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Radiation Induced
Processes at SolidLiquid Interfaces
MIAO YANG
杨淼
KTH Royal Institute of Technology
School of Chemical Science and Engineering
Department of Chemistry
Applied Physical Chemistry
SE-100 44 Stockholm, Sweden
Copyright © Miao Yang
All rights reserved
Paper I © 2013 Elsevier B. V.
Paper II © 2013 Elsevier B. V.
Paper III © 2014 American Chemical Society
Paper IV © 2015 Elsevier B. V.
Paper VIII © The Royal Society of Chemistry 2015
TRITA-CHE Report 2015:39
ISSN: 1654-1081
ISBN: 978-91-7595-656-5
Akademisk avhandling som med tillstånd av KTH Kungliga Tekniska
Högskolan i Stockholm framlägges till offentlig granskning för avläggande
av teknologie doktorsexamen fredagen den 18 September 2015, kl 10.00
i sal F3, KTH, Lindstedtsvägen 26, Stockholm. Avhandlingen försvaras
på engelska. Opponent: Prof. Simon Pimblott från The University of
Manchester, UK.
ii
To my family
If we knew what it was we were doing, it would not be
called research, would it?
Albert Einstein
iii
iv
Abstract
In the thesis, the reactions between water radiolysis products—H2O2,
HO• and O2—with metals and metal oxides utilized in nuclear industry are
studied. The reactions include not only surface reactions, e.g. redox
reactions and catalytic decomposition of H2O2, but also solution reactions
(Haber-Weiss reactions). To study the interfacial reactions, it is crucial to
monitor the dissolution of the solid material, reactivity of H2O2 and
formation of the intermediate hydroxyl radicals.
Hydroxyl radicals are captured by probe (Tris or methanol) to generate
CH2O which can be quantified by the modified Hantzsch method. The
results from γ-irradiation experiments on homogeneous system show that
the conversion yield of CH2O from hydroxyl radicals is affected by O2 and
pH. A mechanism of CH2O production from Tris is proposed.
Besides, the consumption rate of H2O2 in the H2O2/ZrO2/Tris system is
found to be influenced by Tris. A mechanism for the catalytic
decomposition of H2O2 upon ZrO2 surface is proposed which includes
independent surface adsorption sites for H2O2 and Tris. Moreover, it is
demonstrated that the deviation of detected CH2O concentration by the
modified Hantzsch method from actual concentration increases with
increasing [H2O2]0/[CH2O]0.
The inhibition of sulfide on the radiation induced dissolution of UO2 is
confirmed and is dependent on sulfide concentration. And the inhibition
of sulfide is independent to that of H2/Pd.
It is found that the reactivity of H2O2 and dynamics of CH2O formation
are different for the studied materials in the H2O2/MxOy/Probe system.
The kinetic parameters, such as rate constant, activation energy,
frequency factors are determined.
Both surface and solution reactions are observed in the aqueous
W(s)/H2O2/Tris system. It is also demonstrated that Haber-Weiss
reactions which produce HO• continuously are dominating. Furthermore,
it is found that hydroxyl radicals are formed simultaneously during the
dissolution of W in aerobic aqueous system.
The knowledge conveyed by the thesis is relevant to nuclear technology,
as well as the applications related in photocatalysis, biochemistry,
corrosion science, catalysis and optics/electronics.
v
Sammanfattning
I denna avhandling presenteras experimentella studier av reaktioner
mellan radiolysprodukterna H2O2, HO• och O2 och metaller och
metalloxider, förekommande i kärntekniska tillämpningar. Reaktionerna
inkluderar ytreaktioner, t.ex. redoxreaktioner och katalytiskt sönderfall
av H2O2, samt reaktioner i lösning (Haber-Weiss). För att kunna studera
ovanstående ytreaktioner är det viktigt att studera upplösning av det fasta
materialet, försvinnandet av H2O2 och produktion av hydroxylradikaler.
Genom att tillsätta Tris eller metanol kan hydroxylradikalen fångas in
under bildande av en stabil produkt som lätt kan detekteras (formaldehyd)
genom en metod som kallas Hantzschmetoden. Resultat från homogena
gammabestrålningsexperiment visar att utbytet av formaldehyd påverkas
av såväl O2 som pH. En mekanism för produktionen av CH2O från Tris
föreslås.
Dessuton påverkas konsumtionshastigheten av H2O2 på ZrO2 av Tris.
En modifierad mekanism för katalytisk sönderdelning av H2O2 på ZrO2ytor föreslås, vilken omfattar oberoende ytadsorptionsplatser för H2O2
och Tris. Dessutom är det visat att avvikelsen i CH2O-koncentration,
uppmätt med den modifierade Hantzsch metoden, från den verkliga
CH2O-koncentrationen ökar med ökande [H2O2]0/[CH2O]0.
Detta arbete bekräftar även att sulfid inhiberar strålningsinducerad
upplösning av UO2. Effekten beror på sulfidkoncentration. Det har också
visats att denna inhibering är oberoende av H2/Pd.
Det har visat sig att reaktiviteten för H2O2 och dynamiken för bildning
av formaldehyd vid ytkatalys beror på vilken oxid som studeras. De
kinetiska parametrarna, såsom hastighetskonstanter, aktiveringsenergier
och frekvensfaktorer, har bestämts.
Både reaktioner på ytor och reaktioner i lösning har observerats i
W(s)/H2O2/Tris-systemet. Det har visats att Haber-Weissreaktioner
dominerar systemet. Även i system där O2 används för att oxidera W
bildas hydroxylradikaler. Detta beror på att H2O2 bildas som en
intermediär.
Den nya kunskap som denna avhandling förmedlar är relevant inom
kärnteknik såväl som inom tillämpningar av fotokatalys, biokemi,
korrosionsvetenskap, katalys och optik/elektronik.
vi
List of Papers
I.
Inhibition of Radiation Induced Dissolution of UO2 by Sulfide – A
Comparison with the Hydrogen Effect.
Miao Yang, Alexandre Barreiro Fidalgo, Sara Sundin, Mats
Jonsson, Journal of Nuclear Materials, 434(2013)38-42,
DOI: 10.1016/j.jnucmat.2012.10.050
II.
Catalytic Decomposition of Hydrogen Peroxide on Transition
Metal and Lanthanide Oxides.
Claudio M. Lousada, Miao Yang, Kristina Nilsson, Mats Jonsson.
Journal of Molecular Catalysis A: Chemical, 379(2013)178-184.
DOI: 10.1016/j.molcata.2013.08.017
III.
Evaluation of the O2 and pH Effects on Probes for Surface Bound
Hydroxyl Radicals.
Miao Yang, Mats Jonsson. The Journal of Physical Chemistry C,
118(2014)7911-7979, DOI: 10.1021/jp412571p
IV.
Surface Reactivity of Hydroxyl Radicals Formed upon Catalytic
Decomposition of H2O2 on ZrO2.
Miao Yang, Mats Jonsson. Journal of Molecular Catalysis A:
Chemical, 400(2015)49-55, DOI: 10.1016/j.molcata.2015.02.002
V.
The Reactivity and Mechanism of Reactions between H2O2 and
Tungsten Powder.
Miao Yang, Xian Zhang, Alex Grosjean, Inna Soroka, Mats
Jonsson. The Journal of Physical Chemistry C, under review.
VI.
Influence of H2O2 on the Quantitative Determination of
Formaldehyde using Acetoacetanilide as Detection Reagent in the
Hantzsch Method.
Miao Yang, Inna Soroka, Mats Jonsson. Manuscript, to be
submitted.
vii
VII.
Hydroxyl Radical Production in Aerobic Aqueous Solution
containing Metallic Tungsten.
Miao Yang, Inna Soroka,
Communications, under review.
VIII.
Mats
Jonsson.
Catalysis
Kinetics and Mechanisms of Reactions between H2O2 and Copper
and Copper Oxides.
Åsa Björkbacka, Miao Yang, Claudia Gasparrini, Christofer
Leygraf, Mats Jonsson. Dalton Transactions, Accepted
Manuscript, DOI: 10.1039/C5DT02024G.
http://pubs.rsc.org/en/Content/ArticleLanding/2015/DT/C5DT02024G
#!divAbstract
My contribution to the papers
Paper I: I have performed the major part of the experiments, participated
to a large extent in the data evaluation, and written the major part of the
manuscript.
Paper II: I have performed the major part of the experiments,
participated in the data evaluation and the preparation of the manuscript.
Paper III IV VI VII: I have performed the experiments, participated to
a large extent in the data evaluation, and written the manuscript.
Paper V: I have performed the major part of the experiments,
participated to a large extent in the data evaluation, and written the major
part of the manuscript.
Paper VIII: I have performed the experiments for the mechanistic study,
participated in the data evaluation and written the mechanistic part in the
manuscript.
viii
List of abbreviations
SNF – Spent nuclear fuel
G(x) – Radiation chemical yield for the species x
Ċ(x) – Amount of species x produced
Ḋ – Radiation dose rate
ρ – Solvent density
Gy – Gray, equal to 1 J kg−1
k – Reaction rate constant
k1 – First-order rate constant
k2 – Second-order rate constant
t – Time
SA – Surface area of solid
V – Volume of solution
Ea – Arrhenius activation energy
A – Arrhenius pre-exponential or frequency factor
R – Gas constant
T – Absolute temperature
UV-Vis – UV/Vis spectrophotometry;
HPLC – High performance liquid chromatography;
ESR – Electron spin resonance
Tris – Tris(hydroxymethyl) aminomethane
TAPS – N-Tris(hydroxymethyl)methyl-3-aminopropanesulfonic acid
DMPO – 5,5-Dimethyl-1-pyrroline-N-oxide
4-CP – 4-Chlorophenol
DDL – 3,5-Diactyl-1,4-dihydrolutidine
CHD – 1,3-Cyclohexanedione
MA – Methyl acetoacetate
AAA – Acetoacetanilide
AA – Ammonium acetate
ix
Y – Yield of formaldehyde
B.E.T. – Brunauer-Emmett-Teller
XRD – X-ray diffraction
SEM – Scanning electron microscopy
XPS – X-ray photoelectron spectroscopy
ICP-OES – Inductively coupled plasma optical emission spectroscopy
BE – Binding energy
K – Adsorption equilibrium constant
kf – Forward reaction rate constant of equilibrium reaction
kr – Reverse reaction rate constant of equilibrium reaction
Sa – Surface sites for H2O2
Sb – Surface sites for Tris
b2 – Intercept at zero coordinate
Abs – Absorbance
x
Table of Contents
1
2
Introduction ................................................................... 1
1.1
Background ............................................................................... 1
1.2
Water radiolysis ......................................................................... 3
1.3
Chemical kinetics of heterogeneous system ............................. 5
1.4
Oxidative dissolution of metals and metal oxides ..................... 8
1.5
Catalytic decomposition of H2O2 at solid-liquid interfaces ....... 10
1.6
Haber-Weiss reactions ............................................................ 12
1.7
Methodology for the detection of hydroxyl radicals ................. 13
1.8
The scope of this thesis........................................................... 14
Experimental details ................................................... 16
2.1
Instrumentation ........................................................................ 16
2.2
Materials and reagents ............................................................ 17
2.3
Basic methodology .................................................................. 18
2.4
Oxidative dissolution of metals and metal oxides ................... 19
2.4.1
Oxidative dissolution of W by H2O2 .................................... 19
2.4.2
Oxidative dissolution of W by O2 ........................................ 19
2.4.3
Inhibition of radiation induced dissolution of UO2 by sulfide20
2.5
Catalytic decomposition of H2O2 on metal and metal oxide
surfaces ............................................................................................... 20
2.5.1
Metals ................................................................................. 20
2.5.2
ZrO2 .................................................................................... 21
2.5.3
Other lanthanide and transition metal oxides ..................... 22
2.6
Haber-Weiss reactions ............................................................ 22
2.7
Methodology for the detection of hydroxyl radicals ................. 23
2.7.1
The O2 and pH effects on probes for hydroxyl radicals ...... 23
xi
2.7.2 The interference of H2O2 on the modified Hantzsch method
using AAA as derivatization reagents .............................................. 23
3
Results and discussion .............................................. 24
3.1
Methodology for the detection of hydroxyl radicals ................. 24
3.1.1
The effects of O2 and pH on probe chemistry III ................. 24
3.1.2 The effect of Tris on the catalytic decomposition of H2O2 on
ZrO2 surface IV ................................................................................ 32
3.1.3 The interference of H2O2 on the modified Hantzsch method
using AAA as derivatization reagents VI ......................................... 38
3.1.4
3.2
Summary ............................................................................. 39
Inhibition of radiation induced dissolution of UO2 by sulfide I . 40
3.3
Kinetics, mechanistic and simulation study of catalytic
decomposition of H2O2 on metal and metal oxide surfaces ................ 42
3.3.1
Cu, Cu2O and CuO VIII ....................................................... 42
3.3.2
Other lanthanide and transition metal oxides II .................. 43
3.3.3
Summary ............................................................................. 48
3.4
Kinetics and mechanistic study of reactions between W and
H2O2 or O2 ............................................................................................ 48
3.4.1
W and H2O2 V ..................................................................... 48
3.4.2
W and O2 VII ....................................................................... 57
3.4.3
Summary ............................................................................. 60
4
Conclusions ................................................................ 61
5
Future work ................................................................. 63
6
Acknowledgements .................................................... 64
7
References................................................................... 66
xii
1 Introduction
1 Introduction
1.1 Background
Nuclear power has been and will be an increasingly significant
contributor to energy production in several countries due to its
outstanding performance in Greenhouse gas emission, efficiency and
safety. One of the major drawbacks of nuclear power is the generation of
radioactive waste. Regardless of the future of nuclear power, nuclear
waste already exists. The major nuclear waste is the used nuclear fuel,
normally called spent nuclear fuel (SNF) which is highly radioactive. The
radioactivity of SNF will remain at a high level for around 100, 000
years.1 Hence, the SNF must be handled safely for the present and future
generations.
