MECHANISTIC INVESTIGATIONS ON THE ACTIVATION OF PEROXIDES
BY MANGANESE COMPOUNDS
MECHANISTISCHE UNTERSUCHUNGEN ZUR AKTIVIERUNG VON PEROXIDEN
AN MANGANVERBINDUNGEN
Der Naturwissenschaftlichen Fakultät der
Friedrich-Alexander-Universität Erlangen-Nürnberg
zur
Erlangung des Doktorgrades Dr. rer. nat.
vorgelegt von
Sabine Rothbart
aus Nürnberg
Als Dissertation genehmigt von der naturwissenschaftlichen Fakultät der
Friedrich-Alexander-Universität Erlangen-Nürnberg
Tag der mündlichen Prüfung: 26.04.2012
Vorsitzender der Promotionskomission: Prof. Dr. Rainer Fink
Erstberichterstatter: Prof. em. Dr. Dr. h. c. mult. Rudi van Eldik
Zweitberichterstatterin: Prof. Dr. Ivana Ivanović-Burmazović
Die vorliegende Arbeit entstand in der Zeit von September 2007 bis Februar 2012 am Department
Chemie und Pharmazie der Friedrich-Alexander-Universität Erlangen Nürnberg.
DANKSAGUNG
Mein besonderer Dank gilt meinem Doktorvater Prof. em. Dr. Dr. h. c. mult. Rudi van Eldik für
das große Interesse an meiner Arbeit und für zahlreiche wissenschaftliche Diskussionen.
Außerdem möchte ich ihm besonders herzlich für die Möglichkeit danken, Teil seiner einmaligen
Arbeitsgruppe und der Zaubervorlesung sein zu dürfen.
Großen Dank schulde ich auch meinen Eltern, Sigrid und Günther Rothbart, für ihre konstante
Unterstützung und ihren unerschütterlichen Glauben an mich.
Des Weiteren möchte ich mich bei folgenden Leuten bedanken, die zum Gelingen dieser Arbeit
beigetragen haben: Dr. Erika Ember für die Zusammenarbeit, Dr. Ralph Puchta für die DFT
Rechnungen, Dr. Achim Zahl für NMR Messungen, Oliver Tröppner für die
massenspektrometrischen Messungen, sowie Prof. U. Zenneck, Dr. Susanne Mossin und Dr. Jörg
Sutter für die Einführung am EPR-Spektrometer. Prof. D. Chatterjee und seinem Team danke ich
für die Durchführung der HPLC Analyse und gewinnbringende Diskussionen.
Natürlich gebührt auch allen aktuellen und ehemaligen Mitgliedern der Arbeitsgruppe van
Eldik ein herzliches und großes Danke für die angenehme Arbeitsatmosphäre, die stete
Hilfsbereitschaft und den unglaublichen Zusammenhalt. Dies gilt insbesondere für: Ariane,
Christoph, Matthias, Peter, Steffi, Lars, Raquel, Simon und Svetlana. Ihr tragt erheblich dazu bei,
dass mir die Zeit meiner Promotion immer in besonders guter Erinnerung bleiben wird. Außerdem
danke ich den fleißigen Zaubervorlesungs-Mitarbeitern der Arbeitskreise Burzlaff und Ivanović-
Burmazović, sowie der Kaffeeraum-Crew für die gute Atmosphäre und Anita Schmitz für die
schöne Zeit und gute Zusammenarbeit im Praktikum „AC-Explodieren“.
Zu guter Letzt möchte ich besonders meinem Christoph sehr herzlich für die stete
Unterstützung und seelische Kraft danken, die er mir während dieser Zeit gegeben hat.
PUBLICATIONS AND CONFERENCE CONTRIBUTIONS
PUBLICATIONS
1. Erika Ember, Sabine Rothbart, Ralph Puchta and Rudi van Eldik, “Metal – ion
catalyzed oxidative degradation of Orange II by H2O2. High catalytic activity
of simple manganese salts”, New J. Chem., 2009, 33, 34-49. The manuscript
is featured on the cover of the January 2009 issue of the New Journal of
Chemistry.
2. Sabine Rothbart, Erika Ember and Rudi van Eldik, “Comparative study of the catalytic activity of
[MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 in the H2O2 induced oxidation of organic dyes
in carbonate buffered aqueous solution”, Dalton Trans., 2010, 39, 3264-3272.
3. Erika Ember, Hanaa Gazzaz, Sabine Rothbart, Ralph Puchta and Rudi van Eldik, “MnII – a
fascinating catalyst: Mechanistic insight into the catalyzed oxidative degradation of organic
dyes by H2O2”, Appl. Catal. B., 2010, 95, 179-191.
4. Sabine Rothbart, Erika Ember and Rudi van Eldik, “Mechanistic studies on the oxidative
degradation of Orange II by peracetic acid catalyzed by simple manganese(II) salts. Tuning the
lifetime of the catalyst”, New J. Chem. 2012, 36, 732-748.
5. Sabine Rothbart and Rudi van Eldik, “Manganese compounds as versatile catalysts for the
oxidative degradation of organic dyes”, Adv. Inorg. Chem. 2012, 65, submitted.
ONGOING PROJECT
Sabine Rothbart and Rudi van Eldik, “High catalytic activity of a Mn-terpy compound in
oxidative dye degradation by peracetic acid”, in preparation.
CONFERENCES AND WORKSHOPS
POSTER. Oxidative degradation of Orange II by H2O2 catalyzed by divalent transition metals,
Inorganic Reaction Mechanism Group Meeting (IRMG-36), March 2007, York, England.
POSTER. Catalytic activity of [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 in the H2O2 induced
oxidation of organic dyes, Inorganic Reaction Mechanism Group Meeting (IRMG-39), January
2010, Kloster Banz, Germany.
POSTER. MnII – a fascinating oxidation catalyst: Mechanistic insight into the catalyzed oxidative
degradation of organic dyes by H2O2, Inorganic Reaction Mechanism Group Meeting (IRMG-
39), January 2010, Kloster Banz, Germany.
ORAL PRESENTATION. Catalytic activity of [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 in the
H2O2 induced oxidation of organic dyes, Erlangen-Kraków-Workshop on “Understanding the
mechanisms of chemical processes”, May 2010, Kraków.
LIST OF ABBREVIATIONS
A absorbance
AcOH acetic acid
B / mT magnetic field in millitesla
bpy 2,2’-bipyridine
BHT 2,6-di-tert-butyl-4-methylphenol
tBuOH tert-butanol
c concentration
CHES 2-(cyclohexylamino)ethanesulfonic acid
CV cyclovoltammogram
D value of the zero field splitting parameter
DFT density functional theory
DMSO-d6 hexadeuterodimethyl sulfoxide
E / V potential in volts
EPR electron paramagnetic resonance spectroscopy
ES-MS electrospray mass spectrometry
g g-faktor
h Plank constant (6.62606885 · 10-34 J·s)
h hour
HEPES 4-(2-hydroxyethyl)piperazine-1-ethanesulfonic acid
I / A current in ampère
I nuclear spin
K equilibrium constant
k rate constant
kobs observed rate constant
λ wavelength
LMCT ligand-to-metal charge-transfer
nm nanometer
NMR nuclear magnetic resonance
ν frequency
Mo Morin; 2’,3,4’,5,7-Pentahydroxyflavone
M mol/l
Me3tacn 1,4,7-trimethyl-1,4,7-triazacyclononane
OII Orange II; 4-(2-hydroxy-1-naphthylazo)benzenesulfonic acid sodium salt
PAA peracetic acid
phen 1,10-phenanthroline
PNP p-nitrophenol, 1-hydroxy-4-nitrobenzol
ppm parts per million
S electron spin
s second
t time
T temperature
tacn 1,4,7-triazacyclononane
TAPS [Tris(hydroxymethyl)methyl]aminopropanesulfonic acid
TZ Tartrazine; trisodium (4E)-5-oxo-1-(4-sulfonatophenyl)-4-[(4-sulfonato-
phenyl)hydrazono]-3-pyrazolecarboxylate
terpy 2,2':6',2"-terpyridine
TOF turnover frequency (mol of dye oxidized by mol of catalyst per hour)
TRIS 2-amino-2-hydroxymethylpropanediol
UV/Vis ultraviolet-visible spectrophotometry
TABLE OF CONTENTS
1 Introduction 1
1.1 TRANSITION METAL - MEDIATED OXIDATIVE BLEACH CATALYSIS 3
1.1.1 Iron-based bleaching catalysts 4
1.1.2 Manganese-based bleaching catalysts 5
1.2 OBJECTIVES 10
1.3 REFERENCES AND NOTES 12
2 Metal ion - catalyzed oxidative degradation of Orange II by H2O2. 17
2.1 GERNERAL REMARK 17
2.2 INTRODUCTION 17
2.3 RESULTS AND DISCUSSION 20
2.3.1 General observations 20
2.3.2 Complex-formation between Orange II and MnII 22
2.3.3 CV studies on the complex-formation between Orange II and MnII 26
2.3.4 DFT-calculations 27
2.3.5 Kinetic investigations 30
2.3.5.1 Complex-formation between bicarbonate and MnII 30
2.3.5.2 The effect of the total carbonate concentration 34
2.3.5.3 Reactivity profile as function of pH 39
2.3.5.4 Effect of the [MnII] and [H2O2] on the oxidative reaction course 41
2.3.5.5 Stability of the in situ formed catalyst 44
2.3.6 Mechanistic aspects 45
2.4 CONCLUSIONS 48
2.5 EXPERIMENTAL SECTION 49
2.6 REFERENCES AND NOTES 52
3 Comparative study of a MnII-monomer and the corresponding
oxo-bridged Mn2III/IV-dimer 57
3.1 GENERAL REMARK 57
3.2 INTRODUCTION 57
3.3 RESULTS AND DISCUSSION 59
3.3.1 Kinetic measurements of the catalyzed dye degradation with H2O2 59
3.3.2 The reaction of the catalysts with H2O2 in carbonate buffered solution 65
3.3.2.1 EPR-spectroscopic measurements 65
3.3.2.2 UV/Vis spectroscopic measurements 67
3.3.3 In situ formation of the active species in the catalytic oxidation reaction 69
3.3.4 Precursor complex equilibria in solution 71
3.3.5 Mechanistic aspects 73
3.4 CONCLUSION 75
3.5 EXPERIMENTAL SECTION 75
3.6 REFERENCES AND NOTES 77
3.7 SUPPORTING INFORMATION 80
4 Metal ion - catalyzed oxidative degradation of Orange II
by peracetic acid 85
4.1 GENERAL REMARK 85
4.2 INTRODUCTION 85
4.3 RESULTS AND DISCUSSION 87
4.3.1 Peracetic acid formation and its decomposition at higher pH 87
4.3.2 General observations 89
4.3.3 MnII + PAA – Intermediates formed in the absence of substrate 90
4.3.3.1 UV/Vis spectroscopy 90
4.3.3.2 EPR spectroscopy 97
4.3.4 Comparison of reactivity of different high valent oxo-manganese
species with Orange II 100
4.3.5 Reactivity of different in situ formed intermediates towards Orange II 103
4.3.6 MnII catalyzed degradation of Orange II by PAA 107
4.3.7 Mechanistic interpretation 112
4.3.8 Comparison MnII/PAA vs. MnII/HCO4- system 116
4.4 CONCLUSIONS 118
4.5 EXPERIMENTAL SECTION 119
4.6 REFERENCES AND NOTES 121
4.7 SUPPLEMENTARY INFORMATION 126
5 High catalytic activity of a Mn-terpy compound in oxidative dye
degradations with peracetic acid 141
5.1 GENERAL REMARK 141
5.2 INTRODUCTION 141
5.3 RESULTS AND DISCUSSION 143
5.3.1 MnII + terpy in solution 143
5.3.2 MnIIterpy + PAA in the absence of dye substrate 144
5.3.2.1 UV/Vis measurements 144
5.3.2.2 EPR measurements 147
5.3.3 MnIIterpy catalyzed dye degradation with PAA 149
5.3.3.1 General observations 149
5.3.3.2 Kinetics of the MnIIterpy catalyzed dye degradation by PAA 151
5.3.3.3 Readily prepared dimers in the catalytic dye degradation by PAA 156
5.3.4 Mechanistic implications 159
5.4 CONCLUSIONS 162
5.5 EXPERIMENTAL SECTION 163
5.6 REFERENCES AND NOTES 164
6 Summary 169
7 Zusammenfassung 177
1. Introduction
1
1 INTRODUCTION
The history of bleaching has come a long way – the ancient method of bleaching cotton and
linen practiced in Egypt, Asia and Europe comprised simply the exposure of the fabric to sunlight.
Over the centuries it developed to a complicated array of repetitive, empirically found processes,
like soaking in water and sour milk or boiling in alkaline solution followed by air exposure for
several months. In the 18th century bleaching solutions of potash and lye or of dilute sulphuric acid
were used in Holland and France. Yet, the breakthrough of industrial bleach processes came with
the discovery of the powerful properties of chlorine and hydrogen peroxide – the two main
bleaching systems used nowadays.[1]
In general, the main objective of successful bleaching is the whitening of various substrates in
a homogeneous or heterogeneous chemical reaction. This is either achieved by shifting the
wavelength of light absorbtion outside the range of visible light or by chemical decomposition of
the chromophores’ molecular entity and thereby rendering it water-soluble.[2] For this purpose
different bleaching systems are available, which can be divided into two major classes.[2, 3] On one
hand, the reducing bleaching agents that compris free sulforous acid, bisulfate and hyposulforous
acid and on the other side the class of oxidizing ones, including oxygen, ozone, chlorine,
hypochlorites and peroxides.[3] Chlorine containing compounds are partially used for disinfection
or water treatment and bleaching of delignificated Kraft pulp. This technology is characterized by
high efficiency at relatively low costs. Nevertheless, the major disadvantage of this method is still
the formation of hazardous and non-biodegradable chlorinated side products. The chlorine
bleaching of cellulose pulp, for instance, leaves effluents loaded with chlorophenols,
chloroaliphatics and polychlorinated dioxins and furans.[4]
However, the ever increasing ecological awareness causes a soaring global demand for
environmentally benign procedures. Typical common features of these new “green” procedures
are the reduction of energy demand, the minimization of the total amount of chemicals involved
in the overall process, and in particular the avoidance of toxic side products. Therefore, oxygen-
based bleaching agents are considered to be the chemicals of choice for sustainable bleaching
compositions and are gradually replacing chlorine based technologies.
1. Introduction
2
Although the application of atmospheric oxygen is clearly a long term objective as it would
lead to a further decrease in the chemical loading and a reduction in cost, the scope of molecular
oxygen as an oxidant is at present still narrow.[2, 5] Another environmentally benign and widely
employed oxygen carrier is hydrogen peroxide. It is a cheap, readily available chemical, which
forms only non-toxic water and oxygen as side products during the oxidation process. In order to
facilitate transportation for various applications, hydrogen peroxide is substituted by solid
peroxygens. Laundry cleaning formulations contain sodium percarbonate (PCS) or sodium
perborate (PBS). PCS is only considered as an adduct of carbonate salt and hydrogen peroxide with
the general formula Na2CO3 · 1.5 H2O2, although the in situ formation of the peracid, HCO4- has
been shown by various spectroscopic techniques.[6] PBS on the other hand is a real persalt with a
six-membered heterocyclic di-anion structure.[7a-b] Yet, both peroxygen salts yield a weakly basic
solution of hydrogen peroxide upon dissolving in water.[7c] Other commercially relevant secondary
products of hydrogen peroxide are potassium peroxomonosulfate (Oxone®), mainly used for
chloramine elimination in swimming pools,[8a] and peracetic acid, which has a wide scope of
applications. For instance, it is replacing hypochlorite disinfectants in food industry, brewing,
farming, medicine and waste water treatment, but it is also widely used for domestic laundry
cleaning.[2, 8b]
Oxidative bleaching processes are of major economic importance. For example, the annual
world production of hydrogen peroxide is approximately 2.2 million metric tons, of which 50 % is
used for pulp/paper bleaching and 10 % for textile bleaching.[2] These technologies are
omnipresent in numerous industrially relevant procedures, such as bleaching of stains,[9a] raw
wood pulp,[2, 9b] raw cotton fibers[9a] and waste water treatment.[9b-d] Moreover, they represent an
essential component of various cleaning and refinement processes.Yet, with application in laundry
cleaning[2, 9a] or usage in personal care formulations and detergents,[9a] oxidative bleaching
processes also belong to the everyday life of millions of costumers worldwide.
However, as the kinetics of these oxygen-based bleach systems are sluggish, it is necessary to
activate the oxygen source. Therefore, various technologies to enhance the peroxide efficiency are
known.[2] Besides the application of laccases and peroxidases[10a-b] or activated imines,[10c] the
main research focus lies on peroxide activation by transition metal catalysts. From an
environmental point of view these catalysts are mainly based on iron or manganese complexes,
but also cobalt, vanadium and titanium compounds are known.[9b] Despite the larger number of
1. Introduction
3
commercial applications for transition metal mediated peroxide activation, the underlying reaction
mechanisms still remain unclear. Once the mechanism of action is known, the properties of
potential catalytic systems can be tuned particularly to improve their efficiency by choice of the
ideal reaction conditions. Consequently, a detailed mechanistic understanding represents an
indispensable precondition for the improvement of already implemented catalytic systems as well
as for the development of new efficient and environmentally compatible catalysts.
1.1 TRANSITION METAL - MEDIATED OXIDATIVE BLEACH CATALYSIS
In general, two catalytic pathways are conceivable upon reaction of a peroxide with a
transition metal compound. The transiently formed peroxo- or hydroperoxo-intermediate can
undergo either heterolytic or homolytic O-O-bond scission. A homolytic peroxo bond cleavage
leads to the formation of radicals, which is undesired as these radical species lack selectivity in
their subsequent reactions. Another option is a heterolytic scission of the peroxo bond which
results in the formation of high valent metal-oxo species. Both reaction types are depicted for the
example of H2O2 in Scheme 1.1.
Scheme 1.1 Formation of metal-oxo species in the presence of H2O2 by heterolytic O-O-bond cleavage vs. free radical
generation.
Metal-oxo intermediates often have a greater oxidative power compared to the starting
peroxide and are more selective than free radical species. They represent the one- or two-electron
oxidized form of the starting complex, respectively and following oxygen transfer the starting
1. Introduction
4
catalyst is regenerated. To favor metal based two-electron oxidation over radical type one-
electron processes often requires an electron releasing ligand to stabilize the high-valent metal-
oxo species. Depending on the application, in most cases tailor-made ligand systems are needed to
ensure sufficient selectivity and to expand the catalyst lifetime by avoiding ligand self-oxidation.
Despite of decades of academic and industrial research, there are still ongoing efforts to design
new catalytic transition metal-ligand systems. The main target of these research efforts is to find
model compounds that provide selectivity and efficiency which is comparable to natural enzymes.
Peroxidases, oxygenases or oxidases are the common ideal to activate molecular oxygen or
hydrogen peroxide. During their catalytic cycle all these enzymes form high-valent metal-oxo
species which are considered to be responsible for the subsequent oxidation processes. Many of
these enzymes contain manganese or iron centers, which makes these metals ideal for catalytic
model compounds. Moreover, manganese and iron can play an important role due to their very
rich redox chemistry and their environmental compatibility. Yet, the use of such enzyme mimics is
in many cases connected with a big synthetic effort to obtain adequate ligand systems. In the
following some outstanding and commonly used examples for manganese and iron oxidation and
bleach catalysts are presented.
1.1.1 Iron-based bleaching catalysts
Various classes of iron compounds are known to efficiently catalyze oxygenation and bleach
processes. Besides iron complexes with TPA[11a] (tris(pyridin-2-ylmethyl)amine) or different
pentadentate nitrogen-donor ligands,[11b] one of the most succesful iron systems is the class of
tetraamido macrocyclic ligands developed by Collins et al., the so called TAML systems (Scheme
1.2).[12a]
Scheme 1.2 An iron tetraamidate (TAML) bleaching catalyst.
1. Introduction
5
Since these tetraamidate ligands have a formal charge of -4 they are able to stabilize high-
valent metal species as for example a {MnV=O} complex which was characterized by X-ray and
spectroscopic analyses.[12c-f] In this example the electrophilicity of the oxomanganese center can
be increased by replacing the benzene ring with pyridine, and by subsequent binding of cations
such as lithium to the pyridine.[2, 12f]
A broad range of applications are postulated for this class of complexes, mostly connected to
solution bleaching such as waste water treatment and dye transfer inhibition for laundry
cleaning.[2, 12a,b] Different organic dyes relevant for detergents and waste water treatment can be
bleached in solution with hydrogen peroxide as terminal oxidant.[13a] Moreover, stain removal and
enhanced delignification have also been reported for the FeIII-TAML catalyst.[13b] A variation of
ligand substituents results in a significant alteration of the catalytic properties, such as catalytic
lifetime and hydrolytic stability. Whereas the first generation of TAML catalysts contained ethyl
groups next to the carbonyl moieties, the oxidative stability and thereby the catalytic efficiency
could be increased by replacing them with methyl groups.[13a] Hydrolytic stability was additionally
obtained through substitution of the methyl groups by fluorine, or by introduction of electron
withdrawing groups on the phenyl ring.[12a, 13c, d] No indications for free-radical autoxidation
reactions were found.[2]
Despite recent efforts and ongoing controversy among scientists about many mechanistic
issues and putative reactive species of the FeIII-TAML/oxidant system, the underlying mechanism
of action is still not clarified in detail. Recently a TAML based iron(V)-oxo intermediate which
resembles a proposed reactive intermediate in many catalytic hydrogen peroxide applications has
been isolated and characterized.[13e] However, also monomeric and dimeric (TAML)FeIV=O are
frequently stated as reactive species.[13f, g]
1.1.2 Manganese-based bleaching catalysts
The very first bleach catalyst employed in commercial detergent products was a dinuclear oxo-
bridged manganese compound with the Me3tacn ligand (Scheme 1.3).[2, 14] Wieghardt and co-
workers published this compound in 1988 as a model for manganese-containing enzymes.[2, 14b]
The bleaching activity in combination with hydrogen peroxide is very high for the standard tea-
model stain as well as for wine, fruit, and curry stains.[14a] However, the detergent product
containing this catalyst was withdrawn from the market after it was alleged that the product yields
1. Introduction
6
increased fabric- and dye-damage.[2, 14c] Yet, the catalyst-system is still an integral part of various
machine-dishwashing products where it is responsible for superior removal of tea residues.[2]
Scheme 1.3 Dinuclear Mn(Me3tacn) catalyst.
The mechanistic aspects of the manganese-Me3tacn oxidation reaction were examined for
phenolic compounds and catechol as models for tea stains.[15a-c] The observation of a Mn2III/IV-
dimeric species in EPR spectroscopic measurements was explained by the first step of the reaction,
i.e. a one-electron transfer process from the phenolate ion to the EPR-silent starting complex.[16a-c]
This initial reduction is followed by the appearance of mononuclear MnIV species as confirmed by
further EPR and electrospray mass spectrometry (ES-MS) experiments. Since the dinuclear starting
catalyst is stable against hydrolysis, it was concluded that the initial reduction step is required for
monomer formation. The formation of phenoxyl radicals with Trolox, a water-soluble phenol
derivative, as substrate is observed in both the absence and presence of H2O2.[16b] As the second
step, the oxidation to a MnIV or MnV species is assumed, which can be either mono- or dinuclear.[2]
In general it is suggested that both mononuclear and dinuclear manganese species are
operative.[16a] Contrary, more recent ES-MS experiments indicated the sole formation of
monomeric species in the reaction of [Mn2IV(µ-O)3(Me3tacn)2]2+ with H2O2 and 4-
methoxyphenol.[16c,d]
In addition, Gilbert et al. investigated the catalytic activity towards other substrates like azo-
containing dyes.[16] For this reaction the authors claimed the involvement of monomeric {MnV=O}
species and that such an oxo-manganese species gives rise to two consecutive one-electron
oxidation reactions of the azo dye. Therefore, the catalytic system would resemble the reactivity
of heme-peroxidases (Scheme 1.4). No formation of hydroxyl radicals was observed even with the
use of EPR-sensitive spin traps.[16]
1. Introduction
7
Scheme 1.4 Mechanistic cycle postulated for oxidation of azo dyes by the Mn(Me3tacn) catalyst in the presence of
hydrogen peroxide.[16]
On the other hand, Meunier and co-workers concluded that the oxidation is caused by
hydroxyl radicals. They studied the reactivity of [Mn2IV(µ-O)3(Me3tacn)2]2+/H2O2 with catechol in
the absence and presence of mannitol and found that the formed quinone degrades faster in the
absence than in the presence of mannitol.[2, 17] Although no additional EPR spin-trapping
measurements were conducted, the results were interpreted as the trapping of hydroxyl radicals
by mannitol.[17] Therefore, other explanations may be valid, such as binding of mannitol to the
active species that may prevent the oxidative activity of this system.[2]
Besides efficient catalytic stain and dye bleaching the Mn(Me3tacn) catalyst has in addition
been investigated towards the selective epoxidation with H2O2 as oxidant.[16a] Moreover, the two
mononuclear complexes [MnIV(Me3tacn)(CH3O)3]+ and [MnIV(L)]+ (H3L = 1,4,7-tris(hydroxyethyl)-
1,4,7-triazacyclononane) were studied with regard to their epoxidation activity on the water-
soluble olefins 4-vinylbenzoic acid and styrylbenzoic acid.[2, 16a, 18a] In slightly basic solution, both
monomeric compounds exhibit satisfying total turnover numbers for epoxide formation, but
concomitant H2O2 decomposition occured. By using H218O2 it was evidenced that oxygen transfer
occurs from a coordinated hydrogen peroxide or a derivative species. The epoxide product was
fully 18O-labeled.[18b] Hydroxyl radicals were again excluded, since no hydroxylation of the aromatic
styrene ring was observed.
1. Introduction
8
For the manganese Me3tacn species high epoxidation activities have also been reported in
non-aqueous media, which could achieve total turnover numbers as high as 4000 depending on
the applied conditions.[2, 19a-j] When acetone is used as solvent, the decomposition of hydrogen
peroxide can be suppressed and selectivity towards epoxidation can be increased.[2, 19b, d] High
epoxidation activity was also observed in acetonitrile, provided that additives such as oxalate or
ascorbic acid are present, but no information on the active species is available.[2, 19c, e]
Enantioselective epoxidation of substitued styrenes was achieved by Bolm and co-workers by
using a chiral tacn derivative, 1,4,7-S,S,S-tris(2-hydroxypropyl)-1,4,7-triazacyclononane.[2, 19j] Other
oxidation reactions (Scheme 1.5) include cis-dihydroxylation of cyclooctene,[19h, 20a] oxidation of
benzyl alcohol to benzaldehyde,[20b] oxidation of sulfides to sulfoxides and sulfones,[20c-e] allylic C-H
oxidation of cyclohexene to the alcohol and ketone,[20f] and alkane oxidations by radical
reactions.[2, 20g-l]
Scheme 1.5 Examples of substrate conversions catalyzed by the Mn(Me3tacn) system.
Noteworthy, the dimeric [MnIV(Me3tacn)(CH3O)3]+ complex also epoxidizes cinnamic acid in a
buffer/acetone mixture in the presence of H2O2 as observed for the monomeric manganese
compound with tris(hydroxyethyl)-tacn and also yields fully 18O-labeled epoxide.[18b, 21] Again, ES-
MS signals of mononuclear species were detected, which were assigned to [(Me3tacn)MnIV(OH)3]+
and [(Me3tacn)MnIV(O)(OH)]+ where the latter species would imply a penta-coordinate MnIV
species, which is unusual for a low-spin d3 system.[2]
1. Introduction
9
Despite its excellent epoxidation abilities, it is far more surprising, that the monomeric tacn
derivative [MnIV(L)]+ does not catalyze stain bleaching whereas the dimeric [Mn2IV(µ-
O)3(Me3tacn)2]2+ does.[16a] As a consequence it might be concluded that the chemical processes
which underly the oxygen transfer to an alkene on the one side, and stain bleaching of
polyphenolic chromophores on the other side, are basically different. However, differences in the
applied reaction conditions such as pH (pH 8 vs. 10.5), catalyst stability and feasible interactions
with the stain or garments may account for this observation. Be that as it may, these studies led to
the suggestion that the same {MnV=O} species is involved in the catalysis as described for the
oxidation of phenol and styrene.[2] Electron-rich derivatives of cinnamic acid are converted most
efficiently, which is indicative of an electrophilic {Mn=O} center.[2] Although a coherent
mechanistic understanding is still lacking, it is clear that both typical oxygen-transfer and
hydrogen-abstraction reactions are catalyzed. Based on the reactivity of [Mn2IV(µ-
O)3(Me3tacn)2]2+/H2O2 in aqueous solutions, it is assumed that the observed bleaching activity
primarily originates from epoxidation/oxo transfer, but a proof for this hypothesis has still to be
provided.[2]
Another important class of stain bleaching manganese complexes was patented by Ciba,
Clariant and Henkel.[22] Some of these Schiff-base compounds are specially suitable for low-
temperature bleaching and dye-transfer inhibition which sometimes occurs when colored and
white garments are combined and its prevention is achieved by dye oxidation in solution.[2, 22 a-c]
The manganese metal center in these complexes has the oxidation state III. Some examples of the
salen ligand structures, which together with manganese have been patented as bleach catalysts,
are shown in Scheme 1.6. They are less efficient catalysts than the Mn(Me3tacn) complexes, but
according to the patents, they also cause less dye-fading on the cloths and are easier to
synthesize.[2, 22a-c] These catalysts generally activate hydrogen peroxide to yield an improved tea
stain bleaching, and in some cases they outperform the conventional TAED/percarbonate
bleaching system.[2, 22c-d]
Although non of the patents specifies mechanistic information, the reactivity of manganese
Schiff-base compounds has been studied in non-aqueous solution.[2, 23] For the reaction with
hydrogen peroxide, hypochlorite and iodosobenzene, highly enantioselective oxidations have
been reported in the presence of nitrogen donors such as imidazole.[2] The reactivity is ascribed to
an electrophilic {MnV=O} intermediate, which yields a substrate radical intermediate.[2]
1. Introduction
10
Scheme 1.6 Examples of Schiff-base ligands patented as bleach catalysts in combination with manganese.
1.2 OBJECTIVES
Throughout the last decades a broad variety of manganese based transition metal compounds
has been investigated towards their ability to activate peroxide for numerous homogeneous and
heterogeneous applications as summarized in Chapter 1.1. Yet, there are still several aspects that
require further clarification. In this context, the general goal of this thesis was to contribute to a
more detailed mechanistic understanding of these catalytic processes especially in aqueous
solution. As water is nature’s principle solvent, it is the ideal choice for green chemistry.[12a, 13f, 24]
An unique aspect of aqueous transition-metal chemistry that cannot simply be transferred to
other solvents, is the ability to manipulate the reactivity of various species by a simple change in
pH.[25] Having a degree of control over the lifetime of an intermediate facilitates spectroscopic and
kinetic studies and expands the range of methods and techniques that can be used in mechanistic
investigations.[25]
Despite the complexity of the structure and chemistry in catalytic peroxide activation, the
reactivity of complex catalysts can be sometimes reproduced by much simpler models.[25] In this
context, one objective of this work was to find, investigate and improve catalytic systems for the
homogeneous activation of hydrogen peroxide based on simple manganese salts. Chapter 2
1. Introduction
11
describes the study of the MnII ion catalyzed degradation of the commonly used model substrate
Orange II with H2O2, which revealed the crucial role of the carbonate buffer in the formation of
reactive Mn-bicarbonate complexes and the in situ formation of peroxycarbonate, the actual
active oxidant.
Further, it should be clarified to which extent dimeric, oxo-bridged high-valent species
contribute in the studied bleach reaction with H2O2 as terminal oxidant by applying a well known
ligand system. The enhanced complexity caused by the presence of the ligand and by the presence
of more than one metal center in the molecule, naturally leads to a larger number of possible
intermediates which can now include mixed-metal oxidation states,[26 a] bridging and end-on
superoxide[26b, c]/peroxide[27] groups, or isomeric dioxo- vs. μ-peroxo-forms.[25, 28] In this respect,
Chapter 3 reports kinetic investigations of monomeric and dimerc precursors in the H2O2 induced
degradation of different dye substrates. Again a strong influence of the buffer system on the
activation of H2O2 was detected and discussed in the light of the variation of the system compared
to the simple MnII salts in Chapter 2. In addition, EPR-spectroscopic methods proved to be a
helpful tool to elucidate the nature of the predominant species under catalytic conditions.
A further objective was to extend our work to another commonly used oxidant, peracetic acid
(PAA). The application of peracetic acid marks a special challenge for mechanistic studies, since it
always contains an equilibrium amount of H2O2. In order to address this additional complication, a
careful selection of the reaction conditions is required to lay the foundation for one of the few
mechanistic studies on transition metal catalyzed bleaching with PAA as oxidant. In Chapter 4 the
reaction of MnII salts with PAA was studied by detailed UV/Vis- and concomitant EPR-
spectroscopic experiments in combination with selected reactivity studies to disclose the different
in situ formed intermediates and their relevance to the dye degradation process. Furthermore, our
studies revealed the ambiguous role of the equilibrium concentration of H2O2 on the lifetime of
the catalytic system.
To complete the mechanistic aim of this work, Chapter 5 gives some preliminary insight into
the manganese catalyzed activation of PAA in the presence of a chelate. Contrary to the results
described in Chapter 3 for H2O2, our findings point to the involvement of dimeric, oxo-bridged
species in the dye degradation process.
1. Introduction
12
In summary, the general objective was to provide detailed information on putative reactive
intermediates and the mechanism of action for the activation of two widely used oxidants, namely
hydrogen peroxide and peracetic acid, by manganese systems with and without a defined ligand
system, chosen as simplest representatives for the numerous different peroxide activating
catalysts used in bleaching and oxygenation applications.
1.3 REFERENCES AND NOTES
[1] Suess, H. U. Pulp BleachingToday, Walter de Gruyter 2010.
[2] Hage, R.; Lienke, A. Angew. Chem. Int. Ed. 2006, 45, 206-222.
[3] von Georgievics, G. The Chemical Technology of Textile Fibres - Their Origin, Structure,
Preperation, Washing, Bleaching, Dyeing, Printing and Dressing, Scott. Greenwood, London,
1920.
[4] (a) Daube, A. K.; Karlm, M. R.; Dimmel, D. R.; McDonough, T. J.; Banerjee, S. Environ. Sci.
Technol. 1992, 26, 1324-1326; (b) Nakamata, K.; Ohi, H. J. Wood Sci. 2003, 49, 525-530.
[5] Beller, M. Adv. Synth. Catal. 2004, 346, 107-108.
[6] (a) Flanagan, J.; Jones, D. P.; Griffith, W. P.; Skapski, A. C.; West, A. P. J. Chem. Soc., Chem.
Commun. 1986, 20-21; (b) Lane, B. S.; Vogt, M.; DeRose, V. J.; Burgess, K. J. Am. Chem. Soc.
2002, 124, 11946-11954; (c) Richardson, D. E.; Regino, C. A. S.; Yao, H. R.; Johnson, J. V. Free
Radical Biol. Med. 2003, 35, 1538-1550; (d) Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett,
D. A. J. Am. Chem. Soc. 2000, 122, 1729-1739; (e) Bakhmutova-Albert, E. V.; Yao, H.; Denevan,
D. E.; Richardson, D. E. Inorg. Chem. 2010, 49, 11287-11296.
[7] (a) Carrondo, M. A. A. F. d. C. T.; Griffith, W. P.; Jones, P. D.; Skapski, A. C. J. Chem. Soc.,
Dalton Trans. 1977, 2323-2327; (b) Hansson, A. Acta Chem. Scand. 1961, 15, 934-935; (c)
Galwey, A. K.; Hood, W. J. J. Chem. Soc., Faraday Trans. 1 1982, 78, 2815-2827.
[8] (a) Jakob, H.; Lehmann, T.; Jacobi, S.; Gutewort, S. Ullmann´s Encyclopedia of Industrial
Chemistry, Wiley, 2007; (b) Block, S. Peroxygen Compounds, Philadelphia, 1991.
[9] (a) Kirk-Othmer, Encyclopedia of chemical technology, Wiley: New York, 1991; (b) Jones, C. W.
Applications of Hydrogen Peroxide and Derivatives, RSoC: Cambridge, 1999; (c) Aurand, K. Die
Trinkwasserverordnung, Berlin, 1991; (d) Schwarzer, H. Umwelt 1981, 482 - 486.
1. Introduction
13
[10] (a) Solomon, E. I.; Sundaram, U. M.; Machonkin, T. E. Chem. Rev. 1996, 96, 2563-2605; (b)
Hansch, C.; Gao, H. Chem. Rev. 1997, 97, 2995-3059; (c) Pocalyko, D. J.; Coope, J. L.; Carchi, A.
J.; Boen, L.; Madison, S. A. J. Chem. Soc., Perkin Trans. 2 1997, 117-121.
[11] (a) Que Jr., L.; Kim, C.; Kim, J.; Zang, Y. WO-A-9748710, 1996. [Chem. Abstr. 1997, 128,
90357]; (b) Feringa, B. L.; Lubben, M.; Hermant, R. M.; Twisker, R. S.; Que Jr., L.
(Unilever),WO-A-9534268, 1994 [Chem. Abstr. 1995, 124, 205657].
[12] (a) Collins, T. J. Acc. Chem. Res. 2002, 35, 782-790; (c) Collins, T. J.; Gordon-Wylie, S.W. J. Am.
Chem. Soc. 1989, 111, 4511-4513; (b) Beach, E. S.; Malecky, R. T.; Gil, R. R.; Horwitz, C. P.;
Collins; T. J. Catal. Sci. Technol. 2011, 1, 437-443; (d) Collins, T. J.; Powell, R. D.; Slebodnick, C.;
Uffelman, E. S. J. Am. Chem. Soc. 1990, 112, 899-901; (e) Workman, J. M.; Powell, R. D.;
Procyk, A. D.; Collins, T. J.; Bocian, D. F. Inorg. Chem. 1992, 31, 1548-1550; (f) Miller, C. M.;
Gordon-Wylie, S. W.; Horwitz, C. P.; Strazisar, S. A.; Peraino, D. K.; Clark, G. R.; Weintraub, S.
T.; Collins, T. J. J. Am. Chem. Soc. 1998, 120, 11540-11541.
[13] (a) Horwitz, C. P.; Fooksman, D. R.; Vuocolo, L. D.; Gordon-Wylie, S. W.; Cox, N. J.; Collins, T. J.
J. Am. Chem. Soc. 1998, 120, 4867-4868; (b) Collins, T. J. ACS Symp. Ser. 2002, 47; (c) Ghosh,
A.; Ryabov, A. D.; Mayer, S. M.; Horner, D. C.; Prasuhn, D. E. Jr., Gupta, S. S.; Vuocolo, L.;
Culver, C.; Hendrich, M. P.; Rickard, C. E. F.; Norman, R. E.; Horwitz, C. P.; Collins, T. J. J. Am.
Chem. Soc. 2003, 125, 12378-12379; (d) Polshin, V.; Popescu, D.-L.; Fischer, A.; Chanda, A.;
Horner, D. C.; Beach, E. S.; Henry, J.; Qian, Y.-L.; Horwitz, C. P.; Lente, G.; Fabian, I.; Münck, E.;
Bominaar, E. L.; Ryabov, A. D.; Collins, T. J. J. Am. Chem. Soc. 2008, 130, 4497-4506; (e) Tiago
de Oliveira, F.; Chanda, A.; Banerjee, D.; Shan, X.; Mondal, S.; Que Jr., L.; Bominaar, E. L.;
Münck, E.; Collins, T. J. Science 2007, 315, 835-838; (f) Popescu, D.-L.; Vrabel, M.; Brausam,
A.; Madsen, P.; Lente, G.; Fabian, I.; Ryabov, A. D.; van Eldik, R.; Collins, T. J. Inorg. Chem.
2010, 49, 11439-11448; (g) Chanda, A.; Shan, X.; Chakrabarti, M.; Ellis, W. C.; Popescu, D. L.;
Tiago de Oliveira, F.; Wang, D.; Que Jr., L.; Collins, T. J.; Münck, E.; Bominaar, E. L. Inorg.
Chem. 2008, 47, 3669-3678.
[14] (a) Favre, F.; Hage, R.; van der Helm-Rademaker, K.; Koek, J. H.; Martens, R. J.; Swarthoff, T.;
van Vliet, M. R. P. (Unilever), EPB-0458397, 1990 [Chem. Abstr. 1991, 116, 154241]; (b)
Wieghardt, K.; Bossek, U.; Nuber, B.; Weiss, J.; Bonvoisin, J.; Corbella, M.; Vitols, S. E.; Girerd,
J.-J. J. Am. Chem. Soc. 1988, 110, 7398-7411; (c) Verrall, M. Nature 1995, 373, 181.
1. Introduction
14
[15] (a) Hage, R.; J. E. Iburg, J. E.; Kerschner, J.; Koek, J. H.; Lempers, E. L. M.; Martens, R. J.;
Racherla, U. S.; Russell, S. W.; Swarthoff, T.; van Vliet, M. R. P.; Warnaar, J. B.; van der Wolf,
L.; Krijnen, L. B. Nature 1994, 369, 637-639; (b) Gilbert, B. C.; Kamp, N. W. J.; Lindsay-Smith, J.
R.; Oakes, J. J. Chem. Soc. Perkin Trans. 2 1997, 2161-2166; (c) Gilbert, B. C.; Kamp, N. W. J.;
Lindsay-Smith, J. R.; Oakes, J. J. Chem. Soc. Perkin Trans. 2 1998, 1841-1844; (d) Gilbert, B. C.;
Kamp, N. W. J.; Lindsay-Smith, J. R.; Oakes, J. Org. Biomol. Chem. 2004, 2, 1176-1180.
[16] Gilbert, B. C.; Lindsay-Smith, J. R.; Newton, M. S.; Oakes, J.; Pons i Prats, R. Org. Biomol. Chem.
2003, 1, 1568-1577.
[17] Sorokin, A.; Fraisse, L.; Rabion, A.; Meunier, B. J. Mol. Catal. A 1997, 117, 103-114.
[18] (a) Quee-Smith, V. C.; DelPizzo, L.; Jureller, S.H.; Kerschner, J.L.; Hage, R. Inorg.Chem. 1996,
35, 6461-6465; (b) Hage, R.; Kerschner, Trends in Inorg. Chem. 1998, 5, 145-159.
[19] (a) de Vos, D. E.; Bein, T. Chem. Commun. 1996, 917-918; (b) de Vos, D. E.; Bein, T.
J.Organomet.Chem. 1996, 520, 195-200; (c) de Vos, D. E.; Sels, B. F.; Reynaers, M.; Subba Rao,
Y. V.; Jacobs, P. A. Tetrahedron Lett. 1998, 39, 3221-3224; (d) Edwards, J-O.; Sauer, M. V. C. J.
Phys.Chem. 1971, 75, 3004-3011; (e) Berkessel, A.; Sklorz, C. A. Tetrahedron Lett. 1999, 40,
7965-7968; (f) de Vos, D. E.; Meinershagen, J.; Bein, T. Angew. Chem. Int. Ed. 1996, 35, 2211-
2213; (g) de Vos, D. E.; Bein, T. J. Am. Chem. Soc. 1997, 119, 9460-9465; (h) de Vos, D.; de
Wildeman, S.; Sels, B. F.; Grobet, P. J.; Jabobs, P. A. Angew. Chem. Int. Ed. 1999, 38, 980-983;
(i) Rao, S. Y. V.; de Vos, D. E.; Bein, T.; Jacobs, P. A. Chem.Commun. 1997, 355-356; (j) Bolm,
C.; Kadereit, D.; Valacchi, M. Synlett 1997, 6, 687-688.
[20] (a) Brinksma, J.; Schmieder, L.; van Vliet, L.; Boaron, R.; Hage, R.; de Vos, D. E.; Alsters, P. L.;
Feringa, B. L. Tetrahedron Lett. 2002, 43, 2619-2622; (b) Handervan, C.; Hage, R.; Feringa, B. L.
Chem. Commun. 1997, 419-420; (c) Barton, D. H. R.; Choi, S.-Y.; Hu, B.; Smith, J. A.
Tetrahedron 1998, 54, 3367-3378; (d) Brinksma, J.; LaCrois, R.; Feringa, B. L.; Donnoli, M. I.;
Rosini, C. Tetrahedron Lett. 2001, 42, 4049-4052; (e) Barker, J. E.; Ren, T. Tetrahedron Lett.
2004, 45, 4681-4683; (f) Vincent, M.; Rabion, A.; Yachandra, V. K.; Fish, R. H. Angew. Chem.
Int. Ed. 1997, 36, 2346-2349 (g) Metal-catalyzed Oxidations of Organic Compounds (Eds.:
Sheldon, R. A.; Kochi, J. K.), Academic Press, NewYork, 1981; (h) Bennur, T. H.; Sabne, S.;
Deshpande, S. S.; Srinivas, D.; Sivasanker, S. J. Mol. Catal. A 2002, 185, 71; (i) Klopstra, M.;
Hage, R.; Kellogg, R. M.;Feringa, B. L. Tetrahedron Lett. 2003, 44, 4581-4584; (j) Shul’pin, G.
B.; Nizova, G. V.; Kozlov, Y. N.; Pechenkina, I. G. New J. Chem. 2002, 26, 1238-1245; (k)
1. Introduction
15
Lindsay-Smith, J. R.; Shul’pin, G. Tetrahedron Lett. 1998, 39, 4909-4912; (l) Woitiski, C. B.;
Kozlov, Y. N.; Mandelli, D.; Nizova, G. V.; Schuchardt, U.; Shul’pin, G. B. J. Mol. Catal. A 2004,
222, 103-119.
[21] Gilbert, B. C.; Lindsay-Smith, J. R.; Mairata i Payeras, A.; Oakes, J.; Pons i Prats, R. J. Mol. Catal.
A 2004, 219, 265-272.
