet ime of the catalyst was also of interest. Stadtman58 noted that the catalyst lost abou t half of its activity upon rein troduction of hydrogen peroxide into a spent decomposition solution. Stadtman,58 unfortunately, gave no indication as to why the solutions were losing their cata lytic ability. Figure 3-8 shows the kobs versus additions of hydrogen peroxide to a solu tion of manganese(II) and bicarbonate. 0 2 4 6 8 10 12 03691215182124 # of H2O2 Additions103* kobs, s-1 Day 1 Day 2 Day 3 Figure 3-8. Plot of kobs versus # of additions of hydrogen peroxide to a spent solution in the catalyst lifetime study over multiple days. There is a 16 hr delay before addition 16 and 24. As noted by Stadtman,58 the decomposition of peroxi de drops by about one-third upon reintroduction of peroxide, the 2nd data points. The addition of peroxide was studied over multiple additions, and for multiple days. As can be seen in the graph, hydrogen peroxide still decomposes even after the spent solution has been sitting for 16 hrs, data points 16 and 24. These data indica te that the catalyst is able to regenerate simply by the addition of more hydrogen peroxide.

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54 Examining the loss of activity The first suspected reason for the decrease in activity in the catalyst lifetime studies was the loss of bicarbonate from the solutions. The bicarbonate concentration was tested by running a large scale reaction (10 mL) with 1.00 M bicarbonate, 1.00 M hydrogen peroxide and 5 M manganese(II). The reaction was cy cled a total of 10 times, each time bringing the hydrogen peroxide co ncentration back to 1.00 M. After the last reaction had decomposed the hydrogen peroxide, the bicar bonate concentration was analyzed by the standard barium chloride pr ecipitation method. It was found that after 10 cycles, the bicarbonate concentration had dropped from 1.00 M to about 0.50 M. This result indicates that one of the reasons the decomposition of hydrogen peroxide is decreasing in the catalyst lif etime study is that the concentration of bicarbonate is decreasing. Since the bicarbon ate dependence has been observed to be second-order, the loss of bicarbonate will ha ve a dramatic effect on the observed decomposition rate constant. In addition to examining the bicarbonate concentration, the hydrogen peroxide was examined as a po ssible source for the loss of activity. In addition to the loss of bicarbonate, it is known th at hydrogen peroxide is stabilized usi ng tin phosphates.10 The loss of activity may be due to manganese precipitation by the addition of phosphates to the solutions, as was observed when phosphates were used in attempting to contro l the pH of the decomposition solutions. The malachite green assay for phosphates was used to examine the stock 50% hydrogen peroxide solution. It was found that the stock hydrogen peroxide contained approximately 4 mM phosphate. For the cycl es being run, this equates to about 25 M of

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55 phosphate being added to each cycle. This amount of phosphate is enough to begin manganese precipitation. Multiple additions of distilled hydrogen peroxide and solid sodium bicarbonate Once it was determined that bicarbonat e was being lost during each cycle and phosphate being added, anothe r catalyst lifetime study was done. For these reactions, distilled hydrogen peroxide was used. This assured that no phosphates were being added to the solutions. In addition, solid sodium bicarbonate was added before each cycle in an attempt to stabilize the bicarbon ate concentration from cycle to cycle. The results of the catalyst lifetime study using these modi fications are seen in Figure 3-9. 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 03691215 # of H2O2 additions103* kobs, s-1 Day 1 Day 2 Figure 3-9. Plot of kobs versus # of hydrogen peroxide additions for the catalyst lifetime study using distilled hydrogen peroxide and adding solid sodium bicarbonate. The loss of activity is now due only to dilution and the inability to maintain the bicarbonate concentration at a constant value. The reactions are about half as fast as those not using the distilled hydrogen peroxide due to the metal contaminants th at are found in the peroxide. The loss of

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56 activity is now due to diluti on and the inability to mainta in a constant bicarbonate concentration. However, it must be noted that there is not as significant a loss in activity when using distilled peroxide and adding bicarbonate. This indicates that the manganese catalyst is not destroyed during the decompositi on of the peroxide, but can be regenerated by the addition of more hydrogen peroxide. Studies of the Manganese Source In addition to examining the lifetime of the manganese catalyst, the question of whether the source of the manganese was impor tant needed to be answered. For all kinetic experiments, the manganese source was manganese(II) sulfate. Two additional sources of manganese, permanganate and a Mn(IV)-TACN catalyst, were tested to compare their decomposition of hydrogen peroxide to the decomposition with manganese(II) sulfate. Potassium permanganate First, potassium permanganate was te sted at a concentration of 3 and 4 M in the presence of 0.20 M sodium bicarbonate and 0.100 M hydrogen peroxide. As seen in Table 3-1, the observed first-order rate constants for the manganese(II) sulfate and permanganate at the same concentration are well within experimental limits. Table 3-1. Comparison of observed rate constants for the decomposition of hydrogen peroxide, 0.100 M, in 0.20 M s odium bicarbonate with 3 and 4 M manganese(II) and permanganate. Errors reported are to the 95% confidence. Manganese Source Concentration kobs, s-1 Mn(II) 3 M (8.33 0.27) x10-4 MnO4 3 M (9.19 0.31) x10-4 Mn(II) 4 M (1.15 0.04) x10-3 MnO4 4 M (1.18 0.04) x10-3

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57 [MnIV(Me3TACN)(OMe)3]PF6 In addition to permanganate, a Mn(IV)-TAC N catalyst was synthesized for use as the manganese source in the hydrogen peroxide decomposition. In 1996, Kerschner et al.63 synthesized [MnIV(Me3TACN)(OMe)3](PF6), where Me3TACN is 1,4,7-trimethyl1,4,7-triazacyclononane (Figure 3-10). This catalyst wa s capable of oxidizing watersoluble olefins, specifically 4vinylbenzoic acid and styrylacetic acid, in the presence of bicarbonate. The stability of the catalyst wa s demonstrated by the repeated additions of hydrogen peroxide and alkene to produce epoxidized product.63 Figure 3-10. Molecula r structure of [MnIV(Me3TACN)(OMe)3](PF6).63 The Mn(IV) catalyst was synthesized usi ng the procedure reported by Kerschner.63 The compound is obtained by the reaction of manganese(II) chloride with 1,4,7trimethyl-1,4,7-tiazacyclononane in methanol in the presence of sodium peroxide. The complex is crystallized from methanol/w ater as a brown he xafluorophosphate salt. Mn(IV) catalyst stability Initially, the stability of the Mn(IV) cata lyst was examined sp ectrophotometrically in both the presence of bicarbonate and hydr ogen peroxide individua lly. A solution of

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58 1.00 M sodium bicarbonate and 0.108 M Mn(IV) catalyst was dissolved in water, and the catalyst was monitored at 345 nm for 2 hrs, Figure 3-11. The same procedure was also done in the presence of 0.50 M hydrogen peroxide, Figure 3-12. 0 0.2 0.4 0.6 0.8 1 1.2 0120024003600480060007200 Time, secAbs Figure 3-11. Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of 1.00 sodium bicarbonate. 0 0.2 0.4 0.6 0.8 1 1.2 0120024003600480060007200 Time, secAbs Figure 3-12. Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of 0.50 M hydrogen peroxide.

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59 After the catalyst was shown to be stab le in the presence of bicarbonate and hydrogen peroxide alone, experime nts were conducted to test its stability in the presence of bicarbonate and hydrogen peroxide t ogether. A solution of 1.00 M sodium bicarbonate and 0.108 M catalyst was dissolved in water. After about 350 seconds of monitoring, 25 L (0.100 M, final concentration) of hydrogen peroxide was added to the solution (Figure 3-13). Almost instantly, th e catalyst absorbance decays to 0. Within about 3 minutes, a new absorbance is detected and the solution is a yellow color. This new absorbance is very broad, having an absorbance of ~0.7 from ~250 nm to 500 nm, Figure 3-14. This absorbance is most likely due to the metal interaction with the Ndealkylated organic decay products More about the topic of N-dealkylation by Mn(II) with hydrogen peroxide in the presence of bicarbonate will be presented in a later section of this chapter. 0 0.2 0.4 0.6 0.8 1 1.2 1.4 0100200300400500600700 Time, secondsAbsorbance, A Figure 3-13. Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of 25 L (0.100 M, final concentratio n) hydrogen peroxide was done at 350 seconds. The absorbance first decays to 0 and within a matter of minutes, the solution is bright yellow.

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60 Figure 3-14. UV-vis specta of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6). The solid line is the spectrum in the presence of 1. 00 M sodium bicarbonate. The dotted line is the spectrum of the so lution after 1 eq of hydroge n peroxide was added. In a second experiment, 50 L (0.200 M, final concentration) hydrogen peroxide was added to the catalyst to determine whether the developing yellow product(s) was the result of a limited amount of peroxide. A solution of 1.00 M sodium bicarbonate and 0.108 M catalyst was dissolved in water. Th e absorbance at 345 nm was monitored as before (Figure 3-15). At 312 seconds, the pero xide was added to the solution. Unlike the first experiment, the development of the yello w color did not occur. It is therefore apparent that in the presence of only 1 equi valent of hydrogen peroxi de, the catalyst does not completely decay, allowing the yellow colo r to develop. When 2 equivalents of hydrogen peroxide are present, the cata lyst is able to completely decay. 0 0.2 0.4 0.6 0.8 1 1.2 0100200300400500600700 Time, secondsAbsorbance, A Figure 3-15. Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of 50 L (0.200 M, final concentratio n) hydrogen peroxide was done at 312 seconds. Even after 6 minutes, the yellow color does not develop.

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61 Finally, the [MnIV(Me3TACN)(OMe)3](PF6) catalyst was used with bicarbonate to decompose hydrogen peroxide. As wa s done with permanganate, 3 and 4 M [MnIV(Me3TACN)(OMe)3](PF6) was dissolved in water with 0.200 M sodium bicarbonate and the deco mposition was initiated by the addition of 25 L (0.100 M) hydrogen peroxide. As expected, th e observed rate constants for the [MnIV(Me3TACN)(OMe)3](PF6) catalyzed decomposition of hydrogen peroxide are similar to those for the Mn(II) ion decompositi ons (Table 3-2). These results indicate that the [MnIV(Me3TACN)(OMe)3](PF6) catalyst quickly decomposes to release the manganese ion, which then begins catalytic ally decomposing the hydrogen peroxide. If the catalyst did not quickly decompose, a la g in the decomposition of hydrogen peroxide might have occurred, however, th is is not seen experimentally. Table 3-2. Comparison of observed rate constants for the decomposition of hydrogen peroxide (0.100 M, final concentration) in 0.20 M sodium bicarbonate with 3 and 4 M manganese(II) and [MnIV(Me3TACN)(OMe)3](PF6). Errors reported are to the 95% confidence. Manganese Source Concentration kobs, s-1 Mn(II) 3 M (8.33 0.27) x10-4 [MnIV(Me3TACN)(OMe)3](PF6) 3 M (8.79 0.30) x10-4 Mn(II) 4 M (1.15 0.04) x10-3 [MnIV(Me3TACN)(OMe)3](PF6) 4 M (1.11 0.06) x10-3 Cis-trans Isomerization in the Mangan ese(II) Catalyzed Alkene Epoxidation Reactions of cis-alkenes to thei r corresponding epoxi des catalyzed by manganese(II) and hydrogen peroxide in bicarbonate solution will give some indication as to the nature of the oxygen transfer from the active oxygen species to the alkene. For example, in the case of peracid epoxidation of alkenes, cis-alkenes re act to give only the cis-epoxide, as seen in Figure 3-16. Th e oxygen is delivered in a concerted process which retains the stereochemistry of the reactant.

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62 R R' Cl O HOO R R' O H O O R'' RR' O + Cl O HO Figure 3-16. The concerted mechanism for the m-CPBA oxidation of nucleophilic alkenes resulting in the rete ntion of stereochemistry. The reaction of Mn(salen) organometalli c complexes with hydrogen peroxide, on the other hand, do not produce only the cis-e poxide from the cis-alkene, but the transepoxide as well. For these re actions, a two step process occurs where by the C-C sigma bond remains intact, but a carbon radi cal is produced as an intermediate, as seen in Figure 3-17. During the lifetime of the intermed iate carbon radical, before the oxygen bond of the epoxide is formed, the C-C sigma bond has the opportunity to rotate into the more stable trans conformation. In this way, the Mn(salen) catalysts will produce both the cis and trans-epoxides fr om the cis-alkene. R R' Mn(Salen) + H2O2R R' O Mn(salen) O R R' + Mn(salen) Figure 3-17. The stepwise oxida tion of an alkene by Mn(salen ) and hydrogen peroxide is shown. Cis/trans isomerization occurs in the transition state, where the C-C sigma bond is able to rotate into th e more stable trans configuration. Burgess et al.40 noted that under thei r conditions, using water/DMF solutions, the epoxidation of cis-stilbene produ ced both the cis and trans-stil bene oxides, as seen in Figure 3-18. This indicates that the active oxygen species does not add the oxygen in a concerted manner, as do the peracids, for example. If the oxygen were added in a concerted manner, there would be no trans e poxide present. However, Burgess followed

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63 these experiments in mixed solvent. The que stion remains as to whether the cis/trans isomerization will occur in pure aqueous solution. DMF/H2O Mn(II), H2O2, HCO3 -O + O Figure 3-18. The cis/trans isomerization noted Burgess in his epoxidation of stilbene using the Mn(II), H2O2, bicarbonate system using a mixed solvent system of DMF/H2O. (Burgess, 2002)40 To study the cis/trans isomerization in pure water, 4,4’-sulfonated stilbene was the obvious choice, based on Burgess’ use of stilb ene for the reactions in mixed solvent. Unfortunately, all attempts at direct synthesis by sulfonating stilbene using fuming sulfuric acid resulted in black tar. A lit erature search for the preparation of 4,4’sulfonated stilbene resulted in a single paper by van Es.64 The synthetic scheme is illustrated in Figure 3-19. All attempts at sy nthesizing the 4,4’-sulfonated stilbene failed. NH2SO3H Na2CO3NaNO2HCl H2O N SO3H NNaO3S O OHNaOH,CuCl2,H2O 30o C, 2hrs 1000C NaClNaO3S NaSO3 Figure 3-19. Synthetic scheme for synthesi s of 4,4’-sulfonated stilbene. (van Es, 1964)64 Next, two water-soluble alkenes were c hosen, cis and trans2-butene-1,4 diol. These alkenes were chosen for two reasons. Fi rst, they were freely water soluble at the operating pH of 8.4. Second, the epoxides of the alkenes are easily distinguishable by 13C NMR. This made analysis of the reacti ons relatively simple. The cis-alkene is commercially available, while the trans-alkene must be synthesized The trans-2-butene1,4-diol was synthesized using th e method of Schloss and Hartman.65 The synthesis

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64 requires the reduction of 2-butyne-1,4-diol by lithium aluminum hydride in THF. The epoxidation of both the cis a nd trans alkenes were accomp lished by the reaction with mCPBA. The epoxide products’ 1H and 13C NMR were compared to literature values for authentic samples, Figure 3-20. CaO CbOH HOCaCb MeOH cis-2,3-epoxybutane-1,4-diol CaO OH CbHOtrans-2,3-epoxybutane-1,4-diolCaCb MeOH CaO CbOH HOCaCb MeOH cis-2,3-epoxybutane-1,4-diol CaO OH CbHOtrans-2,3-epoxybutane-1,4-diolCaCb MeOH Figure 3-20. 13C NMR of cis and trans-2,3epoxybutane-1,4-diol in D2O using methanol as an internal standard.

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65 Cis/Trans isomerization reactions with cis-2-butene-1,4-diol The cis-2-butene-1,4-diol (0.60 M) was dissolved in D2O along with 1.00 M NaHCO3, and 10 M Mn(II). The reaction was ini tiated by the addition of hydrogen peroxide (final concentration 6.0 M). The reaction was monitored by observing the methylene peak in the 13C NMR using methanol as an internal standard. After 30 minutes, the methylene peak has decreased, bu t the appearance of the epoxide peaks for either the cis or trans epoxide cannot be seen (Figure 3-21, left). After 18 hrs (Figure 321, right), it appears that th e cis-alkene has been oxidized to any number of products, none of which have been identified at this time. From these data, a new water-soluble alkene was needed to examine the cis/tr ans isomerization of the Mn(II)/hydrogen peroxide/bicarbonate oxidation system in water. 30 minutesMeOH CH2 18 hoursMeOH 30 minutesMeOH CH2 30 minutesMeOH CH2 18 hoursMeOH Figure 3-21. Epoxidation of cis-2-butene -1,4-diol (0.60 M) with 1.00 M HCO3 -, Mn(II) (10 M), and H2O2 (6.00 M) after 30 minutes (left) and 18 hrs (right). Cis/Trans isomerization of maleic and fumaric acids Two new alkenes were chosen to study the cis/trans isomerization in pure water. These alkenes are maleic and fumaric acid. The structures of these alkenes at the

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66 operational pH of 8.4 are shown in Figure 3-22. While these may appear to be electrophilic alkenes, the domi nant resonance structure at the operating pH allows these alkenes to react as nucleophilic alkenes. Mo re on this topic will be presented in the following chapter. Once again, the determination of epoxi de products are conveniently made using 1H NMR, as seen in Figure 3-23. O O O O O O O O MaleicAci d Fu m a r icAci d Figure 3-22. The structures of maleic and fumaric acids at the operating pH of 8.4. Figure 3-23. 1H NMR of maleic and fu maric acid oxides in D2O using methanol as an internal standard. When 0.10 M maleic acid was allowed to react with 1.00 M peroxide in the presence of 0.80 M sodium bicarbonate and 10 M Mn(II), a 34% conversion to epoxide was observed by NMR in 15 min, as seen in Figure 3-24.

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67 Figure 3-24. Epoxidation of maleic acid by hydrogen peroxide and manganese(II) in the presence of bicarbonate after 15 min. Reaction conditions: 0.10 M maleic acid, 1.00 M H2O2, 0.80 M NaHCO3, and 10 M Mn(II). Of the 34% epoxide formed, 74% was the fumaric acid oxide and 26% was maleic acid oxide. This result indicates th at even in pure water, the oxyg en is not being added to the

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68 alkene in a concerted mechanism, such as that seen in the oxidation by m-CPBA. This indicates that at some time during the epoxi dation, the C-C sigma bond of the alkene has the opportunity to rotate into the more stable trans conformation before closure of the epoxide ring. This experiment does indicate that radical or car bocation formation is probable in the epoxidation of nucleophilic alke nes, a discussion of possible routes for the addition of the oxygen and rotation into th e trans isomer will be discussed with the possible mechanisms of the reaction. The observation that cis alke nes react with th e active oxygen species to give the trans epoxide indicates that radical chemistry may play a role in the epoxidation of alkenes. This does not, how ever, indicate that free radicals are responsible for the epoxidation. Reaction mechanisms that are similar to those for Mn-salen epoxidation catalysts could also expl ain the rotation about the C-C bond during the epoxidation reaction. In addition, reacti ons involving electrophilic alke nes, which will be presented in the next chapter, led us to question Sych ev’s proposed hydroxyl radical mechanism. In the epoxidation of electrophilic alkenes, as was the case for the nucleophilic alkene epoxidations, the reactions clean ly yielded the epoxide product s with no indication of any radical products.66,67 In an attempt to exclude fr ee radicals as a possible reaction pathway, an examination of the radical traps used by Sychev47 was conducted. Examination of Sychev’s Radical Trap Experiments As discussed in the intr oduction, the combination of hydrogen peroxide and iron(II), Fenton’s reagent, is a useful met hod for the production of hydroxyl radicals. Fenton’s reagent can then be used to oxidize organic molecules. For instance, benzene can be oxidized to form biphenyl and phenol in the presence of Fenton’s reagent, as seen in Figure 3-25.67

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69 Fe2+ + H2O2 Fe3+ + HO + HO-HO + H OH HO OH + H2O HO H H H H H OH + H acceptor + H2O Figure 3-25. Fenton’s reagent can be used to oxidize benzene to phenol and biphenyl. In 1977, Sychev et al.47 began investigating the ro le of manganese(II) in the disproportionation of peroxide in bicarbonate buffered solu tions. His assumption was that manganese(II) reacted similarly to iron(II) ions in Fenton type chemistry. Following this assumption, a hydroxyl radical based mechanism was proposed, shown by Equations (3-4)-(3-14), the sum of which is the deco mposition of hydrogen peroxide to molecular oxygen and water. As with Fenton type chem istry, free hydroxyl and peroxy radicals are formed in this mechanism, along with the carbonate radical anion. [Mn(HCO3)2] + H2O2 [Mn(HCO3)2]+ + OH+ •OH (3-4) [Mn(HCO3)2]+ + H2O2 [Mn(HCO3)2] + H+ + HOO• (3-5) [Mn(HCO3)2] + •OH [Mn(HCO3)2]+ + OH(3-6) [Mn(HCO3)2]+ + HOO[Mn(HCO3)2] + O2 -• + H+ (3-7) [Mn(HCO3)2] + HOO• [Mn(HCO3)2]+ + HOO(3-8) [Mn(HCO3)2]+ + OH[Mn(HCO3)2] + •OH (3-9) •OH + H2O2 O2 -• + H+ + H2O (3-10) 2O2 -• + 2H2O2 2O2 + 2OH+ 2 •OH (3-11) •OH + HOO• H2O + O2 (3-12) •OH + HCO3 CO3 -• + H2O (3-13)

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70 CO 3 -• + H 2 O 2 O 2 -•+ H+ + HCO 3 (3-14) 2H2O2 2H2O + O2 In order to support his clai m of a free hydroxyl radical pathway, Sychev employed a set of experiments using N,N-dimethyl-4-nitrosoaniline (DMNA) as a free hydroxyl radical trap. In his experiment s, he studied the production of O2(g) as a function of time with increasing amounts of DMNA, Figure 3-26. Figure 3-26. The influence of inhibitors on the catalase process in the Mn(II)/HCO3/H2O2 system. [Mn(II)] = 4 x 10-6 M, [H2O2] = 0.10 M, pH 7.0, [HCO3 -] = 0.4 M, and T = 25 C: 0) kinetic cu rve with no inhibitors; 1) 2), 3), and 6) in the presence of DMNA as the inhib itor(at concentrations of 1 x 10-5,1.5 x 10-5, 2 x 10-5,and 4 x 10-5 M respectively; 4)in the presence of tetranitromethane(4 x 10-5 M); 5) in the presen ce of hydroquinone (1.5 x 10-5 M); 7) decomposition of H2O2 without Mn(II) ion (blank e xperiment). (Sychev, 1977)47 As expected for his free hydroxyl radical pathway, Sy chev observed that when DMNA was present, the production of O2(g) was suppressed to the background disproportionation of H2O2 without the addition of metal, Figure 3-26, line 7. As time

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71 progressed, the O2(g) production would begin to increas e back to the purely catalytic production of O2(g), as shown in Figure 3-26, line 0. The conclusion reached by Sychev was that the DMNA was trapping the free hydroxyl radicals produced from the disproporti onation, Equations 3-4, 3-9, and 3-11. Without the presence of free hydroxyl radicals to carry the reaction, O2(g) production would be the same as the uncatalyzed O2(g) production, Figure 3-26, line 7. Eventually, O2(g) production would begin to follow that of the catalyzed reaction as the concentration of DMNA was reduced and an in crease in free hydroxyl radicals occurred. This is seen in Figure 3-26 as all of the inhibited r eactions eventually begin to produce O2(g) at the same rate as the uninhibited reaction, Figure 3-26, line 0. In all of Sychev’s papers, the use of ra dical traps, such as DMNA, provide the entire basis for a free hydroxyl radical mechan ism. In none of his papers, however, did Sychev identify the organic produc ts of the reactions with DM NA. In the current study, it has been hypothesized that instead of a decomposition pathway that requires hydroxyl radicals, the mechanism of peroxide di sproportionation may proceed through a high valent metal oxo species, or by carbonate radi cal anions, which are proposed in Sychev’s model. The data presented for the interruption of hydrogen peroxide decomposition by Sychev could be the result of trapping of carbonate radical ani ons, instead of hydroxyl radicals. A discussion on the reactivity of carbonate radicals will be presented in the proposed mechanism for the oxid ation of the radical traps. The use of a high valent metal oxo species could also explain Sychev’s loss in O2(g) production seen w ith the use of DMNA. Instead of the O2(g) production being inhibited by radical interrup tion, the oxygen normally bei ng released as molecular

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72 oxygen would instead be transferred to DMNA. In order to understand the effect DMNA is having on the hydrogen peroxide decompos ition reaction, the organic products from the reaction need to be identif ied. Once the oxidized orga nic products are identified, a clearer understanding of the reaction mechanism may be possible. A series of experiments were first co nducted to determine what the oxidized products of DMNA could be. Initially, pot assium peroxymonosulfate was employed as the oxidant. This experiment was done to determine what the product of a pure oxygen transfer would be, since pe roxymonosulfate is an excel lent electrophilic oxidant.68,69 When peroxymonsulfate was allowed to react with DMNA in a 1:1 molar ratio, N,Ndimethyl-4-nitroaniline was pr oduced nearly quantitatively af ter 30 minutes, as expected for an oxygen transfer to the nitroso moie ty. A second control experiment was done using H2O2 and HCO3 -, only. When one equivalent of H2O2 was added to 0.40 M chelexed HCO3 -, N,N-dimethyl-4-nitroaniline was produced after 1 hr. In the presence of hydrogen peroxide alone, no reac tion was detected even after 24 hrs. This result indicates that solutions of pe roxycarbonate are able to conve rt the nitroso moiety to the nitro without the addition of any metals. Any products, other than the nitro compound, are then the result of the addition of the metal cations. N N O N NO2HSO5 Figure 3-27. The reaction of N,N-dimethyl-4-nitrosoaniline w ith peroxymonosulfate to yield N,N-dimethyl-4-nitroaniline.

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73 In a second set of reactions DMNA was oxidized using H2O2 in bicarbonate with Mn(II). H2O2 was used as the terminal oxidant in 50x molar excess over the DMNA. The need to increase the H2O2 concentration to such a degree over the DMNA concentration is due to the fact that H2O2 disproportionation is much faster than the oxidation of DMNA. When reactions using only one equivalent of H2O2 were conducted, starting material was the only recovere d organic compound. The final sodium bicarbonate and Mn(II) concentratio ns were set at 0.40 M and 4.0 M, respectively. DMNA oxidations were conducte d by first dissolving the organic substrate in a mixture of CH3CN:H2O (30:70 (v: v)) with a solution of the manganese(II) sulfate. Equilibrated solutions of H2O2 and sodium bicarbonate were then slowly added dropwise over about 30 minutes. Reactions were cons idered to be complete when the production of O2(g) from the H2O2 disproportionation ceased. This was usually 10-15 minutes after the final addition of the H2O2/HCO3 solution. Since the reac tion is highly exothermic, ice was often employed to keep the temperat ure from exceeding 65C. Once the reaction reaches 65C, the H2O2 disproportionation becomes vigorous enough to cause the reaction mixture to boil out of the reaction flask. At the e nd of the reaction, the organic products were extracted into chloroform which was then dried over magnesium sulfate. The solvent was then removed under reduced pressure to give crude product, which was analyzed by 1H NMR, Figure 3-28. The 1H NMR of the crude product shows multiple products, three of which have been identified at this time. Aromatic peaks still remained and retained the characteristic proton signal for a disubstituted aromatic comp ound. It was also noted that new peaks in

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74 the 4 5.5 ppm region had appeared. The methyl peaks of the amine portion of the molecule seemed to have remained, but they were shifted upfield. Figure 3-28. 1H NMR of the crude reaction mixture after an oxidation of N,N-dimethyl-4nitrosoaniline by hydrogen peroxide in the presence of bicarbonate and Mn(II). Reaction conditions: N,N-dimehtyl-4-nitrosoaniline (1 g, 6.66 mmol), 0.400 M sodium bicarbonate, 10 M Mn(II), 6.64 M H2O2, 1 hr. After analysis by 1H NMR, gas chromatography was employed to help determine the number of products obtained from the reacti on. GC analysis showed there were four volatile products formed. It can only be stat ed that these products are volatile, since GC will only separate compounds that can be vol atized and have a low enough boiling point to remain in the gas phase. Both 1H NMR and GC proved that none of the products obtained were that of N,N-dimethyl-4-nitrsoaniline, whic h should have been the case

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75 since the nitroso moiety is easily oxidized by peroxycarbonate. Figure 3-29 is a GC trace using the same method employed for the crude reaction mixtur e, Figure 3-30, for which a peak is not observed at 14.6 min, indicating that all of the N,N-dimethyl-4-nitrsoaniline has been converted to other organic products. 14.637 min 14.637 min Figure 3-29. GC trace for a standard of N,N-dimethyl-4-nitrsoaniline. Non-linear gradient for 30 minutes detection by FID.

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76 15.260 min 17.462 min 19.983 min 24.421 min 15.260 min 17.462 min 19.983 min 24.421 min Figure 3-30. GC trace for the crude react ion material from the oxidation of N,N-dimethyl4-nitrsoaniline from Figure 3-25. Lack of a peak near 14.637 min proves that no starting material remains. GC c onditions: non-linear gradient for 30 minutes, Detection: FID.

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77 After 1H NMR and GC analysis, the reaction material was applied to a silica column employing chloroform as eluant. Frac tions were collected and analyzed by GC. Fractions 4 and 9 were found to contain organic products and their GC traces and 1H NMR are shown in Figure 3-31 and 3-32. The two compounds that we re separated were found to be N,N -dimethyl-4-nitroaniline and 4-ni troaniline, by comparison with authentic samples. These two compounds acco unt for about 80% of the crude reaction mixture, the last 20 % bei ng the peaks at 19.983 and 24.421 min, Figure 3-27. The two remaining organic compounds remained at the top of the column. 17.418 min ppm NO2N HaHbHcHaHbHc 17.418 min ppm NO2N HaHbHcHaHbHc Figure 3-31. GC trace (left figure) and 1H NMR (right figure) for Fraction 4 of the silica column. Identification of the product as N,N-dimethyl-4-nitroaniline was confirmed by comparison with a GC trace and 1H NMR of an authentic sample.

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78 15.182 min HaHbHcCHCl3 NO2NH HcHaHb 15.182 min HaHbHcCHCl3 NO2NH HcHaHb Figure 3-32. GC trace (left figure) and 1H NMR (right figure) for Fraction 9 of the silica column. Identification of the product as 4-nitroaniline was confirmed by comparison with a GC trace and 1H NMR of an authentic sample. Once 4-nitroaniline was iden tified as a product, it was a pparent that the amine was dealkylating. It was then likely that one of the other unidentified peaks was N-methyl-4nitroaniline. A 1H NMR was acquired of a sample of the pure material and compared with the crude reaction material, Figure 3-33. It was found that N-methyl-4-nitroaniline peaks were also present in th e crude reaction mixture. The only peak not identified was the peak at 24.421 min, which accounts for le ss than 5 % of the crude reaction material.

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79 HaHbHcHd NO2NHcHaHbHdCHCl3 HaHbHcHd NO2NHcHaHbHd HaHbHcHd NO2NHcHaHbHdCHCl3 Figure 3-33. 1H NMR of an authentic sample of N-methyl-4-nitroaniline. Comparison with the crude reaction mixture conf irms its presence as a product. For these products to be observed, N-dealkylation must be responsible for cleaving the C-N bond of the amines. In the ca se of DMNA, the cleaved carbon group is formaldehyde based on current literature th at will be discusse d shortly. In the 1H NMR and GC analysis performed on the DMNA react ions, formaldehyde was never detected. This was probably due to the fact that th e high temperature of the reaction volatilized formaldehyde. In order to observe th e aldehyde produced from the reaction, N,N-diethyl4-nitrosoaniline (DENA) was c hosen as the next substrate. Reactions were performed us ing a 50x molar excess of H2O2 over DENA and the final concentrations of bicarbonate and Mn(II ) were 0.40 M and 4.0 M, respectively. As

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80 with the DMNA oxidations, the two majo r products formed were those from Ndealkylation, the N-ethyl-4-nitroaniline and 4-nitroaniline. In order to trap the aldehyde produced the reaction was performed in a multinecked round bottom flask with a stream of nitrogen gas pa ssing over the reaction. As the reaction progressed, the aldehyde would evaporate from the reaction solution and travel with the stream of nitrogen, which was passed th rough an ice cooled trap to condense the aldehyde. An initial Tollen’s test of the condensed solution gave a positive result indicating the presen ce of an aldehyde. GC and 1H NMR analysis of the condensed solution compared with an authentic samp le proved that the pr oduct was acetaldehyde, Figure 3-34. Acetaldehyde Acetonitrile 1.542 min 1.876 min ppm CHCl3 O HaHbHaHb Acetaldehyde Acetonitrile 1.542 min 1.876 min ppm CHCl3 O HaHbHaHb Figure 3-34. The solution collected from the reaction of N,N-diethyl-4-nitrosoaniline was analyzed by Gas Chromatogr aphy (lower figure), whic h was compared to an authentic sample of acetaldehyde. 1H NMR (top figure) of the solution also confirmed that the product was acetaldehyde.

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81 Proposed Mechanism of N -dealkylation The N-methyl-4-nitroaniline, N-ethyl-4-nitroaniline, acetaldehyde, for DENA oxidation, and 4-nitroaniline observed from the oxidation of N,N-dimethyl-4nitrosoaniline (DMNA) and N,N-diethyl-4-nitrosoaniline (D ENA) by hydrogen peroxide catalyzed by manganese(II) in the presen ce of bicarbonate are due to oxidative Ndealkylation. The mechanism of meta l catalyzed oxidative N-dealkylation of amines currently reported in the literature is based on expe riments conducted to determine the mechanism of cytochrome P-450 oxidative N-dealkylation. However, after more than 40 years of study, uncertainties still exis t as to the mechanistic details in the oxidative N-dealkylation of amines by cytochrome P-450.70-74 The current literature recognizes two dominant pathways for the production of the observed products in the presence of cytochrome P450, hydrogen atom abstraction a nd single electron transfer.71-81 In addition, the products obs erved could be the result of oxidation by the carbonate radical anion, however, little is known about th e reactivity of the car bonate radical anion with organic substrates, speci fically, whether the products ar e phenols or biphenyls, such as are observed in reactions with the hydr oxyl radical. Recently, the carbonate radical anion has been proposed to be the active oxida nt in the peroxidase activity of the Cu,ZnSOD in the presence of hydr ogen peroxide and bicarbonate.82-87 The carbonate radical has been implicated in these works to be the source of observed protein damage. Secondorder rate constants of 4.2 x 108 M-1s-1 have been observed for the radical with several organic compounds including, i ndole, and its derivatives,88 but no rate constants were found for reactions with amines.