In light water cooled reactors, the nuclear fuel normally consists of
ceramic UO2. After use in a reactor, UO2 is still the main constituent of
the fuel (95 wt.%) and the remaining 5% comprises of highly radioactive
fission products and transuranium elements.2, 3 Globally, about 10, 000
tonnes of SNF are produced every year and so far, it has been
accumulated to approximately 300, 000 tonnes in pools or dry casks onsite.1 There are different strategies for disposal of SNF, e.g. on-site storage,
space disposal, subseabed disposal, ice-sheet disposal, reprocessing and
geological repository.4 To fulfill the goal of SNF disposal safety, Sweden
chose the KBS-3 model, developed by SKB (Swedish Nuclear Fuel and
Waste Management Company), for direct disposal of SNF in a deep
geological repository.5 As shown in Figure 1,5 there are multiple barriers
to protect the SNF, i.e. copper canister, bentonite clay and crystalline
bedrock. In the worst case scenario, e.g. canister failure, the highly
radioactive SNF will induce radiolysis of the intruding groundwater,
thereby initiating a potentially aggressive chemistry at the surface of
materials used in the repository. The chemistry at the interface of solid
materials and water upon exposure to ionizing radiation is highly
important in nuclear reactors as well as in reprocessing technology.6
1
1 Introduction
Figure 1. Scheme of the SNF deep geological repository.5
There are two fundamental issues for the safety assessment of these
systems (geological repository, nuclear reactor and reprocessing): the
stability of materials and the release rate of radioactive elements.7 Both
processes are to a large extent interfacial processes. However, in-situ
studies are problematic mainly due to the extreme operating temperature,
pressure as well as the very intense ionizing radiation emitted. The
relevant materials in nuclear technological applications are required to be
stable and durable after having withstood these extreme conditions. The
materials stability under high temperature and pressure is primarily
controlled by the chemistry of the system and can be predicted from
experimental data obtained under the much less extreme conditions
based on the tools of thermodynamics.8
The radionuclide release process can be divided into two stages:
instantaneous release from the SNF surface and long-term release due to
dissolution of the fuel matrix.1 The study of the release process is not only
important for geological repository, but also for reprocessing technology.7
The study of heterogeneous systems is not restricted to nuclear
technological applications. In heterogeneous catalysis, biochemistry,
photocatalysis, corrosion science, surface polymerization and many
industrial applications, e.g. water purification, interfacial processes are
2
1 Introduction
also of utmost importance. The ionizing radiation present in nuclear
technological applications makes the problem more intriguing.
The knowledge of the radiation chemistry of water is essential to be
able to control the chemistry of nuclear technological applications. The
homogeneous radiation chemistry of water is well known since many
decades. However, radiation induced processes at solid-liquid interfaces
remain to be fully understood, although they have been investigated to
some extent by both experimental and theoretical studies.6, 9-11
Consequently, better understanding of these processes is a prerequisite
for the development of materials used in nuclear technology.
1.2 Water radiolysis
The interaction between ionizing radiation and water leads to
ionization and electron excitation of water. Electrons emitted can lead to
further ionizations which are also called secondary ionizations. Upon
absorbing energy from fast charged particles or high energy photons,
water is decomposed into its radiolysis products. Within 10−7 s after
energy deposition, the primary radiolysis products are completely formed
in a spur. The products will have diffused away from the spur and the
interactions between the formed products in the spur become negligible.12
It should be noted that other oxidants (e.g. O2•−, HO2• and O2) are also
formed from further reactions between the primary products.13 In pure
water, the sequence of events that follows the deposition of energy is
summarized in Figure 2.12
The radiation chemical yield, also called G-value, is defined as the
amount of specified radiolysis product formed per 100 eV of energy
deposited from ionizing radiation.12 The G-value is expressed in S.I. units
as
𝐺(𝑥) =
𝑛𝑥
(1)
𝛿𝛿
where G(x) is the radiation chemical yield for the radiolysis species x and
nx is the number of moles of x formed per unit of energy (δE) in Joules (J)
deposited in the medium.
3
1 Introduction
Time/s
H 2O
Excitation
Ionization
-
H2O+ +
H2O*
e
10-16
HO + H3O+
H2 + O
H
10-14
+ OH
10-13
-
eaq
Formation of molecular products,
diffusion of molecular species
and radicals out of spur
H2
HO
H 2O 2 H 3O +
HO
-
-
H
eaq
10-7
Figure 2. Time scale of events in water radiolysis resulting in the primary products.
The radiolysis products of water and their chemical yields are well
determined under a multitude of conditions. Generally accepted G-values
for the primary products after 10−7 s are listed in Table 1.
Table 1. Chemical yields (G-values) of radiolysis products in irradiated neutral water.12
Radiation
G(H2)
µmol J−1
G(H2O2)
µmol J−1
G(eaq−)
µmol J−1
G(H•)
µmol J−1
G(HO•)
µmol J−1
G(H3O+)
µmol J−1
γ and fast e−
0.047
0.073
0.28
0.062
0.28
0.28
The G-value for a specific radiolysis product can be used to calculate its
formation rate via the following equation
Ċ(x) = Ḋ × G × ρ
(2)
where Ċ(x) is the production rate in mol dm–3 s–1, Ḋ is the dose rate in Gy
s–1 and ρ is the density of the solvent. The S.I. unit Gray (Gy) represents
the absorbed dose and 1 Gy corresponds to 1 J kg−1.
4
1 Introduction
The powerful oxidizing (HO•) and reducing (eaq−) radiolysis products
have approximately the same chemical yields. The interconversion
between oxidizing and reducing condition can be simply performed by
adding a reactant to the media. By this means, the media can be
alternated to primarily oxidizing or reducing condition. For instance, the
most readily and effective interconversion is that of eaq− to HO• through
the following reactions in N2O-saturated solution.14
−
eaq + N2O → N2 + O
•−
•−
−1
9
O + H2O → HO + HO
•
k3 = 9.1 × 10 L mol
−
6
−1
k4 = 1.8 × 10 L mol
s
s
−1
−1
(3)
(4)
where k is the rate constant for the respective reaction. Under these
conditions, G(HO•) is almost doubled to 6.0 x 10–7 mol J–1 and the system
is alternated to oxidizing media.15
Another type of reagent is secondary alcohol which gives reducing
condition upon reacting with HO• and H•. For instance, 2-propanol reacts
with both radicals forming a strong reducing radical as follows.16 The rate
constant k5 for this reaction is equal to 3.3 × 1012 L mol−1 s−1.16
H3C-CHOH-CH3 + HO (H ) → (CH3)2-C OH + H2O (H2)
•
•
•
(5)
One of the main long-lived products of radiolysis of water is H2O2
which can react with metal/metal oxide surfaces both via redox reactions
and catalytic decomposition.17-19 The relative impact of several radiolytic
oxidants on the dissolution of UO2 was studied and H2O2 was determined
to be the only radiolytical oxidant needed to be accounted for under
geological repository conditions.20 Therefore, the interactions between
H2O2 and metal/metal oxide surfaces have been treated as key reactions
in the radiation induced processes at solid-liquid interfaces and have
been investigated in a number of studies.1, 6, 21-26 By focusing on the key
reactions, the study of the chemistry occurring at the interfaces can be
understood.
1.3 Chemical kinetics of heterogeneous system
The heterogeneous system, involving a solute and a solid phase, has
more steps than the homogeneous system, e.g. diffusion and adsorption.
5
1 Introduction
The reactions in the heterogeneous system can be divided into several
steps as shown in Figure 3.
Figure 3. Schematic representation of reactions in the heterogeneous system. Rsol and Rads
are the reactants in solution and adsorbed on the surface, respectively; TSads is the
transition state for the surface reaction; Pads and Psol represent the products of reaction
adsorbed on the surface and in solution, respectively; kads, kSr and kdes denote the rate
constants for the adsorption, surface reaction and desorption, respectively.
For a non-catalytic process, the surface will undergo alterations, e.g.
corrosion, dissolution, formation of complexes etc. On the other hand, for
catalytic process, the catalyst surface will exhibit no alterations during the
whole process. As can be seen in Figure 3, the process includes several
steps: diffusion of the free reactant to the surface, adsorption onto the
surface followed by the surface reaction, desorption of products into
solution and products in solution (Psol) may react with free reactant (Rsol).
Any of these steps may be rate-limiting.27
In general, the rate of a chemical reaction is dependent on the
concentration of reactants and the rate constant for the reaction. In the
case of a solute and a solid reactant, the reaction rate expression can be
derived as Eq. (6) regarding the consumption rate of the solute and solid,
respectively.
−
d[Solute]
d𝑡
SA
= 𝑘 � � [Solute] or −
V
d𝑛𝑠𝑠𝑠𝑠𝑠
d𝑡
= 𝑘(𝑆𝑆)[Solute]
6
(6)
1 Introduction
where SA, V and k denote the surface area of solid, the solution volume
and the reaction rate constant, respectively.
In the excess of solid material, the kinetics of solute reactant
consumption becomes pseudo first order. The pseudo first order rate
constant can be obtained from the slope when plotting the logarithm of
the solute reactant concentration versus reaction time. By repeating this
procedure for a series of different SA/V, the second order rate constant
can be obtained by plotting the pseudo first order rate constants against
SA/V. The unit of the second order rate constant expressed in this way is
m s−1. Powder suspensions are preferential for studying this
heterogeneous system due to its high surface area to volume ratio and
large flexibility in SA/V.
The activation energy of the reaction can be obtained by determining
the rate constants as a function of temperature according to the
Arrhenius equation
𝑘 = 𝐴𝑒 −𝐸𝐸 ⁄𝑅𝑅
(7)
where Ea represents the activation energy for the reaction; A is the preexponential factor.
If the activation energy is low, the rate constant is mainly determined
by the pre-exponential factor. Under this condition, the reaction rate is
dominated by the rate constant of the diffusion of the reacting species in
the system. Therefore, the pre-exponential factor is sometimes referred to
as the diffusion controlled rate constant and can be considered as the
highest possible rate of reaction. The diffusion rate for a heterogeneous
system is significantly smaller than its homogeneous counterpart for the
same reaction.28
In some systems, both catalytic and non-catalytic reactions may take
place. The products from the non-catalytic process, e.g. corrosion,
dissolution, formation of complexes, may react with the reactants in the
solution, thereby affecting the kinetics of the whole process.
7
1 Introduction
1.4 Oxidative
oxides
dissolution of metals
and metal
The matrix of UO2-based spent nuclear fuel is UO2 (95%) while the
remaining 5% are highly radiotoxic fission products and higher actinides.
The release of these is dependent on the dissolution of the UO2 matrix
after the initial dissolution of readily soluble radionuclides.29 Besides UO2,
the dissolution behavior of other metals, metal oxides utilized in nuclear
technological application is also of importance, such as copper which is
used for the canister in the deep geological repository.5 Hence, it is
worthwhile to understand the mechanism and kinetics of the oxidative
dissolution of these metals/ metal oxides.
H2O2, one of the key long-lived molecular radiolysis products of water,
can react with metal/metal oxide surfaces both via redox reaction and
catalytic decomposition.17-19 If the oxidized metals/metal oxides are
soluble, the redox reaction is oxidative dissolution. For example, metallic
tungsten has been proven to be readily dissolved in aqueous H2O2
solution30, 31 via a step-wise oxidation process from metallic state to W(VI)
via W(III), W(IV) and W(V).32-34
Except for metals, H2O2 also leads to the dissolution of some metal
oxides. For instance, insoluble UO2 can be oxidized to soluble UO22+ via a
two-step single electron oxidation. It should be noted that hydroxyl
radicals can be produced during the single electron oxidation from U(IV)
to U(V) and this step is the rate limiting step of the oxidative dissolution
of UO2.35 The dissolution process can be greatly enhanced in the presence
of carbonate (HCO3-) via formation of soluble complexes.36 The carbonate
concentration in Swedish groundwater is approximately 2 mM.37 It was
proven that the oxidative dissolution rate of UO2 by H2O2 increases
linearly with [HCO3-] from 0 to 1 mM, but becomes independent of
[HCO3-] above 1 mM since the reaction becomes oxidation-limited.23
H2 is another major long-lived molecular radiolysis product of water12
and H2 is also expected to be produced in a large amount as a result of
anaerobic corrosion of the iron insert38 in case of canister damage in the
deep geological repository. H2 was proven to inhibit the radiation induced
8
1 Introduction
dissolution of SNF and the inhibition process is significantly enhanced by
the presence of noble metal nanoparticles of fission products in the SNF
which are normally referred to as ε-particles (containing Mo, Ru, Tc, Pd
and Rh).39-41 The scheme of the oxidative dissolution of UO2 and its
inhibition by H2 is shown in Figure 4.42
Figure 4. Oxidative dissolution of UO2 and its inhibition by H2.42
In the groundwater environment of the SNF geological repository,
sulfide is another common species, the concentration of which is between
1 and 2.1 μM.43 H2S is slightly soluble in water with pKa1 = 7.05, pKa2 =
14.0 at 293 K for HS− and S2−, respectively.44 Therefore, the main sulfide
species in water are H2S and HS- and denote as sulfide throughout the
whole thesis. Sulfide may inhibit the radiation induced dissolution of UO2
due to its reducing ability for U(VI).45, 46 The reaction stoichiometry
determined in the literature46 is shown as follows:
UO2 + HS → UO2 + S + H
2+
-
0
+
(8)
However, HS− can also scavenge radiolytic oxidants.29 Palladium is one
of the fission products39-41 and sometimes used as a model for ε-
9
1 Introduction
particles.25, 47 Palladium can catalyze the inhibition of oxidative
dissolution of UO2 by H2 (see Figure 4).47 It was revealed by previous
studies that palladium can be poisoned by sulfur compounds in aqueous
solutions.48-50 Hence, it is necessary to investigate not only the inhibition
of sulfide but also to study if the poisoning effect exists in the H2/Pd
system.