[22] (a) Reinehr, D.; Metzger, G. (Ciba), WO-A-9719162, 1995 [Chem. Abstr. 1997, 127, 67687]; (b)
Bachmann, F.; Dannacher, J.; Hazenkamp, M.; Schlingloff, G.; Richter, G.; Dbaly, H.; Traber, R.
H. (Ciba), WO-A-200105925, 1999 [Chem. Abstr. 2001, 134, 133329]; (c) Hazenkamp, M.;
Bachmann, F.; Makowka, C.; Dubs, M.-J.; Richter, G.; Schlingloff, G.; Dannacher, J.;
Weingartner, P. (Ciba), WO-A-200053712, 1999 [Chem. Abstr. 2000, 133, 239758]; (d) Nestler,
B. (Clariant), EP-B-869171, 1997 [Chem. Abstr. 1998, 129, 291445]; (e) Blum, H.; Nitsch, C.;
Jeschke, P.; Haerer, J.; Pegelow, U. (Henkel), WO-A-9707191, 1997 [Chem. Abstr. 1997, 126,
239902].
[23] (a) Berkessel, A.; Frauenkron, M.; Schwenkreis, T.; Steinmetz, A.; Baum, G.; Fenske, D. J. Mol.
Catal. A 2002, 113, 321-342; (b) Katsuki, T. Coord. Chem. Rev. 1995, 140, 189-214; (c) Lane, B.
S.; Burgess, K. Chem. Rev. 2003, 103, 2457-2473; (d) Linde, C.; Koliai, N.; Norrby, P.-O.;
Akermark, B. Chem. Eur. J. 2002, 8, 2568-2573; (e) Pietikainen, P. J. J. Mol. Catal. A 2001, 165,
73-79.
[24] Anastas, P. T. Crit. Rev. Anal. Chem. 1999, 29, 167-175.
[25] Bakac, A. Inorg. Chem. 2010, 49, 3584-3593.
[26] (a) Kovaleva, E. G.; Neibergall, M. B.; Chakrabarty, S.; Lipscomb, J. D. Acc. Chem. Res. 2007,
40, 475-483; (b) Egan Jr., J. W.; Haggerty, B. S.; Rheingold, A. L.; Sendlinger, S. C.; Theopold, K.
H. J. Am. Chem. Soc. 1990, 112, 2445-2446; (c) Qin, K.; Incarvito, C. D.; Rheingold, A. L.;
Theopold, K. H. Angew. Chem., Int. Ed. 2002, 41, 2333-2335.
[27] (a) Takahashi, Y.; Hashimoto, M.; Hikichi, S.; Akita, M.; Moro-Oka, Y. Angew. Chem., Int. Ed.
1999, 38, 3074-3077; (b) Ho, R. Y. N.; Roelfes, G.; Hermant, R.; Hage, R.; Feringa, B. L.; Que Jr.,
L. Chem. Commun. 1999, 2161-2162; (d) Rahman, A. F. M. M.; Jackson, W. G.; Willis, A. C.
Inorg. Chem. 2004, 43, 7558-7560; (d) Brausam, A.; Maigut, J.; Meier, R.; Szilagyi, P. A.;
Buschmann, H.-J.; Massa, W.; Homonnay, Z.; van Eldik, R. Inorg. Chem. 2009, 48, 7864-7884.
1. Introduction
16
[28] (a) Halfen, J. A.; Mahapatra, S.; Wilkinson, E. C.; Kaderli, S.; Young Jr., V. G.; Que Jr., L.;
Zuberbuhler, A. D.; Tolman, W. B. Science 1996, 271, 1397-1400; (b) Cramer, C. J.; Tolman, W.
B. Acc. Chem. Res. 2007, 40, 601-608.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
17
2 METAL ION - CATALYZED OXIDATIVE DEGRADATION OF
ORANGE II BY H2O2.
2.1 GERNERAL REMARK
The following chapter is based on the original publication: Metal ion – catalyzed oxidative
degradation of Orange II by H2O2. High catalytic activity of simple manganese salts, Erika Ember,
Sabine Rothbart, Ralph Puchta and Rudi van Eldik, New J. Chem., 2009, 33, 34-49.
2.2 INTRODUCTION
Nowadays, one of the major environmental problems concerns the strong increase in
xenobiotic and organic substances that are persistent in the natural ecosystem. Most of these
compounds have an aromatic structure, which makes them highly stable and thus difficult to
degrade.[1a] A significant source of environmental pollution is industrial dye waste due to their
visibility and recalcitrance, since dyes are highly colored and designed to resist chemical,
biochemical and photochemical degradation.[1b] About half of the global production of synthetic
dyes (700,000 t per year) are classified as aromatic azo compounds that have a -N=N- unit as
chromophore in their molecular structure. Over 15 % of textile dyes are lost in waste water
streams during the dyeing operation.[1c] Azo dyes are known to be largely non-biodegradable
under aerobic conditions and to be reduced to more hazardous intermediates under anaerobic
conditions.[1d] The decolourisation of wastewater has acquired increasingly importance in recent
years, however, there is no simple solution to this problem because the conventional
physicochemical methods are costly and lead to the accumulation of sludges.[1e]
One approach to solve these problems would be to develop low-cost, highly efficient, and
environmentally compatible oxidation catalysts on the basis of transition metal complexes.[2a-d]
Recently, photodegradation methods based on TiO2 as a photocatalyst,[2c] beside Fenton
systems,[2d] emerged as one of the most promising technologies and received increasing attention
due to their practical and potential value in environmental protection. However, they are only
successful under specific pH and temperature conditions.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
18
Several studies were performed during the last few years in order to find good catalysts for the
oxidative degradation of different organic dyes. From an environmental point of view, first row
transition metals are the most challenging. Highly effective Fe,[3, 4] Co,[5] Cr[6] and Mn[2b, 7] based
oxidation catalysts were developed. In combination with different oxidizing agents, the
decomposition of stable organic substances was possible. A novel highly active and
environmentally benign catalytic system based on Fe-TAML (TAML = tetraamido macrocyclic
ligand) was newly reported by Collins et al.[4] However, besides the tremendous synthetic efforts
to obtain an effective catalytic system, the presence of high concentrations of oxidizing agents is
required in many cases.
Among the possible oxidizing agents, H2O2 is one of the most commonly used owing to its eco-
friendly nature. The use of H2O2 as a green oxidizing agent in these reactions is justified by a low
organic content of the wastewater to be treated and a low reaction temperature, thus requiring
the presence of an adequate catalyst due to the high kinetic activation barrier of such reactions.
Commonly used methods for activation of H2O2 include the formation of reactive peroxyacids from
carboxylic acids and peroxycarboximidic acid from acetonitrile (Payne oxidation),[8a] the generation
of peroxyisourea from carbodiimide in the presence of either a weak acid or a mild base,[8b] or the
use of percarbonate, persulfate or perborate in strongly basic solution.[8c] In order to achieve fast
oxidative transformations, the use of large amounts of co-catalyst additives is often required.[9]
Among these, the use of percarbonate, a versatile oxidizing agent, is preferred for environmental
reasons.[10a, b] Oxidation using environmentally benign oxidants has aroused much interest,[10c-e]
because chemical industry continues to require cleaner oxidation, which is an advance over
environmentally unfavoured oxidations and a step up from more costly organic peroxides.[11]
In this report, we propose a fast and clean catalytic oxidative degradation of Orange II as
model substrate by H2O2 in aqueous carbonate solution under mild reaction conditions, pH 8 - 10
and 25 °C, Scheme 2.1.
Scheme 2.1 Catalytic oxidative degradation of Orange II as model substrate by H2O2 in aqueous carbonate solution
under mild reaction conditions, pH 8 - 10 and 25 °C.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
19
Starting from commercially available MnII(NO3)2 in aqueous carbonate solution for catalytic
applications, various aspects of the in situ generation of very reactive high valent manganese
intermediates in the presence of H2O2 were studied.
Baes and Mesmer have shown that manganese salts in aqueous solution are able to form very
reactive aquated intermediates.[12] Moreover, in an alkaline medium, the introduction of a
hydroxy ligand trans to a water ligand is expected to produce more labile OH-Mn-H2O species, and
their formation (Scheme 2.2) is considered to be of major importance for their catalytic activity.
Scheme 2.2 Formation of manganese(II)hexaquacomplex and deprotonation in alkaline medium. L = Cl-, NO3
-, SO4
2-,
CH3COO-.
In the present study, the formation of catalytically inactive MnII(OH)2 species was observed at
higher pH, leading to deactivation of the produced Mn intermediates. The activation of H2O2 in the
presence of manganese salts as a function of pH and carbonate concentration was therefore
monitored using UV/Vis spectrophotometry. In situ formed, high-valent manganese intermediates
are known to be highly unstable in the absence of a spectator ligand. As the study progressed, it
was of importance to investigate the role of the azo dye as a potential coordinating ligand to
stabilize the produced intermediate under different reaction conditions. Electrochemical
measurements and DFT calculations were used to develop a better understanding of the
coordination chemistry of Orange II. The successful implementation of such catalytic systems
becomes a worthwhile objective when issues such as environmental compatibility, high atom
economy, availability, and expenses are considered.[13]
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
20
2.3 RESULTS AND DISCUSSION
2.3.1 General observations
A series of experiments was performed in order to investigate the in situ generation of the
highly reactive manganese catalyst in the oxidative degradation of Orange II by H2O2 under mild
reaction conditions starting with a simple MnII salt. Oxidation reactions are in generally affected by
the protonation state of the substrate, catalyst and oxidant, and the solvent used. It is further
important to note that the studied organic dye (Orange II) can exist in either one of two
tautomeric forms, or in an equilibrium mixture, depending on the process parameters. This kind of
rapid dynamic equilibrium is relevant as one dye species may be more reactive than the other. Azo
dyes containing a hydroxyl group in the ortho position to the azo group within naphthyl or higher
fused ring systems can exist as azo and hydrazone tautomers,[14a] with the relative amounts
varying with reaction parameters such as solvent and temperature.[14b] Furthermore, in aqueous
solution these species are in a pH dependent equilibrium with a common anion, in which the
negative charge is delocalised throughout the molecule (see Scheme 2.3).[14c]
Scheme 2.3 Orange II, pKa = 11.4, λmax = 484 nm.
These are chemically distinct forms which have characteristically different visible spectra, the
azo form absorbs typically at 400 - 440 nm and the hydrazone form at 475 - 510 nm (see Figure
2.1).[14d] The absorption spectrum of Orange II in an aqueous carbonate solution shows under the
selected reaction conditions (Figure 2.1) one main band at 480 nm, which corresponds to the n →
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
21
π* transition of the azo form. The other two bands at 300 and 270 nm are attributed to the π →
π* transition of the benzene and naphthalene rings, respectively.[14e]
Figure 2.1 UV/Vis spectrum of 10-4
M Orange II in carbonate buffer solution at pH 8.5.
Figure 2.2 (A) UV/Vis spectrum of a 5 · 10-5
M Orange II carbonate (0.4 M HCO3-) solution at pH 8.5 with a 5 · 10
-5 M
MnII(NO3)2 and 0.01 M H2O2. (B) Comparison of the absorbance changes at 484 nm during the reaction of (A) (―) and
the uncatalyzed reaction of 5 · 10-5
M Orange II with 0.01 M H2O under identical reaction conditions.
Orange II, due to the presence of aromatic groups, is very stable, and in the presence of a
powerful bleaching agent such as H2O2, degradation of dye solutions occurs slowly under specific
reaction conditions. Surprisingly, the oxidation rate was tremendously accelerated by addition of a
simple manganese salt (Figures 2.2 (A) and (B)). Compared to the uncatalyzed reaction of H2O2
with Orange II, a rate enhancement of several orders of magnitude is found, as becomes obvious
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
22
from the spectral changes at 484 nm in Figures 2.2 (B). The reactivity of the in situ formed
intermediate was comparable with the catalytic activity of some earlier postulated, well known
manganese bleach catalysts[7, 15a] and manganese porphyrines.[15b] In our work, the formation and
stabilization of the active catalyst was studied in a carbonate buffer solution.
2.3.2 Complex-formation between Orange II and MnII
Ortho-hydroxy aromatic azo dyes, which are bidentate complexing agents, are of considerable
practical and theoretical interest because of their ability to form stable chelate complexes with
several metal ions.[16] It is known that Orange II can act as a chelating agent since the hydroxy and
sulfonate groups allow the formation of a stabilized complex.[16b] Addition of Orange II to a freshly
prepared aqueous carbonate solution of a MnII salt results in significant changes in the UV/Vis
spectrum of Orange II as shown in Figure 2.3.
Figure 2.3 (—) UV/Vis spectrum of a 5 · 10-5
M Orange II carbonate (0.1 M HCO3-) solution at pH 8.5 before mixing with
a 5 · 10-5
M MnII(NO3)2 solution at pH 8.5. (—) UV/Vis spectrum recorded directly after mixing (ca. 5 s delay).
UV/Vis spectra recorded before and after mixing (ca. 5 s delay) of 5 · 10-5 M Orange II with 5 ·
10-5 M MnII(NO3)2 showed a significant increase in absorbance at 480, 310 and 228 nm,
respectively. The differences before and after mixing are not profound at low MnII concentrations.
On increasing the MnII concentration, a continuous increase in ΔAbsλ = 480 nm = A(dye + Mn(II)) - Adye was
observed, indicating the formation of an Orange II∙∙∙∙MnII species according to Equation 2.1. It
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
23
should be noted that at higher MnII concentration, a precipitate started to form. The value of Keq
was determined through a constant variation of the MnII concentration.
Figure 2.4 (A) Change in absorbance at 480 nm on addition of different concentrations of MnII to 5 · 10
-5 M Orange II in
aqueous carbonate solution (0.2 M HCO3-) at pH 8.5 and 22 °C. (B) Job plot analysis for complex-formation between
Orange II and MnII in aqueous carbonate solution (0.2 M HCO3
-) at pH 8.5. (C) Spectral changes at 480 nm on addition
of different concentrations of Orange II to a freshly prepared 5 · 10-5
M MnII(NO3)2 carbonate solution (0.2 M HCO3
-) at
pH 8.5 and 22 °C. (D) Job plot analysis for the complex formation in aqueous carbonate solution (0.2 M HCO3-).
For a correct determination of the complex-formation constant, independent measurements
were performed at constant manganese concentration where the Orange II concentration was
continuously varied (see Figure 2.4 (C)). Independent measurements were repeated between five
(A) (B)
(C) (D)
(2.1)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
24
and eight times. Selected data are shown in Figure 2.4 (A), where the solid line represents a fit of
the data to Equation 2.2.
The values of A0 and A∞ represent the absorbances of Orange II and of the complex Orange
II∙∙∙∙MnII, respectively, and Ax is the absorbance at any MnII concentration.
The value of Keq was calculated from Equation 2.2 to be (2.9 ± 0.9) · 104 M-1, indicating a
relatively weak coordination of the dye to the metal center. Experimentally, through addition of a
4 · 10-5 M MnII(NO3)2 solution to a 5 · 10-5 M Orange II aqueous carbonate solution (0.2 M HCO3-), a
decrease in the pH of the solution from 8.5 to 8.3 was observed, which suggests phenolic proton
release due to MnII coordination to Orange II with the formation of a six-membered ring structure
instead of coordination to the terminal sulfonato group.
At higher concentrations (above ca. 1 · 10-3 M) Orange II forms dimers and higher aggregates in
aqueous solutions,[14c, 17a, b] and has a marked effect on the observed spectra, particularly UV/Vis
and NMR.[17c] A Benesi-Hildebrand treatment of the optical data to determine Keq could not be
applied since the concentration of Orange II and MnII were close to each other.[18]
Scheme 2.4 (A) Proposed structure for a 1:1 Orange II∙∙∙∙MnII complex formed in a carbonate buffer solution at a low
concentration of MnII. (B) Proposed structure for a 1:2 Orange II∙∙∙∙Mn∙∙∙∙Orange II complex formed in a carbonate
buffer solution at a high concentration of Orange II.
(A) (B)
(2.2) ][
][
OrangeIIK1
OrangeIIKAbsAbs
eq
eq
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
25
Using Job’s method,[19] the stoichiometry of the formed complex could be determined.
According to the data shown in Figures 2.4 (B) and (D), at lower MnII concentration the formation
of a complex with a stoichiometry of 1:1 can be assumed. Upon a further increase in the Orange II
concentration, complexes with a higher stoichiometry are possibly formed (see Scheme 2.4).
Similar structures have been reported earlier by Nadtochenko and Kiwi when a Fe3+ salt was
added to an Orange II solution in acidic medium.[20a] Bauer also reported a TiIV complex, where TiIV
is coordinated by two oxygen atoms from the sulfonato group and the oxygen of the carbonyl
group of the hydrazone tautomer.[20b] In the enzyme manganese peroxidase, the double role of
Orange II as a stabilizer, forming a complex with MnIII, and as a substrate that permits the
regeneration of MnII, was recently postulated by López et al.[16b] Although, the coordination of
organic dyes, viz. Alizarin, Alizarin S[21] and Orange II,[1c] to several transition metal centres has
been known for years, comparatively little is found on their use as potential stabilizing ligands in
oxidative degradation of organic dyes.
The formed Orange II∙∙∙∙MnII complex was isolated and the validity of its composition was
confirmed by elemental analysis. In control experiments the reactivity of the isolated 1:1 Orange
II∙∙∙∙MnII and 2:1 complexes were studied. The isolated complexes exhibit the same catalytic
activity and stability under the experimental conditions employed for the in situ generation of the
complex. Due to the weak coordination mode of the ligand, no differences between the catalytic
activity of the 1:1 and 2:1 complex were found.
In addition, to gain more information on the activation mode of the catalyst, two further
experimental procedures based on different activation and stabilization modes of the activated
catalyst, were followed. In one, the catalytic active species was generated in situ in the carbonate
buffer solution by addition of the desired amount of H2O2, followed by the addition of the
corresponding quantity of Orange II to the reaction mixture. In the other, Orange II was added to
the manganese solution and the formation of an Orange II∙∙∙∙MnII complex was observed. The
decomposition of the dye was initiated through the subsequent addition of H2O2. It is important to
note that the catalytic oxidation of the dye by H2O2 could only be performed in an aqueous
carbonate buffer solution. No other buffer at the same pH, viz. TRIS, TAPS, HEPES or phosphate,
showed the observed catalytic reaction.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
26
2.3.3 CV studies on the complex-formation between Orange II and MnII
CV measurements of a 4 · 10-5 M MnII solution in the presence of different Orange II
concentrations were performed in order to determine the interaction between the fully aquated
MnII ions and Orange II present in the reaction mixture. Figure 2.5 (A) shows the results of
MnII∙∙∙∙Orange II complex formation in NaCl electrolyte, performed using a standard three
electrode electrochemical setup. To avoid the oxidation of MnII to MnIV, which precipitates as
MnIVO2, the potential scan was discontinued at +1.0 V, after which the reverse scan from +1.0 to -
0.8 V was started. The CVs of MnaqII in the absence of any coordinating substrate exhibit one quasi-
reversible oxidation peak at E = +0.59 V vs. Ag/AgCl and one quasi-reversible reduction peak at E =
+0.35 V, corresponding to the one electron MnIII/MnII redox couple.
In addition, CV measurements on a freshly prepared 4 · 10-5 M Orange II electrolyte solution at
pH 8.5 and 22 °C were performed. Orange II, as it can be seen in Figure 2.5 (B), undergoes two
electrochemically quasi-reversible one-electron reductions with CV half-wave potentials at Ered1 = -
0.19 V and Ered2 = +0.11 V (vs. Ag/AgCl) with a difference between the cathodic and anodic wave of
0.02 and 0.204 V, respectively. As can be seen from Figure 2.5 (B) the reduction potential of MnIII
decreased from +0.35 V to +0.28 V when Orange II was added to the solution, indicating the
stabilization of MnIII ions. In the presence of a chelating substrate, the generated MnIII complex
becomes more stable and the redox potentials attain lower values.[22]
Figure 2.5 CVs of a 4 · 10-5
M MnII solution (0.1 M NaCl electrolyte) in the presence of different Orange II
concentrations at pH 8.5 (adjusted by addition of NaOH) and 22 °C.
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
27
1a 1b
When the concentration of Orange II was increased up to 2 · 10-5 M, the presence of further
reduction peaks along with changes in the oxidation peak intensity were observed, indicating the
formation of other manganese-Orange II species as specified above.
2.3.4 DFT-calculations
To assess the coordination mode of Orange II to the MnII center, DFT (B3LYP/LANL2DZp)
calculations were performed for a series of plausible complexes. Orange II dissociates in aqueous
solution under the selected experimental conditions into an anionic sulfonate group and a cationic
sodium ion. In the presence of an unsolvated SO3- group involving charge transfer from the
electron-rich sulfonate group onto the rest of the molecule, may in general not give satisfactory
DFT results.[14b] Solvent Yellow 14, a model compound for Orange II containing no sulfonate group
was selected for the DFT study of the interaction between the MnII and the chosen azo dye. A
picture of the calculated conformers of the model compound 1 is shown in Figure 2.6.
The optimized geometry of 1a was calculated to be ca. 5.8 kcal mol-1 lower in energy than that
of 1b. Furthermore, the calculated structure of 1a was compared with X-ray structural data of
Solvent Yellow 14.[23] A good agreement between calculated and crystallographically determined
structure was found.
Figure 2.6 Optimized (B3LYP/LANL2DZp) structures of 1a and 1b with a planar geometry and dihedral angles of (a)
180.0° and (b) 178.7° about the azo group, C-N-N-C.
0 kcal mol-1 +5.8 kcal mol-1
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
28
2a 2b
0 kcal mol-1 +2.9 kcal mol-1
2c
ca. +2 kcal mol-1
According to the UV/Vis and electrochemical data presented above, Orange II can coordinate
to a fully aquated MnII center. Different plausible interaction modes of Solvent Yellow 14∙∙∙∙MnII (2)
and Solvent Yellow 14∙∙∙∙MnII∙∙∙∙Solvent Yellow 14 (3) were studied in detail. Optimized structures
of 2 adopting different coordination modes are presented in Figures 2.7 and 2.8. The studied
organic dye can coordinate to aquated MnII ion by forming two new bonds, one between MnII and
the deprotonated phenolic OH-group of 1a and the second between MnII and one of the azo
nitrogen atoms, leading to the formation of either a planar six-membered (2a) or five-membered
(2b) chelate complex.
Figure 2.7 Optimized structures of complexes 2a and 2b (B3LYP/LANL2DZp).
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
29
3
Furthermore, a second interaction mode for 2 involving a hydrogen bond between one
coordinated water molecule and the azo nitrogen atom 2c was taken into consideration. The
calculated energies indicate that 2a is energetically favoured over 2b by about 2 kcal mol-1. The N-
N bond length of 1.30 Å for 2a is nearly identical to that found in the free model molecule 1a (1.28
Å), indicating a weak interaction between the nitrogen atoms and the positively charged
manganese center.
In addition to these structures, DFT calculations were performed for a further possible
interaction of a second dye molecule with the MnII center leading to the formation of chelated
MnII innersphere complexes. Similar transition metal complexes of ortho-hydroxy azo dyes were
prepared and characterised by Drew and Landquist.[16c] The introduction of a second dye molecule
is expected to have certain advantages. In addition to the usual stabilization by the chelate-effect,
the introduction of a second molecule of 1a could result in a protecting effect on the coordination
framework. The optimized structure of 3 adopting different coordination modes is presented in
Figure 2.8.
Figure 2.8 Optimized structure of complex 3 (B3LYP/LANL2DZp).
The calculated structure of 3 shows a C2-symmetry and the axial positions are nearly
equivalent. The calculated Mn-N bond lengths in the equatorial plane for the energetically
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
30
favoured 2a (2.15 Å) and 3 (2.30 Å) are comparable with the X-ray structural data for MnII
complexes with nitrogen containing ligands such as 1,2-bis(imidazol-1-yl)ethane (bim) (2.213 -
2.294 Å),[24a] 2-[N,N-bis(2-pyridylmethyl)-amoniumethyl]-6-[N-(3,5-di-tert-buthyl-2-oxido-benzyl)-
N-(2-pyridylamino)-aminomethyl]-4-methylphenol (H2Ldtb) (2.118 - 2.237 Å)[24b] and 1,4,7-
triazacyclononane (tacn) (2.118 - 2.146).[24c]
As expected, upon coordination of two dye molecules in 3, the N-N bond distance becomes
longer (1.29 Å) than observed in the crystal structure of 1a due to the partial neutralization of the
delocalized negative charge of the nitrogen atom. The elongation of the Mn-O bond trans to the
azo group (Mn-O = 2.38 Å versus 2.06 - 2.27 Å for 2, and Mn-O = 2.27/2.26 Å versus 1.81/2.11 Å for
3) exerts a significant trans influence opposite to the Mn-N bond. The increased lability of the axial
ligand allows the subsequent interaction of the substituted transition metal atom with an oxidant,
leading to the rapid formation of active oxidizing species. Moreover, DFT calculations performed
by Blomberg et al. suggest that in the presence of weak-field ligands for MnII and MnIII, five-
coordination is also accessible whereas MnIV has a much stronger preference for six-
coordination.[25]
2.3.5 Kinetic investigations
2.3.5.1 Complex-formation between bicarbonate and MnII
The reactions between bicarbonate ions (HCO3-) and different manganese species have been
studied for several years, since aquated MnII cations themselves are actually not able to catalyze
H2O2 disproportionation. Depending on the HCO3- concentration in the reaction mixture,
MnII∙∙∙∙HCO3- complexes of different stoichiometry can be formed. Recently, it was suggested that
only the neutral MnII(HCO3-)2 complex can facilitate H2O2 disproportionation.[26] In this study the
complex-formation reaction between MnII and HCO3- was monitored using UV/Vis
spectrophotometric beside CV measurements as a function of carbonate concentration at pH 8.5.
UV/Vis spectra recorded before and after addition of HCO3- to an aqueous MnII solution showed
the formation of a new broad band at 300 nm as illustrated in Figure 2.9 (A). The time course of
the absorption band formation is shown in Figure 2.9 (B).
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
31
Figure 2.9 (A) UV/Vis spectra of an aqueous 4 · 10-4
M MnII solution before (—) and after (—) addition of 0.4 M HCO3
-
at pH 8.5. (B) Time course of the band formation at 300 nm of an aqueous 2 · 10-4
M MnII solution containing different
amounts of HCO3-.
It can be seen from Figure 2.9 (B) that formation of the manganese carbonate intermediate is
enhanced at higher carbonate concentration. The observed first order rate constants following the
induction period in Figure 2.9 (B), are directly proportional to the [HCO3-] in the range 0.01 - 0.5 M
(see Figure 2.10) with a second order rate constant of (3.6 ± 0.2) · 10-2 M-1 s-1 at 25 °C.
Figure 2.10 Plot of observed first order rate constant (kobs) for the formation of MnII∙∙∙∙HCO3
- versus the bicarbonate
concentration in the presence of 4 · 10-4
M MnII at pH 8.5 and 25 °C.
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
32
Moreover, the observed induction period is probably related to the displacement of water
from the first coordination sphere of the fully aquated MnII ion by HCO3- and subsequent
rearrangement of the coordinated ligand, viz. formation of bidentate carbonate complexes. It
should be noted that under these experimental conditions (high carbonate concentration and pH
8.5) insoluble MnIICO3 is formed as a very fine white precipitate at longer reaction times. Its
composition was confirmed by elemental analysis and IR spectroscopy. The reactivity of the
produced intermediate was tested in the oxidative degradation of Orange II by H2O2 at pH 8.5.
During the first 200 s, no change in the reactivity of the in situ formed manganese intermediate
occurs. A significant time dependent loss in catalytic efficiency of the formed MnII∙∙∙∙HCO3-
intermediate was observed after more than 200 s. An irreversible deactivation of the catalyst
occurs within less than 20 min.
On the other hand, no precipitate formation as well as no deactivation of the catalytically
active manganese intermediate could be observed in the presence of a coordinating organic
substrate, i.e. Orange II, over a long period of time (1 - 4 days) in a high carbonate (0.5 M)
containing buffer solution under these conditions. Moreover, this aspect of the stabilization of the
in situ formed active catalyst in the presence of an organic substrate is of considerable practical
interest, because its successful implementation could offer a more efficient alternative for clean
oxidation reactions.
CV measurements of freshly prepared aqueous MnII(NO3)2 solutions were performed in the
presence of different carbonate concentrations in a 0.1 M NaCl electrolyte solution at pH 8.5
(adjusted by careful addition of NaOH) and 22 °C. In the presence of a coordinating substrate, the
displacement of a coordinated water molecule from the manganese coordination sphere takes
place. By coordination of a negatively charged ligand such as HCO3- to a positively charged metal,
the peak potentials are shifted to more negative potentials compared to the fully aquated MnII
(see Figures 2.11 and 2.12).[20b]
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
33
(2.3)
(2.4)
Figure 2.11 (A) Cyclovoltammograms for 4 · 10-5
M MnII in an aqueous solution of 0.1 M NaCl and different
concentrations of NaHCO3. (B) Typical multiple scan CVs of a 4 · 10-5
M MnII solution in the presence of 0.2 M NaHCO3
and 0.1 M NaCl at pH 8.5 and 25 °C.
On increasing the carbonate concentration in solution a decrease in the peak current intensity
occurs concomitantly with peak broadening because of complexation by carbonate. Typical
multiple scan CVs of a 4 · 10-5 M MnII solution in the presence of 0.2 M NaHCO3 and 0.1 M NaCl at
pH 8.5 and 25 °C is presented in Figure 2.11 (B). In the presence of a chelating substrate, the
generated MnIII complex becomes more stable and the redox potentials attain lower values.
Moreover, at higher carbonate concentrations in the reaction mixture the presence of a second
oxidation peak at E = +0.41 V, attributed to the formation of further complexes such as proposed
in Equation 2.3 and 2.4, was observed.
By plotting the peak potential E as a function of the hydrogen carbonate concentration (see
Figure 2.12), the presence of different complex species at different carbonate concentrations is
revealed.
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
34
Figure 2.12 Plot of peak potential E as function of [HCO3-] (E vs. Ag/AgCl electrode). Reaction conditions: [Mn
II] = 4 · 10
-
5 M, [HCO3
-] = 0.1 - 05 M in 0.1 M NaCl electrolyte solution at pH 8.5 and 22 °C.
2.3.5.2 The effect of the total carbonate concentration
The effect of the carbonate concentration on the oxidative degradation of the dye was studied
at a constant pH of 8.5. The total carbonate concentration was varied between 0.05 and 0.5 M. In
the present case, the catalytic reaction leads to a square dependence of kobs on the HCO3-
concentration with a third rate constant (8.3 ± 0.3) · 10-2 M-2 s-1 (see Figure 2.13 (A)), suggesting
that two equivalents of HCO3- are involved in the oxidation mechanism. It is assumed, among
other possibilities, that one equivalent of HCO3- is required for the formation of the more reactive
[MnII(H2O)5(HCO3-)]+ intermediate, and the second equivalent of HCO3
- is necessary for the
formation of the more reactive peroxocarbonate species, known to be a versatile oxidizing agent.
Although the oxidation potential of peroxymonocarbonate and hydrogen peroxide have equal
values (E°(H2O2/H2O) = 1.77 V (vs. NHE), E°(HCO4-/HCO3
-) = 1.8 ± 0.1 V) the higher reactivity of
peroxymonocarbonate compared to that of H2O2 is attributed to carbonate being a better leaving
group than hydroxide.[27d] This is also apparent in the non-catalyzed reaction of H2O2 and Orange II
under variation of the total carbonate concentration (see Figure 2.13 (B)).
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
35
Figure 2.13 (A) Plot of the observed first order rate constants (kobs) for the MnII catalyzed degradation of Orange II vs.
the [HCO3-]. Reaction conditions: 2 · 10
-5 M Mn
II, 5 · 10
-5 M Orange II, 0.01 M H2O2, pH 8.5, 25 °C. (B) Comparison of the
absorbance changes at 480 nm versus time for the spontaneous, non-catalyzed oxidative degradation of 5 · 10-5
M
Orange II by 0.01 M H2O2 at pH 8.5 and different total carbonate concentrations.
The anionic peracid HCO4- forms in water as well as in mixed organic-water solvents at near
neutral pH values.[27] The structure of peroxymonocarbonate, HOOCO2-, is analogous to other
peroxides such as peroxymonosulfate (HOOSO3-), peroxynitrate (O2NOO-), and peracetic acid
(CH3C(O)OOH), which on the other side can only be prepared by the reaction of their
corresponding acids with H2O2 in strongly acidic aqueous solutions.[28] Borate, like bicarbonate, is
an exception and readily forms reactive peroxides in neutral to mildly alkaline solutions.[27, 29]
The peroxymonocarbonate ion has been isolated in various salts and characterized by Raman
and [27a, 30] NMR-spectroscopy.[27a, d] This inorganic peracid is known to be several orders of
magnitude more reactive toward nucleophilic substrates than H2O2 itself[29] and is formed in a
relatively fast pre-equilibrium (K = 0.32 ± 0.02 M-1)[27a] between hydrogen carbonate ions and H2O2
shown in the overall reaction in Equation 2.5.
Moreover, the reaction of H2O2 and HCO3- to form the more electrophilic HOOCO2
- (HCO4-)
occurs rapidly (t1/2 ≈ 300 s) in 1.76:1 (v/v) ethanol/water at 25 °C.[27d] This step is also regarded to
be a key aspect of several oxidation reactions.[27b-d] A more recent kinetic study performed by
(A) (B)
(2.5)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
36
Richardson and coworkers suggests that the mechanism of peroxymonocarbonate formation
proceeds via CO2 as intermediate by reaction of CO2 with H2O2 (perhydration) and its conjugate
base HOO- (base-catalyzed perhydration).[27e]
Figure 2.14 (A) 13
C - NMR spectra recorded for the in situ formation of HCO4- in an aqueous solution of 0.36 M H
13CO3
-
and 1 M H2O2 at pH 8.5 and 25 °C. (B) Time course for the consumption of H13
CO3- and the formation of HCO4 for a
solution of 0.36 M H13
CO3- + 1 M H2O2 and 0.2 M H
13CO3
- + 1 M H2O2. Concentrations were calculated from the relative
peak intensities. Reaction conditions: pH 8.5, 25 °C.
By use of 13C-labeled sodium bicarbonate we were able to follow the in situ formation of
peroxycarbonate in the absence of organic co-solvent. In Figure 2.14 (A) the time dependant
changes of the NMR spectra show a fast appearance of a new signal at 161.7 ppm, attributed to
HOO13CO2-. Figure 2.14 (B) shows the changes in concentration over time for a 0.36 M and 0.2 M
H13CO3- solution, respectively, with 1 M H2O2 at pH 8.5. The concentration values were calculated
from the relative peak intensities of the corresponding NMR-signals. According to Equation 2.6 we
were able to determine the value of the equilibrium constant as K = 0.32 ± 0.02 M-1 and thereby
confirm the literature reported value in purely aquatic solution.[27d]
In the view of these findings we decided to study the influence of carbonate on the manganese
catalyzed oxidation of Orange II by H2O2 and HCO4-, respectively. By performing the oxidation
reactions in the presence of peroxymonocarbonate instead of H2O2 in a 0.5 M carbonate
(A) (B)
(2.6) ]][[
][
3
4
22
-
-
OHHCO
HCOK
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
37
containing buffer solution at pH 8.5, no difference in the reactivity was observed (Figure 2.15 (A)).
The MnII catalyzed oxidative degradation of Orange II by using HCO4- as an oxidizing agent could be
significantly enhanced through increasing the total carbonate concentration in the reaction
mixture (Figure 2.15 (B)). This can be explained in terms of the equilibrium formulated in Equation
2.5. Based on our experimental observations and aspects reported in the literature[27b] for the MnII
catalyzed oxidation reaction by H2O2 in a carbonate containing solution, the reaction sequence
presented in Scheme 2.5 can be suggested to occur.
Figure 2.15 (A). Spectral changes observed at 480 nm for the 2 · 10-5
M MnII(NO3)2 catalyzed oxidative degradation of
2.5 · 10-5
M Orange II in the presence of (-) 0.01 M H2O2 and (-) 0.01 M HCO4-, respectively, at pH 8.5 and 0.5 M total
carbonate concentration. (B) Comparison of the absorbance changes at 480 nm versus time for the 2 · 10-5
M MnII
catalyzed oxidative degradation of 5 · 10-5
M Orange II by 0.01 M H2O2 at pH 8.5 and different carbonate
concentrations.
Addition of H2O2 to hexaaqua MnII in a carbonate solution leads to significant spectral changes
in the UV/Vis spectra during the reaction (see Figure 2.16 (A)). The initial rapid increase of the
intensity of the broad band at 300 nm, as it is illustrated in Figure 2.16 (A), is attributed to the fast
formation of [MnII(H2O)5(HCO3)]+. An isosbestic point at 330 nm suggests the formation of a new
manganese intermediate by addition of an oxidizing agent, i.e. H2O2.
According to our spectroscopic observations the formed complex with an absorption band at
400 nm could be attributed to a MnIV-η2-peroxycarbonate intermediate.[31, 32a] Based on
spectroscopic observations and data reported in the literature,[32a] the formed intermediate can
be most likely regarded to be a high valent manganese complex. Similar Rh,[32b] Pt[32c] and Fe[32a]
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
38
peroxycarbonate complexes have been isolated before and were characterized spectroscopically.
The time course of the absorption band at 400 nm at different pH is illustrated in Figure 2.16 (B).
Scheme 2.5 In situ formation of catalytically active Mn intermediates in the presence of hydrogen peroxide in a
carbonate containing aqueous solution at pH 8.5 and 25 °C.
In the absence of any stabilizing ligand, the formed complex rapidly decomposes with the
formation of catalytically inactive MnIVO2 that precipitates from solution (see Figures 2.16 (A) and
(B)). The decomposition of the active intermediate is accelerated at higher pH (see Figure 2.16
(B)). To ascertain that the formulated reaction steps in Scheme 2.5 are valid under our reaction
conditions, a systematic spectroscopic investigation at different pH values was performed.
Representative data for the reaction course at 400 nm at pH 8.5 and 9.5 are presented in Figure
2.16 (B).
Contrary to our expectations, an increase of one unit in pH resulted in an increase of the
induction period and a decrease in the manganese peroxycarbonate complex formation rate
under the mentioned reaction conditions. This could be partly due to subsequent formation of
MnII(OH)2 precipitates at higher pH and to deprotonation of HCO3- that becomes significant at pH
above 9. This results in a decrease in the HCO3- concentration in the equilibrium presented in
Equation 2.5, reducing the concentration of peroxymonocarbonate present in solution.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
39
Figure 2.16 (A) UV/Vis spectra recorded for the reaction of 2 · 10-4
M MnII(NO3)2 with 10
-3 M H2O2 in a 0.5 M HCO3
-
containing solution at pH 8.4 and 25 °C. (B) Comparison of typical absorbance at 400 nm versus time plots at pH 8.5
(—) and 9.5 (—).
2.3.5.3 Reactivity profile as function of pH
The reactivity of the catalytic system is generally influenced by the protonation state of the
substrate, the catalyst and oxidizing agent. In our work the kinetics were studied in 0.4 M NaHCO3
containing buffer solution in the pH range between 8.0 and 9.5 at 25 °C. The pH of the carbonate
buffer solution was adjusted carefully using small amounts of concentrated NaOH solution to
avoid dilution. A typical manganese catalyzed oxidative degradation of Orange II by H2O2 in a
carbonate buffer solution is presented in Figure 2.17 (A). The catalytic degradation is usually
complete within 1 - 10 min depending on the pH of the solution, the catalyst concentration, and
the H2O2 and carbonate concentrations.
The dye decomposition was followed by monitoring the spectral changes at 484 nm. The
depletion of the band at 484 nm is in general correlated with cleavage (heterolytic or homolytic) of
the azo group leading to colorless oxidation products due to the induced discontinuity in the
conjugation of the π-system in the molecule. The inset in Figure 2.17 (A) shows the first spectrum
of Orange II before the addition of the catalyst and H2O2, and the final spectrum recorded after
250 s. A decrease in the intensity of the two other bands at 270 and 300 nm was observed,
showing that further bleaching also occurs under these reaction conditions. The formation of
small, non-toxic and biodegradable organic molecules, i.e. glyoxalic acid, 4-
hydroxybenzenesulfonic acid or acetic acid through a ring-opening reaction is one of the positive
(B) (A)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
40
aspects of this process.[4] The isolation and characterization of reaction products is extremely
difficult and requires large synthetic efforts, particularly as different reaction intermediates tend
to react further under these experimental conditions. A comparison of the reaction course at
different pH value is shown in Figure 2.17 (B).
Figure 2.17 (A) First and last UV/Vis spectra of a 2 · 10-5
M MnII(NO3)2 catalyzed oxidative degradation of 5 · 10
-5 M
Orange II by 0.01 M H2O2 in a 0.4 M total carbonate containing solution at pH 8.5 and 25 °C. (B) Comparison of
absorbance at 480 nm vs. time plots for the 2 · 10-5
M MnII(NO3)2 catalyzed oxidative degradation of 5 · 10
-5 M Orange
II by 0.01 M H2O2 in a 0.4 M total carbonate containing solution at different pH values and 25 °C.
The MnII catalyzed decolorization and oxidative decomposition of Orange II was found to be
sensitive to the pH of the solution. According to our experimental data, an increase in pH resulted
in a slight decrease in the reaction rate under the above-mentioned reaction conditions and the
highest reactivity is observed at a pH between 8.2 and 8.6. Increasing the pH to > 9 leads to a
decrease in the oxidation rate for the bicarbonate-activated peroxide, which is presumably the
result of the deprotonation of HOOCO2- to form CO4
2-, a less electrophilic oxidant.[33]
At even higher pH, the decomposition of the peroxide is accelerated and may reduce the
oxidation reaction rate. Contrary to our expectations, the observed rate constants for the
decolorization reaction of Orange II are similar to the destruction rate constants of naphthalene
and benzene rings, long-lived intermediates, under the studied conditions (see Figure 2.18). Thus,
for a complete oxidation of these stable molecules higher concentrations of oxidant and catalyst
are required.
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
41
Figure 2.18 Plot of observed rate constant (kobs) calculated for the decolouring reaction followed at 480 and 300 nm,
respectively. Experimental conditions: 2 · 10-5
M MnII(NO3)2, 5 · 10
-5 M Orange II, 0.01 M H2O2, 0.4 M total carbonate
and 25 °C.
A similar screening using MnIICl2, MnII(Ac)2 and MnII(SO4)2 showed identical catalytic activity in
the oxidative degradation of Orange II by H2O2. In all cases, the manganese catalyzed oxidative
degradation of Orange II is favored by moderate alkaline pH values and vanishes completely at
very high or very low (strong acidic) values. According to the experimental observations
mentioned above, the manganese catalyzed oxidative degradation of Orange II by H2O2 in a
carbonate containing solution is considerably inhibited at higher pH values due to the lower
formation of the high valent manganese η2-peroxycarbonate complex (see Figure 2.17 (B)).
2.3.5.4 Effect of the [MnII] and [H2O2] on the oxidative reaction course
To evaluate the effect of the catalyst concentration on the manganese catalyzed oxidative
degradation of Orange II by H2O2 under catalytically relevant experimental conditions, kinetic
studies were performed for solutions in which the carbonate containing water solution with
various amounts of MnII(NO3)2 was added in the presence of 0.01 M H2O2 to a 0.05 M Orange II
solution at 25 °C. The obvious accelerating ability of the HCO3- ions prompted us to study the
catalytic reaction course in more detail at four different carbonate concentrations. The in situ
produced catalyst concentration dependence was studied at 480 nm using in situ UV/Vis
spectroscopic measurements and the kinetic traces could be adequately fitted to a single
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
42
exponential function. Plots of the observed rate constant as a function of [Mn II] at different
carbonate concentrations are presented in Figure 2.19.
Figure 2.19 MnII(NO3)2 concentration dependence of kobs. Reaction conditions: 5 · 10
-5 M Orange II, 0.01 M H2O2, pH
8.5 and 25 °C.
As it is evidenced the [MnII] dependences of the observed rate constants for the manganese
catalyzed oxidative degradation of Orange II by H2O2 in a low carbonate concentration containing
solution (0.1 - 0.3 M HCO3-) are strongly curved (higher K values, see Table 2.1) and reach a limiting
value at higher catalyst concentration. In contrast, similar data at higher carbonate concentrations
(0.4 - 0.5 M HCO3-) result in a less curved dependence of kobs on the catalyst concentration, i.e.
lower K values (see Table 2.1). The observed rate profile can be explained by the general reaction
mechanism proposed in Scheme 2.5 and simplified in Scheme 2.6. The observed rate law for the
proposed reaction steps is given by Equation 2.7. The calculated k and K values from the non-
linear concentration dependences in Figure 2.19 are summarized in Table 2.1.
Scheme 2.6 Proposed reactions steps for the formation of the catalytically active manganese intermediate in the
presence of H2O2 in a carbonate containing solution.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
43
(2.7)
[HCO3-], M k, s-1 K, M-1
0.1 0.0033 34.6 · 103
0.3 0.032 17.6 · 103
0.4 0.051 17.8 · 103
0.5 0.138 15.2 · 103
Table 2.1 The constants k and K for the MnII(NO3)2 catalyzed oxidation of Orange II by H2O2 at pH 8.5 and 22 °C
Figure 2.20 H2O2 concentration dependence of kobs. Reaction conditions: 5 · 10-5
M Orange II, 2 · 10-5
M MnII(NO3)2, pH
8.5 and 25 °C.
The effect of H2O2 on the oxidation reaction course was studied by varying its initial
concentration over a wide range, between 5 – 30 · 10-3 M. At lower H2O2 concentrations (between
1 – 5 · 10-3 M) a fast oxidation reaction occurs in the first few seconds followed by a rapid
consumption of H2O2 resulting finally in a partial and inefficient decolorization of the dye. This
prompted us to study the [H2O2] effect on the catalytic oxidation of the dye at higher
K[Mn(II)]1
kK[Mn(II)]kobs
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
44
concentrations of H2O2. The kobs values were calculated from a single exponential fit to the
absorbance at 480 nm vs. time plots and showed a linear dependence on the initial H2O2
concentration over the studied concentration range (Figure 2.20).
