ADSORPTION AND OXIDATION OF MICROPOLLUATANTS BY MANGANESE OXIDE

THESE
Pour l’obtention du Grade de
DOCTEUR DE L’UNIVERSITE DE POITIERS
(ECOLE NATIONALE SUPERIEURE d’INGENIEURS de POITIERS)
(Diplôme National - Arrêté du 7 août 2006)
Ecole Doctorale : Sciences pour l’Environnement Gay Lussac.
Secteur de Recherche : CHIMIE ET MICROBIOLOGIE DE L’EAU
Présenté par :

Thai Ha TRAN
************************

Adsorption and oxidation of micropollutants by Manganese oxide
************************
Soutenance prévue le 14 Décembre 2015
devant la Commission d’Examen
************************

JURY
Rapporteurs :

M. Philippe BEHRA
M. Emmanuel GUILLON

Examinateurs :

M. Patrick MAZELLIER
M. Sylvain OUILLON

Directeurs de Thèse :

M Hervé GALLARD
M. Jérôme LABANOWSKI

************************


ACKNOWLEDGEMENTS

I wish to thanks my advisors Hervé GALLARD and Jérôme LABANOWSKI for directing
their attentions forward to my work and welfare, and for allowing me freedom and flexibility in
research. Working with them is my honor. Their advices helped me to develop scientific skills and
will be useful for my future career.
I would also like to thank those who encouraged me to enter this work: my father, my mother,
my wife and my little daughter, my whole family. Their supports are always the best thing for me.
A number of fellow students and associates contributed to this work trough their supports,
advice and friendship: Alice TAWK, Virginie SIMON, Pamela ABDALLAH, Rose-Michelle
SMITH, Amal YOUSSOUF IBRAHIM, and many others.
I wish to thank Philippe BEHRA, Emmanuel GUILLON, Patrick MAZELLIER, Sylvain
OUILLON. They kindly served on my examining committees. Their comments and enthusiasm were
appreciated.
The staff of equip E1-Eaux Géochimie Santé and Platform APTEN went to great lengths to
assist with whatever problems arose, especially Audrey ALLAVENA, Sylvie LIU, Marie
DEBORDE.
Financial support from Vietnamese project 911 is gratefully appreciated.

ii


CONTENTS
List of Figures ....................................................................................................................... vi

List of Tables...........................................................................................................................x
General introduction ..............................................................................................................1
Chapter I: Literature review .................................................................................................4
I.1. Fate of organic micropollutants in the environment ...................................................4
I.2. Geochemistry and Manganese oxides mineralogy .....................................................6
I.2.1. Tunnel structure ..................................................................................................8
I.2.2. Layer structure ....................................................................................................9
I.2.2.1.Birnessite .......................................................................................................9
I.2.2.2.Others layer structures .................................................................................11
I.2.3. Other Mn oxides minerals.................................................................................11
I.3. Interactions between Manganese oxides and organic compounds ...........................12
I.3.1. Reaction Models ...............................................................................................13
I.3.1.1.Electron-transfer with bond-formation between metal sites and organic reductant
.......................................................................................................................14
I.3.1.2.Electron-transfer through formation of outer-sphere complex between metal sites
and organic reductant..............................................................................................14
I.3.2. Transformation reactions ..................................................................................15
I.3.2.1.Reactions with model organic compounds ..................................................15
I.3.2.1.1.Phenols..................................................................................................16
I.3.2.1.2.Anilines.................................................................................................17
I.3.2.1.3.Low molecular weight carboxylic acids ...............................................19
I.3.2.2.Reactions with organic contaminants ..........................................................19
I.3.2.2.1.Endocrine discruptors ...........................................................................19

iii


I.3.2.2.2.Antibacterial agents and antibiotics......................................................21
I.3.2.2.3.Other pharmaceuticals and industrial contaminants.............................25
I.3.2.3.Reactions with natural organic matter .........................................................26

I.3.3. Kinetic aspects and influence of reaction conditions........................................28
I.3.3.1.Kinetics and reaction order..........................................................................28
I.3.3.2.Effect of pH .................................................................................................30
I.3.3.3.Effect of mineral constituents......................................................................31
I.3.4. The role of NOM in NOM – micropollutants – MnO2 system .........................32
I.3.4.1.Interactions between NOM and micropollutants.........................................32
I.3.4.2.Effect of NOM on reaction of micropollutants by MnO2 ............................34
I.4. Application in water treatment and decontamination of polluted sites.....................36
Chapter II: Material and methods and preliminary study ..............................................39
II.1.

Manganese dioxide and reagents ......................................................................39

II.1.1.

Preparation of Manganese dioxide and characterization ..............................39

II.1.2.

Natural organic matters and micropollutants................................................40

II.2.
II.2.1.

Protocols ...........................................................................................................43
Kinetic experiments with MnO2 suspensions ...............................................43

II.2.2. Sorption experiments of NIS by NOM .............................................................44
II.3.