1.5 Catalytic decomposition of H 2 O 2 at solid-liquid
interfaces
Except for redox reactions, the catalytic decomposition is the other
type of interfacial reaction between H2O2 and metal/metal oxide surface
which has been studied for decades both experimentally and
theoretically.9, 10, 18, 51 Both reactions occur if the metal in its oxide, such as
U in UO2,21-23, 51 is not in its highest oxidation state.51 In contrast, it was
shown that when the metal is in its highest oxidation state only catalytic
decomposition occurs at the interface9, 51, 52 with the exception of CeO2
which can be reduced to Ce(III) by H2O2.53 54 One of the most studied
metal oxides regarding the catalytic decomposition of H2O2 is ZrO2.10, 11, 51,
52, 55 The 9, 109, 109, 109, 109, 109, 109, 109, 109, 109, 109, 109, 1010, 119, 1010, 1110, 11mechanism is
summarized as follows9, 10:
H2O2(ads) → 2HO (ads)
•
(9)
H2O2(ads) + HO (ads) → H2O(ads) + HO2 (ads)
(10)
2HO2 (ads) → H2O2 + O2
(11)
•
•
•
where ads represents the adsorbed state of the species.
The overall reaction is
H2O2 → H2O + 1/2O2
(12)
It was demonstrated both experimentally and theoretically that the
hydroxyl radical is the primary intermediate product during the catalytic
decomposition of H2O2 at oxide interfaces.10, 51, 52 The half-life of the
produced intermediate surface bound radicals could be sustained
significantly to several hours or even up to few days.56, 57 The stability of
10
1 Introduction
the materials in the presence of H2O2 may be affected by the strong
reactive HO•(ads) at the interface.
Hydroxyl radicals cannot be detected or quantified directly due to its
short half-life and high reactivity. Hence, several probes for HO•, such as
TAPS (N-Tris(hydroxymethyl)methyl-3-aminopropanesulfonic acid),
methanol, Tris (Tris(hydroxymethyl) aminomethane), DMPO (5,5dimethyl-1-pyrroline-N-oxide), benzoic acid and 4-CP (4-chlorophenol)51,
52, 58-61 have been applied to detect the adsorbed hydroxyl radicals via
different reactions and detection techniques (see Table 2).
Table 2. The reaction types and detection techniques of common probes for the hydroxyl
radical.24, 34-38
Probe
Tris/Methanol/TAPS
DMPO
Benzoic acid/4-CP
Reaction type
Detection technique
Oxidation
UV-Vis/HPLC
Addition
ESR
Hydroxylation
HPLC
UV-Vis: UV/Vis spectrophotometry; HPLC: high performance liquid
chromatography; ESR: electron spin resonance
Tris was selected as hydroxyl radical scavenger in a number of
applications, such as interfacial radiation chemistry,52 catalytic reaction,51
photocatalysis62 and biological processes.63 As previously described,52
•
HO abstracts hydrogen from Tris, yielding formaldehyde as one of the
•
products. The rate constant of the reaction between Tris and HO in
homogeneous aqueous solution is 1.1 x 109 M–1 s–1.64 Formaldehyde can
be detected and quantified by a modified version of the Hantzsch
method.65, 66 Similar to Tris, methanol is another scavenger that also
yields formaldehyde upon hydrogen abstraction by HO• with a rate
constant of 9.7 x 108 M–1 s–1 in homogeneous solution.14 However, the rate
constants for Tris and methanol reacting with surface bound hydroxyl
radicals in the heterogeneous system are yet not determined.
Although the kinetic data and mechanism for H2O2 decomposition on
ZrO2 were elucidated, the reaction needs to be studied for a diversity of
metals/metal oxides.
11
1 Introduction
1.6 Haber-Weiss reactions
Besides the redox reactions and catalytic decomposition introduced in
above sections, the reactions between H2O2 and one-electron redox
couple of metal ion(s) become important when the amount of ion resulted
from the oxidative dissolution is sufficient. The reaction is called HaberWeiss reactions and the mechanism is shown as follows67, 68
Mox + H2O2 → Mred + HO2 + H
•
+
Mred + H2O2 → Mox + HO + HO
•
(13)
−
(14)
Mox and Mred represents the oxidized and reduced forms of the oneelectron redox couple in metal complexes (e.g., Fe(III)/Fe(II),
W(VI)/W(V)).
Except in the aqueous phase, the reactions R(13) and R(14) can also
occur in heterogeneous system at the solid-liquid interfaces, 35, 69, 70 e.g.
hydroxyl radical is produced in the oxidative dissolution of UO2 by H2O2
as shown in Figure 4. The rate limiting step is R(13),67 the rate of which
can be influenced by binding to ligands or the formation of
polyperoxometalate.68, 71, 72 R(14) is also called a Fenton-like reaction. The
Haber-Weiss reactions (e.g. H2O2/Fe(III)/Fe(II)) show the behavior of
producing hydroxyl radical continuously without changing total
concentration of iron ions.73, 74 The generation rate of hydroxyl radicals is
governed by the steady-state concentration of Fe(II) which is mainly
controlled by the total iron concentration.73
In the heterogeneous system in the presence of H2O2 and highly soluble
metal/metal oxide, all the three mentioned reactions: oxidative
dissolution, catalytic decomposition and Haber-Weiss reactions, may
occur. A thorough study is needed on a single system including all of
these reactions.
12
1 Introduction
1.7 Methodology for the detection of hydroxyl
radicals
As mentioned in the above section, some probes, e.g. Tris, TAPS and
methanol, can be used to determine the amount of hydroxyl radicals by
quantifying the amount of formed product (CH2O).51, 52 A number of
methods for the determination of formaldehyde have been proposed,
many of which are based on the Hantzsch reaction including the
cyclization of amine, aldehyde and β-diketon to form a dihydropyridine
derivative (Figure 5a).65 A modified version of the Hantzsch method was
introduced by Nash66 using acetylacetone or 2,4-pentanedione to react
with formaldehyde in the presence of ammonia to form a yellow
derivative
of
3,5-diactyl-1,4-dihydrolutidine
(DDL).
Both
spectrophotometry66 and fluorometry75 can be utilized to measure DDL
with high sensitivity. Though, this method is time-consuming and
requires elevated temperature. Several new β-diketon based
derivatization reagents of formaldehyde have been presented, such as 5,5dimethyl-1,3-cyclohexanedione (dimedone),76 4-amino-3-pentene-2-one
(Fluoral-P),77 1,3-cyclohexanedione (CHD),76 methyl acetoacetate (MA)78
and acetoacetanilide (AAA) (Figure 5b)65.
R2
CH +
NH2
O
R3
O
C
H H
+
+2
H +2
N
H H
O
R C
CH2
C O
R1
H O
N C
H 3C
O
R C
R1
N
O
C R + 2 H2O + 1/2 O2
R1
(a)
R3
CH2
C
R2
NH
H 3C
O
O
C
O
C
N
H
NH
+ 2 H2O + 1/2 O2 (b)
CH3
Figure 5. Scheme for the Hantzsch reaction (a) Using β-diketon as derivatization reagent;
(b) Using acetoacetanilide (AAA) as derivatization reagent.
As an excellent derivatization reagent, AAA has been adopted in a
number of studies.10, 11, 51, 52, 79-81 In many of these studies, CH2O is formed
in the presence of relatively high initial concentration of H2O2 (mM level).
The similar reagent CHD is known to suffer positive interference from
13
1 Introduction
H2O2 at H2O2 concentration on µM level.82-85 However, the interference
from H2O2 on CH2O determination using AAA is still unknown.
The yield (Y) of CH2O in a radiolysis experiment is defined (Eq. (15)) as
the ratio between the concentration of scavenged hydroxyl radicals (equal
to the measured formaldehyde concentration) and the accumulated
concentration of hydroxyl radicals formed during the radiolysis of water
which can be calculated based on Eq. (2).
•
Y = [CH2O]/C(HO )
(15)
In N2O-saturated Tris solution, Y was measured to be 35% by using γ
radiation of water as the source of hydroxyl radicals.52 However, the
influencing factors (e.g. O2 and pH) for the yield of formaldehyde as well
the mechanism of hydroxyl radical scavenging Tris were still not fully
understood.
The kinetics of H2O2 decomposition will be changed in the presence of
Tris according to the abovementioned competition.52 However, the
mechanism of catalytic decomposition of H2O2 at interfaces in the
presence of hydroxyl radical scavenger remains to be elucidated.
1.8 The scope of this thesis
The goal of this thesis is for further understanding of the radiation
induced processes at solid-liquid interfaces. The kinetics and mechanism
of the processes are mainly investigated by studying the reactions
between the key water radiolysis product—H2O2 and metals/metal oxides.
The reactions include oxidative dissolution of solid material, catalytic
decomposition of H2O2 and Haber-Weiss reactions. The scheme is shown
in Figure 6.
Specifically, the present thesis is to answer the following questions
which are crucial for better understanding of the interfacial processes.
1.
Is the kinetics and mechanism of catalytic decomposition of H2O2
at interfaces the same for different metal oxides?
2. Does the dissolution of metals/metal oxides affect the mechanism
of the studied heterogeneous system?
14
1 Introduction
3. Do the presence of hydroxyl radical probe, O2 and pH affect the
mechanism of catalytic decomposition of H2O2 at interfaces?
4. What is the limitation of the methodology for quantifying the
production of hydroxyl radicals in the studied heterogeneous
system?
The knowledge is not only relevant to nuclear technological
applications, but also for applications related in photocatalysis,
biochemistry, corrosion science, catalysis and optics/electronics.
Ultimately it is my wish to provide a better vision for material scientists
regarding the reactivity of radiolysis products upon the surface of
materials on the basis of the present thesis.
Figure 6. Scheme of radiation induced processes at solid-liquid interfaces.
15
2 Experimental details
2 Experimental details
2.1 Instrumentation
The specific surface areas of the powders were determined by using the
Brunauer-Emmett-Teller (B.E.T.) method of isothermal adsorption and
desorption of a gaseous mixture consisting of 30% N2, 70% He in a
Micrometrics Flowsorb II 2300 instrument.
The samples were weighed using a Mettler Toledo AT261 Delta Range
microbalance.
Structural studies were done by powder XRD (X-ray diffraction)
method with PANalytical X'Pert PRO diffraction system using BraggBrentano geometry in the 2θ angle from 10° to 100° and CuKα irradiation
(λ = 1.54 Å).
Surface morphology of samples was obtained using SEM (Scanning
electron microscopy) analysis with an FEI-XL 30 Series instrument. The
accelerating voltage was 20 kV.
Elemental analysis was done using XPS (X-ray photoelectron
spectroscopy). The spectra of samples were recorded with a Kratos Axis
Ultra electron spectrometer with a delay line detector. The depth of
analysis of metal oxides/hydroxides was about 6 nm. The element
detection limit was typically 0.1 at. %.
Trace elemental analysis on all solutions were conducted by using
Thermo Scientific iCAP6000 series ICP-OES (Inductively coupled plasma
optical emission spectroscopy). The analysis for tungsten sample was
performed at wavelengths of 207.9, 209.8 and 224.8 nm using ICP
element standard IV from Sigma Aldrich.
Gamma-irradiation was conducted in a MDS Nordion 1000 Elite Cs137 γ-source with a dose rate of 0.15 Gy s–1, quantified by Fricke
dosimetry.12
The powder suspensions were stirred using a Heidolph MR3001K
magnetic stirrer. The temperature was controlled throughout the
16
2 Experimental details
experiments using a Lauda E100 thermostat, calibrated against a Therma
1 Thermometer coupled to a submersible K-type (Ni-Cr-Ni) temperature
probe, with a precision of ± 0.1 K.
UV/Vis spectra were collected by a Thermo Scientific Genesys 20
UV/Vis spectrophotometer.
The pH of solution was measured by Metrohm 713 pH Meter with an
accuracy of ± 0.1, calibrated by standard pH reagents.
The experiments were performed either in open air or in different
gaseous atmosphere, N2O (AGA Gas AB), 20% O2 in N2O (N2O 80 ± 1.0%,
Strandmøllen A/S), N2 (≥ 99.999%, Strandmøllen A/S) and H2 (AGA Gas
AB).
2.2 Materials and reagents
All the solutions used in the thesis were prepared using Millipore MilliQ water (18.2 MΩ cm).
The major information about the powders used is listed in Table 3. All
the powders were used without further purification. The B.E.T. specific
surface area of the powders is the average value of three measurements
obtained by averaging the measured sorption and desorption isotherm
values.
Table 3. The B.E.T. specific surface areas for the powders.
Powder
Information
B.E.T surface area/m2 g−1
ZrO2
CeO2
HfO2
Gd2O3
Fe2O3
CuO
Cu2O
Cu
W
WO3
CAS[1314-23-4], Aldrich 99%
CAS[1306-38-3], Alfa Aesar 99.99%
CAS[12055-23-1], Alfa Aesar 99.95%
CAS[12064-62-9], Aldrich 99.9%
CAS[1309-37-1], Aldrich 99%
CAS[1317-38-0], Aldrich 99.99%
CAS[1317-39-1], Aldrich, 99,9 %
CAS[7440-50-8], Alfa Aesar, 99,9 %
CAS[7440-33-7], Aldrich 99.9%,
CAS[1314-35-8], Aldrich 99.9%
5.0 ± 0.2
14.3 ± 1.0
10.0 ± 0.1
1.7 ± 0.1
9.0 ± 1.0
15.3 ± 0.1
0.46 ± 0.01
0.089 ± 0.002
0.11 ± 0.01
3.51 ± 0.01
17
2 Experimental details
A UO2 pellet supplied by Westinghouse Atom AB was washed by 10
mM NaHCO3 prior to use in the experiments. Pd powder (Aldrich,
average particle size 0.625 μm) was used as the catalyst for H2 in
inhibiting dissolution of UO2.
H2O2 solutions were prepared from a 30% standard solution (Merck).
Reference formaldehyde solution was prepared from a stock solution
(CAS[50-00-0], Aldrich 37 wt% in H2O).
2.3 Basic methodology
The concentration of U(VI) was determined by the Arsenazo III
method86, 87. Arsenazo III (1,8-dihydroxynaphtalene-3,6-disulphonic
acid-2,7-bis[(azo-2)-phenylarsonic acid]) and U(VI) can form a colored
complex which can be quantified by a UV-Vis spectrophotometer at
wavelength 653 nm. The detection limit is 5 x 10−7 M.88
The concentration of H2O2 was determined using the Ghormley
triiodide method89, 90 by mixing samples with the detection reagents
ammoniumdimolybdate (Alfa Aesar, 4% w/v) and potassium iodide
(VWR BDH Prolabo, 99.0%). The absorbance of the product I3- was
measured at the wavelength of 350 nm using a UV-Vis spectrophotometer.
