Metal ion-catalyzed oxidative degradation of Orange II by H 2O2. High catalytic activity of simple manganese salts

Article abstract of DOI:10.1039/b813725k

In an effort to develop new routes for the clean oxidation of non-biodegradable organic dyes, a detailed study of some environmentally friendly Mn(ii) salts that form very efficient in situ catalysts for the activation of H₂O₂ in the

Full text of DOI:10.1039/b813725k

An international journal of the chemical sciences
New Journal of Chemistry
Volume 33 | Number 1 | January 2009 | Pages 1–212
www.rsc.org/njc
ISSN 1144-0546
PAPER
Rudi van Eldik etal.
Metal ion-catalyzed oxidative degradation of Orange II by H2O2.
High catalytic activity of simple manganese salts

PAPER
www.rsc.org/njc | New Journal of Chemistry
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*
Received (in Montpellier) 6th August 2008, Accepted 25th September 2008
First published as an Advance Article on the web 17th November 2008
DOI: 10.1039/b813725k
In an effort to develop new routes for the clean oxidation of non-biodegradable organic dyes,
a detailed study of some environmentally friendly Mn(II) salts that form very efficient in situ
catalysts for the activation of H2O2 in the oxidation of substrates such as Orange II under mild
reaction conditions, was performed. The studied systems have advantages from the viewpoint of
green chemistry in that simple metal salts can be used as very efficient catalyst precursors and
H2O2 is used as a green oxygen donor reagent. Oxidations were carried out in a glass reactor over
a wide pH range in aqueous solution at room temperature. Under optimized conditions it was
possible to degrade Orange II in a carbonate buffer solution in less then 100 s using 0.01 M H2O2
in the presence of only 2 ꢀ 10ꢁ5 M Mn(II) salt. To gain insight into the manganese catalyzed
oxidation mechanism, the formation of the active catalyst was followed spectrophotometrically
and appears to be the initiating step in the oxidative degradation of the dye. High valent
manganese oxo species are instable in the absence of a stabilizing coordinating ligand and lead to
a rapid formation of catalytically inactive MnO2. In this context, the role of the organic dye and
ꢁ
HCO3 as potential stabilizing ligands was studied in detail. In situ UV-Vis spectrophotometric
measurements were performed to study the effect of pH and carbonate concentration of the buffer
solution on the formation of the catalytically active species. Electrochemical measurements and
DFT (B3LYP/LANL2DZp) calculations were used to study the in situ formation of the catalytic
species. The catalytic cycle could be repeated several times and demonstrated an excellent stability
of the catalytic species during the oxidation process. A mechanism that accounts for the
experimental observations is proposed for the overall catalytic cycle.
One approach to solve these problems would be to develop
Introduction
low-cost, highly efficient, and environmental friendly oxida-
tion catalysts on the basis of transition metal complexes.6,7
Nowadays, one of the major environmental problems con-
cerns the strong increase in xenobiotic and organic substances
Recently, photodegradation methods based on TiO2 as a
photocatalyst,8 beside Fenton systems,9 emerged as one of
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.1 A significant
tion due to their practical and potential value in environmental
source of environmental pollution is industrial dye waste due
the most promising technologies and received increasing atten-
protection. However, in some cases they are only successful
to their visibility and recalcitrance, since dyes are highly
under specific pH and temperature conditions.
coloured and designed to resist chemical, biochemical and
Several studies were performed during the last few years in
photochemical degradation.2 About half of the global produc-
order to find good catalysts for the oxidative degradation of
tion of synthetic dyes (700 000 t per year) are classified as
different organic dyes. From an environmental point of view,
aromatic azo compounds that have a –NQN– unit as chromo-
first row transition metals are the most challenging. Highly
effective Fe,10,14 Co,11 Cr12 and Mn13 based oxidation catalyst
phore in their molecular structure. Over 15% of textile dyes
are lost in waste water streams during the dyeing operation.3
were developed. In combination with different oxidizing
Azo dyes are known to be largely non-biodegradable under
agents, the decomposition of stable organic substances was
aerobic conditions and to be reduced to more hazardous
possible. A novel highly active and environmental benign
intermediates under anaerobic conditions.4 The decolorization
catalytic system based on Fe-TAML (TAML = tetraamido
macrocyclic ligand) was recently reported by Chahbane et al.14
of wastewater has acquired increasingly importance in recent
years, however, there is no simple solution to this problem
In many cases tremendous synthetic efforts are required to
because the conventional physicochemical methods are costly
obtain an effective catalytic system and in addition the pre-
and lead to the accumulation of sludges.5
sence of high concentrations of oxidizing agents is needed.
Among the possible oxidizing agents, H2O2 is one of the most
commonly used owing to its eco-friendly nature. The use of
Inorganic Chemistry, Department of Chemistry and Pharmacy,
University of Erlangen-Nurnberg, Egerlandstr. 1, 91058 Erlangen,
¨
Germany
H2O2 as a green oxidizing agent in these reactions is justified
by a low organic content of the wastewater to be treated and a
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34 | New J. Chem., 2009, 33, 34–49