2.3.5.5 Stability of the in situ formed catalyst
Figure 2.21 (A). Spectral changes observed at 480 nm for the repeated addition of 5 · 10-5
M Orange II to a 2 · 10-5
M
MnII(NO3)2 solution in the presence of 0.01 M H2O2 at pH 8.5 and 0.4 M total carbonate concentration. (B) Spectral
changes observed at 480 nm for a new addition of 5 · 10-5
M Orange II and 0.01 M H2O2 to a 48 h old reaction mixture
containing the catalyst solution under the same experimental conditions as mentioned in (A).
In control experiments the stability of the in situ generated catalyst was studied by repeated
addition of dye and H2O2 to a solution of 2 · 10-5 M MnII(NO3)2 at pH 8.5 (0.4 M HCO3-) and 25 °C
(see Figure 2.21 (A) and (B)). As it can be seen in Figure 2.21 (A), the catalytic cycle could be
repeated several times without any significant loss of activity during the oxidation reaction,
indicating an excellent stability of the in situ formed catalyst. After the fifth cycle the reaction
solution containing the active catalyst was allowed to stay at ambient temperature for 48 h.
Subsequently, the catalytic activity of the in situ formed manganese complex was evaluated again
by performing the oxidation reaction in the presence of freshly added Orange II and H2O2. The
experimental results illustrated in Figure 2.21 (B) provide clear evidence for the high efficiency of
the in situ formed catalyst under the above mentioned experimental reaction conditions.
(A) (B)
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
45
2.3.6 Mechanistic aspects
Throughout this study, the oxidation reactions were carried out in a thermostated open glass
reactor vessel at ambient temperature in aqueous hydrogen carbonate containing solutions. The
readily available manganese salts, the mild reaction conditions and the operation simplicity and
practicability allow for an easy and green oxidative degradation of the studied organic dye. In
control experiments the catalytic activity of the in situ generated manganese complex was
investigated under an inert atmosphere. By performing the catalytic reaction in a closed glass
reactor under inert reaction conditions, no change in the decomposition reaction rate was
noticed. A comparison of the reaction course carried out under air and inert atmosphere is
illustrated in Figure 2.22. By performing the reaction under inert reaction conditions no significant
differences in the decomposition reaction rate was observed, indicating that HO∙ or HOO∙ radical
formation is not prevalent for this oxidation reaction.
Figure 2.22 Comparison of typical absorbance at 480 nm versus time plots of a 2 · 10-5
M MnII(NO3)2 catalyzed
oxidative degradation of 5 · 10-5
M Orange II by 0.01 M H2O2 in a 0.4 M HCO3- containing solution at pH 8.5 and
ambient temperature performed in the presence of atmospheric oxygen (—) and inert atmosphere (—), respectively.
Taking into account all obtained spectroscopic and kinetic data, the following reaction
schemes can be proposed for the MnII catalyzed oxidative degradation of Orange II by H2O2 in
carbonate solution under catalytically relevant experimental conditions. A key feature of the
proposed reaction mechanism outlined in Scheme 2.7 is that the overall oxidation of Orange II
occurs in a two electron oxidation step leading to the formation of a relatively stable high-valent
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
46
Mn=O intermediate and transfer of the oxo group to the substrate. Most of the earlier reported
papers on the oxidation reaction catalyzed by several isolated and structurally well defined
manganese complexes have emphasized the formation of a high-valent Mn=O intermediate by the
reaction of manganese with the appropriate oxidant.[11, 34a-e]
According to our observations, HCO3- ions are involved in two catalytically relevant reactions.
HCO3- ions react with aquated MnII present in solution to form a catalytically active Mn-HCO3
-
complex. HCO3- is also involved in a fast equilibrium with H2O2 to form HOOCO2
-, a versatile
heterolytic oxidant. In the following step, through nucleophilic attack of the oxidizing agent on the
MnII center, a MnII-η2-peroxycarbonate complex is formed. The remaining coordination sites in the
first shell will be occupied by water and hydroxyl at a pH between 8 and 10. The principal mode of
the formation of relatively stable high-valent Mn=O intermediates is believed to involve the
heterolytic cleavage of the peroxide bond, as shown in Scheme 2.7. An important role in the
stabilization of the formed Mn=O species is played by the electron donating bicarbonate ions.
Scheme 2.7 Proposed reaction mechanism for the MnII catalyzed oxidative degradation of Orange II by H2O2 in a
carbonate containing aqueous solution at pH between 8 - 9 and 25 °C.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
47
This may also account for the unique requirement of HCO3- in the oxidative decomposition of
Orange II catalyzed by simple manganese salts. The further coordination of the substrate followed
by an oxygen transfer step along with the second electron leads to the formation of several
oxidation products and finally to the regeneration of the catalyst. It must be noted that in the
absence of a catalyst, the oxidative degradation of Orange II by addition of an electrophilic
bleaching agent, HOOCO2-, occurs very slowly under certain reaction conditions. The oxidation
mechanism involves a nucleophilic attack of the dye at the electrophilic oxygen of HOOCO2-. In
aqueous solution, proton transfer can lead to the displacement of HCO3- and the slow formation of
oxidized substrate.
Scheme 2.8 Proposed reaction mechanism involving first substrate coordination to MnaqII in a pre-equilibrium step
during the catalyzed oxidative degradation of Orange II by H2O2 in a carbonate containing aqueous solution at pH
between 8 - 9 and 25 °C.
If substrate binding to MnII occurs before the addition of HOOCO2- to the catalyst solution,
following reactions can be assumed to take place during the reaction cycle under the chosen
experimental conditions. In line with the concerns mentioned above, the first step in Scheme 2.8
involves the prior coordination of Orange II to MnII and formation of MnII-Orange II complexes of
different stoichiometry, followed by nucleophilic attack of the oxidant on the MnII center leading
to the formation of Orange II-MnII-peroxycarbonate species. The subsequent scission of the
peroxo bond leads to the formation of high-valent oxo intermediates, as formulated in Scheme
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
48
2.8. In this case, the formed MnIV=O intermediate is stabilized by Orange II, an electron rich
organic molecule with chelating capacity. The role of Orange II as an axial ligand is also to favor the
homolytic scission of the peroxo bond leading to the MnIV=O intermediate and bicarbonate.
2.4 CONCLUSIONS
A fast and environmentally benign method for the oxidative degradation of a main pollutant
from the dye industry could be achieved using H2O2 in conjunction with catalytic amounts of
relatively non-toxic manganese salts as catalyst precursors in a carbonate containing aqueous
solution under mild reaction conditions. Screening and spectroscopic methods allowed us to study
the catalytic reaction course and to identify some key features of the reaction that reflect upon its
mechanism. Our study revealed that the oxidative degradation of the model substrate Orange II is
catalytic only in carbonate containing aqueous solution. No other buffer containing aqueous
solution could induce the oxidative degradation of Orange II by H2O2 and this led to the implication
of peroxymonocarbonate as a key molecular entity. It was found that the in situ formed high-
valent manganese intermediate possessing one hydrogen carbonate ligand is able to activate
H2O2, but decomposes rapidly with the formation of neutral MnIICO3, which precipitates from
solution as an insoluble white solid. One of the main factors affecting the process efficiency was
the stabilization of the catalytically active Mn complex. Furthermore, by addition of Orange II, the
formation of MnII∙∙∙∙Orange II complexes with different stoichiometry was observed. The
simultaneous σ,π - coordination of the organic dye is well-precedented, and recent DFT studies
support this type of complex formation. The catalytic activity of the formed intermediates was
tested under catalytic reaction conditions.
The kinetic investigations performed at different pH could provide relevant information about
the nature of the oxidizing agent involved in the reaction. It was found that the pH is a critical
issue for the rate of the oxidation process due to its influence on the deprotonation of the
bicarbonate ions, the formation of peroxycarbonate in solution, and the deprotonation of aquated
MnII. The ongoing studies are presently complemented by investigations on different organic
substrates with various functional groups in order to determine the influence of substrate
modification on the catalytic reaction cycle. DFT studies and further kinetic and spectroscopic
investigations should contribute to a better understanding of the catalytic system.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
49
2.5 EXPERIMENTAL SECTION
CHEMICALS. Orange II, certified [Acid Orange 7, C.I. 15510, sodium 4-(2-hydroxy-1-
naphthylazo)benzenesulfonate], 99 % was supplied by Sigma-Aldrich and recrystallized from a
Et2O/H2O mixture at 4 °C. Unless otherwise stated, all other dyes were commercially available
(Acros Organics, Germany) and were used without any further purification. Hydrogen peroxide 35
wt. % and sodium percarbonate were of analytical grade and provided by Acros Organics
(Germany). Different manganese salt hydrates (chloride, sulfate, nitrate, acetate, perchlorate)
were of and purchased from Acros Organics (Germany). Sodium bicarbonate, TAPS, TRIS, HEPES
and NaOH were of analytical grade and provided by Acros Organics (Germany). All buffer solutions
were prepared using Millipore Milli-Q purified water. NaH13CO3 (99 atom % 13C) and DMSO-d6
(99.96 % atom % D) were from Sigma-Aldrich.
13C-NMR MEASUREMENTS. All 13C-NMR measurements were performed by using DMSO-d6 internal
standard in a glass capillary. The concentrations of PAA and AcOH were calculated from the
relative peak intensities using the Lorenz fit obtained by the NMRICMA program (developed at the
“Institut de chimie minérale et analytique” of the University of Lausanne in the group of Prof. A. E.
Merbach) of the NMR data. 13C-NMR spectra were recorded at a frequency of 100 MHz on a
Bruker Advance DRX 400WB spectrometer equipped with a superconducting BC-94/89 magnet
system.
CYCLOVOLTAMMETRIC MEASUREMENTS. Cyclovoltammetric (CV) measurements were performed in a
one-compartment three-electrode cell using a gold working electrode (Metrohm) with a
geometrical surface of 0.7 cm2 connected to a silver wire pseudo-reference electrode and a
platinum wire serving as counter electrode (Metrohm). Measurements were recorded with an
Autolab PGSTAT 30 unit at room temperature. The working electrode surface was cleaned using
0.05 μm alumina, sonicated and washed with water every time before use. The working volume of
10 ml was deaerated by passing a stream of high purity N2 through the solution for 15 min prior to
the measurements and then maintaining an inert atmosphere of N2 over the solution during the
measurements. All CVs were recorded for the reaction mixture with a sweep rate of 50 mV s-1 at
25 °C. Potentials were measured in a 0.5 M NaCl/NaOH electrolyte solution and are reported vs.
an Ag/AgCl electrode.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
50
DFT-CALCULATIONS. Unrestricted B3LYP/LANL2DZp hybrid density functional calculations,[35a-c,] i.e.,
with pseudo-potentials on the heavy elements and the valence basis set[35] augmented with
polarization functions,[36] were carried out using the Gaussian 03[37] suite of programs. The relative
energies were corrected for zero point vibrational energies (ZPE). The resulting structures were
tested for stability and characterized as minima by computation of vibrational frequencies and the
wave functions were tested for stability. Further energy computations were carried out using the
conductor-like polarizable continuum model (CPCM).[38]
ELEMENTAL ANALYSIS. The measurements were carried out on an elemental analyzer Euro EA 3000
instrument from Hekaltech Gmbh. The analytical method is based on the complete instantaneous
oxidation of the sample by “flash combustion” at 1000 °C, which converts all organic and inorganic
substances into combustion products. The resulting combustion gases are swept into the
chromatographic column by the carrier gas (He) where they are separated and detected by a
thermal conductivity detector.
IR MEASUREMENTS. IR spectra were recorded as KBr pellets using a Mattson Infinity FTIR instrument
(60 AR) at 4 cm-1 resolution in the 400 - 4000 cm-1 range.
SPECTROPHOTOMETRIC TITRATION. UV/Vis spectra were recorded on a Shimdazu UV-2101
spectrophotometer at 25 °C. In the experiments concerning the complexation by different dyes
the tandem cuvette with two separate compartments (0.44 cm path length each), was filled with 1
ml 5 · 10-5 M Orange II stock solution in one, and different concentrations of an aqueous
MnII(NO3)2 solution in the other compartment. The cuvette was placed in the thermostated cell
holder of the spectrophotometer for 10 min. UV/Vis spectra were recorded before and after
mixing the solutions.
SYNTHESIS OF INSOLUBLE MnIICO3. In a 150 ml round flask 3.36 g (0.4 M) NaHCO3 were dissolved in
100 ml doubly distilled water and the pH of the solution was set at 8.5 upon addition of small
amounts of concentrated NaOH solution. To the freshly prepared carbonate solution 1 g (0.04 M)
MnII(NO3)2 was added. The mixture was stirred at room temperature for 15 min during which
MnIICO3 ∙ H2O formed as a white precipitate. The product was filtrated and washed several times
with large amounts of water. IR (KBr pellets): ν (cm-1) 3421 (m), 1416 (vs), 862 (s), 725 (m).
Elemental analysis for Mn1C1H2O4: calculated %C 9.03, %H 1.52. Found: %C 9.38, %H 1.52.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
51
SYNTHESIS OF ORANGE II∙∙∙∙MnII COMPLEX. In a 50 ml Schlenk tube 0.014 g (2 · 10-3 M) Orange II were
dissolved in 20 ml doubly distilled water and an aqueous solution of 0.01 g (2 · 10-3 M) MnII(NO3)2
were added dropwise under continuous stirring. The solution mixture was kept for several hours
at room temperature. The formed precipitate was filtered and dried at room temperature. IR (KBr
pellets): ν (cm-1) 3527 (vs), 1619 (s), 1511 (s), 1383 (vs), 1262 (m), 1171 (s), 1120 (s), 1034 (s), 1007
(s), 829 (s), 759 (s), 696 (m), 644 (m), 595 (m). Elemental analysis for Mn1C16H18O10N3S1Na1:
Calculated: %C 36.79, %H 3.47, %N 8.04, %S 6.14, %O 30.63. Found: %C 29.44, %H 3.39, %N 8.05,
%S 4.77, %O 30.31.
SYNTHESIS OF ORANGE II∙∙∙∙MnII∙∙∙∙ORANGE II COMPLEX. An aqueous solution of 0.005 g (1 · 10-3 M)
Mn(NO3)2 was added under continuous stirring to a 0.014 g (2 · 10-3 M) Orange II water solution at
room temperature. The pale yellow precipitate was collected by filtration and dried in air. IR (KBr
pellets): ν (cm-1) 3390 (s), 1619 (s), 1570 (m), 1554 (m), 1520 (vs), 1393 (m), 1260 (m), 1169 (vs),
1119 (vs), 1033 (vs), 1007 (s), 828 (s), 758 (s), 695 (m), 644 (m), 593 (m). Elemental analysis for
Mn1C32H32O15N5S2Na2: Calculated: %C 43.1, %H 3.62, %N 7.23, %S 7.19, %O 26.91. Found: %C
42.99, %H 3.75, %N 7.23, %S 7.00, %O 25.67.
GENERAL PROCEDURE FOR THE H2O2 CATALYZED DYE DEGRADATION REACTIONS. The manganese catalysts were
freshly dissolved in water before use. To a freshly prepared sodium bicarbonate solution, an
adequate amount of NaOH was added to adjust the pH of the solution. Under isothermal
conditions, the desired amount of a concentrated manganese solution was added together with
Orange II, previously dissolved in an aqueous bicarbonate solution, and H2O2. All kinetic data were
obtained by recording time-resolved UV/Vis spectra using a Hellma 661.502 - QX quartz Suprasil
immersion probe attached via optical cables to a 150 W Xe lamp and a multi-wavelength J & M
detector, which records complete absorption spectra at constant time intervals. In a thermostated
open glass reactor vessel equipped with a magnetic stirrer, the freshly prepared catalyst solution
and H2O2 were added to 40 ml of 5 · 10-5 M dye at a pH ranging from 8 to 10 at 25 °C. All kinetic
measurements were carried out under pseudo-first order conditions (i.e. 50 ≤ [oxidant]/[MnII] ≤
1000). The pH of the aqueous solutions was carefully measured using a Mettler Delta 350 pH
meter previously calibrated with standard buffer solutions at two different pH values (4 and 10).
The kinetics of the oxidation reaction was monitored at the λmax of the corresponding dye. First
order rate constants, where possible, were calculated using Specfit/32 and Origin (version 7.5)
software. To estimate the effect of the catalyst and H2O2 concentrations on the catalytic reaction
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
52
at different carbonate concentrations, stopped-flow kinetic measurements were carried out
additionally using an SX.18MV stopped-flow instrument from Applied Photophysics.
2.6 REFERENCES AND NOTES
[1] (a) Mielgo, I.; López, C.; Mareira, M. T.; Feijoo, G.; Lema, J. M. Biotechnol. Prog. 2003, 19, 325-
331; (b) Hastie, J.; Bejan, D.; Teutli-Leon, M. N.; Bunce, J. Ind. Eng. Chem. Res. 2006, 45, 4898-
4909; (c) Park, H.; Choi, W. J. Photochem. Photobiol. 2003, 159, 241-247; (d) Baughman, G. L..;
Weber, E. J. Environ. Sci. Technol. 1994, 28, 267-276; (e) Robinson, T.; McMullan, G.;
Marchant, R.; Nigam, P. Bioresource Technol. 2001, 77, 247-255.
[2] (a) Collins, T. J. Acc. Chem. Res. 1994, 27, 279-285; (b) Hage, R.; Iburg, J. E.; Kerschner, J.;
Koek, J. H.; Lempers, E. L. M.; Martens, R. J.; Racherla, U. S.; Russel, S. W.; Swarthoff, T.; van
Vliet, M. R. P.; Warnaar, J. B.; van der Wolf, L.; Krijnen, B. Nature 1994, 369, 637-639; (c)
Bhattacharyya, A.; Kawi, S.; Ray, M. B. Catal. Today 2004, 98, 431-439; (d) Ramirez, J. H.;
Maldonado-Hodar, F. J.; Perez-Cadenas, A. F.; Moreno-Castilla, C.; Costa, C. A.; Madeira, L. M.
Appl. Catal. B 2007, 75, 312-323.
[3] Ramirez, J. H.; Costa, C. A.; Madeira, L. M.; Mata, G.; Vicente, M. A.; Rojas-Cervantes, M. L.;
López-Peinado, A. J.; Martin-Aranda, R. M. Appl. Catal. B 2007, 71, 44-56.
[4] Chahbane, N.; Popescu, D.-L.; Mitchell, D. A.; Chanda, A.; Lenoir, D.; Ryabov, A. D.; Schramm,
K.-W.; Collins, T. J. Green Chem. 2007, 9, 49-57.
[5] (a) Zhiyong, Y.; Bensimon, M.; Laub, D.; Kiwi-Minsker, L.; Jardim, W.; Mielczarski, E.;
Mielczarski, J.; Kiwi, J. J. Mol. Catal. A 2007, 272, 11-19; (b) Chen, X.; Qiao, X.; Wang, D.; Lin, J.;
Chen, J. Chemosphere 2007, 67, 802-808.
[6] Kim, W.; Hur, S. G.; Swang, S. J.; Park, H.; Choi, W.; Choy, J. H. Adv. Funct. Mater 2007, 17,
307-314.
[7] Wieprecht, T.; Xia, J.; Heinz, U.; Dannacher, J.; Schlingloff, G. J. Mol. Catal. A 2003, 203, 113-
128.
[8] (a) Payne, G. B.; Deming, P. H.; Williams, P. H. J. Org. Chem. 1961, 26, 659-663; (b) Majetich,
G.; Hicks, R. Synlett 1996, 649-651; (c) McKillop, A.; Sanderson, W. R. Tetrahedron 1995, 51,
6145-6166.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
53
[9] Barker, J. E.; Ren, T. Tetrahedron Lett. 2004, 45, 4681-4683.
[10] (a) Oakes, J. Eur. Pat. Appl. 145091, 1985; (b) Feng, X.-M.; Wang, Z.; Bian, N.-S.; Wang, Z.-L.
Inorg. Chim. Acta 2007, 360, 4103-4110; (c) Kim, C.; Chen, K.; Kim, J.; Que Jr., L. J. Am. Chem.
Soc. 1997, 119, 5964-5965; (d) de Vos, D. E.; Bein, T. J. Chem. Soc., Chem. Commun. 1996,
917-918; (e) Battioni; P.; Renaud, J. P.; Bartoli, J. F.; Rein-Artiles, M.; Fort, M.; Mansuy, D. J.
Am. Chem. Soc. 1988, 110, 8462-8470.
[11] Nakayama, N.; Tsuchiya, S.; Ogawa, S. J. Mol. Catal. A 2007, 277, 61-71.
[12] Baes, C. F.; Mesmer, J. R.; Mesmer, R. E. Am. J. Sci. 1981, 281, 935-962.
[13] Liu, S.-Y.; Nocera, D. N. Tetrahedron Lett. 2006, 47, 1923-1926.
[14] (a) Zollinger, H. Color Chemistry, 3rd ed., Wiley-VCH: Weinheim, Germany, 2003; (b) Abbott, L.
C.; Batchelor, S. N.; Oakes, J.; Gilbert, B. C.; Whitwood, A. C.; Smith, J. R. L.; Moor, J. N. J. Phys.
Chem. 2005, 109, 2894-2905. (c) Hodges, G. R.; Lindsay-Smith, J. R.; Oakes, J. J. Chem. Soc.,
Perkin Trans. 2 1998, 3, 617-628; (d) Oakes, J.; Gratton, P. J. Chem. Soc., Perkin Trans. 2 1998,
9, 1857-1864; (e) Feng, W.; Nansheng, D.; Helin, H. Chemosphere 2000, 41, 1233-1238.
[15] (a) Lindsay-Smith, J. R.; Gilbert, B. C.; Mairata i Payeras, A.; Murray, J.; Lowdon, T. R.; Oakes,
J.; Pons i Prats, R.; Walton, P. H. J. Mol. Catal. A 2006, 251, 114-122; (b) Tokuda, J.; Oura, R.;
Iwasaki, T.; Takeuchi, Y.; Kashiwada, A.; Nango, M. Coloration Technol. 2000, 116, 42-47.
[16] (a) Yoshida, T.; Sawada, S. Bull. Chem. Soc. Jpn. 1975, 48, 345-346; (b) López, C.; García-
Monteagudo, J. C.; Moreira, M. T.; Feijoo, G.; Lema, J. M. Enzyme Microbiol. Technol. 2007,
42, 70-75; (c) Drew, H. D. K.; Landquist, J. K. J. Chem. Soc. 1938, 292-304.
[17] (a) Reeves, R. L.; Maggio, M. S.; Harkaway, S. A. J. Phys. Chem. 1979, 83, 2359-2368; (b)
Asakura, T.; Ishida, M. J. Colloid. Interface Sci. 1989, 130, 184-189; (c) Murakami, K. Dyes
Pigm. 2002, 53, 31-43.
[18] Benesi, H.; Hildebrand, H. J. Am. Chem. Soc. 1949, 71, 2703-2807.
[19] Job, P., Compt. Rend. 1925, 180, 928-930.
[20] (a) Nadtochenko, V.; Kiwi, J. J. Chem. Soc., Faraday Trans. 1997, 93, 2373-2778; (b) Bauer, C.;
Jacques, P.; Kalt, A. Chem. Phys. Lett. 1999, 307, 397-406.
[21] Sychev, A. Y.; Pfannmueller, U.; Isak, V. G. Zh. Fiz. Khim. 1983, 8, 1974-1978.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
54
[22] (a) Lesht, D.; Bauman, J. E. Inorg. Chem. 1978, 17, 3332-3334; (b) Dasgupta, J.; Tyryshkin, A.
M.; Kozlov, Y. N.; Klimov, V. V.; Dismukes, G. C. J. Phys. Chem. B 2006, 110, 5099-5111.
[23] Olivieri, A. C.; Wilson, R. B.; Paul, I. C.; Curtin, D. Y. J. Am. Chem. Soc. 1989, 111, 5525-5532.
[24] (a) Zhu, X.; Li, B.; He, X.; Zhang, Y. J. Chem. Crystallogr. 2005, 35, 443-446; (b) Bortoluzzi, A. J.;
Neves, A.; Couto, R. A. A.; Perelta, R. A. Acta Crystallogr., Sect. C: Cryst. Struct. Commun.
2006, C62, m27; (c) Bennur, T.H.; Srinivas, D.; Sivasanker, S.; Puranik, V. G. J. Mol. Catal. A
2004, 219, 209-216.
[25] Blomberg, M. R.; Siegbahn, P. E. M.; Styring, S.; Babcock, G. T.; Aakermark, B.; Korall, P. J. Am.
Chem. Soc. 1997, 119, 8285-8292.
[26] Tikhonov, K. G.; Zastrizhanaya, O. M.; Kozlov, Y. N.; Klimov, V. V. Biochem. 2006, 71, 1270-
1277.
[27] (a) Flanagan, J.; Jones, D. P.; Griffith, W. P.; Skapski, A. C.; West, A. P. J. Chem. Soc., Chem.
Commun. 1986, 20-21; (b) Lane, B. S.; Vogt, M.; DeRose, V. J.; Burgess, K. J. Am. Chem. Soc.
2002, 124, 11946-11954; (c) Richardson, D. E.; Regino, C. A. S.; Yao, H. R.; Johnson, J. V. Free
Radical Biol. Med. 2003, 35, 1538-1550; (d) Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett,
D. A. J. Am. Chem. Soc. 2000, 122, 1729-1739; (e) Bakhmutova-Albert, E. V.; Yao, H.; Denevan,
D. E.; Richardson, D. E. Inorg. Chem. 2010, 49, 11287-11296.
[28] (a) Greenspan, F. P. J. Am. Chem. Soc. 1946, 68, 907; (b) Ball, E.; Edwards, J. O. J. Am. Chem.
Soc. 1956, 78, 1125-1129; (c) Keith, W. G.; Powell, R. E. J. Chem. Soc. A 1969, 90-98.
[29] Swern, D. Organic Peroxides, Ed. Wiley, New York, 1970, 313.
[30] Jones, D. P.; Griffith, W. P. J. Chem. Soc., Dalton Trans. 1980, 2526-2532.
[31] MnIV in the presence of further Mn(II) leads to the formation of Mn(III), a highly unstable
oxidant, according to the reaction: MnIV + MnII 2 MnIII
[32] (a) Hashimoto, K.; Nagatomo, S.; Fujnami, S.; Furutachi, H.; Ogo, S.; Suzuki, M.; Uehara, A.;
Maeda, Y.; Watanabe, Y.; Kitagawa, T. Angew. Chem., Int. Ed. 2002, 114, 1202-1205; (b)
Aresta, M.; Tommasi, I.; Quaranta, E.; Fragale, C.; Mascetti, J.; Tranquille, M.; Galan, F.;
Fouassier, M. Inorg. Chem. 1996, 35, 4252-4260; (c) Hayward, P. J.; Blake, D. M.; Wilkinson,
G.; Nyman, C. J. J. Am. Chem. Soc. 1970, 92, 5873-5878.
[33] Balagam, B.; Richardson, D. E. Inorg. Chem. 2008, 47, 1173-1178.
2. Metal ion - catalyzed oxidative degradation of Orange II by H2O2
55
[34] (a) Khenkin, A. M.; Kumar, D.; Shaik, S.; Neumann, R. J. Am. Chem. Soc. 2006, 128, 15451-
15460; (b) Siegbahn, P. E. M.; Carbtree, R. H. J. Am. Chem. Soc. 1999, 121, 117-127; (c) Kondo,
M.; Mitsui, T.; Ito, S.; Kondo, Y.; Ishigure, S.; Dewa, T.; Yamashita, K.; Nakamura, J.; Oura, R.;
Nango, M. J. Colloid Interface Sci. 2007, 310, 686-689; (d) Shul´pin, G. B.; Matthes, M. G.;
Romakh, V. B.; Barbosa, M. I.; Aoyagi, J. L. T.; Mandelli, D. Tetrahedron 2008, 64, 2143-2152;
(e) Anastasi, A. E.; Walton, P. H.; Lindsay-Smith, J. R.; Sameera, W. M. C.; McGrady, J. E. Inorg.
Chim. Acta 2008, 361, 1079-1086.
[35] (a) Becke, A. D. J. Phys. Chem. 1993, 97, 5648-5652; (b) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev.
B, 1988, 37, 785-789; (c) Stephens, P. J.; Devlin, F. J.; Chabalowski, C. F.; Frisch, M. J. J. Phys.
Chem. 1994, 98, 11623-11627; (d) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 270-283; (e)
Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 284-298; (f) Hay, P. J.; Wadt, W. R. J. Chem.
Phys. 1985, 82, 299-310.
[36] Huzinaga, S. (Ed.), Gaussian Basis Sets for Molecular Calculations, Elsevier, Amsterdam 1984.
[37] Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.;
Montgomery Jr., J. A.; Vreven, T.; Kudin, K. N.; Burant, J. C.; Millam, J. M.; Iyengar, S. S.;
Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani, G.; Rega, N.; Petersson, G. A.;
Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima,
T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.; Li, X.; Knox, J. E.; Hratchian, H. P.; Cross, J. B.;
Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.;
Pomelli, C.; Ochterski, J. W.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg,
J. J.; Zakrzewski, V. G.; Dapprich, S.; Daniels, A. D.; Strain, M. C.; Farkas, O.; Malick, D. K.;
Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul, A. G.; Clifford, S.;
Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.; Martin, R. L.;
Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M.
W.; Johnson, B.; Chen, W.; Wong, M. W.; Gonzalez, C.; Pople, J. A. Gaussian, Inc., Wallingford
CT, 2004.
[38] (a) Barone, V.; Cossi, M. J. Phys. Chem. A 1998, 102, 1995-2001; (b) Cossi, M.; Rega, N.;
Scalmani, G.; Barone, V. J. Comp. Chem. 2003, 24, 669-681.
2. Metal ion-catalyzed oxidative degradation of Orange II by H2O2
56
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
57
3 COMPARATIVE STUDY OF A MNII-MONOMER AND THE
CORRESPONDING OXO-BRIDGED MN2III/IV-DIMER
3.1 GENERAL REMARK
The following chapter is based on the original publication: Comparative study of the catalytic
activity of [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 in the H2O2 induced oxidation of organic
dyes in carbonate buffered aqueous solution, Sabine Rothbart, Erika Ember and Rudi van Eldik,
Dalton Trans. 2010, 39, 3264-3272.
3.2 INTRODUCTION
The development of transition metal complexes as effective catalysts for a wide range of
oxidative transformation reactions and the understanding of the latter processes still represent a
fundamental challenge to inorganic chemists. Inspired by biological systems and under
consideration of environmental aspects, the complexes of choice are bearing an iron or
manganese active site. Manganese complexes with their redox rich chemistry are well-known to
act as outstanding oxygenation catalysts. During recent years many studies have provided
important insight into the relevant manganese catalyzed mechanisms of processes such as
photosynthetic water oxidation,[1] bleaching[2] or epoxidation.[2a, 3] Nevertheless, the nature of the
reactive intermediates in these processes remains to be clarified. Several studies give credence to
the proposal that oxo-bridged MnIV or MnV complexes[1a, 3, 4] are involved as key intermediates. A
fundamental class of complexes among these are µ-oxo-bridged dinuclear complexes of high-
valent MnIII and MnIV, since their involvement in epoxidation,[2a] bleaching[2a] and oxidation of
hydrocarbons[5] and alcohols[5, 6] is known from the literature. As a result, a manifold of applicable
high-valent MnIII and MnIV model complexes have been synthesized, characterized and
investigated towards their catalytic activity in oxygenation reactions.[7] However, there still
remains a vivid debate in the literature regarding the nuclearity of the oxidizing species, since
other studies also endorse monomeric species in higher oxidation states as reactive intermediates
in manganese catalyzed oxidation and epoxidation reactions.[8, 9]
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
58
Previously we reported kinetic and mechanistic studies on the manganese ion catalyzed
oxidation reaction of various highly stable organic dyes by addition of H2O2 to a MnII containing
buffer solution.[10] The results indicated the formation of high valent MnIV=O as a key catalytic
intermediate by addition of H2O2 to a MnII containing buffer solution.[10] In the view of this finding
we extended our work to elucidate the role of bis-µ-oxo-dimanganese(III,IV) species as
intermediates. It is generally acknowledged that these complexes can be formed from MnII
complexes with labile coordination sites in the presence of a stabilizing ligand and an oxidizing
agent like hydrogen peroxide in basic aqueous media.[11] To have a closer look at the nature of
putative intermediates and the rate-limiting step of the reaction sequence, a comparative study
on the reactivity of a [Mn2III/IV(µ-O)2]3+-dimer and its readily accessible mononuclear analogous
MnII complex in the hydrogen peroxide assisted catalyzed oxidation of Orange II was performed.
As the study progressed, it was important to check our results also for other dyes as model
substrates. We selected p-nitrophenol and Morin, a dye of polyphenolic structure (Scheme 3.1),
which is a common element of chromophores present in fruit, vegetable and tea and therefore an
adequate model substrate for bleach processes.[12]
Scheme 3.1 Structure of the model substrates Orange II, Morin and p-nitrophenol.
A commonly used oxidant with proximity to many biological systems and with beneficial
environmental properties, like no additional redundant waste from organic peroxides or it
economically price, is hydrogen peroxide. Although H2O2 is thermodynamically a potent oxidant its
activity is often kinetically hindered under mild experimental conditions as low temperature or
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
59
atmospheric pressure. Hence, several methods for the activation of H2O2 involve the use of
conventional catalysts along with co-catalytic additives. In our previous work we showed the
importance of the in situ formed peroxycarbonate anion,[10] HOOCO2-, which is known to be
several orders of magnitude more reactive than hydrogen peroxide itself.[13] Well-established is
also the application of reactive peroxyacids, formed from carboxylic acids, or the use of persulfate
or perborate in basic solution.[14]
3.3 RESULTS AND DISCUSSION
3.3.1 Kinetic measurements of the catalyzed dye degradation with H2O2
Typical spectral changes observed during the reaction of Orange II as substrate with hydrogen
peroxide in the presence of a catalyst are exemplarily presented by the inset of Figure 3.1. The
decrease of the absorbance with time reflects oxidative degradation of the substrate and can
hence be considered as a measure for the catalytic activity of the complexes. It is noteworthy that
not only decolourization at 484 nm, which corresponds to the dominating hydrazone tautomer n -
π* transition at this pH,[15] was observed but also destruction of the stable naphthalene subunit of
the substrate, shown by the absorbance decrease at 300 nm (see inset of Figure 3.1) with identical
observed rate constants. The rate enhancement in the presence of the catalyst becomes evident
when it is compared with the non-catalyzed oxidation of Orange II by H2O2 under the same
reaction conditions. When the catalysts were used in equimolar concentrations in terms of the
manganese content, surprisingly no difference in the catalytic activity of the µ-oxo-bridged
dinuclear and the mononuclear MnII complex was observed (Figure 3.1). This also holds true under
conditions of catalytic conversion, i.e. tenfold substrate excess compared to the catalyst
concentration (see Figure 3.5). According to our experiments, the observed rate constants for the
oxidation of the naphthalene subunit and the destruction of the azo linkage are very similar. To
verify this correlation the reaction was also performed with Morin and p-nitrophenol as substrate.
For both dyes the reaction course was monitored at a wavelength of 400 nm, which corresponds
to the absorbance of the resorcinol subunit in Morin.[16] In agreement with the results obtained for
Orange II, the oxidation course with Morin and p-nitrophenol as substrate show identical reactivity
for the in situ formed catalyst as well as for the readily prepared catalysts [MnII(bpy)2Cl2] and
[Mn2III/IV(µ-O)2(bpy)4](ClO4)3∙2H2O. In the case of flavonoidic compounds such as Morin, it is known
that upon oxidation an initial increase followed by the oxidative decay at 330 nm can be observed,
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
60
whereas at 400 nm the decay of the phenolic part starts immediately. This is due to the extension
of the π-system of the dye, which results in an initial intensification of the corresponding
absorbance, as it is often observed for flavonoidic dyes. (for typical spectral changes and kinetic
traces see S3.1 and S3.2, Supporting Information). For all three substrates, the [Mn2III/IV(µ-O)2]3+-
dimer and the mononuclear analogous MnII complex gave the same results.
Figure 3.1 Observed spectral changes and kinetic traces recorded at 484 nm for the (·····) uncatalyzed degradation of
Orange II in comparison to the catalyzed reaction. Reaction conditions: 0.1 M HCO3-, 5 · 10
-5 M Orange II, 0.015 M
H2O2, pH 9.0, room temp.. (—) 4 · 10-5
M [MnII(bpy)2Cl2], (—) 2 · 10
-5 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3, (—) 4 · 10
-5 M
[MnII(bpy)2Cl2] in the presence of
tBuOH, (—) 2 · 10
-5 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3 in the presence of
tBuOH.
In order to exclude the involvement of free radical processes the experiment was repeated in
the presence of a strong radical scavenger like tBuOH. It reacts with hydroxyl radicals by
generation of inert intermediates, which cause termination of the radical chain reaction.
Although hydroxyl radicals might be present during the H2O2 induced reaction, they do not
participate in the oxidation process catalyzed by [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3,
since in the presence of tBuOH no negative effect on the reaction course was observed (Figure
3.1).
The influence of pH on the observed rate constant for the decolourization of Orange II (Orange
II: pKa = 11.4[17]) catalyzed by [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 (Figure 3.2), shows
that both catalysts reach a maximum reactivity at a pH of about 8.7, indicating that both catalysts
must involve the same reactive intermediate in the catalytic cycle independent of the pH. At
higher pH the observed rate constants show a drastic decrease. This observation is consistent with
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
61
earlier studies [10, 18] and may be due to deprotonation of HOOCO2- to form the less electrophilic
oxidant CO42-.[18] The additionally performed measurements of the oxidation reaction of p-
nitrophenol at different pH values for either of the catalysts (Figure S3.3, Supporting Information)
confirm the assumption that the observed pH profile is caused by the difference in reactivity of the
oxidizing species and not by substrate effects. Due to the enhanced decomposition of hydrogen
peroxide at lower pH, originating from the parallel catalase-like reaction and leading to incomplete
decolourization of the solution, further investigations were performed at a slightly higher pH of
9.0.
Figure 3.2 Observed rate constants measured at 484 nm as a function of pH. Reaction conditions: 0.1 M HCO3-, 0.01 M
H2O2, 5 · 10-5
M Orange II, room temp., (■) 4 · 10-5
M [MnII(bpy)2Cl2], (▲) 2 · 10
-5 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3.
Earlier investigations on oxidation reactions with H2O2 in bicarbonate buffered solution
revealed the importance of the in situ formed peroxycarbonate anion,[10, 18, 19, 20, 21] HOOCO2-,
which is known to be a stronger oxidant than hydrogen peroxide itself.[13] Despite the low
formation constant of HOOCO2- (viz. K = 0.32 ± 0.02 M-1 [21]), the latter is formed in a relatively fast
pre-equilibrium step.[21] The effect of the carbonate concentration on the oxidation course was
studied in a total carbonate concentration range between 0.01 and 0.4 M at a constant pH of 9.0.
This unambiguously has a strong effect on the observed rate constants for the catalyzed oxidation
of Orange II by either of the complexes (see Figure 3.2). The second order dependence of kobs on
the bicarbonate concentration leads to almost identical rate constants for the µ-oxo bridged and
the mononuclear MnII complexes determined by linear fit of the correlation presented in Figure
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
62
3.3, viz. for MnII(bpy)2Cl2 (15.1 ± 0.4) · 10-2 M-2 s-1 and for [Mn2III/IV(µ-O)2(bpy)4]3+ (16.9 ± 0.5) · 10-2
M-2 s-1. Control experiments at different ionic strengths with additional NaNO3 or NaCl at constant
bicarbonate concentrations showed no influence of ionic strength or the counter ion on the
observed reaction courses (see Figure S3.4, Supporting Information).
Figure 3.3 Observed second order dependence on the total carbonate concentration for the oxidation of Orange II by
both complexes measured at 484 nm. Reaction conditions: 5 · 10-5
M Orange II, 0.01 M H2O2, pH 9.0, room temp., (—
■—) 4 · 10-5
M MnII(bpy)2Cl2, (—▲—) 2 · 10
-5 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3 · 2H2O.
This observation is consistent with earlier findings[10] and can be interpreted in terms of two
equivalents of bicarbonate that are involved in catalytically relevant reaction steps of the
oxidation mechanism of both catalysts. One is the mentioned relatively fast pre-equilibrium
between hydrogen peroxide and bicarbonate,[18, 20, 21] that results in the formation of the
peroxycarbonate anion. Whereas control experiments at pH 9 in the absence of any carbonate
revealed no catalytic effect, which confirms the crucial role of in situ formed peroxycarbonate for
the observed catalyzed oxidation reaction. Therefore the oxidation reaction was additionally
performed under identical reaction conditions without any carbonate present in CHES buffered
solution at pH 9.0. In Figure 3.4 the catalyzed oxidation reaction of Orange II for either of the
catalysts without carbonate present (c) is compared to the same experiments in 0.1 M (b) and 0.3
M (a) carbonate buffered solution under identical reaction conditions. Although the reaction of
the manganese catalysts with HOO- (note pKa of H2O2 > 9.0) could be in general possible, the
tremendous rate enhancing effect of carbonate on the catalyzed oxidation reaction of Orange II
becomes evident for either of the catalysts. This points to the crucial pre-equilibrium of
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
63
peroxycarbonate formation resulting in the second order dependence of the observed reaction
rate on the total carbonate concentration.
Figure 3.4 Kinetic traces recorded for the oxidation of 5 · 10-5
M Orange II by 4 · 10-5
M [MnII(bpy)2Cl2] and 2 · 10
-5 M
[Mn2III/IV
(µ-O)2(bpy)4](ClO4)3 in the presence of 0.01 M H2O2 and (a) 0.1 M HCO3- at pH 9.0 (b) 0.3 M HCO3
- at pH 9.0,
compared to (c) the same reaction in the absence of any carbonate in 0.1 M CHES buffer. Reactions conditions: pH 9.0,
room temp..
The participation of the second equivalent of bicarbonate could be due to coordination of an
electron donating bicarbonate ion to an aquated MnII-bipyridine precursor complex to form a
more reactive MnII-bipyridine-bicarbonate species or to stabilize the oxidatively formed high-
valent MnIV=O intermediate. It was shown before that monocarboxylate ions, i.e. acetate,
bicarbonate and formate, enhance the catalytic activity in the H2O2 induced epoxidation reactions,
but do not considerably disturb the first coordination sphere of well defined MnII complexes.[22]
Whereas dicarboxylate additives such as oxalate have been reported to act as strong ligands to the
MnII catalyst, this leads to lower reactivity in the epoxidation reaction.[22] Consequently, a
displacement of the chelating 2,2’-bipyridine ligands by bicarbonate/carbonate is unlikely, but a
monodentate binding mode is in principle possible if easily exchangeable coordination sites exist,
as in the case of the earlier reported catalytic activity of simple metal salts.[10]
Performing kinetic measurements under variation of the initial hydrogen peroxide
concentration (between 2.5 · 10-3 and 5.0 · 10-2 M) enabled us to evaluate the effect of H2O2 on
the reaction course of the catalytic oxidation of the different substrates. At high catalyst
concentration and a small hydrogen peroxide excess, the strong hydrogen peroxide consumption
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
64
(catalase-like reaction) resulted in an inefficient decolourisation due to the loss of pseudo-first-
order reaction conditions. In order to suppress this effect, we investigated the [H2O2] influence on
the dye degradation at catalyst concentrations of 5 · 10-6 M for [MnII(bpy)2Cl2] and 2.5 · 10-6 M for
[Mn2III/IV(µ-O)2(bpy)4](ClO4)3. Both catalysts again show identical rates within the experimental
error, when used at equimolar manganese content, confirming the suggestion that in both cases
the same catalytically active intermediate is generated. A good linear correlation between the
corresponding kobs value and the oxidant concentration was observed (see Figure 3.5).
Figure 3.5 Plot of the observed rate constants determined at 484 nm as a function of [H2O2] for both studied
complexes. Reaction conditions: 0.1 M HCO3-, 5 · 10
-5 M Orange II, pH 9.0, room temp., (solid lines) 5 · 10
-6 M
[MnII(bpy)2Cl2], (dotted lines) 2.5 · 10
-6 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3 · 2H2O.
Substrate (λmax / nm) k / M-1s-1 (MnII(bpy)2Cl2) k / M-1s-1 ([Mn2III/IV(µ-O)2(bpy)4](ClO4)3)
OII (484) (7.1 ± 0.2) · 10-2 (7.2 ± 0.2) · 10-2
Mo (400) (4.5 ± 0.1) · 10-1 (4.5± 0.1) · 10-1
PNP (400) (1.10 ± 0.06) · 10-2 (1.12 ± 0.05) · 10-2
Table 3.1 Second order rate constants for both studied complexes in the hydrogen peroxide assisted oxidative
degradation of different dyes. Reaction conditions: 0.1 M HCO3
-, 5 · 10
-5 M dye, 5 · 10
-6 M [Mn
II(bpy)2Cl2], 2.5 · 10
-6 M
[Mn2III/IV
(µ-O)2(bpy)4](ClO4)3, pH 9.0, room temperature.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
65
From the slope of the plot with a zero intercept the second order rate constants for the µ-oxo-
bridged dinuclear and the mononuclear MnII complex for all dyes were determined to be identical
within the experimental error (Table 3.1). The values of k for the oxidation of the phenolic group in
Morin are more than one order of magnitude higher than for the oxidation of p-nitrophenol. This
trend can be accounted for in terms of the electronic properties of the benzene core that is
oxidized. Whereas Morin bears two activating, i.e. electron donating, OH-substituents, p-
nitrophenol contains only one OH- and in addition one strongly deactivating NO2-substituent. As a
consequence, p-nitrophenol is less nucleophilic and a less favourable substrate for an oxidizing
species such as an electrophilic, high-valent manganese complex, which is reflected by the large
difference in the second-order rate constants for both dyes (Table 3.1).
