Visible light-driven photooxidation of water at hybrid photoanodes

V I S I B L E L I G H T- D R I V E N P H O T O O X I D AT I O N O F WAT E R ATH Y B R I D P H O T O A N O D E S

D I S S E RTAT I O N

zur Erlangung des Gradeseines Doktors der Naturwissenschaften

Fakultät für Chemie und BiochemieRuhr-Universität Bochum

vorgelegt vonmichał błedowski

November 2013

Błedowski, Michał – 2013Visible light-driven photooxidation of water at hybrid photoanodes

Die vorliegende Arbeit entstand in der Zeit von November 2010 bisNovember 2013 an der Fakultät für Chemie und Biochemie der Ruhr-Universität Bochum unter Anleitung von Herrn Jun.-Prof. Dr. RadimBeránek.

promotionskommission:Jun.-Prof. Dr. Radim BeránekProf. Dr. Wolfgang SchuhmannProf. Dr.

ort:Bochum

datum:

Nihil est veritatis luce dulcius

— Marcus Tullius Cicero “Academici libri priores”

Nonne optimus et gravissimus quisque confiteturmulta se ignorare et multa sibi etiam atque etiam esse discenda?

— Marcus Tullius Cicero “Tusculanae Disputationes”

A C K N O W L E D G M E N T S

I have learned a lot from many amazing individuals I have met dur-ing last three years. I would like to thank all of them for helping meto develop myself. I hope that I was able to repay in a similar way.

I wish to thank Prof. Dr. Radim Beránek for the supervision of thiswork and many, not only scientific, fruitful discussions. I am partic-ular grateful for his trust in me, his advice, generous support of mywork, and magnanimous help every time I needed.

I would like to acknowledge Prof. Dr. Roland A. Fischer and allmembers of Chair of Inorganic Chemistry II for kind support, officeassistance, and interesting discussions during weekly seminars. I amvery grateful towards Prof. Dr. Wolfgang Schuhmann, all membersof Elektroanalytik & Sensorik Group, especially Mrs. Sandra Schmidtfor performing SEM/EDX measurements, and all members of Centerfor Electrochemical Science for generous support and very good at-mosphere in the lab. I am also highly indebted to Mr. Armin Linderand his colleagues form Feinmechanik Werkstatt for help in designand construction of the photoelectrochemical cells. I would like toacknowledge Mr. Jürgen Rotzing and his colleagues for glass work.I would like to thank the Ruhr University Research Schoolplus forsupporting my participation in MRS Spring Meeting 2013 in San Fran-cisco, USA.

I would like to thank my group colleagues: Lidong, Susan, Petra,Ayyappan, Tan, Ewa, Darek, and Oliver for help, useful advices, manycontributions to my work, and amiable atmosphere. Many thanksalso to my friends outside Institute: Edyta, Kinga, Sławek, Agnieszka,friends from PMK Bochum, flatmate from Roncalli Haus, Mike, andfriends form German classes: Antonio and Thomas.

I am very grateful towards my family: my parents, Stanisław andWiesława, and my sisters, Agata and Dorota, for their lifelong loveand encouragement.

Finally, I would like to thank Arleta for sharing all those wonderfulmoments during last three years.

v

C O N T E N T S

i state of the art 11 introduction 32 literature 11

2.1 Photoelectrochemical Cells 112.1.1 Configurations and Efficiencies 112.1.2 Requirements and Trade-offs 16

2.2 Photoanodes Composition 202.2.1 Absorbers 222.2.2 Co-catalysts 22

2.3 Water Oxidation 252.3.1 Water Oxidation in Nature 252.3.2 Mechanism 282.3.3 Trends in Activity 292.3.4 Scaling Relations 29

2.4 Photoanode Improvement Strategies 312.4.1 Nanostructures & Morphology 322.4.2 External Bias 342.4.3 Semiconductor/Co-catalyst Junctions 362.4.4 Semiconductor/Semiconductor Heterojunctions 362.4.5 Sensitization 37

ii experimental details 393 experimental details 41

3.1 Instrumental 413.1.1 Photoelectrochemical Measurements 413.1.2 Dissolved Oxygen Concentration Measurements 42

3.2 Characterization 473.3 Synthesis 47

iii results and discussion : hybrid photoanodes loadedwith different co-catalysts 49

4 iridium-based co-catalysts 514.1 Introduction 514.2 Colloidal Deposition of Iridium Oxide 52

4.2.1 Experimental 524.2.2 Results and Discussion 53

4.3 Photodeposition of Iridium Oxide 554.3.1 Experimental 554.3.2 Results and Discussion 56

4.4 Conclusion 655 cobalt-based co-catalysts 67

5.1 Introduction 67

vii

viii contents

5.2 Photodeposition of Cobalt Phosphate 705.2.1 Experimental 705.2.2 Results and Discussion 70

5.3 Mixtures of Metal Oxides 755.3.1 Cobalt(II,III) Oxide (Co3O4) 755.3.2 Nickel Cobaltite (NiCo2O4) 815.3.3 Cobalt Titanate (CoTiO3) 86

5.4 Conclusion 946 nickel-based co-catalysts 95

6.1 Introduction 956.2 Experimental 966.3 Photodeposition of Nickel Borate 966.4 Conclusion 101

7 manganese-based co-catalysts 1037.1 Introduction 1037.2 Photodeposition of Manganese Oxide Clusters 105

7.2.1 Experimental 1057.2.2 Results and Discussion 105

7.3 Mixtures of Metal Oxides 1077.3.1 Experimental 1077.3.2 Results and Discussion 109

7.4 Conclusion 1098 effect of external bias 111

8.1 Introduction 1118.2 Experimental 1118.3 Results and Discussion 1118.4 Conclusion 118

iv summary 1199 summary 121

references 127

S Y M B O L S

b Tafel slope

Bi borate buffer

e− electron

Eappl external potential (bias) applied

Eg bandgap energy

ERHE potential versus RHE

F Faraday constant

h Planck constant

h+ hole

iph photocurrent density

j current density

j0 exchange current density

jSC short-circuit photocurrent density

kB Boltzmann constant

kHT hole transfer rate constant

krec recombination rate constant

L mean free diffusion length

P light power density

Ptotal illumination power density

Pi phosphate buffer

q elementary charge

S illuminated electrode area

T temperature

t thickness

∗EFn quasi-Fermi level of electrons

∗EFp quasi-Fermi level of holes

ix

η overpotential

hν energy of light

∆G Gibbs free energy

λ wavelength

µ mobility of carrier

ΦH2 rate of hydrogen production

φcatal. catalytic efficiency of water oxidation at the surface

φsepar. charge separation efficiency

τ charge carrier life-time

A C R O N Y M S

ABPE Applied Bias Photon-to-current Efficiency

AM1.5 Air Mass Coefficient (1.5 atmosphere thickness)

BET Brunauer-Emmett-Teller

CB Conduction Band

CD Colloidal Deposition Method

CE Counter Electrode

DFT Density Functional Theory

DL Detection Limit

DO Dissolved Oxygen

DSSC Dye-Sensitized Solar Cell

EDX Energy Dispersive X-ray Spectroscopy

ET Electron Transfer

FE Faradaic Efficiency

FT-IR Fourier Transform Infrared Spectroscopy

FTO Fluorine Tin Oxide

GCN Graphitic Carbon Nitride

HER Hydrogen Evolution Reaction

IPCE Incident Photon to Current Efficiency

x

acronyms xi

ITO Indium Tin Oxide

NHE Normal Hydrogen Electrode

NR Nanorods

BR-NR Branched Nanorods

NT Nanotubes

OER Oxygen Evolution Reaction

XPS X-Ray Photoelectron Spectroscopy

OEC Oxygen Evolving Center (or Catalyst)

PCET Proton-Coupled Electron Transfer

PD Photodeposition Method

PEC Photoelectrochemical

PH Polyheptazine

PSII Photosystem II

SHE Standard Hydrogen Electrode

RHE Reversible Hydrogen Electrode

RE Reference Electrode

STH Solar-to-Hydrogen Efficiency

TOF Turnover Frequency

VB Valence Band

WE Working Electrode

XAS X-Ray Absorption Spectroscopy

XPS X-Ray Photoelectron Spectroscopy

XRD X-Ray Diffraction

Part I

S TAT E O F T H E A RT

1I N T R O D U C T I O N

One of the most fundamental issue that humankind has been strug-gling for centuries is effective conversion and storage of useful formsof energy. Since prehistorical times when biomass was powering hu-man and domestic animals’ muscles, through the age of water andwind powering mills and ships, moving to the age of coal power-ing steam engines, humans have been trying to control natural en-ergy flows, aiming to use them to improve the quality of life. Thereis a obvious relationship, manifested in history, between availabilityof energy and shape of human condition, which implies that futuredevelopment of civilization strongly depends on provision of new en-ergy sources and new efficient methods of energy conversion, storage,and utilization. [1]

Nowadays society is based almost utterly on fossil fuels: coal, nat-ural gas, and petroleum. The total energy consumption rate in 2011was equal to 16.3 TW; 87 % of this value was obtained by burning fos-sil fuels which are non-renewable, non-uniformly distributed in theearth crust, and have a high environmental impact. In 1990 the en-ergy consumption rate was equal to 10.8 TW; in 2030 it is estimatedto reach 22.2 TW, with the highest contribution expected from thenon-OCED growing countries where clean technologies are underde-veloped. [2,3] Therefore, in the new global economy, energy problemhas become a central challenge that humanity has to face in a shortperiod of time.