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 aceto-
nitrile (Payne oxidation),15 the generation of peroxyisourea
from carbodiimide in the presence of either a weak acid or a
mild base,16 or the use of percarbonate, persulfate or perbo-
rate in strongly basic solution.17 In order to achieve fast
oxidative transformations, the use of large amounts of co-
catalyst additives is often required.18
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.24
Experimental
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 mix-
ture at 4 1C. 2,4,6-Tri-tert-butylphenol (TTBP) 96% was
purchased from Sigma–Aldrich and recrystallised several
times from EtOH/H2O (9 : 1) mixtures prior to use. Hydrogen
peroxide 35 wt% as well as different manganese salt hydrates
used in the experiments, were of analytical grade and provided
by Acros Organics (Germany). Carbonate buffer solutions
were prepared using Millipore Milli-Q purified water.
Among these, the use of percarbonate, a versatile oxidizing
agent, is preferred for environmental reasons.19,20 Oxidation
using environmentally benign oxidants has aroused much
interest,7,21 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.22
In this report, we propose a fast and clean catalytic oxida-
tive degradation of Orange II as model substrate by H2O2 in
aqueous carbonate solution under mild reaction conditions,
pH 8–10 and 25 1C, eqn (1).
General procedure
The manganese salts were freshly dissolved in water before
use. To a freshly prepared sodium carbonate 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 carbonate
solution, and H2O2. In typical measurements, 0.01 M H2O2
was prepared from a 35 wt% solution of H2O2. In addition, to
gain more information on the activation mode of the catalyst,
two further experimental procedures based on different activa-
tion 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 IIꢃ ꢃ ꢃMnII 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 phos-
phate, showed the observed catalytic reaction.
ð1Þ
Starting from commercially available Mn(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.23 More-
over, 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 (eqn (2)) is
considered to be of major importance for their catalytic
activity.
Kinetic study of the manganese catalysed oxidative degradation
of Orange II by H2O2
In the present study, the formation of catalytically inactive
Mn(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 man-
ganese 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 inter-
mediate under different reaction conditions. Electrochemical
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 mag-
netic stirrer, a 2 ꢀ 10ꢁ5 M freshly prepared catalyst solution
and 0.01 M H2O2 were added to 40 ml of 5 ꢀ 10ꢁ5 M dye
at a pH ranging from 8 to 10 at 25 1C. All kinetic measure-
ments were carried out under pseudo-first order conditions
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New J. Chem., 2009, 33, 34–49 | 35

(i.e. 50 r [H2O2]/[Mn2+] r 1000). The pH of the aqueous
carbonate solution 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 480 nm. 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 using an SX.18MV
stopped-flow instrument from Applied Photophysics.
DFT calculations
Unrestricted B3LYP/LANL2DZp hybrid density functional
calculations,25a–c i.e., with pseudo-potentials on the heavy
elements and the valence basis set25d–f augmented with
polarization functions,25g were carried out using the Gaussian
0326 suite of programs. The relative energies were corrected
for zero point vibrational energies (ZPE). The resulting struc-
tures were characterized as minima by computation of vibra-
tional frequencies, and the wave functions were tested for
stability.
Synthesis of insoluble MnCO3
Spectrophotometric titration
In a 150 ml round flask 3.36 g (0.4 M) NaHCO3 was 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) Mn(NO3)2 was added. The mixture was stirred at
room temperature for 15 min during which MnCO3ꢃ ꢃ ꢃH2O
formed as a white precipitate. The product was filtered and
washed several times with large amounts of water. Yield:
0.44 g MnCO3, 96.2%. IR (KBr pellets): n (cmꢁ1) 3421 (m),
1416 (vs), 862 (s), 725 (m). Elemental analysis (%) for
MnCH2O4: calc.: C 9.03, H 1.52; found: C 9.38, H 1.52.
UV-Vis spectra were recorded on a Shimadzu UV-2101
spectrophotometer at 25 1C. A 0.88 cm path length tandem
cuvette with two separate compartments (0.44 cm path length
each), was filled with 1 ml 5 ꢀ 10ꢁ5 M Orange II stock solution
in one, and different concentrations of an aqueous Mn(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. The resulting spectrum presents the sum of the
two individual spectra before, and that of the reaction mixture,
after mixing. The observed spectral change is a result of
complex-formation between Mn(II) and Orange II.
Synthesis of Orange IIꢃ ꢃ ꢃMnII complex
In a 50 ml Schlenk tube 0.014 g (2 ꢀ 10ꢁ3 M) Orange II was
dissolved in 20 ml doubly distilled water and an aqueous
solution of 0.01 g (2 ꢀ 10ꢁ3 M) Mn(NO3)2 was 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. Yield:
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 mm 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 sꢁ1 at 25 1C.
Potentials were measured in a 0.5 M NaCl/NaOH electrolyte
solution and are reported vs. an Ag/AgCl electrode.
0.018 g Orange IIꢃ ꢃ ꢃMnII, 87.1%. IR (KBr pellets): n (cmꢁ1
)
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 MnC16H18O10N3SNa:
calc.: 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 IIꢃ ꢃ ꢃMnIIꢃ ꢃ ꢃOrange II complex
An aqueous solution of 0.005 g (1 ꢀ 10ꢁ3 M) Mn(NO3)2 was
added under continuous stirring to a 0.014 g (2 ꢀ 10ꢁ3 M)
Orange II water solution at room temperature. The pale
yellow precipitate was collected by filtration and dried in
air. Yield: 0.017 g Orange IIꢃ ꢃ ꢃMnIIꢃ ꢃ ꢃOrange II, 93.7%. IR
(KBr pellets): n (cmꢁ1) 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 MnC32H32O15N5S2Na2: calc.: 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.
IR measurements
IR spectra were recorded as KBr pellets using a Mattson
Infinity FTIR instrument (60 AR) at 4 cmꢁ1 resolution in
the 400–4000 cmꢁ1 range.
Elemental analysis
The measurements were carried out on an elemental analyzer
Euro EA 3000 instrument from Hekaltech Gmbh. The analy-
tical method is based on the complete instantaneous oxidation
of the sample by ‘‘flash combustion’’ at 1000 1C, 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.
Results and discussion
General observations
A series of experiments were 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 Mn(II) salt.
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36 | New J. Chem., 2009, 33, 34–49