Analysis of micropollutants ..............................................................................44

II.4.

Identification of transformation products .........................................................45

II.5.

Natural Organic matter characterization by Dissolved organic carbon and UV

absorbance analysis and high-pressure size exclusion chromatography ........................46
II.6.

Results from preliminary study.........................................................................49

Chapter III: Adsorption and oxidation of the anthelmintic drug Niclosamide by birnessite 50
III.1.

Introduction.......................................................................................................50

iv


III.2.

Results and discussion ......................................................................................52

III.3.

Conclusion ........................................................................................................68

Chapter IV: Sorption and transformation of Pyrantel pamoate by synthetic birnessite70
IV.1. Introduction............................................................................................................70
IV.2. Results and discussion ...........................................................................................71
IV.3. Conclusion .............................................................................................................81
Chapter V: Oxidative transformation of triketone herbicide, sulcotrione, by manganese
oxide: kinetic, transformation products and impact of natural organic matter ............83
V.1. Introduction.............................................................................................................83
V.2. Results and discussion ............................................................................................85
V.3. Conclusion ..............................................................................................................99
Chapter VI: Degradation of Sulfamethoxypyridazine and cross-coupling reactions mediated
by MnO2 ..............................................................................................................................100
VI.1. Introduction..........................................................................................................100
VI.2. Results and discussion .........................................................................................102
VI.3. Conclusion ...........................................................................................................118
GENERAL CONCLUSION ..............................................................................................120

v


List of Figures
Chapter I
Figure I.1 – Schematic diagram highlighting potential sources and pathways for groundwater
pollution by micropollutants. Adapted from [18]. ................................................................................ 4
Figure I.2 – Cristalline structure of (A) Pyrolusite, (B) Ramsdellite, (C) Hollandite, (D)
Romanecdite, and (E) Todorokite [42] ................................................................................................. 8
Figure I.3 – Cristal structure of (A) Lithiophorite, (B) Chalcophanite, (C) Na-rich Birnessite like
[42]...................................................................................................................................................... 11
Figure I.4 – General view of transformation of organic compound by MnO2 ................................... 13
Figure I.5 – Major reactions involved in the oxidation of phenols by δ-MnO2 and accumulation of

reduced Mn species on the mineral surface. The final aqueous products are shown in blue [32]...... 16
Figure I.6 – Postulated mechanism for oxidative coupling of aniline by reactions with manganese
oxide. The reaction proceeds from a cation radical through coupling products, which then undergo
further oxidation. Adapted from [77].................................................................................................. 18
Figure I.7 – Transformation of citrate by MnO2 ................................................................................. 19
Figure I.8 – Proposed reaction pathway of oxidation of triclosan by MnO2 [63] .............................. 21
Figure I.9 – Proposed Reaction Scheme for Oxidation of FQs by MnO2 [106]................................. 23
Figure I.10 – Proposed mechanism for the oxidative transformation of DCF by manganese oxide
[112].................................................................................................................................................... 26
Figure I.11 – Cross-coupling of sulfamethazine with syringic acid [41]. .......................................... 27
Figure I.12 – Impact of NOM on oxidation and hydrolysis of micropollutant by MnO2. Adapted
from [119] ........................................................................................................................................... 34


Chapter II
Figure II.1 – Zeta potential of Manganese oxide suspension for pH range 2.5 – 6.0 and ionic
strengths of 1 and 10 mM NaNO3....................................................................................................... 40
Figure II.2 – HPSEC calibration curve obtained with polystyrene sulfonate standards..................... 47
Figure II.3 – Typical HPSEC chromatograms for the three different NOM extracts......................... 48
Chapter III
Figure III.1 – Concentration profiles of NIS after centrifugation or reduction by ascorbic acid
([NIS]o = 130 nM , 10 mM acetate buffer pH 5.0). ............................................................................ 53
Figure III.2 – Proposed pathway for the catalytic hydrolysis of NIS by MnO2 ................................. 55
Figure III.3 – Linear adsorption isotherms of Niclosamide at pH 4.0, 4.5, 5.0 and 5.5 ..................... 57
Figure III.4 – First-order kinetic representation of NIS transformation for different MnO2
concentrations (pH 5.0, 10 mM acetate buffer). ................................................................................. 59
Figure III.5 – First-order dependence with respect to MnO2 at pH 5.0 .............................................. 60
Figure III.6 – Dependence of k with respect to H+ (10 mM acetate buffer, 130 nM [NIS]0 and 100
µM [MnO2]0)....................................................................................................................................... 62
Figure III.7 – Evolution of NIS concentration in the absence and presence of CR-NOM ................. 64