The sun offers the best solution by the continuous deliverance of120 000 TW of power to the surface of the earth, an amount going farbeyond (by four orders of magnitude) our current needs. At the sametime the sun energy is extensively available, equally distributed, in-exhaustible, and completely sustainable. The past decades have seenthe rapid development of solar energy conversion systems, especiallysolar-to-thermal and solar-to-electricity ones. However, solar electric-ity accounts presently only for 0.41 % of global electricity produc-tion, [4] essentially because of variable price of Si in standard solarcells and still low efficiency of dye sensitized solar cells. [5] Further-more, energy of the sun impinging on the earth’s surface has diffusenature, fluctuates according to season and weather conditions, andcan be captured only during day-time. Consequently, due to the lackof capable electricity storage systems, especially if automotive appli-cations are considered, relying solely on generation of solar electricitydoes not represent a viable strategy.

3

4 introduction

On this account, in recent years, there has been an increasing in-terest in conversion of solar energy into chemical energy, e.g. highly-energetic hydrogen bonds1 whose energy density per mass unit isapprox. 100 times higher than in batteries and approx. 3 times higherthan in standard liquid fuels. [7] One of the most appealing strategiesto do this, is the use of sunlight-driven devices, in which solar energyis harvested, converted, and directly stored by splitting water intohydrogen and oxygen. [7–9] The overall reaction taking place in suchdevice:

2 H2O hv−→PC

2 H2 + O2 (1.1)

has a4G= + 237 kJ mol-1, indicating that, in contrast to many otherphotocatalytic reaction, e.g. degradation of pollutants, water splittingis thermodynamically uphill.

In a search for a photocatalytic system capable of producing hy-drogen from water, apart from the intensively investigated homoge-neous molecular approach, [10–15] significant attention is attracted bythe “heterogeneous” approach, in which three main strategies can bedistinguished: a) photovoltaic cell combined with an electrolyzer, b)powders photocatalysis, and c) photoelectrochemical cells (Figure 1.1).In the latter strategy, either one (single photoelectrode systems) orboth (tandem configuration) of the water-splitting reactions occur di-rectly at a light-absorbing photoelectrode in contact with aqueouselectrolyte (Figure 1.1 c). [16–18] Typically such photoelectrodes em-ploy semiconductor-based architectures since these enable efficientseparation of photoinduced charges – electrons and holes – that aresubsequently used to drive the water-splitting reactions. Importantly,due to their inherent constructional and functional simplicity, photo-electrochemical cells are potentially more efficient and cheaper thana conventional electrolyzer driven by a photovoltaic cell array (Fig-ure 1.1 a). [19,20]

A substantial simplification of a water-splitting system is possibleif simple dispersion of semiconductor particles instead of photoelec-trodes is used (Figure 1.1 b). However, in such system, separationof produced H2 and O2 is required (in contrast to a photoelectro-chemical cell, in which reduction and oxidation occur in two separatecompartments), and even more importantly, energy conversion effi-ciencies, achieved so far, are rather low, i. e. below 1 %. [21] Neverthe-less, photocatalytic systems using semiconductor powders have beenwidely investigated, [21–24] with remarkable results, e.g. visible-lightdriven overall water spiting using gallium nitride (GaN) and zincoxide (ZnO) solid solution (Ga1-xZnx)(N1-xOx)0.1 > x > 0.4 loaded with

1 Interestingly, the idea of using hydrogen as an energy carrier follows the logic of his-toric development – from hydrogen poor to hydrogen rich, and from solid, throughliquid, to gaseous fuels: [6]

C (coal)→ –CH2– (oil)→ CH4 (natural gas) ?→ H2 (hydrogen)

introduction 5

Figure 1.1: Three main strategies in “heterogeneous” water splitting: a) elec-trolyzer combined with photovoltaic cell, b) semiconductor pow-ders photocatalysis, c) photoelectrochemical cell.

metal and metal oxide islands acting as catalytic sites for both waterreduction and oxidation. [25–31]

The overall solar-to-hydrogen efficiency threshold required for com-mercial implementation of photoelectrochemical cells (Figure 1.1 c) isgenerally considered to be about 10 %. [19] Notably, the feasibility ofachieving even higher efficiencies in photoelectrochemical water split-ting has been demonstrated repeatedly. In 1981 Heller reported 12 %efficiency for an electrically biased p-InP single-crystal photocathodecovered by a thin ruthenium layer. [32] A similar efficiency of 12.1 %for a p-InP epitaxial thin film covered with rhodium nanoparticles

6 introduction

has been reported recently by Lewerenz et al. [33] Turner et al. have re-ported a photoelectrochemical cell with photovoltaic self-bias basedon a p-GaInP2/GaAs multilayer structure showing the efficiency of12.4 %, [34] and an overall efficiency as high as 18.3 % for a multijunc-tion photoelectrochemical cell based on AlGaAs and silicon coupledto RuO2 and Pt black catalysts has been demonstrated by Licht, Trib-utsch et al. in 2001. [35] However, all these high-efficiency cells arebased on very expensive materials and cannot be easily scaled up. Thesuccess of the photoelectrochemical water splitting approach there-fore presupposes the development of novel photoelectrodes based oncheap and abundant materials.

In this context it is also important to realize that it is particularly thewater oxidation reaction that represents a major challenge in photo-electrochemical water splitting. This is because the dioxygen-evolvingreaction is a very complex process requiring a proton-coupled trans-fer of four electrons from two water molecules. Consequently, the ki-netics of oxygen evolution on most inorganic surfaces is slow and as-sociated with considerable overpotential. [37–40] Furthermore, the rela-tively stable metal oxides typically utilized in water-splitting photoan-odes either do not absorb visible light (TiO2), or – if they do absorbin the visible (e.g. α-Fe2O3, WO3) light – they suffer from too posi-tive conduction band (CB) edges (Figure 1.2 b), which translates intoa requirement of a significant external electric bias application or theneed or the need for photocells with tandem configuration to allowfor water reduction at the counter electrode (Figure 2.1 b). [41–43]

Recently, our research group has been developing photoanodesbased on a novel class of visible-light photoactive inorganic/organichybrid materials2 – nanocrystalline TiO2 modified at the surface bya thin (< 1–5 nm) layer of polyheptazine (TiO2-PH) (Figure 1.2 a). [44–49]

Importantly, such photoelectrodes exhibit very different optical andphotoelectrochemical properties from those of both TiO2 and pristinepolyheptazine. The optical absorption edge of the TiO2-PH hybridis strongly red-shifted into the visible (2.3 eV; ∼ 540 nm) as com-pared with the bandgaps of both of the single components, TiO2

(∼ 3.2 eV; ∼ 390 nm) and polyheptazine (∼ 2.9 eV; ∼ 428 nm), whichis due to the formation of an interfacial charge-transfer complex be-tween polyheptazine (donor) and TiO2 (acceptor) (Figure 1.2 b andFigure 1.3). [44] Two further features of these architectures are particu-larly noteworthy. First, polyheptazine (also known as “graphitic car-bon nitride”) [50,51] is a highly robust compound exhibiting thermalstability up to 550 °C in air. [44] Second, the visible light-driven opticalcharge-transfer excitation in the TiO2-PH hybrid leads to photogener-ation of electrons with a relatively negative potential of the conduc-tion band edge of TiO2 [– 0.2 V versus reversible hydrogen electrode

2 Some more detailed investigations of this material class are the topic of the forth-coming dissertation by Lidong Wang

introduction 7

Figure 1.2: a) Schematic view of photoanode consist of TiO2 nanoparticles(gray) modified with polyheptazine (orange – structure proposedin lit. [36] in circle). Energy levels for b) TiO2, WO3, Fe2O3 and c)TiO2 modified with polyheptazine; direct electron transfer formHOMO of polyheptazine to CB of TiO2 in red.

(RHE)], as compared with other mostly studied photoanode materi-als like monoclinic WO3 or α-Fe2O3 (Figure 1.2 b).

Above-listed properties make TiO2-PH hybrids very promising interms of developing new photoanodes at which water splitting canbe possibly carried out without any externally applied bias. However,

8 introduction

Figure 1.3: Diffuse reflectance spectrum (Kubelka–Munk function versuswavelength) of the TiO2-PH hybrid compared to that of anataseTiO2; the inset shows bandgap determination using plots of mod-ified Kubelka–Munk function versus energy (formalism of indi-rect optical transition is assumed).

we have found that complete photooxidation of water to dioxygencannot be experimentally observed using simply TiO2-PH photoan-ode. [44] This changes first after loading of TiO2-PH with a so-called co-catalyst, that is additional metal or metal oxide catalytic sites whichserve as oxygen evolving centers. [44] The main aim of the presentthesis is preparation of TiO2-PH hybrid photoanodes loaded with dif-ferent co-catalysts for water oxidation and their photoelectrochemicalcharacterization, with a focus on determination of material activity byevolved dioxygen measurements. The key questions addressed in thisthesis are:

• Is a TiO2-PH hybrid photoanode loaded with a co-catalyst ableto photooxidize water? Which type of co-catalysts can efficientlymediate the hole transfer process from a TiO2-PH photoelec-trode to water molecule?