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 tre-
mendously accelerated by addition of a simple manganese
salt. The reactivity of the in situ formed intermediate
was comparable with the catalytic activity of some earlier
postulated, well known manganese bleach catalysts13,32 and
manganese porphyrins.33 In our work, the formation and
stabilization of the active catalyst was studied in a carbonate
buffer solution.
Complex-formation between Orange II and Mn2+
Scheme 1 Orange II: R = SO3Na; pKA = 11.4, lmax = 480 nm.
ortho-Hydroxy aromatic azo dyes, which are bidentate com-
plexing agents are of considerable practical and theoretical
interest because of their ability to form stable chelate com-
plexes with some metal ions.34 It is known that Orange II can
act as a chelating agent since the hydroxy and sulfonate groups
allow a stabilized complex to be formed.35 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 Fig. 2.
Oxidation reactions are in general 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,27 with the relative
amounts varying with reaction parameters such as solvent
and temperature.28 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 1).29 These are chemically distinct
forms which have characteristically different visible spectra,
the azo form absorbs typically at 400–440 nm and the hydra-
zone form at 475–510 nm (see Fig. 1).30
UV-Vis spectra recorded before and after mixing (ca. 5 s
delay) of 5 ꢀ 10ꢁ5 M Orange II with 5 ꢀ 10ꢁ5 M Mn(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 Mn2+ concentrations. On
increasing the Mn2+ concentration, a continuous increase in
DAl=480 nm = A(dye+Mn( )) ꢁ Adye was observed, indicating
II
the formation of an Orange IIꢃ ꢃ ꢃMn2+ species according to
eqn (3). It should be noted that at higher Mn2+ concentration,
a precipitate started to form. The value of Keq was determined
through a constant variation of the Mn2+ concentration. For
a correct determination of the complex-formation constant,
independent measurements were performed at constant man-
ganese concentration where the Orange II concentration was
continuously varied (see Fig. 3C). Independent measurements
were repeated between five and eight times. Selected data are
The absorption spectrum of Orange II in an aqueous
carbonate solution shows under the selected reaction condi-
tions (Fig. 1) one main band at 480 nm, which correspond to
the n - p* transition of the azo form. The other two bands at
300 and 270 nm are attributed to the p - p* transition of the
benzene and naphthalene rings, respectively.31 Orange II, due
Fig. 2 (Blue curve) UV-Vis spectrum of a 5 ꢀ 10ꢁ5 M Orange II
carbonate (0.1 M HCO3ꢁ) solution at pH 8.5 before mixing with a
5 ꢀ 10ꢁ5 M Mn(NO3)2 solution at pH 8.5. (red curve) UV-Vis
spectrum recorded directly after mixing (ca. 5 s delay).
Fig. 1 UV-Vis spectrum of 10ꢁ4 M Orange II in carbonate buffer
solution at pH 8.5.
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New J. Chem., 2009, 33, 34–49 | 37

Fig. 3 (A) Change in absorbance at 480 nm on addition of different concentrations of Mn2+ to 5 ꢀ 10ꢁ5 M Orange II in aqueous carbonate
solution (0.2 M HCO3ꢁ) at pH 8.5 and 22 1C. (B) Job plot analysis for complex-formation between Orange II and Mn2+ in aqueous carbonate
solution (0.2 M HCO3ꢁ) at pH 8.5. (C) Spectral changes at 480 nm on addition of different concentrations of Orange II to a freshly prepared
5 ꢀ 10ꢁ5 M Mn(NO3)2 carbonate solution (0.2 M HCO3ꢁ) at pH 8.5 and 22 1C. (D) Job plot analysis for the complex-formation in aqueous
carbonate solution (0.2 M HCO3ꢁ).
shown in Fig. 3A, where the solid line represents a fit of eqn (4)
concentration the formation of a complex with a stoichiometry
of 1 : 1 can be assumed. On increasing the Orange II concen-
tration further, complexes with a higher stoichiometry are
possibly formed (see Fig. 4A and B).
to the data.
Keq
Orange II þ Mn2þ ÐðOrange II ꢃ ꢃ ꢃ MnIIÞ
ð3Þ
Similar structures have been reported earlier by Nadtochenko
and Kiwi when a Fe3+ salt was added to an Orange II solution
in acidic medium.40 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
Ax ꢁ A0 = (AN ꢁ A0)Keq[Orange II]/(1 + Keq[Orange II])
DA = DANKeq[Orange II]/(1 + Keq[Orange II]) (4)
The values of A0 and AN represent the absorbances of
Orange II and of the complex Orange IIꢃ ꢃ ꢃMnII, respectively,
and Ax is the absorbance at any MnII concentration. The value
of Keq was calculated from eqn (4) to be (2.9 ꢄ 0.9) ꢀ 104 Mꢁ1
,
indicating a relatively weak coordination of the dye to the metal
center. Experimentally, through addition of a 4 ꢀ 10ꢁ5
M
Mn(NO3)2 solution to a 5 ꢀ 10ꢁ5 M Orange II aqueous
carbonate solution (0.2 M HCO3ꢁ), a decrease in the pH of
the solution from 8.5 to 8.3 was observed, which suggests
phenolic proton release due to Mn(II) 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. 10ꢁ3 M) Orange II forms dimers and
higher aggregates in aqueous solutions,30,36 and has a marked
effect on the observed spectra, particularly UV-Vis and NMR.37
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.38 Using Job’s method,39 the
stoichiometry of the formed complex could be determined.
According to the data shown in Fig. 3B and D, at lower MnII
Fig. 4 (A) Proposed structure for a 1 : 1 Orange IIꢃ ꢃ ꢃMn2+ complex
formed in a carbonate buffer solution at a low concentration of Mn2+
.
(B) Proposed structure for a 1 : 2 Orange IIꢃ ꢃ ꢃMnꢃ ꢃ ꢃOrange II
complex formed in a carbonate buffer solution at a high concentration
of Orange II.
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38 | New J. Chem., 2009, 33, 34–49