Figure III.8 – Effect of CR-NOM concentration on (a) observed rate constants of NIS transformation
and (b) adsorption isotherms (pH 5.0, 175 µM MnO2 and 130 nM NIS)........................................... 65
Figure III.9 – Changes in 2D fluorescence spectra of CR-NOM with increasing NIS concentration 66
Figure III.10 – Application of Stern-Volmer model for interaction between NIS and NOM (pH
5.0±0.1 and 22±1 °C, 2.2 mg-C L-1, 255/400 nm excitation/emission wavelengths)......................... 67
Figure III.11 – HPSEC/UV chromatograms of 2.0 mgC L-1 NOM before and after contact with 500
µM MnO2 at pH 5.0 ............................................................................................................................ 68

vii


Chapter IV
Figure IV.1 – Behaviour of Pyrantel Pamoate in presence of birnessite ([MnO]2 = 500 µM, [PMA]0
= 260 nM, [PYR]0 = 260 nM 10 mM acetate buffer pH 5.0) ............................................................. 72
Figure IV.2 – Influence of initial concentrations of MnO2 on removal of PMA ([PMA]0 = 260 nM,
pH 5.0, 10 mM acetate buffer)............................................................................................................ 73
Figure IV.3 – First-order dependence with respect to MnO2.............................................................. 73
Figure IV.4 – Dependence of kobs with respect to H+ concentration .................................................. 75
Figure IV.5 – Effect of SR-NOM on oxidation of PMA by MnO2 .................................................... 76
Figure IV.6 – Proposed reaction pathway for the degradation of PMA by MnO2 ............................. 77
Figure IV.7 – Concentration profile of PYR in presence of MnO2 after quenching by ascorbic acid or
centrifugation ...................................................................................................................................... 79
Figure IV.8 – Sorption isotherm of PYR by MnO2 at pH 5.0 ............................................................ 80
Figure IV.9 – Effect of NOM on sorption kinetic of PYR by MnO2 ([MnO2]0 = 1 mM, pH 5.0,
[NOM]0 = 0.5 mgC L-1) ...................................................................................................................... 81
Chapter V
Figure V.1 – Degradation kinetics of triketone herbicides tembotrione and sulcotrione by MnO2 ... 85
Figure V.2 – Influence of quenching mode on kinetics of SCT oxidation by MnO2. ([MnO2]o = 217
µM; [SCT]o = 6.3 µM; pH 5.0 10 mM acetate buffer). ...................................................................... 86
Figure V.3 – First order representation of SCT degradation for different MnO2 concentrations

([SCT]o = 6.3 µM; pH 5.0 10 mM acetate buffer).............................................................................. 87
Figure V.4 – First-order dependence with respect to MnO2 ............................................................... 89
Figure V.5 – Dependence of k with respect to H+ .............................................................................. 90
Figure V.6 – Effect of NOMs isolates on SCT transformation by MnO2........................................... 91
Figure V.7 – Effect of NOM concentration on observed rate constants of SCT transformation ....... 92
Figure V.8 – HPSEC-UV chromatograms of SR-NOM before and after contact with birnessite...... 93

viii


Figure V.9 – Reaction pathway for SCT transformation by MnO2 .................................................... 98
Chapter VI
Figure VI.1 – Influence of quenching mode on the time profile of SMP in presence of MnO2.
([MnO2]0 = 250 µM; [SCT]0 = 1.2 µM; pH 5.0 10 mM acetate buffer) ........................................... 102
Figure VI.2 – Logarithmic representation of SMP transformation by birnessite showing deviation
from first order ([SMP]o = 1.1 µM; pH 5.0 10 mM acetate buffer).................................................. 103
Figure VI.3 – Determination of the order of the initial reaction rate of SMP degradation .............. 106
Figure VI.4 – Retarded model fit for reaction of SMP oxidation by MnO2 ..................................... 109
Figure VI.5 – Effect of different concentrations of CR-NOM on transformation of SMP by MnO2.
(The curves are the retarded model fit)............................................................................................. 111
Figure VI.6 – Effect of syringic additions on SMP degradation by MnO2....................................... 113
Figure VI.7 – Proposed pathway for SMP transformation by MnO2 ............................................... 117
Figure VI.8 – Cross-coupling reaction between SMP and SYR mediated by MnO2 ....................... 118

ix


List of Tables
Chapter I
Table I.1 – Nomenclature and chemical formula of some manganese oxides [42].............................. 7

Table I.2 – Proposed methods for synthesis of Birnessite .................................................................. 10
Table I.3 – Summary of model phenols, anilines and low molecular weight (LMW) acids used to
dissolve manganese oxides [33], [75]–[77]. ....................................................................................... 15
Table I.4 – Reaction order for the transformation of various organic compounds by MnO2. ............ 29
Chapter II
Table II.1 – Physico-chemical properties of micropollutants tested in this study .............................. 42
Table II.2 – Experimental conditions used in this study..................................................................... 43
Table II.3 – Conditions used for the analysis of micropollutants by HPLC/UV................................ 45
Chapter III
Table III.1 – MS spectra of NIS and its transformation products....................................................... 56
Table III.2 – Pseudo-First order kinetic model constant for the transformation of NIS by MnO2 in
10mM acetate buffer ........................................................................................................................... 60
Chapter IV
Table IV.1 – Pseudo-first order kinetic rate constants for the transformation of PMA by MnO2 in 10
mM acetate buffer at pH 5.0 ............................................................................................................... 74
Table IV.2 – MS spectra of PMA and its transformation products .................................................... 78