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82 The mode of attack of the carbonate radica l anion could occur in one of three ways: 1) a single electron transfer from the substrat e to the radical, 2) addi tion of the radical to the substrate, or 3) hydrogen abstraction by the radical. Unfortunately, the organic products for reactions with the carbonate ra dical anion have not been identified. Therefore, it is not know whether oxidative N-dealkylation may be a product of reactions with the carbonate radical anion. The proposed pathways are presented in Figure 3-35, along with potential manganese complexes that are proposed in th is study, carbonate radi cals could also be replaced for any of the manganese species. Sp ecific details about the identity of these manganese species and how they are gene rated in the hydrogen peroxide oxidation system with bicarbonate will follow in a later section. The figure only represents the products, a secondary amine and a molecule of ketone or aldehyde, depending on the substrate, of the first N-dealkylation of a tertiary amine. A tertiary amine can be oxidized a maximum of three times, as ammonia will be the final nitrogen containing compound of the last oxidative N-dealkylation. Also, amines can only be N-dealkylated when a hydrogen is present on the carbon to the nitrogen.

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83 R N R' C R'' H R'''E l e c t r o n T r a n s f e rHyd r ogen A b stra c t i o nR N R' C R'' H R''' R N R' C R'' R''' R N R' C R'' R''' R N R' C R'' R''' OH R NH R' + R''R''' O AB C D E F G [MnIV(O2-)]2+[MnIII(O2-)]+H+[MnIV(OH)]3+"MnIII + HO [MnIII(OH)]2+"MnII + HO "Radical Rebound"caged hydroxy radical" -H+"caged hydroxy radical" -H [MnII(OH)]++OH-+ MnII Figure 3-35. The proposed mechanism for the oxidative N-dealkylation of amines by hydrogen peroxide in the presence of bicarbonate as catalyzed by manganese(II). The secondary amine pr oduced can cycle again as long as it contains a hydrogen on the carbon to the nitrogen. A second molecule of aldehyde or ketone will also be produced. For the hydrogen abstraction pathway, the oxidation of the amine begins with the abstraction of a hydrogen on the carbon to the nitrogen on the substrate by the [MnIV(O2-)]2+ complex, or a carbonate radical. This abstraction may be easier to recognize by using the protonated form of the [MnIV(O2-)]2+ complex, [MnIII(OH)]3+.

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84 This complex can be imagined to be the “caged” hydroxyl radical with a MnIII ion. The abstraction of the hydrogen from the subs trate is accomplished using the “caged” hydroxyl radical to yield the [MnIII(OH)]2+ complex and a carbon centered radical, Figure 3-35, pathway B. The [MnIII(OH)]2+ complex that is produced may also be imagined to be a “caged” hydroxyl radical with a MnII ion, [MnII(OH)]2+. This “caged” hydroxyl radical can then rebound with the carbon cente red radical, Figure 3-35, pathway E, to yield the carbinolamine and a MnII ion. The resulting carbinolamine intermediate then decomposes to yield the observed products, the N-dealkylated amine and a molecule of aldehyde or ketone, depending on the substrate. The carbinolamine intermediate has been isolated under certain conditions depending on the amine used.89-92 The second proposed pathway for the oxidative N-dealkylation of amines occurs via a single electron transf er. Initially, the [MnIV(O2-)]2+ complex, or a carbonate radical, will oxidize the amine by removing a single electron from the lone pair on the nitrogen to form the [MnIII(O2-)]+ complex (a carbonate dianion in th e case of the carbonate radical) and a nitrogen centered radica l cation, Figure 3-35, pathwa y A. The nitrogen centered radical cation can be converted to the car bon centered radical by loss of a proton, at which time the above scheme can be followed to give the oxidative N-dealkylated products. On the other hand, the [MnIII(O2-)]+ complex is again a “caged” hydroxyl radical when protonated. This complex can then abstract a hydr ogen from the nitrogen centered radical cation to yield the imine and [MnII(OH)]+, Figure 3-35, pathway D. Once the imine is formed, attack of a hydroxi de ion on the electrophilic carbon of the imine, Figure 3-35, pathway F, yields th e identical carbinolamine as the hydrogen abstraction pathway, which then decompos es to yield the same final products.

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85 Support for the Single Electron Transfer Pathway Cytochrome P-450 is an important iron por phyrin enzyme responsible for a number of oxidative and reduc tive transformations.72 Much of the current literature on the mechanism of oxidative N-dealkylation by cytochrome P-45 0 indicates that the single electron transfer (SET) mechanism is the more probable pathway over the hydrogen atom transfer (HAT).72-74 Support for the SET mechanism has come from the work of Miwa et al.,73,74 in which experiments were conducted using hydrogen and deuterium isotope labeled substrates and analyzi ng the resulting product distribu tions. A set of experiments using N,N-dimethyl-2-amino-2-methyl-3-phenylpr opane, where the two methyl groups bonded to the nitrogen were deuterated, were conducted to dete rmine the deuterium isotope effect. The measured isotope eff ects were observed in the range of 0.946-1.12 with a mean of 1.04 0.06, a value close to unity. N Figure 3-36. The structure of N,N-dimethyl-2-amino-2-methyl-3-phenylpropane, the substrate used by Miwa et al.73,74 for use in experime nts with cytochrome P450 on the oxidative N-dealkylation mechanism. This small intermolecular deuterium isotope effect indicates that the breaking of the -carbon-hydrogen bond is not th e rate determining step as would be the case in the hydrogen abstraction pathway, Figure 3-35, path way B. If the mechanism were to follow the hydrogen abstraction of Figur e 3-35, pathway B, deuterium isotope values should be much greater than 1, since the observed rate for the substrate with all of the N-bound methyl groups having hydrogens would be larger than the observed ra te where all of the N-bound methyl groups ha d been deuterated.

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86 Solvent Isotope Effect In addition to monitoring the kinetic e ffects of bicarbonate, hydrogen peroxide, and manganese, Bennett46 also observed a large, inverse solv ent isotope effect (0.49 0.05 in 100% deuterium oxide) when the oxidation of nucleophilic alkenes was performed in increasing amounts of deuterium oxide. A prot on inventory study was also carried out to determine the number of protons undergoing ex change during the reaction. From these data, it was determined that only one proton was undergoing ex change during the reaction. Unfortunately, no explanation was give n as to why such a large, inverse isotope effect was found for the epoxidation of th ese alkenes using the manganese catalyzed system. Solvent isotope effect experiments were conducted a second time to confirm their existence and that the value was indeed large and inverse. Sulfonated styrene (1.0 mM) was epoxidized in pure water and deuterium oxi de with 1.00 M sodium bicarbonate, 0.5 M Mn(II), and 0.3 and 1.0 M hydrogen per oxide. The observed first-order rate constants can be found in Table 3-3 a nd are similar to those found by Bennett.46 Table 3-3. Comparison of first-order rate constants for the epoxidation of sulfonated styrene in H2O and D2O. Reaction conditions: 0.001 M SS, 1.0 M Sodium Bicarbonate, 0.50 M Mn(II). [H2O2], M kobs, s-1 (H2O) kobs, s-1 (D2O) Solvent Isotope Effect ( kH/ kD) 0.30 3.67 x10-3 7.65 x10-3 .48 1.0 1.68 x10-3 3.57 x10-3 .47 1.0 2.01 x10-3 4.37 x10-3 .46 “Solvent isotope effect” is the effect of deuterium versus protium solvents often used in discussions of kinetic and equilibrium processes.93 There are three main factors that affect the rates of reactions in these solv ents: 1) The solvent may be a reactant 2) the reactant may rapidly exchange isotopically labeled hydrogens with the solvent or 3) the

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87 nature of the weak solute-solvent inte ractions may change during the activation.94 For any solvent isotope effect, a combination of a ll three factors most likely will contribute to the overall isotope effect. Similar large, inverse isotope effects ar e rare, but have been reported in the literature.95-97 Usually, inverse isotope effects are caused by the protonation of an electronegative atom, such as oxygen, in a rapi d equilibrium prior to the rate determining step. For example, Pritchard and Long95 have investigated th e hydrolysis of simple epoxides in deuterium oxide-water mixtures The proposed mechanism is shown in Figure 3-37. O R'' R''' R R' H+O R'' R''' R R' H Rapid Equilibrium O R'' R''' R R' H OH R'' R''' R R' Rate Determining Step OH R'' R''' R R' H2O OH R'' R ''' R R' OH Fast Figure 3-37. Mechanism of epoxi de hydrolysis in acidic media. Initially, there is a rapid equilibrium be tween the epoxide and a proton, followed by the rate controlling step, cleaving of one of the epoxide C-O bonds to give the carbocation alcohol intermediate. The last step of the reaction is the addition of water to the carbocation and is a fast step. In the case of alkene epoxidation, tw o possible species coul d be protonated in solution, the high valent metal oxo species or carbonate radical anion. As is the case for

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88 the epoxides, the protonated metal oxo species would be a strong acid and the reaction would need to be performed at lower pH for any appreciable c oncentration of the protonated species to be attainable for reacti on. Since the pH of the performed reactions was 8.4, this is not a likely explanation fo r the observed solvent isotope effect. The protonation of the carbonate radical will be discussed shortly. A proton transfer during the ra te determining step could also lead to an inverse isotope effect. In the epoxi dation of an alkene by a metal oxo complex, no proton exchange can be envisioned as shown in Figure 3-38. Mn(II)O O O O O + [Mn(II)CO3]-Mn(IV)O 2+ O + Mn2+Mn(IV)O Mn(III)O O + Mn2+2+ 0 Figure 3-38. Possible epoxidation routes th rough a manganese(IV) oxo complex. None of the envisioned reactions has a proton transfer. A proton transfer could be possible for the epoxidation of a nucl eophilic alkene in which the manganese is only act ing as a Lewis acid, Figure 3-39. This would be the simplest mechanism for the epoxidation of the alkene. In this scenario, the manganese(II) would facilitate the cleaving of the O-O bond and stabilize the formed carbonate leaving group. However, this mech anism does raise questions as to why other

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89 metal cations with similar Lewis acidities do not have similar activities.40 It does, however, satisfy the requirement noted in the epoxidation of styrene in micellar solution that the active species be uncharged. By simply binding a bicarbonate anion ( or another peroxycarbonate) to the metal, the positive ch arge of the complex will be neutralized. Mn(II)O O O O H + Mn(II) O O O O H + O + [Mn(II)HCO3]+ Figure 3-39. Mechanism of oxygen transfer by attack of the alkene on a manganese(II) bound peroxycarbonate. The proton transf er in the transition state may account for the inverse is otope effect observed. It has also been proposed that the carbonate radical anion is form ed in solutions of bicarbonate with hydrogen peroxi de and manganese(II). The carbonate radical is a potent one electron oxidant (1.5 9 V vs SHE at pH 12),98 that is believed to exist as an acid/base couple, although the pKa has yet to be firmly established. There is literature to suggest that the pKa lies in the range of about 7.999 to 9.5.100 However, there has also been the suggestion that it is strong acid (pKa < 0).101 If the protonated carbona te radical were the active oxidizing agent for the epoxidation of nucleophilic alke nes, the epoxidations perfor med in cationic and anionic micellar solutions would suggest that the pKa lies somewhere in the range of 7.9 – 9.5, since no difference in epoxidation rates were observed and the protonated radical is uncharged. The protonated carbonate radi cal would then be the primary oxidizing species. A proposed mechanism is suggested in Figure 3-40. The existence of the protonated radical would explain the inverse solvent isotope eff ect. As noted earlier, the protonated species will be in higher concentr ation in deuterium oxide than in water, therefore, increasing the con centration of the active oxidant With an increase in the

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90 active oxidant, an increase in the rate of the epoxi dation would occur to give the observed inverse effect. R HCO3HCO3 -R HCO3R O O OH R O + H+ + CO2 Figure 3-40. Oxidation of a nucleophilic alkene by two sequential reactions with the carbonate radical. The carbocation interm ediate formed explains the loss of retention observed for cis-alkenes. In addition to measuring the solvent is otope effect on the nucleophilic alkene epoxidation, the effect on the hydrogen per oxide decomposition was also studied. The results of reactions performed in deuterium oxi de are presented in Table 3-4. From these experiments, it is observed th at the hydrogen peroxide deco mposition has a large, normal solvent isotope effect. Table 3-4. Comparison of solvent isotope e ffect for hydrogen peroxide decomposition. Reaction Conditions: 0.40 M HCO3 -, 0.10 M H2O2 [Mn(II)] kH/ kD 3 M 1.47 0.05 4 M 1.49 0.06 Proposed Mechanism Based on the information gleaned from prev ious kinetic experiments, which were reported earlier, on hydrogen peroxide decomposition and nucleophilic alkene epoxidation, in both water and deuterium oxide, the following mechanism is proposed in this work (Figure 3-41). The free hydr oxyl radical mechanism proposed by Sychev seems unlikely given the clean epoxidations of nucleophilic al kenes with a lack of any side products (polystyrene or 2-phenyl etha nol). This does not mean, however, that radical chemistry is not occurring. As disc ussed earlier, carbonate radicals are more than

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91 likely being generated in this system. This was discussed earlier in relation to the products observed in the oxidation of N,N-dimethyl-4-nitrosoaniline and N,N-diethyl-4nitrosoaniline. Chapter 4 will discuss the epoxidation of electrophilic alkenes by hydrogen peroxide in the presen ce of bicarbonate and manganese As with the reactions of nucleophilic alkenes, no side products indicative of fr ee hydroxyl radical chemistry are observed in the reactions that include manganese. Also, it will be shown that the oxidation of electrophilic alkenes occurs exclusively when hydroperoxide and hypochlorite are used. MnIVO2-HO2CO OCO2H OCO2H "A"-S SO + H+ + 2 HCO3 +[MnII(HCO3)]+MnIVHO2CO OCO2H OCO2H OO HCO3 -[Mn(HCO3)(HCO4)] HCO4 -+ [Mn(HCO3)]+H2O2 + HCO3 -MnII + HCO3 -K1K2K3K4K5k6H2O2"B" Radical Decomposition of Hydrogen Peroxide by Carbonate Radical [Mn(II)(HCO3)2(HCO4)]-S SO + 2 HCO3 -S k7SO k8O2 + Mn(II) + 3HCO3 Figure 3-41. Proposed mechanism for hydrogen peroxide decomposition and nucleophilic alkene epoxidation in the presence of bicarbonate catalyzed by Mn(II). The active epoxidation catalyst is proposed to be either the high oxidation state manganese oxo complex or the carbonate radical. High oxidation state metal oxo complexes have been proposed for the ir on(III) tetrakis (pentafluorophenyl) porphyrin complex with hydrogen peroxi de as observed by Traylor.20 The generation of the manganese catalyst would occur via the two el ectron oxidation of manga nese, as seen in

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92 Figure 3-42. A high oxidation state Mn(IV) oxo complex is si milar to the epoxidation route of the Mn-porphyrin102-104 and Mn(salen)105,106 complexes. A discussion of the nucleophilic alkene epoxidation pathway will be presented first, followed by the Fenton reactions of the catalytic species. O O MnIIO O HCO3O3CH MnIVO2-HCO3HCO3HCO3H "A" Figure 3-42. Proposed generation of the active manganese catalyst by from the [MnII(HCO3)2(HCO4)]complex by a 2 electron oxidation of manganese to form a high valent [Mn-O2-]2+ complex. As was noted in Chapter 2, the reaction ra te was not significan tly different when alkene epoxidations were performed in cati on or anion surfactants, indicating that the active epoxidation catalyst must be uncharg ed. These data also indicate that the carbonate radical anion may not be the active epoxidation cata lyst, as a reduction in the rate of the reaction in the presence of an ionic surfactant would be expected, unless the radical is protonated as mentioned earlier. The reaction of the intermediate “A” w ith hydrogen peroxide could form the peroxo intermediate “B”. This complex is similar to that seen for the addition of a molecule of hydrogen peroxide in the oxidation by MTO. Th e alkene can attack one of the electrophilic oxygens of the peroxo ligand to yield the epoxi de. This is analogous to the observation by Espenson in the epoxidati on of olefins by MTO/hydrogen peroxide. As with MTO, this intermediate reacts with olefins at a slower rate than the metal oxo complex. This could explain why a decrea se in the rate of nucleophilic oxidation

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93 catalyzed by manganese(II) is observed at hi gher hydrogen peroxide concentration. As the concentration of hydrogen peroxide increases, so will the concentration of “B”. Since the oxidation of the alkene by “B” is slower, the reaction will slow down as hydrogen peroxide increases, as observ ed in the hydrogen peroxide de pendence. If, however, the carbonate radical is responsible for the oxidation, the genera tion of carbonate radicals will decrease with an increasing hydrogen peroxide concentration. This will be demonstrated in the numerical simulations for the hydrogen peroxide decomposition. In addition to the pathway involving the pe roxycarbonate oxidati on of the metal to generate the two electron metal oxo complex, [MnIV(O2-)]2+, a one electron oxidation begins Fenton type chemistry. However, it is proposed that hydr oxyl radicals are not generated under these conditions since bicarbonate acts as a scavenger for the hydroxyl radicals to generate the carbona te radical anion. As discussed previous ly, there is little known in the current literatu re about the reactivity of the carbonate radical anion oxidation of organic substrates. Numerical Simulation of the Proposed Mechanism Unfortunately, an analytical solution to th e rate expression for this reaction cannot be found. However, numerical si mulation can be used as a tool to model the kinetics of a reaction for which an analytical solution ca nnot be derived. Th erefore, numerical simulation provides a powerful t ool to test kinetic models to determine whether they fit the observed data. While a particular simula tion may fit the experime ntal data, it by no means indicates that this is the true or accurate mechanism of the reaction. Multiple simulations may, in fact, yield good fits to observed data. When a simulation fits the observed data, the simulation may provide a new angle on possible experiments with which to investigate the reaction. If the experiments are performed and the outcome is

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94 contradictory to the simulated data, then that mechanism must be modified or replaced in an attempt to refine the model of the r eaction. Kinetica 2003 was used for all the numerical simulations that will be presented. Kinetica 2003 allows the user to input a se ries of single step reactions with their rate constants to calculate the concentrations of the species with time. It should be noted that only two reactants can combine in a single step, therefore, if three species must react, an intermediate species must be formed firs t and that intermediate then reacts with the third reactant. An exampl e will be given shortly. Initially, since many of the dependencies on the reactants were observed to be the same for the hydrogen peroxide decompositi on and nucleophilic alkene epoxidation, it was thought that these reactions were pro ceeding through the same intermediates, both “A” and “B.” First, the common intermediate “A” must be generated. It is assumed that the generation of this intermediate is fast, since the order of a ddition in the hydrogen peroxide decomposition studies show that the order of addition of hydrogen peroxide, bicarbonate, or manganese(II), has no eff ect on the observed decomposition rate. The lack of a mixed order dependence on any of these species also indicates that the generation of the catalyst is rapid. We propose that hydrogen per oxide and bicarbonate equilibrate to form peroxycarbonate (K1) and that peroxycarbonate combines with bicarbonate and manganese(II) to form the activ e catalyst. This would seem reasonable given the kinetic data suggests a rate = k[Mn(II)][HCO3 -]2[H2O2]x, for the hydrogen peroxide decomposition.

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95 It was thought that the generation of the active manganese compound, “A”, could be generated by the simple equilibration of Mn(II), bicarbonate, and hydrogen peroxide, as seen by Equation (3-15)-(3-17). HCO3 + H2O2 HCO4 + H2O K1 = 0.32 M-1 (3-15) Mn(II) + HCO3 [MnII(HCO3)]+ K2 = 19.05 M-1 (3-16) [MnII(HCO3)]+ + HCO4 A K3 (3-17) Only two of the equilibri um constants are known from literature data. The K1 for the peroxycarbonate formation is 0.32 M-1,32 and the K2 for the generation of the [Mn(HCO3)]+ ion-pair is 19.05 M-1.107 More than likely, manganese(II) is acting as a Lewis acid to rapidly establish the peroxycarbonate equilibrium. Experiments attempting to observe this phenomenon with manganese(II) proved unsuccessful. The paramagnetic nature of manganese(II) made acquiring 13C NMR spectra difficult. Experiments by Yao34 and Albert108 show that the addition of zinc complexes, which act as carbonic anhydrase mimics, do rapidly equilibrate the peroxycarbonate. Therefore, the fo rward and reverse reactions for the production of peroxycarbonate have been accelerated while keeping the overall equilibrium constant the same. The Kinetica 2003 equations used are shown be low, as they appear in the program. The rate constant for a single step is listed after the equation. As noted earlier, only two reactants can be used in any given step, and equilibrium expr essions must be written as two reactions, one for the forw ard reaction and one for the re verse. So, since there are three equilibria in the predicted generation of the active manganese species, six reactions are required in Kinetica to simulate the concen trations over time. Gi ven that the order of

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96 addition of hydrogen peroxide and manganese (II) made no significant impact of the reaction rate in the decomposition data, a rapid equilibrium step is used for the reaction of [Mn(HCO3)]+ + HCO4 -, k = 1 x 107, for both the forward and re verse steps. “A” in the following kinetic steps is the active manganese complex. 1; H2O2 + HCO3 HCO4; 3.800E-02 (3-18) 2; HCO4 HCO3 + H2O2; 1.200E-01 (3-19) 3; Mn(II) + HCO3 Mn(II)HCO3; 1.900E07 (3-20) 4; Mn(II)HCO3 Mn(II) + HCO3; 1.000E06 (3-21) 5; Mn(II)HCO3 + HCO4 A; 1.000E07 (3-22) 6; A Mn(II)HCO3 + HCO4; 1.000E07 (3-23) When a simulation of the above mechanism for the formation of the active manganese catalyst was performed using 1.00 M hydrogen peroxide, 4 M Mn(II), and varying bicarbonate, a secondorder relationship was not f ound for the production of the active catalyst, Figure 3-43, which is th e observed dependence for both the hydrogen peroxide decomposition and nucle ophilic alkene epoxidation. 0.00E+00 1.00E-07 2.00E-07 3.00E-07 4.00E-07 5.00E-07 6.00E-07 7.00E-07 8.00E-07 00.20.40.60.81 [HCO3 -], MConcentration A, M Figure 3-43. Simulation of the dependence on the concentration of the active catalyst with varying bicarbonate. Simulati on conditions: 1.00 M hydrogen peroxide and 4 M Mn(II). y = ((6.55 0.37)x10-7)x, error reported to the 95% confidence.

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97 The predicted second-order dependence on bicarbonate is not observed from these simulations. The reason for the lack of sec ond-order behavior is due to the equilibrium constant for the [Mn(II)(HCO3)]+ ion-pair. Since the equilibri um constant is large, the concentration of [Mn(HCO3)]+ quickly saturates, as seen in Figure 3-44. 0.00E+00 5.00E-07 1.00E-06 1.50E-06 2.00E-06 2.50E-06 3.00E-06 3.50E-06 00.10.20.30.4 [HCO3 -], M[Mn(HCO3)]+, M Figure 3-44. Plot of simu lation results for [Mn(HCO3)]+ versus [HCO3 -]. The [Mn(HCO3)]+ quickly saturates due to the large equilibrium constant of 19.05. Simulation conditions: 1.00 M H2O2, 4 M Mn(II). In order to counteract the saturation eff ect seen in the equi librium of [Mn(HCO3)]+, a third bicarbonate anion is re quired in the generation of th e active catalyst. The steps now required for the generation of th e active catalyst are shown below. 1; H2O2 + HCO3 HCO4; 3.800E-02; (3-24) 2; HCO4 HCO3 + H2O2; 1.200E-01; (3-25) 3; Mn(II) + HCO3 Mn(II)HCO3; 1.900E07; (3-26) 4; Mn(II)HCO3 Mn(II) + HCO3; 1.000E06; (3-27) 5; Mn(II)HCO3 + HCO4 [Mn(II)(HCO3)(HCO4); 1.000E07; (3-28) 6; [Mn(II)(HCO3)(HCO4) Mn(II)HCO3 + HCO4; 1.000E07; (3-29) 7; [Mn(II)(HCO3)(HCO4) + HCO3 A; 1.000E07; (3-30)

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98 8; A [Mn(II)(HCO3)(HCO4) + HCO3; 1.000E07; (3-31) When a simulation is performed with va rying bicarbonate con centration in the presence of 1.00 M hydrogen peroxide and 4 M Mn(II), the graph presented in Figure 345 is obtained, for which a second-order de pendence on bicarbonate is now observed. 0.00E+00 1.00E-07 2.00E-07 3.00E-07 4.00E-07 5.00E-07 00.20.40.60.81 [HCO3 -]2, M2Concentration A, M Figure 3-45. Simulation of the dependence on the concentration of the active catalyst with varying [HCO3]. Simulation Conditions: 1.00 M hydrogen peroxide and 4 M Mn(II). y = ((4.60 0.22)x10-7)x, error reported to the 95% confidence. Originally, before the discovery that car bonate radicals might be involved in the reaction, the nucleophilic alke ne epoxidation studies and proposed rate equations (Equations 3-32 to 3-34) of Bennett46 were used to determine the equilibrium constant for the generation of “B”. The peroxide d ecomposition was ignored since it is slow compared to the epoxidation rate under the studied conditions. H2O2 + HCO3 + Mn(II) ”A” K1 (3-32) “A” + H2O2 ”B” K2 (3-33) “A” + HCO3 + S SO + Mn(II) + 2HCO3 k3 (3-34) Equations (3-35)-(3-45) are th e kinetic steps that have been found which describe the nucleophilic alkene epoxidation.

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99 1; H2O2 + HCO3 HCO4; 3.800E-01 (3-35) 2; HCO4 HCO3 + H2O2; 1.200E00 (3-36) 3; Mn(II) + HCO3 Mn(II)HCO3; 1.900E07 (3-37) 4; Mn(II)HCO3 Mn(II) + HCO3; 1.000E06 (3-38) 5; Mn(II)HCO3 + HCO4 Mn(II)(HCO3)(HCO4); 1.000E07 (3-39) 6; Mn(II)(HCO3)(HCO4) Mn(II)HCO3 + HCO4; 1.000E07 (3-40) 7; Mn(II)(HCO3)(HCO4) + HCO3 A; 1.000E07 (3-41) 8; A Mn(II)(HCO3)(HCO4) + HCO3; 1.000E07 (3-42) 9; A + H2O2 B; 1.100E07 (3-43) 10; B A + H2O2; 1.000E05 (3-44) 11; S + A SO + Mn(II)HCO3 + HCO3 + HCO3; 1.350E05 (3-45) For these kinetic steps, th e hydrogen peroxide dependence was used, initially, since the turnover in the hyd rogen peroxide dependence is only observed in the nucleophilic alkene epoxidation. The fit of the numerical simulations to the hydr ogen peroxide data is shown in Figure 3-46. 0 0.0005 0.001 0.0015 0.002 0.000.200.400.600.801.00 [H2O2], Mkobs, s-1 Figure 3-46. The generated curve for the hydrogen peroxide dependence on nucleophilic alkene oxidation. The po ints represent the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each H2O2 concentration. Reaction a nd simulation conditions: 0.5 M Mn(II), 1.00 M bicarbonate, 0.001 M Sulfonated Styrene (SS).

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100 Once the rate constants were determined to allo w good fits to the observed data, the other dependencies needed to be check to assu re that the kinetic st eps and rate constants will predict the observed trends. Plots of simulations where the rate constants are varied from those above can be found in the Appe ndix. The bicarbonate and manganese dependencies are shown in Figures 3-47 and 3-48, respectively. 0 0.0005 0.001 0.0015 0.002 0.0025 0.003 00.511.5 [HCO3]2, M2kobs, s-1 Figure 3-47. The generated curve for the bi carbonate dependence on nucleophilic alkene oxidation. The points represent the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each [HCO3] concentration. Reaction and Simulation Conditions: 0.5 M Mn(II), 0.10 M hydrogen peroxide, 0.001 M Sulfonated Styrene (SS). 0 0.002 0.004 0.006 0.008 0.01 0.012 0246 [Mn(II)], Mkobs, s-1 Figure 3-48. The generated curve for the ma nganese dependence on nucleophilic alkene oxidation. The points represent the observed kobs experimentally determined by Bennett.46 The line is the simulated kobs at each [Mn(II)] concentration. Reaction and simulation conditions: 1. 00 M bicarbonate, 0.55 M hydrogen peroxide, 0.001 M Sulfonated Styrene.

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101 Once the nucleophilic alkene epoxidation kinetics in the absence of hydrogen peroxide decomposition were complete, simulations attempting to model the mechanism of the hydrogen peroxide decomposition were conducted. As stated earlier, our assumption was that the inte rmediates of the nucleophilic alkene epoxidation would be the same for the hydrogen peroxide decomposition. The simplest decomposition pathway for the decomposition of hydrogen peroxide wo uld be for “B” to decompose to release oxygen and reduce Mn(IV) to Mn(II), as seen in Figure 3-41. The following set of kinetic steps (Equations 3-46 to 3-57) were then used to model this reaction. The alkene concentration (S in Equation 357) is set to 0, since there is no alkene epoxidation occurring in these reactions. 1; H2O2 + HCO3 HCO4; 3.800E-01 (3-46) 2; HCO4 HCO3 + H2O2; 1.200E00 (3-47) 3; Mn(II) + HCO3 Mn(II)HCO3; 1.900E07 (3-48) 4; Mn(II)HCO3 Mn(II) + HCO3; 1.000E06 (3-49) 5; Mn(II)HCO3 + HCO4 Mn(II)(HCO3)(HCO4); 1.000E07 (3-50) 6; Mn(II)(HCO3)(HCO4) Mn(II)HCO3 + HCO4; 1.000E07 (3-51) 7; Mn(II)(HCO3)(HCO4) + HCO3 A; 1.000E07 (3-52) 8; A Mn(II)(HCO3)(HCO4) + HCO3; 1.000E07 (3-53) 9; A + H2O2 B; 1.100E07 (3-54) 10; B A + H2O2; 1.000E05 (3-55) 12; B Mn(II)HCO3 + O2 + HCO3 + HCO3; 1.000E02 (3-56) 13; S + A SO + Mn(II)HCO3 + HCO3 + HCO3; 1.350E05 (3-57)

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102 Unfortunately, the hydrogen peroxide decay curves, for which an example is shown in Figure 3-49, never adequately predicted the observed hydrogen peroxide decay. -5 -4.5 -4 -3.5 -3 -2.5 -2 0100200300400500600700800 Time, secln([H2O2]) Figure 3-49. A typical nume rical simulation plot atte mpting to model the hydrogen peroxide decay curves. Po ints represent observed ln[H2O2] versus time, while the line is the simulated ln([H2O2]) versus time. Reaction and simulation conditions: 0.10 M H2O2, 0.30 M HCO3 -, 4.0 M Mn(II). Since the plots do not adequately predic t the hydrogen peroxide decomposition, a new model was proposed based on Fenton type reactions of metal cations with hydrogen peroxide. This mechanism will be presented shortly. The reaction of the carbonate radical alone does not e xplain the hydrogen peroxide dependence observed in the epoxi dation kinetics. More than likely, an unreactive, or slightly reactive, hydrogen per oxide intermediate (intermediat e “B”) in the presence of a radical species is highly unlikely, therefore, the existence of the intermediate “B” has been excluded at this time. Since both th e carbonate radical anion and the manganese oxo complex have been proposed to be the s ource of oxygen in the alkene epoxidation, the rate of alkene epoxidation will be the su m of the reactions of the alkene with both species. The alkene epoxidation was not at tempted yet in the mechanism in which the carbonate radical is generated.

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103 While Fenton chemistry of metal cations usually generates hydroxyl radicals, no hydroxyl radicals will be produced in this model. Since the concentration of bicarbonate is so much higher than the metal cation, all r eactions that produ ce the hydroxyl and perhydroxyl radicals are replaced with the reaction of a bicarb onate ion to generate water, or hydrogen peroxide, and a carbona te radical anion. While th is radical is known to react with redox active metals, little is know about its reactivity with organic substrates, specifically what the rate constants for reactio ns are. For all reac tions shown, the rate constants used for the reaction of the car bonate radical were taken from the work Mazellier et al.98 Rate constants for the reactions of Mn(II) and Mn(III) with superoxide were used from the work by Pick-Kaplan and Rabani.109 For now, current information av ailable in the literature about the reactivity of the carbonate radical with organic substrates is limited. Future attempts to model the nucleophilic alkene epoxidati on may become possible when new information concerning this radical species becomes available. In th e mean time, kinetic data for the reactivity of the carbonate radical with hydrogen peroxide, and other ac tive oxygen species including the superoxide radical anion, the hydroxyl radical, and the pe rhydroxyl radical is currently available in the literature.98,109 Attempts to model the hydrogen pe roxide decomposition based on Fenton chemistry have continued. Using Sychev’s model of hydrogen peroxide decomposition, Equations 3-4 to 3-14, the hydroxyl radica ls produced were replaced by carbonate radicals by reaction with bicarbonate. Equati ons 3-58 to 3-67 are the reactions generated to describe the decomposition of hydrogen peroxide.