To evaluate the effect of the catalyst concentration on the oxidative degradation of Orange II
by H2O2, kinetic studies over a wide concentration range were performed under catalytically
relevant conditions. Since the concentration of the co-catalyst carbonate has such a severe effect
on the oxidative reaction course, the experiment was also conducted at a higher [HCO3-] and
showed similar tendencies (see Figures S3.5, Supporting Information). The kobs values for
[MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 plotted as a function of the manganese
concentration (note that the dimer has double the Mn content) showed identical saturation
kinetics and reached a limiting value at high catalyst concentrations.
The observed trend of the resulting kobs dependencies of [MnII(bpy)2Cl2] and [Mn2III/IV(µ-
O)2(bpy)4](ClO4)3 at different carbonate content can be attributed to a pre-equilibrium step, i.e. a
rapid equilibration reaction prior to the rate-determining reaction step, involving the
peroxycarbonate anion and one equivalent of manganese, which finally results in a mononuclear
high-valent Mn-oxo species.
3.3.2 The reaction of the catalysts with H2O2 in carbonate buffered solution
3.3.2.1 EPR-spectroscopic measurements
In order to obtain more information on the high-valent manganese species formed in the
presence of H2O2 in bicarbonate containing solution, EPR spectroscopic measurements were
performed. Samples were transferred to a 3 mm EPR quartz tube immediately after mixing the
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
66
reactants and frozen in liquid nitrogen to quench the reaction. The X-band EPR spectra at 6 K of
bicarbonate (0.1 M) and hydrogen peroxide (0.04 M) containing solutions of 1 · 10-4 M
[MnII(bpy)2Cl2] and 5 · 10-5 M [Mn2III/IV(µ-O)2(bpy)4](ClO4)3, exhibit a similar set of features over a
wide range from 50 to 500 mT (see Figure 3.6).
Figure 3.6 X-band EPR spectrum of 0.1 M bicarbonate containing solutions at pH 9.0 of (—) 1 · 10-4
M [MnII(bpy)2Cl2]
and (—) 5 · 10-5
[Mn2III/IV
(µ-O)2(bpy)4](ClO4)3 with 0.04 M H2O2 immediately after mixing. EPR conditions: 8.98 GHz, 6 K,
1 mW microwave power, modulation amplitude 20 mT.
Both spectra show well resolved and slightly anisotropic sextets of the characteristic 55Mn
hyperfine lines (I = 5/2) at g ≈ 2, which is consistent with a high-spin S = 5/2 MnII. In addition,
broad and unresolved signals of lower amplitude at g ≈ 4-5 could be observed for both catalysts,
which is indicative of a mononuclear MnIV complex. Several examples with similar g-values for
mononuclear MnIV species have been reported before[19, 23] and some also show a lower amplitude
feature at g ≈ 2.[24] If more oxidant is used an increase in the characteristic MnIV signal could be
observed (see Figure S3.6, Supporting Information), which is in agreement with our experimental
results for the catalytic oxidation reaction.
In the present case strong EPR spectral features exhibited by the MnIV=O intermediate are
observed in the g ≈ 2 region with only minor components at lower field. In general EPR spectra of
d3 MnIV ions in an axial field (E/D = 0) are often difficult to interpret on account of the dependence
on the magnitude of the zero-field splitting parameters.[25] If the axial zero-field splitting
parameter D is high, the spectrum is dominated by the lower field signal and shows only minor g ≈
2 contribution, whereas small D values result in spectra with inverted signal intensity for these two
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
67
signals. For instance, this is the case for sulfur-containing thiohydroxamate[25b] and
dithiocarbamate[25c] MnIV complexes. In addition we performed EPR experiments in the presence
of tBuOH, i.e. a radical scavenger, to check whether the presence of MnIII or the participation of
radical processes can account for the observed intermediates and therefore the oxidation
reaction. In any case no difference in the observed results was found regardless whether a radical
scavenger was present in the reaction mixture or not (for EPR spectra see Figure S3.6, Supporting
Information).
Throughout our studies no multi-line signals as known for mixed-valent oxo-bridged
manganese complexes could be observed, indicating that these species are not prevalent in the
catalytic solution. This is in agreement with recent results for acetylacetone-based Schiff bases of
manganese which show that in catalytic alkene epoxidation with H2O2 and carbonate present in
acetone/MeOH mixtures, the catalytic centre is a mononuclear Mn complex.[22] The presence of
residual MnIII/IV dimers cannot be absolutely excluded since the excessively present MnII causes a
higher peak intensity at equal concentration levels than the dimer. Nevertheless, the obtained EPR
results in conjunction with the observed reactivity pattern provide a strong indication that a
monomeric MnIV complex is present as potential oxidizing species in the reaction solution and
most notably for both catalysts. The presence of monomeric MnII indicates that in both cases the
key step of the oxidation process is a two-electron oxidation of a monomeric MnII precursor to a
MnIV=O intermediate.
3.3.2.2 UV/Vis spectroscopic measurements
The reaction of [MnII(bpy)2Cl2] or [Mn2III/IV(µ-O)2(bpy)4](ClO4)3 with hydrogen peroxide was also
studied with stopped-flow rapid scan UV/Vis spectroscopy. Since the concentration of possible
high-valent manganese species are rather low in solution, higher concentrations of the manganese
catalysts had to be used to obtain more significant spectral changes. Therefore, a small amount of
freshly prepared catalyst stock solution was added to a 0.1 M bicarbonate containing buffer
solution at pH 9.0 and room temperature, and spectra were recorded at time intervals of 0.2 s
over the first 5 seconds of the reaction after mixing. The spectral changes that accompany the
reaction of [MnII(bpy)2Cl2] with H2O2 are presented in Figure 3.7 and show the formation and
partial decay of an absorption band at 445 nm. It is supposed that this transition arises from a
LMCT process. Similar spectral assignments that the lower energy bands in monomeric MnIV
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
68
complexes arise from LMCT transitions have been reported before for a complex with a tacn
derived ligand[26a] or for [MnIV(bpy)(N3)4].[26b]
Figure 3.7 Spectra recorded for the reaction of 0.2 · 10-3
M [MnII(bpy)2Cl2] with 5 · 10
-3 M H2O2 in 0.1 M HCO3
- solution
at pH 9.0, room temperature. Inset: Corresponding kinetic trace at 445 nm.
In this case it is probably an oxo to MnIV LMCT band, since the band at 450 nm was not
observed in the absence of H2O2. Based on spectroscopic observations, the rapid formation and
decay of a band at 445 nm could be attributed to the intermediate formation of a mononuclear
MnIV-oxo intermediate. The subsequent partial decay of the observed species is attributed to the
parallel oxidation of hydrogen peroxide which is accompanied by gas evolution as no other
substrate was present (Figure 3.7 inset).
The reaction of [Mn2(µ-O)2(bpy)4](ClO4)3 with H2O2/HCO3- was studied in a similar way by use
of rapid scan UV/Vis spectroscopy, although it was not possible to use stopped-flow techniques
with aqueous stock solutions due to the insufficient stability of the µ-oxo-bridged core in aqueous
solution over longer time scale.[27] The spectral changes within the first few seconds after mixing
of the dimer with a H2O2/bicarbonate containing solution were much more intense, but it cannot
be excluded that a similar band at 445 nm underlies the spectral changes of the remaining di-µ-
oxo bridged dimer. (Figure S3.7, Supporting Information). The absence of the characteristic 16-line
signal in the performed EPR experiments indicates that the dimer is no longer the predominant
species under catalytically relevant conditions. Presumably it is rapidly converted to a
mononuclear manganese species due to the presence of bicarbonate/carbonate, which causes the
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
69
dinuclear species to dissociate,[28] and due to the excess hydrogen peroxide being able to reduce
the bis-µ-oxo-bridged dinuclear complex.[9]
With respect to the UV/Vis measurements and the data obtained by means of EPR
spectroscopy, we conclude that a key step of the oxidation reaction for both catalysts is the two
electron oxidation of a monomeric MnII-precursor complex to form an high-valent MnIV=O species,
which could act as potential oxidizing species.
3.3.3 In situ formation of the active species in the catalytic oxidation reaction
In addition, the in-situ formation of the catalytically active species was checked by performing
the oxidation reaction of the different substrates when MnIICl2 and ligand (2,2'-bipyridine) were
added simultaneously to the solution in a molar ratio of 2:1 (bipyridine:MnII). The 2:1
stoichiometry would in principle allow the rapid formation of the µ-oxo-bridged dimers on
stabilizing manganese in its higher oxidation states. Compared to the kinetic traces obtained with
the synthesized catalysts [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3, no difference in the
catalytic reaction course was found for any of the three substrates when such a mixture instead of
the synthesized catalysts was used. This observation, exemplarily shown for Orange II in Figure 3.8
(see Figures S3.1 and S3.2, Supporting Information for the corresponding experiments with Morin
and p-nitrophenol), implies the formation of the same oxidising intermediate regardless whether
the catalyst is formed in situ or used in an isolated form.
The catalytic activity of [MnII(bpy)2Cl2] and [Mn2III/IVO2(bpy)4](ClO4)3 is higher than the activity
of the simple MnIICl2 salt under identical reaction conditions, since complexation of the MnII ion by
a chelating ligand is in general advantageous for the catalytic activity. From the fact that the same
catalytically active form is accessible by in situ preparation of the complex, and the known
catalytic effect of simple MnII salt on the oxidative degradation of various model substrates,[10] the
question arises what the actual MnII precursor form of the active intermediate is that accounts for
the observed oxidation process.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
70
Figure 3.8 Kinetic traces recorded at 484 nm for the catalyzed degradation of Orange II. Reaction conditions: 0.1 M
HCO3-, 5 · 10
-5 M Orange II, 0.015 M H2O2, pH 9.0, room temp., (—) 4 · 10
-5 M [Mn
II(bpy)2Cl2], (—) 2 · 10
-5 M [Mn2
III/IV(µ-
O)2(bpy)4](ClO4)3, (—) 4 · 10-5
M MnIICl2·4H2O and 8 · 10
-5 M 2,2'-bipyridine added to the reaction mixture.
This becomes more evident when the oxidation reaction of Orange II is performed with a MnII
salt in the presence of different concentrations of the bipyridine ligand. Increasing concentrations
of the ligand lead to faster reaction rates as shown in Figure 3.9, which can be ascribed to the
stabilization of the MnII-precursor complex in slightly basic solution, and thereby the stabilization
of a high-valent Mn-oxo intermediate. At higher bipyridine concentration the rapid consumption
of hydrogen peroxide disrupts the oxidative degradation process. This is supported by the
observation that the catalytic reaction can be started again by addition of a fresh amount of
hydrogen peroxide (Figure 3.9), i.e. the disruption is not due to catalyst deactivation, since the
initial oxidation rate remains unchanged at ligand to metal ratios higher than 2:1. The higher
coordination number reached at higher bipyridine concentrations favours decomposition of H2O2
above substrate oxidation. Similar observations were made for the other model substrates Morin
and p-nitrophenol (see Figure S3.8, Supporting Information).
In the case of aquated MnII ions in the presence of bipyridine different complexes, viz. 1:1, 1:2
and 1:3 (M:L), are expected to be formed in solution. Working with a 1:2 stoichiometry seemed
very appropriate for our purpose, since the existence of free coordination sites on the precursor
complex is considered to be of importance for the catalytic activity and allows as a matter of
principle the coordination of hydrogen peroxide and the subsequent formation of µ-oxo-bridged
dimers.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
71
Figure 3.9 Observed kinetic traces recorded at 484 nm for the degradation of Orange II. Reaction conditions: 0.1 M
HCO3-, 0.5-4 % CH3CN, 5 · 10
-5 M Orange II, 0.015 M H2O2, pH 9.0, room temp., 4 · 10
-5 M Mn
IICl2 and different
concentrations of 2,2'-bipyridine from 0 to 1.6 · 10-4
M added to the reaction mixture.
It was also shown before that in the absence of a chelating ligand in bicarbonate containing
solution, catalytically active Mn-bicarbonate complexes of different stoichiometry can be
formed.[10] As a consequence, several differently substituted MnII precursor complexes with
bipyridine and bicarbonate ligands may in principle be able to catalyse the oxygenation reaction.
However, the formal metal-to-ligand ratio in solution and the stoichiometry in the solid state may
differ from the complex species that are finally responsible for the observed catalytic behaviour.
This has to be considered particularly in the case of MnII complexes, since they are known to be
kinetically labile.[29]
3.3.4 Precursor complex equilibria in solution
Consequently, we investigated the formation of potential MnII precursor complexes with
bipyridine in carbonate buffered solution by the method of continuous variation of Job.[30]
Therefore the change in absorbance between before and after mixing of different molar fractions
of MnII and ligand was studied in 0.1 M carbonate buffer solution at pH 9.0 in the presence of a
small amount of acetonitrile (5 %) to improve the solubility of the bipyridine ligand. Although the
absorbance changes are rather small (Figure 3.9, inset), a non-linearity in the Job-plot correlation
(Figure 3.10) at a molar fraction of 0.5 can be observed.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
72
Figure 3.10 Observed spectral changes at 280 nm before and after mixing of different molar fractions of MnIICl2 and
bipyridine, (■) overall concentration 1 · 10-4
M, (●) overall concentration 5 · 10-5
M. Reaction conditions: 5% CH3CN,
0.1 M HCO3-, pH 9.0, room temperature. Inset: Observed spectral changes before and after mixing of Mn
IICl2 and
bipyridine at a molar ratio of 0.65.
This indicates a hidden local extremum as expected for a 1:1 complex. However the broad
maximum at a molar fraction of about 0.7 clearly points to the formation of a 2:1 complex in
solution. Based on this finding the corresponding kinetic data for the MnII catalyzed oxidation
reaction in the presence of different bipyridine concentrations can be interpreted in the following
way: An increase in the ligand concentration shifts the equilibrium from the 1:1 to the 1:2
bipyridine substituted precursor complex which shows a higher catalytic reactivity. Additional
ligand beyond a stoichiometry of 1:2 especially favors the parallel hydrogen peroxide
decomposition and causes disruption of the catalytic cycle. This indicates that the 1:2 form is the
catalytically active precursor species responsible for the oxidation process and emphasizes the
importance of labile coordination sites for a high catalytic activity. In contrast, the 1:3 form favors
parallel decomposition (catalase) of H2O2, which presumably occurs via an outer-sphere electron
transfer process. As expected, UV/Vis experiments verified that the chloride ions are substituted
by water as soon as the [MnII(bpy)2Cl2] complex is dissolved, so that the catalyst precursor can be
described as aqua or hydroxo form of MnII bearing two bipyridine ligands.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
73
3.3.5 Mechanistic aspects
On the basis of the collected data, the following mechanistic conclusions can be drawn to
provide a basis for the understanding of the catalytic activation of H2O2 by manganese complexes
in carbonate buffered aqueous solution. Of significant importance is the observation that in this
catalytic system, i.e. aqueous bicarbonate buffered solution in a pH range of 8-10, it does not
make a difference if a [Mn2III/IV(µ-O)2]3+-dimer or its mononuclear analogous MnII complex is used.
Both catalysts show identical catalytic oxidative reactivity for different kinds of dyes, indicating
that the same oxidizing intermediate is formed under all conditions. Rapid-scan UV/Vis and EPR
spectroscopy provided evidence for this active intermediate being a mononuclear high-valent
MnIV-oxo species. Kinetic and EPR data could be verified in the presence of a radical scavenger,
which rules out an involvement of radical species in the catalytic oxidation process. This in turn
indicates that MnIII species are not relevant for the observed reactivity, since their presence would
be expected in the case of a radical mechanism. The parallel presence of MnII in the EPR spectra
led to the conclusion that the key feature of the reaction of either of the catalysts is a two electron
oxidation process of a MnII precursor complex to form a high-valent MnIV-oxo intermediate.
Accordingly, the main reaction sequence of the studied reaction can be simplified as presented in
Scheme 3.2.
Above all, it is essential to consider the role of peroxycarbonate as a more powerful oxidant
than H2O2 itself.[13] As shown before, this oxidant is formed in situ in an pre-equilibrium process
between hydrogen peroxide and bicarbonate.[18-21] Furthermore, in the case of the Mn catalyzed
substrate oxidation by H2O2, no oxidative reactivity was observed in the absence of a carbonate
buffer.[10] It is proposed that the peroxycarbonate anion results in the formation of a manganese-
η2-peroxycarbonate intermediate as the consequence of a fast nucleophilic attack of the
peroxycarbonate anion on the MnII center. Similar Rh[31] and Fe[32] peroxycarbonate complexes
have been isolated before.
The observed saturation kinetics for the catalytic reaction course implies that this reaction
starts with a pre-equilibrium step. Subsequently, the manganese-η2-peroxycarbonate is believed
to undergo a proton induced heterolytic O-O-bond cleavage to produce a high-valent MnIV=O
intermediate. Since such a MnIV=O intermediate has a high oxidation ability, it will react with
substrate in an oxygen transfer process to lead to the oxidation products and completion of the
catalytic cycle by regeneration of the MnII precursor complex.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
74
Scheme 3.2 Proposed reaction cycle for the catalyzed oxidative degradation of a substrate by H2O2 (S = substrate, L =
water or bicarbonate) by a possible MnII precursor complex formed from [Mn
II(bpy)2Cl2] or [Mn2
III/IV(µ-
O)2(bpy)4](ClO4)3.
Of significant importance is the drastic enhancement of the oxidation activity by the presence
of bicarbonate/carbonate as co-catalyst. Its contribution is firstly related to the in situ formation of
the highly reactive peroxycarbonate and secondly to its capability to serve as electron donating
ligand to form more reactive precursor complexes like it was shown for the manganese ion
catalyzed activation of H2O2 in the absence of any other ligand.[9] Literature findings[22] and the Job
plot analysis support the assumption that the bipyridine ligands in the catalytically active
precursor form are not displaced by bicarbonate. The reactions performed with different
MnII:bipyridine ratios are in favour of an MnII(bpy)2 complex form as precursor for the catalytically
active intermediate in the oxidation process, whereas the two remaining coordination sites are
most likely easily exchangeable bicarbonate and/or aqua and hydroxo ligands.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
75
3.4 CONCLUSION
In summary, we were able to demonstrate by a well studied example that an elaborated oxo-
bridged catalyst in higher oxidation states is not always required for efficient catalysis. Instead, the
same results can be achieved by the use of [MnII(bpy)2Cl2] as catalyst to oxidize various organic
substrates in aqueous solution under mild conditions. Moreover, this catalytic activity is also
accessible by simple in situ preparation of the 1:2 complex from a MnII salt and the ligand within
the reaction mixture. In addition, the present study provides insight into mechanistic aspects and
the nature of the oxidizing species occurring during the reaction of [MnII(bpy)2Cl2] and [Mn2III/IV(µ-
O)2(bpy)4](ClO4)3 with hydrogen peroxide in aqueous carbonate containing solution. As a result we
were able to show that the key feature of both catalysts is a two electron oxidation from a MnII
precursor to a high-valent MnIV=O intermediate, which represents the potential oxidizing species.
3.5 EXPERIMENTAL SECTION
CHEMICALS. Orange II, certified [Acid Orange 7, C.I. 15510, sodium 4-(2-hydroxy-1-
naphthylazo)benzenesulfonate], 99 % was supplied by Sigma-Aldrich and recrystallized from a
Et2O/H2O mixture at 4 °C. Unless otherwise stated, all other dyes were commercially available
(Acros Organics, Germany) and were used without any further purification. Hydrogen peroxide 35
wt. % was of analytical grade and provided by Acros Organics (Germany). Analytical grade
Hydrogen peroxide 30 wt. % was supplied by Sigma-Aldrich. Different manganese salt hydrates
(chloride, sulfate, nitrate, acetate, perchlorate) were of analytical grade and purchased from Acros
Organics (Germany). Sodium bicarbonate and CHES, were of analytical grade and provided by
Acros Organics (Germany). All buffer solutions were prepared using Millipore Milli-Q purified
water. Analytical grade 2,2'-bipyridine, CH3CN, tBuOH and sodium hydroxide were purchased from
Acros Organics (Germany).
MANGANESE CATALYSTS. [MnII(bpy)2Cl2] and [Mn2III/IV(µ-O)2(bpy)4](ClO4)3∙2H2O were synthesized as
reported before[27] and confirmed by elemental analysis. Anal. calculated for [MnII(bpy)2Cl2]: C:
54.82, H: 3.68, N: 12.79; Found: C: 54.91, H: 3.70, N, 12.83. Anal. calculated for [Mn2III/IV(µ-
O)2(bpy)4](ClO4)3∙2H2O: C: 43.64, H: 3.30, N: 10.18; Found: C: 43.62, H: 3.29, N, 10.14. The
oxidation state of [Mn2III/IV(µ-O)2(bpy)4](ClO4)3∙2H2O was verified by EPR spectroscopy.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
76
GENERAL PROCEDURE FOR THE H2O2 CATALYZED DYE DEGRADATION REACTIONS. The manganese catalysts were
freshly dissolved in water (in the case of Mn2III/IV(µ-O)2(bpy)4](ClO4)3∙2H2O the catalyst was freshly
dissolved in acetonitrile) before use. To a freshly prepared sodium bicarbonate solution, an
adequate amount of NaOH was added to adjust the pH of the solution. Under isothermal
conditions, the desired amount of a concentrated manganese solution was added together with
Orange II, previously dissolved in an aqueous bicarbonate solution, and H2O2.
All kinetic data were obtained by recording time-resolved UV-Vis spectra using a Hellma
661.502 - QX quartz Suprasil immersion probe attached via optical cables to a 150 W Xe lamp and
a multi-wavelength J & M detector, which records complete absorption spectra at constant time
intervals. In a thermostated open glass reactor vessel equipped with a magnetic stirrer, the freshly
prepared catalyst solution and H2O2 were added to 40 ml of 5 · 10-5 M dye at a pH ranging from 8
to 10 at 25 °C. The pH of the aqueous solutions was carefully measured using a Mettler Delta 350
pH meter previously calibrated with standard buffer solutions at two different pH values (4 and
10). The kinetics of the oxidation reaction was monitored at the λmax of the corresponding dye.
First order rate constants, where possible, were calculated using Specfit/32 and Origin (version
7.5) software. To estimate the effect of the catalyst and H2O2 concentrations on the catalytic
reaction at different carbonate concentrations, stopped-flow kinetic measurements were carried
out additionally using an SX.18MV stopped-flow instrument from Applied Photophysics.
EPR MEASUREMENTS. EPR spectroscopy was performed using a JEOL continuous wave spectrometer
JES-FA200 equipped with a X-band Gunn diode oscillator, a cylindrical mode cavity and a helium
cryostat. Samples were transferred to a 3 mm EPR quartz tube immediately after mixing the
reactants and frozen in liquid nitrogen to quench the reaction before data collection. Typical
spectrometer conditions were 8.98 GHz microwave frequency, 1 and 2 mW microwave power,
respectively, and 20 mT modulation amplitude.
SPECTROPHOTOMETRIC TITRATION. UV/Vis spectra were recorded on a Shimdazu UV-2101
spectrophotometer at 25 °C. A tandem cuvette with two separate compartments was filled with
stock solutions of MnIICl2 in water/acetonitrile (20:1) solution in one, and bipyridine in 0.2 M
carbonate buffer/acetonitrile (20:1) in the other compartment. The absorbance at 280 nm was
studied for different molar fractions of manganese and bipyridine before and 120 s after mixing of
the reactants.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
77
3.6 REFERENCES AND NOTES
[1] (a) Jacobsen, E. N.; Ojima, I. (Ed.), Catalytic Asymmetric Synthesis, VCH Publishers, New York,
1993, 159; (b) Katsuki, T. Coord. Chem. Rev. 1995, 140, 189-214; (c) Limburg, J.; Vrettos, J. S.;
Liable-Sands, L. M.; Rheingold, A. L.; Crabtree, R. H.; Brudvig, G. W. Science 1999, 283, 1524-
1527; (d) Yachandra, V. K.; Sauer, K.; Klein, M. P. Chem. Rev. 1996, 96, 2927-2950; (e)
Ruettinger, W.; Dismukes, G. C. Chem. Rev. 1997, 97, 1-24.
[2] (a) Hage, R.; Iburg, J. E.; Kerschner, J.; Koek, J. H.; Lempers, E. L. M.; Martens, R. J.; Racherla,
U. S.; Russel, S. W.; Swarthoff, T.; van Vliet, M. R. P.; Warnaar, J. B.; van der Wolf, L..; Krijnen,
B. Nature 1994, 369, 637-639; (b) de Boer, J. W.; Browne, W. R.; Brinksma, J.; Alsters, P. L.;
Hage, R.; Feringa, B. L. Inorg. Chem. 2007, 46, 6353-6372.
[3] (a) Groves, J. T.; Stern, M. K.; J. Am. Chem. Soc. 1987, 109, 3812-3814; (b) Groves, J. T.; Stern,
M. K.; J. Am. Chem. Soc. 1988, 110, 8628-8638.
[4] (a) Limburg, J.; Vrettos, J. S.; Chen, H.; de Paula, J.C.; Crabtree, R. H.; Brudvig, G. W. J. Am.
Chem. Soc. 2001, 123, 423-430; (b) Barton, D. H. R.; Choi, S.-Y.; Hu, B.; Smith, J. A.
Tetrahedron 1998, 54, 3367-3378; (c) Tagore, R.; Chen, H.; Crabtree, R. H.; Brudvig, G. W. J.
Am. Chem. Soc. 2006, 128, 9457-9465.
[5] (a) Tetard, D.; Rabion, A.; Verlhac, J.-B.; Guilhem, J. J. Chem. Soc., Chem. Commun. 1995, 5,
531-532; (b) Romakh, V. B.; Therrien, B.; Suess-Fink, G.; Shul’pin, G. B. Inorg. Chem. 2007, 46,
1315-1331; (c) Vincent, J. M.; Menage, S.; Lambeaux, C.; Fontecave, M. Tetrahedron Lett.
1994, 34, 6287-6290; (d) Fish, R. H.; Fong, R. H.; Oberhausen, K. J.; Konings, M. S.; Vega, M. C.;
Christou, G.; Vincent, J. B.; Buchanan, R. M. New. J. Chem. 1992, 16, 727-733.
[6] Handervan, C.; Hage, R.; Feringa, B. L. Chem. Commun. 1997, 5, 419-420.
[7] (a) Towle, D. K.; Botsford, C. A.; Hodgson, D. J. Inorg. Chim. Acta 1988, 141, 167-168; (b) Oki,
A. R.; Glerup, J.; Hodgson, D. J. Inorg. Chem. 1990, 29, 2435-2441; (c) Schindler, S.; Walter, O.;
Pederson, J. Z.; Toflund, H. Inorg. Chim. Acta 2000, 303, 215-219.
[8] (a) Yin, G.; Buchalova, M.; Danby, A. M.; Perkins, C. M.; Kitko, D.; Carter, J. D.; Scheper, W. M.;
Busch, D. H. Inorg. Chem. 2006, 45, 3467-3474; (b) Adam, W.; Roschmann, K. J.; Saha-Möller,
C. R.; Seebach, D. J. Am. Chem. Soc. 2002, 124, 5068-5073. (c) Finney, N. S.; Pospisil, P. J.;
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
78
Chang, S.; Palucki, M.; Konsler, R. G.; Hansen, K. B.; Jacobsen, E. N. Angew Chem., Int. Ed.
1997, 36, 1720-1723.
[9] Murphy, A.; Stack, T. D. P. J. Mol. Catal. A 2006, 251, 78-88.
[10] (a) Ember, E.; Rothbart, S.; Puchta, R.; van Eldik, R. New J. Chem. 2009, 33, 34-49; (b) Ember,
E.; Gazzaz, H.; Rothbart, S.; Puchta, R.; van Eldik, R. Appl. Catal. B. 2010, 95, 179-191.
[11] Suzuki, M.; Tokura, S.; Suhura, M.; Uehara, A. Chem. Lett. 1988, 3, 477-480.
[12] (a) Blum, H.; Mayer, B.; Pegelow, U. Patent Appl. DE 19620267, 1996; (b) Blum, H.; Kramer, R.
Patent Appl. WO 0032731, 1999.
[13] Swern, D. Organic Peroxides, Wiley, New York, 1970, 313.
[14] Lindsay-Smith, J. R.; Gilbert, B. C.; Mairata i Payeras, A.; Murray, J.; Lowdon, T. R.; Oakes, J.;
Pons i Prats, R.; Walton, P. H. J. Mol. Catal. A 2006, 251, 114-122.
[15] Oakes, J.; Gratton, P. J. Chem. Soc., Perkin Trans. 2 1998, 9, 1857-1864.
[16] (a) Boudet, A. C.; Cornard, J. P.; Merlin, J. C. Spectrochim. Actat A 2000, 56, 829-839; (b)
Castro, G. T.; Blanco, S. E. Spectrochim. Acta A 2004, 60, 2235-2241.
[17] Sabnis, R. W. Handbook of Acid-base Indicators, CRC Press (Taylor and Francis Group),
London, 2008.
[18] (a) Balagam, B.; Richardson, D. E. Inorg. Chem. 2008, 47, 1173-1178; (b) Siegel, A.; Siegel, H.
Metal ions in biological systems, Marcel Dekker, New York, 1999, Vol. 37, 330.
[19] Flanagan, J.; Jones, D. P.; Griffith, W. P.; Skapski, A. C.; West, A. P. J. Chem. Soc., Chem.
Commun. 1986, 20-21.
[20] Lane, B. S.; Vogt, M.; DeRose, V. J.; Burgess, K. J. Am. Chem. Soc. 2002, 124, 11946-11954.
[21] Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett, D. A. J. Am. Chem. Soc. 2000, 122, 1729-
1739.
[22] Stamatis, A.; Doutsia, P.; Vartzouma, C.; Christoforidis, K. C.; Deligiannakis, Y.; Louloudi, M. J.
Mol. Catal. A 2009, 297, 44-53.
[23] (a) Romain, S.; Baffert, C.; Duboc, C.; Leprêtre, J.-C.; Deronzier, A.; Collomb, M.-N. Inorg.
Chem. 2009, 48, 3125-3131; (b) Parsell, T. H.; Behan, R. K.; Green, M. T.; Hendrich, M. P.;
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
79
Borovik, A. S. J. Am. Chem. Soc. 2006, 128, 8728-8729; (c) Campbell, K. A.; Lashley, M. R.;
Wyatt, J. K.; Nantz, M. H.; Britt, R. D. J. Am. Chem. Soc. 2001, 123, 5710-5719.
[24] Saadeh, S. M.; Lah, M. S.; Pecoraro, V. L. Inorg. Chem. 1991, 30, 8-15.
[25] (a) Pedersen, E.; Toftlund, H. Inorg. Chem. 1974, 13, 1603-1612; (b) Pal, S.; Ghosh, P.;
Chakravorty, A. Inorg. Chem. 1985, 24, 3704-3706; (c) Brown, K. L.; Golding, R. M.; Healy, P.
C.; Jessop, K. J.; Tennant, W. C. Aust. J. Chem. 1974, 27, 2075-2081.
[26] (a) Belal, A. A.; Chaudhuri, P.; Fallis, I.; Farrugia, L. J.; Hartung, R.; Macdonald, N. M.; Nuber,
B.; Peacock, R. D.; Weiss, J.; Wieghardt, K. Inorg. Chem. 1991, 30, 4397-4402; (b) Dave, B. C.;
Czernuszewicz, R. S. J. Coord. Chem. 1994, 33, 257-269.
[27] (a) McCann, S.; McCann, M.; Rev. Casey, M.T.; Jackman, M.; Devereux, M.; McKee, V. Inorg.
Chim. Acta 1998, 279, 24-29; (b) Cooper, S. R.; Calvin, M. J. Am. Chem. Soc. 1977, 99, 6623-
6630.
[28] Limburg, J.; Vrettos, J. S.; Crabtree, R. H.; Brudvig, G. W.; de Paula, J. C.; Hassan, A.; Barra, A.-
L.; Duboc-Toia, C.; Collomb, M.-N. Inorg. Chem. 2001, 40, 1698-1703.
[29] Hogg, R.; Wilkins, R. C. J. Chem. Soc. 1962, 341-350.
[30] Job, P. Ann. Chim. 1928, 9, 113.
[31] Aresta, M.; Tommasi, I.; Quaranta, E.; Fragale, C.; Mascetti, J.; Tranquille, M.; Galan, F.;
Fouassier, M. Inorg. Chem. 1996, 35, 4254-4260.
[32] Hashimoto, K.; Nagatomo, S.; Fujnami, S.; Furutachi, H.; Ogo, S.; Suzuki, M.; Uehara, A.;
Maeda, Y.; Watanabe, Y.; Kitagawa, T. Angew. Chem., Int. Ed. 2002, 41, 1202-1205.
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
80
3.7 SUPPORTING INFORMATION
Figure S3.1 Observed spectral changes (inset A) and kinetic traces recorded at 400 nm (A) and 330 nm (B) for the
degradation of Morin. Reaction conditions: 0.1 M HCO3-, 5 · 10
-5 M Morin, 0.015 M H2O2, pH 9.0, room temp., (—) 4 ·
10-5
M [MnII(bpy)2Cl2], (—) 2 · 10
-5 M [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3, (—) 4 · 10
-5 M Mn
IICl2 + 8 · 10
-5 M bpy.
Figure S3.2 Kinetic traces recorded at 400 nm for the degradation of p-nitrophenol. Reaction conditions: 0.1 M HCO3-,
5 · 10-5
M p-nitrophenol, pH 9.0, room temp., (A): 0.015 M H2O2, (—) 4 · 10-5
M [MnII(bpy)2Cl2], (—) 2 · 10
-5 M
[Mn2III/IV
(µ-O)2(bpy)4](ClO4)3, (—) 4 · 10-5
M MnIICl2 + 8 · 10
-5 M bpy. (B): 0.01M H2O2, (—) 5 · 10
-6 M [Mn
II(bpy)2Cl2], (—)
2.5 · 10-6
M [Mn2III/IV
(µ-O)2(bpy)4](ClO4)3, (—) 5 · 10-6
M MnIICl2 + 1 · 10
-5 M bpy and corresponding spectral changes
(inset B).
(A) (B)
(A) (B)
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
81
Figure S3.3 (A) Kinetic traces recorded at 400 nm for the degradation of p-nitrophenol at different pH values. (B) pH
dependence of the initial rate of p-nitrophenol bleaching (■) 4 · 10-5
M [MnII(bpy)2Cl2], (●) 2 · 10
-5 M [Mn2
III/IV(µ-
O)2(bpy)4](ClO4)3. Reaction conditions: 0.1 M HCO3-, 5 · 10
-5 M p-nitrophenol, 0.01 M H2O2, room temp..
Figure S3.4 Kinetic traces recorded for the oxidation of 5 · 10-5
M Orange II by (A) 4 · 10-5
M [MnII(bpy)2Cl2] and 2 · 10
-5
M [Mn2III/IV
(µ-O)2(bpy)4](ClO4)3 in the presence of 0.1 M HCO3- and 0.015 M H2O2, compared to the same reaction in
the presence of 0.2 M NaNO3. (B) 4 · 10-5
M [MnII(bpy)2Cl2] in the presence of 0.1 M NaCl. Reactions followed at 484
nm, pH 9.0, room temp., over 20 min.
(A) (B)
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
82
Figure S3.5 Observed rate constants measured at 484 nm as a function of [MnII(bpy)2(H2O)2]
2+ (A) and [Mn2
III/IV(µ-
O)2(bpy)4]3+
(B) concentration at different total carbonate concentrations. Reaction conditions: 0.01 M H2O2, 5 · 10-5
M
Orange II, pH 9.0, room temp., (■) 0.1 M HCO3-, (▲) 0.5 M HCO3
- .
Figure S3.6 X-band EPR spectra of 0.1 M bicarbonate containing solutions of H2O:tBuOH (1:1). Reaction conditions: pH
9.0, room temp., (A) 1 · 10-4
M [MnIIbpy2Cl2] and (B) 5 · 10
-5 [Mn2
III/IV(µ-O)2(bpy)4](ClO4)3 with (—) 0.01 M and (—) 0.04
M H2O2 immediately after mixing. EPR conditions: 8.98 GHz, 6 K, 1 mW microwave power, modulation amplitude 20
mT.
(A) (B)
(A) (B)
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
83
Figure S3.7 Spectral changes recorded for the reaction of 0.05 · 10-3
M [Mn2(µ-O)2(bpy)4](ClO4)3 with 5 · 10-3
M H2O2 in
0.1 M HCO3- solution. Spectra recorded in time intervals of 0.2 s for the first two seconds, pH 9.0, room temp..
Figure S3.8 Kinetic traces recorded at 400 nm for the in situ oxidation of Morin (A) and p-nitrophenol with 4 · 10-5
M
MnIICl2 and increasing amount of 2,2'-bipyridine added to the reaction mixture. Reaction conditions: 0.1 M
bicarbonate, between 0.5-4 % CH3CN, 5 · 10-5
M dye, 0.015 M H2O2, pH 9.0, room temp..
(A) (B)
3. Comparative study of a MnII-monomer and the corresponding oxo-bridged Mn2III/IV-dimer
84
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
85
4 METAL ION - CATALYZED OXIDATIVE DEGRADATION OF
ORANGE II BY PERACETIC ACID
4.1 GENERAL REMARK
The following chapter is based on the original publication: Mechanistic studies on the oxidative
degradation of Orange II by peracetic acid catalyzed by simple manganese(II) salts. Tuning the
lifetime of the catalyst, Sabine Rothbart, Erika Ember and Rudi van Eldik, New J. Chem. 2012, 36,
732-748.
.
4.2 INTRODUCTION
Persistent and non biodegradable organic waste still marks one of the main environmental
problems of our time. Significant ecological impact is caused by industrial dye waste since over 15
% of textile dyes are lost in waste water streams during the dyeing operation.[1] A recent review
emphasized the high cost of disposing the high volumes of dye effluent and that 128 tons of dyes
are released daily to the global environment.[2] About half of the global production of synthetic
colorants (700.000 t per year) is classified as aromatic azo compounds.[3] Textile dyes in general
are designed to resist chemical, biochemical and photochemical degradation. In aerobic processes,
azo compounds are known to be largely non-biodegradable, whereas under anaerobic conditions
they can be reduced to even more hazardous intermediates.[4] The increasing ecological awareness
stimulated an active field of scientific research dedicated to new and ecologically worthwhile
oxidation processes for the catalytic decomposition of environmental pollutants[5]and dyes[6] in
water. Among the various transition metals for catalytic oxidation, manganese is of particular
interest as one of the most efficient and environmentally benign elements. There are various
manganese complexes with different salen,[7] porphyrin,[8] tacn (1,4,7-triazacyclononane)[6e, 9] or
aromatic N-donor ligands[10] known to efficiently catalyze the oxidation of a wide range of
substrates. However, several limitations including elaborated synthetic methods, long reaction
times and substrate scope, have still to be resolved. Recent attention has therefore been focused
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
86
(4.1)
(4.2)
on metal complexes that use cheap and clean oxidants to bring about efficient oxidation under
mild reaction conditions.[11] Among these, H2O2 or peracetic acid are commonly used green
oxidants due to their eco-friendly nature. The only by-products formed are water and oxygen or
acetic acid, respectively.
In order to develop efficient and simple pre-catalysts for the oxidative degradation of
structurally different organic dyes by H2O2, our earlier work focused on the reactivity of simple
MnII salts under mild reaction conditions (Equation 4.1).[12]
It is known that simple MnII salts are able to form very reactive aquated intermediates in
aqueous solution.[13] Moreover, in slightly alkaline medium, the introduction of a hydroxy ligand
trans to a water ligand is expected to produce more labile HO-Mn-H2O species which are
considered to be of major importance for the catalytic activity.
Detailed mechanistic investigations revealed the in situ formation of percarbonate (HOOCO2-)
as a key molecular entity under the selected experimental conditions.[12, 14] Percarbonate can
coordinate to the MnII center and lead to the formation of a quasi-stable MnII-η2-percarbonate
complex, which subsequently undergoes heterolytic cleavage of the peroxide O-O bond to form
MnIV=O intermediates. The nature of the produced reactive intermediate was confirmed by low
temperature EPR measurements under catalytic reaction conditions.[12a, b] Although simple MnII
ions could efficiently catalyze the oxidative degradation of a large number of organic substrates
under mild reaction conditions, the use of rather high concentrations of pre-catalyst and
bicarbonate in solution was required.[12a, b]
In an attempt to overcome these limitations, we now report our findings for the Mn II catalyzed
degradation of Orange II with peracetic acid (PAA) under mild reaction conditions. PAA is formed
in an equilibrium reaction of hydrogen peroxide and acetic acid catalyzed by sulfuric acid
(Equation 4.2).[15]
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
87
Similar to hydrogen peroxide, PAA is often used as a safe and environmentally friendly oxidant,
since the released side-products, viz. water and acetic acid, are non-toxic unlike organic peroxides
or other oxidants used. In contrast to hydrogen peroxide, mechanistic investigations with
peracetic acid are very rare. Yet, a strong advantage of peracetic acid compared to hydrogen
peroxide is its lower tendency to undergo catalase-like reactions, which could result in rapid
oxidant decomposition under catalytic conditions. These attractive oxidizing properties make the
transition metal catalyzed activation of peracetic acid a worthwhile objective in chemical
oxygenation processes.
4.3 RESULTS AND DISCUSSION
4.3.1 Peracetic acid formation and its decomposition at higher pH
It is generally accepted that at a pH equal to the pKa of PAA (8.2)[16], the hydrolysis of PAA to
acetic acid and H2O2, and the spontaneous decomposition strongly complicate the H2O2-PAA
equilibrium, which causes poor reproducibility of the data.[16a, 17] For a better understanding of the
reactivity and in particular the stability of PAA as oxidant, we investigated its in situ formation and
hydrolysis by means of 13C-NMR spectroscopy (see Supporting Information). When H3C-13COOH
and H2O2 were mixed in water, the slow formation of a new 13C-NMR signal, attributed to H3C-
13C(O)OOH, was observed at 175.3 ppm (see Figure 4.1).
In order to accelerate the formation of 13C-labeled PAA, 0.1 M H2SO4 was added as catalyst.
13C-NMR experiments revealed that under the selected experimental conditions PAA reached its
maximum concentration after approximately 50 h and remained constant over longer periods of
time (see Figures S4.1 (A) and (B), Supporting Information). The observed rate constants at pH = 1
for the formation of PAA is (1.78 ± 0.05) · 10-5 s-1 at 25 °C. To study the hydrolysis of PAA to form
H2O2 and 13C-AcOH, a fully equilibrated solution of 13C-labeled PAA (formed from 2.5 M 13C-AcOH,
2.5 M H2O2 and 0.1 M H2SO4) was diluted in 0.1 M H2SO4 to 25 % of the initial concentration and
the set-in of the new equilibrium position was followed by 13C-NMR spectroscopy (see Figure S4.1
(C) and (D), Supporting Information).
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
88
(4.3)
Figure 4.1 13
C - NMR spectra recorded for the in situ formation of PAA in an aqueous solution of 2.5 M 13
C-AcOH and
2.5 M H2O2 in the presence of 0.1 M H2SO4 at 25 °C. Inset: The development of the H3C-13
C(O)OOH NMR signal at
175.3 ppm as a function of time.
From the determined equilibrium concentrations (x1 = [13C-PAA]eq, x2 = [13C-AcOH]eq) the
average equilibrium constant of K = 1.3 ± 0.1 at 25 °C was calculated for reaction 4.2 according to
Equation 4.3:
Even though there is only limited literature data available for the equilibrium constant and the
reported values differ as a function of temperature, viz. 3.7 (20 °C), 4.0 (30 °C), 4.3 (35 °C),[18a] 1.21
(0 °C), 1.20 (5 °C), 1.29 (35 °C),[18b] 2.10 (20 °C),[18c] our value fits in quite well with the data
published most recently.[18b, c]
An equilibrated solution of 13C-labeled PAA (produced from 2.5 M 13C-AcOH, 2.5 M H2O2 and
0.1 M H2SO4) was diluted to 50 % of the initial concentration with NaOH to a pH of approximately
10 and the decomposition reaction was monitored by 13C-NMR at 25 °C (Figure S4.2 (A) and (B),
Supporting Information). The observed rate constant is (2.3 ± 0.1) · 10-4 s-1 with a half-life of 3000
s. Consequently, it can be concluded that PAA can be considered to be stable in aqueous solution
under the experimental conditions selected for the studied degradation of Orange II, for which the
21220
1201
xx-)O(Hc
xO)(HcxK
][
][
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
89
half-life is at most 100 s. These findings are in agreement with reports in the literature on the
formation and stability of PAA based on redox titrations.[18]
4.3.2 General observations
In order to study the catalytic degradation of Orange II by PAA, a series of measurements was
performed at pH 9.5 and 25 °C. Figure 4.2 (A) shows the UV/Vis spectral changes that accompany
the catalytic degradation reaction.
Figure 4.2 (A) UV/Vis spectral changes observed during the 1 · 10-5
M MnII catalyzed oxidative degradation of 5 · 10
-5 M
Orange II by 0.01 M oxidant at pH 9.5 (0.05 M NaHCO3 buffer) and 25 °C (for the sake of clarity only every tenth
spectrum is shown). (B) Comparison of the absorbance vs. time traces at 484 nm for the oxidative degradation of 5 ·
10-5
M Orange II catalyzed by 1 · 10-5
M MnII and 0.01 M PAA or 0.01 M H2O2 at pH 9.5 (0.05 M NaHCO3 buffer) and 25
°C.
It is obvious that under these reaction conditions the bleaching reaction is not only limited to
the cleavage of the azo-linkage (484 nm), destruction of the more stable aromatic subunits (310
nm) also occurs.[19] Although PAA is a strong oxidant, its spontaneous non-catalyzed reaction with
the dye substrate is about 500 times slower compared to the reactivity in the presence of the Mn II
catalyst, as is apparent from the corresponding observed rate constants, viz. kobs cat. = 6.05 · 10-2 s-1
vs. kobs non-cat. = 1.28 · 10-4 s-1 at 25 °C (see Figure 4.2 (B)).
Consequently, the spontaneous reaction between PAA and Orange II does not significantly
contribute to the determined degradation rates for the MnII catalyzed reactions. Addition of an
excess of EDTA to the non-catalytic reaction mixture only slightly reduced the spontaneous non-
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
90
catalyzed reaction with the dye substrate due to scavenge of possible trace metals (as for instance
adventitious Mn) in the stock solutions (for comparison see Figure S4.3, Supporting Information).