• How the photooxidation of water depends on a method, bywhich the co-catalyst was introduced to the light-absorbing pho-toanode material?

• What is the influence of conditions, e.g. external bias, pH, typeof electrolyte, on the activity of TiO2-PH photoanode loadedwith a co-catalyst in water oxidation reaction?

The present thesis has been organized in following way. The secondChapter will be introductory and will bring basic information about

introduction 9

photoelectrochemical cells, composition of photoanodes, and chem-istry of water photooxidation reaction. In Chapter 3 the general exper-imental methods are described, including design of a completely newexperimental setup used for photogenerated dioxygen measurements.In Chapters 4, 5 ,6 and 7, TiO2-PH hybrid photoanodes with iridium-based, cobalt-based, nickel-based and manganese-based co-catalystsare examined, respectively. Every of Chapters 4–7 begin with a briefintroduction about the particular group of co-catalysts. Afterwards,the photoanodes preparation methods are elaborated, followed byphotoelectrochemical characterization results, discussion, and conclu-sions. In Chapter 8 the effect of external bias is discussed in details.Finally, some concluding remarks are made in Chapter 9.

Some ideas and figures presented in this thesis have appeared pre-viously in the following publications:

1. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek, J. Mater.Res. 2013, 28, 411.

2. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek, Mater.Res. Soc. Symp. Proc. 2012, 1442, Q02.

3. M. Bledowski, L. Wang, A. Ramakrishnan, A. Bétard, O. V. Khavry-uchenko, R. Beranek, ChemPhysChem 2012, 13, 3018.

4. L. Wang, M. Bledowski, A. Ramakrishnan, D. König, A. Ludwig,R. Beranek, J. Electrochem. Soc. 2012, 159 (7), H616.

2L I T E R AT U R E

2.1 photoelectrochemical cells

2.1.1 Configurations and Efficiencies

The first experimental realization of water splitting in a photoelec-trochemical cell was undertaken by Honda and Fujishima [52] whoused crystalline TiO2 (rutile) as a photoanode and platinum wire asa counter electrode (cathode). Illumination of TiO2 with UV light ledto generation of O2 and H2. Half reactions occurring on photoanode(OER) and cathode (HER) under acidic conditions (pH = 0):

OER : 2 H2O + 4 h+ O2 + 4 H+ Eored = +1.23 V vs. NHE (2.1)

HER : 4 H+ + 4 e− 2 H2 Eored = +0.00 V vs. NHE (2.2)

give the overall reaction like in Equation 1.1. As mentioned before,a free energy change 4G = + 273 kJ mol-1 of this reaction correspondsto 4E = – 1.23 V per electron transfered, according to Nernst equa-tion:

4G = −nF4E (2.3)

where n - number of electrons transferred, F - Faraday constant.Energy diagram for a photoelectrochemical cell with a n-type semi-

conducting, photoactive anode and a metal cathode (e.g. Pt) is pre-sented in Figure 2.1 a). In this system (single photoelectrode config-uration) optimal size of photoanode’s bandgap was estimated to ca.620 nm, that is ca. 2.0 eV (see Section 2.1.2 for explanation). The max-imum theoretical energy-conversion efficiencies for different wave-lengths possible to achieve in a system with single absorber wascalculated by Shockley and Queisser. [53] For a device with the sin-gle photoelectrode configuration having the optimal bandgap size(620 nm) the so-called Shockley-Queisser limit restricts the theoreti-cal maximum efficiency to ∼ 23 % . [53] However, the maximum realiz-able efficiency of solar photolysis of water, assuming presence of oneabsorber with the optimal bandgap of 620 nm, was estimated to bearound 17 %. [18,54–56]

The single photoelectrode configuration of photoelectrochemicalcell (Figure 2.1 a) will be also investigated in the present thesis. Semi-conductors which can be used as a photoanode material are detailedin Section 2.2.1.

11

12 literature

Figure 2.1: Examples of energy diagrams for PEC cells configurations un-der illumination: a) single bandgap cell with n- type photoanodeand metal cathode; b) dual bandgap (tandem) cell with n-typephotoanode and p-type photocathode; monolithic device with n-type photoanode biased with integrated: c) standard PV cell ord) dye-sensitized solar cell. Adapted from: van de Krol [57]

2.1 photoelectrochemical cells 13

Conversely to the PEC cell using a n-type photoanode and a metalcathode, configuration with a p-type semiconductor as a photocath-ode and a metal anode is possible. Considerable amount of studiesconcentrate on p-type silicon photoelectrodes, [58,59] often loaded withhydrogen evolving co-catalysts, e.g.: Pt, Ni, or Ni-Mo which signifi-cantly improve photoelectrocatalytic activity. [60] Another promisingmaterial which can be used as a photocathode in PEC cell is copper(I)oxide (Cu2O). The main obstacle limiting its application in photo-cathodes is susceptibility to photocorrosion. Therefore, the researchfocused on formation of TiO2

[61,62] or CuO [63] protective overlayers,and RuO2 hydrogen evolving co-catalyst in order to improve bothstability and catalytic activity. [61]

In Figure 2.1 b) energy diagram for a dual bandgap (tandem) sys-tem which comprise two single bandgap semiconductors in one de-vice is shown. In this approach two photoelectrodes are employed: ann-type semiconductor under illumination acts as a photoanode anddrives water oxidation, a p-type semiconductor under illuminationacts as a photocathode and simultaneously drives water reduction.Both minority carriers: electrons and holes are used, therefore it isnecessary to select semiconductors with appropriate band edge posi-tions – the valence band of the photoanode should be below and theconduction band of the photocathode above oxidation and reductionpotentials of water, respectively.

The dual bandgap configuration can be beneficial, because the com-bination of various semiconductors with smaller bandgaps can pro-vide higher absorption in the visible light range, where the sun hasgreater photon flux (Figure 2.3). Accordingly, for the system involv-ing two bandgaps and absorption of four photons to generate onemolecule of H2, the maximum solar-to-hydrogen (STH) efficiency canbe up to 31 %. [64] Moreover, the dual bandgap configuration can helpto overcome the problem already highlighted in the Introduction:many of n-type semiconductors have conduction band too positiveto reduce protons without application of external bias.

An example of experimental realization of dual bandgap configu-ration is a system consisting of a nanocrystalline n-TiO2 photoanodeand a single crystal p-SiC photocathode which exhibits current den-sity of 0.05 mA cm-2 and STH efficiency of 0.06 % without applicationof external bias. [65] In other work, a photoanode consisting of n-GaAscoated with MnO2 protective layer and a photocathode consisting ofp-InP with a Pt co-catalyst were demonstrated to generate photocur-rent as high as 5 mA cm-2 and STH efficiency of 8.2 %, however insta-bility and high cost of In and Pt restrict a large-scale utilization of thiscell. [66] An another example of tandem cell is the n-/p- type Fe2O3

photoanode/photocathode system, exhibiting unassisted water split-ting with STH efficiency of 0.11 %, limited by the low charge carriertransport rate in p-Fe2O3. [67]

14 literature

Figure 2.2: A schematic energy diagrams of PEC “back-to-back”, wirelessconfigurations: a) photoanode connected through ohmic contactto metal cathode and b) pn-Si junctions coated with metal anodeand cathode, forming thus a “buried” junction. Adapted from:Walter et al. [20]

An alternative way to surmount the problem with not sufficientlynegative positions of CB in the case of n-type semiconductors em-ploys an additional PV cell to provide the extra bias voltage neces-sary to drive reduction at the counter electrode. Figure 2.1 c) depictsthe energy scheme of the PEC cell composed of a photoanode withintegrated PV cell and a metal counter electrode. In this configurationa standard silicon p-n junction solar cell is covered with the photoan-ode material which is in a direct contact with an electrolyte. Suchmonolithically stacked architecture based on semiconductor-liquidjunctions with underlying solid-state junctions were reported in lit-erature. For example, the WO3 photoanode combined with an amor-phous silicon PV cell shows STH efficiency of 3 %. [68] Other studyreported recently that the W-doped BiVO4 photoanode with two un-


derlying amorphous silicon PV cells shows STH efficiency of 4.9 %. [69]

An another example is the compact, crystalline TiO2 photoanode bi-ased by a PV cell which exhibits high photocurrents of 9 mA cm-2

and a considerable H2 generation rate of 37 dm3 h-1 m-1. [70] Insteadof a standard silicon p-n junction, a dye-sensitized solar cell can beused (Figure 2.1 d). [71] For instance, water splitting with 3 % STH ef-ficiency was recently demonstrated in a tandem device consisting ofa WO3 photoanode with an underlying DSSC. [72]