hydrazone tautomer.41 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
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. Furthermore, the reduction potential of
Mn3+ decreased from +0.35 V to +0.28 V when Orange II
was added to the solution, indicating the stabilization of
Mn3+ ions. In the presence of a chelating substrate, the
generated Mn3+ complex becomes more stable and the redox
potentials attain lower values.43 When the concentration of
Orange II was increased up to 2 ꢀ 10ꢁ5 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.
MnII, was recently postulated by Lopez et al.35 Although, the
´
coordination of organic dyes, viz. Alizarin, Alizarin S42 and
Orange II,3 to several transition metal centres has been known
for years, comparatively little has appeared on their use as
potential stabilizing ligands in oxidative degradation of
organic dyes.
The formed Orange IIꢃ ꢃ ꢃMnII complex was isolated and the
validity of its composition was confirmed by elemental ana-
lysis. In control experiments the reactivity of the isolated 1 : 1
Orange IIꢃ ꢃ ꢃMnII 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 coordina-
tion mode of the ligand, no differences between the catalytic
activity of the 1 : 1 and 2 : 1 complex were found.
Beside UV-Vis measurements and DFT calculations,
electrochemical measurements were used to study the in situ
formation of highly reactive MnII catalytic species in
the presence of Orange II under the selected experimental
conditions.
DFT calculations
To assess the coordination mode of Orange II to the Mn(II)
center, DFT (B3LYP/LANL2DZp) calculations were per-
formed for a series of plausible complexes. Orange II in
aqueous solution under the selected experimental conditions
dissociates 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.28 Solvent Yellow 14, a model com-
pound for Orange II containing no sulfonate group was
selected for the DFT study of the interaction between the
Mn(II) and the chosen azo dye. A picture of the calculated
conformers of the model compound 1 is shown in Fig. 6.
The optimized geometry of 1a was calculated to be
ca. 5.8 kcal molꢁ1 lower in energy than that of 1b. Further-
more, the calculated structure of 1a was compared with X-ray
structural data of Solvent Yellow 14.44 A good agreement
between calculated and crystallographically determined struc-
ture was found.
CV studies on the complex-formation between Orange II
and Mn2+
CV measurements of a 4 ꢀ 10ꢁ5 M Mn2+ solution in the
presence of different Orange II concentrations were performed
in order to determine the interaction between the fully aquated
Mn2+ ions and Orange II present in the reaction mixture.
Fig. 5 shows the results of MnIIꢃ ꢃ ꢃOrange II complex forma-
tion in NaCl electrolyte solution, performed using a standard
three electrode electrochemical setup as described above.
To avoid the oxidation of MnII to MnIV, which precipitates
as MnO2, the potential scan was discontinued at +1.0 V, after
which the reverse scan from +1.0 to ꢁ0.8 V was started. The
According to the UV-Vis and electrochemical data pre-
sented above, Orange II can coordinate to a fully aquated
Mn2+ center. Different plausible interaction modes of Solvent
Yellow 14ꢃ ꢃ ꢃMnII (2) and Solvent Yellow 14ꢃ ꢃ ꢃMnIIꢃ ꢃ ꢃSolvent
Yellow 14 (3) were studied in detail. Optimized structures of 2
adopting different coordination modes are presented in Fig. 7.
The studied organic dye can coordinate to aquated Mn2+
ion by forming two new bonds, one between Mn2+ and the
deprotonated phenolic OH-group of 1a and the second
between Mn2+ and one of the azo nitrogen atoms, leading
2+
CVs of Mnaq 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 Mn3+/Mn2+
redox couple. In addition, CV measurements on a freshly
prepared 4 ꢀ 10ꢁ5 M Orange II electrolyte solution at
pH 8.5 and 22 1C were performed. Orange II, as it can be seen
in Fig. 5, undergoes two electrochemically quasi-reversible,
Fig. 5 CVs of a 4 ꢀ 10ꢁ5 M Mn2+ electrolyte (0.1 M NaCl) solution in the presence of different Orange II concentrations at pH 8.5 (adjusted by
addition of NaOH) and 22 1C.
ꢂc
This journal is The Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2009
New J. Chem., 2009, 33, 34–49 | 39

Fig. 6 Optimized (B3LYP/LANL2DZp) structures of 1a and 1b with a planar geometry and dihedral angles of (a) 180.01 and (b) 178.71 about the
azo group, C–N–N–C.
Fig. 7 Optimized structures of complex 2a, b and c (B3LYP/LANL2DZp).
to the formation of either a planar six-membered (2a) or five-
membered (2b) chelate complex. Furthermore, a second inter-
action 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 indi-
cate that 2a is energetically favoured over 2b by about
3 kcal molꢁ1. The N–N bond length of 1.30 A for 2a is nearly
identical to that found in the free model molecule 1a (1.28 A),
indicating a weak interaction between the nitrogen atoms and
the positively charged manganese center.
(2.213–2.294 A),46a 2-[N,N-bis(2-pyridylmethyl)amoniumethyl]-6-
[N-(3,5-di-tert-butyl-2-oxidobenzyl)-N-(2-pyridylamino)amino-
methyl]-4-methylphenol (H2Ldtb) (2.118–2.237 A)46b and
1,4,7-triazacyclononane (tacn) (2.118–2.146 A).46c
As expected, upon coordination of two dye molecules in 3,
the N–N bond distance becomes longer (1.29 A) 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 A vs. 2.06–2.27 A for 2, and Mn–O = 2.27/
2.26 A vs. 1.81/2.11 A for 3) exerts a significant trans influence
In addition to these structures, DFT calculations were
performed for a further possible interaction of a second dye
molecule with the Mn(II) center leading to the formation of
chelated Mn(II) inner-sphere complexes. Similar transition
metal complexes of ortho-hydroxy azo dyes were prepared
and characterised by Drew and Landquist.45 The introduction
of a second dye molecule is expected to have certain advan-
tages. 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 is presented in Fig. 8.
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
favoured 2a (2.15 A) and 3 (2.30 A) are comparable with the
X-ray structural data for Mn(II) complexes with nitrogen
containing ligands such as 1,2-bis(imidazol-1-yl)ethane (bim)
Fig. 8 Optimized structure of 3 (B3LYP/LANL2DZp).
ꢂc
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40 | New J. Chem., 2009, 33, 34–49