Chapter V
Table V.1 – Pseudo-first order kinetic constant for the transformation of SCT by MnO2 in 10 mM
acetate buffer....................................................................................................................................... 88
Table V.2 – Effect of NOM on pseudo-first order kinetic model constant for the transformation of
SCT by MnO2 at pH 5.0 (430 µM MnO2, 6.3 µM SCT) ................................................................... 92
Table V.3 – DOC values in mgC L-1 of NOM solutions before and after incubation with MnO2 (pH
5.0, 500 µM MnO2)............................................................................................................................. 94
Table V.4 – Tranformation products formed from degradation of SCT by MnO2 ............................. 96
Chapter VI
Table VI.1 – Initial rate for the transformation of SMP by MnO2 in 10 mM acetate buffer............ 107
Table VI.2 – Retarded model parameters for the reaction of SMP with MnO2................................ 110
Table VI.3 – Retarded model parameters for the reaction of SMP with MnO2................................ 111

Table VI.4 – Tranformation products formed by SMP oxidation by MnO2 ..................................... 114
Table VI.5 – Cross–coupling products formed from reaction between SMP and SYR mediated by
MnO2 ................................................................................................................................................. 115

xi


General introduction
Manganese (Mn) oxides are naturally present in soils and sediments. These oxides play an
important role in controlling biogeochemical cycles. Interactions of manganese oxides with
organic and inorganic compounds have been the subject of a number of environmental studies.
Their adsorptive and oxidizing properties allow interactions with organic compounds.
Environmental studies conducted with synthetic manganese oxides show that these manganese
oxides can react with model organic compounds such as phenols and anilines. Recent studies have
shown that manganese oxides can contribute to the elimination of organic contaminants such as
pesticides, hormones or antibacterial agents in the environment. These oxidative processes
involving Mn oxides may constitute an important abiotic degradative pathway for organic
compounds in subsurface environment. Additionally some authors have discussed the use of
manganese oxides as a solution for the treatment of emerging contaminants present in waste water
or in polluted sites.
A large number of synthetic organic compounds are used by modern society for a wide
variety of purposes including the production and preservation of food, industrial manufacturing
processes, and human and animal healthcare. The presence of these organic compounds has been
widely reported in many ecosystems and their reactivity in the environment has been the subject of
many studies. Because of most of these compounds are environmentally present at low
concentration ranges from pg L-1 to ng L-1, they are often named micropollutant. The presence and
the concentrations of organic micropollutants are important in the quality evaluation of surface
and ground waters. These pollutants can be persistent, bioresistant compounds or they may
undergo biotransformation, photo-oxidation or/and hydrolysis. In some cases, transformation
products can be more toxic than the parent compound. Similar studies have been conducted to

evaluate their behaviour during water treatment. The use of strong oxidants such as ozone and
adsorbents such as activated carbon can remove micropollutant residues during production of
drinking water. However, the conventional waste water treatment plants using activated sludge


cannot always ensure a complete removal of these compounds. Effluent discharges from waste
water treatment plants, therefore, are the main source of contamination by micropollutants
residues. Tertiary treatments are currently being studied to reduce this contamination. Potential
treatments are ozone, activated carbon or membrane processes such as nanofiltration or reverse
osmosis when water reuse is considered. The implementation of these treatments is costly and is
not considered without new regulations for waste water discharges. Alternative methods are also
being investigated like the use of advanced oxidation processes or oxidants such as ferrate. Some
studies also suggested the use of manganese oxides, alone or in combination with biological
process.
The main objective of this thesis is to investigate the interactions of selected micropollutants
towards synthetic manganese oxide birnessite. The influence of experimental conditions such as
the presence of natural organic matter (NOM) was also evaluated. The interactions between
natural organic matter and oxides were studied because NOM are ubiquitous in the environment
and are present in higher concentrations than micropollutants. A preliminary study was performed
in order to select molecules that will react with manganese oxides. Four molecules were then
choosen for kinetic study and identification of transformation products.
The manuscript is organized into six chapters.
The first chapter is a literature review. This section presents the geochemistry of manganese
oxides and the state of the art describing the transformation of organic compounds by manganese
oxide with a focus on organic micropollutants.
The second chapter described the analytical and experimental methods employed over the
study.
The next four chapters are written as projects of publication.
The third chapter focuses on the adsorption and degradation of the hydrophobic
antihelminthic drug Niclosamide (NIS) by MnO2. The aim was to investigate the role of NOM as

competitor for oxidation and adsorption but also the ability of NOM as sorbent of hydrophobic
micropollutants in presence of a mineral oxide.
2