All PEC cells configurations presented in Figure 2.1 comprise a sep-arate anode and a cathode immersed in an electrolyte – often in twocompartments divided by a porous separator. However, the back-to-back “wireless” design is also possible and has a potential to be a ba-sis of low cost, highly manufacturable, and efficient devices, evendespite significant losses due to reflected or scattered photons on theinterfaces between multiple semiconductor layers. In Figure 2.2 en-ergy levels for two examples of such devices are presented. First ex-ample (a) shows the back-to-back combination of a photoanode anda metal cathode connected via ohmic contact. Second example (b)presents two p-n junctions connected in series coated with a metal an-ode and metal cathode, forming thus a “buried” junction. They haveno direct contact with an electrolyte, hence there is no semiconductor-electrolyte junction. [73] Consequently, this configuration does not rep-resent a true PEC approach, but can be rather classified as a PVcombined with an electrolyzer (Figure 1.1 a). Nevertheless, it shouldbe noted that significant results were achieved using the triple junc-tion amorphous silicon photovoltaic cell interfaced with cobalt-basedoxygen evolving catalyst (Co-Pi, anode) and ternary metal alloy (Ni-MoZn) acting as a hydrogen evolving catalyst (cathode). STH effi-ciency equal to 2.5 % was demonstrated for this water-splitting cellcomprising earth-abundant elements and operating in near-neutral(∼ 9) pH conditions. [74,75]

Solar-to-Hydrogen Efficiency (STH) efficiency frequently cited aboveto evaluate system performance is the most important of all effi-ciency measurements in water splitting research. Therefore, it needsexact definition complemented by description of experimental meth-ods which should be used for its determination. [76]

STH efficiency is defined as “chemical energy output” divided by“solar energy input”, what is represented by the following equation:

STH =

[ΦH2(mmol H2 s−1) × G0

f ,H2(kJ mol−1)

Ptotal (mW cm−2) × S (cm2)

]AM 1.5G

(2.4)

where in the numerator the rate of hydrogen production ΦH2 ismultiplied by the Gibbs free energy of dihydrogen formation4G0

f ,H2=

+ 237 kJ mol-1 and in the denominator the illumination power density

16 literature

Ptotal is multiplied by the illuminated electrode area S. Importantly,the light source should closely match the AM 1.5G spectrum in the in-tensity and spectral distribution. Experiments should be performedin a two electrode setup (short-circuit) under zero bias conditions.If WE and CE are compartmentalized, both electrodes should be im-mersed in the electrolyte with same pH to avoid chemical bias arisingbetween two solutions with different pH. Additionally, the electrolytecannot contain any sacrificial donors or acceptors (e.g. MeOH, H2O2,S2O8

2-), because in this case the reaction would not represent truewater splitting.

The rate of hydrogen evolution should be measured directly by an-alytical methods in order to confirm that it occurs stoichiometrically(H2 : O2 = 2 : 1). If the faradaic efficiency of H2 evolution is known,STH can be calculated using the following equation:

STH =

[jSC (mA cm−2) × E0

H2O/O2(V) × FE

Ptotal (mW cm−2)

]AM 1.5G

(2.5)

where jSC is the short-circuit photocurrent density (photocurrentdivided by the illuminated electrode area), E0

H2O/O2is the water split-

ting potential equal to 1.23 V, Ptotal is the illumination power den-sity, and FE is the faradaic efficiency for H2 evolution, that is theratio calculated on a basis of charge passed during experiment andtrue amount of hydrogen directly quantified by analytical methods.Faradaic efficiency FE = 100 % means that the measured photocur-rent corresponds directly to the molar H2 generation.

If additional bias is applied between WE and CE, conditions forthe STH efficiency determination are not any more fulfilled becauseof the additional electric energy input. Therefore, another useful ef-ficiency, namely applied bias photon-to-current efficiency can be de-fined as:

ABPE =

[jSC (mA cm−2) × (E0

H2O/O2− Eappl)(V)

Ptotal (mW cm−2)

]AM 1.5G

(2.6)

where jSC is the short-circuit photocurrent density obtained underan applied bias Eappl. ABPE measurements can easily give insight intothe functionality and limitations of the device, but high values ofABPE might not directly translate into high values of overall STHefficiency which most accurately characterize PEC cell efficiency, es-pecially in cases where the faradaic efficiency is not known. [76]

2.1.2 Requirements and Trade-offs

The most important issue in design of a photoelectrochemical cell,as already has been pointed out in the Introduction, is a choice of


suitable photoanode material. The following requirements have to befulfilled by a material candidate for application as a photoanode ina solar water splitting system: [42]

1. The light absorption edge of the material must be low enoughto absorb a significant portion of visible light (λ > 400 nm). Opti-mal size of bandgap amounts to Eg = ∼ 2 eV (∼ 620 nm) assum-ing one-bandgap absorber configuration. It comprises 1.23 eVenergy required to split water plus some thermodynamic andkinetic loses estimated to ∼ 0.8 eV. This value (2.0 eV) still en-compasses broad range of the solar spectrum which can be ab-sorbed (Figure 2.3).

Figure 2.3: Solar spectrum reaching the earth surface; shaded area repre-sents photon flux that can be absorbed by the photoanode withthe optimal bandgap of Eg = ∼ 2 eV (∼ 620 nm).

2. All components must be stable. This requirement implies thatboth the light-absorbing semiconductor (absorber) and the co-catalyst have to be robust and resistant to dissolution, decom-position, photocorrosion, or other corrosion reactions.

3. The positions of the quasi-Fermi levels of holes (*EFp) and elec-trons (*EFn), must be positive and negative enough to allow forwater oxidation and reduction, respectively, at best without anyexternal bias. In other words, the electrochemical potentials ofphotogenerated holes *EFn and electrons *EFn have to straddlepotentials of water splitting half reactions (Figure 2.4).

4. The charge transport must be efficient. After an electron-holegeneration, the charge carriers have to reach the semiconduc-tor/electrolyte interface and the back contact. In the absence of

18 literature

Figure 2.4: Schematic representation of energy levels and processes takingplace upon irradiation of photoanode: À generation of e-/h+ pair,Á recombination, Â charge transport within semiconductor parti-cle and Ã to back contact, Ä hole in reactive surface state and co-catalyst-mediated transfer to water molecule. Optimal bandgap:Eg= ∼ 2 eV (∼ 620 nm). *EFp is situated below water oxidationand and*EFn above reduction potential of water (in red).

an external field, ability of the charge carrier to travel a certaindistance is described by the mean free diffusion length L:

L =

√f kBT

qµτ (2.7)

where: f - dimensionality factor, kB - Boltzmann constant, q -charge, T - temperature, µ - mobility of carrier, τ - life-time ofcarrier. In case of the n-type nanostructured semiconductors L isusually higher than particle size, thus overall limiting factor isthe electron transport from an absorber into the external circuit.

5. The multielectron-transfer process required for water oxidationmust be catalyzed efficiently. A suitable co-catalysts (good elec-trocatalyst) must be available to mediate very complex holetransfer to water molecules (see Section 2.2.2).

6. Only low-cost materials can be used. A large-scale applicationimpose utilization of inexpensive and easily manufacturablematerials based only on earth-abundant elements, like rock form-ing elements (green) or major industrial metals (red) (Figure 2.5).


Figure 2.5: Abundance (atom fraction) of the elements in Earth’s upper con-tinental crust. Taken from: Haxel et al. [77]

Satisfying all these requirements is not an easy task, as exemplifiedby only several relatively stable metal oxides which have been stud-ied hitherto (Section 2.2.1). It is also because some of the above-listedrequirements stay in conflict. The following trade-offs can be speci-fied: [57]

• position of band edges vs. stability;

• bandgap vs. stability;

• recombination vs. catalysis;

• absorption vs. charge transport;

• catalysis vs. absorption;

• performance vs. cost.

First two trade-offs relate the position of band edges with stability.The wide-bandgap metal oxide (e.g. TiO2, SrTiO3) are relatively stable,but size of the bandgap is much higher than optimal 2.0 eV, imply-ing low absorption of the sunlight. Moreover, CB for these semicon-ductors is often situated below the potential of water reduction.1 Onthe contrary, non-oxide semiconductors (e.g. CdS, GaP) have smallerbandgaps and positions of CB at more negative potentials, but un-dergo anodic decomposition reactions or form thin layers of oxideson the surface, which hinder both electron and hole transfer acrossthe semiconductor/electrolyte interface. Thus, the general tendencycan be formulated: the wider bandgap entails higher stability, but

1 Both TiO2 (anatase) and SrTiO3 are stable and they have position of CB above thepotential of water reduction, but bandgap of E = 3.2 eV limits their application to UVlight.

20 literature

limits absorption of the solar light; higher position of CB and smallerbandgap allow for reduction of water without external bias under thesolar light illumination, but at the same time lead to lower stability. [57]

Next trade-off highlights the possibility of catalytically active speciesintroduced to a semiconductor to act also as recombination centers.Fourth trade-off is related to absorption of photons by compact semi-conductor layer in different distances from the surface. For semicon-ductors with a low absorption coefficient (ε), the light penetratesdeeper regions and creates e–/h+ pairs far away from the surface.Such e–/h+ pairs, generated far from the surface, recombine beforereaching the semiconductor/electrolyte interface. Fifth trade-off re-flects the possible tendency of the catalytically active centers (co-catalyst) introduced to semiconductor in order to improve water oxi-dation reaction kinetics to block light absorption (see Section 2.2.2). [57]

Ultimately, since most of the state-of-the-art devices comprise ma-terials based on expensive elements and complicated fabrication met-hods, the last trade-off emphasizes the need for novel and very effi-cient, but at same time low-cost, materials and methods for successfulwater-splitting device assembly.