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.47 suggest that in the
presence of weak-field ligands for Mn(II) and Mn(III), five-
coordination is also accessible whereas Mn(IV) has a much
stronger preference for six-coordination.47
Complex-formation between bicarbonate and Mn(II)
The reactions between bicarbonate ions (HCO3ꢁ) and differ-
ent manganese species have been studied for several years,
since aquated Mn2+ cations themselves are actually not able
to catalyze H2O2 disproportionation. Depending on thꢁe
HCO3ꢁ concentration in the reaction mixture, MnIIꢃ ꢃ ꢃHCO3
complexes of different stoichiometry can be formed. Recently,
it was suggested that only the neutral MnII(HCO3ꢁ)2 complex
can facilitate H2O2 disproportionation.48 In this study the
Fig. 10 Plot of observed first order rate constant (kobs) for the
formation of Mn2+ꢃ ꢃ ꢃHCO3 vs. the bicarbonate concentration in
ꢁ
the presence of 4 ꢀ 10ꢁ4 M Mn2+ at pH 8.5 and 25 1C.
in situ formed manganese intermediate occurs. A significant
time dependeꢁnt loss in catalytic efficiency of the formed
MnIIꢃ ꢃ ꢃHCO3 intermediate was observed. 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, 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.
complex-formation reaction between Mn2+ and HCO3 was
ꢁ
monitored using UV-Vis spectrophotometric beside CV mea-
surements as a function of carbonate concentration at pH 8.5.
ꢁ
UV-Vis spectra recorded before and after addition of HCO3
to an aqueous Mn2+ solution showed the formation of a new
broad band at 300 nm as illustrated in Fig. 9A. The time
course of the absorption band formation is shown in Fig. 9B.
It can be seen from Fig. 9B that the rate of formation of the
manganese carbonate intermediate is enhanced at higher
carbonate concentration. The observed first order rate con-
stants following the induction period in Fig. 9B, are directly
proportional to the [HCO3ꢁ] in the range 0.01–0.1 M
(see Fig. 10) with a second order rate constant of (3.6 ꢄ 0.2) ꢀ
10ꢁ2 Mꢁ1 sꢁ1 at 25 1C. Moreover, the observed induction
period is probably related to the displacement of water from
the first coordination sphere of the fully aquated Mn2+ ion by
CV measurements of freshly prepared aqueous Mn(NO3)2
solutions were performed in the presence of different carbo-
nate concentrations in a 0.1 M NaCl electrolyte solution at
pH 8.5 (adjusted by careful addition of NaOH) and 22 1C. 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 HCO3ꢁ to a positively charged metal, the peak
potentials are shifted to more negative potentials compared to
the fully aquated Mn2+ (see Fig. 11A and 12).41 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 ꢀ 10ꢁ5 M Mn2+ solution in the presence of 0.2 M
ꢁ
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 MnCO3
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 and
25 1C. During the first 200 s, no change in the reactivity of the
Fig. 9 (A) UV-Vis spectra of an aqueous 4 ꢀ 10ꢁ4 M Mn2+ solution before (black curve) and after (red curve) addition of 0.4 M HCO3ꢁ at pꢁH
8.5 and 25 1C. (B) Time course of the band formation at 300 nm of an aqueous 2 ꢀ 10ꢁ4 M Mn2+ solution containing different amounts of HCO3
.
ꢂc
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New J. Chem., 2009, 33, 34–49 | 41