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104 H2O2 H+ + HOO(3-58) H2CO3 H+ + HCO3 (3-59) H2CO3 H2O + CO2 (3-60) HCO3 CO3 2+ H+ (3-61) HCO3 + H2O2 HCO4 + H2O (3-62) 2Mn2+ + 2HCO4 2Mn3+ + 2 CO3 -• + 2OH(3-63) Mn2+ + CO3 -• Mn3+ + CO3 2(3-64) Mn3+ + HOO+ HCO3 Mn2+ + H2O2 + CO3-• (3-65) 2Mn3+ + 2CO3 -• 2Mn2+ + O2 + 2CO2 (3-66) CO2(aq) CO2(g) (3-67) Simulations using these reactions were pe rformed in an attemp t to simulate the hydrogen peroxide decomposition. Unlike the simulations using the intermediates “A” and “B”, hydrogen peroxide decay curves, such as the one presented in Figure 3-50, can now be generated. Equation 3-67 must be added since the cataly st lifetime studies proved that bicarbonate is being lost. Equations 3-68 to 3-93 are the equa tions used in Kinetica with the rate constants that were us ed. Rate constants fo r each of the reactions involving protons have been adjusted for pH 8.3. Rate constants for the reactions are from the following sources: Fridovich,110 estimates of Mn(II) and Mn(III) using Mn(II) and Mn(III)porphyrin complex with carbonate radical anion, Mazellier,98 reactions of carbonate radicals with hydrogen pe roxide, Palmer and van Eldik111 and Roughton, 112 carbon dioxide hydration, dehydration, pKa H2CO3, pKa bicarbonate. All other rate

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105 constants have been adjusted to fit the hydrogen peroxide decay curves observed by experiment. 1; H2O2 HOO-; 2.000E-02 (3-68) 2; HOOH2O2; 5.000E01 (3-69) 3; CO2 H2CO3; 3.000E-02 (3-70) 4; H2CO3 CO2; 2.000E01 (3-71) 5; H2CO3 HCO3; 4.000E07 (3-72) 6; HCO3 H2CO3; 7.000E2 (3-73) 7; CO2 HCO3; 1.700E-02 (3-74) 8; HCO3 CO2; 2.000E-05 (3-75) 9; HCO3 CO32-; 3.160E-01 (3-76) 10; CO32HCO3; 5.000E01 (3-77) 11; H2O2 + HCO3 HCO4; 3.800E-01 (3-78) 12; HCO4 H2O2 + HCO3; 1.200E00; (3-79) 13; Mn(II) + HCO3 Mn(II)HCO3; 1.900E07; (3-80) 14; Mn(II)HCO3 Mn(II) + HCO3; 1.000E06; (3-81) 15; Mn(II)HCO3 + HCO3 Mn(II)(HCO3)2; 1.000E07; (3-82) 16; Mn(II)(HCO3)2 Mn(II)HCO3 + HCO3; 1.000E07; (3-83) 17; Mn(III) + HCO3 Mn(III)HCO3; 1.900E07; (3-84) 18; Mn(III)HCO3 Mn(III) + HCO3; 1.000E07; (3-85) 19; Mn(III)HCO3 + HCO3 Mn(III)(HCO3)2 ; 1.000E07; (3-86) 20; Mn(III)(HCO3)2 Mn(III)HCO3 + HCO3; 1.000E07; (3-87) 21; Mn(II)(HCO3)2 + CO3RadAn Mn(III)(HCO3)2 + CO32-; 5.000E07; (3-88)

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106 22; Mn(III)(HCO3)2 + CO3RadAn Mn(IV)CO3(HCO3)2 ; 1.200E07; (3-89) 23; Mn(IV)CO3(HCO3)2+Mn(IV)CO3(HCO3)2 Mn(II)(HCO3)2+Mn(II)(HCO3)2 + O2 + CO2 + CO2; 1.000E07; (3-90) 24; Mn(II)(HCO3)2+HCO4 Mn(III)(HCO3)2 + CO3RadAn ; 4.500E06; (3-91) 25; Mn(III)(HCO3)2 + HOOMn(II)HCO3 + CO3RadAn + H2O2; 4.500E06; (3-92) 26; CO2 CO2g ; 2.500E-03; (3-93) -4.5 -4 -3.5 -3 -2.5 -2 02000400060008000 Time, secln([H2O2]) Experimental Simulation Figure 3-50. Simulation of hydroge n peroxide decay at lower bicarbonate concentration. Points represent data, while the line is the simulated decay. Reaction and simulation conditions: 0.10 M H2O2, 0.10 M HCO3, 4 M Mn(II). From the data presented in Figure 3-46 it appears that th e carbonate radical mechanism accurately predicts the hydrogen pe roxide decomposition in the presence of bicarbonate and manganese(II). Figures 3-51, 3-52, and 3-53 repres ent the bicarbonate, manganese(II), and hydrogen peroxide dependencies for this mechanism.

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107 0 0.001 0.002 0.003 0.004 0.005 0.006 00.050.10.150.20.250.3 [HCO3]2, M2" kobs", s-1 Figure 3-51. A plot of [HCO3]2 versus “kobs” for the hydrogen peroxide decomposition. Reaction and simulation conditions: 0.10 M H2O2, 4 M Mn(II). 0 0.001 0.002 0.003 0.004 0.005 012345 [Mn(II)], M" kobs", s-1 Figure 3-51. A plot of “kobs” versus [Mn(II)] for the hydrogen peroxide decomposition. Reaction and simulation conditions: 0.10 M H2O2, 0.40 M HCO3 0.00E+00 1.00E-04 2.00E-04 3.00E-04 4.00E-04 00.10.20.30.40.5 [H2O2], M" kobs", s-1

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108 Figure 3-53. Simulated hydrogen peroxide dependence for the hydrogen peroxide decomposition reactions. Simulati on conditions: 0.90 M HCO3, 0.5 M Mn(II). The hydrogen peroxide plot show n in Figure 3-53 is an in teresting result from the simulation of this mechanism. Currently, the hydrogen peroxide dependence has never been adequately measured for the hydrogen peroxide decomposition. Another simulation was done to assure that this result occurs at different bicarbonate and manganese(II) concentrations. Figure 3-54 is a simu lation using 0.4 M bi carbonate and 3.0 M Mn(II). This result indicates that the dependence is not only simulated at low metal and high bicarbonate concentration. 0.00E+00 1.00E-03 2.00E-03 3.00E-03 4.00E-03 5.00E-03 6.00E-03 00.10.20.30.40.50.6 [H2O2], M" kobs", s-1 Figure 3-54. Simulated hydrogen peroxide dependence for the hydrogen peroxide decomposition reactions. Simulati on conditions: 0.40 M HCO3, 3.0 M Mn(II). As stated earlier, ln plots for most of the hydrogen peroxide decomposition studies accelerated near the end of the reaction. The hydrogen peroxide dependence explains why this was true. As the decomposition o ccurred, the hydrogen peroxide concentration falls, and in doing so, the rate of the reaction will increase, as seen in Figure 3-55. For example, if a ln plot of one of the higher bicarbonate and hydrogen peroxide reactions is simulated, the accelerati on can be easily seen.

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109 -4.5 -4 -3.5 -3 -2.5 -2 02004006008001000 Time, secln([H2O2]) Figure 3-55. Plot of ln([H2O2]) versus time. The points represent observed data and the line is the simulation. Reaction and simulation conditions:0.10 M H2O2, 0.40 M HCO3, 3.0 M Mn(II). As the plot indicates the reaction accelerates as the decomposition occurs. Conclusions From the results presented, it can be concluded that the mechanism of hydrogen peroxide decomposition and nucleophilic alke ne epoxidation must occur with similar intermediates. The kinetic experiments i ndicate that the dependence of manganese(II) on both reactions is first-order. This indicates that only one meta l ion is present in the active catalyst. The bicarbonate dependence for both reactions is second-order. It would be reasonable to assume that one of the bicarbonat e ions is in the form of peroxycarbonate, while the other is simply a bound bicarbonate or carbonate ion. The nucleophilic alkene epoxidation indicates that at higher hydrogen peroxide concen tration, the reaction rate begins to fall. This could be explained by the hydrogen peroxide dependence that has been simulated, especially if the active oxida nt of the alkene is the carbonate radical anion. The lifetime of the active manganese catal yst has also been examined. From studies conducted where sequent ial additions of hydrogen per oxide are added to the same

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110 solution, the observed rate of hydrogen pe roxide decomposition decreases with each addition. This decrease in activity is due, in part, to the loss of bicarbonate from the solution. When a larger scale reaction was cycled 10 times, the final bicarbonate concentration was measured us ing the standard barium chlo ride method. It was found that the concentration of bicarbon ate had fallen to about half of its original concentration. In addition, experiments using the malachite green assay for phos phate indicted that phosphate was present in the stock hydrogen peroxide. With each addition of hydrogen peroxide, phosphate was being added to the so lution which causes the precipitation of the manganese catalyst. When a second study was conducted usi ng distilled hydrogen peroxide and by adding sma ll amounts of bicarbonate, to stabilize the bicarbonate concentration, the loss of activity was much less than when no bicarbonate was added and the hydrogen peroxide was not distilled. The loss in activity is attributed to the inability to adequately cont rol the bicarbonate concentra tion and from normal dilution effects. The source of manganese for the mangane se dependent decomposition of hydrogen peroxide was also examined. A series of experiments were conducted using three manganese sources. First, manganese(II) sulfate was used. All of the kinetic experiments were conducted using this manga nese source. Second, permanganate was used as the manganese for a series of experiments. Results from these experiments indicate that there is no significant change in the rate in using permanganate as the manganese source. Observed rate constant s for the same concentration of manganese were nearly identical for r eactions using mangane se(II) sulfate and permanganate. The third manganese source was a Mn(IV) complex that has been reported to be an excellent

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111 catalyst for the epoxidation of nucleophilic al kenes by hydrogen peroxide in the presence of bicarbonate. Experiments us ing this catalyst indicate that the catalyst is quickly destroyed upon addition to solutions of bi carbonate and hydrogen peroxide. The most likely decay mechanism is through the N-dealkylation of the triazacyclononane ligand. When equal concentrations of the Mn(IV) catalyst were used to decompose hydrogen peroxide, the observed rate constants were again within error for those found with manganese(II) sulfate. The cis/trans isomerization in pure water has been studied. Reactions of maleic acid, a cis alkene, with hydrogen peroxide a nd manganese in the pr esence of bicarbonate yield the cis and trans epoxide s. This result indicates that during the epoxidation reaction, the C-C sigma bond of the alkene is free to rotate to the more stable trans isomer before the epoxide ring is formed. Base d on the proposed mechan ism, the cis/trans isomerization is more than likely due to the generation of either a carbocation or carbon radical intermediate. If either intermedia te has a long enough life time, rotation along the C-C sigma bond will allow for the generation of the trans epoxide. A reexamination of the oxidation of DM NA has been investigated. Oxidation studies of DMNA and DENA by hyd rogen peroxide and mangan ese in the presence of bicarbonate indicate that oxidative N-dealkylation is occurring, which was not accounted for in Sychev’s proposed mechanism of hydrogen peroxide decomposition in the presence of bicarbonate and manganese. Sy chev’s mechanism, therefore, cannot be completely correct. The loss of O2 (g) production by the addition of DMNA can be supported using a single elec tron transfer mechanism based on work exploring the

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112 reactivity of cytochrome P-450, at least fo r the products that have currently been detected. Solvent isotope effect studies have b een conducted on the manganese catalyzed decomposition of hydrogen peroxide in the pr esence of bicarbonate. A normal isotope effect is observed for these reactions. This is not surprising since th e cleavage of the O-H bond should be faster than that for O-D. The solvent isotope effect has also been examined for the nucleophilic alkene epoxidation. A large, inverse isotope effect is observed for these reactions. The explanati on for the observed isotope effect has been rationalized using the carbonate radical an ion as the oxidizing species. A proposed mechanism for the deco mposition of hydrogen peroxide and nucleophilic alkene epoxidation ha s been presented. The propos ed catalytic species is the product of the reaction of bicarbonate, per oxycarbonate, and manganese. The proposed catalyst can react through a number of path ways to decompose peroxide, epoxidize nucleophilic alkenes, or N-dealkylate aromatic amines. Numerical simulation supports the carbonate radical as the active speci es involved in the hydrogen peroxide decomposition. Future investigations of the hydrogen peroxide decomposition catalyzed by manganese(II) may prove difficult. The formed active metal species is in very low concentration and is formed in situ. If the Mn(IV)-TACN co mplex is any indication, other methods of trying to stabilize the metal center without the use of nitrogen containing ligands will be required due to the decomposition of the ligand.

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113 Experimental Materials and Instrumentation Sodium bicarbonate, sodium acetate, styrene, N,N-dimethyl-4-nitrosoaniline, N,Ndiethyl-4-nitrosoaniline, and manganese(II) sulfate were all analytical grades and obtained from Fisher (Atlan ta, GA). Hydrogen peroxide (35 and 50%) was obtained from Fisher (Atlanta, GA) and standardized often by iodometric ti tration. Water was purified using a Barnstead E-Pure 3-Module De ionization System. Extraneous metal ions from salt solutions were removed by passi ng through a Chelex 100 resin obtained from Aldrich (St. Louis, MO). Sodium bicarbona te solutions were standardized using the method below before use to assure concentration. Proton and 13C NMR spectra were obtained on a Mercury 300 spectrophotometer. The residual solvent peaks were used as an internal standard except deuterated chloroform which used 0.05% TMS as internal standard. All deuterated solvents were obtained from Cambridge Isotope Laboratory, Inc (Andover, MA). UV-Vis kinetic experime nts were obtained usi ng a Hewlett-Packer 8453 spectrophotometer using 1.0 cm quartz cells from Starna Cells, Inc. Temperature was maintained at 25 0.1 C using a Fish er Isotemp 1600S wate r bath circulator. Gas Chromatography experiments were performed using a Varian CP-3800 Gas Chromatograph equipped with a J. W. Scie ntific DB-35MS column. The method for DMNA oxidation analysis consisted of a linear gr adient of 10C/min from 80C to 200C. The temperature was then maintained at 200C for 15 minutes providing for a total analysis time of 27 minutes. The method for acetaldehyde analysis consisted of an isothermal method of 40C for 5 minutes.

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114 Styrene oxidation reac tions were analyzed by High Performance Liquid Chromatography using a Rainin HPLX solven t delivery system on a C-18 reverse phase column. The method consisted of a non-linear gradient of H2O:CH3CN from 75:25 – 5:95 over a 15 minute period. Pr oduct was detected at 221nm. Sulfonated styrene oxidations were analyzed by High Performance Liquid Chromatography using a Varian Prostar syst em. Analysis was performed using a C18 reverse phase column using tetrabutylammonium chloride as an ion-pairing reagent. The method was isocratic eluti on using 80% A: 80:20 CH3CN:H2O B: 0.1 mM in 80:20 CH3CN:H2O Standardization of sodium bicarbonate solutions Solutions of sodium bicarbonate were standardized before each kinetic experiment to assure the concentration eluting from the Chelex 100 column. All solutions were delivered usi ng volumetric pipets. A 10mL aliquot of sodium bicarbonate solution exiting the Chelex 100 column was a pl aced in a clean, dry beaker. An excess amount of sodium hydroxide solution of a known concentration, by titration with potassium hydrogen phthalate, was added to the beaker. The solution was stirred to allow for complete deprotonation of the bicarbonate to form the carbonate dianion. An excess barium chloride solution is then added to precipitate a ll of the carbonate dianion as barium carbonate. Phenolphthalein is then a dded to the mixture to give a pink color due to the residual hydroxide ion. The mixture is titrated us ing a known concentration of hydrochloric acid until th e solution just turns clear. The number of moles of hydrochloric acid added is equal to the excess moles of sodium hydroxide. The difference between the number of moles from the acid titration a nd the number of mo les of hydroxide ion

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115 initially added equals the number of moles of bicarbonate present in the initial 10mL aliquot (Equation 3-10). #moles OHinit-#moles OHexcess = #moles bicarbonate (3-10) The molarity of the solution can then be determined. Hydrogen peroxide decomposition studies The H2O2 decomposition studies were carried out in sodium acetate buffered solutions at a pH 8.4. All salt solutions had been passed through a Chelex 100 column to remove any extraneous metal ions. Hydroge n peroxide/bicarbonate/ buffer solutions were allowed to equilibrate for at least 15 minutes to allow for formation of peroxycarbonate. Kinetic experiments were follo wed for at least 2.5 halflives following the decreasing absorbance at 263 nm. Solutions of mangane se(II) sulfate were always added last to initiate the decomposition of the hydrogen pero xide. Hydrogen peroxi de was held at 0.1 M. The sodium bicarbonate concentration was in the range from 0.0 to 0.55 M. Manganese concentrations we re in the range of 0-15 M. Ionic strength was held constant at 1 M using sodium acetate, the buffering solution. Synthesis of [MnIV(Me3TACN)(OMe)3](PF6) The synthesis is described in the literature by Kerschner.63 To a solution of 1,4,7trimethy-1,4,7-triazacyclononane (0.1 g, 578 mmol) in 8 mL of methanol is added MnCl2 (0.074 g, .578 mmol) predissolved in 2 mL methanol. The solution will turn dark brown in color. The solution is stirred and cooled on ice to 0 C, and Na2O2 (0.046 g, 0.578 mmol) was slowly added to the manganese /TACN solution. After 1 h of stirring at 0 C, the solution was warmed to room temperat ure and stirred a further 1-2 h. Finally, NaPF6 (0.100 g, 5.90 mmol) was added to the so lution. The mixture was filtered through a porous glass frit and the soluti on was neutralized with dilute sulfuric acid. Then 5 mL

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116 of water was added to the neutralized solution, and it was filtered again. The filtrate was concentrated to one-third th e original volume under reduced pressure, and the resulting solution was chilled at 0 C to precipitate the desi red product. UV-Vis (CH3CN): max nm ( M-1cm-1) 326 (13200), 287 (12900), 228 (11500). Mp:160-164 C. Oxidation of N N -dimethyl-4-nitrosoaniline (DMNA) by Oxone To a stirred solution of DMNA (1.0 g, 6.7 mmol) in 30 mL 50:50 CH3CN: H2O (v:v) is slowly added a solution of Oxone (4.0 g, 6.5 mmol) in 20 mL H2O over 30 minutes. During the course of the additi on the starting green color is replaced by a yellow color and a yellow precipi tate begins to float on top of the solution. Addition of a further 20 mL CH3CN after the reaction is over allo ws for the yellow precipitate to dissolve. The entire mixture is placed in a separatory funnel and extracted 3 x 100 mL with CHCl3. The yellow color was transferred to the organic layer. The combined organic washings were dried over anhydrous ma gnesium sulfate. After filtering, the organic solvent was removed under reduced pressure to afford pure N,N-dimethyl-4nitroaniline. 1H NMR (CD3CN): 8.06 (d, J = 9.3Hz, 2H), 6. 69 (d, J = 9.6Hz, 2H), 3.07 (s, 3H). Oxidation of N N -Dimethyl-4-nitrosoaniline (DMNA) by H2O2/HCO3 -/Mn2+ A typical reaction involves stirring a solution of DMNA (0.5 g, 3.3 mmol) with manganese(II) sulfate in a solution of 30:70 CH3CN: H2O (v:v). To this is added an equilibrated solution of H2O2 and sodium bicarbonate slow ly over about 30 minutes. Multiple experiments were performed with H2O2 equal to 1x, 10x, and 50x molar excess over DMNA. Enough sodium bicarbonate and ma nganese(II) sulfate were added to make their final concentrations equal to 0.40 M and 4.0 M, respectively. The results of these experiments were discussed earlier.

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117 Oxidation of N N -diethyl-4-nitrosoaniline (DENA) by H2O2/HCO3 -/Mn2+ To a solution of DENA (1 .00 g, 5.61 mmol) in 70:30 H2O:CH3CN (v:v) is added a solution of manganese(II) sulfate. To th is is slowly added, over 30 minutes, an equilibrated solution of H2O2 and bicarbonate, where the H2O2 concentration is 50x molar excess to DENA. The final bicarbonate and manganese(II) c oncentrations were 0.40 M and 4.0 M, respectively. The acetaldehyde produced during the reaction was trapped by condensation of the gas in an ice cooled trap.

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118 CHAPTER 4 ELECTROPHILIC ALKENE EPOXIDAT ION BY THE PEROXYCARBONATE DIANION Introduction The transfer of an oxygen atom is one of the most important biological transformations and is quite useful in synthetic organic chemistry.113 Generally, oxidation reactions are catalyzed by metals,114 as seen in Chapters 2 and 3 with the manganese(II) catalyzed epoxidation of nucleophi lic alkenes, or main group catalysts.115 Due to the important nature of this oxidative process many catalytically-active transition metals have been extensively st udied, including high-valent d0 metals, such as Mo(VI), Re(VII), Rh(VII), V(V) and Ti (IV), which catalyze the epoxidation of alkenes116-118 by peroxides. For example, hydrogen per oxide reacts with methyltrioxorhenium119,120 generating mono and bis-peroxides.121,122 Both species are efficient oxygen donors, transferring oxygen to various nucleophiles123 allowing significant rate enhancements compared to reactions involvi ng hydrogen peroxide alone. Another example includes the oxygen transfer from the oxo-ligands of oxometalloporphyrin systems, such as MoVI(TPP)(O)2 and FeIV(TPP)(O), which are report ed in the literature.124,125 Some limitations of metal activators include: the toxic nature of the metals, separation of products from reaction mixtures, and th e low solubility of metal catalysts. Organic peracids, such as peracetic acid, have been widely used to oxidize alkenes. One of the most common oxidation reactions involving peracids is the Prileschajew epoxidation.126 The peracid epoxidation of alkene s is influenced by several factors

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119 including: the nature of the carbon-carbon double bond, peraci d substituents, and solvent effects (such as intermolecular H-bonds).127 In 2000, it was reported that peroxycar bonate was an active oxidant of both nucleophilic and electrophilic alkenes.33 Chapters 2 and 3 have focused on the use of manganese(II), which was discovered as a trace contaminant in the bicarbonate salts, as an activator of peroxycarbonate for use in the oxidation of nucle ophilic alkenes and the disproportionation of hydroge n peroxide. These data do not explain how the manganese(II) system also has the ability to oxi dize electrophilic alkenes. The lack of information currently availa ble on the use of peroxycarbonate as an oxidant of electrophilic alkenes requires further investigation. Results and Discussion The oxidation of electrophilic alkenes to their corresponding epoxides was first reported by Weitz and Scheffer128 in 1921 and can be easily ac hieved using a number of different nucleophilic oxid ants including hypochlorite, -OCl, and hydroperoxide,-OOH. Numerous examples of nucleophi lic oxidation of electrophilic alkenes can be found in the literature,129-139 but only a couple of these focu s on the kinetics of the epoxidation.131,132 Figure 4-1 shows the resonance structur e of a common electrophilic alkene, an unsaturated ketone, which provi des a rationalization as to w hy electrophilic alkenes react with nucleophilic oxidants. As can be seen in the figure, the -carbon of the resonance structure has a positive charge and will be the carbon attacked by a nucleophilic oxidant. Figure 4-2, illustrating a r eaction of the hydroperoxide anion, shows the general mechanism by which nucleophilic oxidants can oxidize electr ophilic substrates.

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120 OO Figure 4-1. The resonan ce structure of an electrophilic alkene, an -unsaturated ketone, explains the reactivity with nucleophilic oxidants. The -carbon of the alkene, as seen in the resonance structure, is more electropositive and will be the site of attack by a nucleophilic oxidant. R R' R'' O R R' R'' O OOH O + HO-R'' R' O R HOOFigure 4-2. The mechanism of electrophilic alkene oxidation by th e hydroperoxide anion is illustrated. The nucleophilic oxidant adds at the electrophilic carbon, the carbon. Reformation of the ketone moie ty causes either th e displacement of the hydroperoxide anion, regenerating the starting alkene and hydroperoxide, or ring closure to form the epoxide and the hydroxide anion. Peroxycarbonate can exist as both the anion, HCO4 -, and dianion, CO4 2-, in aqueous solution. While the peroxy carbonate anion is an electr ophilic oxidant as shown by sulfide oxidation,41 the peroxycarbonate dianion will behave as a nucleophilic oxidant in a similar way as the hydroperoxide anion. Figure 4-3 illustrates how the peroxycarbonate dianion will react in a similar way in the oxidation of electrophilic alkenes. The only difference between the hydroperoxide dianion and the peroxycarbonate dianion is that the nucleofuge (leaving group) of the peroxycar bonate dianion is the carbonate dianion. R O R O + CO3 2-O O O O R' R'' OR' OCO3 -R'' R'' O R R' Figure 4-3. The mechanism of electrophili c alkene epoxidation by the peroxycarbonate dianion is illustrated. The mechanism is identical to that of hydroperoxide oxidation, except that the nucleofuge of the peroxycarbonate dianion is the carbonate dianion.

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121 R1O R2R31 R1=CH3, R2=H, R3=H 2 R1=OH, R2=CH3, R3=H 3 R1=OCH2CH3, R2=H, R3=H 4 R1=C6H5, R2=H, R3=C(O)C6H5 Figure 4-4. Electrophi lic alkenes used in this study. Effect of Mn(II) on Electr ophilic Alkene Epoxidation While epoxidation of nucleophilic alkene s by peroxycarbonate are accelerated by the addition of manganese, the same is not true for electrophilic alkenes. For example, the NMR scale epoxidation of 1 (0.10 M) in D2O were allowed to react with 1.00 M NaHCO3 and 0.15 M H2O2 and has a t of 60 min (pH 7.8), Figure 4-5. 3,4-epoxybutanoneHd’exchangablewith solventHa’Hb’Hc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ Hdexchangablewith solventDOH 3,4-epoxybutanoneHd’exchangablewith solventHa’Hb’Hc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ Hdexchangablewith solventDOH Figure 4-5. 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 at 60 min, in the presence of 1.00 M sodium bicarbonate (50% conversion).

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122 When the same epoxidation reaction is performed in the presence of 4 M Mn(II), only a 44% conversion to the epoxi de is obtained in 60 min, as seen in Figure 4-6. The decrease in the rate of epoxida tion is attributed to the meta l assisted disproportionation of H2O2, which has a lower concentration as a result. 3,4-epoxybutanoneHd’exchangablewith solventHa’Hb’Hc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ Hdexchangablewith solventDOH 3,4-epoxybutanoneHd’exchangablewith solventHa’Hb’Hc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ Hdexchangablewith solventDOH Figure 4-6. 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 with 4 M Mn(II) in the presence of 1.00 M sodium bicarbon ate at 60 min (44% conversion).

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123 To further support the eviden ce that trace metal does not affect the reaction of electrophilic alkenes, 1 (0.10 M) in D2O was allowed to react with 1.00 M NaHCO3 and 0.15 M H2O2 with the addition of 5 mM diethylenetriaminepen taacetic acid (DTPA), a metal chelator. The reaction retains the same t of 60 min as seen in the reaction without DTPA, Figure 4-7. 3,4-epoxybutanoneHc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ DOH DTPA 3,4-epoxybutanoneHc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ DOH DTPA Figure 4-7. 1H-NMR of 1 epoxidized by H2O2 at pH 7.8 with 5 mM DTPA in the presence of 1.00 M sodium bicarbon ate at 60 min (50% conversion).

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124 In contrast, when 1 (0.10 M) in D2O is allowed to react with 1.00 M NaHCO3, 0.15 M H2O2, and 10 L saturated NaOH in D2O (final pH 8.6) the t is reduced to 15 min, Figure 4-8. 3,4-epoxybutanoneHc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ DOH Ha’Hb’ 3,4-epoxybutanoneHc’ O O Hd'Hc'Hb'Ha' O HdHaHcHbMethyl Vinyl KetoneHa& HbHcHd,Hd’ DOH Ha’Hb’ Figure 4-8. 1H-NMR of 1 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 8.6 afte r 15 min (50% conversion).

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125 These results illustrate that the addition of manganese(II) only affects electrophilic alkene epoxidation by disproportionating the terminal oxidant, hydrogen peroxide, and thus, reducing the conversion of epoxide. Increasing the pH of the solution, however, affects the conversion of electrophilic alkene s to epoxide by increas ing the equilibrium concentration of the nucle ophilic oxidant, hydroperoxide. Effect of pH on the Oxidation of Electrophilic Alkenes To assure that the results are not substrate specific, 2 was chosen for study. When 2 (0.05 M) is allowed to react with 1.00 M NaHCO3 and 0.30 M H2O2 (pH 7.8) for 24 hrs, there is a 75% conversion to the corresponding epoxide, Figure 4-9. HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ Figure 4-9. 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 7.8 in 24 hrs (75% conversion).

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126 When the same reaction is repe ated with the addition of 10 L saturated NaOH in D2O (final pH 8.6), only a 50% conversion is obtained after 24 hrs, Figure 4-10. HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ Figure 4-10. 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 8.6 at 24 hrs (50% conversion).

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127 The decrease in reactivity of what is a pparently an electrophilic alkene can be rationalized by the predominant resonance fo rm at the operational pH, Figure 4-11. At pH 8.6, methacrylic acid will mainly be deprotonated and will not react with the hydroperoxide anion. Methacrylic acid at pH 7.80 should react more like a nucleophilic alkene, so the addition of Mn(II) should greatly enhan ce the reactivity, based on the evidence presented in Chapter 3. When 2 (0.05 M) is was allowed to react with 1.00 M NaHCO3,0.30M H2O2, and 4 M Mn(II) (pH 7.8), >90% conversion to the epoxide is observed in 15 minutes, Figure 4-12. 0 0.03 0.06 0.09 0.12 0.15 0.18 0.21 0.24 0.27 0.3 01234567891011121314 pHconcentration, M [Methacrylic Acid] [Hydroperoxide] x 103 0 0.03 0.06 0.09 0.12 0.15 0.18 0.21 0.24 0.27 0.3 01234567891011121314 pHconcentration, M [Methacrylic Acid] [Hydroperoxide] x 103 Figure 4-11. Speciation of methacrylic ac id (MAA) at 0.05 M and hydrogen peroxide (0.30 M) as a function of pH. Th e maximum concentration of MAA and hydroperoxide happens to occur at pH 7.8, the buffering pH of sodium bicarbonate.

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128 HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ HO OHaHbHcMethacrylic AcidHbHaHc DOH HO O O Hc'Ha'Hb'Methacrylic Acid OxideHc’Hb’Ha’ Figure 4-12. 1H-NMR of 2 epoxidized by H2O2 in the presence of 1.00 M sodium bicarbonate at pH 7.8 with 4 M Mn(II) in 15 min (>90% conversion). Experiments were then conducted using 3, since the ester moiety provides a similar chemical environment for the double bond as 2, but without allowing for the disruption of the alkene delocalization caused by acid dissociation at the operational pH. When 3 (0.05 M) is allowed to react with 1.00 M NaHCO3 and 0.30 M H2O2 (pH 7.8) for 24 hrs, a 44% conversion to the corresponding epoxide is observed by 1H NMR, Figure 4-13.

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129 O OHaHcHbHdHeEthyl acrylate O O O Hd'He'Hb'Ha'Hc'Ethyl-2,3-epoxypropionateHaHa’HbHcHdHe Hb’Hc’Hd’He’ DOH Figure 4-13. 1H-NMR of 3 epoxidized by H2O2 in 1.00 M sodium bicarbonate at pH 7.8 in 24 hrs (44% conversion). As predicted by the oxidation of 1, when the same reaction is repeated using 10 L of saturated NaOH in D2O (final pH 8.6) for 24 hrs, a 66% conversion to the epoxide is observed, Figure 4-14. This e xperiment proves that if there is no change in the resonance structure of the electrophilic alkene, it will react with the nuc leophilic oxidant as predicted.

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130 O OHaHcHbHdHeEthyl acrylate O O O Hd'He'Hb'Ha'Hc'Ethyl-2,3-epoxypropionateHaHa’HbHcHdHe Hb’Hc’Hd’He’ DOH O OHaHcHbHdHeEthyl acrylate O O O Hd'He'Hb'Ha'Hc'Ethyl-2,3-epoxypropionateHaHa’HbHcHdHe Hb’Hc’Hd’He’ DOH Figure 4-14. 1H-NMR of 3 epoxidized by H2O2 in 1.00 M sodium bicarbonate at pH 8.6 in 24 hrs (66% conversion). Effect of Buffer Choice on Electrophilic Alkene Epoxidation In order to determine the role that the peroxycarbonate dianio n has on the rate of electrophilic alkene oxidation, a study was conducted using NaHCO3 and Na2HPO4 as buffering salts. When solutions of 0.10 M 1, 1.00 M buffer, and 0.15 M H2O2 were allowed to react at varying pH, Table 4-1, it was observed that identical conversions to

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131 the corresponding epoxide were found in the same amount of time and was not dependent upon buffer choice, but only on pH. Table 4-1. The percent conversion of 1 in varying buffer at differing pH. Buffer pH 7.80, 150 min pH 8.60, 90 min pH 8.90, 60 min NaHCO3 74% 92% >99% Na2HPO4 74% 92% >99% Reaction conditions: 0.10 M 1, 1.00 M Buffer and 0.15 M H2O2 This observation leads us to the conclusi on that the peroxycarbona te dianion is at least as reactive as hydroperoxide in the epoxi dation of electrophilic al kenes. This result is not surprising, since simila r results have been seen in the use of the peroxycarbonate dianion in the cleavage of organophosphate esters.140 Similar reactions of 2 and 3 in phosphate and bicarbonate buffer gave identical conversions to epoxide. The explanation for why similar yields of epoxide are obtaine d in the different buffers will be discussed shortly. Electrophilic Alkene Oxidation Kinetics Oxidations of 4 were conducted in aqueous, mi cellar solution unde r second-order conditions using hydrogen peroxi de and sodium hypochlorite as terminal oxidants. The oxidation was followed by monitoring the loss in peak intensity at 320 nm spectrophotometrically. Fo r the reactions with h ydrogen peroxide, 4.00 x 10-4 M substrate and 8.00 x 10-4 M oxidant were used. Reacti ons were performed in aqueous solution consisting of 0.10 M phosphate bu ffer and 0.10 M cetyltrimethylammonium chloride (CTACl) as surfactan t. The reactions were conduc ted over a pH range of 7.5 – 12. The data is presented as the log(kobs) vs. pH in Figure 4-15 to demonstrate that the kobs at lower pH values are not zero. From this data, Equa tion 4-3 was derived using the steady-state approximation for HOO-.