2.2 photoanodes composition

A considerable amount of scientific effort has been devoted world-wide to invent a water splitting material which fulfill all requirementslisted in the previous Section. However, despite decades of researchand more than 130 metal oxides identified to be able to evolve dioxy-gen from water, efficient and stable photoanode material for visible-light driven water oxidation remains undiscovered. [78]

Among many photoanode design strategies, one – particularly at-tractive – has become common in the field: combination of a lightabsorber with electrocatalytic units, so-called co-catalyst, depositedon its surface (Figure 2.6). [20,78] First component is responsible forharvesting the incoming light and charge separation. Second one cat-alyzes four proton-coupled electron transfer events necessary to oxi-dize two water molecules, which have been identified already in theIntroduction as the major obstacle in overall water splitting.

In many cases presence of a co-catalyst is indispensable for wa-ter oxidation reaction even to occur. For the bare semiconductorswhich are able to oxidize water in the pristine form, the depositionof a co-catalyst significantly improves reaction rate by lowering theactivation barrier (overpotential) and also by assisting in the e–/h+

separation. [23,78–80] The latter effect can be explained by the faster ex-traction of charges from the semiconductor, as well as the faster in-jection to the electrolyte of the photogenerated holes which decreasestheir accumulation and recombination with electrons. For example,hole transfer from TiO2 (absorber) to IrO2 (co-catalyst), observed by

2.2 photoanodes composition 21

Figure 2.6: Schematic view of photoelectrochemical cell with a photoanodeconsists of an absorber and a co-catalyst. After charge separa-tion in the absorber, photogenerated electrons are transferred tothe back contact and further through the external circuit to drivewater reduction on the counter electrode, while photogeneratedholes are transferred to the co-catalyst on whose surface the wa-ter oxidation reaction takes place.

transition absorption spectroscopy, occurs with the very high rateconstants of ∝ 10-5 s-1. [81] Such fast transfer already spatially sepa-rates charges, retarding thus recombination. On the other hand, a co-catalyst-semiconductor interface may also act as a recombination cen-ter, but in most of the cases this undesired effect is overcompensatedby improved kinetics.

Typically, the photocurrent arising from photoelectrochemical wa-ter oxidation at the photoanode can be expressed by the equation: [82–84]

JH2O = Jabs × φseparation × φcatalysis (2.8)

where Jabs is the photon absorption rate expressed as current, φseparation

is the charge separation efficiency, and φcatalysis is the catalytic effi-ciency of water oxidation at the surface. Accepting the above-discussedreasons, the presence of a co-catalyst improves both φcatalysis andφseparation. It is apparent, therefore, that the effective coupling betweenco-catalyst and absorber is a key issue in invention of the efficient oxygenevolving photoanode. A better understanding of how charge trans-port across the semiconductor/co-catalyst interface depends on ma-terial’s electronic and structural features may open possibility to itsrational design and manipulation. Consequently, careful semiconduc-tor/co-catalyst interface engineering can prevent accumulation of thephotogenerated holes, improve charge separation, help to match reac-tion kinetics to the incoming solar light flux, and in this way drasti-cally increase the overall device efficiency. [85]

22 literature

2.2.1 Absorbers

An absorber is typically an n-type semiconductor, such that minor-ity carriers (holes) can be driven towards by an in-built gradient ofenergy levels within the space charge layer. Most of the studies con-centrate on metal oxide semiconductors (often nanostructured) dueto the diversity of their optical properties, chemical and thermal sta-bility, good electronic transport properties, easy fabrication methods,and relatively low cost. The following metal oxides have been partic-ularly investigated:

• binary: α-Fe2O3, [86–103] WO3, [104–111] ZnO, [112–115] and TiO2[52,116–122]

often doped with metal (Fe, Cu, Cr, Nb) [123–130] or non-metal (N,C, S, B) [130–136] impurities to enable response in the visible light;

• ternary: SrTiO3[137–139] and BiVO4. [140–144]

Each oxide has its own limitations according to the trade-offs de-scribed in the previous Section, for instance: α-Fe2O3, WO3 and BiVO4

have too positive position of CB, ZnO is unstable under the illumina-tion, BiVO4 and α-Fe2O3 have bad charge transport properties, TiO2,SrTiO3 and ZnO have too large bandgap, etc.

Numerous studies have been also investigating mixed-oxide sys-tems, for example: WO3-TiO2, [145] α-Fe2O3-WO3, [91] WO3-BiVO4, [146]

and many other. [147] A combinatorial approach has been also em-ployed to screen mixed oxides in order to find a composition withthe highest activity. [148–152]

Detailed discussion of advantages and disadvantages, limitations,developments, trends and prospects for water oxidation with semi-conductor metal oxides can be find in the large volume of publishedliterature. [20,41,42,78,79,153,154]

2.2.2 Co-catalysts

As mentioned before, presence of a co-catalyst can improve the kinet-ics of water oxidation by providing a redox pathway with low over-potential. Therefore, it should be a perfect electrocatalyst (in “dark”catalysis) characterized by the low slope (b) and the high exchangecurrent density (j0), the two constants in the Tafel equation:

η = b log(j/j0) (2.9)

which relate: j - observed current density and η - overpotential.It has been known for decades, that noble metal oxides: RuO2 andIrO2 are particularly active towards water oxidation. Besides, differ-ent oxides with spinel (Co3O4, NiCo2O4) or pervoskite (NiLa2O4,LaCoO3) structure have been investigated. [155,156] Importantly, useful-ness of electrocatalysts like RuO2, IrO2, Co3O4, NiCo2O4 and Mn2O3

2.2 photoanodes composition 23

in photochemical processes was confirmed already in 80’s by exper-iments demonstrating dioxygen evolution from colloidal solutionsof these oxides containing photosynthesizer (Ru2+(bipy)3) and sac-rificial electron acceptor (S2O8

2-). [157] Recently, non-stoichiometric ox-ides/hydroxides of the first-row transition metals - CoOx, NiOx, MnOx

attracted significant attention because of considerable activity, lower(in comparison to Ir or Ru) prices and easy preparation methods.

Detailed literature review of materials based on Ir, Co, Ni and Mn,which can act as a co-catalyst in composite photoanode (absorber +co-catalyst) can be find in Introductions to individual Chapters inPart III: Results and Discussion of the present thesis. Here some gen-eral remarks, trends and trade-offs will be discussed.

A composite photoanode in most cases outperforms an absorberalone, however amount of a co-catalyst has to be well-tuned in orderto prevent the blockage of the light absorption. Figure 2.7 a) illustratesthe trade-off between the co-catalyst loading and reactivity. At small

Figure 2.7: a) Trade-off between the efficiency of the photooxidation reac-tion and amount of loaded co-catalyst. Adapted from: Maeda [21]

b) Illustration of incoming light attenuation by the co-catalystlayer; Fsol – solar flux incident on the co-catalyst surface andFcat – solar flux transmitted through the co-catalyst layer (inci-dent on the semiconductor (absorber) surface). Adapted from:Trotochaud et al. [158]

amounts of a loaded co-catalyst, the number of active sites is very low,thus low reactivity (water oxidation reaction rate) is observed. Whenamount of a co-catalyst increases, the photoreaction rate increases aswell, because more active sites for water oxidation are available on thesurface of an absorber, but only to the maximum point, where amountof a co-catalyst is optimal. After this point, an increase of the amountof co-catalyst leads to decrease in photoreactivity associated with re-

24 literature

flection and/or scattering of the incoming light by the co-catalyst. Inother words, excess loading with the co-catalyst blocks incoming pho-tons which cannot reach the absorber surfaces. To avoid hindering ofthe light absorption, the amount of the co-catalyst as well as its distri-bution have to be optimal. Consequently, the co-catalysts in the formof nanoparticles or very thin layers of metal clusters are preferred.

Ideally, a co-catalyst should be transparent for incoming light andat the same time exhibit high catalytic activity. Recently, this issuehave been addressed by a study which proposed an optocatalyticmodel describing coupling between the semiconductor and coloredco-catalyst films. It accounts for the attenuation of light reaching theabsorber surface by the co-catalyst ( Fcat

Finc) and lowering of overpotential

(η) due to the catalytic film thickness (t) increase (Figure 2.7 b). [158] In-situ spectroelectrochemical measurements combined with theoreticalcalculations for electrodeposited IrOx, CoOx, NiOx metal oxides/hy-droxide clusters have shown that rather ultra-thin layers are appro-priate in such systems – the optimal thickness for IrOx is equal to8.8 nm, for CoOx to 2.9 nm, and for NiOx to only 0.57 nm (∼ 3 mono-layers). [158]

It is apparent, therefore, that the co-catalyst preparation as wellas loading method determine photoelectrochemical properties andthereby overall efficiency of the photoanode. Impregnation, colloidaldeposition, and in-situ photo(electro)deposition are methods of choicefor interfacing co-catalysts with light-absorbing materials. High-tem-perature methods may cause decomposition of semiconductor or in-duce side reactions etc., hence are not preferred. Another major draw-back of treatment at elevated temperature is that co-catalysts havehigher crystallinity (and also stability), but lower number of defectswhich act as the active sites. In contrary, defective, amorphous films,prepared without calcination, have more catalytic sites available, there-fore often exhibit increased activity. [20,159,160]

Most studies in the field of OER electrocatalyst have been con-ducted using simple trial-and-error approach. There has been littlediscussion, so far, about structural and electronic factors at molecularlevel, which can help in rational design of OER catalyst. It is not sur-prising if one realizes, that even the mechanism of water oxidation isstill unclear. In the next Section, attempts to establish basic electronicand mechanistic relationships which can guide further research onwater photooxidation will be discussed.