When the results are compared to the MnII catalyzed oxidative degradation of Orange II by H2O2 as
oxidant under identical experimental conditions (see Figure 4.2), it is obvious that the rate
constant in the presence of PAA as oxidant is several orders of magnitude higher (kobs = 6.05 · 10-2
s-1 for PAA, and kobs = 7.92 · 10-4 s-1 for H2O2[12a] at 25 °C). This surprising result clearly suggests
fundamental differences in the activation of the two peroxides by MnII ions. Since the
experimental data did not allow the determination of pseudo-first-order rate constants under all
experimental conditions, the initial rate method was used in most cases to further quantify the
obtained results. If not stated otherwise, the MnII salt used in this study was MnIICl2 · 2H2O, since
no influence of the anion on the reaction course was observed, as is obvious from a comparison of
the Orange II degradation rate for different MnII salt catalysts (see Figure S4.4, Supporting
Information).
The possible contribution of free radicals to the remarkable efficiency of the MnII catalyzed
degradation of Orange II by PAA could be excluded by the use of tBuOH and BHT as radical
scavenger. Although hydroxyl radicals might be present during the PAA induced reaction, they do
not participate in the studied oxidation process, since in the presence of tBuOH or BHT no negative
effect on the reaction course was observed (see Figure S4.5, Supporting Information). In further
control experiments it could be confirmed that neither acetate (since commercial PAA is a
equilibrated mixture of H2O2 and acetic acid), nor the counter ion of the MnII salts affects the
performance of the catalyzed dye degradation (see Figure S4.4, Supporting Information). The
influence of ionic strength was found to be negligible over a wide concentration range (see Figure
S4.6, Supporting Information).
4.3.3 MnII + PAA – Intermediates formed in the absence of substrate
4.3.3.1 UV/Vis spectroscopy
In order to follow the speciation of Mn during the catalytic cycle, we analyzed the catalytic
solution in detail in the absence of added substrate by means of UV/Vis and EPR spectroscopy.
From the observed spectral changes and the corresponding EPR spectra recorded at different time
intervals some preliminary conclusions can be drawn. As shown in Figure 4.3, the reaction of
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
91
aquated MnII proceeds in two different phases. In the first phase of the reaction (Figure 4.3 (A),
inset) a new species with weak absorbance bands at approximately 405 and 470 nm is formed.
This intermediate persists for several seconds. Similar spectra with a band at 470 nm are often
attributed to MnIII species.[20] However, absorbance bands in the region of approximately 450 nm
are also considered to be a result of an oxo to MnIV charge transfer transition of a possible MnIV=O
intermediate.[12b, c, 21] In the second phase of the reaction, the first intermediate is rapidly
converted to permanganate (band at 525 nm) and colloidal MnIVO2 (band at 350 nm) (Figure 4.3
(B)).
Figure 4.3 UV/Vis spectra recorded for the reaction of 1 · 10-4
M MnII with 2.5 · 10
-2 M PAA in a 0.05 M NaHCO3
containing buffer solution at pH 9.5 and 25 °C. (A) Reactions studied over 300 s. Inset: first 25 s. (B) Biphasic behavior
of the kinetic traces at 350 nm and 525 nm.
Due to the broad absorbance of colloidal MnIVO2 over the whole spectral range (with
characteristic bands at 312 and 350 nm[22]), the exact quantification of the generated products
became difficult. However, the generation of the pink permanganate ion is evident from the five
characteristic absorbance bands in the range from 500-570 nm. In general, the formation of
permanganate is known to also involve short-lived MnVI and MnV intermediates, which tend to
undergo disproportionation to permanganate and MnIVO2 in the absence of a large excess of OH-
.[23, 24] This stands in direct contrast to the behavior of the MnII/H2O2/HCO3- system under
comparable conditions, in which the oxidation state of manganese did not exceed that of Mn IV=O
and no permanganate was formed.[12]
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
92
The observation of sigmoid-shaped kinetic traces as shown in Figure 4.3 (B), which are marked
by the beginning of the second reaction phase, i.e. formation of colloidal MnO2 (which partially
precipitates at higher [Mn] during the course of the reaction in phase II) and permanganate,
remind of typical autocatalytic pathways that are common features of chemical reactions involving
higher oxidation states of manganese.[25, 26, 27] It was found that during the initial phase the parallel
existence of MnVIIO4- and MnII, as well as MnIII and MnIV[27] during the MnII catalyzed
decomposition reaction of PAA, resulted in the formation of permanganate and Mn IVO2 depending
on the initial reactant concentration. In neutral to basic media, the autocatalytic contribution
occurring in oxidation reactions by permanganate are generally attributed to the existence of Mn IV
species as autocatalytic reaction product.[25, 28, 29]
Catalyst and oxidant influence on the biphasic reaction behavior. In the absence of substrate the
studied reaction is very sensitive to variation of the reactant concentration. Since the use of
stopped-flow techniques is hardly possible due to the formation of MnIVO2 precipitates and the
way the catalytic mixture is prepared via a pH jump, reliable results can only be obtained with in
situ UV/Vis spectroscopy. Nevertheless, upon variation of the metal to catalyst ratio some trends
in the biphasic reaction behavior could be observed. A comparison of the spectral changes for
different [MnII] at constant [PAA] (Figure 4.4), clearly shows that the length of the first reaction
phase at pH 9.5 strongly depends on the [MnII] present in solution.
Figure 4.4 Kinetic traces recorded at 350 nm for the reaction of 0.01 M PAA with different [MnII]. Inset: 1 · 10
-4 M Mn
II
with different [PAA]. Reaction conditions: 0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
93
If the pre-catalyst concentration is kept constant (Figure 4.4, inset), an increase in the [PAA]
leads only to minor changes in the biphasic behavior, but the amount of the species formed during
phase I increased, as indicated by the higher absorbance. Higher [MnII] on the other hand resulted
in a much shorter reaction time during phase I (see Figure 4.4). A higher overall [Mn] probably
causes a faster consumption or decomposition of the equilibrium content of H2O2, omnipresent in
peracetic acid solutions, which as a consequence shortens the induction period. The reducing
character of H2O2 towards high valence metal species also represents another likely explanation
for the biphasic reaction behavior. As long as H2O2 is present in the reaction mixture it maintains a
steady-state concentration of the lower valence Mn species, i.e. MnIII and MnIV, by reduction of
the rapidly formed higher valence manganese intermediates. Once all H2O2 has depleted, the lack
of the back reaction causes immediate accumulation of colloidal MnIVO2 and permanganate as the
only stable high valence reaction products. This suggestion prompted us to further investigate the
role of hydrogen peroxide in the catalytic system.
H2O2 as reductant. Commercially available PAA is not a pure peroxide, but an equilibrated mixture
of acetic acid (45 %), hydrogen peroxide (6 %) and water with sulfuric acid as catalyst. Thus, PAA
solutions always contain a significant amount of hydrogen peroxide. The effect of an extra amount
of hydrogen peroxide is also apparent in the UV/Vis spectral changes observed in the absence of a
substrate. If MnII and PAA react in the presence of additional 2 · 10-3 M of H2O2, the first phase of
the reaction is extended, while the formation of colloidal MnIVO2 and permanganate is delayed for
50 to 100 seconds (see Figure 4.5).
This behavior can be interpreted in terms of a continuous fast reduction of the rapidly in situ
formed high valence manganese species (MnVIIO4- and colloidal MnIVO2) as long as H2O2 is present
in the catalytic reaction mixture. Once the available hydrogen peroxide has been used, the
concentration of the high valence Mn species increases. However, since MnVIO42- and MnVO4
3- are
very unstable below pH 14, the known disproportionation chemistry will apply and lead to the
accumulation of colloidal MnIVO2 and MnVIIO4- as final products (Scheme 4.1).
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
94
Figure 4.5 Absorbance changes at 350 nm for the reaction of 2 · 10-5
M MnII with 0.01 M PAA (0.05 M NaHCO3 buffer,
pH 9.5, 25 °C) without and with different amounts of added H2O2.
Scheme 4.1 Reactions suggested to account for the disproportionation of high valence Mn-oxo anions.
Transiently formed colloidal manganese dioxide is also known to be reduced by hydrogen
peroxide.[30] However, depending on the experimental conditions, either partial or total reduction
of MnIVO2 occurs. Whereas under acidic conditions total reduction of MnIVO2 to MnII is possible, it
is reported that under neutral to slightly basic conditions only partial reduction of the colloid
occurs.[31] It is generally accepted that these processes proceed via non-stoichiometric hydrated
MnIVO2 with interstitial metal ions of lower valence like MnIII, leading to an overall increased
amount of MnIII present in the catalytic mixture.[31, 32, 33] According to studies by Perez-Benito and
co-workers, partial reduction of MnIVO2 only takes place on the surface of the colloid.[31] The result
is a monomolecular MnII oxide layer on the surface of a mixed MnIVO2-MnIIO colloid. Both the MnIV
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
95
and MnII oxides in a mixed MnIVO2-MnIIO colloid can react to form MnIII oxide according to Scheme
4.2, which may also account for the formation of the absorbance shoulder at 470 nm during the
initial stage of the reaction (see Figure 4.3 (A), inset).[31]
Scheme 4.2 Possible reduction pathway of MnIV
O2 with H2O2.
In order study the role of hydrogen peroxide in the catalytic cycle, we investigated the
reduction of the reaction products MnVIIO4- and colloidal MnIVO2 with H2O2 as a function of pH
under reaction conditions comparable to those for the catalytic degradation reaction. Upon
addition of H2O2 to a colloidal MnIVO2 containing solution at pH 9.5, a fast decrease in the broad
and characteristic absorbance between 300 and 400 nm was observed (see Figure S4.7 (A),
Supporting Information). As expected, the reduction of colloidal MnIVO2 was not complete and the
extent to which the reduction proceeds depends strongly on the pH (see Figure S4.7 (B),
Supprorting Information). This is due to the behavior described above and the formation of a
precipitate at higher pH, which results in incomplete reduction. In contrast, the reaction of H2O2
with MnVIIO4- was found to be accelerated with increasing pH (see Figure S4.8, Supporting
Information), which presumably originates from the increase in the formation of HOO- on
approaching the pKa of hydrogen peroxide (pKa = 11.67)[34] and the thereby increased reducing
ability. Although oxidation reactions by permanganate, often autocatalytic, have been studied for
quite a long time,[35, 36] they are not fully understood. Nevertheless, the results clearly indicate the
occurrence of a reaction of the in situ formed intermediates with hydrogen peroxide, which most
likely represents a parallel degradation reaction in the presence of a dye.
Role of pH. Another factor that drastically influences the biphasic reaction behavior of MnII with
PAA, is the pH of the solution. The maximum rate of permanganate and MnIVO2 formation was
found at a pH around 8.5, where no induction period is observed. Surprisingly, a pH of 8.5 is not
where the best reactivity is observed, which implies that another oxidizing intermediate than
permanganate is responsible for the excellent reactivity at pH ≈ 9.5. Increasing the pH from 8.8 to
10.1 causes a drastic delay in the formation of the high valence oxidation products (see Figure
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
96
S4.9, Supporting Information). This could, however, also be due to the enhanced reducing
properties of H2O2 at higher pH since it cumulatively dissociates to form HOO- (pKa = 11.67)[34], i.e.
another potential reducing agent for higher valence manganese species such as MnVIIO4-.
Influence of the total carbonate concentration. Earlier studies on the MnII ion catalyzed
degradation of Orange II by H2O2 revealed the in situ formation of percarbonate (HOOCO2-) from
bicarbonate and hydrogen peroxide as a key molecular entity under the selected reaction
conditions.[12] Furthermore, the role of bicarbonate as possible ligand that facilitates the
heterolytic O-O-bond scission and thereby MnIV=O formation was discussed.[12] Since PAA is
strongly acidic, and HCO3-/CO3
2- was used as buffer system, the formation of possible
intermediates was also investigated as a function of the total carbonate concentration.
Control experiments already excluded any influence of ionic strength (see Figure S4.6) so that
under identical experimental conditions the difference in the formation of intermediates is directly
attributed to the presence of bicarbonate and carbonate. Figure 4.6 shows the UV/Vis spectral
changes that accompany the reaction of MnII with PAA at two different [total carbonate], viz. 0.05
M (Figure 4.6 (A) and 0.4 M (Figure 4.6 (B)).
Figure 4.6 UV/Vis spectral changes that accompany the reaction of 2 · 10-5
M MnII with 0.01 M PAA at pH 9.5 and 25 °C
in (A) 0.05 M NaHCO3 buffer and (B) 0.4 M NaHCO3 buffer. Inset: UV/Vis spectra recorded at t = 300 s for the reaction
as a function of [total carbonate].
On following the spectral changes at 350 nm (colloidal MnIVO2) and 525 nm (MnVIIO4-) (see
Figure S4.10, Supporting Information) under variation of the [total carbonate], only minor changes
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
97
in the biphasic behavior were observed. Yet, it is evident that the distribution of the in situ formed
high-valance manganese intermediates changes drastically as a function of [HCO3-/CO3
2-]. The
inset in Figure 4.6 (B) presents UV/Vis spectra recorded at a reaction time of 300 s for a [total
carbonate] range from 0.025 to 0.4 M. Whereas colloidal MnIVO2 is favored at lower buffer
concentration, the formation of MnVIIO4- is enhanced at higher [total carbonate]. A reasonable
explanation might be that HCO3-/CO3
2- can coordinate to the metal center and thereby facilitate
either a further oxidation by PAA or open up alternative reaction pathways leading to the
enhanced formation of permanganate. The outstanding role of bicarbonate ligands in the
beneficial manipulation of the MnIII/MnII redox couple by the formation of highly reactive Mn-
bicarbonate complexes has been discussed before.[12a, 37] The present results provide strong
evidence that a change in [HCO3-/CO3
2-] should result in drastic changes in the degradation
reaction of the dye (see further discussion).
4.3.3.2 EPR spectroscopy
The above described information on how to tune the biphasic behavior of the reaction of MnII
with PAA, enabled us to select appropriate reaction conditions that allowed sufficient time to
investigate the nature of the formed intermediates in more detail. In an attempt to further
characterize the possible reactive intermediates, samples were taken directly from the
concomitant UV/Vis measurements at different reaction times (marked by the arrows in Figure 4.7
(A)), immediately frozen to quench the reaction and analyzed by perpendicular EPR spectroscopy
at 10 K.
The initial MnII catalyst containing buffer solution without any oxidant present, showed the
typical six-line pattern of MnII (I = 5/2, S = 5/2) centered at g ≈ 2 (see Figure 4.7 (B), first
spectrum). Formation of hydroxo species in basic medium and complex-formation with
bicarbonate further reduces the symmetry of the ligand field around the MnII ion and thereby
causes a weaker intensity of the six-line EPR signal compared to the more symmetric fully aquated
MnII ion as described before.[12b] Figure 4.7 (B) shows the successive development of the X-band
EPR spectrum upon addition of MnII to a PAA/H2O2 containing buffer solution at pH 9.8. The
decrease in the overall spectral features compared to the spectrum in the absence of oxidant
indicates that a small amount of the Mn catalyst is in an EPR-silent form, which according to the
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
98
UV/Vis spectral changes is most likely a MnIII species. However, neither Mn2III/III nor Mn2
IV/IV
dimeric species are expected to give an EPR spectrum.[38]
Figure 4.7 (A) UV/Vis spectra recorded during the reaction of 1 · 10-4
M MnII with 2.5 · 10
-2 M PAA and 5 · 10
-3 M H2O2
in 0.05 M NaHCO3. The first phase of the reaction was extended by the addition of 5 · 10-3
M H2O2. Inset: kinetic traces
at 350 and 525 nm. Arrows mark the time at which different samples were taken for the EPR spectra in (B). Reaction
conditions: 1 · 10-4
M MnII with 2.5 · 10
-2 M PAA and 5 · 10
-3 M H2O2 at pH 9.8 and 25 °C. (B) X-band EPR spectra
recorded at 10 K for 1 · 10-4
M MnII in carbonate buffer and at different time intervals after the addition of 2.5 · 10
-2 M
PAA and 5 · 10-3
M H2O2 in 0.05 M NaHCO3. EPR conditions: 8.95 GHz, 10 K, 1 mW microwave power, modulation
amplitude 400 mT.
In the first spectrum recorded 20 s after mixing, still some MnII is found. The lack of the typical
well resolved MnII sextet at g ≈ 2 is attributed to the enhanced formation of manganese-hydroxo
species, which results in a broader and weaker transition of the MnII signal at g ≈ 2. Moreover, two
further signals at g ≈ 2 and g ≈ 4 can be observed. On one hand the sharp g ≈ 2 transition might be
due to a very low [MnVIO42-], since the spectrum slightly resembles the one obtained for a readily
prepared, dark green MnVIO42- solution (d1, S = 1/2) (see Figure S4.11, Supporting Information).
However, the lack of the characteristic absorbance band at 630 nm for MnVIO42- in the
corresponding UV/Vis spectra of the intermediate during the initial phase implies that only a
minor [MnVIO42-] is present due to the equilibrium concentration of H2O2 omnipresent in PAA. On
the other hand, the characteristic signals at g ≈ 2 and g ≈ 4 can also originate from a high-spin MnIV
(S = 3/2) species in an octahedral environment having D < hν, which also has a higher amplitude
feature at g ≈ 2.[39, 40] In general, EPR spectra of d3 MnIV ions in an axial field (E/D = 0) are often
difficult to interpret, since they strongly depend on the magnitude of the zero-field splitting
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
99
parameters.[41] A large axial zero-field splitting parameter D accounts for a spectrum dominated by
the g ≈ 4 signal, as for complexes with hard oxygen-rich catecholate[42a] and sorbitalate[42b] ligands.
If D is small, the g ≈ 2 signal dominates with relatively weak low field signals. This is seen for
example for sulfur-containing thiohydroxamate[41b] and dithiocarbamate[43] manganese(IV)
complexes. The assignment of a MnIV intermediate also correlates with the observed UV/Vis
spectra observed during the initial reaction stage, in which the absorbance bands in the region of
450 nm can be interpreted as a result of an oxo to MnIV LMCT band.[12b, c, 21] Consequently, it is
more likely to be the result of an in situ formed MnIV species showing g ≈ 2 and g ≈ 4 signals during
the initial phase of the reaction of MnII with PAA.
From the subsequent EPR spectra (Figure 4.7 (B)) taken with a delay of 75 and 120 s for the
same reaction mixture, it is apparent that the concentration of the in situ formed MnIV
intermediate increases during the time of the first reaction phase. With the beginning of the
second reaction phase (130 s) this MnIV intermediate is not detected anymore, whereas the
corresponding UV/Vis spectra show the formation of MnVIIO4- and precipitation of colloidal
MnIVO2. Although the UV/Vis and EPR results indicate the existence of monomeric MnIII and
monomeric MnIV species during the initial phase of the reaction, no evidence for the formation of
bis µ-oxo bridged MnIII/MnIV dimers was found. For such a strongly coupled dimer (MnIII(µ-
O)2MnIV) a very characteristic 16-line EPR signal at g ≈ 2 would be expected.[44] The EPR spectra
recorded in the presence of tBuOH as radical scavenger show a similar behavior, indicating that
free radical processes do neither interfere in the catalytic dye degradation nor in the reaction of
the MnII catalyst with PAA in the absence of substrate (see Figure S4.12, Supporting Information).
Very similar results were obtained when the reaction was carried out at a higher [total carbonate]
(see Figure S4.13, Supporting Information).
For a better understanding of the influence of the equilibrium content of H2O2 on the in situ
formed intermediates, EPR spectroscopic experiments were also performed in the presence of an
excess of H2O2. As is evident from the UV/Vis spectral changes recorded for the reaction between
MnII and PAA in the presence of a large excess of H2O2 (see Figure 4.8 (A)), formation of MnVIIO4-
and colloidal MnIVO2 does not occur on a catalytically relevant time scale. Instead, the
disappearance of the characteristic MnII sextet at g ≈ 2 and the concomitant formation of the MnIV
species are delayed for several seconds compared to the experiments in the absence of an excess
H2O2. Moreover, the intensity of the above described MnIV signals at g ≈ 2 and g ≈ 4 is decreased,
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
100
implying that less MnIV is formed (see Figure 4.8 (B)). Consequently, there is a reductive influence
of the excess of H2O2 on the in situ formed MnIV intermediate, which shifts the complex
oxidation/reduction equilibria in solution to the side of the MnII pre-catalyst and thereby causes
delayed formation of higher valence-species, such as the MnIV intermediate.
Figure 4.8 (A) UV/Vis spectra recorded for the reaction of 1 · 10-4
M MnII with 0.01 M PAA and 0.05 M H2O2 in a 0.05 M
NaHCO3 containing buffer solution at pH 9.8 and 25 °C. Inset: kinetic traces at 350 nm and 525 nm, arrows mark the
time at which samples were taken for EPR spectroscopic measurements. (B) X-band EPR spectra recorded at 10 K for 1
· 10-4
M MnII in carbonate buffer and at different time intervals after the addition of 0.01 M PAA and 5 · 10
-2 M H2O2 in
0.05 M NaHCO3. EPR condidions: 8.95 GHz, 10 K, 1 mW microwave power, modulation amplitude 400 mT.
4.3.4 Comparison of reactivity of different high valent oxo-manganese species
with Orange II
In the absence of PAA oxidant. Since the reaction of peracetic acid with aquated manganese ions
in slightly basic solution leads to the rapid in situ generation of several high-valence manganese-
oxo intermediates, we tested the different species separately for their degradation ability towards
the dye substrate in the absence of an oxidant. Following the spectral changes during the
stoichiometric (1:1) reaction of Orange II with the freshly prepared O3MnV(OH)2-, MnVIO42- and
MnVIIO4- species (Figure 4.9) under the same experimental conditions, it is obvious that the activity
of the different oxo complexes decreases drastically (Figure 4.9, inset) with decreasing oxidation
state of the metal center.
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
101
Figure 4.9 UV/Vis spectra of 2 · 10-5
M high valence manganese-oxo anions. Inset: comparison of the spectral changes
at 484 nm during the stoichiometric reaction of 2 · 10-5
M Orange II with 2 · 10-5
M O3MnV(OH)
2-(a), Mn
VIO4
2- (b) and
MnVII
O4- (c).
[45] Reaction conditions: 0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
This is not surprising since Pode and Waters already reported that the order of reactivity of
high valence Mn-oxo anions towards organic substrates decreases markedly along the series
MnVIIO4- > MnVIO4
2- > MnVO43-.[46] Although permanganate is a strong oxidant, its full oxidative
potential in the studied reaction with Orange II develops at higher pH values as can be seen from
the pH dependence in (Figure S4.14, Supporting Information). If the degradation reaction is
performed with manganate(VI) and hypomanganate(V) under identical experimental conditions, it
becomes clear that permanganate is the stronger oxidant. Without any primary oxidant present,
one equivalent of Orange II is oxidized by two equivalents of MnVIIO4- (see the remaining
absorbance of Orange II at 484 nm, Figure 4.9, inset).
Moreover, during the first 50 s of the oxidation by manganate(VI), a build-up of the
absorbance at 484 nm can be observed. This is due to the background disproportionation reaction
during which permanganate is formed. At the same time the Orange II degradation starts as soon
as sufficient permanganate is formed in the disproportionation process. The time scale of these
absorbance changes fits well with the disproportionation of manganate(VI) in the absence of
substrate under identical experimental conditions (see Figure S4.15 (A), Supporting Information).
Moreover, a comparison of the pH dependence of the disproportionation of manganate(VI) in the
stoichiometric reaction with Orange II clearly shows that no reaction between MnVIO42- and
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
102
Orange II occurs under conditions where manganate(VI) disproportionation is avoided, i.e. at
higher pH (see Figure S4.15 (B), Supporting Information).
Upon addition of hypomanganate(V), no absorbance decrease was observed, probably
because of the high instability of O3MnV(OH)2- in non-basic media (pKa = 13.7[23]). Control
experiments at higher pH confirmed this order of reactivity. If the pH is high enough to ensure the
stability of hypomanganate(V) and manganate(VI) for several seconds, no reaction with the dye
substrate occurred (see Figure S4.16, Supporting Information). Consequently, it is assumed that
despite their possible transient presence, hypomanganate(V) and manganate(VI) play a negligible
role as active species in the studied MnII catalyzed dye degradation of Orange II with PAA.
As expected, colloidal MnIVO2 showed no reaction with the substrate (Figure S4.17, Supporting
Information). Yet, oxo-manganese(IV) compounds are known to be versatile homogeneous
oxygenation species,[12, 47] but in the absence of a stabilizing ligand they rapidly agglomerate to
give insoluble MnIVO2 and thus lose their oxygenation ability. Hence, the in situ generation of
transient, soluble MnIV-oxo species was investigated by stoichiometric reduction of MnVIIO4- with
sulfite and hydrogen peroxide, and immediate addition of dye substrate as soon as the MnVIIO4-
band vanished. Figure S4.18 (A) (Supporting Information) presents the UV/Vis spectra of MnVIIO4-
after reduction with sulfite, and after immediate addition of Orange II to the obtained Mn IV
species. Although MnVIIO4- is no longer present after reduction, a decrease in the characteristic
absorbance of Orange II is observed in the reaction with substrate (see Figure S4.18 (B),
Supporting Information). The rather moderate reactivity of the in situ formed transient MnIV
species towards Orange II is attributed to the rapid formation of a MnIVO2 precipitate. Yet, it
exhibits oxygenating capabilities and can represent a catalytically active species.
In the presence of PAA oxidant. The reactivity of freshly prepared Mn-oxo species towards
Orange II was tested in the presence of PAA. Surprisingly, the dye degradation catalyzed by the
same amount of O3MnV(OH)2-, MnVIO42-, MnVIIO4
- and MnII under identical experimental
conditions, showed no significant difference within the experimental error limits (see Figure 4.10).
This strongly suggests that the rate-limiting step of the catalytic dye degradation process does not
involve the underlying disproportionation chemistry of high valence Mn-oxo anions. Moreover,
the latter species must be rapidly reduced by the equilibrium amount of H2O2 present in PAA.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
103
Figure 4.10 Comparison of the absorbance changes observed at 484 nm during the 1 · 10-5
M Mn catalyzed
degradation of 5 · 10-5
M Orange II by 5 · 10-3
M PAA with different Mn complexes as catalyst. Reaction conditions:
0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
The catalytic activity of colloidal MnIVO2 as starting material in the presence of PAA was found
to be significantly lower than that of MnVIIO4-, MnVIO4
2- and MnII (Figure 4.10). Hence, the surface
activation of the oxidant by coordination to colloidal MnIVO2 particles is unlikely to be responsible
for the outstanding dye degradation reactivity of the MnII/PAA system.
4.3.5 Reactivity of different in situ formed intermediates towards Orange II
The careful selection of reaction conditions enabled us to perform comparative measurements
on the catalytic performance of different in situ produced reactive intermediates. In the
experimental set-up the addition of the substrate Orange II to the reaction mixture of MnII and
PAA was carried out at different points of time during the reaction course under catalytic
conditions. A comparison of the observed initial rates for the degradation of Orange II yielded
some helpful information about the redox states through which the catalytic system cycled during
the course of the reaction. Figure 4.11 (A) presents the UV/Vis spectral changes observed for the
reaction of MnII with PAA in the absence of substrate, and the corresponding kinetic traces
recorded at 350 nm (colloidal MnIVO2) and 525 nm (MnVIIO4-) at pH 9.6 (inset in Figure 4.11 (A))
with the earlier mentioned biphasic behavior. If Orange II is added after different time intervals,
the initial oxidation rates change significantly. It is obvious that the catalytic mixture containing
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
104
different Mn-oxo intermediates formed during the reaction of MnII with PAA, shows different
reactivity (Figures 4.11 (B) and 4.12).
Figure 4.11 (A) Spectral changes recorded for the reaction of 1 · 10-5
M MnII with 0.01 M PAA. Inset: kinetic traces at
350 and 525 nm. (B) Identical reaction conditions as in (A) but with different delay times for substrate addition (5 · 10-5
M Orange II) followed at the maximum absorbance of Orange II at 484 nm. Reaction conditions: 0.05 M NaHCO3
buffer, pH 9.6 and 25 °C.
A plot of the initial rate of dye degradation vs. the delay time for the addition of Orange II is
shown in Figure 4.12 (■) along with the corresponding absorbance changes at 350 and 525 nm in
the inset of Figure 4.12. At first an increase in reactivity is observed up to the point shortly before
the end of the first phase. As a matter of fact, about 185 s after the start of the reaction (under the
selected experimental conditions), the initial degradation activity is three times higher than at the
beginning. As soon as colloidal MnIVO2 and MnVIIO4- accumulate in the reaction mixture, the
catalytic performance almost vanishes, although a large excess of oxidant is still present in
solution. Thus, it can be assumed that the most reactive intermediates are formed during the first
phase of the reaction, more precisely near to its end, which in turn coincides with the amount of
hydrogen peroxide present in the reaction mixture, as discussed above. If the experiment is
repeated in the presence of a small aliquot of additional H2O2 right from the beginning of the
reaction, the initial phase is extended, but the above described behavior remains to be identical
(Figure 4.12 ▲).
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
105
Figure 4.12 Comparison of the initial degradation rates determined at 484 nm for the 1 · 10-5
M MnII catalyzed
reaction with 0.01 M PAA with different delay times for the addition of 5 · 10-5
M Orange II and different amounts of
H2O2. Inset shows the corresponding absorbance changes at 350 and 525 nm in the absence of substrate. Reaction
conditions: 0.05 M NaHCO3 buffer, pH 9.6 and 25 °C.
If H2O2 is added shortly before the beginning of the second phase (Figure 4.12 ●), the further
oxidation of the catalyst is again avoided and delayed such that the reactivity increases at the end
of the first phase (t ≈ 880 s) and decreases with the occurrence of permanganate and Mn IVO2. This
is a strong indication for the reductive influence of H2O2 during the first phase of the reaction.
In order to elucidate what causes the drop in reactivity as well as the insufficient substrate
bleaching in the second phase when MnVIIO4- and colloidal MnIVO2 are formed, it was attempted to
selectively reduce the latter species by addition of H2O2. In the presence of substrate, the sluggish
activity of the catalytic system at the stage where MnVIIO4- and colloidal MnIVO2 are formed could
be partially reactivated by the addition of a small aliquot of H2O2 (see Figure S4.19, Supporting
Information). In addition, selective reduction of most of the intermediates present in the second
phase (MnIVO2 or MnVIIO4-) is also achieved in the absence of substrate (Figure 4.13 (A)).
A comparison of the catalytic reactivity at different points of time of the reaction of MnII with
PAA revealed a drastic decrease as soon as MnVIIO4- and colloidal MnIVO2 are formed during the
second phase (Figures 4.13 (A) and (B), 1 and 2). However, if under identical reaction conditions
the Orange II degradation is performed after the partial selective reduction of these intermediates
with H2O2, the catalytic activity of the system is almost fully restored (Figure 4.13 (B), 3).
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
106
Figure 4.13 (A) Absorbance changes at 350 and 525 nm for the reaction of 1 · 10-5
M MnII with 0.01 M PAA, followed
by reduction with 2 · 10-3
M H2O2 after 250 s. Arrows mark the addition of 5 · 10-5
M Orange II for the corresponding
reactivity test in (B). (B) Absorbance changes at 484 nm for the different in situ formed intermediates in phase 1 at
150 s, phase 2 at 250 s and after reduction with 2 · 10-3
M H2O2 at 330 s. Reaction conditions: 0.05 M NaHCO3 buffer,
pH 9.60 and 25 °C.
In order to study the influence of H2O2 on the in situ formed intermediates in the second phase
of the reaction, we investigated the UV/Vis spectral changes that accompany the reduction. To the
reaction mixture containing 2 · 10-5 M MnII and 0.01 M PAA, 2 · 10-3 M H2O2 (pH 9.6, 25 °C) was
added in the second phase of the reaction after 50 s where the formation of colloidal MnIVO2 and
MnVIIO4- was complete. The UV/Vis spectra in Figure 4.14 show the rapid and complete reduction
of permanganate as well as the reduction of approximately 50 % of the colloidal MnIVO2 by H2O2
under the selected experimental conditions. The incomplete reduction of MnIVO2 by H2O2 is due to
the already discussed pH effects and causes a loss in catalytically active Mn and thus a slight
decrease in reactivity.
Nevertheless, the result is surprising since one would expect the highest reactivity at the point
of complete formation of MnVIIO4-, the strongest oxidant among the high valence Mn-oxo anions.
Furthermore, it is obvious from the corresponding reactivity test that the catalytic activity of the
system is restored after the reduction by H2O2. Consequently, the reactivity drop at this stage is
not caused by the depletion of the PAA oxidant, but by the over-oxidation of MnII to MnVIIO4- due
to the depletion of H2O2 as reducing agent present in stock solutions of PAA. The oxidation of the
catalyst to the MnVIIO4- state negatively affects the catalytic performance. This also implies that
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
107
the in situ formed MnVIIO4- is not the most reactive species in the oxidative degradation of Orange
II, and that for efficient catalysis the oxidation state of manganese should not exceed that of MnIV.
Figure 4.14 UV/Vis spectrum recorded after 2 · 10-3
M H2O2 was added to the in situ formed intermediates during the
reaction of 2 · 10-5
M MnII with 0.01 M PAA before and after the addition. Reaction conditions: 0.05 M NaHCO3 buffer,
pH 9.5 and 25 °C. Inset shows the corresponding absorbance changes at 350 and 525 nm.
4.3.6 MnII catalyzed degradation of Orange II by PAA
pH dependence. The oxidative degradation of Orange II by PAA was studied in the pH range 7.0 to
12.6 (± 0.1) to investigate the catalytic activity of the in situ formed high-valence Mn intermediates
as a function of pH. Figure 4.15 reports the dependence of the initial rate of the degradation
reaction on the pH of the solution.
The results show a bell shaped profile with a maximum reactivity at a pH between 9 and 10. As
could be expected, very low reaction rates are observed at pH 7.5 where peracetate is present in
its protonated form and can only weakly interact with MnII. The first inflection at pH ≈ 8.5
coincides rather well with the pKa of PAA, viz. 8.2,[16] which suggests the necessity of PAA to form
CH3C(O)OO- to reach a high catalytic activity. In addition, it may be related to the faster reduction
of permanganate, rapidly formed in the presence of PAA, by H2O2 to form the actual catalytic
species, viz. MnIV=O, which shows an apparent pK value around 8.8 (see Figure S4.8 (B)). Upon
increasing the pH above pH 9.5, the decrease in catalytic activity can be ascribed to the formation
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
108
of catalytically inactive Mn intermediates such as Mn(OH)2, MnCO3 and insoluble MnIVO2.
Alternatively, it may be due to the disproportionation of manganate(VI) that shows a similar
decrease in reactivity on increasing the pH with an apparent pK value around 10.5.
Figure 4.15 pH dependence of the initial rate for the 1 · 10-5
M MnII catalyzed oxidative degradation of 5 · 10
-5 M
Orange II by 5 · 10-3
M PAA in 0.05 M NaHCO3 at 25 °C.
Variation of the [pre-catalyst] and [PAA]. To obtain more information on the underlying reaction
mechanism, the initial rate for the degradation of Orange II was studied under optimal conditions
(pH 9.5 and 25 °C) as a function of the [PAA] for two pre-catalyst concentrations. For this purpose
the initial [PAA] was varied in the range of 0.5 – 50 · 10-3 M. The observed rate profile (see Figure
4.16 (A)) shows typical saturation behavior and reaches a limiting rate at high oxidant
concentration. The saturation effect at higher [PAA] implies the formation of a pre-equilibrium
between the oxidant and the MnII precursor, prior to the rate-limiting step of the oxidative
degradation of Orange II. This is in line with a mechanism involving coordination of the
deprotonated peroxide to the MnII center and transformation of the thereby formed peroxo
complex into a catalytically relevant species.
At a constant [PAA], a good linear dependence of the initial rate on the initial pre-catalyst
concentration was observed (see Figure 4.16 (B)). The slope of the plot (with a zero intercept) was
found to be (1.47 ± 0.05) · 10-1 s-1 at 25 °C, which corresponds to a turnover frequency of 530 h-1
under the selected reaction conditions.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
109
Figure 4.16 (A) Dependence of the initial rate of the Mn(II) catalyzed degradation reaction of Orange II on the [PAA] at
2 · 10-5
and 2 · 10-6
M MnII. (B) Dependence of the initial rate on the [Mn
II]. Reaction conditions: 0.05 M NaHCO3, pH
9.5, 5 · 10-5
M Orange II, 25 °C, (A) 2 · 10-5
M and 2 · 10-6
M MnII, (B) 5 · 10
-3 M PAA.
Influence of bicarbonate buffer. In our earlier work on the MnII catalyzed decomposition of
organic dyes by H2O2, the role of the bicarbonate buffer turned out to be of major importance.[12]
The second order dependence of the observed reaction rate on the [total bicarbonate] suggested
the involvement of HCO3- in two different ways. The first was the role of bicarbonate in the in situ
generation of peroxycarbonate, whereas the second was ascribed to the formation of highly
reactive manganese(II)-bicarbonate complexes under the selected experimental conditions (pH 8.5
- 9).[12a, b] However, when PAA is used as primary oxidant, this situation changes completely. Figure
S4.20 (Supporting Information) shows the dependence of the initial reaction rate on the total
carbonate concentration for PAA as oxidant. The oxidative degradation is obviously decelerated at
higher [total bicarbonate], which on one hand could be caused by enhanced formation of
insoluble MnCO3 at this pH. But on the other hand, if the results are compared to the reaction of
MnII with PAA in the absence of Orange II, an unequivocal tendency is observed. As described
above, higher [total bicarbonate] resulted in the formation of more MnVIIO4-. This suggests that the
more MnVIIO4- is formed, the less reactive the catalytic system is, which in turn emphasizes that
permanganate is not the actual catalytically oxidizing species.
Influence of [H2O2]. Figure 4.17 illustrates the change in the initial bleaching rate upon
addition of H2O2 to the catalytic reaction mixture.
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
110
Figure 4.17 Dependence of the initial rate of the MnII catalyzed degradation reaction of Orange II by PAA on the added
[H2O2]. Inset: Corresponding kinetic traces recorded at 484 nm for the reaction of 1 · 10-5
M MnII with 0.01 M PAA and
5 · 10-5
M Orange II at different concentrations of added H2O2. Reaction conditions: 1 · 10-5
M MnII, 0.01 M PAA, 5 · 10
-5
M Orange II, 0.05 M NaHCO3, pH 9.5 and 25 °C.
The results show that higher [H2O2] inhibit the bleach reaction. However, it was shown that
small quantities of H2O2 reactivated the catalytic system by acting as reducing agent for MnVIIO4-
and colloidal MnIVO2 formed during the second, less reactive phase of the reaction. Thus, on the
one hand the catalytic degradation reaction is hindered by excess of additional H2O2, but on the
other hand it benefits from the presence of a low [H2O2]. This seemingly inconsistent behavior can
be accounted for by the EPR spectroscopic experiments for the reaction of Mn II with PAA in the
presence of an excess H2O2 (see Figures 4.8 (A) and (B)). It was evidenced that an excess of H2O2
delayed the in situ formation of the MnIV intermediate, while the MnII precursor persists longer in
solution most likely in a catalase-like reaction of the MnIV species to MnII. However, a small
amount of H2O2 avoids the visible formation of higher-valence species such as permanganate.
Hence, the ambivalent role of H2O2 can be interpreted in terms of stabilizing a low steady-state
concentration of a reactive MnIV=O species, whereas an excess of H2O2 pushes the complex
oxidation by PAA/reduction by H2O2 equilibrium back to the side of the MnII starting compound.
Catalytic cycles. In control experiments the stability of the in situ generated catalyst was
studied by repeated addition of substrate to a solution of MnII with an excess of PAA (see Figure
4.18). The catalytic cycle could be repeated about five times by sole addition of new portions of
dye substrate without any significant loss of activity. Yet, at the stage of the sudden formation of
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
111
colloidal MnIVO2 and MnVIIO4-, the catalytic degradation performance decreased drastically,
although a large excess of PAA oxidant was still present in solution. Addition of small aliquots of
H2O2 as reducing agent restored the catalytic activity. In this manner, several cycles could be
completed up to the depletion of PAA.
Figure 4.18 Absorbance changes recorded at 484 nm during the repeated addition of 5 · 10-5
M Orange II to a 1 · 10-4
M MnII and 0.025 M PAA containing solution. Asterisks mark the reactivation by addition of 5 · 10
-3 M H2O2. Reaction
conditions: 0.05 M NaHCO3, pH 9.5 and 25 °C.
As a first step to the possible use for larger substrate quantities, we performed a preliminary
experiment with a higher substrate to catalyst ratio, as required for putative applications. Within
only 100 s complete oxidative degradation of the 50 fold excess of Orange II (5 · 10-4 M) was
achieved with only 10 µM MnII catalyst, which implies an even higher TOF of about 1800 h-1. The
experimental results provide clear evidence for a highly efficient catalytic turnover during the
reactive first phase of the MnII/PAA system and demonstrate its possible application for larger
substrate quantities (see Figure S4.21, Supporting Information).
In order to gain further mechanistic insight, the reaction intermediates and final products formed
during the reaction process were examined by HPLC and mass spectrometric methods.HPLC
product analysis was performed immediately after decoloration of the reaction solution
containing MnII, PAA and Orange II (typical experimental conditions as in Figure 4.2 (A), for which
the reaction is over in 100 s). The selected HPLC protocol was very similar to that used for the
analysis of the degradation of Orange II by H2O2 in the presence of [RuIII(edta)H2O]- as catalyst.[48]
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
112
A typical chromatogram is shown in Figure S4.22 (Supporting Information). The broad peak at 3.00
- 3.82 min can be assigned to an overlap of short-chain organic acids including phthalic acid, oxalic
acid and glycolic acid that have retention times of 3.13, 3.45 and 3.79 min, respectively, under the
selected experimental conditions. Although we were not able to resolve the broad signal at 3.82
min any further by variation of the chromatographic parameters, these results are in agreement
with that found in earlier studies.[48, 6d]
We further performed mass spectrometric analysis to compare the degradation products of
Orange II with that of other well studied catalytic systems. The corresponding spectral changes
during this reaction at 484 nm are shown in Figure S4.23 (Supporting Information). Oxalic acid and
glycolic acid could not be detected by this method since they are known to decompose at higher
temperatures. However, the presence of phthalic acid and 4-hydroxybenzosulfonate could be
demonstrated (see Figure S4.24, Supporting Information), which is again in good agreement with
literature reported degradation products.[6d]
4.3.7 Mechanistic interpretation
Although from a mechanistic point of view, the actual catalytically active intermediates are
quite difficult to pin down, some valid conclusions can be drawn on the basis of the available
experimental results. The activation of PAA certainly involves a labile manganese aqua-hydroxo
species and the deprotonated form of the peroxide, H3CCOOO-, as can be concluded from the pH
sensitive rate profile for the overall dye degradation reaction.
The reaction of these components results in the transient formation of an acetylperoxo-MnII
intermediate. Such peroxo intermediates are often proposed to precede the formation of high-
valence manganese-oxo species invoked as active oxidants for synthetic manganese catalysts.
Recent developments suggest that more attention should be paid to the role of these metal-
peroxo species in general, since they may be more important than considered before.[49] However,
in the case of the MnII catalyzed oxidative degradation of Orange II, the involvement of catalytic
active MnII-peroxo species is excluded, since the EPR experiments show only negligible presence of
MnII species during the reactive phase. In general, two possible reaction pathways are conceivable
for a MnII-peroxo intermediate. A heterolytic cleavage of the peroxo bond will lead to a MnIV-oxo
species, whereas a homolytic bond cleavage will yield a MnIII intermediate and an organic radical.
Since the formation of both MnIII and MnIV species could be evidenced in the first phase of the
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
113
reaction by UV/Vis and EPR spectroscopy, it has to be considered that both reactions may occur.
Nevertheless, the interference of free radical processes in the degradation of the dye was
excluded by the use of tBuOH and BHT as radical scavengers, which had no negative effect on the
degradation reaction. The presence of MnIII at this stage could also be due to the manifold dis- and
synproportionation reactions of manganese ions in the absence of a stabilizing ligand. Control
experiments showed no degradation activity of MnIII itself towards the studied dye substrate (see
Figure S4.25, Supporting Information). It is known that the reduction of colloidal MnIVO2 in non-
acidic media results only in partial reduction on the colloidal surface, which is a monomolecular
MnII-oxide layer on the surface of a mixed MnIVO2-MnIIO colloid.[28] Finally, this mixed colloid may
form to some extent MnIII-oxide, which causes the formation of an absorbance shoulder at 470
nm. Taking all this into account, it is suggested that the first oxidation step mainly proceeds via O-
O bond heterolysis to give a high-valence MnIV=O species, which was evidenced by EPR
spectroscopy. Furthermore, a slight increase in the UV/Vis absorbance, as well as in the MnIV EPR
signal, indicate a small increase in the concentration of this MnIV intermediate during the first
reaction phase.
The presence of an excess of PAA would allow further oxidation of the intermediate MnIV
species and thereby enable the formation of higher-valence species like MnV intermediates,
MnVIO42- and MnVIIO4
-. The manganese-oxo anions of MnV and MnVI are highly unstable at pH
values below 14,[23] so that they rapidly disproportionate to yield permanganate and colloidal
MnIVO2 as final products as shown by UV/Vis spectroscopy.
Despite the presence of an excess of PAA, the degradation rate slows down drastically as soon
as the visible formation of permanganate occurs, which points to the requirement of lower-
valence manganese species rather than permanganate for efficient catalysis. Consequently, the
catalytically active intermediate must be formed in the initial reactive phase. During this reaction
stage only MnIV and a small amount of MnIII and MnII were detected. Moreover, the increasing
reactivity towards the dye substrate along with the concomitant increase in the [MnIV
intermediate], strongly suggest that the latter is the actual catalytically active intermediate formed
during the reaction of MnII with PAA. This is in agreement with earlier studies on the peroxide
activation ability of MnII ions, where the observed reactivity was also attributed to a transiently
formed MnIV=O species.[12a, b]
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
114
Although MnIV=O species are commonly postulated as reactive intermediates in numerous
oxygenation reactions, they tend to yield MnIVO2 precipitates in the absence of a stabilizing ligand.