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132 H2O2 H+ + HOOKa (4-1) S + HOOSO + OHk2 (4-2) )) K ] [H ( (1 [S] ] O [H t [S]a 0 2 2 2 k (4-3) The rate law indicates that there are two pa rameters that can be fit to the data, the rate constant and the pKa of the oxidant. Non-linear re gression was used to provide values for the rate constant and pKa of the hydrogen per oxide oxidation of 4, presented as the solid line in Figure 4-15. From the fit, the value of k2 = 660 40 M-1s-1 and the pKa of hydrogen peroxide is 11.70. -1.2 -0.7 -0.2 0.3 0.8 1.3 1.8 2.3 2.8 7.008.009.0010.0011.0012.00 pHlog( kobs) Experimental Predicted Figure 4-15. Plot of log( kobs) vs pH for the oxidation of 4 by H2O2. Since hypochlorite is also a nucleophilic oxidant, similar resu lts should be found for reactions with 4. Reaction conditions were identical to that with hydrogen peroxide, except the hypochlorite c oncentration was 2.40 x 10-3 M and the pH range was from 5 – 10. A plot of the log(kobs) vs pH is shown in Figure 4-16. Non-linear regression analysis

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133 using the derived rate law is shown in Figure 416 as the line. From the fit, the value of k2 = 118 2 M-1s-1 and the pKa of hypochlorous acid is 8.26. Figure 4-16. Plot of log( kobs) vs pH for the oxidation of 4 by -OCl. Discussion of the Second-Order Rate Constants As stated in the introduction, only two references in th e literature have presented detailed kinetic data on the topic of electrophilic epoxidati on by nucleophilic oxidants.131,132 From the work of Rosenblatt and Broome,131 the mechanism of the epoxidation has been attributed to the slow addition of th e nucleophilic oxidant to the electrophilic alkene followed by a rapid internal cyclization to form the epoxide, as seen in Figure 4-17. O RR'' R' HOO-rate determining O R R'' R' OOH R' O + OH-fast R O R'' Figure 4-17. Reaction mechanism determ ined by Rosenblatt for the nucleophilic oxidation of electrophilic alkenes.131 -1.50 -1.00 -0.50 0.00 0.50 1.00 1.50 2.00 2.50 5.006.007.008.009.0010.00 pHlog( kobs) Experimental Predicted

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134 Rosenblatt and Broome131 studied the reaction of o chlorobenzylidenemalononitrile, Figure 4-18, with hydroperoxide and hypochlorite at 25 C in CH3CN:H2O (1:99 (v:v)) mixtures. The secondorder rate constant for the reaction in the presence of hydroperoxide was 4.00 x105 M-1s-1, while the second-order rate constant for hypochlorite was 2.20 x104 M-1s-1. This indicates that the hydroperoxide anion is about 20 times more reactive than hypochlorite. CN CN Cl Figure 4-18. Structure of the substr ate used by Rosenblatt and Broome,131 o chlorobenzylidenemalononitrile. The explanation for the observed reactivity is due to the basicity of the corresponding nucleophiles. Since hydrogen peroxide is a weaker acid (pKa =11.65)141 than hypochlorous acid (pKa = 7.40)142, the hydroperoxide anion is a stronger base than hypochlorite. Since hydroperoxide is the stronger base, it will react with th e electrophilic carbon of the alkene at a faster rate than hypochlorite. Since the rate determ ining step of the reaction is the nucl eophilic addition to th e electrophilic carbon, the reactivity of the nucleophilic anion will determine the rate of addition to a particular substrate. The nucleophile basicity explains why there was no observed difference in the yields of the reactions of 1 in sodium bicarbonate and s odium phosphate presented in Table 4-1. Sin ce hydroperoxide, pKa 11.65,141 and peroxycarbonate have similar pKa values, their reactivity toward electrophilic alkenes would be nearly identical. Since they are both strong bases they will add at approxima tely the same rate to the same substrate giving rise to the identical c onversions seen in th e epoxidation data seen in Table 4-1.

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135 The reactivity of peroxycarbonate and hydrope roxide also confir ms the mechanism of Rosenblatt and Broome.131 Since the fast step of the reaction has been proposed to be the internal cyclization to form the epoxide and the release of the leaving group, the stability of the leaving group is irrelevant. If th e slow step of the epoxidation were the ring closure, the identity of the leaving group would become the dominant characteristic in the determination of reactivity, as seen in Figure 4-19. If the ri ng closure were the rate determining step, peroxycarbonate would react faster than hydroper oxide since carbonate is the more stable leaving group over hydroxide. Since there is no increase in reactivity of peroxycarbonate over hydroperoxide in the data presented in Table 4-1, the rate limiting step must be the addition of the nuc leophilic oxidant to the electrophilic carbon, supporting the conclusion of Rose nblatt and Broome, Figure 4-20. Reactants ProductsEne r gyReactionCoo r d inateR OR' R'' OZ R OR' R'' OZ O R O Z R' R'' Figure 4-19. A reaction coordina te diagram for the addition of a nucleophilic oxidant to an electrophilic alkene. If the ring closure of th e epoxide is the rate determining step, as seen in the diagram by a larger energy barrier for the formation of the epoxide, the identity of the nucleofuge (Z) will determine the reactivity of the oxidant. The more st able the nucleofuge, the more reactive the oxidant will be.

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136 Reactants ProductsEne r gy R eactionCoo r d inateR OR' R'' OZ R OR' R'' OZ O R O Z R' R'' Figure 4-20. A reaction coordina te diagram for the addition of a nucleophilic oxidant to an electrophilic alkene. If the addition of the oxidant is the rate determining step, as seen in the diagram by a larger energy barrier to the formation of the intermediate, the identity of the nucle ofuge (Z) will no longer determine the reactivity of the oxidant. The more basic oxidant will react faster. Bunton and Minkoff132 studied the epoxidation of mesityl oxide and ethylideneacetone by the hydroperoxide anion in aqueous solution. They observed second-order rate constants of 1.37 x 10-2 and 7.70 x 10-2 M-1s-1, respectively, for the oxidation at 0 C. Using the general rule that reaction rates d ouble for every 10 C increase in temperature, the second-o rder rate constants would be 8.22 x 10-2 and 4.62 x 10-1 M-1s-1, respectively, at 25 C. Given that hydroperoxide reacts faster than hypochlorite, the second-order rate constants observ ed by Bunton and Minkoff132 are orders of magnitude smaller than th at observed by Rosenblatt and Broome’s,131 4.00 x 105 M-1s-1.

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137 O Mesityl oxide O Ethylideneacetone Figure 4-21. Substrates used by Bunton and Minkoff to study the oxidation of electrophilic alkenes by th e hydroperoxide anion. The apparent discrepancies in rate cons tant observed can only be rationalized by the particular substrate being epoxidized. In the case of Bunton and Minkoff, both substrates were -unsaturated ketones, Figure 4-21, while in the case of the work by Rosenblatt and Minkoff, the substrate unde r investigation was a dicyano compound, Figure 4-18. The addition of two cyano groups on the -carbon makes the -carbon more electrophilic than the -carbon for the substrates used by Bunton and Minkoff. The presence of a more electrophilic carbon then e xplains why such varyi ng rate constants are observed for the reactions with hydroperoxide. In the work presented in this report, th e second-order rate constants observed for the epoxidation of 4 by hydroperoxide and hypochlorit e were 660 40 and 118 2 M-1s1, respectively. While the epoxidation by the hydroperoxide anion in this study is 5.6 times greater than that for hypochlorite, Rosenblatt and Broome131 observed a 20 times greater reactivity for hydroperoxide ov er hypochlorite in the oxidation of ochlorobenzylidenemalononitrile. The increas ed reactivity of hydroperoxide may depend on the substrate involved, as seen in the co mparison with the secondorder rate constants observed by Bunton and Minkoff132 for the epoxidation of mesityl oxide and ethylideneacetone. Since the substrate used for this study has a second carbonyl group to the alkene, this may account for the increa sed reactivity observed. The surfactant may also account for the difference in the second-order rate constants.

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138 There are a couple of factors that might a ffect the reaction of electrophilic alkenes with nucleophilic oxidants in surfactant soluti ons. First, since a cationic surfactant was used for this study, the cationic head group c ould stabilize the enolate intermediate. By neutralizing the negative char ge on the oxygen atom, the ener gy of the intermediate would be lowered. By lowering the energy of the intermediate, the rate of the reaction should be slower. In order to determine whet her the surfactant is having the postulated effect, experiments would need to be c onducted to observe the second-order rate constants in the absence of surfactant. Si nce the substrate is not soluble in aqueous solution, mixed solvent systems would n eed to be used, which may introduce new variables. Second, the intr oduction of the cationic surfacta nt may increase the observed second-order rate constants by concentrating the active oxidant at the micelle surface. In addition to the differences in th e second-order rate constants, the pKa values for the two oxidants were also fit using Equation 43. It is known that micellar media affects the acid-base dissociat ion of organic dyes,57,143-150 the substrates of choice for monitoring acid-base equilibria in micellar solution, si nce they can be easily monitored by their characteristic absorbance by Uv-vis. It ha s been suggested that effects on acid-base equilibria may be related to the repulsion of protons by cationic surfactants, and their attraction to anionic surfactan ts. Fernandez and Fromherz143 have attempted to compute these attractive and repulsive forces using th e micelle as a submicroscopic solvent, along with a term to account for the su rface potential of the micelle. To further complicate the prediction of aci d-base equilibria in micellar media, attempts to understand the affect of added salt s to the micellar media have been attempted by Romsted149 and Quina.150 In their work, equations have been developed to describe all

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139 of the ionic exchanges between the bulk wate r and the micellar surface. For all of these studies, the acid-base equilibria of any give n acidic substrate may increase or decrease upon incorporation in a micelle, depending on the particular substrate and the concentration of added salts.151 Table 4-2 lists the pKa values for a group of dyes in aqueous and micellar (CTACl) solution. Table 4-2. pKa values in water and CTACl for several different dyes.151 water CTACl dye pKa1 pKa2 pKa3 pKa1 pKa2 pKa3 fluorescein 2.14 4.45 6.80 0.60 6.41 7.17 sulfonefluorescein 3.22 6.76 2.33 7.00 2,7-dichlorofluorescein 0.35 4.00 5.19 <-0.5 5.50 5.79 eosin -2.0 2.81 3.75 1.83 5.76 ethyl eosin 1.91 1.11 Further kinetic studies on the epoxidati on of electrophilic al kenes in both pure water and micellar media by nucleophilic oxida nts need to be conducted to determine which of these two possibilitie s are more or less respons ible for the observed secondorder rate constants. Conclusions Unlike epoxidations of nucleophilic alkene s, electrophilic alke ne oxidation is not catalyzed by manganese(II) with hydrogen per oxide as the terminal oxidant in the presence of bicarbonate. In fact, by addi ng manganese(II) to the reactions, the hydrogen peroxide is disproportionated, thus decreasing the concentrati on of the active oxidant, the hydroperoxide anion. Further evidence that ma nganese(II) does not affect the reaction is observed when 5 mM DTPA is added to a reaction of 1. The half-life of this reaction is unaltered from that of a reaction with no metal added. This indicat es that electrophilic alkenes do not react with either the proposed high valent manganese species, or the carbonate radical anion.

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140 Reactions of methacrylic acid indicate that the dominant species of both the oxidant and substrate present in solution at the operating pH are important. This was demonstrated by plotting the concentrations of both methacrylic acid and hydroperoxide as a function of pH, Figure 4-11. From the plot, the greatest react ivity will occur where the methyacrylic acid and hydrope roxide concentrations are ma ximized. This occurs at a pH value of 7.8. For the r eactions observed, the deprotona ted alkene, methacrylate, does not react as an electrophilic alkene but ha s a reactivity that resembles a nucleophilic alkene. The addition of manganese(II) to thes e solutions, therefore, should increase the reactivity. In fact, when 4 M manganese(II) was allowed to react with methacrylic acid at pH 7.8 in the presence of 1.00 M sodium bicarbonate, the reaction was observed to be more than 90% complete in 15 min. Finally, in order to assure that the reactiv ity seen with methyl vinyl ketone was not substrate-specific, reactions of ethyl acrylate were conducte d. Since ethyl acrylate does not contain the ester moiety, the problems en countered with methacrylic acid, namely the deprotonation of the acid, should not be pres ent. When ethyl acr ylate was allowed to react with hydrogen peroxide in the presence of bicarboante at pH 7.8, a 44% conversion to the epoxide was observed in 24 hrs. At pH 8.6, however, a 66% conversion to the epoxide was observed in 24 hrs, proving that the increase in pH is responsible for the increased reactivity. Experiments with methyl vinyl ketone we re also conducted to determine whether the peroxycarbonate dianion is an effective oxidant of electrophilic alkenes. Reactions were performed under identical conditions with sodium bicarbonate and sodium phosphate as the buffering salts. Conversions of methyl vinyl ketone at the same pH

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141 values gave identical conversi ons in the different buffers. This result indicates that the choice of buffer does not make a difference, only the pH of the solution. Also, this result indicates that peroxycarbonate is nearly equal to hydroperoxi de at oxidizing electrophilic alkenes. These data support the me chanism proposed by Rosenblatt and Broome131 for the epoxidation of electrophilic alkenes by nuc leophilic oxidants. The rate determining step is the attack of the nucleophilic oxidant at the -carbon of the electrophilic alkene. Therefore, the hydroperoxide anion and pe roxycarbonate dianion should have similar reactivities, which is supported by the data in Table 4-1. Kinetic experiments were conducted for th e reaction of dibenzoylethylene with hydroperoxide and hypochlorite. The second-orde r rate constant for the reaction with hydrogen peroxide is 660 40 M-1s-1 and the pKa of hydrogen peroxide is 11.70. For the reaction of hypochlorous acid with dibenzoylethy lene, the second-order rate constant is 118 2 M-1s-1 and the pKa of hypochlorous acid was determined to be 8.26. Limited kinetic data available in the lit erature for the reaction of electrophilic alkenes with nucleophilic oxidants make interp retation of the observed second-order rate constants difficult. From the data that has been reported, the iden tity of the substrate seems to make the greatest impact on the react ivity with nucleophilic oxidants. This can be examined by comparing the second-order rate constants observed by Rosenblatt131 and Bunton.132 The substrate used by Rosenblatt and Broome, ochlorobenzylidenemalononitrile, has a second-or der rate constant that is orders of magnitude greater than the substrates us ed by Bunton and Minkoff, mesityl oxide and ethylideneacetone. Since ochlorobenzylidenemalononitrile has two cyano groups on the carbon of the alkene, the -carbon is more electrophilic than the -carbon of the

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142 substrates of Bunton and Minkoff, which only have one electron withdrawing group bonded to the carbon of the alkene. Limited data using surfactants for th e nucleophilic oxidati on of electrophilic alkenes has also made interpretation of the obs erved second-order rate constants difficult. For this work, the cationic surfactant used ma y contribute to the la rger second-order rate constants observed by concentrating the act ive oxidant at the micellar surface, which would increase the observed second-order ra te constants. The cationic surfactant may also affect the second-order rate cons tants by stabilizing th e oxygen atom of the intermediate enolate. By stabilizing the in termediate, however, the rate constants would decrease, since the intermediate would be lower in energy. Materials and Instrumentation Hydrogen peroxide (35%) was purchas ed from Sigma-Aldrich and was standardized often by iodome tric titration. Sodium hypoc hlorite was purchased from Fisher and its concentration was checked sp ectrophotometrically by observing the peak at 236 nm ( = 100 M-1cm-1). Methyl vinyl ketone (Aldri ch), methacrylic acid (Aldrich), ethyl acrylate (Aldrich), and dibenzoylethyl ene (Acros) were purchased as their highest grade and used without further purificati on. Sodium bicarbonate, sodium phosphate dibasic, and sodium phosphate monobasic (Fis her) were purchased as analytical grade and used without further purification. Diethylenetriaminepentaacetic acid (SigmaAldrich), cetyltrimethylammonium chloride (Sigma-Aldrich), and manganous sulfate (Fisher) were purchased and used without fu rther purification. Water was purified by a Barnstead E-Pure 3-Module Deionization System. 1H-NMR spectra were obtained on a Mercury 300 or VXR 300 spectrophotometer in D2O (Cambridge Isotope) and the residual proton peak (DOH) used as internal standard. UV-vis kine tic experiments were

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143 obtained using a Hewlett-Packer 8453 spectroph otometer using 1.0 cm polystyrene cells from Fisher. The reaction temperatures were maintained at 25 0.1 C using a Fisher Isotemp 1600S water bath circulator. Experimental Electrophilic Alkene Epoxidation Electrophilic alkene reacti ons were performed in D2O in both sodium phosphate (1.00 M) and sodium bicarbonate (1.00 M) solutions at pH 7.8. Reactions which were performed at pH 8.6 had 10 L of added saturated NaOH dissolved in D2O. Alkene concentration was 0.10 or 0.05 M, as indicate d in the text. Hydrogen peroxide was 0.15 M for reactions with methyl vi nyl ketone and 0.30 M for reac tions with methacrylic acid and ethyl acrylate. Reac tions were monitored by 1H NMR by observing the disappearance of the alkene protons and the appearance of the epoxide protons. Percent conversions were determined using the integr ations of the alkene and epoxide protons. NMR spectra were recorded on a VXR 300 MHz or Mercury 300 MHz instrument. Dibenzoylethylene Kinetics Reactions of dibenzoylethylene were performed in 0.10 M CTACl, 0.10 M phosphate buffer, and either 8.00 x 10-4 M hydrogen peroxide, or 2.40 x 10-3 M sodium hypochlorite at 25 C. Dibenzoylethylene was dissolved in ethanol. LReactions were initiated by addition of 72 of dibenzoylethyle ne solution. Total organic solvent in the reactions was less than 2.5%. Reaction progress was monitored by loss of peak intesity at 320 nm using a Hewlett-Packer 8453 spect rophotometer with multi-cell capabilities.

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144 CHAPTER 5 GENERAL CONCLUSIONS Hydrogen peroxide decomposition and nucle ophilic alkene epoxidation are easily achieved using bicarbonate and manganese(II) as catalysts. The formation of a high valent metal oxo species initiates the reaction for both the decomposition and nucleophilic alkene epoxidation. This oxi dant can deliver an oxygen atom to other organic species as indicated by th e products of amine oxidation. In Chapter 2, large scale epoxidations in micellar media have been shown to be effective for producing large amounts of epoxi de, although significant hydrolysis of the product epoxide has been detected when the reactions are performed in the absence of metal. The long reaction times can be overcome by performing the reaction in the presence of Mn(II). Surfactants have b een used as a means of dissolving hydrophobic substrates into the aqueous solutions. From the observed kinetic data, the choice of surfactant, either cationic or anionic, does not have a significant impact of the reaction rate. Also, the source of manganese, either from a bulk stock of the ion or paired with the surfactant, has no significant im pact on the reaction rate for the epoxidation. Purification of the resulting epoxide from the reacti on solutions has been accomplished by using liquid-liquid extraction. This method provides a useful alternative to normal extraction techniques that in the presence of the surf actants produce emulsions. Interesting results of the kinetics of the reactions of nucleophi lic alkenes in micella r solution led to the investigation of the background hydrogen peroxide decomposition.

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145 Chapter 3 describes the investigations of the background metal-assisted decomposition of peroxide in bicarbonate so lution. The kinetics of the reaction have been measured and the results are similar to wo rk that has been previously reported. The dependence of the reaction on the metal cation has been observed to be first-order. This indicates that only one metal cation is pr esent in the active metal complex. The bicarbonate dependence for the reaction ha s also been found. A second-order dependence was observed which is consiste nt with the results observed for the nucleophilic alkene epoxidation. This result indicates that two bicarbonate anions are present in the active metal sp ecies. The hydrogen peroxide dependence has also been observed for these reactions. For the hydroge n peroxide decomposition, the dependence is linear through a concentration of 0.5 M. This is in contra st to the inverse relationship observed for the oxidation of nucleophilic alkenes. The lifetime of the active metal catalyst wa s also conducted. Results from these experiments indicate that the active species is easily regenerated from spent solutions by the reintroduction hydrogen peroxi de. A loss in the catalytic rate is observed over time. Experiments were conducted to determine what factors were contributing to the loss of activity. Large scale cycling of a decompos ition reaction was used to show that the concentration of bicarbonate is falling thr oughout the reaction. Since the reaction has a second-order dependence on bicarbonate the lo ss of bicarbonate has the most dramatic effect on the loss of activity. In addition to examining the bicarbonate concentration, the hydrogen peroxide was examined as a possible so urce of contaminants that would lead to the loss of activity on the catalyst lifetime. Results using the malachite green assay for phosphates indicate that the stock hydrogen pe roxide solutions ar e stabilized with

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146 phosphates. The introduction of these phosphate s has a deleterious effect on the reaction by precipitating the metal cation. Reactions performed using distilled hydrogen peroxide and adding bicarbonate after each reaction was shown to help prevent the loss in the active catalyst. The inability to full control the bicarbonate is still a factor in the loss of activity. In addition to examining the lifetime of the catalyst, the source of the manganese ion was also investigated. Manganese(II) sulfate, potassium permanganate, and a Mn(IV) TACN complex were each used to examine the decomposition of hydrogen peroxide. When reactions were performed using the sa me concentration of manganese(II) sulfate and potassium permanganate, identical observed ra te constant, within error, are observed. This is not surprising since permanganate is known to react with hydrogen peroxide to generate Mn(II) cations. A Mn(IV)-TACN complex that has been tout ed as an excellent epoxidation catalyst for nucleophilic alkenes using hydrogen peroxi de as the terminal oxidant was also examined for its ability to catalyze the deco mposition of hydrogen peroxide. Initially, the stability of the catalyst was examined sp ectrophotometrically in the presence of bicarbonate and hydrogen peroxide. When a solution of the ca talyst is dissolved in a bicarbonate solution, the catalyst appears to be stable, as the absorbance of the complex does not change with time. Upon addition of 1 equivalent of hydrogen peroxide, the catalyst’s absorbance rapidly drops to near zero. Af ter time, however, a broad absorbance is detected and solutions turn br ight yellow. When 2 equivalents of hydrogen peroxide were used, a different result was obs erved. Instead of a yellow color developing after the addition of the hydrogen peroxide, the absorban ce remained close to zero for

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147 several minutes. It is believed that th e decomposition of the catalyst is being accomplished by oxidative N-dealkylation of the TACN ligand. In the presence of only 1 equivalent of hydrogen peroxide, the decompos ition of the catalyst is incomplete and the generation of metal-amine coordination compou nds explains the resulting yellow color. On the other hand, when two equivalents of peroxide were used, the catalyst is completely decomposed, and no metal-amine in termediates are produced to give rise to the yellow color observed. When the same concentration of the Mn(IV)-TACN complex was used for the decomposition of hydrogen peroxide in bicarbonate solution, similar observed rate constants to the manganese(II) sulfate reactions were found. This result indicates that the Mn(IV)-TACN complex qui ckly decomposes to release a manganese ion which decomposes the hydrogen peroxide at the same rate as if the metal ion had been introduced as a salt. Cis/trans isomerization was also examined in this study in an attempt to understand how the active oxygen catalyst was deliveri ng the oxygen atom. If the oxygen is being delivered via a concerted mechanism, rotati on of the cis alkene upon oxidation should not be observed, as is the case for peracid epoxi dation. However, if rotation to the more stable trans-epoxide is observed, the result would indicate that the C-C sigma bond has the ability to rotate in the transition state to the more stable trans conformation. A numbe of different substrates were attempted to be used for this study, but maleic and fumaric acid were eventually used in these experiments. When the cis acid, maleic acid, was epoxidized using the manganese system, both the cis and trans epoxides were observed by 1H NMR. This result, as mentioned earlier, indicates that rotation to the more stable trans conformation in the tran sition state is occurring.

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148 Examination of the oxidation of the radical traps used in other studies to confirm the presence of hydroxyl radicals was also conducted. Results from these experiments indicate that the oxidation system can reac t with amines to yield their oxidative Ndealkylation products. The proposed mechanism for the metal catalyzed oxidative Ndealkylation is based on current work investig ating the reactivity of cytochrome P-450, an iron containing enzyme. Current literature proposes that the mechanism proceeds through a single electron transfer. Support fo r this mechanism has come from work involved with amines whose N-alkyl groups have been deut erated. Results of oxidation with these substrates yield deuterium isotope effects of near unity. These data suggest that the breaking of the al pha-carbon hydrogen bond is not rate limiting. If the reaction were progressing via a hydrogen abstra ction, a normal isotope effect, kH/kD >1, would be expected. Solvent isotope effect studies have also been conducted for the nucleophilic alkene epoxidation and hydrogen peroxide decompositi on. A large, inverse solvent isotope effect has been observed for the epoxidation reaction, while a normal isotope effect was measured for the decomposition reaction. The result of the inverse isotope effect for the alkene epoxidation reaction indi cates that there is a proton tr ansfer in the transition state of the reaction that proceeds faster in solutio ns of deuterium oxide than water. Possible reaction mechanisms have been proposed in an attempt to explain the observed effect. Two possibilities exist in this system. First, the metal may simply be acting as a Lewis acid to coordinate and activat e the peroxycarbonate. The s econd proposed explanation is that the reaction is proceeding by reaction with the carbonate radi cal anion. Both of these

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149 reactants can be uncharged and would rati onalize the rates of epoxidation observed in cationic and anionic surfactants. Finally, a mechanism is proposed for the activation of peroxycarbonate to form a high oxidation metal oxo complex which could be responsible for the observed reactions. In addition, a radical oxida tion route is also proposed fo r the observed reactions of nucleophilic alkenes and hydrogen peroxide de composition. Numerical simulation has been used in an attempt to either confirm or reject possible reaction mechanisms. This work indicates that the generation of carbona te radicals and a high oxidation state metal oxo complex are reasonable. In addition, simulations using the carbonate radical mechanism indicate that the hydrogen per oxide decomposition is inhibited with increasing concentrations of hydrogen peroxide. This interesting result might explain the downturn in the hydrogen peroxide depende nce observed for nucleophilic alkene epoxidation. Chapter 4 discussed the use of the pe roxycarbonate dianion as a nucleophilic oxidant of electrophilic alkenes. Experime nts conducted in the pres ence of metal indicate that the reaction has no dependen ce of metal, and in fact, the reaction is inhibited by the addition of metal by decomposing the active oxi dant. Further support that the metal has no effect on the epoxidation of electrophilic alkenes has come from experiments using metal chelating agents. When DTPA is added to the reactions, no change in t is observed. Therefore, the mechanism describe d in Chapter 3 is not responsible for the epoxidation. The pH of the solutions was found to affect the rate of epoxidation of electrophilic alkenes. This would support the idea that a nucleophilic al kene is responsible for the

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150 epoxidation. As the pH is raised, the c oncentration of the hydroperoxide anion would increase. Since this is proposed to be the active epoxidizing agent, a greater concentration of the oxidant will yield more epoxide. The predominant resonance structure at th e operating pH also plays an important role in the epoxidation of el ectrophilic alkenes. For example, the epoxidation of methacrylic acid is inhibited at higher pH. This is due to the fact that the acid will be deprotonated at the operating pH, and this alkene will react more similarly to a nucleophilic alkene. This was shown to be true when a reaction was conducted with the addition of manganese(II). An epoxidation th at was only 50% complete in 24 hrs, was >90% complete with the addition of the metal. Experiments using the similar ethyl acrylate, where the ester moiety replaces the aci d, resulted in trends that are consistent with the epoxidation of elect rophilic alkenes by nucleophilic oxidants. The ester moiety provides a similar chemical environment to that of the methacrylic acid, but deprotonation is no longer an issue. The kinetics of the epoxidation of dibe nzoylethylene were conducted in cationic micellar solutions. The reaction was found to fit an expression that describes the nucleophilic attack of the oxida nt on the substrate. Hydr ogen peroxide and hypochlorite were the oxidants chosen for study. The s econd-order rate consta nt for the epoxidation by hydroperoxide was observed to be 660 40 M-1s-1, with a pKa of hydrogen peroxide being 11.7. Hypochlorite reac tions resulted in the observa tion of a second-order rate constant of 118 2 M-1s-1 and the pKa of hypochlorous acid being 8.26. In conclusion, it has been shown that th e oxidation of organic substrates can be accomplished using peroxycarbonate. For substr ates which are more nucleophilic, metal

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151 activators are required to increase the rate of product formation. A mechanism based on a high oxidation state metal oxo complex and ca rbonate radicals are implicated in the reaction. On the other hand, electrophilic substrates can react with the peroxycarbonate dianion directly. The peroxycarbonate diani on has been shown to be the active species and has a reactivity similar to th at of the hydroperoxide anion.

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152 APPENDIX VARIATIONS IN NUCLEOPHILIC AL KENE EPOXIDATION AND HYDROGEN PEROXIDE RATE CONSTANTS The following figures were generated us ing the full mechanism of nucleophilic alkene epoxidation and hydrogen peroxide d ecomposition as presented in Chapter 3. Only the rate constant or equilibrium co stant indicated has been changed from the constants listed in Chapter 3. 0 0.0004 0.0008 0.0012 0.0016 0.002 00.20.40.60.81 [H2O2], Mkobs, s-1 Experimental k = 135000 k = 121000 k = 149000 Figure A-1. Variation in the S + A Products rate constant.

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153 0 0.0004 0.0008 0.0012 0.0016 0.002 00.20.40.60.81 [H2O2], Mkobs, s-1 Experimental K = 110 K = 99 K = 121 Figure A-2. Variation in the e quilibrium constant for A + H2O2 B. 0 0.001 0.002 0.003 0.004 0.005 00.20.40.60.81 [H2O2], Mkobs, s-1 Experimental K = 1 K = .1 K = 10 Figure A-3. Variation in th e equilibrium constant fo r the formation of “A”.

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154 -3.8 -3.6 -3.4 -3.2 -3 -2.8 -2.6 -2.4 -2.2 0100200300400500600700800 Time, secln([H2O2]) Experimental k = 9000 k = 8100 k = 9900 Figure A-4. Variation in the rate constant for A radicals.

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163 BIOGRAPHICAL SKETCH Andrew P. Burke was born March 31st, 1977, in Fort Walton Beach, FL, as the youngest of three children to A ndrew J. and Cynthia A. Burke. It was during his high school years at Choctawhatchee High School that Andrew developed his interest in chemistry. After graduating from high school in 1995, A ndrew went to Furman University in Greenville, SC. From the summer of 1998 until graduation, Andrew conducted undergraduate research under the direction of Dr. Laura Wright investigating Pt(0) diimine complexes. It was during this time that Andrew decided to pursue a doctoral degree in chemistry. After graduating from Furman University in 2000 with Bachelor of Science degrees in chemistry and computer science, A ndrew began pursuing a doctoral degree in inorganic chemistry at the Univ ersity of Florida in Gainesvi lle, FL, under the direction of Dr. David E. Richardson. In December 2002, Andrew married Erin Ringus, a member of the biochemistry division of the Chemistr y Department. After graduating with his doctoral degree, Andrew will be moving to Charleston, SC, to join High-Purity Standards.


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Permanent Link: http://ufdc.ufl.edu/UFE0011480/00001

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Title: Hydrogen Peroxide Disproportionation and Organic Compound Oxidation by Peroxycarbonate Catalyzed by Manganese(II): Kinetics and Mechanism
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Copyright Date: 2008

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Title: Hydrogen Peroxide Disproportionation and Organic Compound Oxidation by Peroxycarbonate Catalyzed by Manganese(II): Kinetics and Mechanism
Physical Description: Mixed Material
Copyright Date: 2008

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Source Institution: University of Florida
Holding Location: University of Florida
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HYDROGEN PEROXIDE DISPROPORTIONATION AND ORGANIC COMPOUND
OXIDATION BY PEROXYCARBONATE CATALYZED BY MANGANESE(II):
KINETICS AND MECHANISM















By

ANDREW P. BURKE


A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT
OF THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY

UNIVERSITY OF FLORIDA


2005

































Copyright 2005

by

ANDREW P. BURKE

































to my wife and parents















ACKNOWLEDGMENTS

The work presented here would not have been possible without the help and

support of a number of people. I would like to acknowledge these people individually for

their contributions in making the following document possible.

First, I would like to thank my advisor, Dr. David Richardson, for all of his help

and support during these past five years. Dr. Richardson has helped to make me a better

scientist. My presentation and writing skills have vastly improved under his advisement,

and they will prove useful in all my future endeavors. I have also had the opportunity to

learn from Dr. Richardson the proper method for performing chemical kinetics.

I would also like to thank the members of the Richardson group, both past and

present, for all of their help during the years. Without the fun environment they created,

working in the lab would have been much less enjoyable. I would like to thank my

partner in crime, Dan Denevan. It was always nice having Dan to make jokes with and to

have around to complain to about everything going wrong in the lab. I would also like to

thank Dr. Ana Ison for all of her support during this process. Ana was always around to

discuss ideas about projects. I would also like thank her for the help she provided in

acquiring the GC data. I would especially like to thank Dr. Celeste Regino. Celeste

taught me the intricacies of HPLC and that in order for it to do what you want you must

coddle it at all times.

I would also like to thank Pat Butler for all of her hard work during my elementary

education. As my SLD teacher, Mrs. Butler worked with me constantly for many years









to help me cope with a disability I never thought I would be able to overcome. Now, as I

finish my dissertation to achieve my Ph.D., I appreciate even more all of the techniques

she taught me to help me achieve my goals.

I could not have accomplished this goal without the support of my family,

especially my parents. My parents have always supported me in all of the decisions I

have made and attending graduate school was no exception. Without their support, I

would never have had the courage to face new challenges and persevere in the face of

opposition.

I would also like to thank my wife, Erin. I never expected to meet my wife in

graduate school, nor did I expect her to be a chemistry graduate student. She has been a

constant support these past 5 years, and I would have given up this dream long ago were

it not for her constant vigilance in driving me toward my goal.

Most of all, I would like to thank God for all of His love and support through not

only my graduate career, but my entire life. Through Him all things are possible.
