The fouth chapter presents the reactivity of birnessite towards the deworming agent, pyrantel
pamoate.
The fifth chapter presents study on the transformation of triketone herbicide, Sulcotrione
(SCT) by manganese oxide. The oxidation kinetics of SCT was followed for different initial
conditions and transformation products were identified by LC/MSn. Degradation mechanism of
SCT has been proposed.
The

sixth

chapter

deals

with

the

transformation

of

sulfonamide

antibiotic,

sulfamethoxypyridazine (SMP), by manganese oxide. Transformation products were identified by
LC/MSn. Degradation mechanism of SMP has been proposed. A retarded kinetic model including
the influence of the natural organic matter acting as a competitor in the system is presented. The
cross-coupling reaction between SMP and Syringic acid, a model NOM constituent, mediated by
MnO2 was also presented.
Finally, this paper ends with a general conclusion summarizing the essential advances
obtained in our study.

3


Chapter I: Literature review
I.1.

Fate of organic micropollutants in the environment
In the last few decades, the occurrence and fate of organic micropollutants in water bodies

including veterinary drugs, endocrine disrupting chemicals, pharmaceuticals, personal care
products and antibiotics have been extensively studied due to their possible adverse effects to
wildlife and humans [1]–[5]. It is now established that these compounds enter the environment
from a number of sources: livestock activities including waste lagoons and manure application to
soil [6]–[8]; subsurface storage of household and industrial wastes [9] [10], as well as indirectly
through the process of groundwater–surface water exchange [11]; wastewater effluents from
municipal treatment plants [12]–[15]; septic tanks [16], [17]; hospital effluents [1]. All described
pathways were illustrated in Figure I.1.

Figure I.1 – Schematic diagram highlighting potential sources and pathways for groundwater
pollution by micropollutants. Adapted from [18].

4


Once a micropollutant is released into the environment, key physico-chemical properties –
such as water solubility and volatility – will influence its behavior. In surface water, the main
elimination processes are hydrolysis, biodegradation, sorption, and photodegradation.
Organic micropollutants can be degraded through biotic transformations in soils and water.
Generally, these processes reduce the potency of the drugs. However, some transformation
products have similar toxicity to their parent compound [19], [20]. Bacteria and fungi are the two
major groups of microorganisms responsible of organic compounds biotransformation. Fungi are
particularly important in soils, but do not usually play an important role in the aquatic
environment. Therefore, in sewage treatment plants, surface and ground water bacteria are
assumed to be responsible for most of the biodegradation processes. Even though microbial
degradation is enhanced in waste water treatment systems due to the higher bacterial density,
recalcitrant micropollutants were detected in secondary waste water effluents and polishing
treatments such as ozone and activated carbon are proposed to fully eliminate these compounds
and reduced the toxicity of effluents [3]. For example, more than 20 antibiotics representing the
most important groups of antibiotics were found not to be readily biodegradable [1], [21], [22].
Sorption may also have an impact on the spread and bio-availability of micropollutants in
the environment [23]–[29]. Sorption coefficients (Kd) for micropollutants in soils and sediments
range from 0.2 for chloramphenicol in marine sediment to 5610 L kg-1 for enrofloxacin in soil [6].
Values vary considerably for a given compound in different soils [6]. Mechanisms other than
hydrophobic partitioning, such as cation exchange, cation bridging at clay surfaces, surface
complexation and hydrogen bonding play a role in the sorption of organic micropollutants [25].
Also, the sorption of selected micropollutants may depend heavily on pH and ionic strength [26].
The highest values of Kd,solid ranged between 70 and 5000 L kg-1 for tetracycline and quinolone
carboxylic acid antibiotics. According to a classification of pesticide mobility in soil [25], these
micropollutants can be considered to be immobile. Intermediate Kd,solid for avermectin, tylosin, and
efrotomycin ranged between 7 and 300 L kg-1. Low values of Kd,solid ranging between 0.2 and 2 L
kg-1 are for micropollutants having low sorption affinity to soil particles such as olaquindox,

5


sulfamethazine, sulfathiazole, metronidazole, chloramphenicol. The latter two groups of
micropollutants can be considered to be low to slightly mobile (50 L kg-1 > Kd,solid > 5 L kg-1) and
medium to highly mobile (Kd,solid < 5 L kg-1), respectively [25].
Many micropollutants are also expected to be photoactive because many of these
compounds feature aromatic rings, heteroatoms, and other functional groups that can either absorb
solar radiation or react with photogenerated transient species in natural waters. Thus, photolysis
and photochemical processes should be considered as important removal mechanisms of
micropollutants in surface waters [30], [31].
In recent years, the sorption and abiotic transformation of organic micropollutants by metal
oxides has been in the scope of several studies [32]–[34]. The ability of manganese oxides to
participate in the fate of micropollutants in the environment has been extensively studied [35]–
[37]. The use of manganese oxides in water treatment to remove organic micropollutants [38],
[39] and to decontaminate polluted soil [40], [41] has then be proposed. From this point of view,
the reactivity between manganese oxides and organic micropollutants will be the object of this
literature review. General information on the geochemistry and mineralogy of manganese oxides
will first be given in the next paragraphs.
I.2.