Since the first initial phase with seemingly constant intermediate distribution is very unlikely to be
the result of a stable MnIV=O compound in the absence of a stabilizing ligand, there has to be an
ongoing back reaction that avoids the formation of MnIVO2 as well as the accumulation of MnVIIO4-.
It was evidenced that this back reaction is caused by the equilibrium concentration of H2O2
omnipresent in PAA solutions. H2O2 acts as reducing species for the higher-valence manganese
intermediates such as colloidal MnIVO2 and MnVIIO4- under the experimental conditions of this
study. Thereby, a small [H2O2] keeps the complex oxidation/reduction/disproportionation
equilibria of manganese on the side of the lower-valence MnIII and MnIV species during the initial
phase of the reaction of MnII with PAA. This is supported by the fact that in the UV/Vis spectra
recorded during the initial reaction stage no indication for the formation of O3MnV(OH)2-, MnVIO42-
or MnVIIO4- was found. If the system is selectively reduced by H2O2 at the stage of the second
phase, its catalytic performance is almost completely restored, which strongly supports the
observation that the Mn species produced in the initial phase are crucial for the catalytic activity.
While small [H2O2] are required to maintain the catalytic activity of the system, larger
concentrations (5 - 40 · 10-3 M) of additional H2O2 negatively affect the catalytic performance. This
observation was further clarified by detailed EPR measurements on the reaction of MnII with PAA
in the presence of an excess of additional H2O2. This resulted in a delayed formation of MnIV
species during which the MnII precursor persisted much longer in solution. Furthermore, gas
evolution was observed in the absence of Orange II and with excess H2O2. It is suggested that this
is due to a catalase-like reaction between the in situ formed MnIV species and hydrogen peroxide.
This suggestion also provides an explanation for the increasing reactivity towards the dye
substrate during the initial phase of the reaction. With progressive H2O2 consumption, the back
reaction to MnII becomes less important and the concentration of the reactive MnIV=O slightly
increases. However, as soon as H2O2 has been depleted, MnVIIO4- and colloidal MnIVO2 accumulate
and the catalytic reactivity is lost. Hence, the role of H2O2 in the MnII/PAA system is an ambivalent
one. On the one hand, at low [H2O2] it benefits the catalytic dye degradation by preventing two
contra productive processes. The catalytic life-time is extended by avoiding the rapid over-
oxidation of the Mn catalyst to MnVIIO4- and by preventing catalytic deactivation when MnIVO2 is
formed. On the other hand, higher [H2O2] inhibit the catalytic dye degradation. It pushes the
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
115
complex oxidation by PAA/reduction by H2O2 equilibrium back to the side of the MnII starting
compound, which thereby represents a parallel reaction to the desired dye degradation. A
simplified mechanistic scheme to account for the described observations is presented in Scheme
4.3.
Scheme 4.3 Simplified reaction pathways leading to the formation of reactive MnIV
=O species and permanganate in
the MnII catalyzed oxidative degradation of Orange II by PAA (H3CC(O)OO
-) including the role of the equilibrium
content of H2O2 as reductive species (S = Orange II).
Although reactivity tests with synthetic samples of hypomanganate(V), manganate(VI) and
permanganate(VII) under comparable experimental conditions revealed that permanganate is the
most reactive species for stoichiometric oxidation, they are not relevant to the catalytic process
that occurs during the reaction of the MnII catalyst with PAA. Consequently, it has to be concluded
that the remarkably high reactivity of the MnII/PAA system towards the dye substrate Orange II at
pH 9.5 is due to a well balanced sequence of oxidation (by PAA) and reduction (by H2O2) reactions
to maintain an ideal steady state concentration of an highly reactive MnIV-oxo intermediate for
efficient catalysis of dye degradation. This sequence of events is outlined schematically in Scheme
4.3.
It is known that ortho-substituted azo dyes may act as ligands for metal centers and their
chelating abilities strongly depend on the nature and number of substituents especially adjacent
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
116
to the azo linkage.[12a, 50] Earlier studies reported that a specific complexation mode of ortho-
dihydroxy substituted azo dyes to MnII was required for efficient oxidation catalysis.[50b, d]
However, mono-ortho or unsubstituted dyes showed no reaction which was attributed to their
lower metal binding constants.[50b, d] Consequently, two further azo dyes were selected to clarify to
what extent dye coordination is an essential prerequisite for dye oxidation. Methyl Orange
contains no ortho substituent and thus represents a monodentate ligand, and Calmagite is an o,o’-
dihydroxy azo dye with high binding affinity. Under identical reaction conditions, a very definite
tendency was observed for these dyes (see Figure S26, Supporting Information). Whereas for
Methyl Orange almost no oxidation was observed, the introduction of one hydroxo group adjacent
to the azo bridge in Orange II already resulted in an enhanced performance. Furthermore, a
markedly faster decrease in the characteristic azo absorbance occurred in the case of Calmagite.
These findings are in line with earlier results and strongly suggest that coordination of the dye to
the metal center occurs during the reaction course and is indispensable for efficient catalysis.[12,
50b, d]
Whether the dye substrate coordinates to the MnII precursor or a high-valence form such as
the reactive MnIV=O intermediate, remains to be resolved. If these effects require dye
coordination to manganese in its pre-catalyst MnII form, no reaction or a slower reaction should
occur when the substrate is added to the MnII/PAA containing reaction mixture with a delay of a
several seconds, where our experiments show that MnII is no longer present in the reaction
mixture (see EPR experiments). In fact, the reaction rate for the degradation of Orange II increases
as more MnII is converted to MnIV=O (see Figure 4.12), suggesting that coordination occurs to the
MnIV intermediate. In general, the tendency of the electron rich dye molecule to bind to a metal
centre is expected to increase with increasing electrophilicity of the metal center. Thus it can be
argued that the coordinating influence of the different ortho-substituted dyes can be ascribed to
complex-formation with a high-valence MnIV intermediate rather than with the MnII precursor
form of the catalyst.
4.3.8 Comparison MnII/PAA vs. MnII/HCO4- system
In earlier investigations on the MnII catalyzed oxidative degradation of organic dyes by H2O2, it
was found that a crucial aspect of the catalytic system was the in situ formation of
peroxycarbonate as actual oxidant.[12] Peroxycarbonate formation is known to proceed in a rapid
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
117
equilibration process, however, with an unfavorable formation constant, viz. K = 0.32 ± 0.02 M-1
for solvent mixtures[14a] and K = 0.33 ± 0.02 M-1 for pure aqueous solution.14b Therefore, high
[H2O2] and [total carbonate] are required to obtain an adequate equilibrium concentration of
HCO4- in the catalytic dye degradation with MnII/H2O2/HCO3
- at pH 8.5. Furthermore, it was shown
that the crucial reaction step involves a two electron oxidation of the MnII precursor to a MnIV-oxo
intermediate, but the oxidation state of the metal did not exceed MnIV, and MnVIIO4- formation
never occurred.[12] EPR studies on the MnII/HCO4- system disclosed the presence of a very low
concentration of MnIV, while the main spectroscopic features pointed to the dominance of MnII
species in the catalytic reaction mixture (see Figure 4.19).
Figure 4.19 Comparison of the EPR spectra for MnII/H2O2/HCO3
-12b vs. Mn
II/PAA. Inset: Amplification of the g = 4 signal
for MnII/H2O2/HCO3
-. Conditions for Mn
II/H2O2/HCO3
-: 1 · 10
-4 M Mn
II, 0.02 M H2O2 in 0.4 M HCO3
-, directly after mixing
at pH 8.5 and 25 °C. EPR spectra: 7 K, 9.4 GHz, 2 mW microwave power. Conditions for MnII/PAA: 1 · 10
-4 M Mn
II, 2.5 ·
10-2
M PAA and 5 · 10-3
M H2O2 at pH 9.8 and 25 °C. EPR spectra: 8.96 GHz, 10 K, 1 mW microwave power.
On the other hand, the use of the readily accessible peroxide PAA enables a significantly faster
catalytic degradation at lower [catalyst], [oxidant] and [buffer], during which the formation of
higher-valence Mn-oxo species such as permanganate is observed. Moreover, the MnII/PAA
system suffers less from the undesired catalase-like parallel reaction due to the lower [H2O2]
present. Yet, it was shown that H2O2 is indispensable as reducing agent for efficient catalysis in the
PAA/MnII system, which in turn emphasizes the decisive role of the MnIV species in the catalytic
degradation of Orange II by PAA.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
118
A comparison of the EPR results obtained for both systems, i.e. MnII/H2O2/HCO3- [12b] vs.
MnII/PAA, points to a much higher concentration of the transiently formed MnIV species when PAA
is used. For MnII/H2O2/HCO3- the excess H2O2 used, which is required for the more efficient
formation of HCO4-, enhances the reduction of the in situ formed MnIV=O to MnII, such that only a
very low [MnIV] is constantly present in the catalytic solution. The MnII/PAA system benefits from
the readily accessible peroxoacetic acid, so that more MnIV=O is formed and the lower equilibrium
concentration of the reducing H2O2 keeps it at a higher steady-state concentration as compared to
the MnII/H2O2/HCO3- system. Thus, it is suggested that both reaction systems, i.e. MnII/H2O2/HCO3
-
and MnII/PAA, involve the same catalytically active MnIV=O intermediate, and that through the
described effects of H2O2 the different steady state concentrations of MnIV=O account for the
observed difference in catalytic activity.
4.4 CONCLUSIONS
In conclusion, we were able to gain more insight into the complex chemical and mechanistic
behavior, and the reactivity of the different in situ formed species that occur during the reaction of
simple MnII salts with PAA in weakly basic solution. The kinetics of the catalytic oxidative dye
degradation was investigated in detail for Orange II as substrate. Our kinetic studies lead to the
suggestion that a highly reactive mixture is formed upon addition of PAA to a weakly basic solution
of MnII. The reaction shows a biphasic behaviour which is very sensitive to the selected conditions
and concentrations of the complex catalytic system. By careful selection of the reaction
conditions, we were able to closely investigate the different reaction steps by means of UV/Vis
and EPR spectroscopy, and thereby disclose the different intermediates and their relevance to the
catalytic dye degradation process. Selective reactivity studies of the different in situ formed
intermediates, as well as readily prepared high-valence manganese species, provided further
information on the possible reactive species. Moreover, the omnipresent equilibrium content of
H2O2 in PAA solutions was shown to play a crucial role as reductive species in the catalytic cycle,
presumably by avoiding precipitation of inactive MnIVO2 and over-oxidation of the catalyst,
thereby extending the lifetime of the catalytic system.
On the basis of the experimental results in the presence and absence of substrate, a simplified
reaction scheme to account for the observed behaviour is presented. The key feature of the Mn II
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
119
catalyzed degradation of Orange II by PAA is suggested to involve a complex oxidation by
PAA/reduction by H2O2 reaction sequence in which a steady state equilibrium of the reactive
MnIV=O intermediate sets in. A comparison of the results of the presented detailed kinetic
investigations led to the conclusion that in both systems the same catalytically active intermediate
accounts for the oxidative degradation of the substrate Orange II. The much higher degradation
rate found in the presence of PAA can be ascribed to the more efficient formation of Mn IV=O than
in the presence of HCO4-, as supported by the EPR measurements.
4.5 EXPERIMENTAL SECTION
CHEMICALS. Orange II, certified [Acid Orange 7, C.I. 15510, sodium 4-(2-hydroxy-1-
naphthylazo)benzenesulfonate], 99 % was supplied by Sigma-Aldrich and recrystallized from an
EtOH/H2O mixture at 4 °C.[51 ] Peracetic acid 39 wt. %, H2O2 30 %, as well as CH3-13COOH were of
analytical grade and provided by Sigma-Aldrich. Analytical grade tBuOH, BHT (2,6-di-tert-butyl-4-
methylphenol), Calmagite (1-(1-hydroxy-4-methyl-2-phenylazo)-2-naphthol-4-sulfonic acid), and
Methyl Orange (4-[4-(dimethylamino)phenylazo]benzenesulfonic acid, sodium salt) were used. All
other chemicals were of the highest purity commercially available and used without any further
purification. Carbonate buffer solutions were prepared using Millipore Milli-Q purified water.
Stock solutions of manganate(VI)[52] and colloidal MnIVO2[25] were prepared according to the
literature and monitored by UV/Vis (MnVIO42-: ε610 nm = 1500 M-1cm-1 [23]) and EPR spectroscopy,
respectively. Hypomanganate(V) solutions were prepared by careful reduction of manganate(VI)
with Na2SO3 as reported earlier.[23] The concentration of the hypomanganate(V) solutions was
estimated by the molar extinction coefficient as published before.[45]
GENERAL PROCEDURE AND PH JUMP TECHNIQUE. The different manganese salts (MnIICl2 · 2H2O, MnIISO4 ·
H2O, MnII(NO3)2 · 4H2O, MnII(ClO4)2 · 4H2O and MnII(O2CCH3)2 · 4H2O) were freshly dissolved in
water before use. To a freshly prepared 0.05 M sodium bicarbonate solution, an adequate amount
of NaOH was added to adjust the pH in a way that the subsequent addition of a specific [PAA] gave
the desired pH. The reaction was started under isothermal conditions by addition of small aliquots
of a concentrated manganese stock solution together with Orange II to the PAA containing buffer
solution. The catalytic reaction was followed by in situ UV/Vis spectroscopy.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
120
INSTRUMENTATION AND EQUIPMENT. All kinetic data were obtained by recording time-resolved UV/Vis
spectra using a Hellma 661.502 – QX quartz Suprasil immersion probe attached via optical cables
to a 150 W Xe lamp and a multi-wavelength J & M detector, which records complete absorption
spectra at constant time intervals. Kinetic measurements were carried out under pseudo-first-
order conditions. The pH of the PAA containing aqueous carbonate solution was carefully
measured and adjusted using a Mettler Delta 350 pH meter previously calibrated with standard
buffer solutions at two different pH values (4 and 10). The kinetics of the Orange II degradation
reaction was monitored at 484 nm. First-order rate constants, where possible, were calculated
using Specfit/32 and Origin (version 7.5) software.
13C-NMR measurements were performed by using DMSO-d6 as internal standard in a glass
capillary. The [PAA] and [AcOH] were calculated from the relative peak intensities using the Lorenz
fit obtained by the NMRICMA program (developed at the Institut de chimie minérale et analytique,
University of Lausanne, in the group of Prof. A. E. Merbach) from the NMR data. 13C-NMR spectra
were recorded at a frequency of 100 MHz on a Bruker Advance DRX 400WB spectrometer
equipped with a superconducting BC-94/89 magnet system.
Perpendicular mode EPR spectra were recorded on an X-band Joel Jes Fa 200 spectrometer
equipped with a cylindrical mode cavity and a liquid helium cryostat. Samples were taken from the
investigated solutions and immediately frozen to quench the reaction. The EPR measurements
were performed in quartz tubes at 10 K (9.45 GHz, 1 mW microwave power). Data analyses were
done with the Jes-Fa Series software package.
HPLC analysis of the oxidation products of Orange II was performed using a Waters (M 515 &
PDA) HPLC equipped with a photodiode array detector and a Symmetry C18 (5 m, 100 A) column.
Resultant solutions collected after 4 min and 60 min of the reaction were subjected to HPLC
analysis using a mobile phase of HPLC grade water-acetonitrile mixture (70:30 v/v) at a flow rate of
0.5 mL/min. HPLC parameters were quantified and optimized with authentic samples of naphthalic
acid, glycolic acid and oxalic acid prior to the analysis.
Mass spectrometric analysis of the oxidation products of ORII were performed on an UHR-TOF
Bruker Daltonik (Bremen, Germany) maXis, an ESI-TOF mass spectrometer capable of a resolution
of at least 40 000 fwhm used by the group of Prof. Ivana Ivanović-Burmazović at the University of
Erlangen-Nürnberg. The mass spectrometric detection was carried out in the negative-ion mode
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
121
with a source voltage of 5500 V, a flow rate of 500 μL/h, and the drying gas (N2) was kept at 180
°C. The instrument was calibrated prior to every experiment via direct infusion of the Agilent ESI-
TOF low concentration tuning mixture. The sample was taken after completion of the reaction
performed under the standard conditions (1 · 10-5 M MnII, 0.01 M PAA, 0.05 M NaHCO3 buffer, pH
9.5). The HCO3-/CO3
2- buffer was removed by acidification with HCl and the pH was readjusted
with NaOH prior to the analysis.
4.6 REFERENCES AND NOTES
[1] (a) Park, H.; Choi, W. J. Photochem. Photobiol. A 2003, 159, 241-247; (b) Brown, D.,
Laboureur, P. Chemosphere 1987, 12, 397-404; (b) Goszczynski, S.; Paszczynki, A.; Pasti-
Grigsby, M. B.; Crawford, R. L.; Crawford, D. L. J. Bacteriol. 1994, 176, 1339-1347; (c)
Chivukula, M.; Spadaro, J. T.; Renganathan, V. Biochemistry 1995, 34, 7765-7772; (d) Spadaro,
J. T.; Renganathan, V.; Arch. Biochem. Biophys. 1994, 312, 301-307; (e) Spadaro, J. T.; Isabelle,
L.; Renganathan, V. Environ Sci & Tech. 1994, 28, 1389-1393.
[2] (a) Zollinger, H. Color Chemistry, 3rd ed., Wiley-VCH: Weinheim, Germany, 2003; (b) Beach, E.
S.; Malecky, R. T.; Gil, R. R.; Horwitz, C. P.; Collins; T. J. Catal. Sci. Technol. 2011, 1, 437-443.
[3] Baughman, G. L.; Weber, E. J.; Environ. Sci. Technol. 1994, 28, 267-276.
[4] Hastie, J.; Bejan, D.; Teutli-Leon, M.; Bunce, N. J. Ind. Eng. Chem. Res. 2006, 45, 4898-4909.
[5] (a) Ellis, W. C.; Tran, C. T.; Roy, R.; Rusten, M.; Fischer, A.; Ryabov, A. D.; Blumberg, B.; Collins,
T. J. J. Am. Chem. Soc. 2010, 132, 9774-9781. (b) Kluson, P.; Drobek, M.; Krejcikova, S.; Krysa,
J.; Kalaji, A.; Cajthaml, T.; Rakusan, J. Appl. Catal. B 2008, 80, 321-326. (c) Wieprecht, T.; Xia,
J.; Heinz, U.; Dannacher, J.; Schlingloff, G. J. Mol. Catal. A: Chem. 2003, 203, 113-128. (d)
Ramirez, J. H.; Maldonado-Hódar, F. J.; Pérez-Cadenas, A. F.; Moreno-Castilla, C.; Costa, C. A.;
Madeira, L. M. Appl. Catal. B 2007, 75, 312-323. (e) Hage, R.; Iburg, J. E.; Kerschner, J.; Koek, J.
H.; M. Lempers, E. L.; Martens, R. J.; Racherla, U. S.; Russell, S. W.; Swarthoff, T.; van Vliet, M.
R. P.; Warnaar, J. B.; van der Wolf, L.; Krijnen, B. Nature 1994, 369, 637-639.
[6] (a) Oakes, J.; Gratton, P. J. Chem. Soc., Perkin Trans. 2 1998, 1857-1864. (b) Nadtochenko, V.;
Kiwi, J. J. Chem. Soc., Faraday Trans. 1997, 93, 2373-2378. (c) Oakes, J.; Gratton, P; Weil, I. J.
Chem. Soc., Dalton Trans. 1997, 3805-3809. (d) Chahbane, N.; Popescu, D.-L.; Mitchell, D. A.;
Chanda, A.; Lenoir, D.; Ryabov, A. D.; Schramm, K.-W.; Collins, T. J. Green Chem. 2007, 9, 49-
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
122
57. (e) Gilbert, B. C.; Lindsay-Smith, J. R.; Newton, M. S.; Oakes, J.; Ponts i Parts, R. Org.
Biomol. Chem., 2003, 1, 1568-1577.
[7] (a) Song, Y. J.; Lee, S. H.; Park, H. M.; Kim, S. H.; Goo, H. G.; Eom, G. H.; Lee, J. H.; Lah, M. S.;
Kim, Y.; Kim, S.-J.; Lee, J. E.; Lee, H.-I.; Kim, C. Chem. Eur. J. 2011, 17, 7336-7344. (b) Song, F.
Wang, C.; Falkowski, J. M.; Ma, L.; Lin, W. J. Am. Chem. Soc. 2010, 132, 15390-15398. (c)
Palucki, M.; Finney, N. S.; Pospisil, P. J.; Guler, M. L.; Ishida, T.; Jacobsen, E. N. J. Am. Chem.
Soc., 1998, 120, 948-954.
[8] (a) Arunkumar, C.; Lee, Y.-M.; Lee, J. Y.; Fukuzumi, S.; Nam, W. Chem. Eur. J. 2009, 15, 11482-
11489. (b) Crestoni, M. E.; Fornarini, S.; Lanucara, F. Chem. Eur. J. 2009, 15, 7863-7866. (c)
Song, W. J.; Seo, M. S.; de Beer George, S.; Ohta, T.; Song, R.; Kang, M.-J.; Tosha, T.; Kitagawa,
T.; Solomon, E. I.; Nam, W. J. Am. Chem. Soc. 2007, 129, 1268-1277. (d) Rose, E.; Andrioletti,
B.; Zrig, S.; Quelquejeu-Etheve, M. Chem. Soc. Rev., 2005, 34, 573-583.
[9] (a) Wieghardt, K.; Schmidt, W.; Nuber, B.; Weiss, J. Chem. Ber. 1979, 112, 2220-2230. (b)
Romakh, V. B.; Therrien, B.; Süss-Fink, G.; Shul’pin, G. B. Inorg. Chem. 2007, 46, 1315-1331; (c)
Mandelli, D.; Kozlov, Y. N.; Golfeto, C. C.; Shul'pin, G. B. Catal. Lett. 2007, 118, 22-29; (d)
Sibbons, K. F.; Shastri, K.; Watkinson, M. Dalton Trans. 2006, 645-661; (e) Gilbert, B. C.;
Lindsay Smith, J. R.; Mairata y Payeras, A.; Oakes, J. Org. Biomol. Chem. 2004, 2, 1176-1180;
(f) Bennur, T. H.; Srinivas, D.; Sivasanker, S.; Puranik, V. G. J. Mol. Cat. A 2004, 219, 209-216;.
(g) Hage, R.; Kerschner, J. Trends in Inorg. Chem. 1998, 5, 145-159.
[10] De Boer, J. W.; Browne, W. R.; Brinksma, J.; Alsters, P. L.; Hage, R.; Feringa, B. L. Inorg. Chem.
2007, 46, 6353-6372.
[11] (a) Shilov, A. E.; Shul’pin, G. B. Chem. Rev. 1997, 97, 2879-2932; (b) Sheldon, R. A. J. Chem.
Technol. Biotechnol. 1997, 68, 381-388.
[12] (a) Ember, E.; Rothbart, S.; Puchta, R.; van Eldik, R. New J. Chem. 2009, 33, 34-49; (b) Ember,
E.; Gazzaz, H. A.; Rothbart, S.; Puchta, R.; van Eldik, R. Appl. Catal. B. 2009, 95, 179-191; (c)
Rothbart, S.; Ember, E.; van Eldik, R. Dalton Trans. 2010, 39, 3264-3272.
[13] Baes, C. F.; Mesmer, R. E. Am. J. Sci. 1981, 281, 935-962.
[14] (a) Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett, D. A. J. Am. Chem. Soc. 2000, 122, 1729-
1739. (b) Bakhmutova-Albert, E. V.; Yao, H.; Denevan, D. E.; Richardson, D. E. Inorg. Chem.
2010, 49, 11287-11296.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
123
[15] Silbert, L. S.; Organic Peroxides (Ed:D . Swern), Wiley-Interscience, New York 1970.
[16] (a) Evans, D. F.; Upton, M. W. J. Chem. Soc., Dalton Trans. 1985, 1151-1158 (b) Koubek, E.;
Haggett, M. L.; Battaglia, C. J.; Ibne-Rasa, K. M.; Pyun, H. Y.; Edwards, J. O. J. Am. Chem. Soc.
1963, 85, 2263-2268 (c) Yuan, Z. ; Ni, Y.; Heiningen, A. R. P. Can. J. Chem. Eng. 1997, 75, 37-
41.
[17] (a) Ball, R. E.; Edwards, J. O. J. Am.Chem. Soc. 1956, 78, 1125-1129; (b) Awad, M. I.; Harnoode,
C.; Tokuda, K.; Ohsaka, T. Anal. Chem. 2001, 73, 1839-1843.
[18] (a) Shapilov, O. D. Doctoral (Pharm.) Dissertation, Moscow, 1978; (b) Doguzova, I. A. Cand.
Sci. (Chem.) Dissertation, Moscow, 2000. (c) Dul’neva, L. V.; Moskvin, A. V. Russian Journal of
General Chemistry, 2005, 75, No. 7, 1125-1130.
[19] Feng, W.; Nansheng, D.; Helin, H. Chemosphere 2000, 41, 1233-1238.
[20] (a) Jee; J.-E.; Pestovsky, O.; Bakac; A. Datlton Trans. 2010, 39, 11636-11642. (b) Wells, C. F.,
Davies, G. J. Chem. Soc. A. 1967, 1858-1861.
[21] (a) Wei, Z.; Cady, C. W.; Brudvig, G. W.; Hou, H. J. M. J. Photochem. Photobiol. B 2011, 104,
118-125; (b) Chen, H.; Tagore, R.; Das, S.; Incarvito, C.; Faller, J. W.; Crabtree, R. H.; Brudvig,
G. W. Inorg. Chem. 2005, 44, 7661-7670; (c) Gamelin, D. R.; Kirk, M. L.; Stemmler, T. L.; Pal, S.;
Armstrong, W. H.; Penner-Hahn, J. E.; Solomon, E. I. J. Am. Chem. Soc. 1994, 116, 2392-2399;
(d) Belal, A. A.; Chaudhuri, P.; Fallis, I.; Farrugia, L. J.; Hartung, R.; MacDonald,N. M.; Nuber,
B.; Peacock, R. D; Weiss, J.; Wieghardt, K. Inorg. Chem. 1991, 30, 4397-4402.
[22] Lin, Z.-X.; Huang, T. T. Ind. Eng. Chem. Res. 1987, 26, 2148-2151.
[23] Rush, J. D.; Bielski, B. H. J. Inorg. Chem. 1995, 34, 5832-5838.
[24] Leavason, W.; McAuliffe, C. A. Coord. Chem. Rev. 1972, 2, 353-384.
[25] (a) Perez-Benito, J. F. J. Phys. Chem. C 2009, 113, 15982-15991. (b) Abel, E. Monatheft für
Chemie 1949, 80, 455-462.
[26] (a) Bourgougnon, A. J. Am. Chem. Soc. 1889, 11, 94-98; (b) Doona, C. J.; Schneider, F. W. J.
Am. Chem. Soc. 1993, 115, 9683-9686; (c) Li, Y.; Zhao, Y.; Zhu, Z.; Anal. Sci., 2011, 27, 37-41;
(d) Ladbury, J. W.; Cullis, C. F. Chem. Rev. 1958, 58, 403-438.
[27] Popov, E.; Eloranta, J.; Hietapelto, V.; Vuorenpalo, V.-M.; Aksela, R.; Jäkärä, J. Holzforschung
2005, 59, 507-513.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
124
[28] Perez-Benito, J. F.; Arias, C.; Amat, E. A. J. Collo. Interf. Scienc. 1996, 177, 288-297.
[29] Perez-Benito, J. F. J. Collo. Interf. Scienc. 2002, 248, 130-135.
[30] Petlicki, J.; Palusova, D.; van de Ven, T. G. M. Ind. Eng. Chem. Res. 2005, 44, 2002-2010.
[31] Perez-Benito, J. F.; Arias, J. Collo. Interf. Scienc. 1992, 152, 70-84.
[32] Baral, S.; Lume-Pereira, C.; Janata, E.; Henglein, A. J Phys. Chem. 1985, 89, 5779-5783.
[33] Kanungo, S.B.; Parida, K. M.; Sant, B. R. Electrochim. Acta 1981, 26, 1157-1167.
[34] Evan, M. G.; Uri, N Trans. Faraday Soc. 1949, 45, 224-236.
[35] Simoyi, R. H.; Keeper, P. D.; Epstein, I. R.; Kustin, K. Inorg. Chem. 1986, 25, 583-542.
[36] (a) Casado, J.; Lizaso-Lamsfus, J. An. R. Soc. Es. Fis. Quim. 1971, 63, 739-748. (b) Senent, S.;
Casado, J.; Lizaso, J. An. R. Soc. Es. Fis. Quim. 1971, 67, 1133-1144.
[37] (a) Kozlov, Y. N.; Zharmukhamedov, S. K.; Tikhonov, K. G.; Dasgupta, J.; Kazakova, A. A.;
Dismukes; G. C.; Klimov, V. V. Phys. Chem. Chem. Phys. 2004, 6, 4905-4911; (b) Dasgupta, J.;
Tyryshkin, A. M.; Kozlov, Y. N.; Klimov, V. V.; Dismukes, G. C. J. Phys. Chem. B 2006, 110, 5099-
5111.
[38] Sarneski, J. E.; Didiuk, M.; Thorp, H. H.; Crabtree, R. H.; Brudvig, G. W.; Faller, J. W.; Schulte,
G. K. Inorg. Chem. 1991, 30, 2833-2835.
[39] (a) Saadeh, S. M.; Lah, M. S.; Pecoraro, V. L. Inorg. Chem. 1991, 30, 8-15. (b) John, R. P.;
Sreekanth, A.; Prathapachandra Kurup, M. R.; Fun, H.-K. Polyhedron 2005, 24, 601-610.
[40] Lane, B. S.; Vogt, M.; DeRose, V. J.; Burgess, K. J. Am. Chem. Soc. 2002, 124, 11946-11954.
[41] (a) Duboc, C.; Collomb, M.-N. Chem. Commun. 2009, 2715-2717; (b) Pedersen, E.; Toftlund, H.
Inorg. Chem. 1974, 13, 1603-1612; (c) Pal, S.; Ghosh, P.; Chakravorty, A. Inorg. Chem. 1985,
24, 3704-3706; (d) Hempel, J. C.; Morgan, L. O.; Lewis, W. B. Inorg. Chem. 1970, 9, 2064-2072.
[42] (a) Magers, K. D.; Smith, G. C.; Sawyer, D. T. Inorg. Chem. 1980, 19, 492-496. (b) Richens, D.
T.; Sawyer, D. T. J. Am. Chem. Soc. 1979, 101, 3681-3683.
[43] Brown, K. L.; Golding, R. M.; Healy, P. C.; Jessop, K. J.; Tennant, W. C. Aust. J. Chem. 1974, 27,
2075-2081.
[44] Cooper, S. R.; Dismukes, G. C.; Klein, M. P.; Calvin, M. J. Am. Chem. Soc. 1978, 100, 7248-7252.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
125
[45] Note that according to ref. [23] MnV-oxo anion is protonated below its pKa of pH 13.7 and in
the form of O3MnV(OH)2-.
[46] Pode, J. S. F.; Waters, W. A. J. Chem. Soc. 1956, 717-725.
[47] (a) Garcia-Bosch, I.; Company, A.; Cady, C. W.; Styring, S.; Browne, W. R.; Ribas, X.; Costas, M.
Angew. Chem. Int. Ed. 2011, 50, 5648-5653; (b) Sawant, S. C.; Wu, X.; Cho, J.; Cho, K.-B.; Kim,
S. H.; Seo, M. S.; Lee, Y.-M.; Kubo, M.; Ogura, T.; Shaik, S.; Nam, W. Angew. Chem. Int. Ed.
2010, 49, 8190-8194; (c) Adam, W.; Roschmann, K. J.; Saha-Möller, C. R.; Seebach, D. J. Am.
Chem. Soc. 2002, 124, 5068-5073; (d) Finney, N. S.; Pospisil, P. J.; Chang, S.; Palucki, M.;
Konsler, R. G.; Hansen, K. B.; Jacobsen, E. N. Angew. Chem. Int. Ed. 1997, 36, 1720-1723; (e)
Murphy, A.; Stack, T. D. P. J. Mol. Cat. A 2006, 251, 78-88.
[48] Chatterjee, D.; Ember, E.; Pal, U.; Ghosh, S.; van Eldik, R. Dalton. Trans. 2011, 40, 10473-
10480.
[49] (a) Mukherjee, A.; Cranswick, M. A.; Chakrabarti, M.; Paine, T. K.; Fujisawa, K.; Münck, E.; Que
Jr., L. Inorg. Chem. 2010, 49, 3618-3628; (b) Geiger, R. A.; Chattopadhyay, S.; Day, V. W.;
Jackson, T. A. J. Am. Chem. Soc. 2010, 132, 2821-2831; (c) Yin, G.; Buchalova, M.; Danby, A.
M.; Perkins, C. M.; Kitko, D.; Carter, J. D.; Scheper, W. M.; Busch, D. H. Inorg. Chem. 2006, 45,
3467-3474.
[50] (a) Yoshida, T.; Sawada, S. Bull. Chem. Soc. Jpn. 1975, 48, 345-346; (b) Oakes, J.; Gratton, P. J.
Chem. Soc., Dalton Trans. 1998, 9, 3805-3809. (c) Madeira, M. T.; Feijoo, G.; Lema, J. M.
Enzyme Microbiol. Technol. 2007, 42, 70-75; (d) Sheriff, T. S.; Cope, S.; Ekwegh, M. Dalton
Trans. 2007, 44, 5119-5122.
[51] Theodoridis, A.; Maigut, J.; Puchta, R.; Kudrik, E. V.; van Eldik, R. Inorg. Chem. 2008, 47, 2294-
3013.
[52] Carrington, A; Symons, M. C. R. J. Chem. Soc. 1956, 3373-3380.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
126
4.7 SUPPLEMENTARY INFORMATION
Figure S4.1 Time course for the consumption of AcOH (A) and the formation of PAA (B). Experimental conditions: 2.5
M AcOH, 2.5 M H2O2, 0.1 M H2SO4 and 25 °C. (C) Time course for the reformation of AcOH and the consumption of
PAA (D). Experimental conditions: 0.041 M PAA (produced from 2.5 M AcOH, 2.5 M H2O2, 0.1 M H2SO4 at 25 °C), 0.1 M
H2SO4 and 25 °C. The concentration profiles were calculated from the Lorenz fit of the 13
C-NMR signals at 179.2 ppm
(H3C-13
COOH) and 175.3 ppm (H3C-13
C(O)OOH).
(A) (B)
(C) (D)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
127
Figure S4.2 Time courses for the formation of AcOH (A) and decomposition of PAA (B) at pH ≈ 10 and 25 °C. The
concentration profiles were calculated from the Lorenz fit of the 13
C-NMR signals at 179.2 ppm (H3C-13
COOH) and
175.3 ppm (H3C-13
C(O)OOH).
Figure S4.3 Comparison of the absorbance changes at 484 nm in the uncatalyzed reaction of 5 · 10-5
M Orange II with
5 · 10-3
M PAA (—), with 5 · 10-3
M PAA and 5 · 10-3
M EDTA (—), and in the presence of 1 · 10-5
M MnII catalyst (—).
Reaction conditions: 0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
128
Figure S4.4 (A) Influence of acetate concentration on the reactivity of 1 · 10-5
M MnII with 0.01 M PAA and 5 · 10
-5 M
Orange in 0.05 M NaHCO3 buffer at pH 9.5 and 25 °C. Absorbance change followed at 484 nm. (B) Comparison of
different MnII salts used as starting material for the catalytic decomposition of 5 · 10
-5 M Orange II with 1 · 10
-5 M Mn
II
and 0.01 M PAA in 0.05 M NaHCO3 buffer at pH 9.5 and 25 °C. Absorbance change followed at 484 nm.
Figure S4.5 Comparison of the reactivity of 1 · 10-5
M MnII with 0.01 M PAA and 5 · 10
-5 M Orange II in the absence (—)
and presence of 0.02 M tBuOH (—) and 0.02 M BHT (—) as radical scavenger followed at 484 nm. Reaction conditions:
0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
129
Figure S4.6 Effect of ionic strength (adjusted with different concentrations of NaNO3 on the catalytic decomposition of
5 · 10-5
M Orange II with 1 · 10-5
M MnII and 0.01 M PAA in 0.05 M NaHCO3 buffer at pH 9.5 and 25 °C. Absorbance
change followed at 484 nm.
Figure S4.7 (A) UV/Vis spectral changes that accompany the reaction of 2 · 10-5
M colloidal MnIV
O2 with 2 · 10-3
M H2O2
(0.05 M NaHCO3 buffer) at pH 9.5 and 25 °C. Inset: Absorbance changes at 350 nm. (B) Observed rate constants for the
reduction of 2 · 10-5
M colloidal MnIV
O2 by 1 · 10-3
M H2O2 followed at 350 nm. Reaction conditions: 0.05 M
bicarbonate or phosphate buffer at 25 °C.
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
130
Figure S4.8 (A) Absorbance change at 525 nm for the reaction of 2 · 10-5
M MnVII
O4- with 5 · 10
-4 M H2O2 as a function
of pH. (B) Corresponding observed rate constants as a function of pH (induction period was neglected). Reaction
conditions: stopped-flow experiments, 0.05 M bicarbonate or phosphate buffer at 25 °C.
Figure S4.9 Absorbance changes at 350 nm (A) and 525 nm (B) for the reaction of 2 · 10-5
M MnII with 0.01 M PAA in
0.05 M NaHCO3 as a function of pH at 25 °C.
(A) (B)
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
131
Figure S4.10. Absorbance changes at 350 nm (A) and 525 nm (B) for the reaction of 2 · 10-5
M MnII with 0.01 M PAA at
different total carbonate concentrations. Reaction conditions: pH 9.5 and 25 °C.
Figure S4.11. X-band EPR spectra recorded at 10 K for 1 · 10-4
M MnVI
O42-
in 10 M KOH (—) and for the first sample
taken after mixing 1 · 10-4
M MnII with 2.5 · 10
-2 M PAA and 5 · 10
-3 M H2O2, 0.05 M carbonate buffer at pH 9.8 and 25
°C (—). EPR conditions: 8.95 GHz, 10 K, 1 mW microwave power, modulation amplitude 400 mT for (—) and 40 mT for
(—).
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
132
Figure S4.12 (A) X-band EPR spectra recorded at 10 K for 1 · 10-4
M MnII in carbonate buffer (—) and at different time
intervals after the addition of 0.025 M PAA and 5 · 10-3
M H2O2 in 0.05 M NaHCO3 with 20 % of tBuOH. (B) Kinetic
traces at 350 and 525 nm. Arrows mark the time at which different EPR samples were taken. Reaction conditions: 1 ·
10-4
M MnII, 0.025 M PAA, 5 · 10
-3 M H2O2, pH 9.8 and 25 °C. EPR conditions: 8.95 GHz, 10 K, 1 mW microwave power,
modulation amplitude 400 mT.
Figure S4.13 (A) X-band EPR spectra recorded at 10 K for 1 · 10-4
M MnII at different time intervals after the addition of
0.025 M PAA and 5 · 10-3
M H2O2 in 0.2 M carbonate buffer (B) Kinetic traces at 350 and 525 nm. Arrows mark the time
at which different EPR samples were taken. Reaction conditions: 1 · 10-4
M MnII, 0.025 M PAA, 5 · 10
-3 M H2O2, pH 9.8
and 25 °C. EPR conditions: 8.95 GHz, 10 K, 1 mW microwave power, modulation amplitude 400 mT.
(A) (B)
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
133
Figure S4.14 (A) Absorbance changes recorded at 484 nm for the reaction of 5 · 10-5
M MnVII
O4- with 5 · 10
-5 M Orange
II as a function of pH. (B) Initial rate as a function of pH. Reaction conditions: 0.05 M NaHCO3 buffer and 25 °C.
Figure S4.15 (A) UV/Vis spectral changes recorded for the disproportionation of 2 · 10-5
M MnVI
O42-
. Inset: kinetic
traces at 525 (—) and 610 nm (—). Reaction conditions: 0.05 M NaHCO3 buffer, pH ≈ 9.5 and 25 °C. (B) Initial rate of
the reaction of 2 · 10-5
M MnVI
O42-
with 5 · 10-5
M Orange II as a function of pH. Inset: observed rate constants for the
disproportionation of 2 · 10-5
M MnVI
O42-
as a function of pH. Reaction conditions: 2 · 10-5
M MnVI
O42-
, 0.05 M NaHCO3
buffer, 25 °C.
(A) (B)
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
134
Figure S4.16 Comparison of the absorbance changes at 465 nm during the reaction of 5 · 10-5
M Orange II with 5 · 10-5
M MnVII
O4- (), Mn
VIO4
2- () and O3Mn
V(OH)
2- () at pH 12.0. Reaction conditions: 0.05 M NaHCO3 and 25 °C.
Figure S4.17 Absorbance change at 484 nm for the reaction of 1 · 10-5
M colloidal MnIV
O2 with 5 · 10-5
M Orange II.
Reaction conditions: 0.05 M NaHCO3 buffer, pH 9.5 and 25 °C.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
135
Figure S4.18 (A) UV/Vis spectra of 1 · 10-4
M MnO4- (—) after the reduction by 4 · 10
-4 M SO3
2- (—) and immediately
after addition of 5 · 10-5
M Orange II (—). (B) Absorbance changes at 484 nm that accompany the reaction of the in situ
generated MnIV
via reduction of MnVII
O4-. Reaction conditions: 1 · 10
-4 M Mn
VIIO4
-, 0.05 M NaHCO3 buffer, pH 9.5 and
25 °C.
Figure S4.19 UV/Vis absorbance changes recorded at 484 nm showing the reactivation of the catalytic system by
addition of 1 · 10-3
M H2O2. Reaction conditions: 1 · 10-5
M MnII, 0.01 M PAA, 5 · 10
-5 M Orange II, 0.05 M NaHCO3
buffer at pH 9.6 and 25 °C.
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
136
Figure S4.20 Dependence of the initial rate of the MnII catalyzed degradation reaction on the total carbonate
concentration. Reaction conditions: 1 · 10-5
M MnII, 0.01 M PAA, 5 · 10
-5 M Orange II, 25 °C and pH 9.5.
Figure S4.21 (A) UV/Vis spectral changes recorded during the reaction of 5 · 10-4
M Orange II with 2.5 · 10-2
M PAA
catalyzed by 1 · 10-5
M MnII. (B) Corresponding absorbance vs. time plot at 484 nm. Reaction conditions: 0.05 M
NaHCO3 buffer at pH 9.5 and 25 °C (During the first 40 s of the measurement the absorbance remains above the
detection limit of the detector due to the high Orange II concentration).
(B) (A)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
137
Figure S4.22 HPLC product analysis for the degradation of Orange II by PAA in the presence of MnII. Experimental
conditions: 5 · 10-5
M Orange II, 1 · 10-5
M MnII and 0.01 M PAA at pH 9.5 and 25 °C.
Figure S4.23 Absorbance vs. time traces at 484 nm for the oxidative degradation Orange II catalyzed by 1 · 10-5
M MnII
and 0.01 M PAA at pH 9.5 (0.05 M NaHCO3 buffer) and 25 °C. Substrate addition was carried out in five consecutive
steps each consisting of 5 · 10-5
M Orange II.
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
138
Figure S4.24 Mass spectra of mono-protonated phthalate (m/z = 165,0184) (A) and 4-hydroxybenzosulfonate (m/z =
172,9904) (B) as products of the degradation reaction in Figure S4.23.
165.0184
166.0217
-MS, 2.5-7.3min #(208-605)
165.0193
166.0227
C8O4H5, M ,165.020
10
20
30
40
50
Intens.
[%]
0
20
40
60
80
100
[%]
164.5 165.0 165.5 166.0 166.5 167.0 m/z
172.9904
173.9935174.9871
-MS, 2.5-7.3min #(208-605)
172.9914
173.9947174.9873
C6SO4H5, M ,172.990
10
20
30
40
Intens.
[%]
0
20
40
60
80
100
[%]
172.5 173.0 173.5 174.0 174.5 175.0 175.5 m/z
(A)
(B)
Experiment
Simulation
Experiment
Simulation
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
139
Figure S4.25 Absorbance changes at 484 nm in the reaction of 1 · 10-4
M MnIII
(OAc)3 with 5 · 10-5
M Orange II. Reaction
conditions: 0.05 M NaHCO3, pH 9.5 and 25 °C.
Figure S4.26 Comparison of the absorbance changes recorded at λmax of the corresponding dye during the reaction of
5 · 10-5
M dye with 1 · 10-5
M MnII (A) and 1 · 10
-6 M Mn
II (B), respectively, and 5 · 10
-3 M PAA containing solution for
Methyl Orange 465 nm (), Orange II 484 nm () and Calmagite 612 nm (). Reaction conditions: 0.05 M NaHCO3,
pH 9.5 and 25 °C.
(A) (B)
4. Metal ion - catalyzed oxidative degradation of Orange II by peracetic acid
140
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
141
5 HIGH CATALYTIC ACTIVITY OF A MN-TERPY COMPOUND IN
OXIDATIVE DYE DEGRADATIONS WITH PERACETIC ACID
5.1 GENERAL REMARK
The following chapter is based on ongoing work: High catalytic activity of manganese
terpyridine in the oxidative catalytic dye degradation with PAA, Sabine Rothbart and Rudi van
Eldik, in preparation.