Two linear calibration curves were obtained by plotting the H2O2
concentration against the absorbance of I3– in the concentration range
0.01 – 1.0 mM (sample volume 0.2 mL) and 1.0 – 5.0 mM (sample
volume 0.02 mL) respectively. The uncertainty associated with the
determination of the concentration of H2O2 from the actual concentration
was < 2%.
Produced hydroxyl radicals in the experiments were scavenged by
Tris(hydroxymethyl) aminomethane (CAS[77-86-1], BDH Chemicals,
99%) or methanol (CAS[67-56-1], Aldrich, 99.9%) generating
formaldehyde. Formaldehyde was quantified by the modified version of
the Hantzsch method65, 66 using derivatization reagents acetoacetanilide
(AAA) (CAS[102-01-2], Aldrich, ≥98%) and ammonium acetate (AA)
(CAS[631-61-8], Aldrich, ≥98%). Specifically, 1 mL AAA (0.2 M, dissolved
in ethanol (99.5%)), 2.5 mL AA (4 M) and 1.5 mL sample were mixed and
18
2 Experimental details
incubated at 40°C for 15 min in a testing tube so as to form the DDL
derivative (Figure 5b) which can be measured at the wavelength of 368
nm by a UV-Vis spectrophotometer. A linear calibration curve was
obtained by plotting the absorbance of derivative against formaldehyde
concentration (in bulk),52 ranging from 0.15 µM to 1 mM. As reported
earlier, the uncertainty of this method is < 2%.52
All samples in suspensions were filtered through a 0.2 μm cellulose
acetate filter and pH of solutions was adjusted to 7.5 by HCl/NaOH in the
presence of Tris unless otherwise specified.
2.4 Oxidative
oxides
2.4.1
dissolution of metals
and metal
Oxidative dissolution of W by H2O2
To study the reaction between W and H2O2, W powder (SA/V = 550 –
8800 m–1) was submerged into H2O2 (5 mM) and/or Tris (100 mM)
solution at room temperature. In parallel, the dissolution behavior of
WO3 was also examined as a comparison to metallic W powder. The
experiments were conducted in the dark at ambient temperature, N2purged solution and continuous magnetic stirring at a rate of 600 rpm.
The pH of suspension was adjusted to 7.5 by HCl/NaOH after mixing Tris
and W powder. In the case without Tris, no pH adjustment was done for
the suspensions. After 30 min pre-mixing, H2O2 was added to the
suspension to initialize the reaction. The amount of dissolved tungsten in
the samples filtered at selected time points was measured via ICP-OES.
2.4.2
Oxidative dissolution of W by O2
Similar to the experiments for the oxidative dissolution of W by H2O2,
the air-saturated suspension of W powder and Tris or methanol (“free”
pH) were also undertaken the ICP-OES measurements for the
concentration of dissolved tungsten.
19
2 Experimental details
2.4.3
Inhibition of radiation induced dissolution of UO2 by sulfide
The inhibition of radiation induced dissolution of UO2 by sulfide was
studied in experiments performed at room temperature by using a UO2
pellet exposed to γ irradiation in aqueous solution containing sulfide and
NaHCO3. The U(VI) concentration of samples taken at selected time
points was measured.
Besides the pellet experiments, the inhibition of sulfide was studied
further by mixing U(VI) solution and sulfide directly. The mixture of
palladium powder (11 mg/100 mL) and 4 mM NaHCO3 was purged with
N2 for 20 minutes before adding to the autoclave. After purging with N2
for another 10 minutes and then pressurizing with pure N2 or H2 (20 bar),
the autoclave was left overnight with continuous stirring (60 rpm).
UO2(NO3)2 solution and H2S saturated solution were added to initialize
the reaction. The obtained initial concentrations of U(VI) and sulfide
(represents total H2S/HS− throughout the whole thesis) were 150 μM and
1 or 2 mM, respectively. The inhibition of sulfide was compared to H2, the
present experiments were divided into several groups: 1) N2 and sulfide, 2)
N2, sulfide and Pd, 3) H2 and Pd, 4) H2 and sulfide, 5) H2, Pd and sulfide.
The U(VI) concentrations of filtered samples taken at selected time points
were determined.
2.5 Catalytic decomposition of H 2 O 2 on metal and
metal oxide surfaces
2.5.1
Metals
Copper and tungsten powder were used to study the catalytic
decomposition of H2O2 at solid-liquid interfaces. In the thesis, the
mechanistic study of such reaction on metal surface is the primary task.
Both the H2O2 and CH2O concentrations were measured throughout the
whole reaction (room temperature) in the suspension containing Cu
(SA/V = 4800 m−1)/H2O2 (5 mM)/methanol (100 mM) or W (SA/V = 550
- 8800 m−1)/H2O2 (5 mM)/Tris (100 mM).
20
2 Experimental details
2.5.2
2.5.2.1
ZrO2
Effect of Tris
The experiments were performed in N2-saturated H2O2/ZrO2/Tris
suspensions in sealed vials with magnetic stirring at a rate of 750 rpm at
room temperature. SA(ZrO2)/V was fixed at 4.5 x 105 m−1 which is in line
with the previous study.10 The initial concentration of Tris varied from 20
mM to 400 mM while that of H2O2 varied from 0.5 mM to 5 mM. The
sample was only taken and filtered when the reaction was complete (after
16 h). The concentration of CH2O in the filtered sample was measured.
The model for H2O2 decomposition on the surface of ZrO2 in the
presence of Tris was built up using the kinetic simulation software Gepasi
3 which has been used in kinetic simulation studies.91, 92 All the relevant
kinetic parameters in the model were adapted to the reported
experimental works.52, 93
2.5.2.2
Effect of O2
For investigating the O2 effect on the catalytic decomposition of H2O2
on ZrO2 surface, the experiments were performed with N2-purged (O2 free)
or air-saturated suspensions for methanol (20 mM, pH 7.5), Tris (20 mM,
pH 7.5) and a mixture of the two scavengers (20 mM methanol + 5 mM
Tris, pH 7.5). The initial concentration of H2O2 and SA(ZrO2)/V were fixed
at 5 mM and 4.5 x 105 m−1, respectively. Both the concentrations of H2O2
and CH2O were monitored during the reaction.
2.5.2.3
Effect of pH
For the pH effect, the experiments were only performed in the presence
of Tris under the conditions: pH 7.0 – 9.0 (within buffering range) and
air-saturated. The other procedures were the same as the above section
on the effect of O2.
21
2 Experimental details
2.5.3
2.5.3.1
Other lanthanide and transition metal oxides
Kinetic study
The reaction suspension contains 0.5 mM H2O2 and different oxides
with different masses at different temperatures (298 – 353 K): Fe2O3 (0.2
– 1.5 g), HfO2 (0.1 - 0.75 g), Gd2O3 (0.25 - 1.0 g), CuO (0.0025 - 0.1 g). A
volume of 0.2 ml of filtered sample at selected time point was used for the
measurement of H2O2 concentration.
2.5.3.2
Mechanistic study
The reactions between the oxides and H2O2 were performed at room
temperature in the presence of 50 mL N2-purged Tris solution (20 mM).
The suspension contains H2O2 (5 mM), Tris (20mM) and different
amount of oxides: Fe2O3 (1.5 g) or CeO2 (1.6 g) or HfO2 (2.25 g) or Gd2O3
(3.0 g) or CuO (0.06 g). Both the concentrations of H2O2 and CH2O were
measured in the filtered samples at selected time points.
2.6 Haber-Weiss reactions
For the study of the homogeneous Haber-Weiss reactions between
H2O2 and dissolved tungsten, the experiment procedure was the same as
above (Section 2.5.1) in the heterogeneous system before the filtration. At
predetermined time point, a volume (10 or 20 mL) of sample was
extracted and then stored in a dark vial. The concentrations of H2O2 and
CH2O in the filtered solution as well as in the original suspension were
monitored and compared.
In order to further investigate the homogeneous reactions, extra H2O2
was added to the abovementioned filtered solution when H2O2 was
consumed completely. The concentrations of H2O2 and CH2O were
monitored. Besides, the amount of dissolved tungsten prior to and after
the addition of H2O2 was examined.
22
2 Experimental details
2.7 Methodology for the detection of hydroxyl
radicals
2.7.1
The O2 and pH effects on probes for hydroxyl radicals
The O2 and pH effects on the scavenging yield (Eq. (2)) of the probes
(Tris and/or methanol) for hydroxyl radicals were investigated by varying
dissolved O2 concentration and pH value of the solution. The experiments
were performed in N2O- or (80% N2O + 20% O2)- saturated aqueous Tris
(20 mM) or methanol (20 mM) solution which were subjected to γ
irradiation. The pH effect was only examined in the Tris case and the pH
was varied from 7.0 to 9.0. The G-value of hydroxyl radical is doubled to
6.0 x 10–7 mol · J–1 since all the solvated electron was converted to
hydroxyl radical by reacting with N2O according to R(3)and R(4).14, 15 The
accumulated concentration of CH2O in the sample taken at selected time
point was measured, which can be used to calculate the corresponding
yield based on Eq. (2).
2.7.2
The interference of H2O2 on the modified Hantzsch method
using AAA as derivatization reagents
The interference of H2O2 on the modified version of the Hanztsch
method52 was investigated by adding concentrated H2O2 solution to a
reference CH2O solution with known concentration. The CH2O
concentration in the mixed solution (H2O2 + CH2O) was measured by the
procedure described in section 2.3. Additional experiments were
conducted by adding concentrated H2O2 to the mixed solution (H2O2 +
CH2O) after the incubation time (15 min). The absorbance of the formed
derivative after the incubation was monitored after the addition of H2O2.
23
3 Results and discussion
3 Results and discussion
A number of metals and metal oxides (W, Cu, UO2, ZrO2, CeO2, HfO2,
Gd2O3, CuO, Cu2O and Fe2O3) were investigated regarding the interaction
with H2O2 in the aqueous system which includes oxidative dissolution,
catalytic decomposition and Haber-Weiss reactions. Depending on the
oxidation state as well as the solubility of the materials, one or more of
these processes can occur in the heterogeneous system containing H2O2
and metal/metal oxide powder. The kinetics, mechanism and the factors
influencing the processes were investigated.
Furthermore, since the hydroxyl radical is the intermediate product of
all of these three processes, the studies on the probes for hydroxyl
radicals are crucial. Here, the focus was on factors influencing the
scavenging yield, the scavenging mechanism, the influence of the probes
on the mechanism of the studied system and the influence of H2O2 on the
accuracy of the Hanztsch method for the quantification of hydroxyl
radical.
3.1 Methodology for the detection of hydroxyl
radicals
3.1.1
3.1.1.1
The effects of O2 and pH on probe chemistry III
The effect of O2
The experiments for studying the O2 effect on the homogeneous
aqueous system were performed under different conditions. The
dissolved O2 concentration (DO) in this work is kept constant at 0.26 mM
(Henry’s law constant is 769 atm dm3 mol–1 at 298 K).94 Hydroxyl radicals
are produced by using γ-irradiation on the solution. The concentration of
produced formaldehyde as a function of reaction time in the presence and
absence of O2 for both methanol and Tris is presented in Figure 7.
24
3 Results and discussion
0.6
[CH2O] / mM
0.5
0.4
0.3
0.2
0.1
0
0
1000
2000
3000
4000
5000
6000
Time / s
Figure 7. Concentration of CH2O as a function of reaction time for the γ-radiolysis of water
in different reacting solutions: methanol (20 mM, N2O/O2, unadjusted pH) (
mM, N2O, unadjusted pH) (
), Tris (20 mM, N2O/ O2, pH 7.5) (
), methanol (20
), Tris (20 mM, N2O, pH
7.5) ( ). Hydroxyl radicals formed in the radiolysis of water calculated from Eq.(2) (
American Chemical Society
). ©
As can be seen in Figure 7, the concentrations of CH2O are higher in
the presence of O2 both in the methanol and Tris cases. Specifically, the
presence of O2 leads to a five-fold (14% to 68%) and two-fold (16% to 29%)
increase in the formaldehyde yield in methanol and Tris, respectively.
After evaluating the O2 effect on the homogeneous system, the effect on
the heterogeneous system was also investigated under different
conditions. The production of CH2O as a function of reaction time is
shown in Figure 8.
25
3 Results and discussion
0.14
0.12
[CH2O] / mM
0.1
0.08
0.06
0.04
0.02
0
0
5000
10000
15000
20000
25000
30000
35000
40000
Time / s
Figure 8. Concentration of CH2O as a function of reaction time for the reaction of H2O2
(initial concentration 5 mM) with ZrO2 (4.5 g) under different conditions: methanol (20 mM,
air, pH 7.5) (
), methanol (20 mM, N2, pH 7.5) (
mM, N2, pH 7.5) (
), Tris (20 mM, air, pH 7.5) (
), mixture (20 mM methanol + 5 mM Tris, air, pH 7.5) (
mM methanol + 5 mM Tris, N2, pH 7.5) (
American Chemical Society
), Tris (20
), mixture (20
) and Tris (5 mM, air, pH 7.5) (
). © 2014
As can be seen in Figure 8, the production of CH2O in the Tris case is
enhanced in the presence of O2. However, this enhancement is
insignificant in the methanol case and not detectable in the mixture. As
compared to the O2-dependence experiments in the homogeneous
solution (see Figure 7), the O2 effect observed in the heterogeneous
system is much weaker, especially for methanol. The reason for this
inconsistence in the methanol case should be that H2O2 in the present
system competes with O2 to react with •CH2OH forming CH2O.95, 96 The
hydroxymethyl radical is an O2-sensitive precursor for CH2O and plays an
important role in both the Tris and methanol case. However, the O2-free
system is not oxidant-free. Therefore, addition of O2 only has marginal
effect on the production of CH2O. In general, the production of CH2O
increases in the order: methanol < Tris/methanol mixture < Tris. The
production of CH2O from the mixture is approximately the sum of the
26
3 Results and discussion
individual contributions from its constituents. Hence, stabilizing the pH
by addition of Tris, does not appear to influence the yield of formaldehyde
from methanol. The higher production in the presence of Tris than
methanol may be attributed to the difference in affinity towards ZrO2
surface. The difference between their affinities towards TiO2 has already
been illustrated in a photocatalysis system.62, 97
Firstly introduced in Paper II, a plot was employed to distinguish the
effect of different oxides on the production of CH2O as a function of H2O2
conversion. A slight revision was made by using normalized evolution of
CH2O ([CH2O]t/[CH2O]f) instead. [CH2O]t and [CH2O]f represent the
concentration of CH2O at a selected time point and the final
concentration, respectively. [CH2O]t/[CH2O]f is plotted against H2O2
conversion for different scavengers and the results are shown in Figure
9.