Geochemistry and Manganese oxides mineralogy
Manganese is one of the most abundant elements in the earth and the second transition metal

after iron present in the earth’s crust [42]–[44]. Humanity has started using manganese for
thousand years as pigment in cave and to clarify glass. Nowadays manganese is mostly used as
catalysts and for the production of battery [45].
There are more than 30 Mn oxides/hydroxide minerals, and many of them occur abundantly
in a wide variety of geological settings. In addition to being important as ores of Mn metal, they
also play an active role in the environmental geochemistry at the Earth’s surface. Manganese

oxides and hydroxides are important constituents of the soil and sediments, and because they are
highly chemically active and strong scavengers of heavy metals, they exert considerable

6


influences on the composition and chemical behavior of the sediments, soils and associated
aqueous systems. They participate in numerous chemical reactions with constituents of the soil
and/or groundwater. On the other hand, manganese is an essential nutrient for plants and animals.
Manganese participates in the synthesis of enzyme and is a cofactor involved in carbohydrate and
nitrogen metabolism [46]. It was reported that at high concentrations manganese is toxic to plant
growth [47].
Many studies were carried out to understand the crystallography of mineral phases and the
geochemistry of manganese [42]–[44], [48]. In the environment, manganese exists in three
different oxidation states: +2, +3, +4. The oxides can be regrouped in three types: tunnel, layer
structure and others. The basic structural units for Mn oxides are MnO6 octahedra. Listed in Table
I.1 are the informations for some of the most important Mn oxide minerals and their chemical
formula.

Table I.1 – Nomenclature and chemical formula of some manganese oxides [42]
Name

Symbol

Chemical formula

Structure

Birnessite

δ-MnO2

(Na,Ca)Mn7O14.2.8H2O

Layer

ZnMn3O7.3H2O

Layer

Chalcophanite

4+

3+

Coronadite

α-MnO2

Pbx(Mn ,Mn )8O16

Tunnel

Cryptomelane

α-MnO2

Kx(Mn4+,Mn3+)8O16

Tunnel

Hollandite

α-MnO2

Bax(Mn4+,Mn3+)8O16

Tunnel

LiAl2(Mn24+,Mn3+)O6(OH)6

Layer

Lithiophorite
Manganite
Nsutite

γ-MnO2
γ-MnO2

MnOOH
Mn(O,OH)2

Tunnel

MnO2

Tunnel

Ramsdellite

MnO2

Tunnel

Romanechite

Ba0.66Mn4+3.68Mn3+1.32O10x1.34H2O

Tunnel

Todorokite

(Na,Ca,K)x(Mn4+,Mn3+)6O12.3.5H2O Tunnel

Pyrolusite

ρ-MnO2
β-MnO2

7


I.2.1. Tunnel structure
The tunnel Mn oxides are constructed of single, double, triple chains of edge-sharing MnO6
octahedra, and the chains share corners with each other to produce framework that have tunnel
with square or rectangular cross section. The larger tunnels are partially filled with water
molecules and/or cations. Pyrolusite, Ramsdellite, Hollandite, Romanecdite, and Todorokite
belongs to this group of Mn oxides. Figure I.2 illustrated different types of crystalline tunnel

structure.

Figure I.2 – Cristalline structure of (A) Pyrolusite, (B) Ramsdellite, (C) Hollandite, (D)
Romanecdite, and (E) Todorokite [42]
Hollandite group includes hollandite, cryptomelane, coronadite, and manjiroite. This group
is constructed of double chains of edge-sharing MnO6 octahedra, but they are linked in such a way
as to form tunnels with square cross sections and two octahedra for each side (Figure I.2). The
tunnels are partially filled with large uni- or divalent cations and, in some cases, water molecule.
These manganese oxides differ by the presence of cation K, Ba and Pb inside the structure [49].
Cryptomelane is generally synthesized by dropwise addition of HCl to KMnO4. Then the solution