5.2 INTRODUCTION
Significant ecological impact is caused by industrial dye waste since over 15 % of textile dyes
are lost in waste water streams during the coloring operation.[1] A recent review emphasized the
high cost involved in disposing the high volumes of dye effluent and that 128 tons of dyes are
released daily to the global environment.[2] About half of the global production of synthetic
colorants (700.000 t per year) are classified as aromatic azo compounds which are often designed
to be resistant to chemical, biochemical and photochemical degradation.[3] The increasing
ecological awareness stimulated an active field of scientific research dedicated to new and
ecologically worthwhile oxidation processes for the catalytic decomposition of environmental
pollutants[4] and dyes[5] in water. Among the various transition metals for catalytic oxidation,
manganese is of particular interest as one of the most efficient and environmental benign
elements. There are various manganese complexes with different salen,[6] porphyrin,[7] tacn[8] or
aromatic N-donor ligands[9] known to efficiently catalyze the oxidation of a wide range of
substrates. However, several limitations including elaborated synthetic methods, long reaction
times and substrate scope, have still to be resolved. Recent attention has therefore been focused
on metal complexes that use cheap and clean oxidants like H2O2 or peracetic acid (PAA) to bring
about efficient oxidation under mild reaction conditions.[10]
Our earlier work focused on the reactivity of simple MnII salts and Mn complexes in the H2O2
induced catalytic degradation of various organic dye substrates.[11] Detailed kinetic and
mechanistic investigations revealed the in situ generation of percarbonate (HCO4-)][12] as a key
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
142
molecular oxidant from H2O2 and HCO3-, which leads to the formation of a reactive MnIV=O
species.[11a, b] Despite the high efficiency combined with the striking simplicity of the system, the
use of rather high concentrations of oxidant and bicarbonate was required.[11 a, b]
As reported recently these limitations could be overcome by the use of PAA.[13] By a detailed
UV/Vis and EPR spectroscopic investigation of the different reaction steps and intermediates, it
was evidenced that the key feature of the MnII ion catalyzed dye degradation by PAA involves a
complex oxidation by PAA vs. reduction by H2O2 (omnipresent in commercial PAA) reaction
sequence in which a steady state equilibrium of the reactive MnIV=O intermediate sets in. The
much higher degradation rate found in the presence of PAA was ascribed to the higher steady
state availability of the reactive MnIV=O compared to that in the use of HCO4- or H2O2/HCO3
-.
Scheme 5.1 Structure of the model substrates Orange II (OII), Tartrazine (TZ)and p-nitrophenol (PNP).
We now report an extension of the MnII/PAA system by application of the simple organic
chelate terpyridine (terpy). The much higher catalytic activity in the presence of the terpyridine
chelate in the Mn catalyzed oxidative degradation by PAA, was studied for various different
organic dyes (Scheme 5.1) to gain deeper insight into the mechanism of action of this intriguing
reaction. Besides Orange II (OII), a commonly used model system, we also chose highly stable dye
substrated as Tartrazine (TZ), which is one of the top three dyes by total sales and widely
employed as food colorant, in textile and paper applications[2] ,and p-nitrophenol (PNP) as
benchmark for the catalytic performance.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
143
(5.1)
(5.2)
5.3 RESULTS AND DISCUSSION
5.3.1 MnII + terpy in solution
In general the tridentate ligand terpyridine is able to form mono (1:1) and bis (1:2) complexes
with various transition metals (Equations 5.1 and 5.2). The equilibrium constant K1 for the
formation of the mono complex has been reported before.[14] We additionally performed a
spectrophotometric titrations under reaction conditions (phosphate buffer, pH 7.0) which reflect
those of the catalytic degradation reaction (results not shown) and the thereby obtained
equilibrium constant K1 = 1.1 · 104 M-1 fits rather well within the literature known range for K1 ((2.5
- 3.0) · 104 M-1)[14]. The fact that K1 appears to be smaller in our case is attributed to the use of the
H2PO4-/HPO4
- buffer, since phosphate anions could also form relatively stable metal-phosphate
complexes causing deviations in the known equilibrium of Equation 5.1. Since a second
coordination (K2) can not be excluded, a typical Job-plot analysis was performed to gain more
information on the distribution of the different species in solution.
The inset in Figure 5.1 exemplarily shows spectral changes before and after mixing of a MnII
solution with terpyridine in a phosphate containing buffer solution. A plot of the maximum
absorbance changes vs. the molar fraction reveals that the maximum is about at a molar fraction
of 0.58, which indicates that in solution a mixture of mono- and bis-terpyridine-MnII complexes is
formed (Figure 5.1). The deviation of the linear behavior at about a molar fraction of 0.5 is
attributed to an underlying local maximum of the mono compound. The lack of labile coordination
sites in a possible bis-MnIIterpy compound would be disadvantageous for the activation of the
oxidant, thus it is assumed that only the in situ generated mono terpyridine complex is relevant to
the catalytic reaction.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
144
Figure 5.1 Absorbance change at two wavelengths for different molar fractions of MnII and terpy at pH 7.0. Inset:
UV/Vis spectral changes before (MnII in water with terpy in buffer/CH3CN) and after mixing at a
([terpy]/([terpy]+[MnII]) of 0.3. Reaction conditions: [Mn
II]+[terpy] = 1 · 10
-4 M, 0.05 M phosphate buffer with 10 %
CH3CN, pH 7.0, 25 °C.
5.3.2 MnIIterpy + PAA in the absence of dye substrate
5.3.2.1 UV/Vis measurements
Upon following the UV/Vis spectral changes of the reaction of MnIIterpy with PAA in the
absence of substrate (Figure 5.2 (A)) a rapid reaction is observed which seems to consist of several
parallel reaction steps. The fact that no isosbestic points are observed already points to the
complexity of this reaction. However, a closer look at the UV/Vis spectra directly after mixing
implies the presence of oxo-bridged high valence manganese species which have been reported
before in the reaction of Mn precursors with oxone (HSO5-).[15] UV/Vis characteristics of both
species, i.e. µ-oxo-bridged Mn2III/IV and Mn2
IV/IV can be found, as is evident from a comparison with
the spectral features of the readily prepared dimers (see Figures 5.3 (A)).[16, 17]
Similar to literature findings, the decomposition of these intermediates to MnVIIO4- (five
characteristic bands at 500-575 nm) was observed [15] after different reaction times depending on
the [PAA] and [Mn] (results not shown). Since the formation of free permanganate involves
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
145
decomplexation of the terpy ligand and no significant amounts of free ligand is found after the
reaction (see 296 nm), it can be assumed that ligand degradation also occurs during the reaction
of MnIIterpy with PAA.
Figure 5.2 (A) UV/Vis spectral changes recorded during the reaction of MnIIterpy (Mn
II:terpy = 1:2) with PAA. Inset:
aborbance vs. time plots recorded at different wavelengths. Reaction conditions: 2 · 10-5
M MnII, 4 · 10
-5 M terpy, 2 ·
10-3
M PAA, 0.05 M phosphate buffer, pH 7.0, 25 °C. (B) Plot of kobs vs. the delay time of the addition of 5 · 10-5
M OII
to a mixture of MnIIterpy and PAA at 5 °C. Inset: aborbance vs. time plot recorded at 336 nm for the same reaction as
in (A) at 5 °C. Reaction conditions: 2 · 10-5
M MnII, 4 · 10
-5 M terpy, 2 · 10
-3 M PAA, 5 · 10
-5 M OII, 0.05 M phosphate
buffer, pH 7.0, 5 °C.
Figure 5.3 (A) UV/Vis spectra of 5 · 10-5
M readily prepared [Mn2III/IV
(µ-O)2terpy2(H2O)2](NO3)3 (—) and [Mn2IV/IV
(µ-
O)2terpy2(SO4)2] (—) in 0.05 M phosphate buffer at pH 7.0 and 25 °C. (B) UV/Vis spectra of the intermediate species
formed during the reaction of MnIIterpy (Mn
II:terpy = 1:2) with PAA, reaction conditions: 2 · 10
-5 Mn
II, 4 · 10
-5 M terpy,
2 · 10-3
M PAA, phosphate buffer, pH 7.0, 25 °C.
(A) (B)
(A) (B)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
146
In order to clarify what the catalytically relevant phases of the reaction are, we performed an
experiment where a dye substrate (OII) is added after different delay times in the reaction of
MnIIterpy with PAA. A comparison of the thereby obtained rates could give some helpful
information on the different reaction phases and their relevance to the catalytic dye degradation
reaction. To slow down the reaction and gain sufficient time for several of these reactivity tests,
we chose a reaction temperature of 5 °C. The results presented in Figure 5.2 (B) show that with
the decomposition of the intermediates of the first reaction phase a concomitant loss in the
catalytic activity towards the dye substrate is observed. This is attributed to the above described
MnVIIO4- formation and ligand degradation, and leads to the suggestion that the first reaction
phase is more important to the studied reaction of the catalytic dye degradation with PAA.
Contrary to our earlier studies on the activation of PAA by MnII ions, we were not able to
resolve the complex behavior in particular during the first phase of the reaction by variation of the
different reaction parameters, i.e. [catalyst], [PAA], [H2O2], pH or temperature (results not shown).
This, unfortunately, made it impossible to obtain stable reaction phases for more detailed
investigations and thereby disclose the mechanism of action of the dimer formation upon reaction
of the MnII precursor with PAA. However, it is well established that the reaction of a subequivalent
of oxidant with a 1:1 mixture of terpy in aqueous solution results in the formation of [Mn2III/IV(µ-
O)2terpy2(H2O)2]3+, in a way that it can be isolated and crystalographically characterized (see
Scheme 5.2).[16]
Scheme 5.2 Structure of the proposed µ-oxo dimers with the Mn oxidation states either III/IV (overall charge 3+) or
IV/IV (overall charge 4+).
Furthermore, the reaction of [Mn2III/IV(µ-O)2terpy2(H2O)2]3+ in the presence of an excess of the
two-electron oxygen-atom donor HSO5- has extensively been studied before and serves as a model
system for the oxygen-evolving complex which catalyzes the conversion of water to dioxygen.[15, 16,
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
147
18, 22] In order to elucidate if there are analogies in the reaction with PAA and HSO5-, the
experiments were repeated with KHSO5/H2O2 under comparable conditions to mimic the presence
of the [H2O2]eq in PAA. The observed UV/Vis spectral changes of the reaction of the MnII precursor
with KHSO5/H2O2 (Figures 5.4 (A) and (B)) clearly resemble the ones observed with PAA as oxidant.
Figure 5.4 (A) UV/Vis spectral changes during the reaction of MnIIterpy (Mn
II:terpy = 1:2) with a HSO5
-/H2O2 mixture
and (B) corresponding absorbance vs. time plots at 336 nm and 410 nm. Reaction conditions: 2 · 10-5
M MnII, 4 · 10
-5 M
terpy, 2 · 10-3
M HSO5-, 2.5 · 10
-4 M H2O2, 0.05 M phosphate buffer, pH 7.0, 25 °C.
The small decay following the first increase in absorbance can be attributed to the
disproportionation of the [H2O2]eq in PAA by the high-valence di-µ-oxo core.[11c, 19, 20] In general,
oxo-bridged Mn2III/IV and Mn2
IV/IV dimers can be considered as oxidized forms of typical catalase
model compounds and are known to cause H2O2 dismutation.[19] Consequently, it is concluded,
that the reaction of the MnII precursor with PAA proceeds in analogy to the oxidation by KHSO5
besides the fact that the reaction is complicated by the [H2O2]eq in PAA. Since it was hardly
possible from the UV/vis data alone to determine whether the dimeric form is predominantly
Mn2III/IV or Mn2
IV/IV under catalytic conditions, further EPR spectroscopic experiments were
performed.
5.3.2.2 EPR measurements
The initial MnII catalyst containing buffer solution without any oxidant present showed an
expected signal for mononuclear MnII centered at g = 2 (see Figure 5.5 (A), first spectrum). The
lack of the typical six-line pattern (I = 5/2, S = 5/2) is attributed to the formation of different
(B) (A)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
148
substituted aqua-species, a mixture of mono- and bis-terpyridine compounds (see Chapter 5.3.1),
and complex-formation with phosphate buffer under the selected experimental conditions. The
second spectrum, taken directly after addition of PAA is shown in Figure 5.5 (B).
Figure 5.5 X-band EPR spectra at 10 K of (A) MnIIterpy (Mn
II:terpy = 1:2) in solution and after mixing with PAA at
different time intervals. (B) Amplification of the first spectrum in the presence of oxidant at approximately 15 s
reaction time. Reaction conditions: 1 · 10-4
M MnII, 2 · 10
-4 M terpy, 0.01 M PAA, 0.05 M phosphate buffer, pH 7.0, 25
°C. EPR conditions: 9.45 GHz, 1 mW microwave power, modulation amplitude 400 mT.
The drastic decrease in the overall spectral features compared to the spectrum in the absence
of oxidant emphasizes the fact that the main species in solution under catalytic conditions is EPR
silent. On the other hand a small amount of a 16-line signal at g = 2 is observed (Figure 5.5 B,
amplification of the second spectrum recorded 15 s after mixing of the reactants). These spectral
features are characteristic for an antiferromagnetically exchange-coupled Mn2III/IV dimer.[21] An
estimated quantification implies that µ-oxo bridged Mn2III/IV dimer, content is approximately ≤ 10
%, whereas the remaining ≈ 90 % are EPR inactive. On the basis of the UV/Vis spectroscopic results
this EPR silent form is assumed to be mostly a Mn2IV/IV dimer which is in fact in very good
agreement with the detailed mechanistic investigations of the intermediates under catalytic
conditions for the [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/ HSO5
- system by Brudvig/Crabtree et al., who
evidenced that the predominant species in the catalytic solution is the EPR silent Mn2IV/IV dimer.[22]
(B) (A)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
149
5.3.3 MnIIterpy catalyzed dye degradation with PAA
5.3.3.1 General observations
In order to study the catalytic dye degradation by PAA, a series of measurements were
performed at pH 7.0 and 25 °C. Figures 5.6 (A), (B) and (C) depict the UV/Vis spectral changes
during the catalytic degradation reaction and the concomitant absorbance changes at λmax for the
different model substrates: OII (λmax = 484 nm), TZ (λmax = 420 nm) and PNP (λmax = 400 nm).
Figure 5.6 UV/Vis spectral changes observed for the catalyzed oxidative degradation of OII (A), TZ (B) and PNP (C) by
PAA, insets show the absorbance vs. time plots at the corresponding λmax of the dye. Reaction conditions: (A) and (B) 1
· 10-6
M MnII, 2 · 10
-6 M terpy, 5 · 10
-3 M PAA, 5 · 10
-5 M dye, 0.05 M buffer, pH 7.0, 25 °C; (C) 5 · 10
-6 M Mn
II, 1 · 10
-5 M
terpy, 5 · 10-3
M PAA, 5 · 10-5
M dye, 0.05 M buffer, pH 7.0, 25 °C.
(A) (B)
(C)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
150
When the results for TZ are compared to the use of chlorine or hypochlorite in basic media,
which is a well known technology for bleaching dyes,[2b, 23] it is obvious that under these reaction
conditions the bleaching reaction is not only limited to the cleavage of the azo-linkage (420 nm for
TZ) as evidenced by the UV/Vis absorbance after 500 s. Thus, it is concluded that the more stable
aromatic rings and primary dye destruction products which result from the azo-cleavage are also
extensively degraded.[2b, 24] These results highlight the fact that at least in the studied reactions the
absorbance decrease of the dye at λmax provides a satisfactory basis for the evaluation of the
performance of the dye treatment method. In addition, PNP, a highly stable phenolic model
substrate, is efficiently degraded by only minor amounts of oxidant and catalyst (see Figure 5.6).
Based on the excellent catalytic activity of the MnIIterpy/PAA system at lower catalyst and oxidant
concentrations compared to the simple MnII salt,[13] the reaction conditions were refined to
evaluate its putative applicability in the catalytic degradation of various organic dyes with different
structural motives by determining the catalytic turnover frequency. Table 5.1 contains the
calculated turnover frequencies, i.e. mol of dye oxidized by mol of catalyst per hour.
Dye substrate OII TZ PNP
Initial rate (Ms-1) 9,92 · 10-6 5,74 · 10-6 4.43 · 10-7
TOF (h-1) 7142 4133 319
Table 5.1 Summary of initial rates and turn over frequencies for the degradation reactions of OII, TZ and PNP
determined at the corresponding λmax of the dye. Reaction conditions: 5 · 10-6
M MnII, 1 · 10
-5 M terpy, 5 · 10
-5 M dye
substrate, 5 · 10-3
M PAA, 0.05 M phosphate buffer, pH 7.0, 25 °C.
Numerous control experiments were performed (results not shown) to exclude different
influencing factors on the overall dye degradation catalysis, viz. the spontaneous, non-catalyzed
reaction of PAA with the dye substrates, a possible contribution of free radicals, acetate or
phosphate buffer concentration. In organic buffers suitable for this pH range such as MES and Bis-
Tris no reaction with dye substrate was observed (results not shown). This is probably due to the
fact that these organic molecules could serve as substrate to the highly reactive Mn intermediates
formed during the reaction so that the dye degradation is suppressed by the excess content of
buffer required for reliable pH jump experiments.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
151
Apart from the fact that the catalytic degradation of the studied organic dyes occurred
efficiently, the optimization of the process by analyzing the effect of several parameters such as
pH, [oxidant] and [catalyst], was investigated in more detail.
5.3.3.2 Kinetics of the MnIIterpy catalyzed dye degradation by PAA
pH dependence. The kinetics of the manganese catalyzed oxidative degradation of the selected
dyes was studied in 0.05 M buffer (phosphate, acetate and bicarbonate), in a pH range from 2.0 to
10.0 at 25 °C. The decomposition of the dyes was followed by monitoring the absorbance change
at the characteristic λmax of each dye substrate. Figure 5.7 shows the changes of the
experimentally determined observed first order rate constants for the MnIIterpy catalyzed
degradation for the different substrates as a function pH.
Figure 5.7 Plots of the observed first order rate constants for the degradation of OII (—), TZ (—) and PNP (—) as a
function of pH determined at the corresponding λmax. Reaction conditions: 5 · 10-6
M MnII, 1 · 10
-5 M terpy, 5 · 10
-3 M
PAA, 5 · 10-5
M dye substrate, 0.05 M acetate or phosphate buffer, 25 °C.
For all systems the rate of oxidation increases with increasing pH and goes through a
maximum at a pH between 6.5 and 7.0, suggesting that the same reactive manganese species and
the same in situ formed oxidizing agent is responsible for the decomposition of the various
substrates. The decrease in oxidation rate at higher pH is partly due to the subsequent formation
of Mn(OH)2 precipitates at higher pH, which negatively affect the stability of the MnIIterpy
precursor and thereby the amount of available catalyst. The deviation observed at a pH of about
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
152
4.5 is attributet to the pKa of the terpyridine ligand (pKa = 4.7),[25] which interferes in the
equilibrium described in Equation 5.1 and thereby lowers the available precursor concentration
for the catalytic reaction.
Effect of [terpy] on the catalytic oxidative degradation of various dyes. Since the complex-
formation constants of MnII and terpyridine are known to be relatively unfavourable, the dye
degradation reaction was studied as a function of the terpyridine ligand concentration for OII, TZ
and PNP (Figure 5.8).
Figure 5.8 Plots of the observed first order rate constants for the degradation of OII (—), TZ (—) and PNP (—) as a
function of the terpy:Mn ratio determined at the corresponding λmax. Reaction conditions: 5 · 10-6
M MnII, 1 · 10
-3 M
PAA, 5 · 10-5
M dye substrate, 0.05 M phosphate buffer, pH 7.0, 25 °C.
The results clearly show that complexation between the σ-donor and π-acceptor terpyridine
as ligand and the metal, is indispensable for efficient catalytic performance. With increasing ligand
concentration the maximum rate increases and reaches limiting values at higher excess
concentrations of terpyridine ligand for all the studied dye substrates.
For this result two possible explanations are plausible. On the one hand the low binding
constant of terpyridine to manganese may require higher [ligand] to assure complex-formation
between terpyridine and manganese. On the other hand, the UV/Vis spectral changes during the
catalytic degradation reaction imply that partial ligand degradation also occurs, so that a higher
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
153
[terpy] is required for efficient catalytic performance. In any case, all further measurements were
performed at a MnII:terpy ration of 1:2.
Critical role of [H2O2]eq. Previous results established that the [H2O2]eq plays an important role in
the overall catalytic oxidation reaction by PAA (commercially available PAA always contains
H2O2).[13] It was evidenced that a low [H2O2]eq was required to avoid MnIVO2 precipitaition and
over-oxidation to permanganate. Moreover, in the presence of excess [PAA] the catalytic system
could be reactivated by only minor amounts of [H2O2]. However, larger [H2O2]eq negatively
affected the catalytic performance by delaying the formation of active Mn=O intermediates.
Figure 5.9 Plots of the inverse observed first order rate constants for the degradation of OII (—), TZ (—) and PNP (—)
as a function of the additional [H2O2] determined at the corresponding λmax. Reaction conditions: 5 · 10-6
M MnII, 1 · 10
-
5 M terpy, 5 · 10
-3 M PAA, 5 · 10
-5 M dye substrate, 0.05 M phosphate buffer, pH 7.0, 25 °C.
In the case of the MnIIterpy catalyzed degradation of dyes, no reactivation of the catalytic
system by addition of small aliquots of H2O2 was found. The influence of the [H2O2]eq on the
catalytic degradation reaction was investigated by performing the reaction in the presence of an
additional amount of H2O2 in the concentration range from 0.25 · 10-3 to 5 · 10-3 M (equal to
[PAA]). It shows that a higher [H2O2]eq inhibits the reaction. If the corresponding inverse observed
rate constant is plotted against the additional [H2O2] for OII, TZ and PNP (Figure 5.9), a linear
dependence with no intercept is observed. The slopes of these linear fits were found to be 36.4 ∙
103 s M-1 for OII, 61.2 ∙ 103 s M-1 for TZ and 162.7 ∙ 103 s M-1 for PNP, respectively. To account for
this observation it is suggested that H2O2 reacts with the rapidly formed oxo-bridged dimeric
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
154
intermediates (Mn2III/IV and Mn2
IV/IV) in an unproductive parallel catalase-like pathway during
which oxygen is formed and the dimers are reduced or partially cleaved. Similar observations have
been reported before.[11c] This represents a kind of back reaction that depending on the [H2O2]eq
lowers the steady state availability of the dimeric, oxo-bridged species, which in turn emphasizes
their role as key intermediates in the MnIIterpy catalyzed dye degradation with PAA.
Effect of [PAA] on the MnIIterpy catalyzed oxidative degradation of various dyes. In order to
investigate the effect of the oxidant concentration on the MnIIterpy catalyzed degradation of the
different dyes, the [PAA] was varied from (0.5 - 15) · 10-3 M at a catalyst concentration of 5 · 10-6
M. Figure 5.10 shows the plot of the observed first order rate constants as a function of [PAA] for
Orange II, TZ and PNP.
Figure 5.10 Plots of the observed first order rate constants for the degradation of OII (—), TZ (—) and PNP (—) as a
function of the [PAA] determined at the corresponding λmax. Reaction conditions: 5 · 10-6
M MnII, 1 · 10
-5 M terpy, 5 ·
10-5
M dye substrate, 0.05 M phosphate buffer, pH 7.0, 25 °C.
Despite of the difference in reactivity towards the different dyes, in all cases the system
reaches limiting values at higher [PAA]. Although this saturation behaviour is in line with Michelis-
Mentin kinetics, which involves a pre-equilibrium reaction prior to the rate-limiting step of the
studied reaction, it has to be considered in light of the findings concerning the effect of [H2O2]eq,
that it might rather be related to the equilibrium content of H2O2 described above. With increasing
[PAA] the omnipresent [H2O2]eq also accelerates the back reaction which in turn slows down the
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
155
observed dye degradation rates, so that the observed rate constants for the overall degradation
reaction only show an apparent saturation effect.
Effect of [MnIIterpy] on the catalytic oxidative degradation of various dyes. To evaluate the
effect of the catalyst concentration on the oxidative degradation of different organic dyes by PAA,
kinetic studies were performed under variation of the MnIIterpy pre-catalyst concentration.
Surprisingly, in the present case, the catalytic reaction leads to a square dependence of kobs on the
[MnIIterpy], which implies a possible involvement of two equivalents of the MnIIterpy precatalyst
in the overall catalytic reaction as expected when Mn-dimers participate. The plots of the
corresponding observed first order rate constants vs. [catalyst]2 for OII, TZ and PNP (for two
different oxidant concentrations) are presented in Figures 5.11 (A) and (B), respectively. From the
slope of these dependencies the third order rate constants for all substrates were calcultated and
are summarized in Table 5.2.
Figure 5.11 Plots of the observed first order rate constants for the degradation of OII (—), TZ (—) and PNP (—) as a
function of [catalyst]2 determined at the corresponding λmax for (A) 1 · 10
-3 M PAA and (B) 5 · 10
-3 M PAA. Reaction
conditions: MnII:terpy = 1:2, 5 · 10
-5 M dye substrate, 0.05 M phosphate buffer, pH 7.0, 25 °C.
OII TZ PNP
k (M-2s-1) for 1 · 10-3 M PAA 3.74 · 109 2.31 · 109 6.23 · 107
k (M-2s-1) for 5 · 10-3 M PAA 7.25 · 109 4.36 · 109 11.65 · 107
Table 5.2 Third order rate constants for two different [PAA] calculated from the linear fit of the observed first order
rate constants in Figures 5.11 (A) and (B).
(A) (B)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
156
Since the observed second order dependence highlights again the perticipation of dimeric
manganese species in the overall catalytic degradation reaction, we decided to additionally repeat
some experiments for the readily prepared oxo-bridged Mn2III/IV and Mn2
IV/IV dimers as catalysts.
5.3.3.3 Readily prepared dimers in the catalytic dye degradation by PAA
A direct comparison of the catalytic reactivity at indentical manganese content revealed a
slightly slower degradation rate when the dimers are used (Figure 5.12 (A)).
Figure 5.12 (A) Comparison of the depedencies of the observed first order rate constants for the degradation of OII on
the [catalyst]2 for monomeric Mn
IIterpy (―), dimeric [Mn2
III/IV(µ-O)2terpy2(H2O)2](NO3)3 (―) and Mn2
IV/IV(µ-
O)2terpy2(SO4)2 (―) at identical [total Mn]. Reaction conditions: 5 · 10-3
M PAA, 5 · 10-5
M OII, 0.05 M phosphate
buffer, pH 7.0, 25 °C. (B) Comparison of the traces recorded at 484 nm that accompany the catalytic OII degradation of
(―)1 · 10-6
M MnII + 2 · 10
-6 M terpy and (―) 0.5 µM [Mn2
III/IV(µ-O)2terpy2(H2O)2](NO3)3 + 1 · 10
-6 M terpy. Reaction
conditions: 5 · 10-3
M PAA, 5 · 10-5
M OII, 0.05 M phosphate buffer, pH 7.0, 25 °C.
This observation can be explained by the effect of the two-fold Mn:Ligand ratio described
above. Whereas all prior experiments were performed with a excess of terpyridine ligand to
assure the complexation despite the low binding constant and occurring ligand degradation, the
dimers contain only a 1:1 ratio of Mn:terpy, which accounts for the slightly lower activity. In fact, if
a comparison is made between the monomeric and the dimeric starting compound with the same
amount of terpyridine ligand present, no difference in the catalytic performance is observed
(Figure 5.12 (B)).
(B) (A)
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
157
In general, oxo-bridged manganese complexes are often involved as key intermediates in
several catalytic oxygenation processes such as photosynthetic water oxidation,[16, 26, 27]
bleaching[28], epoxidation[29] or oxidation of hydrocarbons[30] and alcohols[31]. However, it is known
from literature that these species are not prevalent in catalytic solution when H2O2 is used as
oxygen source.[11c, 32] Moreover, the lability of the [Mn2III/IV(µ-O)2terpy2(H2O)2]3+ compound has
been shown before.[11c, 33] Presumably it is rapidly converted to a mononuclear manganese species
by reduction to lower-valence bridged dimanganese species and eventually concomitant cleavage
of the bis-µ-oxo motive by H2O2.[11c, 19, 20] This could also account for the observation that no
significant difference is found, whether the MnII starting compound or the readily prepared
Mn2III/IV or Mn2
IV/IV µ-oxo dimers are used in the catalytic dye degradation with PAA, since the oxo-
bridged core of the dimers is reduced or partially cleaved by the [H2O2]eq.
Figure 5.13 Comparison of the absorbance vs. time plots at 484 nm in the stoichiometric reaction of 5 · 10-5
M
[Mn2III/IV
(µ-O)2terpy2(H2O)2](NO3)3 (―) and 5 · 10-5
M Mn2IV/IV
(µ-O)2terpy2(SO4)2 (―) with 5 · 10-5
M OII. Reaction
conditions: 0.05 M phosphate buffer, pH 7.0, 25 °C.
However, also in the absence of oxidant the readily prepared Mn2III/IV or Mn2
IV/IV dimers react
with the dye substrate. Figure 5.13 shows the absorbance vs. time plots recordedat the dye’s λmax
upon reaction with a stoichiometric amount of [Mn2III/IV(µ-O)2terpy2(H2O)2](NO3)3 and Mn2
IV/IV(µ-
O)2terpy2(SO4)2. Both dimers show rather sluggish activity towards the dye substrate. Whereas in
the reaction with the Mn2III/IV dimer only minor changes are observed, addition of Mn2
IVIV causes a
rapid decrease of the chacracteristic absorbance. It is clear that a stoichiometric amount of
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
158
Mn2III/IV and Mn2
IV/V yields incomplete dye degradation, which is reflected by the fact that only 15
to 40 % of the azo-bond absorbance band is decreased.
As a consequence, it has to be considered that both in situ formed dimmer species contribute
at least partially to the observed reactivity in the overall dye degradation. However, this also
implies that there is another, more reactive intermediate formed.
The fact that [Mn2III/IV(µ-O)2terpy2(H2O)2]3+ with its easily exchangeable water ligands is rapidly
formed when an excess of PAA oxidant is present, gives rise to further potential reactive “Mn=O”
intermediates. These could involve the potential dimeric species with a MnV=O or MnIV-O• moiety
as postulated before.[23, 27, 34] Since it is known that the latter reaction steps to even higher valence
species than Mn2III/IV and Mn2
IVIV, ´which afford the labile terminal coordination sites occupied by
two water ligands in the terpy system[35] (see Scheme 5.2), a comparison of the reactivity pattern
of the corresponding bidentate bpy or phen ligand systems could reveal some helpful information
on further putative catalytically relevant species. Both oxo-bridged dimers of the bpy and phen
(metal:bpy/phen ligand ratio in complex = 2:1) lack the essential feature of the terminal water
ligands. These labile coordination sites of the terpyridine compound are believed to result in
enhanced lability of di-µ-oxo dimer and give rise to reaction pathways involving the potential high-
valence MnV=O or MnIV-O• intermediates.[35]
Some preliminary eperiments with the use of phen instead of terpy already indicate that this
motive might be relevant. Figure 5.14 shows the UV/Vis spectral changes at λmax during the
catalytic degradation reaction of Orange II in the presence phen. Although two different
MnII:ligand ratios were used to exlude the interference of the MnII:ligand precursor equilibrium
(related to Equations 5.1 and 5.2), it is evident that the so-formed potential dimeric phen species
are much less reactive.
Similar observations on the influence of the terminal binding site on the rate of oxygen
evolution from the [Mn2III/IVO2]3+/oxidant have been reported before.[35] It is concluded that the
lack of the labile coordination site in the corresponding in situ formed bpy and phen dimers is
responible for the decrease in the catalytic dye degradation activity. Hence, we suggest that the
MnIIterpy + PAA system indeed comprises parallels to the intermediates reported for the water
oxidizing model system [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
- and that therefore the postulated
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
159
MnV=O or MnIV-O• containing intermediates might account for the extraordinarily high activity in
the catalytic dye degradation with PAA.
Figure 5.14 UV/vis spectral changes at 484 nm that accompany the catalytic degradation of OII by PAA with different
MnIIphen precursors (Mn
II:phen = 1:20 (―) and Mn
IIphen = 1:4 (―)). Reaction conditions: 1 · 10
-6 M Mn
II, 2 · 10
-5 M
phen (―), 4 · 10-6
M phen (―), 5 · 10-5
M OII, 5 · 10-3
M PAA, 0.05 M phosphate buffer, pH 7.0, 25 °C.
Although the exact role of the high-valence Mn2III/IV and Mn2
IV/IV dimers or higher oxidized
dimers with MnV=O or MnIV-O• moieties in this efficient catalytic degradation reaction remains to
be clarified, it has to be considered that there are several catalytically relevant species formed in
the reaction with MnIIterpy and PAA that contribute to the observed overall dye degradation.
5.3.4 Mechanistic implications
Although the existence of free coordination sites in the precursor complex is considered to be
of general importance for the catalytic activity, our experiments on the MnIIterpy catalyzed dye
degradation by PAA point to the necessity of an excess [terpy], which in turn may favour the
formation of a coordinatively saturated bis-terpy compound. This seemingly contradictive
behaviour can be accounted for under consideration of two phenomena: On one hand, the
binding constant of terpy with MnII is not ideal, especially in phosphate containing solution it
probably interferes in the complexation equilibrium of the MnIIterpy precursor by formation of
Mn-phosphate complexes and thereby lowers the available [pre-catalyst]. On the other side,
ligand degradation also occurs during the reaction with oxidant, which results in the formation of
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
160
free permanganate. Consequently, an excess ligand is required to suppress the effect of the
decreasing ligand concentration.
Since over the applied pH range no catalytic degradation reaction with MnII and PAA is
observed in the absence of the chelate ligand due to insufficient interaction between the
reactants, it is concluded that the complex-formation by terpyridine is indispensable.
Furthermore, it is suggested that a reaction with PAA in its protonated form (pH 7) is in the first
place enabled by the strong π-acceptor properties of the terpyridine ligand and the proximity to
the pKa of PAA.
The UV/Vis spectral changes in the absence of dye imply a rapid formation of µ-oxo-bridged
high-valence Mn2III/IV and Mn2
IV/IV dimers. This particularly complex reaction behaviour (see
absorbance vs. time plots) could be satisfactorily copied by the use of HSO5-/H2O2 under conditions
comparable to those used in the PAA system, which emphasizes the role of the [H2O2]eq
omnipresent in PAA. The [H2O2]eq gives rise to a potential back reaction with the rapidly formed
high-valence Mn-intermediates, which thereby lowers the catalytic dye degradation activity. In
general, oxo-bridged Mn2III/IV and Mn2
IV/IV dimers can be considered as oxidized forms of typical
catalase model compounds and are known to cause H2O2 dismutation.[11c, 19, 20] The involvement of
dimeric manganese intermediates is in addition supported by EPR spectroscopic measurements
and the second-order dependence of the catalytic dye degradation on the [catalyst]. Although the
exact mechanism of action of dimer formation remains to be clarified, it is well known that the
synthesis and isolation of [Mn2III/IV(µ-O)2terpy2(H2O)2]3+ is already achieved by a subequivalent of
oxidant.[16] On the other hand, numerous reports on the reaction between the [Mn2III/IV(µ-
O)2terpy2(H2O)2]3+ and an excess HSO5- oxidant revealed by the use of various spectroscopic
methods that the major species under catalytic conditions is an Mn2IV/IV dimer.[22]
MnIIterpy shows comparable reaction behavior with PAA as well as with HSO5-/H2O2, and our
own EPR spectroscopic measurements of the intermediates formed from MnIIterpy and PAA
confirmed the presence of Mn2III/IV and mainly Mn2
IV/IV dimers. Therefore, it can be concluded that
there are fundamental parallels to the intermediates formed during the well studied reaction of
[Mn2III/IV(µ-O)2terpy2(H2O)2]3+. Brudvig/Crabtree et. al. investigated the conversion of a [MnIII(μ-
O)2MnIV]3+ complex to its one-electron-oxidized product, [MnIV(μ-O)2MnIV]4+, by the two-electron
oxidant HSO5-. According to literature reports, a simplified mechanistic sequence to account for
these putative intermediates is outlined in Scheme 5.3.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
161
This oxidation is believed to pass through an intermediate two-electron-oxidized form of the
[Mn2III/IV(µ-O)2terpy2(H2O)2]3+, with either a MnV=O or a MnIV-O• moiety, which through
disproportionation with [MnIII(μ-O)2MnIV]3+ mainly yields the corresponding Mn2IV/IV dimer.[34, 35]
The binding site for this second oxidation step is supposed the MnIII ion since it is expected to
show comparable or even higher substitution rates than the MnIV binding site, due to the
enhanced lability in the Jahn-Teller distorted MnIII ion.[34]
Scheme 5.3 Proposed mechanism for the reaction between [Mn2III/IV
(µ-O)2terpy2(H2O)2]3+
and oxygen-atom transfer
reagents (XO) acting as a water-oxidation catalyst according to references [27] and [34].
Although in the absence of oxidant, both dimers, viz. [Mn2III/IVO2terpy2(H2O)2](NO3)3 and
Mn2IV/IV(µ-O)2terpy2(SO4)2, show slight activity toward the dye substrate, it has to be considered
that the excess of oxidant present allows further oxidation steps as in [Mn2III/IV(µ-
O)2terpy2(H2O)2]3+ / HSO5- (outlined in Scheme 5.3). However, to what extend the differént in situ
formed µ-oxo dimers Mn2III/IV and Mn2
IV/IV, respectively, or the postulated reactive intermediates
and mechanistic ideas of the water oxidizing model system [Mn2III/IV(µ-O)2TERPY2(H2O)2]3+/HSO5
-
apply to the observed dye degradation reactivity, remains to be clarified.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
162
5.4 CONCLUSIONS
In summary, we presented an extraordinarily efficient method for the catalytic oxidative
degradation of different highly stable organic dyes, such as Orange II, Tartrazine and p-nitrophenol
as reflected by the high catalytic turnover frequencies. Screening and spectroscopic methods
allowed us to study the catalytic reaction course and to shed light on some key mechanistic
features. The second order dependence of the observed first order rate constants for the overall
dye degradation on the pre-catalyst concentration implies the involvement of dimeric
intermediates in the catalytic reaction. The omnipresent [H2O2]eq in the PAA stock solution was
shown to result in an inverse dependence on the catalytic reaction, most likely by causing a
catalase-like H2O2 decomposition, which also becomes noticeable at higher [PAA]. Consequently,
the saturation kinetics observed for higher oxidant concentrations can be ascribed to the thereby
increased [H2O2]eq.
UV/Vis spectral and EPR spectroscopic measurements in the absence of dye substrate
confirmed that the intermediates formed under catalytic conditions and in the absence of
substrate mainly consist of dimeric species with a [Mn2III/IV(µ-O)2]3+- and [Mn2
IV/IV(µ-O)2]4+-core.
These results strongly resemble earlier findings on the water oxidizing model system [Mn2III/IV(µ-
O)2terpy2(H2O)2]3+/HSO5-. Both these dimers show rather moderate activity towards the dye
substrate. Preliminary experiments were performed with another ligand system than terpy that
does not allow the formation of dimers with a labile coordination site like [Mn2III/IV(µ-
O)2terpy2(H2O)2]3+. These results already imply that the latter dimer might be a pre-stage to even
higher oxidized dimeric species, which could be relevant to the extraordinary performance in the
MnIIterpy catalyzed dye degradation. We are currently trying to clarify to what extend the
different in situ formed µ-oxo dimers and the mechanistic ideas of the water oxidizing model
system [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
- and its potential dimeric intermediates with either a
MnV=O or a MnIV-O• moiety, account for the observed reactivity in the MnIIterpy catalyzed dye
degradation with PAA.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
163
5.5 EXPERIMENTAL SECTION
CHEMICALS. Orange II, certified [Acid Orange 7, C.I. 15510, sodium 4-(2-hydroxy-1-
naphthylazo)benzenesulfonate], 99 % was supplied by Sigma-Aldrich and recrystallised from a
Et2O/H2O mixture at 4 °C. Peracetic acid 39 wt. %, H2O2 30 %, as well as 2,2’:6’,2’’-terpyridine,
1,10-phenanthroline, 2-bis(2-hydroxyethyl)amino-2-(hydroxymethyl)-1,3-propanediol] and 2-(N-
morpholino)ethanesulfonic acid were of analytical grade and provided by Sigma-Aldrich. The
synthesis of [Mn2III/IVO2terpy2(H2O)2](NO3)3 · 6H2O was performed according to literature.[16]
[Mn2IV/IVO2terpy2(SO4)2] · 6H2O was synthesized as described before and crystallographically
characterized (Crystal data: monoclinic, space group: C2/c, a(Å) = 27.060, b(Å) = 9.020, c(Å) =
18.236, V(Å3) = 3539.29, α(°) = 90, β(°) = 127.331, γ(°) = 90, Z = 4.).[26a] All other chemicals were
commercially available (Acros Organics, Sigma-Aldrich) and were used without any further
purification.
GENERAL PROCEDURE AND PH JUMP TECHNIQUE. The manganese terpyridine compound was freshly
produced by dissolving the appropriate ratio of MnII:ligand in water with 10 % CH3CN before use
(MnII:terpy = 1:2). Stock solutions of [Mn2III/IVO2terpy2(H2O)2](NO3)3 and [Mn2
IV/IVO2terpy2(SO4)2]
were prepared in water and used immediately. All solutions were prepared in Millipore Milli-Q
purified water. To a freshly prepared 0.05 M sodium phosphate solution, an adequate amount of
NaOH was added to adjust the pH in a way that the subsequent addition of a specific
concentration of PAA gave the desired pH within less than one minute. The reaction was started
under isothermal conditions by addition of small aliquots of a concentrated manganese stock
solution together with Orange II to the PAA containing buffer solution. The catalytic reaction was
followed by in situ UV/Vis spectroscopy.
INSTRUMENTATION AND EQUIPMENT. All kinetic data were obtained by recording time-resolved UV-Vis
spectra using a Hellma 661.502 – QX quartz Suprasil immersion probe attached via optical cables
to a 150 W Xe lamp and a multi-wavelength J & M detector, which records complete absorption
spectra at constant time intervals. All kinetic measurements were carried out under pseudo-first-
order conditions. The pH of the PAA containing aqueous carbonate solution was carefully
measured and adjusted using a Mettler Delta 350 pH meter previously calibrated with standard
buffer solutions at two different pH values (4 and 10). The kinetics of the degradation reaction was
monitored at 484 nm. First order rate constants, where possible, were calculated using Specfit/32
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
164
and Origin (version 7.5) software. Perpendicular mode EPR spectra were recorded on an X-band
Joel Jes Fa 200 spectrometer equipped with a cylindrical mode cavity and a liquid helium cryostat.
Samples were taken from the investigated solutions and immediately frozen to quench the
reaction. The EPR measurements were performed in quartz tubes at 10 K (9.45 GHz, 1 mW
microwave power). Data analyses were done with the Jes-Fa Series software package.
5.6 REFERENCES AND NOTES
[1] (a) Park, H.; Choi, W. J. Photochem. Photobiol. A 2003, 159, 241-247; (b) Brown, D., Laboureur
P. Chemosphere 1987, 12, 397-404; (b) Goszczynski, S.; Paszczynki, A.; Pasti-Grigsby, M. B.;
Crawford, R. L.; Crawford, D. L. J. Bacteriol. 1994, 176, 1339-1347; (c) Chivukula, M.; Spadaro,
J. T.; Renganathan, V. Biochem. 1995, 34, 7765-7772; (d) Spadaro, J. T.; Renganathan, V. Arch.
Biochem. Biophys. 1994, 312, 301-307; (e) Spadaro, J. T.; Isabelle, L.; Renganathan, V. Environ
Sci. Tech. 1994, 28, 1389-1393.
[2] (a) Zollinger, H. Color Chemistry, 3rd ed., Wiley-VCH: Weinheim, Germany, 2003; (b) Beach, E.
S.; Malecky, R. T.; Gil, R. R.; Horwitz, C. P.; Collins; T. J. Catal. Sci. Technol. 2011, 1, 437-443.
[3] Baughman, G. L.; Weber, E. J. Environ. Sci. Technol. 1994, 28, 267-276.
[4] (a) Ellis, W. C.; Tran, C. T.; Roy, R.; Rusten, M.; Fischer, A.; Ryabov, A. D.; Blumberg, B.; Collins,
T. J. J. Am. Chem. Soc. 2010, 132, 9774-9781. (b) Kluson, P.; Drobek, M.; Krejcikova, S.; Krysa,
J.; Kalaji, A.; Cajthaml, T.; Rakusan, J. Appl. Catal. B 2008, 80, 321-326. (c) Wieprecht, T.; Xia,
J.; Heinz, U.; Dannacher, J.; Schlingloff, G. J. Mol. Catal. A 2003, 203, 113-128. (d) Ramirez, J.
H.; Maldonado-Hódar, F. J.; Pérez-Cadenas, A. F.; Moreno-Castilla, C.; Costa, C. A.; Madeira, L.
M. Appl. Catal. B 2007, 75, 312-323. (e) Hage, R.; Iburg, J. E.; Kerschner, J.; Koek, J. H.; M.
Lempers, E. L.; Martens, R. J.; Racherla, U. S.; Russell, S. W.; Swarthoff, T.; van Vliet, M. R. P.;
Warnaar, J. B.; van der Wolf, L.; Krijnen, B. Nature 1994, 369, 637-639.
[5] (a) Oakes, J.; Gratton, P. J. Chem. Soc., Perkin Trans. 2 1998, 1857-1864. (b) Nadtochenko, V.;
Kiwi, J. J. Chem. Soc., Faraday Trans. 1997, 93, 2373-2378. (c) Oakes, J.; Gratton, P; Weil, I. J.
Chem. Soc., Dalton Trans. 1997, 3805-3809. (d) Chahbane, N.; Popescu, D.-L.; Mitchell, D. A.;
Chanda, A.; Lenoir, D.; Ryabov, A. D.; Schramm, K.-W.; Collins, T. J. Green Chem. 2007, 9, 49-
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
165
57. (e) Gilbert, B. C.; Lindsay-Smith, J. R.; Newton, M. S.; Oakes, J.; Ponts i Parts, R. Org.
Biomol. Chem. 2003, 1, 1568-1577.
[6] (a) Song, Y. J.; Lee, S. H.; Park, H. M.; Kim, S. H.; Goo, H. G.; Eom, G. H.; Lee, J. H.; Lah, M. S.;
Kim,Y.; Kim, S.-J.; Lee, J. E.; Lee, H.-I.; Kim, C. Chem. Eur. J. 2011, 17, 7336-7344. (b) Song, F.;
Wang, C.; Falkowski, J. M.; Ma, L.; Lin, W. J. Am. Chem. Soc. 2010, 132, 15390-15398. (c)
Palucki, M.; Finney, N. S.; Pospisil, P. J.; Guler, M. L.; Ishida, T.; Jacobsen, E. N. J. Am.Chem.