% of [CH2O]t/[CH2O]f
100
80
60
40
20
0
0
20
40
60
80
100
% of H2O2 converted
Figure 9. [CH2O]t/[CH2O]f as a function of conversion of H2O2 under different conditions (all
with ZrO2 (4.5g): methanol (20 mM, air, pH 7.5) (
Tris (20 mM, air, pH 7.5) (
), methanol (20 mM, N2, pH 7.5) (
), Tris (20 mM, N2, pH 7.5) (
mM Tris, air, pH 7.5) ( ) and mixture (20 mM methanol + 5 mM Tris, N2, pH 7.5) (
2014 American Chemical Society
27
),
), mixture (20 mM methanol + 5
). ©
3 Results and discussion
It can be seen from Figure 9 that there is no clear oxygen-dependent
variation for any of the scavengers. This similarity indicates that O2 has
minor influence on the evolution of formaldehyde as a function of the
consumption of H2O2
3.1.1.2
The effect of pH
A thorough study regarding the pH effect in homogeneous aqueous
solution was performed only for Tris since methanol cannot buffer the
solution pH. The experiment was performed under the conditions: 20
mM Tris, pH = 7.0 – 9.0 (pH buffering range of Tris), 20% O2 in N2O and
room temperature. The accumulated concentrations of CH2O are plotted
as a function of reaction time in Figure 10.
0.3
[CH2O] / mM
0.25
0.2
0.15
0.1
0.05
0
0
1000
2000
3000
4000
5000
6000
Time / s
Figure 10. Concentration of CH2O as a function of reaction time for the γ-radiolysis of water
in Tris (20 mM) with pH 7.0 ( ), 7.5 ( ), 8.0 ( ), 8.5 ( ) and 9.0 ( ). All the
experiments were purged with mixed N2O/O2. © 2014 American Chemical Society
It is clearly seen from Figure 10 that the yield of formaldehyde
increases with increasing pH and is doubled by raising pH from 7.0 to 9.0.
This fairly strong pH-dependence may indicate that the reaction yielding
formaldehyde is to some degree base-catalyzed.
28
3 Results and discussion
In addition, the pH effect on the heterogeneous system was
investigated with varied pH. The initial and final pH of each experiment
was measured and it was observed that the pH values were stable in all
cases except of the lowest pH. However, the changes in pH during the
course of the experiments are negligible. The production of CH2O is
plotted as a function of reaction time in Figure 11.
0.4
0.35
[CH2O] / mM
0.3
0.25
0.2
0.15
0.1
0.05
0
0
5000
10000
15000
20000
25000
30000
Time / s
Figure 11. Concentration of CH2O as a function of reaction time for the reaction of H2O2
(initial concentration 5 mM) with ZrO2 (4.5g) in Tris (20 mM) with pH: 7.0 (
( ), 8.5 ( ) and 9.0 (
Chemical Society
), 7.5 (
), 8.0
). All the experiments were exposed to air. © 2014 American
It can clearly be seen from Figure 11 that the production of CH2O
increases with increasing pH but the pH effect is more pronounced than
in the homogeneous solution (see Figure 10). It should be noted that a
4-fold increase in the production of CH2O is displayed when increasing
pH from 7.0 to 7.5. The CH2O production at pH 9.0 is almost one order of
magnitude higher than the lowest one (pH 7.0). Solution pH can not only
influence the yield of CH2O as in the homogeneous system, but also
change the surface charge and the adsorption of the scavenger in the
29
3 Results and discussion
heterogeneous system, thereby altering the relative coverage of reactants
on the surface.
The normalized evolution of CH2O as a function of the conversion of
H2O2 in the pH range under study is shown in Figure 12.
% of [CH2O]t/[CH2O]f
100
80
60
40
20
0
0
20
40
60
80
100
% of H2O2 converted
Figure 12. [CH2O]t/[CH2O]f as a function of conversion of H2O2 (initial concentration 5 mM)
in Tris (20 mM) with pH: 7.0 (
Chemical Society
), 7.5 (
), 8.0 (
), 8.5 (
) and 9.0 (
). © 2014 American
As can clearly be seen in Figure 12, pH does not induce a change in
the relative dynamics of the formation of CH2O.
3.1.1.3
Formaldehyde production mechanism by Tris and hydroxyl
radical
Based on the experimental results, a formaldehyde production
mechanism by Tris and hydroxyl radical is proposed as in Figure 13. The
scheme consists of two independent routes: O2- and pH- dependent
routes. Each route contains several possible reacting pathways.
30
3 Results and discussion
HO
O
(a)
+ HO
HO
HO
HO
NH
HO
HO
H
O
NH
C
H
HO
NH2
+
H
HO
HO
NH2
HO
(b)
β-scission
O
+ HO
HO
(Tris)
NH2
O
+
C
NH2
H
OH
HO
H
1, 2
-H
sh
ift
+ HO
HO
(c)
HO
+O2
HO
HO
HO
NH2
HO
O
-HO2
CHOO
NH2
NH2
HO
NH2
HO
HO
(d)
Cβ-Cγ
cleavage
HO
OH
+
+O2
CH2 OH
-HO2
O
OOCH2OH
C
H
H
Figure 13. Scheme for formaldehyde production mechanism by Tris. © 2014 American
Chemical Society
As shown in the scheme, aminyl, alkoxyl and hydroxyalkyl radical are
formed via hydrogen abstraction from Tris by hydroxyl radicals. The
alkoxyl radical can undergo 1, 2-hydrogen shift reaction and be converted
to a hydroxyalkyl radical.98, 99
Formaldehyde can be formed from the aminyl radical via a fivemembered cyclic transition state and this route is base-catalyzed and O2independent (Reaction (a)).100 In addition, formaldehyde can be formed
directly from the alkoxyl radical via an O2-independent β-scission
reaction (Reaction (b)). Furthermore, the hydroxyalkyl radical can
undergo two O2-dependent pathways (Reaction (c) and (d)) producing
aldehyde and formaldehyde, respectively. To produce formaldehyde, the
radical may undergo a fragmentation reaction by the cleavage of the Cβ-Cγ
bond (Reaction (d)) yielding •CH2OH—the precursor for formaldehyde in
O2-containing systems. The Cβ-Cγ cleavage reaction of alkyl radicals has
been reported101, 102 and may also take place in the present system.
31
3 Results and discussion
The effect of Tris on the catalytic decomposition of H2O2 on
ZrO2 surface IV
3.1.2
The two competing reactions: O2 and H2O2 oxidizing •CH2OH and H2O2
and Tris reacting with HO•, were evaluated under different conditions.
The final production of CH2O is plotted against [H2O2]0 in Figure 14.
0.40
0.35
2
f
[CH O] / mM
0.30
0.25
0.20
0.15
0.10
0.05
0.00
0
1
2
3
[H2O2]0 / mM
4
5
Figure 14. Final production of CH2O as a function of initial concentration of H2O2 under
different conditions: (N2 and 100 mM Tris) (black squares), (Air and100 mM Tris) (red
circles), (N2 and 20 mM Tris) (blue up triangles), (Air and 20 mM Tris) (cyan down triangles).
© 2015 Elsevier B. V.
In Figure 14, it can be seen that at a given [H2O2]0, higher [Tris]0
leads to higher [CH2O]f both in the presence and absence of O2. With a
given [Tris]0, [CH2O]f increases with increasing [H2O2]0. Nevertheless, it
is apparent that a 5-fold increase in [Tris]0 or [H2O2]0 does not result in a
5-fold increase in [CH2O]f (only by a factor of 2~3 in the studied
concentration range). This obvious inconsistency is due to the fact that
formaldehyde is formed via a surface reaction and the competition
between hydroxyl radicals reacting with Tris and H2O2 is governed by
relative surface coverage rather than relative bulk concentrations.
32
3 Results and discussion
Based on the assumption that ZrO2 behaves similarly as TiO2,62 i.e.,
H2O2 has higher affinity than Tris to ZrO2, a modified mechanism (R16)R20)) must be employed which includes two different surface sites: Sa
and Sb for H2O2 and Tris, respectively. This mechanism constitutes the
basis for numerical simulation of the kinetics of the H2O2/ZrO2/Tris
system. All kinetic parameters are assigned on the basis of experimental
data or reasonable estimations and listed in Table 4. The detailed
deduction process of some parameters is presented in Paper IV.11
Table 4. Modified mechanism of H2O2 decomposition on ZrO2 surface in the presence of
Tris and the related parameters. © 2015 Elsevier B. V.
No.
Reaction
Constant
Initial Conc./mM
-2
-1
-1
R16
H2O2 + 2Sa ⇌ H2O2(Sa)
k16f = 5.4 x 10 M s
k16r = 9 x 10-7 s-1
K16 = 6 x 104 M-1
R17
H2O2(Sa) →2HO•(Sa)
K17 = 3 x 10-4 s-1
R18
R19
R20
Tris + Sb ⇌ Tris(Sb)
k18f = 5 x 10-5 M-1 s-1
k18r = 5 x 10-7 s-1
K18 = 1 x 102 M-1
HO•(Sa) + H2O2(Sa) → H2O + HO2• + 3Sa
K19 = 1 x 108 M-1 s-1
HO•(Sa) + Tris(Sb) → CH2O + Sa + Sb
[H2O2]0 = 0.5-5
[Sa]0 = 2
[Tris]0 = 20, 100
[Sb]0 = 1
K20 = 1 x 108 M-1 s-1
As shown in Table 4, the sensitivity of the systems at the reactant
concentrations and adsorption capacity used in the simulations can be
summarized as: the surface is saturated with H2O2 and Tris under the
present conditions. In this case, increasing the adsorption equilibrium
constants (K16 and K18) has negligible effect on the dynamics of the system.
However, decreasing K16 results in higher final production of
formaldehyde whereas decreasing K18 results in lower final production of
formaldehyde. This is directly resulted from the competition between
H2O2 and Tris for surface bound hydroxyl radicals. In addition, the ratio
between k19 and k20 also has a direct impact on the final production of
formaldehyde.
Additional simulations were performed by using the kinetic simulation
software Gepasi 3.091, 92 to further evaluate the model. The initial Tris and
33
3 Results and discussion
5
0.10
4
0.08
3
0.06
2
0.04
1
0.02
0
0
[CH2O] / mM
[H2O2] / mM
H2O2 concentration in the simulation were kept the same as in the
experiments. The experimental dynamics of H2O2 decomposition is
presented in Figure 15, consisting of the time-resolved H2O2
concentration and accumulated CH2O concentration. The dynamics of
simulation is shown in Figure 16a and Figure 16b. In addition to [H2O2]
and [CH2O], the time-resolved [H2O2(Sa)] and [Tris(Sb)] in the simulations
are also presented.
0.00
5000 10000 15000 20000 25000 30000 35000 40000
Time / s
Figure 15. Concentration of H2O2 (black squares) and the production of CH2O (red circles)
as a function of reaction time in suspension with H2O2 (initial concentration 5 mM), ZrO2
(SA/V = 4.5 x 105 m-1) and 20 mM Tris at pH 7.5. © 2015 Elsevier B. V.
By comparing Figure 16a to Figure 15, it can be observed that the
dynamics in the simulation resembles the experimental counterpart.
More specifically, a fast decrease of the concentration of H2O2 in the
initial period can be seen in both cases indicating that H2O2 adsorbs onto
ZrO2 surface rapidly. The adsorption step is more readily observed in
Figure 16c and Figure 16d where the logarithmic axis is applied.
34
3 Results and discussion
4
[CH2O]
[H2O2]
[H2O2](Sa)
[Tris](Sb)
b
4 4
5
4
3
3 3
3
2
2 2
2
1
1 1
1
0
0
5
Concentration / mM
5 5
5000
10000
0 0
20000
0
15000
[CH2O]
[H2O2]
[H2O2](Sa)
[Tris](Sb)
c
4
5 5
5000
10000
0
20000
15000
[CH2O]
[H2O2]
[H2O2](Sa)
[Tris](Sb)
d
4 4
5
4
3
3 3
3
2
2 2
2
1
1 1
1
0
100
101
102
103
104
0 0
0
100
Time / s
Concentration / mM
Concentration / mM
[CH2O]
[H2O2]
[H2O2](Sa)
[Tris](Sb)
Concentration / mM
a
5
101
102
103
104
Time / s
Figure 16. The simulated concentrations of different species as a function of reaction time:
[CH2O] (black), [H2O2] (red), [H2O2(Sa)] (blue), [Tris(Sb)] (cyan). (a) 20 mM Tris, (b) 100 mM
Tris, (c) 20 mM Tris and logarithmic x axis, (d) 100 mM Tris and logarithmic x axis. © 2015
Elsevier B. V.
According to the higher adsorption constant (K16) and forward reaction
rate constant (k16f), H2O2 can reach the maximum surface coverage
instantly (~ 10 s) and remain as the dominating species for approximately
6000 s. Although slower than the adsorption of H2O2, Tris can saturate
the surface in approximately 500 s at a concentration of 100 mM and the
saturated surface concentration ([Tris(Sb)]) is approximately equal to the
saturated surface concentration of H2O2 ([H2O2(Sa)] = 1 mM).
In general, the H2O2 concentration exhibits a fast initial drop due to the
adsorption upon ZrO2 surface. Then the decomposition of H2O2 becomes
the rate limiting step, following first-order kinetics. Not surprisingly, in
the last period when [H2O2] in solution approaches zero, CH2O is still
produced until H2O2 on the surface diminishes completely.