8


is washed, dried, and iginited at 400 oC for 60 hours, and again washed with water. This gave
cryptomelane containing 7.2% K, with a surface area of 58 m2 g-1 [50].
Pyrolusite is the most stable form of Mn oxide minerals with a tunnel structure. In pyrolusite
(β-MnO2), single chains of edge-sharing MnO6 octahedra share corners with neighboring chains to
form a framework structure containing tunnels with square cross sections. Each square cross
section is one by one octahedron (1x1) on a side (Figure I.2) [42].
The ramsdellite β-MnO2 is one rare structure of MnO2. In the ramsdellite structure the MnO6
octahedra are linked into double chains, each consisting of two adjacent single chains that share
octahedral edges. The double chains, in turn, link corners with each other to form a framework
having tunnels with rectangular shaped cross sections that are 1x2 octahedra on a side (Figure I.2)
[42], [44], [50].
Romanechite is a valuable ore of manganese, which is used in steelmaking. The romanechite
structure is constructed of double and triple chains of edge-sharing MnO6 octahedra that link to
form large tunnels with rectangular cross sections, measuring two by three octahedra (Figure I.2).
The tunnels are filled with Ba cations and water molecules in a 1:2 ratio, and the charges on the
tunnel cations are balanced by substitution of Mn3+ for some of the Mn4+ [42], [51], [52].

Todorokite is a rare complex hydrous manganese oxide mineral. Todorokite is made up of
MnO6 octahedra that share edges to form triple chains. Nsutite (γ-MnO2) is an important cathodic
material for use in dry-cell batteries [42]. Although classified as a mineral, nsutite is actually an
intergrowth between pyrolusite and ramsdellite.
I.2.2. Layer structure
Layer Manganese oxides as birnessite consist of stacked sheets of edge-sharing Mn
octahedral.
I.2.2.1.

Birnessite

Birnessite was first described as a natural phase from Birness [42], and since then it has been
recognized that birnessite and birnessite-like minerals occur in a wide variety of geological
9


settings. It is a major phase in many soils and an important component in desert varnishes and
other coatings and in ocean Mn nodules [53]. It readily participates in oxidation-reduction and
cation-exchange reactions and therefore plays a significant role in soil and groundwater chemistry.
Birnessite has a two-dimensional layered structure that consists of edgeshared MnO6
octahedra with water molecules and alkali metal cations or protons occupying the interlayer space.
The interlayer spacing is typically about 0.7 nm [54], [55]. The heterovalent Mn cations (i.e., Mn3+
and Mn4+) result in a net negative charge for the MnO2 basal layers, which is balanced by the
interlayer cations. The chemical composition of birnessite can be expressed by the general formula
AxMnO2 · yH2O, where A is H+ or a metal cation such as K+, Na+, and Ca2+[56]. Birnessite
materials have received increasing attention in recent years, owing to their wide range of
applications as ion-exchangers, selective adsorbents, catalysts, electrode materials [56]. Table I.2
presented several methods to synthesize birnessite in laboratory conditions.

Table I.2 – Proposed methods for synthesis of Birnessite

Methods
Oxidation
manganous
hydroxide

Procedure

Ref.

of A mixture of 0.4 moles of manganous sulfate and 5.5 moles of described
potassium hydroxide in two litres of water was cooled to 5°C by
and oxidized by bubbling oxygen for 5 hours. This produced a McKenzie
black birnessite containing 9.0% K, with a surface area of 75 [50]
m2g-1

Reduction
of Two moles of concentrated HCl was added dropwise to a
potassium
boiling solution of one mole of potassium permanganate in
permanganate
2.5 litres of water with vigorous stirring. After boiling for a
further ten minutes, the precipitate was filtered and washed.

described
by
McKenzie
[50]

This gave a brown birnessite with a potassium content of
9.5% K, and a surface area of 32 m2g-1.

Reduction
of Eight mmoles of MnO4- and 16 mmoles of NaOH were added
potassium
in 3 litre deionized water, followed by a dropwise addition of
permanganate
1.5 litre MnCl2 solution (12 mM). The MnO2 particles were
collected by decantation, centrifugation and the supernatant
was replaced by deionized water several times until the
conductivity was significantly reduced. This birnessite has
pHzpc of 2.25 with surface area value of 270 m2g-1.
10

described
by
Murray
[53]


I.2.2.2.

Others layer structures

The lithiophorite structure consists of a stack of sheets of MnO6 octahedra alternating with
sheets of Al(OH)6 octahedra in which one-third of the octahedral sites is vacant (Figure I.3). In the
ideal formula, Li cations fill the vacant sites in the Al layer, and charge balance is maintained by
substitution of an equal number of Mn3+ for Mn4+ cations [42]. The layers are cross-linked by
hydrogen bonds between hydrogen of the hydroxyl groups on the Al/Li layer and O atoms in the
Mn sheet.
Chalcophanite is a common weathering product in many Mn-bearing base metal deposits. Its
structure consists of sheets of edge-sharing MnO6 octahedra that alternate with layers of Zn

cations and water molecules (Figure I.3). One of the seven octahedral sites in the Mn layer is
vacant, and the Zn cations are above and below the vacancies [42]. The water molecules form a
hexagonal close-packed layer with one of the seven molecules absent.