2.3 water oxidation 25

2.3 water oxidation

2.3.1 Water Oxidation in Nature

In recent years there has been an increasing interest in water oxida-tion in Nature, motivated, among other things, by conviction thatbetter understanding of the process occurring during natural pho-tosynthesis in PSII can accelerate progress in artificial photosynthe-sis. [39,161,162] Consequently, research concentrates not only on chem-istry of the oxidation process [163,164] and structure of PSII, [165,166] butalso on the synthesis of bio-mimetic catalysts, [167,168] or comparisonof energy conversion [169,170] and mechanistic [171] pathways in artifi-cial and natural photosynthesis.

Water oxidation is well-known example of reaction involving Proton-Coupled Electron Transfer (PCET), which is a redox process, requiringsimultaneous transfer of proton and electron from different orbitalsin the donor to different orbitals in the acceptor, in one, concertedelementary step or sequential, step-wise electron transfer followed byproton transfer (ET-PT) or the other way round (PT-ET). PCET canbe very often met in essential biological processes. A classic exampleis Kok cycle (biological water oxidation) where in S0 to S1 transitionelectron is transferred from phenolic group of tyrosine (Yz) – aminoacid directly taking part in OER in PSII – to chlorophyll (P680), whileproton is simultaneously transferred to His190 redox cofactor. Accord-ingly, water oxidation should be considered as multi-electron/multi-proton process in which both electron and proton coupled move-ments are essential. The extreme complexity of reaction requiringfour PCET events highlights the need of further fundamental stud-ies that can improve understanding of PCET mechanism, particularlyin context of artificial water oxidation. [172–174]

It is apparent that studying structure, operating principles andmechanism can be beneficial especially if one takes into account supe-rior properties of PSII in comparison to artificial systems. For exam-ple, the overpotential for the water oxidation reaction on rutile TiO2

(110) was estimated to η ~ 0.78 V, [175] whereas for the CaMn4O5 clus-ter in PS II to only η ~ 0.3 V. [176] The Oxygen Evolving Center (OEC)in PSII outperforms most artificial electrocatalysts also with respectto Turnover Frequency (TOF).2 For active cluster in PSII, TOF was es-timated to 100–400 s-1, while for one of the best electrocatalsyts, thatis IrO2 and RuO2 to only 5–40 s-1. [80] In addition, OEC in nature op-erates at mild conditions and it consists of earth abundant elements.On the contrary, IrO2 or RuO2 are very expensive and work best instrong acidic or alkaline pH conditions. For these reasons, it seemsappealing to identify and analyze in more details lessons given byNature, which can be useful in artificial system design. [177,178]

2 Turnover frequency (TOF) = Number of molecules reacted per active site in unit time

26 literature

It is fairly certain that OEC in nature owes its outstanding abil-ity to mediate transition between intermediates to unique structureand higher flexibility. The inorganic core of OEC in PSII, consists offour manganese and one calcium atom connected with µ-oxo bridges(CaMn4O5). It has a very interesting geometry, resembling distorted,asymmetric cuban (Figure 2.8). This structural motif (distorted cuban;µ-oxo bridges) was also found in very active artificial electrocatalysts:CoOX (Co-Pi), NiOx (Ni-Bi) and MnOx (MnCat) (see Section 5.1, Sec-tion 6.1, Section 7.1) rationalizing usefulness of design guidelines pre-sented by Nature.

Figure 2.8: Structure of Oxygen Evolving Center in PSII – CaMn4O5 in-organic core surrounded by amino acids and water molecules.Taken from: Umena et al. [166]

There are four water molecules and several amino acids in directproximity of the inorganic core, whereas hundreds of amino acidsand approximately one thousand of water molecules are present inwhole PSII. Amino acids do not only serve as redox mediators, but to-gether with water molecules govern proton, water, and oxygen move-ments by the extensive network of hydrogen bonds. Importance ofproton management is obvious when necessity of fourfold proton ex-traction from two water molecules in PCET mechanism is kept inmind. However, it is often disregarded in artificial water oxidationsystems design. [177,178]

Another critical issue, not only for water photooxidation, but for allphotocatalytic systems, is charge separation. Because of the absenceof internal electric fields in the form of space charge layers in smallmetal clusters, the charge separation is governed only by kinetic fac-tors, similarly to e.g. nanostructured semiconductors sensitized withdyes. [179] It can be clearly seen in analysis of the sequence and kinet-ics of electron transfers (ETs) between chlorophyll, redox cofactors,and OEC.


Figure 2.9: a) Electron transfer pathway in PSII. Taken from: McEvoy andBrudvig [163] b) Energy diagram of intermediate states after exci-tation of chlorophyll P680. Adapted from: Dau and Zaharieva [176]

In the first step, the red light incidents chlorophyll (P680) resultingin the electron excitation followed by the rapid transfer to the firstredox cofactor, namely pheophytin (Phe). Relaxation of P680 to theground state is very fast – occurs in nanoseconds, but transfer to Pheis even faster – it takes only 3 ps. Afterwards, electron is transferedfrom Phe to plastoquinone QA, followed by the next ET from QA toanother plastoquinon QB. The reduced QB diffuses away and hole inchlorophyll is filed via tyrosine (Yz) by electron extracted from watermolecules in the OEC (Figure 2.9 a). Similarly to the first ET, each sub-sequent ET is much faster (2–3 orders of magnitude) than competingrecombination pathway to the ground state. In other words, there isa series of very rapid (much faster than recombination) ETs over shortdistances, resulting in spatial separation of charges, which convert the

28 literature

short-lived molecular excited state (P*680), into the long-lived charge

separated state (OEC+QB–). The spatial separation through series of

cascade electron transfers increases life-time of charge carriers by fac-tor of ca. 106 (from ns to ms/s). Of course, multiplication of the life-time due to cascade ETs from states with higher, to states with lowerenergy, has an energetic cost which in the discussed case amountsto ca. 0.8 V (Figure 2.9 b). Consequently, it can be concluded that in-creased carrier life-time (decreased recombination) is preferable evenif electrochemical potential has to be sacrificed. [80,176]

2.3.2 Mechanism

The research efforts devoted to unlock water electrooxidation mecha-nism at inorganic catalysts have been reviewed in literature, empha-sizing partially conflicting or incomparable results, due to, for exam-ple, different catalyst preparation methods or electrochemical mea-surements conditions. [39] In order to determine the exact sequence ofelementary steps (and the rate determining one), advanced surface-sensitive tools, like in-situ XPS and XAS or Raman spectroscopy etc.,combined with electrochemical characterization, should be more of-ten employed. [39] An example of such tools is also rapid scan FT-IR spectroscopy which was recently used to identify hydroperoxo-(HOO-) intermediate as one of the surface-adsorbed oxygen speciesduring water oxidation at iridium oxide clusters under visible lightirradiation. [180]

Considerably more is known about water photooxidation mecha-nism, especially at TiO2 surface. It is a four electron/hole process, [40]

most likely with three intermediates. The early mechanistic work pro-posed photooxidation of Ti-OH– as a first step. Resulting Ti-OH• rad-icals were assumed to recombine producing H2O2 molecule which isfurther photooxidize by holes to give dioxygen. [181] This mechanism

Figure 2.10: Mechanism of water photooxidation on TiO2. Adapted from:Nakato et al. [182]


was challenged by Nakato et al., [182] who propose that reaction is ini-tiated by the nucleophilic attack of water molecule on the surfacetrapped hole. In following steps adsorbed hydroxo-, oxo- and peroxo-groups are formed (Figure 2.10).

2.3.3 Trends in Activity

The vast majority of experimental research on electrocatalysts for wa-ter oxidation concentrate only on general properties, but do not eval-uate microscopic factors leading to increased activity of particularmaterials. Nevertheless, a few specifically designed studies have at-tempted to establish relationships which can help in rational designof new electrocatalysts.

In the first example, different instrumental methods combined withDFT calculations were used to investigate amorphous 3d metal hy-dro(oxy)oxides: MOx (M = Ni, Co, Fe, Mn) deposited on the Pt single-crystal in alkaline media. [183] The activity in “dark” electrocatalysisfor these materials was correlated to the molecular-level descriptor,namely M-OHad bond strength (oxophilicity). It was found, that thecatalytic activity decreases in the sequences: NiOx > CoOx > MnOx

> FeOx, according to increasing M-OHad bond strength. Indeed, thedecrease in the activity in such sequence was confirmed by carefulexperimental comparison of NiOx, CoOx, MnOx, and FeOx electrocat-alysts conducted in another study. [184]

In the second example, the relationship between catalytic activityand eg orbital occupancy of transition metal in perovskite oxide hasbeen investigated. [185] The results of this study allowed to create thevolcano-shaped plot with the maximum activity predicted for the eg

filling factor close to unity. The selection of eg orbital occupancy asthe activity descriptor was motivated by its participation in σ bondingbetween transition metal and oxygen-related intermediated species.Drawing on this – derived from on molecular orbital principles – pre-diction, Ba0.5Sr0.5Co0.8Fe0.2O3-x perovskite oxide was synthesized. Itexhibits very high activity, reported to be even higher than one of thebest known electrocatalysts, i.e. IrO2. [185]

These two examples of descriptor-guided investigations indicatepivotal role of the fundamental, atomic scale understanding of sur-faces processes in engineering and optimization of active sites forwater oxidation.