TABLE OF CONTENTS

page

A C K N O W L E D G M E N T S ................................................................................................. iv

LIST OF TABLES ............. ...... ......... ............. ... ............................... ix

LIST OF FIGURES ............................... ... ...... ... ................. .x

A B S T R A C T .........x.................................... ....................... ................. xx

CHAPTER

1 IN TR OD U CTION ............................................... .. ......................... ..

G general O xidation ........... .................................................................. ......... . ...
R active O xygen Species ............................................................. ....................... 3
Hydrogen Peroxide .................. ............................................. .. ...... .. .4
A ctivation of Hydrogen Peroxide......................................................................6
U V A ctivation ....................... ........................ ...... ........... .....6..
Strong B ase A ctiv ation ............................................................... .....................7
Strong A cid A ctiv ation ............................................................... .....................8
Acyl Hydroperoxides.................. ...... .............................8
Iron(II) A ctivation ......................................... .......................... .. ........ .. .. .9
Transition-metal Organometallic Complexes..........................................10
M ethyltrioxorhenium ................................................. .. ......... .............. 11
A sy m m etric O x idation .................................................................. .. ...................... 12
Sharpless Oxidation of Allylic Alcohols................................................. 12
M n(III)-salen Epoxidation Catalysts .... .......... ........................................ 13
Chiral K etone Epoxidation Catalysts ...................................... ............... 15
P eroxycarbonate ................ .... .................................................... 16
Transition-metal Peroxycarbonate Complexes.......................... ............... 19
Transition-metal Activation of Peroxycarbonate in Solution.............................22
Scope of the D issertation ............................................................................. .. ... 23

2 OXIDATION OF NUCLEOPHILIC ALKENES IN AQUEOUS MICELLAR
M E D IA ........ .. ........ ......... .................................................... 2 5

Introduction .............. ...... .............. .................................. 25
R results and D discussion .................. .. .... ........................ .. ........... .. ....... ....28
Styrene Oxidation in Micellar Media in the Absence of Mn(II).........................28









Large Scale Styrene Oxidation............................................................. 28
Styrene Oxidation in Micellar Media in the Presence of Mn(II).........................31
R action K in etics................... .......................... .... ............ .......... ... ...... 34
Dependence of Styrene Oxidation on Surfactant Identity ...............................35
Dependence of Styrene Oxidation on the Manganese(II) Source .....................36
B icarbonate D ependence .......................................................... ... .............37
E x p erim en tal .................................................................................................... 3 8
M materials and Instrumentation ........... ...... ........... .... ............... 38
Standardization of Sodium Bicarbonate Solutions............................................39
Styrene Oxidation Reactions ................... .................... ............... 40
Large Scale Styrene Oxidations ........................................ ....... ............... 40
Synthesis of M n(D S)2 ..................................................... .. ........... .............. ... 41
Styrene Oxidation in SDS with Mn(II) and Mn(DS)2 .............. ............... 41

3 KINETIC INVESTIGATIONS OF THE MANGANESE(II) CATALYZED
DISPROPORTIONATION OF HYDROGEN PEROXIDE IN THE PRESENCE
OF BICARBONATE AND THE COMPARISON TO NUCLEOPHILIC
A L K EN E E PO X ID A TIO N ............................................................. ..................... 42

In tro d u ctio n ......................................................................................4 2
Results and D discussion ................. .... .... .......... .. ... .. ............. 44
Kinetics of Hydrogen Peroxide Decomposition...............................................44
M anganese(II) D ependence...................................... ........................ 47
B icarbonate D ependence......................................................... ..................48
Comparison of Hydrogen Peroxide Reaction Kinetics to Nucleophilic Alkene
Epoxidation Kinetics............... .... ......... ...... .. .... .. ............ 50
Manganese dependence on nucleophilic alkene epoxidation.....................50
Bicarbonate dependence on nucleophilic alkene epoxidation....................51
Hydrogen peroxide dependence on nucleophilic alkene epoxidation ..........52
Catalyst Lifetim e Studies ............................................................................. 53
Exam ining the loss of activity .................................................................54
Multiple additions of distilled hydrogen peroxide and solid sodium
bicarbonate................................................................... ......... 55
Studies of the M anganese Source............................................... .................. 56
Potassium perm anganate ................................. ....................................... 56
[M nv(M e3TACN)(OM e)3]PF6........................................................... 57
M n(IV ) catalyst stability ............................ .............. ............... .... 57
Cis-trans Isomerization in the Manganese(II) Catalyzed Alkene Epoxidation ...61
Cis/Trans isomerization reactions with cis-2-butene-l,4-diol ...................65
Cis/Trans isomerization of maleic and fumaric acids ..............................65
Examination of Sychev's Radical Trap Experiments ........................................68
Proposed M echanism of N-dealkylation .................................. ............... 81
Support for the Single Electron Transfer Pathway....... ...... ....... ...........85
S olv ent Isotop e E effect .............................................................. .....................86
Proposed Mechanism..................... .................................... 90
Numerical Simulation of the Proposed Mechanism ................. ................93
C o n clu sio n s.................................................... ................ 10 9









E xperim ental ............................................................................... 113
M materials and Instrum entation................................ ................ ............... 113
Standardization of sodium bicarbonate solutions..................................... 114
Hydrogen peroxide decomposition studies ...............................................115
Synthesis of [MnlV(Me3TACN)(OMe)3](PF6) ...................... ...............115
Oxidation of N,N-dimethyl-4-nitrosoaniline (DMNA) by Oxone ...........116
Oxidation ofN,N-Dimethyl-4-nitrosoaniline (DMNA) by H202/HCO3
/M n2+ .............................. ... ... ................................................. 1 16
Oxidation of N,N-diethyl-4-nitrosoaniline (DENA) by H202/HCO3-/Mn2+ 17

4 ELECTROPHILIC ALKENE EPOXIDATION BY THE
PEROXYCARBONATE DIANION..................................................................118

Introduction ...................................... ............................... ......... ...... 118
Results and D discussion ....................... ...... ........ ................ .... .......... 119
Effect of Mn(II) on Electrophilic Alkene Epoxidation ................................... 121
Effect of pH on the Oxidation of Electrophilic Alkenes .................................125
Effect of Buffer Choice on Electrophilic Alkene Epoxidation ......................130
Electrophilic Alkene Oxidation Kinetics .............. ...................... ...............131
Discussion of the Second-Order Rate Constants..............................................133
C o n clu sio n s....................................................... ................ 13 9
M materials and Instrum entation ......................................... ........................ 142
Experimental ............... .... .... ...................143
Electrophilic Alkene Epoxidation ........................................ ............... 143
D ibenzoylethylene K inetics ................................................... ................. 143

5 G EN ER A L CON CLU SION S .............................................................. ...............144

APPENDIX: VARIATIONS IN NUCLEOPHILIC ALKENE EPOXIDATION AND
HYDROGEN PEROXIDE RATE CONSTANTS ...............................................152

LIST OF REFEREN CE S ......... .................................. ........................ ............... 155

B IO G R A PH ICA L SK ETCH ......... ................. ...................................... .....................163















LIST OF TABLES


Table page

1-1 Some comm on reactive oxygen species................................. ....................... 3

2-1 Comparison of Styrene Oxidation in CTAC1 and SDS for the Mn(II) catalyzed
epoxidation. Reaction conditions: 0.05 M Styrene, 0.100 M CTAC1 or SDS,
0.25 M NH4HCO3, 1.00 M H202, and 10 [tM Mn(II). Errors are reported to the
95% confidence. .................................................... ................. 36

2-2 Comparison of observed rate constants for differing manganese sources for
micellar styrene oxidation. Reaction conditions: 0.05 M Styrene, 0.100 M SDS,
0.25 M NH4HCO3, 1.00 M H202, and 10 [M Mn(II) or Mn(DS)2. Errors are
reported to the 95% confidence ................ ..... ......... .....................37

3-1 Comparison of observed rate constants for the decomposition of hydrogen
peroxide, 0.100 M, in 0.20 M sodium bicarbonate with 3 and 4 [tM
manganese(II) and permanganate. Errors reported are to the 95% confidence.......56

3-2 Comparison of observed rate constants for the decomposition of hydrogen
peroxide (0.100 M, final concentration) in 0.20 M sodium bicarbonate with 3
and 4 [tM manganese(II) and [MnIV(Me3TACN)(OMe)3](PF6). Errors reported
are to the 95% confidence. .............................................. ............................. 61

3-3 Comparison of first-order rate constants for the epoxidation of sulfonated
styrene in H20 and D20. Reaction conditions: 0.001 M SS, 1.0 M Sodium
B icarbonate, 0.50 M M n(II) ....................................................... .....................86

3-4 Comparison of solvent isotope effect for hydrogen peroxide decomposition.
Reaction Conditions: 0.40 M HC03-, 0.10 M H202............... ................90

4-1 The percent conversion of 1 in varying buffer at differing pH..............................131

4-2 pKa values in water and CTAC1 for several different dyes.51............................139















LIST OF FIGURES


Figure p

1-1 The sulfate dianion .......... .... ........ .. .... ......... ....... ........... ..

1-2 The oxidation of organic molecules is defined as formation of bonds to carbon
with atoms that are more electronegative than carbon. Reduction is the loss of
bonds to more electronegative atoms and bond formation with hydrogen. .............3

1-3 Superoxide dismutase enzymatically oxidizes the superoxide anion and two
protons to hydrogen peroxide, another reactive oxygen species. Hydrogen
peroxide is the disproportionate by catalase to yield water and molecular
oxygen ............................................................................... 4

1-4 The AO-process for the industrial production of hydrogen peroxide. ......................5

1-5 Illustration of a nucleophilic attack on hydrogen peroxide. The use of a general
acid facilitates the proton transfer to yield the oxidized nucleophile and water........6

1-6 The reactivity of olefins with hydroxyl radicals.................................... 7

1-7 Polymerization of olefins by hydroxyl radical. ...................................... .......... 7

1-8 Reactivity of electrophilic olefins with nucleophilic oxidants, such as
hydroperoxide, react to produce the epoxide plus the oxidants' corresponding
leaving group, in this case hydroxide. ......................... .......... ............. .................. 7

1-9 The reaction of an alkene with OH+ generates an intermediate carbocation. A
general base can then deprotonate the oxygen of the intermediate which results
in ring closure to form the epoxide. ........................................ ....................... 8

1-10 A lkene oxidation by m -CPB A ......................................................... ............... 9

1-11 Activation of iron(III) tetrakis(pentafluorophenyl) porphyrin by hydrogen
peroxide to produce a high oxidation state iron complex. ..............................11

1-12 The two dominant forms in the MTO/H202 system under acid conditions. The
diperoxorhenium adduct reacts slightly slower than the monoperoxorhenium
co m p lex .23 .................................................................................12









1-13 Nucleophilic attack of an olefin on the electrophilic oxygen of the hydrogen
peroxide activated methyltrioxorhenium yields the oxidized nucleophile and
regenerates MTO. Attack of a nucleophile on the diperoxo complex generates
the oxidized nucleophile and the monoperoxorhenium complex.24 ......................12

1-14 Illustration of the asymmetric epoxidation using the Sharpless method. Use of
the (+) or (-)-tartrate allows for the oxygen atom to be added to only one face of
th e ally lic alc o h o l.25............................................................................................ 13

1-15 A salen ligand ........................................................................................................ 13

1-16 spiro[2H-1-benzopyran-2,1'-cyclohexane]......... ...... ... ...... ..... .......... 14

1-17 Asymmetric epoxidation of alkenes can be easily achieved using
peroxymonosulfate to generate a dioxirane in situ.30 ...............................................15

1-18 Structure of 1,2:4,5-di-O-isopropylidene-D-erythro-2,3-hexodiuro-2,6-pyranose
used by Shi30 for the asymmetric epoxidation of alkenes using
peroxymonosulfate to generate a dioxirane in situ ...............................................16

1-19 The equilibrium formation of bicarbonate and peroxycarbonate proceeds
through CO2 as an intermediate.34........... ............................... 17

1-20 Fe(qn)2(02C(0)O]Ph4P 1.5MeOH-0.5 (CH3)2NCHO.38 .........................................18

1-21 Nucleophilic attack on the peroxycarbonate anion. An intramolecular proton
transfer in the transition state allows for release of bicarbonate instead of
hydroxide as in the case of hydrogen peroxide ................................. ............... 19

1-22 Generation of a metal peroxycarbonate (LnM(CO4)Xm) from its parent 02
complex, LnM(O2)Xm, by passing CO2 through a dry solution of the parent
c o m p le x .4 2 .......................................................................................................... 1 9

1-23 Structure of the (Ph3P)2Pt(C04) complex of Nyman.45.........................................20

1-24 Routes for the oxidation of PR3 by (PEt2Ph)3RhCl(C04).43 Route A shows the
solution chemistry where a solvent molecule displaces a phosphine before it is
oxidized. Route B shows the solid state chemistry where coordination of
ethylene occurs first with the displacement of a phosphine ligand followed by
oxidation of the ligand...... .... ........................................... ........................... ..... 2 1

1-25 Structure of products of styrene oxidation by [(PEt2Ph)3RhCl(C04)] under a
CO2/02 atmosphere that indicate a radical mechanism.4 ....................................22

1-26 The structure of diethylenetriaminepentaacetic acid (DTPA)..............................22









2-1 The structures of three common surfactants. Cetyltrimethylammonium chloride
is a cationic surfactant, while sodium dodecylsulfate is anionic. Triton X-100 is
a non-ionic surfactant. ........................ .............. ................... ......... 26

2-2 The structure of a micelle with a concentration greater then the cmc....................27

2-3 The graphical representation of an alkene dissolved in a micelle............................27

2-4 The reaction scheme for the oxidation of styrene by hydrogen peroxide in the
presence of bicarbonate and cetyltrimethylammonium chloride (CTAC1) without
the presence of Mn(II). Hydrolysis of the product epoxide forms the
corresponding diol. Reaction conditions: 0.05 M Styrene, 0.10 M CTAC1, 2.00
M H202, 1.00 M NH4HCO3, 3 days ............................................... 28

2-5 A picture of a lighter-than-water liquid-liquid extractor .................... ...............30

2-6 Reaction scheme used by Burgess40 in the mixed solvent epoxidation of styrene. .32

2-7 Schematic representation for the oxidation of styrene in surfactant with
hydrogen peroxide and bicarbonate catalyzed by manganese(II). Reaction
conditions: 50 mM styrene, 0.10 M CTAC1, 2.00 M H202, and 1.00 M
N H 4H C O 3, 30 m minutes. ................................................................. .....................32

2-8 HPLC chromatograms for the initial reaction (top panel) and after 30 minutes
(bottom panel) for the oxidation of styrene with H202, HC03-, and Mn(II) in the
presence of surfactant (CTAC1). HPLC performed using a C18 reverse phase
column using a non-linear gradient for 12 minutes. Mobile Phase: 25%:75%
(v:v) CH3CN:H20 95% :5% CH3CN:H20 ........... ..............................................33

2-9 Styrene area disappearance versus time from the HPLC analysis of styrene
oxidation by hydrogen peroxide in micellar media in the presence of bicarbonate
and Mn(II). Reaction conditions: 0.05 M Styrene, 0.100 M CTAC1, 0.25 M
NH4HCO3, 1.00 M H202, 10 [tM Mn(II). ................................ ............... 34

2-10 In(styrene area) versus time to find the first-order rate constant. The line is the
linear regression to the data at the 95% confidence. The kobs is the negative
slope of the line. ........................................................................35

2-11 Structure of manganese(II) bisdodecylsulfate .......................................................36

2-12 Graph of kobs versus [NH4HCO3] for the styrene oxidation in the presence of
0.100 M CTAC1. Reaction conditions: 0.05 M Styrene, 0.100 M CTAC1, 1.00
M H202, and 10 M M n(II). .......................................................................38

3-1 Hydrogen peroxide decomposition in the presence of manganese(II) and
bicarbonate. .......................................... ............................ 45









3-2 Plot of the ln([H202]) versus time for varying bicarbonate concentration. The
0.20 and 0.30 M bicarbonate reactions are typical of the accelerations noticed
for these reactions. Reaction conditions: 0.10 M H202, 4 M Mn(II) ....................46

3-3 The dependence of kobs on the [Mn(II)]. Reaction conditions: 0.10 M H202, 0.4
M HC03-, varying [Mn(II)]. y = ((7.98 0.62) x10-4), error reported to the
95% confidence. ........................... ...... ... ... .. ...... .......... .... 47

3-4 Plot of kobs versus [NaHCO3]2. Reaction conditions: 0.10 M H202 and 4 tLM
Mn(II). y = ((2.08 0.25) x10-3)x, error reported to the 95% confidence ............49

3-5 Plot of kobs versus [Mn(II)] observed for nucleophilic alkene epoxidation
(Bennett, 2002)46 y = ((2.09 0.25) x10-3)x, error reported to the 95%
confi dence. ....................................................... .................. 50

3-6 Plot of kob versus [HCO3-]2 which shows a second-order dependence. Reaction
conditions: 0.001 Mp-vinyl benzene sulfonate, 1.00 [tM Mn2+ (+) 0.10 M H202
y = ((2.62 + 0.17) x10-3)x (m) 0.50 M H202y = ((1.19 0.23) x10-3)x (A) 0.75
M H202 y = ((8.33 0.76) x104)X, errors reported to the 95% confidence.
(Bennett, 2002)46 .................................................................. .. ......... 51

3-7 Plot of kobs on the [H202]. Reaction conditions: 0.001 M p-vinyl benzene
sulfonate (A) 1.00 M NaHCO3, 0.50 [tM Mn2+ (m) 0.75 M NaHCO3, 0.50 [tM
Mn2+ (+) 1.00 M NaHC03, trace metal catalysis (Bennett, 2002).46........................52

3-8 Plot of kobs versus # of additions of hydrogen peroxide to a spent solution in the
catalyst lifetime study over multiple days. There is a 16 hr delay before addition
16 an d 2 4 ........................................................ ................. 5 3

3-9 Plot of kobs versus # of hydrogen peroxide additions for the catalyst lifetime
study using distilled hydrogen peroxide and adding solid sodium bicarbonate.
The loss of activity is now due only to dilution and the inability to maintain the
bicarbonate concentration at a constant value....... ... ........................................ 55

3-10 Molecular structure of [MnlV(Me3TACN)(OMe)3](PF6).63........... ......................57

3-11 Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of
1.00 sodium bicarbonate. ............................................. .............................. 58

3-12 Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence of
0.50 M hydrogen peroxide. ......................................................... ......................58

3-13 Monitoring of 0.108 M [MnlV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of
25 p.L (0.100 M, final concentration) hydrogen peroxide was done at 350
seconds. The absorbance first decays to 0 and within a matter of minutes, the
solution is bright yellow ............................... .............. .............. ............. 59









3-14 UV-vis specta of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6). The solid line is the
spectrum in the presence of 1.00 M sodium bicarbonate. The dotted line is the
spectrum of the solution after 1 eq of hydrogen peroxide was added....................60

3-15 Monitoring of 0.108 M [MnlV(Me3TACN)(OMe)3](PF6) at 345 nm. Addition of
50 p.L (0.200 M, final concentration) hydrogen peroxide was done at 312
seconds. Even after 6 minutes, the yellow color does not develop.......................60

3-16 The concerted mechanism for the m-CPBA oxidation of nucleophilic alkenes
resulting in the retention of stereochem istry. ...........................................................62

3-17 The stepwise oxidation of an alkene by Mn(salen) and hydrogen peroxide is
shown. Cis/trans isomerization occurs in the transition state, where the C-C
sigma bond is able to rotate into the more stable trans configuration..................62

3-18 The cis/trans isomerization noted Burgess in his epoxidation of stilbene using
the Mn(II), H202, bicarbonate system using a mixed solvent system of
DM F/H20. (Burgess, 2002)40................ ..................... ... ...................................63

3-19 Synthetic scheme for synthesis of 4,4'-sulfonated stilbene. (van Es, 1964)64 .........63

3-20 13C NMR of cis and trans-2,3-epoxybutane-l,4-diol in D20 using methanol as
an internal standard. ..................... .................... ................. ..........64

3-21 Epoxidation of cis-2-butene-1,4-diol (0.60 M) with 1.00 M HC03-, Mn(II) (10
[tM), and H202 (6.00 M) after 30 minutes (left) and 18 hrs (right) .......................65

3-22 The structures of maleic and fumaric acids at the operating pH of 8.4.................66

3-23 1H NMR of maleic and fumaric acid oxides in D20 using methanol as an
internal standard. ......................................................................66

3-24 Epoxidation of maleic acid by hydrogen peroxide and manganese(II) in the
presence of bicarbonate after 15 min. Reaction conditions: 0.10 M maleic acid,
1.00 M H202, 0.80 M NaHCO3, and 10 pM Mn(II). .............................................67

3-25 Fenton's reagent can be used to oxidize benzene to phenol and biphenyl .............69

3-26 The influence of inhibitors on the catalase process in the Mn(II)/HCO3/H202
system. [Mn(II)] = 4 x 10-6M, [H202] = 0.10 M, pH 7.0, [HCO3-] = 0.4 M, and T
= 25 C: 0) kinetic curve with no inhibitors; 1), 2), 3), and 6) in the presence of
DMNA as the inhibitor(at concentrations of 1 x 10-5,1.5 x 10-5, 2 x 105,and 4 x
10-5 M respectively; 4)in the presence of tetranitromethane(4 x 10-5 M); 5) in the
presence of hydroquinone (1.5 x 10-5 M); 7) decomposition of H202 without
M n(II) ion (blank experiment). (Sychev, 1977)47 ........................ ......... ...... 70

3-27 The reaction ofN,N-dimethyl-4-nitrosoaniline with peroxymonosulfate to yield
N ,N -dim ethyl-4-nitroaniline. .......................................................... .....................72









3-28 1H NMR of the crude reaction mixture after an oxidation of NN-dimethyl-4-
nitrosoaniline by hydrogen peroxide in the presence of bicarbonate and Mn(II).
Reaction conditions: N,N-dimehtyl-4-nitrosoaniline (1 g, 6.66 mmol), 0.400 M
sodium bicarbonate, 10 [tM Mn(II), 6.64 M H202, 1 hr. ............................74

3-29 GC trace for a standard ofN,N-dimethyl-4-nitrsoaniline. Non-linear gradient for
30 minutes, detection by FID ..... ......................................................................75

3-30 GC trace for the crude reaction material from the oxidation ofN,N-dimethyl-4-
nitrsoaniline from Figure 3-25. Lack of a peak near 14.637 min proves that no
starting material remains. GC conditions: non-linear gradient for 30 minutes,
Detection: FID ............................ .... ............................... ..........76

3-31 GC trace (left figure) and 1H NMR (right figure) for Fraction 4 of the silica
column. Identification of the product as N,N-dimethyl-4-nitroaniline was
confirmed by comparison with a GC trace and 1H NMR of an authentic sample. ..77

3-32 GC trace (left figure) and 1H NMR (right figure) for Fraction 9 of the silica
column. Identification of the product as 4-nitroaniline was confirmed by
comparison with a GC trace and 1H NMR of an authentic sample ..........................78

3-33 H NMR of an authentic sample of N-methyl-4-nitroaniline. Comparison with
the crude reaction mixture confirms its presence as a product. ............................79

3-34 The solution collected from the reaction ofN,N-diethyl-4-nitrosoaniline was
analyzed by Gas Chromatography (lower figure), which was compared to an
authentic sample of acetaldehyde. 1H NMR (top figure) of the solution also
confirmed that the product was acetaldehyde. ................................. ............... 80

3-35 The proposed mechanism for the oxidative N-dealkylation of amines by
hydrogen peroxide in the presence of bicarbonate as catalyzed by
manganese(II). The secondary amine produced can cycle again as long as it
contains a hydrogen on the carbon a to the nitrogen. A second molecule of
aldehyde or ketone will also be produced. .................................... ............... 83

3-36 The structure of N,N-dimethyl-2-amino-2-methyl-3-phenylpropane, the substrate
used by Miwa et al.73'74 for use in experiments with cytochrome P450 on the
oxidative N-dealkylation mechanism ........................................... ............... 85

3-37 Mechanism of epoxide hydrolysis in acidic media. ............................................87

3-38 Possible epoxidation routes through a manganese(IV) oxo complex. None of the
envisioned reactions has a proton transfer. ................................... ............... 88

3-39 Mechanism of oxygen transfer by attack of the alkene on a manganese(II) bound
peroxycarbonate. The proton transfer in the transition state may account for the
inverse isotope effect observed. ........................................ ......................... 89









3-40 Oxidation of a nucleophilic alkene by two sequential reactions with the
carbonate radical. The carbocation intermediate formed explains the loss of
retention observed for cis-alkenes ........................................ ........................ 90

3-41 Proposed mechanism for hydrogen peroxide decomposition and nucleophilic
alkene epoxidation in the presence of bicarbonate catalyzed by Mn(II)..................91

3-42 Proposed generation of the active manganese catalyst by from the
[Mn"(HCO3)2(HCO4)]- complex by a 2 electron oxidation of manganese to form
a high valent [M n-O2-]2+ complex. ........................................ ........................ 92

3-43 Simulation of the dependence on the concentration of the active catalyst with
varying bicarbonate. Simulation conditions: 1.00 M hydrogen peroxide and 4
[M Mn(II). y = ((6.55 0.37)x10-7)x, error reported to the 95% confidence. ........96

3-44 Plot of simulation results for [Mn(HCO3)+ versus [HCO3-]. The [Mn(HCO3)+
quickly saturates due to the large equilibrium constant of 19.05. Simulation
conditions: 1.00 M H202, 4 tM M n(II). ...................................... ............... 97

3-45 Simulation of the dependence on the concentration of the active catalyst with
varying [HCO3]. Simulation Conditions: 1.00 M hydrogen peroxide and 4 [LM
Mn(II). y = ((4.60 + 0.22)x10-7)x, error reported to the 95% confidence ...............98

3-46 The generated curve for the hydrogen peroxide dependence on nucleophilic
alkene oxidation. The points represent the observed kobs experimentally
determined by Bennett.46 The line is the simulated kobs at each H202
concentration. Reaction and simulation conditions: 0.5 [MM Mn(II), 1.00 M
bicarbonate, 0.001 M Sulfonated Styrene (SS). ............. ......................... .......... 99

3-47 The generated curve for the bicarbonate dependence on nucleophilic alkene
oxidation. The points represent the observed kobs experimentally determined by
Bennett.46 The line is the simulated kobs at each [HC03] concentration. Reaction
and Simulation Conditions: 0.5 [MM Mn(II), 0.10 M hydrogen peroxide, 0.001 M
Sulfonated Styrene (SS). ............................................. ............................... 100

3-48 The generated curve for the manganese dependence on nucleophilic alkene
oxidation. The points represent the observed kobs experimentally determined by
Bennett.46 The line is the simulated kobs at each [Mn(II)] concentration.
Reaction and simulation conditions: 1.00 M bicarbonate, 0.55 M hydrogen
peroxide, 0.001 M Sulfonated Styrene.................. ..............100

3-49 A typical numerical simulation plot attempting to model the hydrogen peroxide
decay curves. Points represent observed ln[H202] versus time, while the line is
the simulated ln([H202]) versus time. Reaction and simulation conditions: 0.10
M H202, 0.30 M HC03-, 4.0 [M Mn(II)....... .... ........................102









3-50 Simulation of hydrogen peroxide decay at lower bicarbonate concentration.
Points represent data, while the line is the simulated decay. Reaction and
simulation conditions: 0.10 M H202, 0.10 M HCO3, 4 |tM Mn(II). ..................106

3-51 A plot of [HCO3]2 versus "kobs" for the hydrogen peroxide decomposition.
Reaction and simulation conditions: 0.10 M H202, 4 jtM Mn(II)..........................107

3-51 A plot of "kobs" versus [Mn(II)] for the hydrogen peroxide decomposition.
Reaction and simulation conditions: 0.10 M H202, 0.40 M HCO3 ........................107

3-53 Simulated hydrogen peroxide dependence for the hydrogen peroxide
decomposition reactions. Simulation conditions: 0.90 M HCO3, 0.5 [tM Mn(II). 108

3-54 Simulated hydrogen peroxide dependence for the hydrogen peroxide
decomposition reactions. Simulation conditions: 0.40 M HCO3, 3.0 [LM Mn(II). 108

3-55 Plot of ln([H202]) versus time. The points represent observed data and the line is
the simulation. Reaction and simulation conditions:0.10 M H202, 0.40 M HCO3,
3.0 [tM Mn(II). As the plot indicates the reaction accelerates as the
decom position occurs .................. ............................. .... .... .. ............ 109

4-1 The resonance structure of an electrophilic alkene, an a,P-unsaturated ketone,
explains the reactivity with nucleophilic oxidants. The 0-carbon of the alkene,
as seen in the resonance structure, is more electropositive and will be the site of
attack by a nucleophilic oxidant .................................................. ......... ...... 120

4-2 The mechanism of electrophilic alkene oxidation by the hydroperoxide anion is
illustrated. The nucleophilic oxidant adds at the electrophilic carbon, the P-
carbon. Reformation of the ketone moiety causes either the displacement of the
hydroperoxide anion, regenerating the starting alkene and hydroperoxide, or ring
closure to form the epoxide and the hydroxide anion. ........................................ 120

4-3 The mechanism of electrophilic alkene epoxidation by the peroxycarbonate
dianion is illustrated. The mechanism is identical to that of hydroperoxide
oxidation, except that the nucleofuge of the peroxycarbonate dianion is the
carbonate dianion. ........................... ..... ............. .. .............. 120

4-4 Electrophilic alkenes used in this study. ..................................... ............... 121

4-5 1H-NMR of 1 epoxidized by H202 at pH 7.8 at 60 min, in the presence of 1.00
M sodium bicarbonate (50% conversion). .................................. ............... 121

4-6 H-NMR of 1 epoxidized by H202 at pH 7.8 with 4 [tM Mn(II) in the presence
of 1.00 M sodium bicarbonate at 60 min (44% conversion).............................122

4-7 H-NMR of 1 epoxidized by H202 at pH 7.8 with 5 mM DTPA in the presence
of 1.00 M sodium bicarbonate at 60 min (50% conversion). .............................123









4-8 1H-NMR of 1 epoxidized by H202 in the presence of 1.00 M sodium bicarbonate
at pH 8.6 after 15 min (50% conversion). .................................... ...............124

4-9 1H-NMR of 2 epoxidized by H202 in the presence of 1.00 M sodium bicarbonate
at pH 7.8 in 24 hrs (75% conversion).......... ................................................. 125

4-10 1H-NMR of 2 epoxidized by H202 in the presence of 1.00 M sodium bicarbonate
at pH 8.6 at 24 hrs (50% conversion) ........ ................................................ 126

4-11 Speciation of methacrylic acid (MAA) at 0.05 M and hydrogen peroxide (0.30
M) as a function of pH. The maximum concentration of MAA and
hydroperoxide happens to occur at pH 7.8, the buffering pH of sodium
b icarb on ate. ...................................................... ................. 12 7

4-12 1H-NMR of 2 epoxidized by H202 in the presence of 1.00 M sodium bicarbonate
at pH 7.8 with 4 [tM Mn(II) in 15 min (>90% conversion). ............................128

4-13 1H-NMR of 3 epoxidized by H202 in 1.00 M sodium bicarbonate at pH 7.8 in 24
hrs (44% conversion). ........................ ...... ................ ............... .... .......... 129

4-14 1H-NMR of 3 epoxidized by H202 in 1.00 M sodium bicarbonate at pH 8.6 in 24
hrs (66% conversion). ........................ ...... ................ ............... .... .......... 130

4-15 Plot of log(kobs) vs pH for the oxidation of 4 by H202 ........................................ 132

4-16 Plot of log(kobs) vs pH for the oxidation of 4 by OC1. ...................................... 133

4-17 Reaction mechanism determined by Rosenblatt for the nucleophilic oxidation of
electrophilic alkenes.31 ..................................... ...... .. ........... ............ 133

4-18 Structure of the substrate used by Rosenblatt and Broome,131 o-
chlorobenzylidenem alononitrile .................................................. ........... .... 134

4-19 A reaction coordinate diagram for the addition of a nucleophilic oxidant to an
electrophilic alkene. If the ring closure of the epoxide is the rate determining
step, as seen in the diagram by a larger energy barrier for the formation of the
epoxide, the identity of the nucleofuge (Z) will determine the reactivity of the
oxidant. The more stable the nucleofuge, the more reactive the oxidant will be..135

4-20 A reaction coordinate diagram for the addition of a nucleophilic oxidant to an
electrophilic alkene. If the addition of the oxidant is the rate determining step,
as seen in the diagram by a larger energy barrier to the formation of the
intermediate, the identity of the nucleofuge (Z) will no longer determine the
reactivity of the oxidant. The more basic oxidant will react faster.....................136

4-21 Substrates used by Bunton and Minkoff to study the oxidation of electrophilic
alkenes by the hydroperoxide anion ................... ................. .............................137


xviii









A-i Variation in the S + A -* Products rate constant. ..................................................152

A-2 Variation in the equilibrium constant for A + H202 B ..................................153

A-3 Variation in the equilibrium constant for the formation of"A"..........................153

A-4 Variation in the rate constant for A -* radicals. .................................................154















Abstract of Dissertation Presented to the Graduate School
of the University of Florida in Partial Fulfillment of the
Requirements for the Degree of Doctor of Philosophy

HYDROGEN PEROXIDE DISPROPORTIONATION AND ORGANIC COMPOUND
OXIDATION BY PEROXYCARBONATE CATALYZED BY MANGANESE(II):
KINETICS AND MECHANISM

By

Andrew P. Burke

August 2005

Chair: David E. Richardson
Major Department: Chemistry

The investigation of the mechanism of hydrogen peroxide disproportionation and

alkene epoxidation in aqueous solutions of bicarbonate at near neutral pH as catalyzed by

manganese(II) is described. Current literature proposes that a free hydroxyl radical

pathway based on Fenton chemistry is responsible for the hydrogen peroxide decay. This

proposed mechanism does not adequately explain the unique requirement of bicarbonate

in these reactions. Also, the proposed free hydroxyl radical mechanism does not explain

why no radically coupled products in the oxidation of nucleophilic alkenes are detected.

We suggest that manganese(II) is activated by peroxycarbonate, a hydrogen peroxide and

bicarbonate adduct, to form a high oxidation state manganese(IV) complex. In addition,

it is proposed that the carbonate radical anion is also a product of the reaction of

peroxycarbonate in the presence of metal cations. Both the carbonate radical anion and

the high oxidation state manganese(IV) complex are believed to be the main reactive









oxygen donors in this system responsible for the observed reactivity. Numerical

simulations of the hydrogen peroxide and nucleophilic alkene epoxidation by hydrogen

peroxide in solutions of bicarbonate and manganese(II) have also been conducted.