Fig. 11 (A) Cyclovoltammograms for 4 ꢀ 10ꢁ5 M Mn2+ solution in an aqueous solution of 0.1 M NaCl and different concentrations of HCO3
(B) Typical multiple scan CVs of a 4 ꢀ 10ꢁ5 M Mn2+ solution in the presence of 0.2 M NaHCO3 and 0.1 M NaCl at pH 8.5 and 22 1C.
.
ꢁ
Fig. 13 Second-order carbonate concentration dependence of kobs
.
Experimental conditions: 2 ꢀ 10ꢁ5 M Mn(NO3)2, 5 ꢀ 10ꢁ5 M Orange II,
0.01 M H2O2, pH 8.5, 25 1C.
Fig. 12 Plot of peak potential E as function of [HCO3ꢁ]; E vs.
Ag/AgCl electrode, [Mn2+] = 4 ꢀ 10ꢁ5 M, [HCO3ꢁ] = 0.1–05 M
in 0.1 M NaCl electrolyte solution at pH 8.5 and 22 1C.
In the present case, the catalytic reaction leads to a square
dependence of kobs on the HCO3ꢁ concentration (Fig. 13) with
a third rate constant (8.3 ꢄ 0.3) ꢀ 10ꢁ2 Mꢁ2
s
ꢁ1, suggesting
NaHCO3 and 0.1 M NaCl at pH 8.5 and 25 1C is presented in
Fig. 11B. In the presence of a chelating substrate, the gener-
ated Mn3+ 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 eqn (5), was
observed.
ꢁ
that 2 equivalents of HCO3 are involved in the oxidation
mechanism. It is suggested, among other possibilities, that one
equivalent of HCO3ꢁ is required for the formation of the more
reactive [MnII(H2O)5(HCO3ꢁ)]+ intermediate, and the second
equivalent of HCO3ꢁ is required for the formation of the more
reactive peroxocarbonate species, known to be a versatile
oxidizing agent. It should also be noticed that no oxidation
of Orange II by H2O2 was observed in the absence of a
carbonate buffer. 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 HCO4ꢁ, respectively.
K1
½MnIIðH2OÞ ꢅ2þ þHCO3ꢁ нMnIIðH2OÞ ðHCO3Þꢅþ
K2
6
5
½MnIIðH2OÞ5ðHCO3ÞꢅþþHCO3ꢁ нMnIIðH2OÞ4ðHCO3Þ2ꢅ
ð5Þ
The reaction of carbonate with H2O2 at pH 8.5 and 25 1C
Peroxycarbonate ions, known to be several orders of magni-
tude more reactive toward nucleophilic substrates than H2O2
By plotting the peak potential E as a function of the
hydrogen carbonate concentration (see Fig. 12), the presence
of different complex species at different carbonate concentra-
tions is revealed.
itself,49 are formed in
a relatively fast pre-equilibrium
(K = 0.32 ꢄ 0.02 Mꢁ1 40
)
between hydrogen carbonate ions
and H2O2 shown in eqn (6).
Carbonate concentration dependence of the catalytic reaction
course
ꢁ
ꢁ
HCO3 + H2O2 " HCO4 + H2O
(6)
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.
Moreover, the reaction of H2O2 and HCO3ꢁ to form the more
electrophilic HOOCO2ꢁ (HCO4ꢁ) occurs rapidly (t1/2 E 300 s)
at near neutral pH and 25 1C.50 This step is also regarded to be
a key aspect of several oxidation reactions.51,52 The higher
ꢂc
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42 | New J. Chem., 2009, 33, 34–49

Fig. 14 (A) Spectral changes observed at 480 nm for the 2 ꢀ 10ꢁ5 M Mn(NO3)2 catalyzed oxidative degradation of 2.5 ꢀ 10ꢁ5 M Orange II in the
presence of (black curve) 0.01 M H2O2 and (red curve) 0.01 M HCO4ꢁ, respectively, at pH 8.5 and 0.5 M total carbonate concentration.
(B). Comparison of the absorbance changes at 480 nm vs. time for the 2 ꢀ 10ꢁ5 M Mn2+ catalyzed oxidative degradation of 5 ꢀ 10ꢁ5 M Orange II
by 0.01 M H2O2 at pH 8.5 and different carbonate concentrations.
reactivity of peroxycarbonate compared to that of H2O2 is
attributed to carbonate being a better leaving group than
hydroxide.40
observations the formed complex with an absorption band at
400 nm could be attributed to a MnIV–Z2-peroxycarbonate
intermediate.53 Based on spectroscopic observations and data
reported in the literature,54 the formed intermediate can be
regarded as most likely to be a high valent manganese com-
plex. Similar Rh,55 Pt56 and Fe57 peroxycarbonate complexes
have been isolated before and were characterized spectro-
scopically. The time course of the absorption band at 400 nm
at different pH is illustrated in Fig. 15B.
By performing the oxidation reactions in the presence of
peroxycarbonate instead of H2O2 in a 0.5 M carbonate con-
taining buffer solution at pH 8.5, no difference in the reactivity
was observed (Fig. 14A).
The Mn(II) catalyzed oxidative degradation of Orange II by
ꢁ
using HCO4 as an oxidizing agent could be significantly
enhanced through increasing the total carbonate concentra-
tion in the reaction mixture. This can be explained in terms of
the equilibrium formulated in eqn (6). Based on our experi-
mental observations and aspects reported in the literature41 for
the Mn(II) catalyzed oxidation reaction by H2O2 in a carbo-
nate containing solution, the reaction sequence presented in
Scheme 2 can be suggested to occur.
In the absence of any stabilizing ligand, the formed complex
rapidly decomposes with the formation of catalytically
inactive MnIVO2 that precipitates from solution (see Fig. 15B).
The decomposition of the active intermediate is accelerated at
higher pH (see Fig. 15B). To ascertain that the formulated
reaction steps in Scheme 2 are valid under our reaction
conditions, a systematic spectroscopic investigation at differ-
ent pH values was performed. Representative data for the
reaction course at 400 nm at pH 8.5 and 9.5 are presented in
Fig. 15B. 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 forma-
tion rate under the mentioned reaction conditions. This could
be partly due to subsequent formation of Mn(OH)2 pꢁrecipi-
tates at higher pH and to deprotonation of HCO3 that
becomes significant at pH above 8 to 9.51 This results in a
Complex-formation between Mn2+ and H2O2 in an aqueous
carbonate solution
Addition of H2O2 to hexaaqua Mn2+ in a carbonate solution
leads to significant spectral changes in the UV-Vis spectra
during the reaction (see Fig. 15A). The initial rapid increase
of the intensity of the broad band at 300 nm, as it is illustra-
ted in Fig. 15A, 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
ꢁ
decrease in the HCO3 concentration in the equilibrium
presented in eqn (6), reducing the concentration of peroxy-
carbonate present in solution.
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 oxidiz-
ing agꢁent. In our work the kinetics was studied in 0.4 M
HCO3 containing buffer solution in the pH range between
8.0 and 9.5 at 25 1C. The pH of the carbonate buffer solution
was adjusted carefully using small amounts of concentrated
NaOH solution to avoid dilution. A typical manganese cata-
lyzed oxidative degradation of Orange II by H2O2 in a
carbonate buffer solution is presented in Fig. 16A. The
catalytic degradation is usually complete within 1–10 min
Scheme 2 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 1C.
ꢂc
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New J. Chem., 2009, 33, 34–49 | 43