2.3.4 Scaling Relations

Previous Sections provided evidences reported in the literature thatinteractions between oxygen-based intermediates and the catalyst sur-face are of great importance. Thermodynamics of binding of these in-

30 literature

termediates on metal and metal oxides surfaces was investigated usingDensity Functional Theory by Rossmeisl et al. [186,187]

In Figure 2.11 an example of diagram plotting chemisorption en-ergies (∆G) of adsorbed intermediates for four elementary steps ofwater oxidation vs. reaction coordinate is shown. At potentials bellow

Figure 2.11: a) Three adsorbed intermediates during oxidation of water.b) Exemplified plot of Gibbs free chemisorption energies of ad-sorbed intermediates vs. reaction coordinate in underpotential(η < 0), reversible potential (η = 0) and overpotential (η > 0) con-ditions for ideal (red) and real catalyst (blue). Adapted from:Dau et al. [39]

the reversible potential (η < 0; underpotential conditions) all steps areuphill, therefore reaction cannot occur. For the ideal catalyst (red line)all chemisorption energies are equal, whereas for a real catalyst (blueline) there is an imbalance between chemisorption energies. In thisgeneral example, ∆G3 is higher than other chemisorption energies,meaning that peroxo group is bound to the surface stronger thanother oxygen-based intermediates. Because of this imbalance, for thereal catalyst (blue line) at reversible potential (η = 0) there is still theenergetic barrier which hinders the reaction to occur. In the case ofthe ideal catalyst (dotted red line), due to the equality of the all steps,there is straight line indicating that there is no energetic barrier whichcan hinder reaction. Therefore, it can occur already at the reversiblepotential of 1.23 V vs. RHE (η = 0). For the real catalyst, only whenhigher potential than reversible potential is applied (η > 0; overpoten-

2.4 photoanode improvement strategies 31

tial conditions) all steps become downhill (dotted blue line) and thereaction can occur.

In some respect it is an illustration of the old Sabatier principleformulated for reactions with one absorbed intermediate, saying that:„interactions between the catalyst and the substrate should be neither toostrong nor too weak“. The too strong interaction hinder desorption ofintermediate, on the other hand, too weak interaction hinder adsorp-tion of intermediate. [188]

However, in the discussed case of multi-step OER there is anothercritical issue – all chemisorption energies are linked to each otherby scaling relations. [186,187] It means, that changing ∆G3, for exampleby surface engineering, will change also other chemisorption ener-gies: ∆G1, ∆G2 and ∆G4. Consequently, equidistant steps (like in theideal case) are unachievable, ergo: high overpotential on well-defined,flat, and rigid surfaces of metals and metal oxides is unavoidable. Thistheoretical depiction of energetic relationships between intermediateswas originally developed for dark electrocatalysis, but can be appliedto photo-processes as well. [175] The scaling relation between interme-diates, resulting in minimal, unavoidable thermodynamic overpoten-tial was confirmed for many rutile and perovskite oxides. [189] Conse-quently, a key experimental challenge for OER is to find new class ofcatalysts that circumvents or at least minimize undesirable energeticrelationships between intermediates. [190] It can be achieved by utiliza-tion of three dimensional, flexible structures, which can change andadopt optimal geometry itself during the catalytic cycle, just like theenzyme in Photosystem II which drives water oxidation in green leafs.Indeed, it was suggested in literature that “active cluster in enzyme haslarger flexibility of oxidation site and geometry which cannot be matched byflat and rigid metal oxide surfaces”. [191]

Investigations using the same theoretical framework were also con-ducted for graphitic carbon nitride materials functionalized with tran-sition metals. They exhibit scaling relations similar to those of metaloxides. [192] However, two other conclusions formulated in this studyare noteworthy. Firstly, the catalytic activity of transition metals intro-duced to organic structures depends on chemical nature of surround-ing C and N atoms indicating the possibility of chemical modificationof such inorganic/organic structures in order to achieve lower overpo-tentials. Secondly, the best catalytic behavior was found for elementsfrom groups 7–9, especially Rh, Ir and Co.

2.4 photoanode improvement strategies

There is general agreement in literature that one of the main reasonsof all losses and the main limitation in photocatalysis is recombina-tion of charge carriers. It is also the case of metal oxides most studiedin context of photoelectrochemical water splitting: α-Fe2O3, TiO2, and

32 literature

Figure 2.12: Schematic diagram of energy levels of a semiconductor in con-tact with electrolyte. The rate constant of hole transfer to watermolecule (kHT) is usually much lower than the recombinationrate constant (krec). Adapted from Barroso et al. [193]

WO3. Although for these oxides the hole transfer from semiconduc-tor to water molecule occurs with rate constants in range of kHT =101–10-3 s-1, the recombination is much faster – rate constants are inrange of krec = 1012–106 s-1 (Figure 2.12). [80] In other words, hole in-jection from semiconductor’s valence band to water molecule takesfrom milliseconds to seconds, but hole life-time is in only pico- tomicroseconds long. It is apparent, therefore, that enhancement of life-times of charge separated states by 103–109 is needed in order to usephotogenerated charges efficiently. [80]

The present Section will give an account of material strategies whichcan be used to improve photoanodes in terms of charge separation aswell as two other factors determining overall efficiency, i.e. the lightabsorption and the rate of surface reaction.

2.4.1 Nanostructures & Morphology

Size and morphology are fundamental parameters which should betaken into consideration during photoanode design because they of-ten determine the material’s performance. There are many examplesof dramatic improvement of certain material properties due to nanos-tructured nature of the photocatalyst or change in morphology.

The first reason why nanosized photocatalyst can be beneficial isshortened charge carrier collection distance. [20] In case of nanostructredmaterials the distance which charge carrier must travel until it willreach the surface is remarkably reduced, especially in comparison tothick and flat compact films (Figure 2.13 a). According to Equation 2.7,the free diffusion length depends on the carrier mobility, hence short-ened pathway of collection is substantially advantageous in materialswith very low mobility of charges, e.g. α-Fe2O3. Indeed, nanostruc-tured Fe2O3 exhibit two order of magnitude higher photocurrents in


comparison to single-crystal-like α-Fe2O3. Activity enhancement dueto nanostructuring was observed also in case of many co-catalystsand other semiconductors, e.g. WO3 and BiVO4. [79]

Figure 2.13: A comparison of the absorption-related effects in compact andnanostructured photoanodes: a) shorter charge carrier collec-tion distance, b) improved light distribution due to successiveabsorption of the scattered light (“light-trapping”) c) lower ab-sorbed photon flux per unit area. Adapted from: Osterloh [79]

Besides the shorter charge collection pathway, improvements ob-served in nanostructures can be attributed to larger specific surface area,which promotes charge transfer across the semiconductor/electrolyteinterface. It is particularly significant for reactions with very slow ki-netics of the surface reaction, as in the case of water oxidation. Onthe other hand, high surface area can give rise to recombination atdefects at the surface or at boundaries between particles. One moredisadvantage of nanostructures may be slower electron transport be-tween nanoparticles if they are not sufficiently fused together. In thiscase, electrons have to travel from nanoparticle to nanoparticle and toback contact through series of hopping or tunneling events, increas-ing therefore energetic barriers that need to be overcome.

The next reason why nanostructures can be rewarding is improvedlight distribution. Nanoscale surface structuring of the photoanodemay cause the horizontal scattering of incoming light which increasesthe overall absorption. In contrary, from flat surface of the compactphotoanode some amount of light is directly reflected, and thereforelost (Figure 2.13 b).

34 literature

If the light scattering is weak and the better distribution of lightdoes not contribute significantly to overall absorption, a negative ef-fect, that is reduced thermodynamic driving force available for wateroxidation because of lower absorbed photon flux, might decrease over-all efficiency. Because of the higher roughness (and consequently thehigher surface area), solar flux per unit junction area between nanos-tructured material and electrolyte is lower in comparison with the flatsurface of the compact photoanode (Figure 2.13 c). This unfavorableeffect reduces open circuit voltage available for water oxidation. [20]

From analysis above, it is apparent that there are several trade-offsassociated with reducing the size of particles, however, positive ef-fects usually offset negative. Apart from reducing the size of par-ticles, significant improvements can be also achieved by controllingmaterial morphology on a nanometer scale. [154] It can be illustrated bycomparing four TiO2 morphologies commonly investigated as pho-toanodes for water oxidation: nanoparticles (NP), nanotubes (NT),nanorods (NR), and branched nanorods (BR-NR). The photocurrentdensities for water oxidation, recorded under identical experimentalconditions, increase in the following sequence – lowest values for NP,via medium for NT and NR, to the highest values for BR-NR. Theobserved superior behavior of BR-NR was attributed to high liquidjunction area, outstanding light distribution (light-trapping) and bet-ter electron transport to back-contact properties (Figure 2.14). [154,194]

Ultimately, it is essential to point out, that charge carrier trans-port properties, recalled several times in this Section, are of greatimportance, because the thickness of the space charge layer whichgoverns charge separation in compact, single-crystal-like photoelec-trodes, is significantly reduced in case of nanocrystalline photoelec-trodes. [195,196] In absence of such internal mechanism, charge sepa-ration is determined by efficiency of ETs to back contact, thereforein the case of slow charge transport kinetics, e–/h+ recombination issubstantial.

