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

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