While the epoxidation of nucleophilic alkenes by hydrogen peroxide in bicarbonate

solutions is catalyzed by manganese(II), the same is not true for the epoxidation of

electrophilic alkenes. Investigations have been conducted using several water soluble

alkenes. For these reactions, the addition of manganese(II) has been shown to inhibit the

oxidation by decomposing the active oxidant, the hydroperoxide anion. Kinetic

investigations of the oxidation of dibenzoylethylene in micellar media by hydroperoxide

and hypochlorite will also be presented. The observed second-order rate constant for the

oxidation by hydroperoxide is 660 40 M-^s^- and that for hypochlorite is 118 2 M^s^-.















CHAPTER 1
INTRODUCTION

General Oxidation

Oxidation-reduction redoxx) reactions are some of the most important chemical

reactions. These reactions are responsible for the formation of compounds from their

elements, the generation of electricity, and combustion reactions, some of which produce

energy at the cellular level. Redox reactions are always coupled, and the number of

electrons transferred must be equal in number between the oxidation and reduction half-

reactions. Redox reactions can be easily determined by identifying the oxidation states of

the atoms in the ions and molecules involved in the reaction. Lewis structures provide an

easy convention by which oxidation states may be assigned to atoms. Typically, all

bonds must be assumed to be completely ionic, and the more electronegative atom of the

bonded pair is allocated the pair of electrons. For example, consider the sulfate dianion,

S042-




.'*'o*
22-



*O S 2*



.0.

Figure 1-1. The sulfate dianion

The sulfur-oxygen bonds are polar covalent, polarized toward the oxygen atoms.

Each oxygen atom is given an oxidation number of -2, eight valence electrons versus six









for the free atom. The sulfur atom is given an oxidation number of +6, zero valence

electrons versus six for the free atom. The charge of the ion is given by the sum of all the

formal oxidation charges, in this case 6-8 = -2.

The use of oxidation states can now be used to easily identify redox reactions. As

an example, consider the reaction of sulfite and permanganate anions in acidic solution to

yield the sulfate anion and manganese(II).

5S032- + 2MnO4- + 6H+ D 5S042- + 2Mn2+ + 3H20 (1-1)

In this example, the sulfur atom of sulfite begins in the +4 oxidation state and is in

the +6 oxidation state on the product side, as is seen in the Equation 1-1, a loss of two

electrons. By the definition of oxidation and reduction, this is an oxidation. By

definition, oxidation reactions must be coupled to a reduction reaction, the permanganate

must be gaining electrons. In the example above, this is seen to be true since the reactant

manganese of permanganate is in the +7 oxidation state and a gain of 5 electrons yields

manganese(II), as seen on the product side of the reaction.

In organic chemistry, the assignment of oxidation states is not as simple as the

above example with the sulfate dianion.1'2 It has been the traditional method, therefore,

to define oxidation in organic chemistry as the "loss" of electrons by forming bonds with

elements that are more electronegative than carbon, such as oxygen or nitrogen.

Reduction then is the "gain" of electrons by breaking bonds with more electronegative

atoms and forming bonds with hydrogen.1,2 For example, ethanol can be transformed to

form acetaldehyde. In this process, ethanol looses a bond to hydrogen and gains a bond

to oxygen, an oxidation. Similarly, acetic acid can be converted to acetaldehyde. In this

process, acetic acid looses a bond to oxygen and gains a bond to hydrogen, a reduction.









Oxidized 0
OHH


0 0
0 Reduced 0

OH H
Figure 1-2. The oxidation of organic molecules is defined as formation of bonds to
carbon with atoms that are more electronegative than carbon. Reduction is the
loss of bonds to more electronegative atoms and bond formation with
hydrogen.

Reactive Oxygen Species

Molecular oxygen and reactive oxygen species (ROS) are the main oxygen sources

for oxidation processes and are highly reactive oxygen donor molecules with the ability

to react with a wide variety of substrates.3 Table 1-1 lists some of the most commonly

encountered radical and non-radical reactive oxygen species.

Table 1-1. Some common reactive oxygen species
Radical Non-radical
Superoxide, 02- Hydrogen Peroxide, H202
Hydroxyl, OH- Hypochlorous acid, HOCl
Peroxyl, ROO- Alkyl Hydroperoxide, ROOH
Alkoxyl, RO-
Hydroperoxyl, HOO-

In the human body, for example, the effects of reactive oxygen species have been

measured by examining the oxidative stress on cells.4 Oxidative stress is defined as the

imbalance between the cellular production of reactive oxygen species and the antioxidant

mechanisms in existence to remove them.4 The effects of these reactive oxygen species

have been linked to chronic disease and aging.5'6

The human body has several mechanisms, including enzymes and radical

scavengers, that can intervene with reactive oxygen species.4 For example, superoxide









dismutase converts the superoxide radical anion plus two protons to hydrogen peroxide.

The product hydrogen peroxide, which is yet another reactive oxygen species, can then

be enzymatically decomposed by catalase to water and oxygen. The human body also

uses a number of radical scavengers, such as ascorbate, urate, and tocopherol, to rid cells

of high concentrations of reactive oxygen species.

2H
02' 2 H202
Superoxide Dismutase

2 H202 Catalase 2 H20 + 02
Figure 1-3. Superoxide dismutase enzymatically oxidizes the superoxide anion and two
protons to hydrogen peroxide, another reactive oxygen species. Hydrogen
peroxide is the disproportionate by catalase to yield water and molecular
oxygen.

Hydrogen Peroxide

Hydrogen peroxide (H202) is a common reactive oxygen species which is

environmentally friendly due to its decomposition to molecular oxygen and water. It is a

weak, nonspecific, electrophilic oxidant with an E = 1.77 V vs. NHE7 that has been used

as a bleaching agent for over a century.8 Hydrogen peroxide is only a weak oxidant

under mild conditions. Recent interest in the use of H202 as a terminal oxidant has come

from increasing pressure in the industrial sector to find more environmentally friendly

oxidation reagents.8 Many industries are beginning to use hydrogen peroxide in the

treatment of wastewater. Recently, hydrogen peroxide was shown to remove cyanide

from thermoelectric power station wastewater.9

Hydrogen peroxide is also an important commercial chemical in the production of

epoxides. In this case, hydrogen peroxide is used to generate peracids that are then used

in the epoxidation of numerous alkenes.10 The activation of hydrogen peroxide to form

peracids will be presented shortly in the introduction.









Hydrogen peroxide is produced commercially by the AO-Process,10 which involves

the hydrogenation of a 2-alkyl-9,10-anthraquinone to the corresponding hydroquinone.

The hydroquinone produced is then oxidized with oxygen, or air, to regenerate the

anthraquinone and produce hydrogen peroxide (Figure 1-4). The hydrogen peroxide

produced is extracted with water, while the organic components can be recycled back

through the hydrogenation step.

H2
H2 Catalyst



O OH

R R






O OH





H202 02
Figure 1-4. The AO-process for the industrial production of hydrogen peroxide.

Hydrogen peroxide is an excellent environmental choice for two reasons. First, the

decomposition products are molecular oxygen and water. Second, due to its relative

inactivity, specific methods of activation must be used which can tune the reactivity to

the particular oxidative process required. Figure 1-5 illustrates the nucleophilic attack of

a substrate on hydrogen peroxide. A general acid can act as a proton transfer agent to

assist in the cleaving of the peroxide bond to form the oxidized nucleophile and water.









H
0 o H

Nu: Nu H
+HA


NuOH + OH H H



Nud H


NuO + H20
Figure 1-5. Illustration of a nucleophilic attack on hydrogen peroxide. The use of a
general acid facilitates the proton transfer to yield the oxidized nucleophile
and water.

Activation of Hydrogen Peroxide

UV Activation

While hydrogen peroxide may be used for some types of oxidations, activation is

required for use in a wider variety of reactions. For instance, solutions of hydrogen

peroxide can be irradiated using UV radiation to homolytically cleave the peroxide bond

to form two hydroxyl radicals, Equation 1-2.

hv
H202 2 HO- (1-2)

The hydroxyl radical is a potent, nonspecific, one-electron oxidant that can readily

react with alkenes (Figure 1-6) by addition to the double bond.11 The resulting organic

radical can then react with another hydroxyl radical to form the diol, or in the presence of

iron(II) and acid, the alcohol. Depending on the nature of the double bond, radical

polymerization can also occur, as seen in Figure 1-7.









H

HO--R'


H HO* HO R"
H R'
HH

HO- RR'
HO*+

H R" H R" Re H
H +- HO- R'

H R"

H
Figure 1-6. The reactivity of olefins with hydroxyl radicals.1

HO* HO
R -R

R
HO R R HO -- -- Polymer

Figure 1-7. Polymerization of olefins by hydroxyl radical.1

Strong Base Activation

Other methods for the activation of hydrogen peroxide are known. The reaction of

hydrogen peroxide with a strong base generates the hydroperoxide anion, as seen in

Equation 1-3, which is an effective nucleophilic oxidant. The hydroperoxide anion can

epoxidize an electrophilic alkene as seen in Figure 1-8.

H202 + OH OOH + H20 (1-3)


O
o R' O- 0 R'
SHOO- R I OOH R R' + OH-
R R" R R" R R"
Figure 1-8. Reactivity of electrophilic olefins with nucleophilic oxidants, such as
hydroperoxide, react to produce the epoxide plus the oxidants' corresponding
leaving group, in this case hydroxide.









Strong Acid Activation

Hydrogen peroxide is also activated by strong acids. Protonation of one of the

oxygens in hydrogen peroxide results in polarization of the 0-0 bond to generate OH+, a

strong electrophilic oxidant that can react with nucleophiles, such as alkenes. Water is

the other product of the reaction (Equation 1-4).

H+
H202 HO + + H20 (1-4)

0
OH+ + A- + HA

R R
Figure 1-9. The reaction of an alkene with OH+ generates an intermediate carbocation. A
general base can then deprotonate the oxygen of the intermediate which
results in ring closure to form the epoxide.

Acyl Hydroperoxides

Acyl hydroperoxides, a broad category of oxidants including the organic peracids,

are electrophilic oxidants often used for the heterolytic oxidation of organic substrates.

Peracids are synthesized from the acid-catalyzed equilibrium between hydrogen peroxide

and the acid form of the peracid, as seen in Equation 1-5.12 In the absence of a catalyst,

the equilibrium is slow. In order to isolate the peracid from the equilibrium mixture,

continuous distillation or an extraction step must be used.

H
RCO2H + H202 RCO3H + H20 (1-5)

A common example of a peracid used in organic oxidations is m-

chloroperoxybenzoic acid (m-CPBA). The mechanism of the reaction of an organic

nucleophile, an alkene, proceeds as shown in Figure 1-10. This example uses m-CPBA

as the oxidant. The rate of peracid epoxidation of alkenes is influenced by three main

factors. First, the reaction is dependent on the type of double bond. Second, the









substituents of the peracid affect its ability to oxidize an alkene. Third, the rate of

reaction is reduced in the presence of coordinating solvents, such as ethers, which form

intermolecular H-bonds.13 The kinetic aspects of peracid oxidations are as follows: 1) the

reaction is second order, 2) the reaction is stereospecific, meaning that cis-alkenes will

react to give cis-epoxides and trans-alkenes will react to give trans-epoxides, and 3) the

rate of reaction is increased with increasing strength of the formed acid.13

o
R O 0 C
R H -
R L0

R -R + HO C1
R' R


Figure 1-10. Alkene oxidation by m-CPBA

Iron(II) Activation

One of the best known and most studied processes for the activation of hydrogen

peroxide is by iron(II) salts, the combination of which is known as Fenton's reagent. The

most accepted mechanism for the activation, introduced by Haber and Weiss14'15 and

studied extensively by Barb et al.,16-19 involves the redox cycle ofiron(II). The

mechanism is shown in Equations (1-6)-(1-10).

Fe2+ + H202 Fe3+ + OH + OH- (1-6)

Fe2++ OH -Fe3+ + OH- (1-7)

'OH + H202 HOO' + H20 (1-8)

Fe2++ HOO' Fe3+ + HOO- (1-9)

Fe3+ + HOO* Fe2++ 02 + H+ (1-10)









Equation 1-6 represents the initiation reaction by which hydrogen peroxide is

activated to form a free hydroxyl radical. Further chain propagation reactions involving

the hydroxyl radical ultimately produce peroxyl radicals, as shown in Equation 1-8 with a

hydrogen abstraction. The hydroxyl and peroxyl radicals are reactive oxygen species

known to be powerful oxidizing agents and have been implicated in aging and chronic

disease. In the chain termination step, a peroxyl radical reacts with iron(III) to yield a

proton, molecular oxygen, and regeneration of iron(II) to complete the redox cycle.

Transition-metal Organometallic Complexes

Another method for hydrogen peroxide activation is through the use of transition-

metal cation complexes with organo ligands, including porphyrins. There are a wide

variety of metal complexes described in the literature with any number of differing metal

cations including Cu(I), Cu(II), Ni(II), Co(II), Co(III), Fe(II), Fe(III), Mn(II), and

Mn(III).8 The mechanisms by which these metal porphyrin complexes activate hydrogen

peroxide vary and include production of free hydroxyl radicals as well as formation of

high valent metal-oxo species.8 In the case of free hydroxyl radical formation,

mechanisms based on Fenton type chemistry are proposed, as seen in the [Cu(phen)2]

example in Equation 1-11. In the case of high valent metal-oxo formation reactions, such

[Cu(phen)2]+ + H202 [Cu(phen)2]2+ + HO + HO- (1-11)

as seen by Traylor et al20 in the activation of hydrogen peroxide by iron(III)

tetrakis(pentafluorophenyl) porphyrin, the mechanism of oxidation is believed to occur

via an oxygen transfer from a high valent iron complex.









F
F F

F F

F F F F
N H202
F FeN F "Fe=O"
/ N
F F A F F

F F

F F
F
Figure 1-11. Activation of iron(III) tetrakis(pentafluorophenyl) porphyrin by hydrogen
peroxide to produce a high oxidation state iron complex.20

Methyltrioxorhenium

Methyltrioxorhenium (MTO), first introduced by Beattie and Jones in 1979,21 in

combination with hydrogen peroxide provides a useful system for the oxidation of

alkenes and other substrates such as alkynes and ketones, as reported by Herrmann.22

Considerable research on the kinetics of the reaction of MTO with hydrogen peroxide has

been studied by Espenson, who has shown that two predominant species exist in the

MTO/H202 system, shown in Figure 1-12, that are more stable under acidic conditions.23

Both are efficient oxygen donors to substrates such as phosphines and sulfides. In the

reaction of alkenes with MTO/H202, both of the peroxide adducts react to form epoxide,

but the monoperoxide tends to react slightly faster than the diperoxide.24 There are three

major drawbacks to the use of MTO as an alkene epoxidation catalyst: 1) MTO has a low

stability in the presence of peroxide 2) MTO is both expensive and difficult to synthesize

and 3) because reactions are conducted under acidic conditions, ring-opening of acid

sensitive epoxides is possible.23










+H202, -H20 +H202, -H20 0 0
O-Re 0 O Re Re
-H202,+H20 0 -H202,+H20 O
O O0
0 0 0
MTO monoperoxorhenium(VII) diperoxorhenium(VII)
Figure 1-12. The two dominant forms in the MTO/H202 system under acid conditions.
The diperoxorhenium adduct reacts slightly slower than the
monoperoxorhenium complex.23

S +H1202 O, +H202 O
O=Re=0 I Re=0.- I Re I
II -H202 0 II -H202 0 II 0
Nu: N


NuO NuO
Figure 1-13. Nucleophilic attack of an olefin on the electrophilic oxygen of the hydrogen
peroxide activated methyltrioxorhenium yields the oxidized nucleophile and
regenerates MTO. Attack of a nucleophile on the diperoxo complex generates
the oxidized nucleophile and the monoperoxorhenium complex.24

Asymmetric Oxidation

In addition to activating hydrogen peroxide for use as an oxidizing agent, it is also

advantageous to control the addition of the oxygen atom to substrates, for instance,

alkenes, to generate only one epoxide enantiomer. A number of different methods have

been proposed in the literature and have used a number of different techniques to assure

that the oxygen atom only adds to one face of the alkene. Three of these methods are

discussed in detail below.

Sharpless Oxidation of Allylic Alcohols

A simple and relatively inexpensive method for asymmetric epoxidation of allylic

alcohols using titanium(IV) tetraisopropoxide, tetrabutyl hydroperoxide (TBHP), and (+)

or (-)-diethyl tartrate was introduced by Sharpless in 1980.25 The use of the (+) or (-)-

diethyl tartrate facilitates the addition of the oxygen to one face of the alkene as shown in

Figure 1-14. One of the most interesting aspects of this epoxidation system is the ability









to add the oxygen to a particular face, depending on which tartrate enantiomer is used,

regardless of the substitution pattern of the alkene.




R1

R iR"
HO /



Figure 1-14. Illustration of the asymmetric epoxidation using the Sharpless method. Use
of the (+) or (-)-tartrate allows for the oxygen atom to be added to only one
face of the allylic alcohol.25

Mn(III)-salen Epoxidation Catalysts

Chiral manganese(III)-salen complexes (the salen ligand is illustrated in Figure 1-

15) are another example of an asymmetric alkene epoxidation catalysts. Although the

Sharpless method used tartrate as an additive for achieving asymmetric epoxidation,

manganese(III)-salen catalysts rely on the chirality of the complex to provide for the

asymmetric addition of the oxygen atom. The popularity of the epoxidation system

comes from the ease in synthesis of the catalyst and the use of cheap, readily available

oxidants, such as iodosylbenzene and hypochlorite.




H ""H
N N


R OH HO / R


R R
Figure 1-15. A salen ligand.









Hydrogen peroxide can also be used as the terminal oxidant, although

decomposition of the oxidant by the catalyst has been observed. Katsuki26 in 1994

reported that hydrogen peroxide could be used with Mn(III)-salen catalysts for the

epoxidation of chromene. The yields were low (17-53 %), but with good ee (93-96 %).

It was noted by Katsuki that in order for the reaction of Mn(III)-salen catalysts to

epoxidize alkenes with hydrogen peroxide, an axial ligand was required. For the

epoxidations of chromene, N-methylimidazole was used as the axial ligand. It has also

been noted that the use of carboxylate salts as additives are useful in the epoxidation of

alkenes by Mn(III)-salen catalysts.27 In 1998, Pietikainen found that the use of

manganese(III)-salen with 30 % hydrogen peroxide, along with ammonium acetate,

oxidized spiro[2H-1-benzopyran-2,1'-cyclohexane] in 90 % yield and an ee of 91 %.28








Figure 1-16. spiro[2H-1-benzopyran-2, '-cyclohexane]

Enantiomeric excess (ee) provides a method for reporting the yield of one

enantiomer in comparison to the other. In the case above for the oxidation by

Pietikainen, a 91% ee was reported. This means that 9 % (100% 91%) of the product is

racemic, implying that the remaining mixture is 4.5 % of each enantiomer. The total

yield of the predominant enantiomer is then 95.5 % (91 % + 4.5 %), while the other

enantiomer is 4.5 %.









Chiral Ketone Epoxidation Catalysts

The use of chiral dioxiranes for the epoxidation of alkenes was first reported by

Curci et al. in 1984.29 Shi et al.30'31 has observed that chiral ketones are effective

catalysts for the asymmetric epoxidation of alkenes by the in situ generation of dioxiranes

using potassium peroxymonosulfate (Oxone), as seen in Figure 1-17.

R,
0 0 HS05


RI R2



RR<

R R <^ HS04

RI
Figure 1-17. Asymmetric epoxidation of alkenes can be easily achieved using
peroxymonosulfate to generate a dioxirane in situ.30

These reactions are performed in mixed solvent systems, usually 1,2-

dimethoxymethane and water. The alkene is soluble in the organic solvent, while the

Oxone is soluble in the aqueous layer. The ketone catalyst used is soluble in both water

and the organic solvent, thus allowing the ketone to act as a phase transfer catalyst. The

ketone is oxidized by peroxymonosulfate in the aqueous layer to form the dioxirane,

which then transfers to the organic layer where it can oxidize the alkene and regenerate

the starting ketone. The asymmetric epoxidation is facilitated by the chiral nature of the

ketone. As the generated dioxirane nears the alkene, the oxygen is transferred to only

one face of the alkene. The ee's that have been reported using 1,2:4,5-di-O-

isopropylidene-D-erythro-2,3-hexodiuro-2,6-pyranose, whose structure is shown in

Figure 1-18, range from 12 98%. The lowest ee's are for the epoxidation of cis-alkenes









(12 56.2%), while an ee of 76.4 98% has been observed for the trans-alkenes. The

difference in the ee between cis and trans-alkenes has been attributed to the approach the

catalyst can make to the alkene in the transition states for the two alkenes.30




0
0O 0










Figure 1-18. Structure of 1,2:4,5-di-O-isopropylidene-D-erythro-2,3-hexodiuro-2,6-
pyranose used by Shi30 for the asymmetric epoxidation of alkenes using
peroxymonosulfate to generate a dioxirane in situ.

Peroxycarbonate

Recent work in our group has found that bicarbonate is an effective activator of

H202,32,33 known as BAP, bicarbonate activated peroxide. Equilibrium between

bicarbonate and H202 produces the peroxycarbonate anion (HC04-), as seen in Equation

1-12.

HCO3- + H202 HC04 + H20 (1-12)

The mechanism by which this equilibrium occurs has been determined by Yao34

and has been found to proceed through carbon dioxide as an intermediate. The presence

of carbonic anhydrase or the carbonic anhydrase model complex 1,4,7,10-

tetraazacyclododecanezinc(II) accelerates the equilibrium reaction through catalysis of

the dehydration of bicarbonate and possible catalysis of the perhydration pathway.34 The

complete equilibrium processes is shown in Figure 1-19.



















Ka (H2CO4) Ka (H2CO3)

0H02H OH


O OOH
HO 0 OH HO OH


HO2- + H+ H202 C2 H20 -OH + H

Figure 1-19. The equilibrium formation of bicarbonate and peroxycarbonate proceeds
through CO2 as an intermediate.34

Peroxycarbonate is a strong oxidant with an Eo (HCO4-/HCO3-) of 1.8 0.1 V vs.

NHE.35 Inorganic salts and metal complexes of peroxycarbonate have been isolated and

analyzed by X-ray crystallography and vibrational spectroscopy.36'37 The analysis

indicates that peroxycarbonate is a true peroxide with a structure of HOOC2-. Recently,

an iron(III) complex has been isolated and characterized by X-ray crystallography, as

seen in Figure 1-20.38 Synthesis of metal complexes of peroxycarbonate will be

presented in the next section of this introduction. Peroxycarbonate should not be

confused with sodium percarbonate, which is simply the cocrystallite of sodium

carbonate and H202 (Na2C03*1.5 H202).

Peroxycarbonate is a moderately active oxidant for organic substrates, including

sulfides and alkenes.39-42 The increase in reactivity over hydrogen peroxide can be










attributed to the nature of the leaving group during a nucleophilic attack of the

electrophilic oxygen of peroxide and peroxycarbonate. In the case of hydrogen peroxide,

a general acid is required as a proton transfer agent, as seen in Figure 1-5. In the case of

peroxycarbonate, however, an intramolecular proton transfer can release bicarbonate as

the leaving group, as seen in Figure 1-21. Because bicarbonate is a weaker base than

hydroxide, peroxycarbonate is a stronger electrophile over hydrogen peroxide by a factor

of about 300 based on studies with sulfides.41




CO




C1 C7

901











as 02
Figure 1-20. Fe(qn)2(02C(O)O]Ph4P-1.5MeOH-0.5 (CH3)2NCHO.38

Typical reactions of substrates with peroxycarbonate are slow, but still much faster

than background reactions with hydrogen peroxide alone. For instance, epoxidations of

water soluble alkenes in the absence of bicarbonate yields negligible products in 24

hours. With the addition of bicarbonate, however, NMR analysis shows 90% conversion

to the corresponding epoxide in 15 hours. A similar trend is also seen in sulfide

oxidation.43'44










Nu.


o
O O t H- O- NuO + HCO3

Nu:
Figure 1-21. Nucleophilic attack on the peroxycarbonate anion. An intramolecular
proton transfer in the transition state allows for release of bicarbonate instead
of hydroxide as in the case of hydrogen peroxide.

Transition-metal Peroxycarbonate Complexes

Transition-metal complexes containing the peroxycarbonate dianion ligand, C042,

are known. The general formula for these peracids are LnM(CO4)Xm, where L = an

ancillary ligands, n = 2 or 3, M = Pd, Pt, Rh, or Ir, X = a halogen, and m = 0 or 1.42 The

peroxycarbonate complexes are generally synthesized by passing carbon dioxide gas

through a solution of the LnM(O2)Xm parent complex dissolved in a dry solvent. Two

possible mechanisms for the formation of the peroxycarbonate complexes are shown in

Figure 1-22.

Oxygen label studies43 indicate that the carbon dioxide does not insert into the M-O

bond (pathway 2), but instead inserts into the 0-0 bond (pathway 1). These complexes

are classified as heterolytic oxidants and are good electrophilic oxidants which react with

nucleophiles such as alkenes and phosphines.43

/ / c02 O O
M or M M 0
1 0* 0*

M
0 o-* o o

0


M
0

Figure 1-22. Generation of a metal peroxycarbonate (LnM(CO4)Xm) from its parent 02
complex, LnM(O2)Xm, by passing CO2 through a dry solution of the parent
complex.42









For example, Nyman et al.44'45 observed that when carbon dioxide was passed

through a dry benzene solution of (Ph3P)2PtO2, the platinum peroxycarbonate complex,

whose structure is shown in Figure 1-23, was obtained. The complex was identified

based on its infrared spectrum, chemical properties, and elemental analysis.

0--0
Ph3P t /
Pt

Ph3P
Figure 1-23. Structure of the (Ph3P)2Pt(C04) complex of Nyman.45

Estimation of the peroxide content or oxidation power of the complex was

attempted, but it was stated that no completely satisfactory method was found.44 In

general, the (Ph3P)2Pt(C04) complex in the presence of acidified iodide solutions did

produce an immediate color change attributed to the release of iodine, but the color faded.

The loss of color was thought to occur via the oxidation of the triphenylphosphine, but no

data were presented indicating that the oxidation products were identified.

Unfortunately, attempts to oxidize organic species were not undertaken.

In 2001, Aresta et al.43 examined the reactivity of (PEt2Ph)3RhCl(C04) which was

synthesized by passing CO2 through the parent 02 complex. She describes both the

solution and solid state oxidation of one of the phosphine ligands. In solution, a solvent

molecule displaces a phosphine ligand from the coordination sphere of the metal. The

phosphine is proposed to act as a nucleophile and attack the electrophilic oxygen of the

peroxycarbonate ligand. This reaction then yields the corresponding carbonato Rh

complex and the oxidized phosphine (Figure 1-24, Route A).

In the solid state reaction, the presence of ethylene and the Rh complex does not

yield ethylene oxide. The ethylene displaces a phosphine ligand from the coordination









sphere (Figure 1-24, Route B). The displaced phosphine then attacks the

peroxycarbonate ligand and is oxidized. From these experiments it was concluded that

the mechanism of oxidation does not occur by an intramolecular oxygen transfer from the

peroxycarbonate ligand directly to the phosphine. For this particular complex, the

phosphine must be displaced first before it can be oxidized.

R3P
0\
-PR3 Solv \/ O 0
--Rh O + R3P=O
Solution Cl / O

R3P 0-0 A R3P
R3P Rh /
C1/ O
Cl 00PR3
PR3B -PR3 0

Rh O + R3P=O
Solid State Cl / O
R3P
Figure 1-24. Routes for the oxidation of PR3 by (PEt2Ph)3RhCl(C04).43 Route A shows
the solution chemistry where a solvent molecule displaces a phosphine before
it is oxidized. Route B shows the solid state chemistry where coordination of
ethylene occurs first with the displacement of a phosphine ligand followed by
oxidation of the ligand.

The (PEt2Ph)3RhCl(C04) complex has also been observed to oxidize more reactive

olefins, such as styrene.43 When styrene (0.1 mL, 0.873 mmol) in 2 mL THF was

allowed to react with [(PEt2Ph)3RhCl(C04)] (0.100g, 0.16 mmol) under a CO2/02

atmosphere (10:1 v:v), benzaldehyde, phenylacetaldehyde, phenyl methyl ketone, and

styrene oxide were observed by GC/MS in a ratio of 1:3:3:5, respectively. The presence

of benzaldehyde, phenylacetaldehyde, and phenyl methyl ketone suggest that the

mechanism of oxidation occurs via radical chemistry as opposed to a simple oxygen

transfer in which case styrene oxide would be the only product.










0 0

H




Benzaldehyde Phenylacetaldehyde Phenyl methyl ketone
Figure 1-25. Structure of products of styrene oxidation by [(PEt2Ph)3RhCl(C04)] under a
CO2/02 atmosphere that indicate a radical mechanism.43

Transition-metal Activation of Peroxycarbonate in Solution

Recent work by Burgess et al.39'40 has shown that the addition of certain transition-

metal cations increases the catalytic rate of alkene epoxidation in solutions of hydrogen

peroxide and bicarbonate in mixed solvent systems. Of the inorganic metal salts tested,

manganese(II) sulfate produced the greatest increase in the epoxidation reaction. Along

with an increase in the rate of epoxidation, the addition of Mn(II) to solutions of H202

and bicarbonate also enhances the rate of H202 disproportionation. The rate of

disproportionation is enhanced to such a degree that methods must be employed to deal

with the excessive amount of heat evolved. Recent studies by Bennett46 have shown that

the addition of the chelating agent diethylenetriaminepentaacetic acid (DTPA) inhibits

the oxidation of alkenes, but only in some cases. This has been attributed to the removal

of extraneous metal cations from the bicarbonate salts.46

O

0 OH
OH

N N
N

HO OH O


0 OH 0


0
Figure 1-26. The structure of diethylenetriaminepentaacetic acid (DTPA)









In 1977, Sychev et al.47 reported that hydrogen peroxide is rapidly

disproportionate in the presence of bicarbonate and free Mn(II). In a series of papers

from 1977 to 1984,47-56 an investigation on the reaction mechanism of H202

decomposition with manganese(II) and bicarbonate was conducted. His proposed

mechanism assumes that Mn(II) follows the Fenton type chemistry of Fe(II) with its

reaction with H202 and therefore, proceeds via a free hydroxyl radical pathway. His

work, however, does not provide adequate detail into the necessity of the bicarbonate ion

in this reaction. Addition of similar anions, such as acetate, phosphate, oxalate, or borate,

does not result in H202 disproportionation when Mn(II) is introduced. Also, the

explanation provided by Sychev does not enlighten us in the observation that alkenes are

cleanly oxidized to epoxide without detection of usual radical coupled products, as seen

in the work by Aresta.43

Scope of the Dissertation

The goal of this current study is to further understand the reactive nature of the

peroxycarbonate anion and dianion. Chapter 2 will discuss the reaction of Mn(II) and

bicarbonate in the oxidation of styrene in micellar media. Both small and large scale

oxidations of styrene were attempted and the results of these experiments will be

presented. Questions arising from these experiments led us to investigate the hydrogen

peroxide disproportionation reaction further.

The importance of peroxycarbonate in the Mn(II) catalyzed disproportionation of

hydrogen peroxide will be the focus of Chapter 3. Kinetic investigations of the reaction

have been conducted, and the results of these experiments will be presented. The lifetime

of the catalyst was also investigated. The similarities between hydrogen peroxide

disproportionation and nucleophilic alkene oxidation using Mn(II) was also of interest









during this study. A proposed mechanism for the hydrogen peroxide decomposition and

alkene epoxidation will be introduced and numerical simulations of various proposed

models for the disproportionation and alkene epoxidation will be presented.

Chapter 4 will discuss the use of the peroxycarbonate dianion as a nucleophilic

oxidant for epoxidation of electrophilic alkenes. The results of experiments with the

peroxycarbonate dianion will be compared with kinetic measurements using other

nucleophilic oxidants. A summary and discussion of possible future work will comprise

Chapter 5.














CHAPTER 2
OXIDATION OF NUCLEOPHILIC ALKENES IN AQUEOUS MICELLAR MEDIA

Introduction

Prior work by Bennett46 and Yao33 has shown that alkenes can be oxidized to the

corresponding epoxides using bicarbonate-activated peroxide. Studies of hydrophobic

alkenes were conducted in mixed solvent systems to all

ow for solubility of the alkene, while water-soluble alkenes were chosen for pure

water studies. In general, epoxidations were found to be slow using peroxycarbonate

solutions, but faster than background oxidations by hydrogen peroxide alone. It was also

found that reactions proceed faster in pure water than in mixed solvent systems. The

faster reaction in water has been attributed to a proton transfer, which proceeds faster in

pure water than in mixed solvent systems.32 It was also noted by Burgess et al.40 that the

addition of transition metal salts in the presence of hydrogen peroxide and bicarbonate in

H20/DMF solutions accelerated the oxidation of alkenes. Of the transition-metals tested,

manganese(II) sulfate produced the most dramatic increase in oxidation. In order to take

full advantage of the use of transition-metal salts for alkene oxidation, surfactants were

used to allow for the oxidation of hydrophobic alkenes in aqueous solution in the

presence and absence of manganese(II) salts. The use of surfactants is also advantageous

since the work by Bennett46 indicates that the oxidation of alkenes tends to proceed faster

in pure water.

Surfactants are long chain alkanes with hydrophobic tails and polar head groups,

which allow for solubility in water. Three examples of common surfactants,









cetyltrimethylammonium chloride (CTAC1), sodium dodecylsulfate (SDS), and Triton X-

100, are found in Figure 2-1.



C CO-S-O0 Na
O
\ /15 1 0
0
cetyltrimethylammonium chloride sodium dodecylsulfate



(H3C)3CCH2(H3C)2C- -- OCH2CH2OH

x=10
Triton X-100
Figure 2-1. The structures of three common surfactants. Cetyltrimethylammonium
chloride is a cationic surfactant, while sodium dodecylsulfate is anionic.
Triton X-100 is a non-ionic surfactant.

When dissolved in water, surfactants will begin to organize themselves into

micelles after the concentration reaches a crucial level known as the critical micelle

concentration (cmc),57 which is a unique value for each surfactant and depends on the

ionic strength of the solution. The detection of micellization can be accomplished by

observation of the surface tension, refractive index, or conductivity (for ionic

surfactants).7

At the cmc, the hydrophobic tails will begin to congregate, expelling water from

the forming micelle's core, while the polar head groups will arrange to allow for the

maximum interaction with water, as seen in Figure 2-2. The Stern layer is defined as the

area around the micelle where the polar head groups are located, as are their counter
ions.57
ions.