The ratio between final production of formaldehyde and initial
concentration of H2O2 in the simulations as well as in the experiments
35
3 Results and discussion
0.12
1.2
0.10
1.0
0.08
0.8
0.06
0.6
0.04
0.4
0.02
0.2
[CH2O]f/[H2O2]0
[CH2O]f/[H2O2]0
(under N2 atmosphere) is plotted against [H2O2]0 and the results are
presented in Figure 17.
0.0
0.00
0
1
2
3
[H2O2]0 / mM
4
5
Figure 17. The experimental and simulation [CH2O]f/[H2O2]0 as a function of [H2O2]0 under
different conditions: 100 mM Tris, experimental (black squares), 20 mM Tris, experimental
(black circles), 100 mM Tris, simulation (red up triangles), 20 mM Tris, simulation (red down
triangles). © 2015 Elsevier B. V.
It can be seen from Figure 17 that, the trend in the simulation case is
similar to the trend in the experimental case. [CH2O]f/[H2O2]0 decreases
with increasing [H2O2]0 which is due to the competition between Tris and
H2O2 in the reaction with the surface bound hydroxyl radical.
Nonetheless, it should be noted that the experimental results are
approximately one order of magnitude lower than the CH2O productions
in the simulation (8% and 12% for 20 and 100 mM Tris, respectively).
This difference can be attributed to the low experimental yield in
converting hydroxyl radical to CH2O. As being determined in Paper III,
16% is the yield in homogeneous system.55 In the simulations the yield is
set as 100%. However, [CH2O]f/[H2O2]0 is higher at 100 mM than 20 mM
Tris in both cases.
36
3 Results and discussion
[CH2O]t/[CH2O]f is plotted against H2O2 conversion for the simulation
data and compared to the corresponding experimental curve55 (as
presented in Figure 18).
100
[CH2O]t/[CH2O]f / %
80
60
40
20
0
0
20
40
60
Converted H2O2 / %
80
100
Figure 18. The experimental (black) 55 and simulation (red) normalized evolution of CH2O
as a function of conversion of H2O2. [Tris]0 in both cases are set as 20 mM. © 2015 Elsevier
B. V.
As can be seen, the curve based on simulation data is similar to the
experimental one. In the initial adsorption-governed period, the reaction
between hydroxyl radicals and the adsorbed H2O2 (R(19)) is the primary
reaction which is attributed to the slower adsorption of Tris. As a result,
the initial rate of CH2O-formation is relatively low. However, the rate in
the experimental case is non-zero which is attributed to the fact that Tris
was added 30 min prior to the addition of H2O2. As H2O2(Sa) and Tris(Sb)
both reaches a steady-state concentration on the surface, the formation
rate of CH2O increases and remains constant until the steady-state
concentrations diminish.
37
3 Results and discussion
3.1.3
The interference of H2O2 on the modified Hantzsch method
using AAA as derivatization reagents VI
Experiments were performed regarding the influence of H2O2 on the
detection of CH2O at different [H2O2]0/[CH2O]0 ratios. The results are
presented in Figure 19. The experimental deviation (Dexp) is defined in
Eq. (21).
Dexp = 100% x ([CH2O]0 – [CH2O])/[CH2O]0
(21)
40
Deviation / %
30
20
10
0
1
10
100
1000
[H2O2]0/[CH2O]0
Figure 19. The experimental deviation (Dexp, black squares) as a function of [H2O2]0/[CH2O]0.
As shown in Figure 19, Dexp
[H2O2]0/[CH2O]0 and the deviation
[H2O2]0/[CH2O]0 reaches 20.
increases
becomes
with increasing
significant after
In order to figure out the reason for the observed deviation, additional
experiments were performed by adding concentrated H2O2 to the mixed
solution (H2O2, CH2O and derivazation reagents) after the incubation
38
3 Results and discussion
time (15 min). The absorbance of the formed derivative is monitored
before and after the addition of extra H2O2 (represented as Abs0 and Abs).
The result is shown in Figure 20 where ln(Abs/Abs0) is plotted as a
function of reaction time.
0.0
-0.2
-0.4
ln(Abs/Abs0)
-0.6
-0.8
-1.0
-1.2
-1.4
-1.6
0
200
400
600
800
1000
Time / s
Figure 20. ln(Abs/Abs0) of the derivative as a function of reaction time (black squares).
It can clearly be seen from Figure 20 that ln(Abs/Abs0) is linearly
correlated to the reaction time. Therefore, it is worthwhile to point out
that H2O2 reacts with the produced derivative and the reaction follows
first order kinetics under the present conditions. The observed deviation
of CH2O in Figure 19 should mainly be due to this irreversible reaction
consuming the produced derivative during the incubation.
3.1.4
1.
Summary
The yields of formaldehyde in the reaction between the hydroxyl
radical and methanol and Tris are both enhanced by increasing
39
3 Results and discussion
O2 concentration. For Tris, the yield also increases with
increasing solution pH. This is not observed for methanol.
2. The production of CH2O in the heterogeneous H2O2/ZrO2/Probe
system is affected by O2 and solution pH. Though, the dynamics
of the whole process is not likely influenced. Based on the results,
a mechanism for the CH2O production by Tris and HO• is
proposed including O2- and pH-dependent pathways.
3. Tris competes with H2O2 on reacting with hydroxyl radical,
thereby affecting the mechanism of the catalytic decomposition of
H2O2 on ZrO2 surface. A modified mechanism is proposed
including two independent surface adsorption sites for H2O2 and
Tris on ZrO2 surface. The simulated results based on the model
are in good agreement with the experimental results.
4. H2O2 reacts with the formed derivative by the modified Hantzsch
method, thereby lowering the accuracy of the method.
3.2 Inhibition of radiation induced dissolution of
UO 2 by sulfide I
The experiments were performed by using a UO2 pellet exposed to γ
irradiation in aqueous solution containing sulfide and NaHCO3. The
concentration of U(VI) is plotted against the reaction time as presented in
Figure 21.
As can be seen in Figure 21, the rate of radiation induced dissolution
of UO2 is reduced by the addition of sulfide and the reduction is
dependent on the sulfide concentration. Specifically, the inhibition lasts
for a given time after which the formation of U(VI) accelerates. The
period is longer with higher sulfide concentration and the length of the
periods roughly corresponds to the times required to produce equal
amounts of radiolytic oxidants under the present conditions (dose rate).
40
3 Results and discussion
10.0
[U(VI)] / µM
8.0
6.0
4.0
2.0
0.0
0
5000
10000
15000
20000
25000
30000
35000
Time / s
Figure 21. Concentration of U(VI) plotted against irradiation time with different sulfide
concentrations: 0 mM (
), 1 mM (
), 2 mM (
). © 2013 Elsevier B. V.
The inhibition of sulfide was confirmed. Furthermore, the effect was
compared to that of H2/Pd by studying their capacity on the reduction of
U(VI) in aqueous solution. The results are shown in Figure 22.
From Figure 22, it is observed that the reduction of U(VI) by H2 in the
absence of Pd is negligible in the present work which is in line with
previous observations.25, 47 Moreover, the reduction of sulfide appears to
be completely independent of the presence of H2 without a catalyst.
The reduction rate in H2/Pd system is much higher than sulfide within
the concentrations used in the present work. The reduction rate in the
H2/Pd system in the absence and presence of 1 mM sulfide are identical
which indicates that there is no poisoning effect of sulfide on Pd catalyst
in the present work. On the other hand, there is no synergy of sulfide and
H2/Pd on the reduction of U(VI).
41
3 Results and discussion
180
160
[U(VI)] / μM
140
120
100
80
60
40
20
0
0
5000
10000
15000
20000
25000
30000
35000
Time / s
Figure 22. Concentration of U(VI) plotted against reaction time at 20 bar H2 or 1.5 bar N2
under different conditions: 20 bar H2 (
mM sulfide (
sulfide + Pd (
); 20 bar H2 and 1 mM sulfide (
); 1.5 bar N2 and 2 mM sulfide (
); 20 bar H2 and Pd (
); 20 bar H2 and 2
); 20 bar H2 + 1 mM
). © 2013 Elsevier B. V.
3.3 Kinetics, mechanistic and simulation study of
catalytic decomposition of H 2 O 2 on metal and
metal oxide surfaces
3.3.1
Cu, Cu2O and CuO VIII
Cu, Cu2O and CuO were investigated under the same conditions
([H2O2]0 = 5 mM, [CH3OH]0 = 100 mM) and similar SA/V (4600 ± 200
m-1). Since the solubilities of Cu2O and CuO under the present condition
are relatively low (3 and 10 µM, respectively),103 the solution reactions are
assumed to be negligible. In Figure 23, the H2O2 consumption and the
CH2O production as function of reaction time is plotted for the three
materials.
42
3 Results and discussion
Not surprising, as can be seen in Figure 23, the final production of
formaldehyde in the CuO case is the highest since catalytic decomposition
is the dominating surface reaction pathway for consuming H2O2 in this
system. While the lower CH2O production in the Cu and Cu2O case is
because smaller fractions of H2O2 are decomposed to hydroxyl radical as
a result of the competition between surface oxidation and catalytic
decomposition on consuming H2O2.
0.20
0.08
(a)
(b)
5
0.07
4
0.06
4
0.12
3
[CH2O] / mM
[H2O2] / mM
[H2O2] / mM
0.16
0.05
3
0.08
2
1
0.04
0
0.00
5000
0
1000
2000
3000
4000
0.03
0.02
1
0.01
0
0
Time / s
500
1000
1500
2000
2500
0.00
3000
Time / s
0.7
(c)
5
0.04
2
[CH2O] / mM
5
0.5
3
0.4
0.3
2
0.2
1
[CH2O] / mM
[H2O2] / mM
0.6
4
0.1
0.0
0
0
5000
10000
15000
20000
Time / s
Figure 23. Concentration of H2O2 (black squares) and CH2O (red squares) as a function of
reaction time for (a) Cu, (b) Cu2O, and (c) CuO powder. © The Royal Society of Chemistry
2015
3.3.2
Other lanthanide and transition metal oxides II
A number of transition and lanthanide oxides were studied, including
CeO2, Fe2O3, HfO2, Gd2O3 and CuO powder. The kinetic studies were
performed by following the concentration of H2O2 with reaction time. The
results are shown in Figure 24.
43
3 Results and discussion
1.0
[H2O2]/[H2O2]0
0.8
0.6
0.4
0.2
0.0
0
5000
10000
15000
20000
25000
Time / s
Figure 24. Normalized H2O2 concentration ([H2O2]0 = 0.5mM in 50 mL) as a function of
reaction time for the reaction with CuO (
), CeO2 (
), HfO2 (
), Gd2O3 (
) and Fe2O3
( ) at T = 298.15 K SA: Fe2O3 (4.5 m2), CeO2 (7.5 m2); CuO (0.3 m2); HfO2 (7.5 m2); Gd2O3
(1.7 m2). © 2013 Elsevier B. V.
It can be seen from Figure 24 that, except of Gd2O3, the other oxides
show only slightly different reactivity towards H2O2 under the present
conditions.
The first-order rate constants k1 obtained from the data of Figure 24
are given in Table 5. The second order rate constants and the intercepts
at the zero coordinate (b2) of the least squares fits for the plots of k1 as a
function of SA/V are given in Table 5.
Table 5. Obtained k1, k2 and b2 for decomposition of H2O2 (0.5mM; 50 mL) catalyzed by
different oxides at T = 298 K. k1 was obtained with the following SA of oxides: Fe2O3 (4.5
m2), CeO2 (7.5 m2); CuO (0.3 m2); HfO2 (7.5 m2); Gd2O3 (1.7 m2). © 2013 Elsevier B. V.
Material
Fe2O3
CeO2
CuO
HfO2
Gd2O3
k1 (s-1)
k2 (m·s-1)
-4
(2.1 ± 0.4) × 10
(1.7 ± 0.5) × 10-4
(1.90 ± 0.05) × 10-4
(4.3 ± 0.9) × 10-4
(3.6± 0.3) × 10-5
b2 (s-1)
-9
(3.0 ± 0.06) × 10
(2.80 ± 0.07) × 10-8
(1.23 ± 0.06) × 10-9
(2.78 ± 0.02) × 10-9
(9.4 ± 1) × 10-10
44
2 × 10-4
5 × 10-6
6 × 10-6
1 × 10-5
6 × 10-6
3 Results and discussion
As can be seen in Table 5, the values of k2 deviate significantly from
the expected diffusion controlled reaction rate constants for particles of
this range of sizes (in the order of 10-5 m·s-1).104 In addition, the value of
b2 for Fe2O3 is significantly higher than for the other oxides. This
indicates that for Fe2O3, a homogeneous Fenton reaction in the bulk is
also occurring due to the presence of dissolved Fe2+. This extra reaction
has increased significance on the overall reaction kinetics when compared
with the other materials.105
The temperature intervals for determining the Arrhenius activation
energies were set as T = (298–334) K for Fe2O3, CuO, HfO2 and T = (298–
353) K for CeO2 and Gd2O3. The obtained Arrhenius energies are shown
in Table 6.
Table 6. Arrhenius activation energies (Ea) and frequency factors (A) for the decomposition
of H2O2 catalyzed by different oxide materials. © 2013 Elsevier B. V.
Material
Ea (kJ·mol-1)
A (s-1)
Particle size
Fe2O3
CeO2
CuO
HfO2
Gd2O3
47 ± 1
40 ± 1
76 ± 1
60 ± 1
63 ± 1
2.2 × 103
1.4 × 103
3.5 × 109
1.1 × 107
3.4 × 106
< 5 μm
14 μm
< 50 nm
44 μm
12 nm
It can be seen from Table 6 that the obtained Ea values differ for the
different oxides studied. This indicates that the activation energies are
most likely determined by microstructural properties of the particles, e.g.
the type of atoms present at the catalytically active surface sites and the
extent of hydroxylation at these sites. The differences in the values of
frequency factors indicate that surface catalytic capacities of the oxides
are very different. The differences may be attributed to several reasons.
The most obvious one is that the number of active surface sites (e.g.
adsorption sites) differs for different oxides. Other possible reasons are
the different numbers of catalytically active sites at the interfaces and the
strength of bonding between decomposition products upon H2O2 and
these active sites.
The dynamics of formation of the hydroxyl radicals during catalytic
decomposition of H2O2 on the oxides were studied by using the probe Tris.