ꢁ
Fig. 15 (A) UV-Vis spectra recorded for the reaction of 2 ꢀ 10ꢁ4 M Mn(NO3)2 with 0.01 M H2O2 in a 0.5 M HCO3 containing solution at
pH 8.4 and 25 1C. (B) Comparison of typical absorbance at 400 nm vs. time plots at pH 8.5 (black curve) and 9.5 (red curve).
Fig. 16 (A) UV-Vis spectra of a 2 ꢀ 10ꢁ5 M Mn(NO3)2 catalyzed oxidative degradation of 5 ꢀ 10ꢁ5 M Orange II by 0.01 M H2O2 in a 0.4 M total
carbonate containing solution at pH 8.5 and 25 1C. The inset in Fig. 16A shows the first spectrum of Orange II before the addition of the catalyst
and H2O2, and the final spectrum recorded after 250 s. (B) Comparison of absorbance at 480 nm vs. time plots for the 2 ꢀ 10ꢁ5 M Mn(NO3)2
catalyzed oxidative degradation of 5 ꢀ 10ꢁ5 M Orange II by 0.01 M H2O2 in a 0.4 M total carbonate containing solution at different pH values
and 25 1C.
depending on the pH of the solution, the catalyst concentra-
tion, and the H2O2 and carbonate concentrations. The decom-
position of the dye was followed by monitoring the spectral
changes at 480 nm. The depletion of the band at 480 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 p-system in
the dye molecule. The inset in Fig. 16A 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 isolation and
characterization of reaction products is extremely difficult and
requires large synthetic efforts, particularly as different reac-
tion intermediates tend to react further under experimental
conditions. A comparison of the reaction course at different
pH values is shown in Fig. 16B.
(see Fig. 17). Increasing the pH to 49 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 CO42ꢁ, a less electrophilic oxidant.58
ꢁ
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 Fig. 17). Thus, for a complete oxidation of these stable
molecules higher concentrations of oxidant and catalyst are
required.
A similar screening using MnCl2, Mn(Ac)2 and Mn(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 Z2-peroxycarbonate complex (see Fig. 15B).
The Mn(II) catalyzed decolorization and oxidative decom-
position 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
ꢂc
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44 | New J. Chem., 2009, 33, 34–49

Table 1 The constants k and K for the Mn(NO3)2 catalyzed oxidation
of Orange II by H2O2 at pH 8.5 and 25 1C (see Scheme 3)
[HCO3ꢁ]/M
k/sꢁ1
10ꢁ3K/Mꢁ1
0.1
0.3
0.4
0.5
0.0033
0.032
0.051
0.138
34.6
17.6
17.8
15.2
Scheme 3 Proposed reactions steps for the formation of the cata-
lytically active manganese intermediate in the presence of H2O2 in a
carbonate containing solution.
Fig. 17 Plot of observed rate constant (kobs) calculated for the
decoloring reaction followed at 480 and 300 nm, respectively. Experi-
mental conditions: 2 ꢀ 10ꢁ5 M Mn(NO3)2, 5 ꢀ 10ꢁ5 M Orange II,
0.01 M H2O2, 0.4 M total carbonate and 25 1C.
degradation of Orange II by H2O2 in a low carbonate con-
ꢁ
centration containing solution (0.1–0.3
M
HCO3
)
are
strongly curved (higher K values, see Table 1) and reach a
limiting value at higher catalyst concentration. In contrast,
similar data at higher carbonate concentrations (0.4–0.5 M
HCO3ꢁ) result in a less curved dependence of kobs on the
catalyst concentration, i.e. lower K values (see Table 1). The
observed rate profile can be explained by the general reaction
mechanism proposed in Scheme 2 and simplified in Scheme 3.
The observed rate law for the proposed reaction steps in
Scheme 3 is given by eqn (7). The calculated k and K values
from the non-linear concentration dependences in Fig. 18 are
summarized in Table 1.
Effect of the manganese concentration 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 Mn(NO3)2 was added in the presence of 0.01 M H2O2 to a
5 ꢀ 10ꢁ5 M Orange II solutiꢁon at 25 1C. 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 con-
centration dependence was studied at 480 nm using in situ
UV-Vis spectroscopic measurements and the kinetic traces
could be adequately fitted to a single exponential function.
Plots of the observed rate constant as a function of [Mn2+] at
different carbonate concentrations are presented in Fig. 18.
As it is evidenced in Fig. 18, the [Mn2+] dependences of the
observed rate constants for the manganese catalyzed oxidative
kK½MnðIIÞꢅ
kobs
¼
ð7Þ
1 þ K½MnðIIÞꢅ
Effect of the H2O2 concentration on the manganese-catalyzed
oxidative degradation of Orange II
The effect of H2O2 on the oxidation reaction course was
studied by varying its initial concentration over a wide range,
between 5 and 30 mM (Fig. 19). At lower H2O2 concentrations
(1 and 5 mM) 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 concentration effect on
the catalytic oxidation of the dye at higher concentrations of
H2O2. The kobs values were calculated from a single exponen-
tial 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.
Stability of the in situ formed catalyst
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 ꢀ 10ꢁ5 M Mn(NO3)2 at pH 8.5 (0.4 M HCO3
ꢁ
)
and 25 1C (see Fig. 20A and B).
As it can be seen in Fig. 20A, 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
Fig. 18 Mn(NO3)2 concentration dependence of kobs. Reaction con-
ditions: 5 ꢀ 10ꢁ5 M Orange II, 0.01 M H2O2, pH 8.5 and 25 1C.
ꢂc
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New J. Chem., 2009, 33, 34–49 | 45