Normally, hydrophobic molecules are unable to dissolve is aqueous solution.

However, if micelles are present, a hydrophobic molecule, such as an alkene, is able to

penetrate into the hydrophobic core of the micelle and dissolve, as seen in Figure 2-3.

Polar Head
Group C
Group Counterion

Hydrophobic Core


Figure 2-2. The structure of a micelle with a concentration greater then the cmc.


Figure 2-3. The graphical representation of an alkene dissolved in a micelle.









Results and Discussion

Styrene Oxidation in Micellar Media in the Absence of Mn(II)

Initially, styrene was oxidized in micellar media using the BAP method without the

introduction of manganese(II) salts (Figure 2-4). Styrene (50 mM), CTAC1 (100 mM),

H202 (2.00 M), and ammonium bicarbonate (1.00 M) in a volume of 250 mL were

allowed to react in water for 3 days in the dark with stirring.

Analysis of the products by HPLC showed that approximately 90% of the starting

styrene had reacted to yield the corresponding epoxide (-90 %), although significant

hydrolysis to the corresponding diol has also been detected (-10%).

o OH


H202, HCO3 H20 O
CTAC1 OH

Figure 2-4. The reaction scheme for the oxidation of styrene by hydrogen peroxide in the
presence of bicarbonate and cetyltrimethylammonium chloride (CTAC1)
without the presence of Mn(II). Hydrolysis of the product epoxide forms the
corresponding diol. Reaction conditions: 0.05 M Styrene, 0.10 M CTAC1,
2.00 M H202, 1.00 M NH4HCO3, 3 days

Large Scale Styrene Oxidation

In addition to small scale epoxidations of styrene in micellar media, large scale

epoxidations were attempted. For a typical large scale epoxidation, 5 mL of styrene (175

mM), 230 mL CTAC1 (350 mM), 2.00 M H202, 39.6 g NH4HCO3 (1.00 M), and enough

water (to bring the volume to 500 mL) were allowed to react in a total volume of 500 mL

with stirring for 3 days in the dark at room temperature. For the small scale epoxidations,

purification of the unreacted styrene and the styrene oxide product was unnecessary due

to the use of HPLC for the analysis of the reaction products. For the large scale

epoxidations, however, a purification method was required.









First, extraction with methylene chloride was attempted for the isolation of the

styrene and styrene oxide. Since the surfactant would probably produce an emulsion, it

was thought that if the reaction solution were diluted the emulsion would dissipate within

a short amount of time. Unfortunately, this was not the case. The emulsion formed, even

when the reaction was diluted to 2.5 L.

On occasion, allowing the emulsion to stand overnight would allow for small

amounts of methylene chloride to be isolated from the extraction. Upon drying and

removal of solvent, styrene oxide, the surfactant, and water were observed by 1H NMR

analysis. While the surfactant, CTAC1 in this case, is a cationic species and should not

dissolve in organic solvents, the surfactant could form a reverse micelle, where the polar

head groups now surround a small amount of water and the hydrophobic tails extend into

the organic solvent.5 This would explain why both surfactant and water are observed by

1HNMR.

Given the unsatisfactory results from extraction, another method for purification of

the organic products was required. The second method used for the purification of large

scale styrene oxidations was liquid-liquid extraction. This purification method uses the

same principal as extraction, but without the tendency to form emulsions. At first, for the

large scale epoxidations, ether was used for the liquid-liquid extractions since only a

lighter-than-water liquid-liquid extraction apparatus was immediately available. The

setup of the apparatus is shown in Figure 2-5.

For the process of liquid-liquid extraction, a constant volume of organic solvent can

be used to extract the organic product from the aqueous layer. This provides a

convenient method for extracting slightly soluble organic products with a minimum








amount of organic solvent, as opposed to normal extraction procedures which typically

require larger volumes of organic solvent.

I


p
N


Condenser


I. .01


Reflux Arm



Return Arm






Distilling Organic_
Solvent *l i


I Funnel


Organic Layer



Aqueous Layer


Figure 2-5. A picture of a lighter-than-water liquid-liquid extractor.

Initially, the organic solvent is layered over the aqueous solution. In addition, a

small amount of organic solvent is placed in a round bottom flask connected to the reflux

arm. Once the organic solvent in the round bottom has begun refluxing, it will be









liquefied in the condenser and collects in the funnel. As the organic solvent collects in

the funnel, small amounts of the solvent will be pushed out the end of the funnel into the

aqueous layer. As the organic solvent rises through the aqueous layer, a small amount of

organic product will diffuse into the droplet. After a few minutes, enough organic

solvent will have added to the solvent layered over the aqueous solution to allow the

organic solvent to drip down the return arm back to the refluxing solvent. In this way,

the organic product will slowly accumulate in the round bottom flask. After the

extraction is complete, the solvent in the round bottom can be dried and the solvent

removed to give the desired product, as would be done for a normal extraction process.

The time of completion for the extraction must be determined experimentally for the

unique conditions in which the extractor is being used. Heavier than water liquid-liquid

extractions also exist and extract the organic product in a similar way.

For the large scale epoxidations, it was found that the highest yield of epoxide was

observed after 3 days of liquid-liquid extracting. In addition to the styrene oxide, the

corresponding diol was also present (-15%), when analyzed by HPLC. This is

reasonable since after 3 days of reacting a 10% conversion to the diol is observed. The

additional 3 days of extraction accounts for the additional 5% conversion of the epoxide

to the diol product. On occasion, surfactant and water were still observed in the purified

product, even when care was taken to assure that no emulsions were formed when

layering the organic solvent over the aqueous layer in the extractor.

Styrene Oxidation in Micellar Media in the Presence of Mn(II)

Recent work by Burgess et al.39'40 has shown that the introduction of transition-

metal salts to oxidations of hydrophobic alkenes by H202 and bicarbonate in mixed

solvent systems of dimethyl formamide (DMF) and water show a significant rate









enhancement, Figure 2-6.39,40 Upon addition of a transition-metal salt, epoxidations with

reaction times greater than 48 hours were decreased to only 16 hours. The main

drawback, specifically the long reaction times, to Burgess' method has been the slow

addition of the H202 and bicarbonate solutions to DMF to minimize precipitation.40 This

rate enhancement is greatest when the inorganic salt added is manganese(II) sulfate.

0


H202, HC03, N 2+
DMF/H20

Figure 2-6. Reaction scheme used by Burgess40 in the mixed solvent epoxidation of
styrene.

Our current method of alkene epoxidation using micellar media offers a better

alternative to the mixed solvent system employed by Burgess, Figure 2-7. The main

drawback seen by Burgess, namely the slow addition of the H202/bicarbonate solution, is

not an issue in micellar media. The organic substrate is dissolved by the micelle, and all

of the remaining reactants are freely soluble in water, so precipitation is no longer of

concern. When a test reaction was performed using styrene (50 mM), CTAC1 (100 mM),

H202 (2.00 M), NH4HCO3 (1.00 M), and only 10 tM MnSO4, the epoxidation was

complete in less than 30 minutes as seen by the HPLC chromatograms in Figure 2-8, as

opposed to 3 days in the absence of manganese(II).

0


H202, HC03, Mn2+
Surfactant, H20

Figure 2-7. Schematic representation for the oxidation of styrene in surfactant with
hydrogen peroxide and bicarbonate catalyzed by manganese(II). Reaction
conditions: 50 mM styrene, 0.10 M CTAC1, 2.00 M H202, and 1.00 M
NH4HCO3, 30 minutes.













9,414 miin


I i

I I


I 'i


S Styrene Peak at Time 0


-z.


S8303 min


I I


Stb rene Oide Peak- at 30 Minutes


Figure 2-8. HPLC chromatograms for the initial reaction (top panel) and after 30 minutes
(bottom panel) for the oxidation of styrene with H202, HC03-, and Mn(II) in
the presence of surfactant (CTAC1). HPLC performed using a C18 reverse
phase column using a non-linear gradient for 12 minutes. Mobile Phase:
25%:75% (v:v) CH3CN:H20 95%:5% CH3CN:H20


AVO -









Reaction Kinetics

Kinetic experiments were conducted to determine the dependence of various

conditions, including the identity of the surfactant, the source of the manganese, and the

bicarbonate concentration, on the manganese(II) catalyzed oxidations of styrene in

micellar media. For these reactions, aliquots of reaction solutions were removed over

time and added to a solution of bovine catalase to destroy any remaining H202 and

therefore, quench the reaction. The aliquots were then diluted with acetonitrile and

analyzed by HPLC. Figure 2-9 shows a representative graph demonstrating the

disappearance of styrene versus time. From a plot of the In(styrene area) versus time, the

first-order rate constant can be determined from the slope of line, as shown in Figure 2-

10.


4000000
3500000
a 3000000 -
S2500000
S2000000
S1500000
m 1000000 -
500000 -
0 *
0 700 1400 2100 2800 3500
Time, seconds


Figure 2-9. Styrene area disappearance versus time from the HPLC analysis of styrene
oxidation by hydrogen peroxide in micellar media in the presence of
bicarbonate and Mn(II). Reaction conditions: 0.05 M Styrene, 0.100 M
CTAC1, 0.25 M NH4HCO3, 1.00 M H202, 10 iM Mn(II).










16
15 y y=-0.0014x+ 15.385
kobs= (1.43 + 0.10) x103, s-1
S14
5 13
12
11
10
0 500 1000 1500 2000 2500 3000 3500
Time, seconds

Figure 2-10. In(styrene area) versus time to find the first-order rate constant. The line is
the linear regression to the data at the 95% confidence. The kobs is the
negative slope of the line.

Dependence of Styrene Oxidation on Surfactant Identity

For all reactions previously described, the surfactant used was

cetyltrimethylammonium chloride, a cationic surfactant. In order to determine whether

the active species is charged, the oxidation of styrene was performed under the same

conditions as described except for the substitution of sodium dodecylsulfate (SDS) for

CTAC1. If the active catalytic species is positively charged, the reaction should proceed

faster in the anionic micelle due to attractive forces between the micelle and the active

oxidant. If the active oxidant is negatively charged, the reaction should be slower in the

anionic micelle due to the repulsive force between the micelle and the active oxidant.

Conversely, if the active catalyst is uncharged, no difference in the rate of the reaction

should be observed. From the data presented in Table 2-1, the observed rate constants for

the oxidation of styrene in the two surfactants give the same observed first-order rate

constants (within error), therefore, the conclusion must be that the active manganese

oxidant is uncharged.









Table 2-1. Comparison of Styrene Oxidation in CTAC1 and SDS for the Mn(II)
catalyzed epoxidation. Reaction conditions: 0.05 M Styrene, 0.100 M CTAC1
or SDS, 0.25 M NH4HCO3, 1.00 M H202, and 10 [M Mn(II). Errors are
reported to the 95% confidence.
Surfactant kobs, s-1
Cetyltrimethylammonium chloride (1.39 0.10) x10-3
Sodium dodecylsulfate (1.50 + 0.15) x10-3

Dependence of Styrene Oxidation on the Manganese(II) Source

In addition to testing whether the surfactant made an impact on the reaction, the

addition of the metal was also examined. When bulk manganese(II) is added to the

solution containing SDS, the manganese(II) ions must exchange with sodium ions at the

surface of the micelle.

Manganese(II) bisdodecylsulfate (Mn(DS)2) was synthesized to allow for the metal

to be added already bound to the surfactant. Mn(DS)2 is precipitated by the addition of

saturated sodium dodecylsulfate and saturated manganese(II) chloride. Since the

manganese(II) is bound to the surfactant, all of the manganese(II) should be bound to the

micelle surface.

O O

110 5-0-Mn--0-

O O
Figure 2-11. Structure of manganese(II) bisdodecylsulfate.

When the reaction was performed using 10 [LM Mn(DS)2, an identical observed

first-order rate constant, within error, was observed in comparison with the reaction with

the addition of bulk manganese(II). The results of this experiment indicate that with the

addition of bulk manganese(II), the metal is in rapid equilibrium with the sodium ions at

the micelle surface. So, in the case of Mn(DS)2, the metal is not remaining at the micellar

surface, but is rapidly being released into the bulk solution by exchange with sodium









ions. It is, therefore, unnecessary to add the metal already bound to the surfactant, since

the bulk metal rapidly exchanges with sodium ions at the micellar surface.

Table 2-2. Comparison of observed rate constants for differing manganese sources for
micellar styrene oxidation. Reaction conditions: 0.05 M Styrene, 0.100 M
SDS, 0.25 M NH4HCO3, 1.00 M H202, and 10 [tM Mn(II) or Mn(DS)2.
Errors are reported to the 95% confidence.
Manganese Source kobs, S-1
Bulk Manganese (Mn2+) (1.50 + 0.15) x10-3
Mn(DS)2 (1.43 + 0.07) x10-3

Bicarbonate Dependence

The bicarbonate dependence on the styrene oxidations in micellar media was also

investigated. When a plot ofkobs versus [HCO3-] was made, as seen in Figure 2-12, the

rate has saturated by 0.25 M bicarbonate. This finding was unexpected, since previous

work on sulfide and alkene oxidations did not show saturation with bicarbonate at

concentrations similar to those used here.32,41 It is apparent, therefore, that the addition of

the surfactant is concentrating the active oxidant near the micelle surface in order for the

oxidation to be saturating. Since the active oxidant had not been examined at this time, it

was impossible to make any conclusive remarks about how the micelle was affecting the

epoxidation of styrene by manganese(II) and bicarbonate with hydrogen peroxide in the

presence of surfactant. In order to further understand the use of the manganese(II) as a

catalyst for the epoxidation of alkenes, the hydrogen peroxide disproportionation reaction

needed to be better understood. The hydrogen peroxide disproportionation can be

examined in the absence of surfactant, since all of the reactants are fully water soluble. If

the active oxidant can be identified, further experiments can then be designed to probe

the nature of the oxidant in micellar oxidations of alkenes. The results of experiments

with the disproportionation reaction are presented in the next chapter.













2.5
2.0
1.5 -
1.0
0.5
0.0 ,
0 0.2 0.4 0.6 0.8 1

[NH4HC03], M

Figure 2-12. Graph of kobs versus [NH4HCO3] for the styrene oxidation in the presence of
0.100 M CTAC1. Reaction conditions: 0.05 M Styrene, 0.100 M CTAC1, 1.00
M H202, and 10 [LM Mn(II).

Experimental

Materials and Instrumentation

Sodium bicarbonate, sodium acetate, styrene, and manganese (II) sulfate were all

analytical grades and obtained from Fisher (Atlanta, GA). Cetyltrimethylammonium

chloride was obtained from Aldrich (St. Louis, MO). Hydrogen peroxide (35 and 50%)

was obtained from Fisher (Atlanta, GA) and standardized often by iodometric titration.

Water was purified using a Barnstead E-Pure 3-Module Deionization System.

Extraneous metal ions from salt solutions were removed by passing the solutions through

a Chelex 100 resin obtained from Aldrich (St. Louis, MO). Sodium bicarbonate solutions

were standardized using the method below before use to assure concentration.

UV-vis kinetic experiments were obtained using a Hewlett-Packer 8453

spectrophotometer using 1.0 cm quartz cells from Starna Cells, Inc. Temperature was

maintained at 25 (+ 0.1) C using a Fisher Isotemp 1600S water bath circulator.









Styrene oxidation reactions were analyzed by High Performance Liquid

Chromatography (HPLC) using a Rainin HPLX solvent delivery system with a C-18

reverse phase column. The method consisted of a non-linear gradient of mobile phase

A:H20 and mobile phase B:CH3CN from 75:25 A:B to 5:95 over a 12 minute period.

Products were detected at 221 nm.

Standardization of Sodium Bicarbonate Solutions

Solutions of sodium bicarbonate were standardized before each kinetic experiment

to determine the concentration eluting from the Chelex 100 column. All solutions were

delivered using volumetric pipets. A 10 mL aliquot of sodium bicarbonate solution

exiting the Chelex 100 column was a placed in a clean, dry beaker. An excess amount of

sodium hydroxide solution of a known concentration, by titration with potassium

hydrogen phthalate, was added to the beaker. The solution was stirred to allow for

complete deprotonation of the bicarbonate to form the carbonate dianion. An excess

barium chloride solution is then added to precipitate all of the carbonate dianion as

barium carbonate. Phenolphthalein is then added to the mixture to give a pink color due

to the residual hydroxide ion. The mixture is titrated using a known concentration of

hydrochloric acid until the solution just turns clear. The number of moles of hydrochloric

acid added is equal to the excess moles of sodium hydroxide. The difference between the

number of moles from the acid titration and the number of moles of hydroxide ion

initially added equals the number of moles of bicarbonate present in the initial 10 mL

aliquot (Equation 2-1). The molarity of the solution can then be determined.

#moles OHinit-#moles OHexcess = #moles bicarbonate (2-1)









Styrene Oxidation Reactions

Kinetic experiments of styrene oxidations were performed in micellar solutions and

analyzed by the decreasing area of the styrene peak by HPLC. Reactions were performed

using 0.05 M styrene, varying ammonium bicarbonate, 1.00 M H202, and 0.10 M

surfactant, where the surfactant was either cetyltrimethylammonium chloride or sodium

dodecylsulfate (CTAC1 and SDS, respectively). For reactions using manganese(II), as

either manganese sulfate of Mn(DS)2, 10 [LM was added to the reactions. Aliquots (100

ItL) of the reaction mixture were taken over time and quenched using a catalase solution,

which converts any excess H202 to 02 and H20, thus removing the terminal oxidant.

Each of the aliquots is then diluted with CH3CN to an appropriate concentration for

HPLC analysis. The kobs is calculated using pseudo-first order plots of the In(styrene

area) vs. time. First-order plots were linear for the conditions studied.

Large Scale Styrene Oxidations

Large scale styrene epoxidations were conducted in micellar solutions. Reactions

were performed using 175 mM Styrene (5 mL), 350 mM CTAC1, 1 M NH4HCO3, and 2.0

M H202 in a total volume of 500 mL, where the remaining volume is water. The reaction

was allowed to stir in the dark for 3 days. The reaction mixture was then diluted to 1 L

and poured into a lighter than water liquid-liquid extractor. Ether was then layered on top

of the aqueous reaction solution, and the extractor was allowed to extract for 3 days.

After cooling the receiving flask after 3 days, magnesium sulfate was added to dry the

solvent. After filtering off the magnesium sulfate and removing the solvent under

reduced pressure, the organic residue was analyzed by 1H NMR.









Synthesis of Mn(DS)2

Mn(DS)2 was synthesized by mixing equal parts saturated manganese(II) chloride

and saturated sodium dodecylsulfate in water to form a while solid. The solid was then

filtered and washed with ice-cold water. Calculated for C24H50S2OsMn4H20 Calc: C:

43.82% H:8.89% S:9.75% 0:29.19%Mn:8.35% Found C: 43.65% H:9.10%

Styrene Oxidation in SDS with Mn(II) and Mn(DS)2

Styrene oxidations were performed using 0.05 M styrene, 0.500 M ammonium

bicarbonate, 1.00 M H202, and 0.10 M sodium dodecylsulfate. Manganese(II) was added

as either 10 [LM bulk manganese(II) or 10 [LM Mn(DS)2. Aliquots of the reaction mixture

were taken over time and quenched using a catalase solution, which converts any excess

H202 to 02 and H20. Each aliquot is then diluted with CH3CN to an appropriate

concentration for HPLC analysis. The kobs are calculated using pseudo-first order plots of

the In(styrene area) vs. time. First order plots were linear for the conditions studied.














CHAPTER 3
KINETIC INVESTIGATIONS OF THE MANGANESE(II) CATALYZED
DISPROPORTIONATION OF HYDROGEN PEROXIDE IN THE PRESENCE OF
BICARBONATE AND THE COMPARISON TO NUCLEOPHILIC ALKENE
EPOXIDATION

Introduction

Investigations on the use of manganese(II) in the catalysis of alkene epoxidation

presented in the preceding chapter raise some interesting questions about the hydrogen

peroxide disproportionation under the reaction conditions. Namely, what is the active

manganese species? Why is bicarbonate necessary for the reaction to proceed?

In 1977, Sychev et al.47 reported that hydrogen peroxide disproportionate rapidly

in the presence of bicarbonate and free manganese(II). In a series of papers from 1977 to

1984,47-56 Sychev studied the mechanism of peroxide decomposition. His proposed

mechanism assumes that manganese(II) follows the Fenton chemistry of iron(II) in its

reaction with peroxide, and therefore, proceeds via a free hydroxyl radical pathway. His

work, however, does not provide adequate detail into the unique catalytic ability of the

bicarbonate ion in this reaction or the identity of the active manganese species.

In 1990, Stadtman et al.58 also investigated the use of the manganese(II) catalysis as

a method for oxidation of amino acids. The decomposition of hydrogen peroxide by

manganese(II) with bicarbonate was also examined during their studies. As with the

work of Sychev, Stadtman does not provide an explanation for the unique reactivity of

bicarbonate or speculate about the active manganese species responsible for the catalysis.









In this chapter, kinetic investigations on the manganese(II) catalyzed

decomposition of hydrogen peroxide in the presence of bicarbonate will be presented.

The bicarbonate, manganese(II), and hydrogen peroxide dependencies measured during

this study are similar to those reported by Sychev,47 but differ slightly from those

observed by Stadtman.58 The observed differences in the dependence are probably due to

the conditions under which the reaction was examined. For this work and for that of

Sychev, the hydrogen peroxide and bicarbonate concentrations were in the 100-500 mM

range, with a 0-10 pM Mn(II) concentration. For the studies performed by Stadtman, the

hydrogen peroxide and bicarbonate concentrations were in the 30 mM range, with much

larger Mn(II) concentrations of 0.10 mM. The dependencies measured during this study

are also similar to those observed by Bennett46 for the epoxidation of nucleophilic

alkenes by hydrogen peroxide in the presence of bicarbonate catalyzed by manganese(II).

Investigations on the lifetime of the catalyst have also been conducted. While the

catalyst is still active upon addition of hydrogen peroxide to spent solutions, the reactivity

is about half of the original activity. These results are consistent with those reported by

Stadtman. Two factors have been identified to explain the loss of activity.

Experiments have also been conducted to examine whether the manganese source

has an impact on the reaction. Three different manganese sources were used in this study

and include manganese(II) sulfate, potassium permanganate, and a Mn(IV)-TACN

complex. The results of these experiments indicate that the source of the manganese has

no effect on the observed rate of hydrogen peroxide decomposition.

Experiments were also conducted using cis-alkenes to discern information about

the mechanism of oxygen transfer from the active catalyst. Results from these









experiments indicate that the oxygen atom is not being transferred in a concerted fashion,

since cis/trans isomerization is observed in the epoxidation of nucleophilic alkenes.

Investigations of the radical traps used in the work by Sychev led to the interesting

result that amines can react with this system to yield N-dealkylated products. Reports on

oxidative N-dealkylation are not common in the literature. Current knowledge of

oxidative N-dealkylation comes from studies with cytochrome P-450. The most accepted

mechanism for the oxidative N-dealkylation of amines begins with a single electron

transfer from the amine to a metal in a high oxidation state. A mechanism using

manganese species that will be proposed in this study will be presented.

Solvent isotope effects have been measured for the nucleophilic alkene epoxidation

and hydrogen peroxide decomposition. A large, inverse isotope effect was observed for

the alkene oxidation, while a normal isotope effect was observed for the hydrogen

peroxide decomposition. The results of the nucleophilic alkene oxidation are consistent

with those observed by Bennett. An attempt to justify the difference in the observed

solvent isotope effects will be presented.

Finally, a mechanism based on the work from this study will be presented in an

attempt to explain both the nucleophilic alkene epoxidation and hydrogen peroxide

decomposition. Numerical simulation has been employed as a tool in the attempt to

understand the reaction kinetics. Several models will be presented and discussed.

Results and Discussion

Kinetics of Hydrogen Peroxide Decomposition

Hydrogen peroxide disproportionation, Figure 3-1, was followed

spectrophotometrically by monitoring the decreasing peak intensity at 263 nm. The

hydrogen peroxide concentration was held constant at 0.10 M, while the sodium









bicarbonate was varied in concentration from 0.0-0.60 M. The manganese(II)

concentration was varied in the range from 0-30 [tM.

Mn(II)
2 H202 2 H20 + 02
NaHCO3, buffer
Figure 3-1. Hydrogen peroxide decomposition in the presence of manganese(II) and
bicarbonate.

In order to maintain a constant pH of 8.4 and an ionic strength of 1.00 M, two

buffering systems were tried. Initially, sodium phosphate buffer was used, but it was

discovered that the phosphate anion causes the manganese to precipitate. Also, the

phosphate buffer did not maintain the pH. Even with higher concentrations, the pH

would continue to rise during the experiment.

Eventually, sodium acetate was employed to control ionic strength and stabilize

pH. While sodium acetate is not actually a buffer at pH 8.4, the addition of the salt in the

solutions allowed for the stabilization of the pH. To exclude acetate as participating in

the mechanism, a set of experiments were conducted in the absence of acetate. While the

pH did change slightly during these reactions, the observed rate constants were nearly

identical to those in the presence of acetate.

Solutions of hydrogen peroxide, bicarbonate, and sodium acetate were allowed to

equilibrate for at least 15 minutes before kinetic experiments were performed. Addition

of a solution of the manganese(II) was always preformed last to initiate the reaction. To

ensure that the order of reagent addition does not have an effect on the reaction,

experiments were conducted which were initiated by the addition of hydrogen peroxide.

These experiments gave nearly identical observed rate constants as those initiated by

manganese(II), therefore, the order of reagent addition makes no significant difference to

the overall reaction rates.










The disproportionation reaction was monitored for greater than two half-lives.

Linear first-order plots were only obtained under certain conditions from the absorbance

decay by plotting the equation ((At-Ao)/(Ao- Ao)) versus time, as seen in Figure 3-2. In

most cases, however, linear first-order plots were not obtained. For these instances, the

first-order plots were used as an attempt to approximate an observed first-order rate

constant. In all cases, the reaction accelerates near the end of the reaction. The

hypothesis is that hydrogen peroxide is an inhibitor of the reaction. As the concentration

of hydrogen peroxide drops, an acceleration in the peroxide decomposition will occur.

Further details concerning this aspect of the reaction will be presented in the section

discussing the numerical simulations.


01
-0.5


S-1.5



-2.5
0 2000 4000 6000 8000
Time, seconds

0.1MNaHCO3 U 0.2MNaHCO3 A 0.3MNaHCO3

Figure 3-2. Plot of the ln([H202]) versus time for varying bicarbonate concentration. The
0.20 and 0.30 M bicarbonate reactions are typical of the accelerations noticed
for these reactions. Reaction conditions: 0.10 M H202, 4 [tM Mn(II).

The slope of the line produced by linear regression, of those plots that are linear, is

the negative of kobs. For reactions that were not linear, pseudo "kobs" were attempted to

be used to analyze the data for these reactions. Background disproportionation of

peroxide in bicarbonate and acetate buffered solutions was negligible for the time scale of

the catalytic disproportionations.










Manganese(II) Dependence

The plot of kobs versus Mn(II) concentration gives a straight line with a y-intercept

of 0 for the range of Mn(II) from 0-6.0 [aM, as seen in Figure 3-3, indicating a first-order

dependence of Mn(II) on the reaction in that range. Since the hypothesis is that the

catalytic disproportionation of peroxide is metal dependent, a y-intercept of 0 is expected.

The straight line plot indicates that only one manganese ion is found in the transition state

complex. These data indicate that a multiple metal center complex is not involved in the

disproportionation reaction under aqueous conditions. This is significant since much of

the current literature on metal catalyzed peroxide disproportionation focuses on

complexes with multiple metal centers.59-62 Scatter in the Mn(II) data begins to appear at

about 6.0 [M. The turnover in the data is attributed to precipitation of manganese(II)

carbonate.


8.00E-03

6.00E-03 -

I 4.00E-03

2.00E-03

0.00E+00 /
0 5 10 15 20 25 30
[Mn2+], M

Figure 3-3. The dependence ofkobs on the [Mn(II)]. Reaction conditions: 0.10 M H202,
0.4 M HC03-, varying [Mn(II)]. y = ((7.98 + 0.62) x10-4)x, error reported to
the 95% confidence.

The solubility-product constant (Ksp) is defined as the equilibrium constant between

the ions in solution and the precipitated solid. An example using manganese(II)

carbonate is given below.









MnCO3(s) Mn2+(aq) + CO32-(aq) (3-1)


Ksp =[Mn2+]* [C32-] (3-2)

The Ksp is reached at -6.50 [tM Mn(II), given a constant bicarbonate concentration of

0.40 M, a pH of 8.4, and a Ksp of 2.72 x10-7, the conditions under which the Mn(II)

dependence was measured. This corresponds well with the turnover in the manganese

dependence, Figure 3-3. This turnover was also observed by Sychev47 for his

experiments on peroxide disproportionation in bicarbonate solutions with manganese(II),

but further data as to why the turnover occurred were not presented. The explanation

given in the text simply referred to possible precipitation of manganese salts.

Bicarbonate Dependence

Studies of the dependence of sodium bicarbonate on the rate law indicate a second-

order dependence, as seen in Figure 3-4. The plot of kobs versus [HCO3-]2 produces a

straight line with a y-intercept of 0. The data begins to scatter at higher bicarbonate

concentration for two reasons. First, the reactions become very fast and getting accurate

data becomes more difficult. Second, the higher bicarbonate reactions have a more

pronounced curve in the In plots (as discussed earlier), therefore the "kobs" reported is not

a true first-order rate constant. As with Mn(II), since the disproportionation of peroxide

is dependent on the bicarbonate concentration, a y-intercept of 0 is expected. For the

sodium bicarbonate, however, a second-order dependence reveals that two bicarbonate

ions are present in the transition state complex of the rate determining step. Presumably,

one of these bicarbonate ions is in the form of peroxycarbonate, while the other may

simply be coordinated to the metal center. Similar results were found for the nucleophilic

alkene epoxidations studied by Bennett46 and the hydrogen peroxide decomposition









studies of Sychev.47 Stadtman,58 on the other hand, reports a third-order dependence on

the bicarbonate concentration. The difference in the order of the reaction could be due to

the reaction conditions that were used, since the hydrogen peroxide and bicarbonate were

much lower than this study and the manganese was much higher. Currently, with the

proposed mechanisms that will be presented later, a third-order dependence has not be

observed in the numerical simulations.



8.00E-03
7.00E-03
6.00E-03 -
S5.00E-03 *
S4.00E-03
S3.00E-03
2.00E-03
1.00E-03
0.00E+00
0 0.05 0.1 0.15 0.2 0.25 0.3 0.35

[NaHC3]2, M2



Figure 3-4. Plot of kobs versus [NaHCO3]2. Reaction conditions: 0.10 M H202 and 4 [LM
Mn(II). y = ((2.08 0.25) x10-3)x, error reported to the 95% confidence.

The overall rate equation defined by the results above is given in Equation 3-3. At

this time, no reliable method has been found to study the hydrogen peroxide dependence,

so, its dependence for the hydrogen peroxide decomposition has yet to be observed.

Numerical simulations which will be presented later indicate that there is an inverse

dependence of the hydrogen peroxide concentration. It is for this reason that the In plots

curve near the end of the reactions. As the hydrogen peroxide decays, the reactions begin

to accelerate.









v = kobs[Mn(II)][HCO3-]2[H202]x (3-3)

Comparison of Hydrogen Peroxide Reaction Kinetics to Nucleophilic Alkene
Epoxidation Kinetics

In order to better understand the nature of the active catalyst in the manganese

dependent hydrogen peroxide disproportionation, the results of the kinetic investigation

of hydrogen peroxide decomposition need to be compared to the results observed by

Bennett46 for the manganese(II) dependent nucleophilic alkene epoxidation observed in

pure water using sulfonated styrene and 4-vinylbenzoic acid. A summary of her results

for the study of sulfonated styrene are presented here.

Manganese dependence on nucleophilic alkene epoxidation

Bennett46 examined the manganese dependence on the oxidation of sulfonated

styrene in bicarbonate solution (1.00 M) with 1.00 M hydrogen peroxide. The

dependence on manganese was shown to be linear in the range of 0 5.0 [tM, Figure 3-5.



0.01

0.008

S0.006

0.004

0.002



0 1 2 3 4 5

[Mn2+], iM

Figure 3-5. Plot of kobs versus [Mn(II)] observed for nucleophilic alkene epoxidation
(Bennett, 2002)46 y = ((2.09 + 0.25) x103)x, error reported to the 95%
confidence.









These data are consistent with the manganese(II) dependence observed for the

hydrogen peroxide decomposition presented earlier and support the proposal that only

one manganese ion is present in the active catalyst.

Bicarbonate dependence on nucleophilic alkene epoxidation

Bennett46 also examined the bicarbonate dependence on the oxidation of sulfonated

styrene with manganese(II). The dependence of HC03O on the oxidation was shown to be

second-order (Figure 3-6), which is also seen in the manganese dependent hydrogen

peroxide decomposition. Once again, these data indicate that two bicarbonate ions are

present in the active catalyst. Presumably, one of these bicarbonate ions is present as a

peroxycarbonate ion while the other bicarbonate may simply be coordinated to the metal

center.





0.0030
0.0025
-; 0.0020
0.0015
S0.0010
0.0005
0.0000
0.00 0.20 0.40 0.60 0.80 1.00 1.20

[HC 3-]2, M2


Figure 3-6. Plot of kobs versus [HCO3-]2 which shows a second-order dependence.
Reaction conditions: 0.001 Mp-vinyl benzene sulfonate, 1.00 pM Mn2 (*)
0.10 M H202 y = ((2.62 0.17) x10-)x (-) 0.50 M H202y = ((1.19 0.23)
x10-3)x (A) 0.75 M H202 y = ((8.33 0.76) x104)X, errors reported to the
95% confidence. (Bennett, 2002)46










Hydrogen peroxide dependence on nucleophilic alkene epoxidation

Bennett46 also observed the hydrogen peroxide dependence for the oxidation of

sulfonated styrene with manganese(II). The dependence of H202 on the oxidation was

shown to have an inverse relationship with increasing peroxide concentration (Figure 3-

7).