45
3 Results and discussion
The concentration of formed CH2O is plotted against reaction time in
Figure 25.
0.6
[CH2O] / mM
0.5
0.4
0.3
0.2
0.1
0
0
10000
20000
30000
40000
Time / s
Figure 25. Concentration of CH2O as a function of reaction time under the conditions: 5 mM
H2O2, 20 mM Tris, 50 mL for different oxides. Fe2O3(
CuO (
), CeO2 (
), HfO2 (
), Gd2O3 (
),
). © 2013 Elsevier B. V.
It can be seen from Figure 25 that the dynamics of formation of CH2O
for the different oxides vary considerably.
In Figure 26 the concentration of CH2O as a function of relative
conversion of H2O2 is shown. However, this can be optimized by using
normalized CH2O concentration ([CH2O]t/[CH2O]f) instead. (see Paper
III and IV) The initial reaction conditions are the same in all
experiments ([H2O2]0= 5 mM and [Tris]0= 20 mM).
As can be seen in Figure 26, it is obvious that the CH2O production is
very low up to a certain conversion of H2O2 for a few of the oxides. These
curves resemble the one shown in above section (Figure 18) and the low
CH2O production period is the time needed for H2O2 and Tris to reach a
steady state on the oxide surface. When the steady state for both H2O2
46
3 Results and discussion
and Tris is obtained, the formation rate of CH2O starts to increase,
displayed as an inflection point in the figure.
0.6
0.5
[CH2O] / mM
0.4
0.3
0.2
0.1
0
0
20
40
60
80
100
% of H2O2 consumed from solution
Figure 26. Concentration of CH2O present in the reaction system as a function of the
percentage of H2O2 consumed from solution during reaction with different oxide materials.
CeO2 ( ), Gd2O3 ( ), Y2O3 ( ), ZrO2 (
UO2 ( ). © 2013 Elsevier B. V.
), Fe2O3 (
), HfO2 (
), CuO (
), TiO2 ( ) and
For some oxides, e.g. TiO2 and Gd2O3, a clear inflection point is never
reached under the present conditions and very small amounts of
formaldehyde are formed. It may be attributed to that the relative surface
coverage of H2O2 is much higher than Tris, thereby most of H2O2 are
consumed via reacting with HO• other than reacting with Tris to produce
CH2O.
For other more extreme cases such as HfO2, an inflection point appears
when the H2O2 is almost completely consumed form the solution.
However, it is not possible to attribute the position of the inflection points
as the result of a single effect for all the materials.
In the UO2 case, the production of CH2O reaches a plateau
corresponding to 20% of H2O2 consumed. The CH2O concentration does
not increase afterwards regardless of the continuous consumption of
47
3 Results and discussion
H2O2 from solution. H2O2 can react with UO2 both via catalytic
decomposition and oxidative dissolution.35 The hydroxyl radicals
produced in the oxidative dissolution of UO2 cannot be scavenged by
Tris.106 The observed plateau of CH2O may be attributed to that catalytic
decomposition is hindered since the regenerated fresh surface is
dissolved according to the oxidative dissolution.
3.3.3
1.
Summary
The competition between surface oxidation and catalytic
decomposition on consuming H2O2 is the main reason for the lower
production of CH2O in the Cu and Cu2O cases than that in the CuO
case.
2. H2O2 can be catalytically decomposed on the surface of a number of
oxides. The reactivity of H2O2 and dynamics of CH2O formation is
not the same for all the studied materials.
3.4 Kinetics and mechanistic study of reactions
between W and H 2 O 2 or O 2
3.4.1
W and H2O2 V
The XRD, SEM and XPS measurements were conducted to investigate
the crystalline structure and composition of the tungsten powder before
and after treatment. A typical X-ray diffraction pattern is shown in
Figure 27a.
It can be seen from Figure 27a that there are XRD peaks attributed to
two crystalline phases present in the W powder. One group of peaks
belongs to metallic W which possess bcc (base-center-cubic) structure,107
the other group belongs to the so-called A15-W phase where the tungsten
unit cell is stabilized by oxygen.108, 109 Note that, the XRD patterns
recorded from the W powder after different treatments, considered in the
current study, look quite similar to those for the untreated powder. This
observation is supported by our SEM studies (see Figure 27b and
48
3 Results and discussion
Figure 27c) where the surface morphology of W powder before and after
the reaction with H2O2/Tris is presented.
Figure 27. (a) X-ray diffraction pattern of untreated tungsten powder. Two types of
crystalline structures of tungsten and their crystalline phases are labelled on the graph.
SEM image of (b) untreated tungsten powder; (c) tungsten powder after reaction with
Tris/H2O2.
It can be seen from Figure 27b and Figure 27c that the W powder
consists of two types of particles (faceted crystal and porous). These two
phases are most probably corresponding to metallic W and A-15 WOx
determined by XRD analysis. Moreover, no significant changes in the
surface morphology of W powder before and after treatment were
49
3 Results and discussion
observed. Also, no other phases than metallic tungsten and A-15 have
been detected by XRD and SEM.
The composition of tungsten powder before and after reaction with Tris
and H2O2 was studied by means of XPS. Since the penetration depth of
photoelectrons is about 6 nm, mostly surface composition was studied.
The results are shown in Table 7.
Table 7. XPS results about atomic concentration (AC) of different W species on the surface.
C 1s
O 1s
W4f7/2
BE,
eV
W
ref.,
at %
W (2g)
and
H2O2,
at. %
W (2g)
and
Tris,
at %
W
(2g)
and
Tris/H2O2
, at %
W
(1g)
and
Tris/H2O2
, at %
Chemical
binding,
compounds
284.5
16.3
11.5
15.2
11.9
19.9
C-(C,H)
286.3
2.9
5.2
6.7
5.5
7.9
C-OH
289.0
1.1
2.5
2.9
2.2
4.7
COOH
530.9
49.4
47.2
45.1
46.2
33.6
W-O,
531.9
5.8
8.0
7.1
6.6
6.3
W-OH, C=O
533.1
3.2
3.8
3.1
3.6
5.0
C-OH
31.4
3.9
5.2
3.7
7.7
8.7
W metal110
32.0
1.6
1.9
1.8
2.7
4.0
W-O, W-C111
35.9
15.8
14.8
14.5
13.6
10.0
WO3112
BE: Binding energy
As seen in Table 7, the surface of W powder consists mostly of metallic
tungsten and WO3. No significant changes on the amount of metallic
tungsten and WO3 were observed after being treated with H2O2 and Tris.
Generally, XPS, XRD and SEM studies reveal that W powder does not
undergo significant changes in the composition and structure when react
with Tris and H2O2.
The kinetics of H2O2 decomposition and formation of CH2O in the
W(s)/H2O2/Tris system at varied amount of tungsten powder was
investigated. The results are shown in Figure 28.
50
3 Results and discussion
6
(a)
[H2O2] / mM
5
4
3
2
1
0
0
4000
8000
12000
16000
Time / s
20000
24000
8000
12000
16000
Time / s
20000
24000
0.35
(b)
0.30
[CH2O] / mM
0.25
0.20
0.15
0.10
0.05
0.00
0
4000
Figure 28. (a) Concentration of H2O2 and (b) CH2O as a function of reaction time with
varied amounts of tungsten powder: SA/V = 8800 m–1 (black squares), SA/V = 4400 m–1
(red circles), SA/V = 2200 m–1 (green up triangles), SA/V = 1100 m–1 (blue down triangles),
SA/V = 550 m–1 (cyan diamonds). All cases are under the condition: [H2O2]0 = 5 mM, [Tris] =
100 mM, N2 and pH 7.5.
It can be seen from Figure 28a that the consumption of H2O2 appears
to be almost independent of the amount of tungsten powder except in the
case of the lowest amount (550 m–1). Generally, the H2O2 concentration
appears as a slow initial decrease followed by accelerated consumption.
This acceleration indicates that the reaction is autocatalyzed. As be seen
in Figure 28b, the final production of formaldehyde decreases with
increasing the amount of tungsten powder. Moreover, an initial slow
phase followed by an acceleration of CH2O evolution can be observed in
all cases. This slow process producing CH2O (Haber-Weiss reactions) will
be discussed below and the period corresponds to the time required to
build up a sufficiently high concentration of W(V). The increasing CH2O
production rate indicates that H2O2 is initially consumed through
oxidation of W. The produced dissolved tungsten species is capable of
catalytically decomposing H2O2 to hydroxyl radicals.
51
3 Results and discussion
6
(a)
[H2O2] / mM
5
4
3
2
1
0
0
4000
8000
12000
16000
20000
24000
Time / s
[CH2O] / mM
0.30
(b)
0.25
0.20
0.15
0.10
0.05
0.00
0
4000
8000
12000
16000
20000
24000
Time / s
0.25
(c)
[W] / mM
0.20
0.15
0.10
0.05
0.00
0
4000
8000
12000
16000
20000
24000
Time / s
Figure 29. Concentration of (a) H2O2, (b) CH2O and (c) dissolved tungsten species as a
function of reaction time in non-filtered (black squares) and filtered (red circles) system.
Filtration was done at 1860 s. The experiment was done under the condition: SA/V = 4400
m−1, [H2O2]0 = 5 mM, [Tris] = 100 mM, N2 and pH 7.5.
Considering the possibility of reactions between H2O2 and tungsten
ions in homogeneous solution, the solid was removed by filtration from
the W(s)/H2O2/Tris system. The concentrations of H2O2 and CH2O in the
filtered solution as well as in the original non-filtered solution were
monitored and compared. The concentration of dissolved tungsten
species was also measured at the same sampling time points. W(VI) is the
dominating stable ionic tungsten species in solution under basic
condition according to the reported Eh-pH diagram of the W-H2O
system.113 The results are presented in Figure 29. The standard redox
52
3 Results and discussion
potential of W(VI)/W(V) couple is -0.03 V which is much lower than that
of Fe(III)/Fe(II) (+0.77V).67 Hence, it is more difficult for H2O2 to reduce
W(VI) to W(V) via R(13) than to reduce Fe(III) to Fe(II). In the
H2O2/Fe(III) system, the concentration of Fe(II) is insignificant as
compared to that of Fe(III).28 Similar as for Fe(III)/Fe(II), it can be
assumed that [W(V)] << [W(VI)] ≈ [W]total.
It can be seen from Figure 29a that in the non-filtered solution the
consumption of H2O2 is accelerated with reaction time while in the
filtered solution no acceleration is observed. Nevertheless, H2O2 is also
consumed in the filtered solution which implies that the product
responsible for the autocatalysis is only formed in the presence of solid W.
Moreover, the major part of the H2O2 is consumed by the soluble Wspecies.
From Figure 29b, it can be observed that H2O2 is catalytically
decomposed in both the filtered and the non-filtered samples. The
formaldehyde production accelerates more in the non-filtered sample
than in the filtered sample. This is consistent with the observed
acceleration in the H2O2 consumption for the non-filtered sample. The
initial slow phase producing CH2O exactly corresponds to the period to
build up a sufficiently high concentration of W(V), while in the nonfiltered case the phase is shorter. Similar as in the H2O2/Fe(III)/Fe(II)
system where the Fe(II) concentration is determined by the total iron
concentration,73 sufficient W(V) is readily obtained in the non-filtered
samples due to the continuous dissolution of W.
It is observed in Figure 29c that the W concentration in solution
increases almost linearly with time during the initial phase of the reaction
and then reaches a maximum level. By comparison of Figure 29c and
Figure 29b, it can be seen that the production of formaldehyde remains
fairly low during the initial phase when W is being released to solution.
This indicates that the catalytic decomposition of H2O2 is mainly
attributed to solution reactions.
The kinetics of H2O2 consumption at different W-concentrations in
solution was studied in filtered samples extracted from the
53
3 Results and discussion
W(s)/H2O2/Tris system. The experimental details can be seen in Paper V.
The results are presented in Figure 30.
0.0
(a)
ln([H2O2]/[H2O2]0)
-0.5
-1.0
-1.5
-2.0
-2.5
-3.0
-3.5
0
5000
10000
15000
20000
25000
Time/s
7.0x10-4
6.0x10-4
5.0x10
kobs / s-1
(b)
y= (1.50 +/- 0.01)x + 9.96E-6
R2 = 97.1%
-4
4.0x10-4
3.0x10-4
2.0x10-4
1.0x10-4
0.0
0.0
1.0x10-4
2.0x10-4
3.0x10-4
4.0x10-4
[W] / M
Figure 30. (a) ln([H2O2]/[H2O2]0) as a function of reaction time in different experiments:
experiment NO. 1 (black squares), 2 (red circles), 3 (green up triangles), 4 (blue down
triangles), 5 (cyan diamonds), 6 (magenta left triangles), 7 (yellow right triangles), 8 (dark
yellow hexagons) and 9 (navy stars). (b) Observed first-order rate constants of experiments
NO. 1- 9 as function of [W] (black squares).
From Figure 30a, it is clear that H2O2 is consumed by W species in
solution. This reaction follows first order kinetics.
It is also shown in Figure 30b that the first order rate constant is
linearly dependent on the W-concentration corresponding to a second
order rate constant of (1.50 ± 0.01) M−1 s−1. The intercept obtained is
close to zero which implies that the only reaction responsible for the
consumption of H2O2 is the reaction with W in solution. Given the
concentrations of the reactants, first order kinetics would not be observed
if the solute W-species was consumed as well. Hence, the solute Wspecies must be a catalyst for the decomposition of H2O2.
54
3 Results and discussion
It has been proven in Paper V that the addition of Tris can significantly
improve the decomposition rate of H2O2 in the presence of dissolved
tungsten.