By performing the reaction under inert reaction conditions
no significant differences in the decomposition reaction rate was
observed, indicating that HOꢆ or HOOꢆ radical formation
is not prevalent for this oxidation reaction. This is further
supported by the observation that addition of radical traps
such as TTBP had no effect on the reaction course (see Fig. 21).
Taking into account all obtained spectroscopic and kinetic
data, the following reaction schemes can be proposed for the
Mn2+ 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 4 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 MnQO intermediate and transfer
of the oxo group to the substrate. Most of the earlier reported
papers22,59 on the oxidation reaction catalyzed by several
isolated and structurally well defined manganese complexes
have emphasized the formation of a high-valent MnQO
intermediate by the reaction of manganese with the appro-
Fig. 19 H2O2 concentration dependence of kobs. Reaction conditions:
5 ꢀ 10ꢁ5 M Orange II, 2 ꢀ 10ꢁ5 M Mn(NO3)2, pH 8.5 and 25 1C.
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 evalu-
ated again by performing the oxidation reaction in the
presence of freshly added Orange II and H2O2. The experi-
mental results illustrated in Fig. 20B provide clear evidence for
the high efficiency of the in situ formed catalyst under the
above mentioned experimental reaction conditions.
ꢁ
priate oxidant. According to our observations, HCO3 ions
ꢁ
are involved in two catalytically relevant reactions. HCO3
ions react with aquated MnII present in solutioꢁn to form a
ꢁ
catalytically active Mn–HCO3 complex. HCO3 is also in-
ꢁ
volved in a fast equilibrium with H2O2 to form HOOCO2
,
Mechanistic aspects of the manganese-catalyzed oxidative
degradation of Orange II by H2O2 in carbonate solution
a versatile heterolytic oxidant. In the following step, through
nucleophilic attack of the oxidizing agent on the MnII center, a
MnII–Z2-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 MnQO inter-
mediates is believed to involve the heterolytic cleavage of the
peroxide bond, as shown in Scheme 4. An important role in
the stabilization of the formed MnQO species is played by the
electron donating bicarbonate ions. This may also account for
Throughout this study, the oxidation reactions were carried
out in a thermostated open glass reactor vessel at ambient
temperature in aqueous hydrogen carbonate containing solu-
tions. 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 different experimental conditions is illustrated in
Fig. 21.
ꢁ
the unique requirement of HCO3 in the oxidative decom-
position 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
Fig. 20 (A) Spectral changes observed at 480 nm for the repeated addition of 5 ꢀ 10ꢁ5 M Orange II to a 2 ꢀ 10ꢁ5 M Mn(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 ꢀ 10ꢁ5 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.
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46 | New J. Chem., 2009, 33, 34–49

Scheme 5 Proposed reaction mechanism involving first substrate
coordination to MnII 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 1C.
Fig. 21 Comparison of typical absorbance at 480 nm vs. time plots of
a 2 ꢀ 10ꢁ5 M Mn(NO3)2 catalyzed oxidative degradation of 5 ꢀ 10ꢁ5 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 atmos-
pheric oxygen (black curve), inert atmosphere (red curve) and TTBP
(blue curve), respectively.
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 4. In this case, the
formed MnIVQO intermediate is stabilized by Orange II, an
electron rich organic molecule with chelating capacity. The
importance of Orange II as an equatorial ligand is also to
favor the heterolytic scission of the peroxo bond leading to the
MnIVQO intermediate and bicarbonate.
Conclusions
A fast and environmentally benign method for the oxidative
degradation of Orange II could be achieved using H2O2 in
conjunction with catalytic amounts of relatively non-toxic
manganese salts as catalyst precursors in a carbonate contain-
ing aqueous solution under mild reaction conditions. Screen-
ing 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 peroxycarbonate as a key
molecular entity. The reported experimental data suggests that
the in situ formed high-valent manganese intermediate posses-
sing one hydrogen carbonate ligand is able to activate
H2O2, but decomposes rapidly with the formation of neutral
MnCO3, 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 forma-
tion of MnIIꢃ ꢃ ꢃOrange II complexes with different stoichio-
metry was observed. The simultaneous s,p-coordination of
the organic dye is well-precedented, and recent DFT studies
support this type of complex formation.28 The catalytic
Scheme 4 Proposed reaction mechanism for the Mn(II) catalyzed
oxidative degradation of Orange II by H2O2 in a carbonate containing
aqueous solution at pH between 8–9 and 25 1C.
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 nucleophilic attack of the dye
at the electrophilic oxygen of HOOCO2ꢁ. In aqueous solꢁution,
proton transfer can lead to the displacement of HCO3 and
the slow formation of oxidized substrate.
If substrate binding to MnII occurs before the addition of
HOOCO2ꢁ 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 5 involves the prior coordination of Orange II to MnII
and formation of MnII–Orange II complexes of different
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New J. Chem., 2009, 33, 34–49 | 47

activity of the formed intermediates was tested under catalytic
reaction conditions.
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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 deprotona-
tion of aquated Mn2+. The ongoing studies are presently
complemented by investigations on different organic sub-
strates with various functional groups in order to determine
the influence of substrate modification on the catalytic
reaction cycle. DFT studies beside further kinetic and spectro-
scopic investigations should contribute to a better understanding
of the catalytic system.
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Acknowledgements
The authors kindly acknowledge fruitful discussions with
Dr Anette Nordskog and Dr Wolfgang von Rybinski, Henkel
KGaA, Dusseldorf, Germany.
¨
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