16.00
14.00 A
12.00 -
10.00 A
8.00 -
A M
S6.00 A
M A
4.00 A
2.00 *
0.00 II
0.00 0.20 0.40 0.60 0.80 1.00 1.20
[H202], M

Figure 3-7. Plot of kobs on the [H202]. Reaction conditions: 0.001 M p-vinyl benzene
sulfonate (A) 1.00 M NaHCO3, 0.50 [M Mn2+ (m) 0.75 M NaHCO3, 0.50 |iM
Mn2+ (+) 1.00 M NaHCO3, trace metal catalysis (Bennett, 2002).46

Two possibilities exist for the downward trend observed in the hydrogen peroxide

dependence. One explanation involves the reaction of hydrogen peroxide to form a less

reactive intermediate of the manganese catalyst. This explanation is plausible given the

work by Espenson on the oxidation of nucleophilic alkenes by methyltrioxorhenium

(MTO).24 In the case of MTO, the addition of a second hydrogen peroxide molecule

generates a diperoxo complex which has a slightly lower epoxidation rate constant than

does the monoperoxide intermediate.

The other explanation for the downturn in the reaction has come from work in this

study using numerical simulation to model the decomposition and epoxidation kinetics.

From this work, it appears that hydrogen peroxide may actually inhibit its own










decomposition at higher concentrations. Further details about this possibility will be

presented with the numerical simulations.

Catalyst Lifetime Studies

In addition to examining the kinetics of the manganese(II) catalyzed decomposition

of hydrogen peroxide, the lifetime of the catalyst was also of interest. Stadtman58 noted

that the catalyst lost about half of its activity upon reintroduction of hydrogen peroxide

into a spent decomposition solution. Stadtman,5 unfortunately, gave no indication as to

why the solutions were losing their catalytic ability. Figure 3-8 shows the kobs versus

additions of hydrogen peroxide to a solution of manganese(II) and bicarbonate.



12 -
10 -
S8 *
6 $l
4 4
2
0
0 3 6 9 12 15 18 21 24
# of H202 Additions

Day 1 Day 2 ADay 3

Figure 3-8. Plot of kobs versus # of additions of hydrogen peroxide to a spent solution in
the catalyst lifetime study over multiple days. There is a 16 hr delay before
addition 16 and 24.

As noted by Stadtman,58 the decomposition of peroxide drops by about one-third

upon reintroduction of peroxide, the 2nd data points. The addition of peroxide was

studied over multiple additions, and for multiple days. As can be seen in the graph,

hydrogen peroxide still decomposes even after the spent solution has been sitting for 16

hrs, data points 16 and 24. These data indicate that the catalyst is able to regenerate

simply by the addition of more hydrogen peroxide.










Examining the loss of activity

The first suspected reason for the decrease in activity in the catalyst lifetime studies

was the loss of bicarbonate from the solutions. The bicarbonate concentration was tested

by running a large scale reaction (10 mL) with 1.00 M bicarbonate, 1.00 M hydrogen

peroxide and 5 [tM manganese(II). The reaction was cycled a total of 10 times, each time

bringing the hydrogen peroxide concentration back to 1.00 M. After the last reaction had

decomposed the hydrogen peroxide, the bicarbonate concentration was analyzed by the

standard barium chloride precipitation method. It was found that after 10 cycles, the

bicarbonate concentration had dropped from 1.00 M to about 0.50 M.

This result indicates that one of the reasons the decomposition of hydrogen

peroxide is decreasing in the catalyst lifetime study is that the concentration of

bicarbonate is decreasing. Since the bicarbonate dependence has been observed to be

second-order, the loss of bicarbonate will have a dramatic effect on the observed

decomposition rate constant. In addition to examining the bicarbonate concentration, the

hydrogen peroxide was examined as a possible source for the loss of activity.

In addition to the loss of bicarbonate, it is known that hydrogen peroxide is

stabilized using tin phosphates.10 The loss of activity may be due to manganese

precipitation by the addition of phosphates to the solutions, as was observed when

phosphates were used in attempting to control the pH of the decomposition solutions.

The malachite green assay for phosphates was used to examine the stock 50% hydrogen

peroxide solution. It was found that the stock hydrogen peroxide contained

approximately 4 mM phosphate. For the cycles being run, this equates to about 25 [tM of









phosphate being added to each cycle. This amount of phosphate is enough to begin

manganese precipitation.

Multiple additions of distilled hydrogen peroxide and solid sodium bicarbonate

Once it was determined that bicarbonate was being lost during each cycle and

phosphate being added, another catalyst lifetime study was done. For these reactions,

distilled hydrogen peroxide was used. This assured that no phosphates were being added

to the solutions. In addition, solid sodium bicarbonate was added before each cycle in an

attempt to stabilize the bicarbonate concentration from cycle to cycle. The results of the

catalyst lifetime study using these modifications are seen in Figure 3-9.


5
4.5 $
4 .- $ m
3.5 -
M3M

2.5
S2 -
1.5
1-
0.5
0 -
0 3 6 9 12 15
# of H202 additions

Day 1 m Day 2

Figure 3-9. Plot of kobs versus # of hydrogen peroxide additions for the catalyst lifetime
study using distilled hydrogen peroxide and adding solid sodium bicarbonate.
The loss of activity is now due only to dilution and the inability to maintain
the bicarbonate concentration at a constant value.

The reactions are about half as fast as those not using the distilled hydrogen

peroxide due to the metal contaminants that are found in the peroxide. The loss of









activity is now due to dilution and the inability to maintain a constant bicarbonate

concentration. However, it must be noted that there is not as significant a loss in activity

when using distilled peroxide and adding bicarbonate. This indicates that the manganese

catalyst is not destroyed during the decomposition of the peroxide, but can be regenerated

by the addition of more hydrogen peroxide.

Studies of the Manganese Source

In addition to examining the lifetime of the manganese catalyst, the question of

whether the source of the manganese was important needed to be answered. For all

kinetic experiments, the manganese source was manganese(II) sulfate. Two additional

sources of manganese, permanganate and a Mn(IV)-TACN catalyst, were tested to

compare their decomposition of hydrogen peroxide to the decomposition with

manganese(II) sulfate.

Potassium permanganate

First, potassium permanganate was tested at a concentration of 3 and 4 [tM in the

presence of 0.20 M sodium bicarbonate and 0.100 M hydrogen peroxide. As seen in

Table 3-1, the observed first-order rate constants for the manganese(II) sulfate and

permanganate at the same concentration are well within experimental limits.

Table 3-1. Comparison of observed rate constants for the decomposition of hydrogen
peroxide, 0.100 M, in 0.20 M sodium bicarbonate with 3 and 4 [tM
manganese(II) and permanganate. Errors reported are to the 95% confidence.
Manganese Source Concentration kobs, s-1
Mn(II) 3 |tM (8.33 0.27) x10-4
MnO4- 3 tM (9.19 0.31) x10-4
Mn(II) 4 tM (1.15 0.04) x10-3
MnO4- 4 tlM (1.18 0.04) x10-3










[MnIV(Me3TACN)(OMe)3]PF6

In addition to permanganate, a Mn(IV)-TACN catalyst was synthesized for use as

the manganese source in the hydrogen peroxide decomposition. In 1996, Kerschner et

al.63 synthesized [MnIV(Me3TACN)(OMe)3](PF6), where Me3TACN is 1,4,7-trimethyl-

1,4,7-triazacyclononane (Figure 3-10). This catalyst was capable of oxidizing water-

soluble olefins, specifically 4-vinylbenzoic acid and styrylacetic acid, in the presence of

bicarbonate. The stability of the catalyst was demonstrated by the repeated additions of

hydrogen peroxide and alkene to produce epoxidized product.63

C4


C3 C45

C54




Mnl N5



C C65

C6
Figure 3-10. Molecular structure of [MnlV(Me3TACN)(OMe)3](PF6).63

The Mn(IV) catalyst was synthesized using the procedure reported by Kerschner.63

The compound is obtained by the reaction of manganese(II) chloride with 1,4,7-

trimethyl-l,4,7-tiazacyclononane in methanol in the presence of sodium peroxide. The

complex is crystallized from methanol/water as a brown hexafluorophosphate salt.

Mn(IV) catalyst stability

Initially, the stability of the Mn(IV) catalyst was examined spectrophotometrically

in both the presence of bicarbonate and hydrogen peroxide individually. A solution of









1.00 M sodium bicarbonate and 0.108 M Mn(IV) catalyst was dissolved in water, and the

catalyst was monitored at 345 nm for 2 hrs, Figure 3-11. The same procedure was also

done in the presence of 0.50 M hydrogen peroxide, Figure 3-12.


1.2



0.8

0.6

0.4

0.2

0 -
0 1200 2400 3600 4800 6000 7200

Time, sec

Figure 3-11. Stability of the [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence
of 1.00 sodium bicarbonate.


1.2



0.8

0.6

0.4

0.2

0
0 1200 2400 3600 4800 6000 7200

Time, sec

Figure 3-12. Stability of the [MnlV(Me3TACN)(OMe)3](PF6) at 345 nm in the presence
of 0.50 M hydrogen peroxide.









After the catalyst was shown to be stable in the presence of bicarbonate and

hydrogen peroxide alone, experiments were conducted to test its stability in the presence

of bicarbonate and hydrogen peroxide together. A solution of 1.00 M sodium

bicarbonate and 0.108 M catalyst was dissolved in water. After about 350 seconds of

monitoring, 25 ptL (0.100 M, final concentration) of hydrogen peroxide was added to the

solution (Figure 3-13). Almost instantly, the catalyst absorbance decays to 0. Within

about 3 minutes, a new absorbance is detected, and the solution is a yellow color. This

new absorbance is very broad, having an absorbance of -0.7 from -250 nm to 500 nm,

Figure 3-14. This absorbance is most likely due to the metal interaction with the N-

dealkylated organic decay products. More about the topic of N-dealkylation by Mn(II)

with hydrogen peroxide in the presence of bicarbonate will be presented in a later section

of this chapter.


1.4
1.2
< 1-
a 0.8
0.6 -
0.4
0.2
0
0 100 200 300 400 500 600 700
Time, seconds


Figure 3-13. Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm.
Addition of 25 pL (0.100 M, final concentration) hydrogen peroxide was done
at 350 seconds. The absorbance first decays to 0 and within a matter of
minutes, the solution is bright yellow.













0.8

0.2


Figure 3-14. UV-vis specta of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6). The solid line
is the spectrum in the presence of 1.00 M sodium bicarbonate. The dotted line
is the spectrum of the solution after 1 eq of hydrogen peroxide was added.

In a second experiment, 50 ptL (0.200 M, final concentration) hydrogen peroxide

was added to the catalyst to determine whether the developing yellow products) was the

result of a limited amount of peroxide. A solution of 1.00 M sodium bicarbonate and

0.108 M catalyst was dissolved in water. The absorbance at 345 nm was monitored as

before (Figure 3-15). At 312 seconds, the peroxide was added to the solution. Unlike the

first experiment, the development of the yellow color did not occur. It is therefore

apparent that in the presence of only 1 equivalent of hydrogen peroxide, the catalyst does

not completely decay, allowing the yellow color to develop. When 2 equivalents of

hydrogen peroxide are present, the catalyst is able to completely decay.


1.2


0.8
0.6 -
I 0.4
0.2 -


0 100 200 300 400 500 600 700
Time, seconds

Figure 3-15. Monitoring of 0.108 M [MnIV(Me3TACN)(OMe)3](PF6) at 345 nm.
Addition of 50 ptL (0.200 M, final concentration) hydrogen peroxide was done
at 312 seconds. Even after 6 minutes, the yellow color does not develop.









Finally, the [MnIV(Me3TACN)(OMe)3](PF6) catalyst was used with bicarbonate to

decompose hydrogen peroxide. As was done with permanganate, 3 and 4 [tM

[MnIV(Me3TACN)(OMe)3](PF6) was dissolved in water with 0.200 M sodium

bicarbonate and the decomposition was initiated by the addition of 25 p.L (0.100 M)

hydrogen peroxide. As expected, the observed rate constants for the

[MnIV(Me3TACN)(OMe)3](PF6) catalyzed decomposition of hydrogen peroxide are

similar to those for the Mn(II) ion decompositions (Table 3-2). These results indicate

that the [MnIV(Me3TACN)(OMe)3](PF6) catalyst quickly decomposes to release the

manganese ion, which then begins catalytically decomposing the hydrogen peroxide. If

the catalyst did not quickly decompose, a lag in the decomposition of hydrogen peroxide

might have occurred, however, this is not seen experimentally.

Table 3-2. Comparison of observed rate constants for the decomposition of hydrogen
peroxide (0.100 M, final concentration) in 0.20 M sodium bicarbonate with 3
and 4 [iM manganese(II) and [MnIV(Me3TACN)(OMe)3](PF6). Errors
reported are to the 95% confidence.
Manganese Source Concentration kobs, S-1
Mn(II) 3 tM (8.33 0.27) x10-4
[MnIV(Me3TACN)(OMe)3](PF6) 3 [tM (8.79 0.30) xl0-4
Mn(II) 4 iM (1.15 0.04) x10-3
[MnIV(Me3TACN)(OMe)3](PF6) 4 MtM (1.11 0.06) xl0-3


Cis-trans Isomerization in the Manganese(II) Catalyzed Alkene Epoxidation

Reactions of cis-alkenes to their corresponding epoxides catalyzed by

manganese(II) and hydrogen peroxide in bicarbonate solution will give some indication

as to the nature of the oxygen transfer from the active oxygen species to the alkene. For

example, in the case of peracid epoxidation of alkenes, cis-alkenes react to give only the

cis-epoxide, as seen in Figure 3-16. The oxygen is delivered in a concerted process

which retains the stereochemistry of the reactant.









0
Cl 0

R \ H- O H C1
..." "'O HO
HOG
+ --
R--O R"
R'

Figure 3-16. The concerted mechanism for the m-CPBA oxidation of nucleophilic
alkenes resulting in the retention of stereochemistry.

The reaction of Mn(salen) organometallic complexes with hydrogen peroxide, on

the other hand, do not produce only the cis-epoxide from the cis-alkene, but the trans-

epoxide as well. For these reactions, a two step process occurs where by the C-C sigma

bond remains intact, but a carbon radical is produced as an intermediate, as seen in Figure

3-17. During the lifetime of the intermediate carbon radical, before the oxygen bond of

the epoxide is formed, the C-C sigma bond has the opportunity to rotate into the more

stable trans conformation. In this way, the Mn(salen) catalysts will produce both the cis

and trans-epoxides from the cis-alkene.

Mn(Salen) + H202
R R, o R,

SMn(salen) -* + Mn(salen)
R' R'! )R'


Figure 3-17. The stepwise oxidation of an alkene by Mn(salen) and hydrogen peroxide is
shown. Cis/trans isomerization occurs in the transition state, where the C-C
sigma bond is able to rotate into the more stable trans configuration.

Burgess et al.40 noted that under their conditions, using water/DMF solutions, the

epoxidation of cis-stilbene produced both the cis and trans-stilbene oxides, as seen in

Figure 3-18. This indicates that the active oxygen species does not add the oxygen in a

concerted manner, as do the peracids, for example. If the oxygen were added in a

concerted manner, there would be no trans epoxide present. However, Burgess followed










these experiments in mixed solvent. The question remains as to whether the cis/trans

isomerization will occur in pure aqueous solution.


0

DMF/H20 +
d\ Mn(II), H202, HCO3


Figure 3-18. The cis/trans isomerization noted Burgess in his epoxidation of stilbene
using the Mn(II), H202, bicarbonate system using a mixed solvent system of
DMF/H20. (Burgess, 2002)40

To study the cis/trans isomerization in pure water, 4,4'-sulfonated stilbene was the

obvious choice, based on Burgess' use of stilbene for the reactions in mixed solvent.

Unfortunately, all attempts at direct synthesis by sulfonating stilbene using fuming

sulfuric acid resulted in black tar. A literature search for the preparation of 4,4'-

sulfonated stilbene resulted in a single paper by van Es.64 The synthetic scheme is

illustrated in Figure 3-19. All attempts at synthesizing the 4,4'-sulfonated stilbene failed.

N

NH2 N+ o
N A
Na2CO3 OH
NaNO2 NaO3S 1000C NaO3S
HC1 NaOH,CuC12,H20 NaC1
H20 300 C, 2hrs
SO3H SO3H
Figure 3-19. Synthetic scheme for synthesis of 4,4'-sulfonated stilbene. (van Es, 1964)64

Next, two water-soluble alkenes were chosen, cis and trans-2-butene-1,4 diol.

These alkenes were chosen for two reasons. First, they were freely water soluble at the

operating pH of 8.4. Second, the epoxides of the alkenes are easily distinguishable by

13C NMR. This made analysis of the reactions relatively simple. The cis-alkene is

commercially available, while the trans-alkene must be synthesized. The trans-2-butene-

1,4-diol was synthesized using the method of Schloss and Hartman.65 The synthesis








requires the reduction of 2-butyne-1,4-diol by lithium aluminum hydride in THF. The

epoxidation of both the cis and trans alkenes were accomplished by the reaction with m-

CPBA. The epoxide products' H and 13C NMR were compared to literature values for

authentic samples, Figure 3-20.


0

HO C- OH
Cb
cis-2,3 -epoxybutane- 1,4-diol

Ca




Cb


0

HO C OH
Cb
trans-2,3-epoxybutane- 1,4-diol

Ca


SCb


MeOH


80 60 40 80 50 40
Figure 3-20. 13C NMR of cis and trans-2,3-epoxybutane-1,4-diol in D20 using methanol
as an internal standard.









Cis/Trans isomerization reactions with cis-2-butene-1,4-diol

The cis-2-butene-1,4-diol (0.60 M) was dissolved in D20 along with 1.00 M

NaHCO3, and 10 [LM Mn(II). The reaction was initiated by the addition of hydrogen

peroxide (final concentration 6.0 M). The reaction was monitored by observing the

methylene peak in the 13C NMR using methanol as an internal standard. After 30

minutes, the methylene peak has decreased, but the appearance of the epoxide peaks for

either the cis or trans epoxide cannot be seen (Figure 3-21, left). After 18 hrs (Figure 3-

21, right), it appears that the cis-alkene has been oxidized to any number of products,

none of which have been identified at this time. From these data, a new water-soluble

alkene was needed to examine the cis/trans isomerization of the Mn(II)/hydrogen

peroxide/bicarbonate oxidation system in water.



CH2
MeOH
/ MeOH









8Y 4 80 a 40
30 minutes 18 hours
Figure 3-21. Epoxidation of cis-2-butene-l,4-diol (0.60 M) with 1.00 M HCO3-, Mn(II)
(10 [M), and H202 (6.00 M) after 30 minutes (left) and 18 hrs (right).

Cis/Trans isomerization of maleic and fumaric acids

Two new alkenes were chosen to study the cis/trans isomerization in pure water.

These alkenes are maleic and fumaric acid. The structures of these alkenes at the










operational pH of 8.4 are shown in Figure 3-22. While these may appear to be

electrophilic alkenes, the dominant resonance structure at the operating pH allows these

alkenes to react as nucleophilic alkenes. More on this topic will be presented in the

following chapter. Once again, the determination of epoxide products are conveniently

made using 1H NMR, as seen in Figure 3-23.

0



0 Y---v 0 0 0",/ 0



0 0 0
Maleic Acid Fumaric Acid

Figure 3-22. The structures of maleic and fumaric acids at the operating pH of 8.4.

o

0 -"-k \
DOH 0 O O- DOH 0 0

0 0 0
O O O
Maleic Acid Oxide Fumaric Acid Oxide







MeOH
MeOH




Figure 3-23. H NMR of maleic and fumaric acid oxides in D20 using methanol as an
internal standard.

When 0.10 M maleic acid was allowed to react with 1.00 M peroxide in the

presence of 0.80 M sodium bicarbonate and 10 [tM Mn(II), a 34% conversion to epoxide

was observed by NMR in 15 min, as seen in Figure 3-24.













DOH

MeOH







MA








FAO


MAO


LA\


6


Figure 3-24. Epoxidation of maleic acid by hydrogen peroxide and manganese(II) in the
presence of bicarbonate after 15 min. Reaction conditions: 0.10 M maleic
acid, 1.00 M H202, 0.80 M NaHCO3, and 10 gM Mn(II).
Of the 34% epoxide formed, 74% was the fumaric acid oxide and 26% was maleic acid

oxide. This result indicates that even in pure water, the oxygen is not being added to the


-1 .. .









alkene in a concerted mechanism, such as that seen in the oxidation by m-CPBA. This

indicates that at some time during the epoxidation, the C-C sigma bond of the alkene has

the opportunity to rotate into the more stable trans conformation before closure of the

epoxide ring. This experiment does indicate that radical or carbocation formation is

probable in the epoxidation of nucleophilic alkenes, a discussion of possible routes for

the addition of the oxygen and rotation into the trans isomer will be discussed with the

possible mechanisms of the reaction.

The observation that cis alkenes react with the active oxygen species to give the

trans epoxide indicates that radical chemistry may play a role in the epoxidation of

alkenes. This does not, however, indicate that free radicals are responsible for the

epoxidation. Reaction mechanisms that are similar to those for Mn-salen epoxidation

catalysts could also explain the rotation about the C-C bond during the epoxidation

reaction. In addition, reactions involving electrophilic alkenes, which will be presented

in the next chapter, led us to question Sychev's proposed hydroxyl radical mechanism. In

the epoxidation of electrophilic alkenes, as was the case for the nucleophilic alkene

epoxidations, the reactions cleanly yielded the epoxide products with no indication of any

radical products.66'67 In an attempt to exclude free radicals as a possible reaction

pathway, an examination of the radical traps used by Sychev47 was conducted.

Examination of Sychev's Radical Trap Experiments

As discussed in the introduction, the combination of hydrogen peroxide and

iron(II), Fenton's reagent, is a useful method for the production of hydroxyl radicals.

Fenton's reagent can then be used to oxidize organic molecules. For instance, benzene

can be oxidized to form biphenyl and phenol in the presence of Fenton's reagent, as seen

in Figure 3-25.67








Fe2+ + H202 Fe3+ + HO+ HO-
OH
H OH
HIOO +- HO*
HO'+ HO+ H20


HH H HO~ X H'acceptor + H
+0 H \H/ _0 /+H20
H---HH -0
Figure 3-25. Fenton's reagent can be used to oxidize benzene to phenol and biphenyl.

In 1977, Sychev et al.47 began investigating the role of manganese(II) in the

disproportionation of peroxide in bicarbonate buffered solutions. His assumption was

that manganese(II) reacted similarly to iron(II) ions in Fenton type chemistry. Following

this assumption, a hydroxyl radical based mechanism was proposed, shown by Equations

(3-4)-(3-14), the sum of which is the decomposition of hydrogen peroxide to molecular

oxygen and water. As with Fenton type chemistry, free hydroxyl and peroxy radicals are

formed in this mechanism, along with the carbonate radical anion.

[Mn(HCO3)2] + H202 [Mn(HC03)2]+ + OH- + 'OH (3-4)

[Mn(HCO3)2]+ + H202 [Mn(HCO3)2] + H+ + HOO" (3-5)

[Mn(HCO3)2] + 'OH [Mn(HCO3)2] + OH- (3-6)

[Mn(HCO3)2] + HOO- [Mn(HCO3)2] + 02-' + H+ (3-7)

[Mn(HCO3)2] + HOO' [Mn(HCO3)2] + HOO- (3-8)

[Mn(HCO3)2] + OH- -- [Mn(HCO3)2] + 'OH (3-9)

'OH + H202 02-* + H+ + H20 (3-10)

202-" + 2H202 202 + 20H- + 2 'OH (3-11)

'OH + HOO'-H20 + 02 (3-12)

'OH + HCO3- CO3- + H20 (3-13)









CO3- + H22 -- 02'+ H+ + HC03- (3-14)

2H202 2H20 + 02

In order to support his claim of a free hydroxyl radical pathway, Sychev employed

a set of experiments using N,N-dimethyl-4-nitrosoaniline (DMNA) as a free hydroxyl

radical trap. In his experiments, he studied the production of 02(g) as a function of time

with increasing amounts of DMNA, Figure 3-26.


V02, mi











// /






BOI /Iff 30jl seconds
Figure 3-26. The influence of inhibitors on the catalase process in the Mn(II)/HCO3/H202
system. [Mn(II)] = 4 x 10-6 M, [H202] = 0.10 M, pH 7.0, [HCO3-] = 0.4 M,
and T = 25 C: 0) kinetic curve with no inhibitors; 1), 2), 3), and 6) in the
presence of DMNA as the inhibitor(at concentrations of 1 x 10-5,1.5 x 10-5, 2 x
10-,and 4 x 10-5 M respectively; 4)in the presence of tetranitromethane(4 x
10-5 M); 5) in the presence of hydroquinone (1.5 x 10-5 M); 7) decomposition
of H202 without Mn(II) ion (blank experiment). (Sychev, 1977)47

As expected for his free hydroxyl radical pathway, Sychev observed that when

DMNA was present, the production of 02(g) was suppressed to the background

disproportionation of H202 without the addition of metal, Figure 3-26, line 7. As time









progressed, the 02(g) production would begin to increase back to the purely catalytic

production of 02(g), as shown in Figure 3-26, line 0.

The conclusion reached by Sychev was that the DMNA was trapping the free

hydroxyl radicals produced from the disproportionation, Equations 3-4, 3-9, and 3-11.

Without the presence of free hydroxyl radicals to carry the reaction, O2(g) production

would be the same as the uncatalyzed 02(g) production, Figure 3-26, line 7. Eventually,

02(g) production would begin to follow that of the catalyzed reaction as the concentration

of DMNA was reduced and an increase in free hydroxyl radicals occurred. This is seen

in Figure 3-26 as all of the inhibited reactions eventually begin to produce 02(g) at the

same rate as the uninhibited reaction, Figure 3-26, line 0.

In all of Sychev's papers, the use of radical traps, such as DMNA, provide the

entire basis for a free hydroxyl radical mechanism. In none of his papers, however, did

Sychev identify the organic products of the reactions with DMNA. In the current study,

it has been hypothesized that instead of a decomposition pathway that requires hydroxyl

radicals, the mechanism of peroxide disproportionation may proceed through a high

valent metal oxo species, or by carbonate radical anions, which are proposed in Sychev's

model. The data presented for the interruption of hydrogen peroxide decomposition by

Sychev could be the result of trapping of carbonate radical anions, instead of hydroxyl

radicals. A discussion on the reactivity of carbonate radicals will be presented in the

proposed mechanism for the oxidation of the radical traps.

The use of a high valent metal oxo species could also explain Sychev's loss in

02(g) production seen with the use of DMNA. Instead of the 02(g) production being

inhibited by radical interruption, the oxygen normally being released as molecular









oxygen would instead be transferred to DMNA. In order to understand the effect DMNA

is having on the hydrogen peroxide decomposition reaction, the organic products from

the reaction need to be identified. Once the oxidized organic products are identified, a

clearer understanding of the reaction mechanism may be possible.

A series of experiments were first conducted to determine what the oxidized

products of DMNA could be. Initially, potassium peroxymonosulfate was employed as

the oxidant. This experiment was done to determine what the product of a pure oxygen

transfer would be, since peroxymonosulfate is an excellent electrophilic oxidant.68'69

When peroxymonsulfate was allowed to react with DMNA in a 1:1 molar ratio, N,N-

dimethyl-4-nitroaniline was produced nearly quantitatively after 30 minutes, as expected

for an oxygen transfer to the nitroso moiety. A second control experiment was done

using H202 and HC03-, only. When one equivalent of H202 was added to 0.40 M

chelexed HCO3-, N,N-dimethyl-4-nitroaniline was produced after 1 hr. In the presence of

hydrogen peroxide alone, no reaction was detected even after 24 hrs. This result

indicates that solutions of peroxycarbonate are able to convert the nitroso moiety to the

nitro without the addition of any metals. Any products, other than the nitro compound,

are then the result of the addition of the metal cations.

0
O N NO2


HS05




N N

Figure 3-27. The reaction of N,N-dimethyl-4-nitrosoaniline with peroxymonosulfate to
yield N,N-dimethyl-4-nitroaniline.









In a second set of reactions, DMNA was oxidized using H202 in bicarbonate with

Mn(II). H202 was used as the terminal oxidant in 50x molar excess over the DMNA.

The need to increase the H202 concentration to such a degree over the DMNA

concentration is due to the fact that H202 disproportionation is much faster than the

oxidation of DMNA. When reactions using only one equivalent of H202 were conducted,

starting material was the only recovered organic compound. The final sodium

bicarbonate and Mn(II) concentrations were set at 0.40 M and 4.0 aM, respectively.

DMNA oxidations were conducted by first dissolving the organic substrate in a

mixture of CH3CN:H20 (30:70 (v: v)) with a solution of the manganese(II) sulfate.

Equilibrated solutions of H202 and sodium bicarbonate were then slowly added dropwise

over about 30 minutes. Reactions were considered to be complete when the production

of 02(g) from the H202 disproportionation ceased. This was usually 10-15 minutes after

the final addition of the H202/HC03- solution. Since the reaction is highly exothermic,

ice was often employed to keep the temperature from exceeding 65C. Once the reaction

reaches 65C, the H202 disproportionation becomes vigorous enough to cause the

reaction mixture to boil out of the reaction flask. At the end of the reaction, the organic

products were extracted into chloroform which was then dried over magnesium sulfate.

The solvent was then removed under reduced pressure to give crude product, which was

analyzed by 1H NMR, Figure 3-28.

The 1H NMR of the crude product shows multiple products, three of which have

been identified at this time. Aromatic peaks still remained and retained the characteristic

proton signal for a disubstituted aromatic compound. It was also noted that new peaks in









the 4 5.5 ppm region had appeared. The methyl peaks of the amine portion of the

molecule seemed to have remained, but they were shifted upfield.


9 8 7 6 5 4 3 2 1
Figure 3-28. H NMR of the crude reaction mixture after an oxidation of N,N-dimethyl-4-
nitrosoaniline by hydrogen peroxide in the presence of bicarbonate and
Mn(II). Reaction conditions: N,N-dimehtyl-4-nitrosoaniline (1 g, 6.66 mmol),
0.400 M sodium bicarbonate, 10 [tM Mn(II), 6.64 M H202, 1 hr.

After analysis by 1H NMR, gas chromatography was employed to help determine

the number of products obtained from the reaction. GC analysis showed there were four

volatile products formed. It can only be stated that these products are volatile, since GC

will only separate compounds that can be volatized and have a low enough boiling point

to remain in the gas phase. Both 1H NMR and GC proved that none of the products

obtained were that of N,N-dimethyl-4-nitrsoaniline, which should have been the case









since the nitroso moiety is easily oxidized by peroxycarbonate. Figure 3-29 is a GC trace

using the same method employed for the crude reaction mixture, Figure 3-30, for which a

peak is not observed at 14.6 min, indicating that all of the N,N-dimethyl-4-nitrsoaniline

has been converted to other organic products.

14.637 min




































Figure 3-29. GC trace for a standard ofN,N-dimethyl-4-nitrsoaniline. Non-linear
gradient for 30 minutes, detection by FID.









15.260 min


17.462 min


-JL ;


19.983 min







i


24.421 min


II-i I I '11 lli -- T F P T
T- ~ N- C4 N C% CN (N I CN CN c%;
Figure 3-30. GC trace for the crude reaction material from the oxidation ofN,N-dimethyl-
4-nitrsoaniline from Figure 3-25. Lack of a peak near 14.637 min proves that
no starting material remains. GC conditions: non-linear gradient for 30
minutes, Detection: FID.


J









After 1H NMR and GC analysis, the reaction material was applied to a silica

column employing chloroform as eluant. Fractions were collected and analyzed by GC.

Fractions 4 and 9 were found to contain organic products and their GC traces and 1H

NMR are shown in Figure 3-31 and 3-32. The two compounds that were separated were

found to be N,N -dimethyl-4-nitroaniline and 4-nitroaniline, by comparison with

authentic samples. These two compounds account for about 80% of the crude reaction

mixture, the last 20 % being the peaks at 19.983 and 24.421 min, Figure 3-27. The two

remaining organic compounds remained at the top of the column.




17.418 min NO2 H

Ha



Hb

N He















S -- F .-' F -
S7 s 4 3 2 1 ppm

Figure 3-31. GC trace (left figure) and 1H NMR (right figure) for Fraction 4 of the silica
column. Identification of the product as N,N-dimethyl-4-nitroaniline was
confirmed by comparison with a GC trace and 1H NMR of an authentic
sample.











CHC13
15.182 mi NO

Ha Ha


Hb

NH










H




8 7 B 5 a

Figure 3-32. GC trace (left figure) and 1H NMR (right figure) for Fraction 9 of the silica
column. Identification of the product as 4-nitroaniline was confirmed by
comparison with a GC trace and 1H NMR of an authentic sample.

Once 4-nitroaniline was identified as a product, it was apparent that the amine was

dealkylating. It was then likely that one of the other unidentified peaks was N-methyl-4-

nitroaniline. A 1H NMR was acquired of a sample of the pure material and compared

with the crude reaction material, Figure 3-33. It was found that N-methyl-4-nitroaniline

peaks were also present in the crude reaction mixture. The only peak not identified was

the peak at 24.421 min, which accounts for less than 5 % of the crude reaction material.












NO2 Hd

Ha



Hb

Hd NH,



Hb


H



CHC13



He


8 7 6 5 4 3
Figure 3-33.1H NMR of an authentic sample of N-methyl-4-nitroaniline. Comparison
with the crude reaction mixture confirms its presence as a product.

For these products to be observed, N-dealkylation must be responsible for cleaving

the C-N bond of the amines. In the case of DMNA, the cleaved carbon group is

formaldehyde based on current literature that will be discussed shortly. In the 1H NMR

and GC analysis performed on the DMNA reactions, formaldehyde was never detected.

This was probably due to the fact that the high temperature of the reaction volatilized

formaldehyde. In order to observe the aldehyde produced from the reaction, N,N-diethyl-

4-nitrosoaniline (DENA) was chosen as the next substrate.

Reactions were performed using a 50x molar excess of H202 over DENA and the

final concentrations of bicarbonate and Mn(II) were 0.40 M and 4.0 pM, respectively. As