Soc. 1998, 120, 948-954.
[7] (a) Arunkumar, C.; Lee, Y.-M.; Lee, J. Y.; Fukuzumi, S.; Nam, W. Chem. Eur. J. 2009, 15, 11482-
11489. (b) Crestoni, M. E.; Fornarini, S.; Lanucara, F. Chem. Eur. J. 2009, 15, 7863-7866. (c)
Song, W. J.; Seo, M. S.; DeBeer George, S.; Ohta, T.; Song, R.; Kang, M.-J.; Tosha, T.; Kitagawa,
T.; Solomon, E. I.; Nam, W. J. Am. Chem. Soc. 2007, 129, 1268-1277. (d) Rose, E.; Andrioletti,
B.; Zrig, S.; Quelquejeu-Etheve, M. Chem. Soc. Rev. 2005, 34, 573-583.
[8] (a) Wieghardt, K.; Schmidt, W.; Nuber, B.; Weiss, J. Chem. Ber. 1979, 112, 2220-2230; (b)
Romakh, V. B.; Therrien, B.; Süss-Fink, G.; Shul’pin, G. B. Inorg. Chem. 2007, 46, 1315-1331; (c)
Mandelli, D.; Kozlov, Y. N.; Golfeto, C. C.; Shul'pin, G. B. Catal. Lett. 2007, 118, 22-29; (d)
Sibbons, K. F.; Shastri, K.; Watkinson, M.; Dalton Trans., 2006, 645-661; (e) Gilbert, B. C.;
Lindsay-Smith, J. R.; Mairata y Payeras, A.; Oakes, J. Org. Biomol. Chem. 2004, 2, 1176-1180;
(f) Bennur, T. H.; Srinivas, D.; Sivasanker, S.; Puranik, V. G. J. Mol. Cat. A , 2004, 219, 209-216;.
(g) Hage, R.; Kerschner, J. Trends in Inorg. Chem. 1998, 5, 145-159.
[9] De Boer, J. W.; Browne, W. R.; Brinksma, J.; Alsters, P. L.; Hage, R.; Feringa, B. L. Inorg. Chem.
2007, 46, 6353-6372.
[10] (a) Shilov, A. E.; Shul’pin, G. B.; Chem. Rev. 1997, 97, 2879-2932; (b) Sheldon, R. A. J. Chem.
Technol. Biotechnol. 1997, 68, 381-388.
[11] (a) Ember, E.; Rothbart, S.; Puchta, R.; van Eldik, R. New J. Chem. 2009, 33, 34-49; (b) Ember,
E.; Gazzaz, H. A.; Rothbart, S.; Puchta, R.; van Eldik, R. Appl. Catal. B. 2009, 95, 179-191; (c)
Rothbart, S.; Ember, E.; van Eldik, R. Dalton Trans. 2010, 39, 3264-3272.
[12] (a) Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett, D. A. J. Am. Chem. Soc. 2000, 122, 1729-
1739. (b) Bakhmutova-Albert, E. V.; Yao, H.; Denevan, D. E.; Richardson, D. E. Inorg. Chem.
2010, 49, 11287-11296.
[13] Rothbart, S.; Ember, E. E.; van Eldik, R. New J. Chem. 2012, DOI: 10.1039/C2NJ20852K.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
166
[14] Mohr, R.; van Eldik, R. Inorg. Chem. 1985, 24, 3396-3399.
[15] Limburg, J.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 1997, 119, 2761-2762.
[16] Limburg, J.; Vrettos, J. S.; Liable-Sands, L. M.; Rheingold, A. L.; Crabtree, R. H.; Brudvig, G. W.
Science 1999, 283, 1524-1527.
[17] Baffert, C.; Romain, S.; Richardot, A.; Leprêtre, J.-C.; Lefebvre, B.; Deronzier, A.; Collomb, M.-
N. J. Am. Chem. Soc. 2005, 127, 13694-13704.
[18] Cady, C. W.; Crabtree, R. H.; Brudvig, G. W. Coord. Chem. Rev. 2008, 252, 444-455.
[19] Dubois, L.; Pécaut, J.; Charlot, M.-F.; Baffert, C.; Collomb, M.-N.; Deronzier, A.; Latour J.-M.
Chem. Eur. J. 2008, 14, 3013-3025.
[20] Collomb, M.-N.; Deronzier, A. Eur. J. Inorg. Chem. 2009, 14, 2025-2046.
[21] Thorp, H. H.; Brudvig, G. W. New J. Chem. 1991, 15, 479-490.
[22] Chen, H.; Tagore, R.; Olack, G.; Vrettos, J. S.; Weng, T.-C.; Penner-Hahn, J.; Crabtree, R. H.;
Brudvig, G. W. Inorg. Chem. 2007, 46, 34-43.
[23] Oakes, J; Gratton, P. J. Chem. Soc., Perkin Trans. 2 1998, 2201-2206.
[24] Feng, W.; Nansheng, D.; Helin, H. Chemosphere 2000, 41, 1233-1238.
[25 ] Offenhartz, P. O.; George, P.; Haight, G. P. J. Phys. Chem. 1963, 67,116-118.
[26] (a) Barton, D. H. R.; Choi, S.-Y.; Hu, B.; Smith, J. A.; Tetrahedron 1998, 54, 3367-3378; (b)
Tagore, R.; Chen, H.; Crabtree, R. H.; Brudvig, G. W. J. Am. Chem. Soc. 2006, 128, 9457-9465;
(c) Jacobsen, E. N. in: I. Ojima (Ed.), Catalytic Asymmetric Synthesis, VCH Publishers, New
York, 1993; (d) Katsuki, T. Coord. Chem. Rev. 1995, 140, 189-214; (e) Yachandra, V. K.; Sauer,
K.; Klein, M. P. Chem. Rev. 1996, 96, 2927-2950; (f) Ruttinger, W.; Dismukes, G. C. Chem. Rev.
1997, 97, 1-24.
[27] Limburg, J.; Vrettos, J. S.; Chen, H.; de Paula, J. C.; Crabtree, R. H.; Brudvig, G. W. J. Am. Chem.
Soc. 2001, 123, 423-430.
[28] (a) Hage, R.; Iburg, J. E.; Kerschner, J.; Koek, J. H.; Lempers, E. L. M.; Martens, R. J.; Racherla,
U. S.; Russel, S. W.; Swarthoff, T.; van Vliet, M. R. P.; Warnaar, J. B.; van der Wolf, L.; Krijnen,
B. Nature, 1994, 369, 637-639; (b) de Boer, J. W.; Browne, W. R.; Brinksma, J.; Alsters, P. L.;
Hage, R.; Feringa, B. L. Inorg. Chem. 2007, 46, 6353-6372.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
167
[29] (a) Groves, J. T.; Stern, M. K. J. Am. Chem. Soc. 1987, 109, 3812-3814; (b) Groves, J. T. Stern,
M. K. J. Am. Chem. Soc. 1988, 110, 8628-8638.
[30] (a) Tetard, D.; Rabion, A.; Verlhac, J.-B.; Guilhem, J. J. Chem. Soc., Chem. Commun., 1995, 531-
532; (b) Menage, S.; Collomb-Dunand-Suathier, M.-N.; Lambeaux, C.; Fontecave, M. J. Chem.
Soc., Chem. Commun. 1994, 16, 1885-1886; (c) Vincent, J. M.; Menage, S.; Lambeaux, C.;
Fontecave, M. Tetrahedron Lett. 1994, 34, 6287-6290.
[31] Handervan, C.; Hage, R.; Feringa, B. L. Chem. Commun. 1997, 5, 419-420.
[32] Stamatis, A.; Doutsia, P.; Vartzouma, C.; Christoforidis, K. C.; Deligiannakis, Y.; Louloudi, M. J.
Mol. Cat. A 2009, 297, 44-53.
[33] (a) Limburg, J.; Vrettos, J. S.; Crabtree, R. H.; Brudvig, G. W.; de Paula, J. C.; Hassan, A.; Barra,
A.-L.; Duboc-Toia, C.; Collomb, M.-N. Inorg. Chem. 2001, 40, 1698-1703; (b) Wei, Z.; Cady, C.
W.; Brudvig, G. W.; Hou, H. J. M. J. Photochem. Photobio. B 2011, 104, 118125; (c) Zhang, F.;
Cady, C. W.; Brudvig, G. W.; Hou, H. J. M. Inorg. Chim. Acta 2011, 366, 128-133, (d) Lieb, D.;
Zahl, A.; Shubina, T. E.; Ivanović -Burmazović, I. J. Am. Chem. Soc. 2010, 132, 7282-7284.
[34] Tagore, R.; Crabtree, R. H.; Brudvig, G. W. Inorg. Chem. 2008, 47, 1815-1823.
[35] Tagore, R.; Crabtree, R. H.; Brudvig, G. W. Inorg. Chem. 2007, 46, 2193-2203.
5. High catalytic activity of a Mn-terpy compound in oxidative dye degradations with peracetic acid
168
6. Summary
169
6 SUMMARY
Although during the last decades a wide variety of manganese compounds has been
investigated towards their ability to activate peroxide for numerous homogeneous and
heterogeneous applications (summarized in Chapter 1.1), there are several aspects that require
further clarification. In fact, there is still an ongoing controversy among scientists on many
mechanistic issues in this area of research. In this context, the general goal of this thesis was to
contribute to a more detailed mechanistic understanding of these processes in aqueous solution
with the help of kinetic methods and spectroscopic techniques, such as UV/Vis- and EPR-
spectroscopy.
In Chapter 2 a fast and environmentally benign method for the oxidative degradation of a
typical industrial dye waste compound was introduced using H2O2 in conjunction with catalytic
amounts of simple manganese salts as catalyst precursors in a carbonate containing aqueous
solution under mild reaction conditions. The results provide an innovative, viable and simplistic
solution for efficient and clean oxidative degradation of highly stable organic dyes and shed light
on some key mechanistic features of this intriguing catalytic system.
The choice of the right buffer turned out to be a decisive factor. It was shown that the
oxidative degradation of the model substrate Orange II is only catalytic in carbonate containing
aqueous solution, since no other buffer solution showed comparable degradation rates. In the
absence of bicarbonate buffer MnIIaq is not able to activate H2O2 for the studied reaction under
identical reaction conditions (pH, ionic strength, concentrations etc.). By the use of excess
bicarbonate buffer a rapid pre-equilibrium between HCO3- and H2O2 leads to the formation of the
deprotonated peracid, HCO4-, as a key molecular entity and the actual oxidizing agent. The
generation of peroxymonocarbonate was further supported by 13C-NMR measurements on the
reaction of H2O2 with H13CO3-, which confirmed the known equilibrium constant for this reaction in
purely aquatic solution. However, the second order dependence of the observed overall dye
degradation rate on the total carbonate concentration in solution implies the involvement of
HCO3- ions in two catalytically relevant reaction steps. According to CV and UV/Vis spectroscopic
measurements the requirement of a second equivalent of HCO3- can be attributed to its potential
6. Summary
170
use as stabilizing ligand. Its electron donating abilities facilitate the formation and stabilization of
catalytically relevant, high-valent Mn-oxo species.
Kinetic investigations performed at different pH could provide relevant information about the
nature of the species involved in the catalytic reaction, viz. the formation of peroxycarbonate and
the deprotonation of the aquated MnII starting catalyst. On the basis of the systematic
investigation presented in Chapter 2, a reaction sequence for the underlying reaction mechanism
was proposed: In the first step a labile manganese-hydroxo precursor reacts with the in situ
formed nucleophilic peroxycarbonate in a rapid pre-equilibrium reaction. The so formed transient
intermediate, most likely a MnII-η2-peroxycarbonate complex, undergoes cleavage of the peroxo
bond to yield a high valent manganese-oxo compound. As a consequence of a heterolytic O-O-
bond scission a MnIV=O species is formed, as indicated by UV/Vis spectroscopic measurements and
confirmed by complementary EPR spectroscopy. Radical degradation pathways which are
expected for one-electron reactions could be excluded. Hence, a key feature of the proposed
reaction mechanism is that the overall oxidation of Orange II with the MnII precursor occurs in a
two-electron oxidation step and leads to the formation of a MnIV=O intermediate. This high valent
MnIV=O species represents a potent key catalytic intermediate which subsequently transfers the
oxo group to the substrate to reform the MnII catalyst and thereby close the catalytic cycle.
Another interesting aspect of a metal ion catalyst is the possible direct interaction of the latter
with the dye molecule itself. In titration experiments the formation of complexes with different
stoichiometries between the azo dye Orange II and MnII was depicted. This precedented
simultaneous σ,π - coordination of the organic dye was underlined by the presented additional
DFT studies and may significantly contribute to the stabilization of the MnII ion pre-catalyst in
slightly basic solution (Chapter 2).
In Chapter 3 the speciation of the catalytic hydrogen peroxide activation for homogeneous dye
degradation by MnII compounds was studied to answer the question to what extent µ-oxo-bridged
manganese dimers of mixed, higher valence nature contribute as reactive intermediates in the
catalytic cycle, since they represent frequently postulated intermediates for this type of reaction.
To have a closer look at the nature of putative intermediates, a comparative study on the
reactivity of a [Mn2III/IV(µ-O)2]3+ dimer and its readily accessible mononuclear analogous MnII
complex in the hydrogen peroxide assisted catalyzed oxidation of Orange II was performed
(Chapter 3). As the study progressed, verification of the obtained results for other dyes as model
6. Summary
171
substrates became increasingly relevant. Consequently p-nitrophenol, an extremely stable
aromatic dye, and Morin as polyphenolic representative were chosen to yield a broader spectrum
of potential substrates.
In summary, it was demonstrated by a well studied example that elaborated catalysts in higher
oxidation states are not always required for efficient catalysis. Instead, the same results are
achieved by the application of [MnII(bpy)2Cl2] as catalyst to oxidize various organic substrates.
Moreover, detailed kinetic investigations revealed that both catalysts show identical reactivity,
implying that the same reactive intermediate is formed regardless of the starting material used.
This was further confirmed by EPR-spectroscopic measurements, which disclosed the generation
of a high-valent MnIV=O intermediate along with a MnII precursor for the [Mn2III/IV(µ-O)2]3+-dimer
and the analogous mononuclear MnII complex upon addition of H2O2 in carbonate containing
solution. The large excess concentrations of H2O2, indispensable for the in situ formation of
peroxycarbonate, is responsible for the rapid reduction and cleavage of the dimer core. Thus, it is
concluded that both catalysts, i.e. [Mn2III/IV(µ-O)2(bpy)4]3+ and the analogous MnII monomer,
exhibit the same two-electron oxidation from a MnII precursor to a high-valent MnIV=O
intermediate.
In addition, it was possible to attain the same catalytic reactivity by simple in situ preparation
from a 1:2 ratio of MnII salt and the bipyridine ligand within the reaction mixture. A comparison
with the results in Chapter 2 confirmed that complexation of the MnII ion by a chelating ligand
generally benefits the reactivity. On the other hand, the stronger coordination at higher bipyridine
concentrations favors the undesired parallel decomposition of H2O2 above substrate oxidation.
Through UV/Vis-titration experiments it was determined that the dominating complex
stoichiometry reflects an 1:2 Mn:bpy compound under the applied reaction conditions.
Furthermore, the detailed investigations revealed very similar kinetics and intermediates as for
the simple MnII ion catalysis in the absence of a ligand, as described above for Chapter 2. Hence,
the general rate enhancement by the addition of the bipyridine ligand is attributed to a potential
stabilization of the reactive high-valent MnIV=O intermediate. However, this implies that the
observed second-order dependence of the observed dye oxidation rate on the total carbonate
concentration is not necessarily the result of bicarbonate coordination to the Mn II-center. In fact,
another plausible explanation might be the occurence of general acid-catalysis of the heterolytic
O-O bond cleavage in which the proton is provided by bicarbonate.
6. Summary
172
Chapter 4 describes the extension of the MnII ion system in Chapter 2 by another peroxide -
peracetic acid or peracetate (PAA), which in contrast to HCO4- is readily accessible. Preliminary
NMR measurements with 13C-labelled acetic acid permitted in situ monitoring of the formation
and dissociation of PAA and in particular of its decomposition at higher pH. The latter is often
assumed to intervene in sensible reproducibility of kinetic data, but the time scale of this
decomposition turned out to be negligible compared to the much faster reaction between the
MnII salt and PAA. In conclusion, the results led to the suggestion that a highly reactive mixture is
formed upon addition of PAA to a weakly basic solution of MnII and that free radical pathways do
not account for the observed performance. Thereby the application of the readily prepared
peroxide that is structurally related to peroxycarbonate, results in an enhanced reactivity towards
the dye substrate Orange II compared to H2O2 (HCO4-). This is a rare example of an investigation on
peracetic acid activation and most notably the first to specifically address the role of the
equilibrium amount of H2O2 omnipresent in PAA.
In the absence of dye substrate the reaction between MnII and PAA shows a biphasic reaction
behavior which is very sensitive to the applied experimental conditions. By careful choice of the
reaction parameters, it was possible to closely investigate the different reaction stages, the
composition of intermediates and their relevance to the catalytic dye degradation reaction by
means of UV/Vis and EPR spectroscopy. It was shown that during the initial stage a MnIV=O
compound is formed, whereas the second phase of the reaction is characterized by the sudden
formation of permanganate and colloidal manganese dioxide. A direct comparison of the catalytic
reactivity of the different intermediates towards the oxidative dye degradation provided more
detailed information on the nature of possible reactive species. Especially the initial reaction
phase shows extraordinary reactivity, whereas a tremendous loss of reactivity is observed
beginning with the formation of MnVIIO4-/MnIVO2. Further assignment of potential reactive species
was achieved by a comparison of the oxidative reactivity between Orange II and several high-
valent Mn-oxo anions, viz. manganate(VI) and hypomanganate(V), in the presence and absence of
PAA, since these oxo-anions often represent postulated intermediates in permanganate
oxidations. In addition, measurements with readily prepared solutions of colloidal MnIVO2 of
known particle size confirmed that heterogeneous pathways do not account for the catalytic
course.
6. Summary
173
However, the omnipresent equilibrium content of H2O2 in the PAA oxidant turned out to play a
decisive role as reductive species in this biphasic reaction. During the initial reaction stage it avoids
accumulation of the less reactive Mn-species, viz. MnVIIO4- and MnIVO2, and thus keeps a favorable
steady state concentration of the suggested reactive MnIV=O intermediate as long as H2O2 is
present in the reaction mixture. The corresponding reactivity tests and EPR measurements
evidenced the unequivocally higher reactivity in the first reaction phase, while a tremendous drop
in the catalytic performance was observed at the stage of MnVIIO4-. Moreover, selective reduction
of the less reactive intermediates in the second phase was achieved by application of small
amounts of additional H2O2 as reductant which thereby restored the catalytic activity. In this way
the lifetime of the catalytic system can be tuned by avoiding two undesired secondary processes,
namely precipitation of inactive MnIVO2 and the over-oxidation of the catalyst to MnVIIO4-. On the
other hand, a large initial excess content of H2O2 was shown to interfere with the in situ formation
of the MnIV=O intermediate, as well as the reactivity towards the dye substrate, by shifting the
complex “oxidation by PAA vs. reduction by H2O2” sequence back to the side of the MnII precursor.
Detailed kinetic studies of the oxidative catalytic degradation of the model substrate Orange II
in conjunction with the results described above provided the basis for a simplified underlying
reaction mechanism. Hence, for the MnII ion catalyzed peroxide activation a major mechanistic
conclusion can be drawn: For the MnII/H2O2/HCO3- system (Chapter 2) the excess H2O2 content,
which is required for more efficient HCO4- formation, enhances the back reaction of the in situ
formed MnIV=O to MnII, thus only a minor amount of catalytically active MnIV=O is constantly
present in the catalytic reaction mixture. The MnII catalyzed activation of PAA (Chapter 4) benefits
from the application of the prefabricated peroxide, viz. peracetate, so that more of the reactive
oxidant is present in solution. This results in a more efficient formation of the active Mn IV=O while
the minor equilibrium content of reducing H2O2 keeps it at a higher steady-state concentration as
compared to the combination of MnII/H2O2/HCO3-. In the case of PAA, a further oxidation to
higher-valent species along with enhanced deactivation of the transient MnIV=O by MnIVO2
precipitation is observed. For this reason small amounts of H2O2 are required to avoid these
secondary processes and thereby extend the catalytic lifetime. Consequently, both reaction types,
viz. MnII/H2O2/HCO3- vs. MnII/PAA, are characterized by reaction sequences of rapid oxidation and
reduction processes which basically lead to the same reactive MnIV=O intermediate. The striking
differences in the catalytic degradation reactivity apparently reflect the differing steady-state
availability of the reactive MnIV=O species.
6. Summary
174
Chapter 5 reports preliminary investigations on an efficient method for the catalytic oxidative
degradation of different highly stable organic dyes, such as Orange II, Tartrazine and p-nitrophenol
by using PAA and a MnIIterpy starting compound. Contrary to our former studies, it shows the
highest activity at pH ≈ 7 and the extraordinary efficiency is reflected by the catalytic TOFs which
were determined for the different substrates. The excess ligand concentration required for
catalysis is needed for a more efficient precursor formation. Moreover, it can be assumed that also
a partial degradation of the terpy ligand occurs during the catalytic reaction. Screening and
spectroscopic methods allowed us to study the reaction course and to shed light on some
mechanistic features. The omnipresent [H2O2]eq in PAA stock solution was shown to result in an
inverse dependence on the catalytic reaction, by causing a catalase-like H2O2 decomposition,
which also becomes noticeable at higher [PAA]. Consequently, the saturation kinetics observed for
higher oxidant concentrations can be ascribed to the thereby increased [H2O2]eq. The second order
dependence of the observed first order rate constants for the overall dye degradation on the
catalyst concentration implies the involvement of dimeric intermediates in the catalytic reaction.
This is confirmed by the UV/Vis and EPR measurements of the species formed in the absence
of substrate which strongly resemble earlier findings concerning the water oxidizing model system
[Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
-. In analogy, it was found that the main species during the
catalytic phase is a µ-oxo bridged Mn2IV/IV dimer with only minor (≤ 10 %) amounts of the
corresponding µ-oxo Mn2III/IV dimer. Unfortunately it was not possible to resolve the complex
reaction behavior which leads to the mixture of different dimers by variation of the numerous
reaction parameters, i.e. [catalyst], [PAA], [H2O2], pH or temperature. However, very similar
UV/Vis spectral characteristics were observed when a combination of HSO5- and H2O2 at similar
conditions as those of the PAA system was used, which already implies parallels in the use of PAA
and HSO5- if the [H2O2]eq is not considered.
Both readily prepared dimer species with a [Mn2III/IV(µ-O)]3+ and [Mn2
IV/IV(µ-O)2]4+ core showed
rather sluggish reactivity towards the dye substrate, which leads to the suggestion that the excess
oxidant might result in the formation of even higher oxidized intermediates as known for the
[Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
- system. There is evidence from preliminary experiments with
phen as ligand system that there are indeed parallels to the intermediates previously postulated
for [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
-. Contrary to terpy, the in situ formed µ-oxo bridged phen
dimers lack the essential feature of the terminal water ligands such that further oxidation is
6. Summary
175
hindered. However, to what extent the different in situ formed µ-oxo dimers and mechanistic
suggestions from the water oxidizing model system [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
- (i.e. its
potential dimeric intermediates with either a MnV=O or a MnIV-O• moiety) account for the
extraordinary reactivity in the MnIIterpy catalyzed dye degradation with PAA, remains to be
clarified.
In summary, this work provides further insights into the mechanism of action, intermediates
and characteristic features of the activation of peroxides by manganese compounds with and
without a defined ligand system which serve as simplest representatives for the numerous
manganese-based bleaching and oxygenation catalysts.
6. Summary
176
7. Zusammenfassung
177
7 ZUSAMMENFASSUNG
Obwohl in den letzen Jahrzehnten eine Vielzahl verschiedenster Manganverbindungen auf ihre
Fähigkeiten zur Aktivierung von Peroxiden für zahlreiche homogene und heterogene
Anwendungen untersucht worden sind, sind viele Einzelheiten bisher noch nicht geklärt. Vielmehr
bietet dieser Bereich nach wie vor Raum für kontroverse Diskussionen innerhalb der
wissenschaftlichen Gemeinde bezüglich verschiedenster mechanistischer Details. In diesem
Zusammenhang war das allgemeine Ziel der vorliegenden Arbeit mittels kinetischer und
spektroskopischer Methoden, wie UV/Vis- und EPR-Spektroskopie, zu einem tieferen
mechanistischen Verständnis dieser katalytischen Prozesse, im Besonderen in wässriger Lösung,
beizutragen.
In Kapitel 2 wurde eine schnelle und umweltfreundliche Methode zum oxidativen Abbau eines
beispielhaften Abfallprodukts der Farbenindustrie unter milden Reaktionsbedingungen mittels
H2O2 und katalytischen Mengen des simplen MnII-Salzes vorgestellt. Die Ergebnisse stellen eine
innovative, praktikable Lösung zum oxidativen Abbau stabiler Farbstoffe dar und geben Aufschluss
über wesentliche mechanistische Merkmale dieser verblüffenden Reaktion.
Die Verwendung des richtigen Puffers erwies sich als ausschlaggebend in diesem katalytischen
System. Es konnte gezeigt werden, dass die oxidative Zersetzung des Modellsubstrates Orange II
nur in wässriger HCO3-/CO3
2--Lösung katalytisch verläuft, da kein anderes Puffersystem
vergleichbare Zersetzungsgeschwindigkeiten zeigte und aquatisierte MnII-Ionen alleine nicht in der
Lage sind, H2O2 unter vergleichbaren Bedingungen (pH, Ionenstärke, Konzentrationen) für die zu
untersuchende Reaktion zu aktivieren. Durch die Verwendung eines Überschusses an HCO3-/CO3
2-
hingegen wird die rasche Einstellung eines vorgelagerten Gleichgewichts zwischen HCO3- und H2O2
ermöglicht, was zur Bildung von HCO4-, der einfach deprotonierten Form der Perkohlensäure, als
eigentlichem Oxidationsmittel führt. Die Peroxycarbonat-Bildung wurde außerdem mittels 13C-
NMR-Messungen der Reaktion zwischen H2O2 und H13CO3- untersucht, wobei die Ergebnisse die
bekannte Gleichgewichtskonstante der Reaktion für eine rein wässrige Lösung bestätigen. Jedoch
impliziert eine Abhängigkeit zweiter Ordnung der beobachteten Geschwindigkeit der katalytischen
Bleichreaktion vom Gesamtgehalt an HCO3-/CO3
2- eine Beteiligung des Hydrogencarbonat an zwei
katalytisch relevanten Reaktionsschritten. Cyclovoltammetrische sowie UV/Vis-spetroskopische
7. Zusammenfassung
178
Messungen deuten auf HCO3- als potentiellen Liganden hin, da es durch seine elektronischen
Eigenschaften die Bildung und Stabilisierung katalytisch relevanter, hochvalenter Mn-oxo-
Intermediate erleichtert.
Durch detaillierte kinetische Untersuchungen bei verschiedenen pH-Werten konnten
grundlegende Informationen über die Natur der oxidierenden Spezies gewonnen werden. Auf
Basis der in Kapitel 2 beschriebenen Ergebnisse wurde ein Reaktionsschema vorgeschlagen, das
den zugrundeliegenden Mechanismus reflektiert: Im ersten Schritt reagiert eine labile MnII-
Hydroxo-Spezies mit dem in situ gebildeten nukleophilen Peroxycarbonat in einem raschen
vorgelagerten Gleichgewicht. Die dadurch gebildete Verbindung, wahrscheinlich ein MnII-η2-
Peroxycarbonat-Komplex, ergibt nach der Spaltung der Peroxo-Bindung ein hochvalentes Mn-oxo-
Intermediat. Als Konsequenz eines heterolytischen O-O-Bindungsbruchs ergibt sich eine MnIV=O-
Spezies, wie aus den UV/Vis-spektroskopischen Messungen und den zusätzlichen EPR-Ergebnissen
zu schließen ist. Reaktionswege, welche auf die Beteiligung freier Radikale hindeuten, wie sie im
Falle einer 1e--Reaktion zu erwarten sind, konnten ausgeschlossen werden. Folglich ist der
Schlüsselschritt des katalytischen, oxidativen Abbaus von Orange II die 2e--Oxidation einer MnII-
Vorstufe zu einem MnIV=O-Intermediat. Dieses kurzlebige MnIV=O-Intermediat stellt eine
potentiell reaktive Spezies dar, welche in der Lage ist, den Sauerstoff auf das Substrat zu
übertragen und somit den katalytischen Zyklus unter Rückbildung der MnII-Vorstufe zu schließen.
Ein weiterer interessanter Aspekt eines Metallionenkatalysators ist die mögliche
Wechselwirkung des freien Ions mit dem Farbstoffmolekül selbst. Durch UV/Vis-
spektrophotometrische Titrationen konnte die Bildung von MnII-(Orange II)-Komplexen
verschiedener Stöchiometrien gezeigt werden. Die wohlbekannte, gleichzeitige σ,π –Bindung des
organischen Farbstoffes wurde zusätzlich durch DFT-Rechnungen einer derartigen
Koordinationsform, welche auch eine mögliche Rolle bei der Stabilisierung der freien MnII-Ionen in
schwach basischem Medium spielt, bestärkt.
Kapitel 3 schildert die weitere Ausführung der homogenen Aktivierung von Wasserstoffperoxid
durch MnII-Verbindungen, wobei die Frage geklärt werden sollte, inwieweit höhervalente, µ-oxo
verbrückte Mn-Verbindungen als reaktive Spezies in den katalytischen Zyklus involviert sind, da
diese häufig postulierte Intermediate in derartigen Reaktionen darstellen. Hierfür wurde eine
vergleichende Untersuchung der Reaktivität eines [Mn2III/IV(µ-O)2]3+-Dimers und des
entsprechenden mononuklearen MnII-Analogons als Katalysatoren für den oxidativen Abbau von
7. Zusammenfassung
179
Orange II durchgeführt. Im Verlauf der Studie gewann es zunehmend an Bedeutung, die
erhaltenen Ergebnisse auch für andere Farbstoffarten als Modellsubstrate zu verifizieren. Mit der
Wahl von p-Nitrophenol, einem hoch stabilen aromatischen Farbstoff, und Morin, exemplarisch
für polyphenolische Farbstoffe, ergab sich ein breites Spektrum potentieller Substrate.
Insgesamt war es möglich, anhand dieser wohlbekannten Verbindungen zu zeigen, dass
effiziente Katalyse nicht immer derartig komplexer Katalysatoren bedarf. So können
verschiedenste organische Substrate ebenso mittels des einfachen [MnII(bpy)2Cl2]-Katalysators
oxidativ zersetzt werden. Die detaillierten kinetischen Untersuchungen ließen vielmehr erkennen,
dass bei gleichem Mangangehalt beide Katalysatoren identische Reaktivität aufweisen, was
wiederum auf die vom Ausgangsmaterial unabhängige Bildung der gleichen reaktiven Spezies
hindeutete. Dies konnte mittels EPR-Spektroskopie bestätigt werden, welche die Bildung eines
MnIV=O-Intermediates neben der Präsenz einer mononuklearen MnII-Verbindung unter den
entsprechenden Reaktionsbedingungen für beide Systeme ([Mn2III/IV(µ-O)2]3+-Dimer und der
analogen MnII-Verbindung) offenbarte. Der für die in situ Bildung des Peroxycarbonats
unverzichtbare große Überschuss an H2O2 zeichnet verantwortlich für die rasche Reduktion und
Spaltung des Dimer-Kerns. Durch die Verwendung von tBuOH als Radikalfänger konnte die
Beteiligung freier Radikale ausgeschlossen werden. Daraus folgt, dass beide Katalysatoren, also
[Mn2III/IV(µ-O)2(bpy)4]3+ und das analoge MnII-Monomer, das selbe katalytische Hauptmerkmal
aufweisen, nämlich eine 2e--Oxidation der MnII-Vorstufe zu einem hochvalenten MnIV=O-
Intermediat.
Des Weiteren wurde eine identische katalytische Aktivität durch simple in situ Herstellung des
Katalysators aus einem 1:2 Verhältnis von MnII-Salz und Bipyridin-Ligand erzielt. Der Vergleich mit
den Ergebnissen für unkomplexiertes MnII zeigt, dass die Komplexierung des MnII-Ions durch einen
Chelatliganden allgemein von Vorteil ist, da höhere katalytische Zersetzungsgeschwindigkeiten
beobachtet werden. Andererseits resultiert aus einem größeren bpy-Gehalt eine höhere
Komplexierung, und damit einhergehend eine Verstärkung der H2O2-Zersetzung als unerwünschte
Nebenreaktion zu Lasten des oxidativen Farbstoffabbaus. Durch UV/Vis-Titration konnte gezeigt
werden, dass die dominierende Komplexstöchiometrie unter den Reaktionsbedingungen der
Studie, die einer 1:2 Mn:bpy-Verbindung ist.
Die Untersuchungen lassen weiterhin auf sehr ähnliche Intermediate und ein ähnliches
mechanistisches Verhalten wie im Falle der in Kapitel 2 geschilderten einfachen MnII-Ionen ohne
7. Zusammenfassung
180
Chelatliganden schließen. Als Folge wird die allgemeine Steigerung der katalytischen Aktivität
durch Zusatz des idealen Verhältnisses an Bipyridin als Ligand der dadurch ermöglichten
Stabilisierung des reaktiven MnIV-Intermediates zugeschrieben. Somit ergibt sich jedoch, dass die
Abhängigkeit zweiter Ordnung der beobachteten Geschwindigkeitskonstanten der katalytischen
Farbstoffzersetzung nicht zwingend das Resultat der Koordination von HCO3- an das MnII-Zentrum
ist. Tatsächlich stellt die Möglichkeit der allgemeinen Säurekatalyse, welche durch Bicarbonat als
Protonenquelle die heterolytische Spaltung der Peroxo-Bindung begünstigt, einen weiteren
Erklärungsansatz dar.
Kapitel 4 schildert die Erweiterung des MnII-Ionen-Systems aus Kapitel 2 um ein weiteres
Peroxid ― Peroxyessigsäure bzw. Peroxyacetat (PAA) ― welches im Gegensatz zu HCO4-
kommerziell zugänglich ist. Vorausgehende NMR-Messungen mit 13C-markierter Essigsäure
ermöglichten zusätzlich eine in situ Analyse der Bildung und Dissoziation, aber vor allem der
Zersetzung von PAA bei höherem pH. Letztere wird häufig als Grund für die schlechte
Reproduzierbarkeit kinetischer Messdaten genannt, doch der zeitliche Rahmen dieser Zersetzung
erwies sich als vernachlässigbar gegenüber der Geschwindigkeit der untersuchten Reaktion von
PAA mit dem MnII-Salz. Die Ergebnisse lassen auf eine nicht-radikalische Bildung eines
hochreaktiven Gemisches durch Zugabe eines MnII-Salzes zu einer schwach basischen Lösung von
PAA schließen. Die Verwendung eines bereits zugänglichen Peroxids, welches Peroxycarbonat
strukturell ähnelt, ermöglicht eine um mehrere Größenordnungen gesteigerte Reaktivität
gegenüber H2O2 (bzw. HCO4-) unter idealen Reaktionsbedingungen. Die Ergebnisse stellen eines
der seltenen Beispiele einer Untersuchung der Aktivierung von PAA dar und behandeln erstmals
ausführlich die Rolle des in PAA allgegenwärtigen H2O2, so dass ein tieferes mechanistisches
Verständnis der komplexen Wechselwirkung mit MnII erhalten wurde.
Die Reaktion zwischen MnII und PAA zeigt in Abwesenheit eines Substrates einen zweiphasigen
Reaktionsverlauf, der sich als äußerst empfindlich bezüglich der verwendeten
Reaktionsbedingungen und Konzentrationen erwies. Durch eine vorsichtige Wahl dieser
Bedingungen war es möglich, mittels UV/Vis- und EPR-Spektroskopie die verschiedenen
Reaktionsphasen, die entsprechende Intermediatzusammensetzung sowie deren Relevanz für die
katalytische Farbstoffzersetzung genauer zu bestimmen. Die Analyse zeigte die Bildung einer
MnIV=O-Spezies innerhalb der ersten Reaktionsphase, während die zweite Reaktionsphase durch
abrupte Bildung von kolloidalem Braunstein und Permanganat gekennzeichnet ist. Eine gezielte
7. Zusammenfassung
181
Untersuchung der Reaktivität bezüglich des oxidativen Farbstoffabbaus der verschiedenen
Reaktionsintermediate erlaubte eine genauere Eingrenzung möglicher reaktiver Formen. Hierbei
zeigt im Besonderen die erste Phase sehr hohe katalytische Aktivität, wohingegen ein erheblicher
Reaktivitätsverlust mit Beginn der Permanganat/MnIVO2-Bildung einhergeht. Ein Vergleich der
Reaktivität mit und ohne PAA gegenüber dem Modellsubstrat für die häufig als Intermediate in
Permanganat Oxidationen postulierten, hochvalenten Mangan-oxo-Anionen Manganat(VI) und
Hypomanganat(V) lieferte eine weitere Eingrenzung der reaktiven Spezies. Zusätzlich konnte ein
heterogener katalytischer Verlauf durch Experimente mit kolloidalen MnIVO2-Lösungen
ausgeschlossen werden.
Als für den zweiphasigen Reaktionsverlauf von MnII und PAA ausschlaggebende Komponente
offenbarte sich die in PAA-Lösungen enthaltene Gleichgewichtskonzentration an H2O2. Diese
verhindert als Reduktionsmittel während der ersten Reaktionsphase nachweislich die Bildung von
Permanganat und kolloidalem MnIVO2, wodurch ein stabiler stationärer Zustand der vermutlich
reaktiven MnIV=O-Spezies erreicht wird, solange H2O2 in der Reaktionsmischung zugegen ist. Die
entsprechenden Reaktivitätsstudien und EPR-Untersuchungen zeigten eindeutig die Korrelation
zwischen den intermediären Spezies und der anfänglich katalytisch aktiven Phase des MnII/PAA-
Systems auf, während ein erheblicher Aktivitätsverlust mit dem Beginn der zweiten Phase
einhergeht. Darüber hinaus war es möglich, die weniger aktiven Spezies der zweiten
Reaktionsphase durch Zugabe kleinerer Mengen an H2O2 selektiv zu reduzieren und somit bei
entsprechend großem PAA-Überschuss die katalytische Aktivität des Systems wieder herzustellen.
Auf diese Art und Weise konnten die unerwünschten Sekundärprozesse, nämlich Desaktivierung
durch MnIVO2-Bildung und „Überoxidation“ des Katalysators zu MnVIIO4-, verhindert werden, was
eine verbesserte Feineinstellung der Lebensdauer des katalytischen Systems ermöglicht.
Andererseits konnte unter Beweis gestellt werden, dass größere Mengen an zusätzlichem H2O2
sowohl die Bildung des vermutlich reaktiven MnIV=O als auch die katalytische Abbaureaktion
behindern, indem die komplexe Reaktionssequenz „Oxidation durch PAA vs. Reduktion durch
H2O2“ weiter auf die Seite der MnII-Ausgangsverbindung geschoben wird.
Detaillierte kinetische Untersuchungen der katalytischen, oxidativen Abbaureaktion des
Modellsubstrats Orange II lieferten in Verbindung mit bereits geschilderten Ergebnissen die Basis
für eine vereinheitlichte mechanistische Vorstellung der zugrundeliegenden Reaktionen. Somit
ergibt sich folgender größerer mechanistischer Zusammenhang für die Peroxidaktivierung durch
7. Zusammenfassung
182
MnII-Salze: Im Falle des MnII/H2O2/HCO3--Systems begünstigt der für die Peroxycarbonat-Bildung
benötigte große Überschuss an H2O2 die Rückreaktion des in situ gebildeten MnIV=O zu MnII, so
dass nur sehr geringe Mengen des katalytisch aktiven MnIV=O dauerhaft in der Reaktionsmischung
verfügbar sind. Dagegen profitiert die MnII-katalysierte Aktivierung von PAA (Kapitel 4) von der
Verwendung des vorgefertigten Peroxides PAA, indem durch den geringeren H2O2-Gehalt die
stationäre Verfügbarkeit des MnIV=O erhöht ist. Andererseits erfolgt im Falle des PAA auch eine
merkliche Weiteroxidation zu höhervalenten Spezies sowie eine Desaktivierung des intermediären
MnIV=O durch Bildung von Braunstein, weshalb geringe Mengen H2O2 benötigt werden, um diese
Sekundärprozesse zu unterbinden und somit die katalytische Lebensdauer zu verlängern.
Zusammengefasst handelt es sich daher bei beiden Reaktionstypen, also MnII/H2O2/HCO3- und
MnII/PAA, um eine komplexe Reaktionsfolge rascher Oxidations- und Reduktionsprozesse, welche
zu demselben aktiven Intermediat führt und in denen lediglich die unterschiedliche stationäre
Verfügbarkeit dieser Spezies für die unterschiedliche Reaktivität verantwortlich zeichnet.
Kapitel 5 beschreibt die anfänglichen Untersuchungen einer neuen effizienten Methode des
katalytischen oxidativen Abbaus verschiedener sehr stabiler organischer Farbstoffe unter
Verwendung von PAA und einer MnIIterpy-Verbindung. Im Gegensatz zu unseren früheren Studien
zeigt das System höchste Aktivität bei pH ≈ 7, und die außergewöhnliche Effizienz wird durch die
für die verschiedenen Substrate bestimmten Wechselzahlen untermauert. Die
Überschusskonzentration an terpy-Ligand, welche für die Abbaukatalyse benötigte wird, lässt sich
der damit einhergehenden effizienteren Bildung der mono-terpy-Katalysatorvorstufe zuschreiben.
Darüber hinaus wird angenommen, dass während des katalytischen Reaktionsverlaufs ebenfalls
eine oxidative Zersetzung des Liganden stattfindet. Verschiedene Arten der spektroskopischen
Reaktionsverfolgung ermöglichten es einige mechanistische Teilaspekte näher zu beleuchten. Es
konnte gezeigt werden, dass die allgegenwärtige [H2O2]eq des Oxidationsmittels PAA in einer
inversen Abhängigkeit der katalytischen Abbaureaktion resultiert. Dies erfolgt durch eine
katalaseähnliche H2O2-Zersetzung, welche sich auch bei höheren [PAA] bemerkbar macht. Damit
ergibt sich auch, dass die festgestellte Sättigungskinetik bei höheren Konzentrationen des
Oxidationsmittels eigentlich der damit erhöhten [H2O2]eq zuzuschreiben ist. Eine mögliche
Teilnahme dimerer Intermediate an entscheidenden Reaktionschritten des katalytischen
Farbstoffabbaus wird durch die Abhängigkeit zweiter Ordnung der beobachteten
Geschwindigkeitskonstanten nahegelegt.
7. Zusammenfassung
183
Diese Vermutung wurde durch UV/Vis- und EPR-Messungen bestätigt, welche sehr stark an die
Ergebnisse des [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
--Modellsystems zur katalytischen
Wasseroxidation erinnern. Analog hierzu fanden wir, dass die Hauptspezies während der
katalytisch aktiven Phase ein µ-oxo verbrücktes Mn2IV/IV-Dimer ist. Des Weiteren wurde ein
geringerer Prozentsatz (≤ 10 %) des entsprechenden µ-oxo-verbrückten Mn2III/IV-Dimers gefunden.
Leider gelang es nicht durch eine Variation der unterschiedlichen Reaktionsparameter, wie [Kat],
[PAA], [H2O2], pH oder Temperatur, das komplexe Reaktionsverhalten, welches zu Bildung der
Mischung beider Dimere führt, besser aufzulösen. Jedoch zeigte eine Kombination aus HSO5- und
H2O2 bei vergleichbaren Konzentrationen sehr ähnliche UV/Vis-Charakteristika, was auf Parallelen
in der Verwendung von PAA und HSO5- schließen lässt, sofern die [H2O2]eq außer Betracht gelassen
wird.
Die bereits zuvor hergestellten dimeren Spezies mit einem [Mn2III/IV(µ-O)]3+- bzw. [Mn2
IV/IV(µ-
O)2]4+-Kern zeigten beide eine eher geringe Aktivität gegenüber dem Farbstoffsubstrat, was die
Vermutung nahelegt, dass der Überschuss an Oxidationsmittel die Bildung weiterer hochvalenter
Intermediate, wie vom [Mn2III/IV(µ-O)2terpy2(H2O)2]3+/HSO5
- System bekannt, ermöglicht.
Tatsächlich lassen einige vorläufige Experimente mit phen als Ligandensystem darauf schließen,
dass Parallelen zu den hierfür postulierten Intermediaten bestehen. Im Gegensatz zu den terpy-
Verbindungen verfügen die entsprechenden in situ gebildeten phen-Dimere nicht über die
ausschlaggebenden terminalen H2O-Liganden, so dass ein weiterer Oxidationsschritt in ihrem Falle
erschwert ist. Inwieweit genau die in situ gebildeten µ-oxo-Dimere bzw. die mechanistischen
Vorstellungen des Modellsystems zur katalytischen Wasseroxidation, d.h. insbesondere die
potentiellen intermediären Dimere mit einer MnV=O-oder MnIV-O•-Einheit, für die ausgesprochen
hohe Reaktivität im MnIIterpy-katalysierten Farbstoffabbau mit PAA verantwortlich zeichnen, muss
hierbei noch geklärt werden.
Zusammengenommen liefert diese Arbeit weitere Erkenntnisse über Intermediate, besondere
Merkmale und den zugrundeliegenden Reaktionsmechanismus der Peroxidaktivierung durch
Manganverbindungen mit und ohne definiertes Ligandensystem, welche als einfache Modelle der
zahlreichen Mangan basierten Bleich- und Oxidationskatalysatoren dienen.
184