Figure I.3 – Cristal structure of (A) Lithiophorite, (B) Chalcophanite, (C) Na-rich Birnessite like
[42]
I.2.3. Other Mn oxides minerals
Hausmannite [Mn2+Mn3+2O4] has a spinel-like structure with Mn2+ in the tetrahedral and
Mn3+ in the octahedral sites. Hausmannite and bixbyite [(Mn,Fe)2O3] typically are found in
hydrothermal or metamorphic deposits. The amount of Fe that can be accommodated into the

11


bixbyite structure is a function of temperature [57], and therefore the mineral is an important
geothermometer in some ore deposits. The crystal structure of pyrochroite [Mn(OH)2] consists of
stacked sheets of Mn2+(OH)6 octahedra, and manganosite (MnO) is isostructural with halite. Both
minerals are relatively rare, typically occurring in low-temperature hydrothermal veins in Mn-rich
deposits [50].
I.3.

Interactions between Manganese oxides and organic compounds
Manganese oxides rank among the strongest natural oxidant in soil and sediments [32], [53].

The standard reduction potential of MnO2 at pH 7 and 25°C is 1.29 V [53] (equation (I.1)). With
large surface area up to 270 m2 g-1 (Table I.2), manganese oxides can sorb and further transform
organic micropollutants via direct oxidation [33] and/or surface catalysis [58].
1
1
MnO2(s) + 2H+ + e- → Mn2+(aq) + H2O

2
2

(I.1)

As shown in Figure I.4, organic micropollutants are hypothezised to be transformed by
manganese oxides via the following steps:
(1) Sorption of organic compounds onto oxide surfaces, (2) formation of a precursor
surface complex, (3) hydrolysis and/or electron transfer inside the complex to form transformation
producs and Mn2+, (4) the formed product can be sorbed or desorbed and possibly further
transformed by manganese oxides.
The sorption capacity of manganese oxides without further transformation was documented
for several organic compounds (chlorpheniramine [59], ciprofloxacin [60] and benzoic acid [61]).
Both adsorption and transformation of macrolides [62], triclosan and chlorophen [63], tetracycline
[64], [65] were also reported. Reduction of manganese oxides by ascorbic or oxalic acids and
separation by centrifugation or filtration are used to determine the amount of organic compounds
adsorbed onto MnO2 surface and transformed by MnO2. The reduction of MnO2 allows
determining the total amount of organic compounds i.e. in solution and sorbed onto MnO2 surface.
Centrifugation allows determining the amount of organic compound in solution. The difference
between both methods gives the organic amount sorbed onto MnO2 surface. Compared to

12


centrifugation, reduction of MnO2 by ascorbic acid releases the adsorbed organic compound from
oxide surface and is thus a strategy to figure out the rate-limiting step. Electron transfer can be
considered as the rate-limiting step if unreacted sorbed organic compounds are detected on MnO2
surface. This behaviour was reported for transformation of triclosan [63], tetrabromobisphenol
[66], N-oxides [67]. Conversely, if unreacted sorbed organic is determined with no detectable
amount, then adsorption is the rate-limiting step. Transformation of lincosamide [68], tetracycline

[69] showed this behaviour.

Figure I.4 – General view of transformation of organic compound by MnO2

I.3.1. Reaction Models
Two mechanisms are usually considered to describe the oxidation of organics by metal
oxides: bonded and nonbonded electron transfer [33].

13


I.3.1.1.

Electron-transfer with bond-formation between metal sites and organic

reductant
This mechanism involves direct bonding between the metal center and the organic reductant
i.e. inner-sphere complex formation prior to electron transfer. Electron transfer at the oxide
surface complex is represented by the following reactions:

≡MnOH + HA

1a
→

←

k
k

−1 a

(I.2)

≡MnA + H2O

k2
≡MnA →
Products

(I.3)

HA is the organic reductant, ≡MnOH is a free oxide surface site, and ≡MnA an inner-sphere
surface complex. When this mechanism occurs, relative rates of reaction with different organic
substrates will reflect, in part, bonding requirements of Mn oxide surface sites. An organic
reductant with high affinity for surface sites will react more quickly than one with low affinity,
providing that rates of electron transfer are similar. This mechanism implies that the organic
reductant is chemically bonded to reactive surface sites in the precursor complex. In previous
studies catechol [70], pentachlorophenol [71], glyphosate [72] were reported to form inner-sphere
complex with birnessite. In the inner sphere complex, the organic is partially dehydrated and
directly bound to the surface, whereas the outer sphere complex, the organic retains its
hydradation sphere and attaches to the surface [73].

I.3.1.2.

Electron-transfer through formation of outer-sphere complex between

metal sites and organic reductant
Reaction via a non-bonded mechanism, in contrast, involved outer-sphere complex
formation prior to electron transfer, and no direct bond between oxidant and reductant is formed.

For example acid orange 7 [74] , glyoxylic acid [75] would form outer-sphere complex with
manganese oxides. In the following scheme, (≡MnOH, HA) represents the outer-sphere precursor
complex.

14