1.0
(a)
[W] / mM
0.8
0.6
0.4
0.2
0.0
[W] / mM
0
0.30
0.25
0.20
0.15
0.10
0.05
0.00
4000
6000
8000 10000 12000 14000 16000
Time / s
(b)
0
rdiss. / 10-8 M s-1
2000
8
7
6
5
4
3
2
1
0
2000 4000 6000 8000 100001200014000160001800020000
Time / s
(c)
0
1000 2000 3000 4000 5000 6000 7000 8000 9000 10000
SA/V / m-1
Figure 31. Concentration of dissolved tungsten species ([W] is used throughout the thesis
to represent total concentration of dissolved tungsten species) as a function of the reaction
time under different conditions: (a) WO3 + H2O (black squares), WO3 + H2O + Tris (red
circles), WO3 + H2O + Tris + H2O2 (green up triangles), SA/V = 140400 m–1. (b) W + H2O
(black squares), W + H2O + Tris (red circles), W + H2O + H2O2 (green up triangles), W +
H2O + Tris + H2O2 (blue down triangles), SA/V = 4400 m–1. (c) Dissolution rate (rdiss.) of
55
3 Results and discussion
tungsten species (black squares) as a function of tungsten powder amount (SA/V) under the
condition: [H2O2]0 = 5 mM, [Tris] = 100 mM, N2 and pH 7.5.
Experiments on oxidative dissolution of W powder by H2O2 (5 mM)
were performed in anoxic aqueous solution with H2O2 and/or Tris (100
mM). XPS result of W powder (Table 7) shows that trace amount of WO3
was found on the surface of tungsten powder, so parallel experiments
were also performed for WO3. The results are shown in Figure 31.
As can be seen in Figure 31a, the dissolution of W from WO3 in water
is enhanced by Tris. In this process, the presence of H2O2 is insignificant.
It is clearly seen from Figure 31b that both H2O2 and Tris influence
the dissolution of W. The combined effect between H2O2 and Tris is most
significant which indicates that oxidative dissolution of W in the presence
of Tris is the key-process.
In Figure 31c, it is also clear that the dissolution rate of W depends on
the amount of solid W.
Based on all results, a mechanism (Figure 32) is proposed regarding
H2O2 reacts with metallic tungsten powder in the presence of anoxic Tris
solution, including both heterogeneous and homogeneous reaction
pathways.
H 2O 2
W(s)
OH
H 2O 2
OH
(ii)
(i)
W(n)(s)
(iv)
W(VI)(aq)
(0 < n < 6)
HO2
H 2O 2
W(V)(aq)
(Tris)
(v)
H 2O 2
(iii)
OH
H 2O 2
OH
(vi)
OH
Tris
CH2O
Figure 32. Reactions between H2O2 and metallic tungsten in anoxic Tris solution.
The whole process is initiated when H2O2 is added to the suspension of
tungsten powder and Tris. Firstly, H2O2 adsorbs on the surface of
56
3 Results and discussion
tungsten powder triggering the oxidative dissolution of W (i) and (ii) and
catalytic decomposition of H2O2 on surface (iii). The production of
hydroxyl radicals is accompanying these two surface reactions, although
the production is probably negligible as compared to that in the solution
reactions (see Figure 28). Hydroxyl radicals produced during the
oxidative dissolution has been reported in previous studies.69, 70 Though,
it is impossible to distinguish among the surface reactions, (i), (ii) and
(iii). The dissolved tungsten (mostly W(VI)) reacts with H2O2 via the
Haber-Weiss reactions in the solution ((iv) and (v)) which form hydroxyl
radicals continuously. This process is facilitated by the presence of Tris.
Regardless if produced via surface reactions ((i), (ii) and (iii)) or solution
reactions ((iv) and (v)), the hydroxyl radicals will be scavenged by probes
to produce formaldehyde (vi).
In the W(s)/H2O2/Tris system, the overall consumption rate of H2O2 can
be defined as:
−
𝑑[𝐻2 𝑂2 ]
𝑑𝑑
= 𝑘2𝑠𝑠𝑠𝑠
𝑆𝑆
𝑉
[𝐻2 𝑂2 ] + 𝑘2𝑠𝑠𝑠 [𝑊][𝐻2 𝑂2 ]
(22)
where k2surf and k2sol represents the second-order rate constant for the
surface and solution reactions, respectively. The latter has already been
determined to (1.50 ± 0.01) M−1 s−1 from the slope of Figure 30b. Using
this rate constant and the concentration of W and H2O2, the rate of H2O2
consumption in the homogeneous phase can be calculated. These rates
are very close to the corresponding slopes of Figure 28a, again proving
that the H2O2 consumption is dominated by the solution reactions.
3.4.2
W and O2 VII
The experiments were performed in the aerobic or anaerobic Tris and
methanol solution. The results are presented in Figure 33.
As can be seen from Figure 33a and Figure 33d, both concentrations
of dissolved tungsten and CH2O are significantly enhanced by O2 in both
the Tris and methanol cases. This indicates that hydroxyl radicals are
produced during the oxidative dissolution of W. This is in line with the
similar systems where H2O2 is produced and then decomposed to
57
3 Results and discussion
hydroxyl radicals via the Haber-Weiss reactions.71, 114-117 (Paper V). The
detected concentration of H2O2 is marginal which should be due to the
fact that H2O2 is consumed rapidly in the present system. Furthermore, a
slight delay in the production of CH2O is observed in Figure 33c and 33d
for both Tris and methanol. This may be attributed to the delay in
producing surface catalytic sites for H2O2 decomposition as was shown in
Paper V.
500
500
400
400
300
300
200
200
100
100
0
0
4000
300
16000
0
(c)
250
[CH2O] / uM
8000
12000
Time / s
4000
(d)
0
8000 12000 16000 20000 24000
Time / s
300
250
200
200
150
150
100
100
50
50
0
0
4000
8000
12000
Time / s
16000
0
4000
[W] / uM
(b)
[CH2O] / uM
[W] / uM
(a)
0
8000 12000 16000 20000 24000
Time / s
Figure 33. Concentration of dissolved tungsten ([W]) and CH2O in aerobic (Air, black
squares) and anaerobic (N2, red circles) as a function of reaction time under the condition:
(a) and (c) Tris (100 mM, pH 7.5), (b) and (d) methanol solution (100 mM, unadjusted pH).
The concentration of formed CH2O was plotted against the
concentration of dissolved tungsten species in Figure 34. Every point in
the figure corresponds to [CH2O] and [W] measured from the same
sample (same reaction time). The figure is based on the data already
presented in Figure 33.
58
3 Results and discussion
250
y = (1.02 +/- 0.04)x - 181.36
R2 = 99.1%
[CH2O] / uM
200
150
Tris
Methanol
100
y = (0.33 +/- 0.01)x - 4.83
R2 = 99.4%
50
y = 0.16x - 1.20
R2 = 100%
0
0
100
200
300
[W] / uM
400
500
Figure 34. Concentration of formaldehyde in W (s)/air/Tris (black squares) and
W(s)/air/methanol system (red circles) as a function of the amount of dissolved tungsten.
It is observed in Figure 34 that the slope of [CH2O]/[W] in the Tris
case is half of that in the methanol case for low [W] (< 200 µM). During
this period, the difference in the slope is virtually identical to the
difference in the conversion yield Y reported in Paper III (29% and 68%
for Tris and methanol in the presence of O2, respectively).55 This implies
that the production rate of HO• is the same for both systems at fairly low
[W] where surface reactions are expected to dominate. Interestingly, the
slope of [CH2O]/[W] increases at higher [W] for Tris. The slope becomes
three times higher than that in the methanol case. A plausible explanation
for this is that the homogeneous Haber-Weiss reactions become
dominating at high [W] in the presence of Tris. This is in line with the
observations on the impact of Tris and methanol. The homogeneous
Haber-Weiss reactions are facilitated by Tris but not by methanol. (Paper
V and VII).
59
3 Results and discussion
3.4.3
1.
Summary
Homogeneous Haber-Weiss reactions are mainly responsible for
the consumption of H2O2 and formation of CH2O in the aqueous
W(s)/H2O2/Tris system. Another two reactions in this system are
oxidative dissolution of W and catalytic decomposition of H2O2
occurring at the interface. A mechanism including these reactions
in the present system was proposed.
2. A slow phase of CH2O formation in both filtered and non-filtered
system corresponds to the time to build up sufficient W(V) to
accelerate the homogeneous Haber-Weiss reactions to produce
hydroxyl radicals continuously.
3. Oxidative dissolution of W by H2O2 is evident and can be
facilitated in the presence of Tris.
4. Oxidative dissolution of W by O2 is also evident and hydroxyl
radicals are formed during the dissolution. The key step is the
formation of H2O2.
60
4 Conclusions
4 Conclusions
The aim of the thesis was to further investigate radiation induced
processes in heterogeneous systems including surface reactions (between
H2O2 and metal/metal oxide surfaces, i.e. redox reactions and catalytic
decomposition) as well as solution reactions (Haber-Weiss reactions).
Based on the studies presented in the thesis, the following conclusions
can be drawn:
•
H2O2 oxidatively dissolves W. The reaction is facilitated by the
presence of Tris. In the W(s)/H2O2/Tris aqueous system, after a
short initial period of dissolution, solution reactions (homogeneous
catalysis) become the dominating reactions consuming H2O2. The
later reaction is accompanied by formation of hydroxyl radicals
(monitored via the formation of CH2O).
•
Hydroxyl radicals are formed also during the O2-induced
dissolution of W. This is the result of the Haber-Weiss reactions
between the formed intermediate H2O2 and dissolved W producing
HO• continuously.
•
The competition between the two surface reactions on consuming
H2O2 is the main reason for the lower production of CH2O in the Cu
and Cu2O cases than that in the CuO case. The catalytic
decomposition of H2O2 is the only surface reaction in the CuO case.
•
Sulfide inhibits the radiation induced dissolution of UO2 and the
inhibition is weaker than H2/Pd in the present work. However,
sulfide does not interfere with the inhibition process by H2/Pd.
•
The mechanism of catalytic decomposition of H2O2 on ZrO2 surface
is not likely affected by changing the dissolved O2 concentration
and solution pH. However, the consumption rate of H2O2 is
affected by the competition between Tris and H2O2 reacting with
surface bound hydroxyl radical.
•
CH2O production by probes (Tris and methanol) and HO• include
both O2-dependent and O2-independent pathways. There is a pH-
61
4 Conclusions
dependent pathway for Tris however not for methanol. The
accuracy of the modified Hanztsch method for detection of CH2O is
affected by the presence of H2O2 via the reaction with the formed
derivative.
62
5 Future work
5 Future work
To the best of my knowledge, this thesis presents some of the first
systematic studies of the kinetics and mechanism of radiation induced
processes in heterogeneous system. In addition, the methodology for the
detection of hydroxyl radicals is also studied to a large extent. Based on
the results, the experimental design of kinetics and mechanistic studies
on the heterogeneous system can be further developed.
For some metals and metal oxides (e.g. Fe2O3, UO2) which can be
oxidatively dissolved by H2O2, filtrations at selected time points are
desirable. The modification is in order to compare the reactivity of H2O2
and production of CH2O in both filtered and non-filtered samples. The
time-resolved dissolution of studied materials can also be monitored.
Consequently, the mechanism which may include both surface and
solution reactions can be elucidated.
The kinetics, the dependence on the concentration of dissolved species
and the effect of ligands on the solution reactions in the filtered
homogeneous system could be interesting to study.
The radiation induced processes on the aqueous W(s)/Tris system may
be studied by using γ-irradiation. The surface structure and composition
of tungsten powder and the formation of CH2O after irradiation would be
interesting to study.
Furthermore, the study on the temperature dependence for the
formation of CH2O in both homogeneous and heterogeneous
H2O2/MxOy/Probe systems would give valuable information for
understanding the interfacial reactions.
63
7 Acknowledgements
6 Acknowledgements
The PhD journey would not have been so wonderful without the help
and support on both academic and daily life from many enthusiastic
individuals in the last four years. I would like to express my most sincere
gratitude and appreciation to all of you:
Prof. Mats Jonsson, my supervisor. I really appreciate you for
providing me the opportunity to work in this project, leading me the way
into the academia and orienting me during the journey with full support,
endless patience and continuous encouragement.
Associate Prof. Susanna Wold, my co-supervisor. Thank you so much
for sharing your professional knowledge and support.
Prof. An Lin, Prof. Fuxin Gan, Prof. Dihua Wang, my former
supervisors in Wuhan University, China. Thank you for introducing me
into the scientific world and supporting me to study abroad. Special
thanks for Prof. An Lin for immense help and suggestions on my career.
Prof. Daqing Cui at Stusvik. I really appreciate you for sharing
professional knowledge and your immense help on my future career.
Prof. Lars Kloo. I really appreciate you for sharing professional
knowledge and your invaluable suggestions on my thesis.
Sara Sundin, Claudio Miguel Lousada. Thank you very much for
guiding me into my doctoral study, teaching me the basic knowledge and
the experimental methodology. Also thanks for the collaboration and
invaluable discussions.
Åsa Björkbacka, Karin Norrfors, Alexandre Barreiro Fidalgo,
Veronica Diesen, Björn Dahlgren, Mats Jansson, Prof. Gábor Merényi,
Prof. Johan Lind, Majid Safdari, Annika Maier, Elin Toijer, Joakim
Halldin Stenlid, Camilla Gustafsson. I am sincerely grateful for your
kindly help and for sharing the wonderful moments with me. For Åsa and
Alex, special thanks for your immense scientific input and invaluable
discussions on our collaborations.
64
7 Acknowledgements
Inna Soroka, Prof. Christofer Leygraf, Kristina Nilsson, Alex Grosjean,
Claudia Gasparrini, my co-authors. Thanks for the scientific input and
invaluable discussions. Special thanks Inna for the kindly help on the
thesis.
Bo Xu, Jing Huang, Hong Chen, Yinan Yang, Lei Wang, Jiajia Gao,
Haining Tian, Qian Gao, Yong Hua, Juan Sun, Ming Cheng, Cheng Chen,
Qingpeng Meng, Wen Zhang, Wangzhong Mu, Lan Hu, Yifeng Yun.
Great thanks for all of you for making my everyday life in Sweden
enjoyable and unforgettable.
China Scholarship Council (CSC) and Swedish Nuclear Fuel and Waste
Management Co (SKB) are gratefully acknowledged for financial support.
My family, both in Yang’s and Zhang’s. Thank you so much for your
endless support and understanding all these years.
My parents, thank you for your endless love and full support.
My wife Xian Zhang, I would not have become a better man and a Ph.D
without your immense encouragement, remarkable consideration and
endless love.
My daughter Ruoyu Alice Yang, you are the most amazing gift in my
life, thank you.
65
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