2.4.2 External Bias

In the case of macro-structured, compact photoanodes, the magni-tude of electric fields generated in the space charge layer can beenlarged by application of external bias (Eappl). Owing to increasedwidth and potential drop in the space charge layer, the photoholes aremuch more strongly driven to the surfaces and electrons to the bulkof the semiconductor, resulting in higher charge separation yields(Figure 2.15). The beneficial effect of external bias can be determinedby transient absorption spectroscopy measurements, which recentlyhave shown three orders of magnitude increase in the life-time of thephotogenerated holes in α-Fe2O3, when 400 mV of additional poten-tial is applied. Importantly, the longer lifetime of the photogenerated


Figure 2.14: Schematic view of three different morphologies of nanos-tructured photoanodes: a) nanoparticels, b) nanorods, andc) branched nanorods deposited on conductive substrate –ITO/FTO (back-contact; in gray). Also distance which electronhas to travel in semiconductor can be reduced using conduc-tive nanowires covered by thin layer of semiconductor (e) in-stead of semiconductor nanorods (d). Adapted from: van deKrol [57], Cho et al. [194]

holes translates directly into higher photocurrents from water oxida-tion. [197]

36 literature

Figure 2.15: Schematic illustration of increased band bending due to the ap-plication of external bias. Increased band bending bring abouthigher charge separation yields.

Although nanoparticles are too small to develop a considerablespace charge layer, external bias has beneficial effect on charge sepa-ration in nanocrystalline photoelectrodes as well. In this case, applica-tion of additional potential increases driving force pulling electronsinto back-contact. In other words, external bias is considered to de-plete nanoparticles from electrons, in this way again limiting recom-bination. [198]

Despite the fact, that the application of external bias significantlyimprove the photoanode performance, it is of course, technologicallyundesirable because represents an extra energy input which shouldbe avoided since it reduces the overall conversion efficiency (see Equa-tion 2.6). [76]

2.4.3 Semiconductor/Co-catalyst Junctions

Detailed discussion of semiconductor/co-catalyst junctions can befound in Section 2.2.2.

2.4.4 Semiconductor/Semiconductor Heterojunctions

Semiconductor/semiconductor junctions utilize rapid ET through theinterface between two different semiconductors to provide spatialseparation of charges, which decreases recombination. They belong toredox heterojunctions, characterized by a series of cascade ETs downto the states with lower energy, without significant participation ofthe space charge layer. Another example of redox heterojunctions aredye-sensitized materials (see next Section) which use ultra-fast elec-tron injection from LUMO of a dye (donor) to CB of TiO2 (accep-


tor), in this way generating spatial separation of charges. It is oppo-site to semiconductor/electrolyte or semiconductor/semiconductorhomojunctions between p- and n- type doped regions of the samematerial,3 where charges are separated owing to gradient of energylevels in the space charge layer.

Figure 2.16: Schematic diagram of energy levels for WO3/SnO2/BiVO4 com-posite material at pH= 0. Adapted from: Saito et al. [199,200]

Semiconductor/semiconductor junctions will be discussed on theexample of a WO3/SnO2/BiVO4 photoanode (Figure 2.16). [199,200] Af-ter generation of e–/h+ pair in BiVO4 directly exposed to irradiation,rapid (faster than recombination) ET to SnO2 interlayer occurs fol-lowed by next ET to CB of WO3. This series of cascade ETs resultsin spatial separation of charges because of very low probability of re-combination over the high distance and through two interfaces. Dueto the increased life-time of charges, the multicomponent photoanodeexhibits 2–3 times higher photocurrents in comparison to the single-component (BiVO4 or WO3) photoanode and yields high dioxygenamounts under AM 1.5G irradiation with stable photocurrent for atleast 5 h. Nevertheless, cascade ETs from CB of BiVO4 to CB of WO3

via SnO2 are associated with energy cost equal to the difference be-tween electrochemical potentials of photogenerated electrons in CBof BiVO4 and WO3.

2.4.5 Sensitization

An interesting approach for water photooxidation is represented byphotoanodes utilizing an electron collector with relatively negativeconduction band edge (typically TiO2) sensitized by an organic com-

3 Like in conventional np-Si solar cells

38 literature

pound (typically a dye) coupled to a colloidal co-catalyst. [201] In thistype of materials, photoinduced ET from LUMO of a dye to CB of thesemiconductor takes place, resulting in spatial separation of charges.Furthermore, this approach is highly interesting as it combines twofeatures characteristic for water oxidation in nature: the kinetic chargeseparation, [179] and the flexibility of the oxidation site, which has beensuggested to be responsible for excellent kinetics of water oxidationin PSII. In addition, it is possible to use stable, wide-bandgap semi-conductors, because the light absorption is not determined by semi-conductor’s bandgap, but by absorption characteristic of a dye whichcover the visible range. For these reasons, dye-sensitized photoanodeshave attracted significant attention in context of water oxidation. [202]

However, the main obstacle in utilizing such assemblies is typicallythe instability of the organic component in harsh conditions duringwater oxidation, as observed, for example, in case of TiO2 sensitizedby the ruthenium dye coupled to IrO2 colloidal nanoparticles. [203]

Taking into account instability of a dye in sensitized photocatalystsfor water oxidation, TiO2-PH hybrid photoanodes, investigated in thepresent thesis, seem to be particularly attractive alternative due to therobustness of polyheptazine and strong coupling with semiconductor(direct ET from HOMO of PH to CB of TiO2).

Part II

E X P E R I M E N TA L D E TA I L S

3E X P E R I M E N TA L D E TA I L S

3.1 instrumental

3.1.1 Photoelectrochemical Measurements

The photoelectrochemical setup for monochromatic photocurrent mea-surements consisted of a SP-300 BioLogic potentiostat (operated byEC-Lab V10.23 software) and a three-electrode cell using a platinumcounter electrode and Ag/AgCl (3 M KCl) reference electrode (EAg/AgCl= 0.207 V vs. NHE). In the majority of cases, potentials were reportedwith respect to the Reversible Hydrogen Electrode (RHE). The poten-tial values were recalculated using the following equation:

ERHE = 0 + Eappl + 0.207 V + 0.059 pH (3.1)

where ERHE is potential versus RHE reference electrode, Eappl is exter-nal potential (bias) applied vs. Ag/AgCl reference electrode. The pho-toelectrodes were pressed against an O-ring leaving an irradiated areaof 0.5 cm2. The electrodes were irradiated from the rear side (throughthe ITO or FTO glass) by tunable monochromatic light source (In-stytut Fotonowy) provided with a 150 W Xenon lamp and a gratingmonochromator with a bandwidth of ∼ 10 nm. Appropriate cut-offfilters were used to eliminate second-order diffraction radiation.

Figure 3.1: The experimental setup for monochromatic photoelectrochemi-cal measurements.

41

42 experimental details

The photoaction spectra were recorded under intermittent irradi-ation. The value of photocurrent density was taken as a differencebetween current density under irradiation and in the dark. The in-cident photon-to-current conversion efficiency (IPCE) value for eachwavelength was calculated according to equation:

IPCE [%] =iphhcλPq

× 100 (3.2)

where iph is the photocurrent density, h is Planck’s constant, c velocityof light, P the light power density, λ is the irradiation wavelength, andq is the elementary charge. The spectral dependence of lamp powerdensity was measured by the NOVA II optical power meter equippedwith a PD300-UV silicon photodiode (Ophir Optronics).

3.1.2 Dissolved Oxygen Concentration Measurements

3.1.2.1 General remarks

In order to assess ability of a photocatalyst to oxidize water, accuratemeasurements of evolved dioxygen are necessary. The photoelectro-chemical cell used for this measurements must be extremely airtightand inlets completely sealed, which makes the direct O2 detectionnot trivial. Consequently, careful selection of the detection methodand PEC cell design are essential to enable the proper quantification.In many instances PEC vessels consist of two compartments (anodicand cathodic) divided by a porous separator, though one compart-ment vessels in the two electrode setup are common as well. In thelatter case reduction of evolved oxygen on the counter electrode (typ-ically Pt) and mixing with hydrogen can be disadvantageous. On theother hand, in the case of the two-compartment cell, the complexityof the system is increased (two argon circulation systems, separator).

There are two general concepts for determination of evolved dioxy-gen. In the first one, quantity of dioxygen is measured in a headspace(in a gas phase above electrolyte), usually by means of gas chro-matography, [204–206] fluorescence-based sensor, [207] or mass spectrom-etry [208]. In the second one, dissolved dioxygen is measured in liquidphase (directly in the electrolyte) by means of a Clark electrode [209,210]

or a fluorescence-based sensor (as in the present dissertation). Sec-ond approach has several advantages, for example, a sensor can bevery close to the photoelectrode, the method is very sensitive (DL= 15 ppb), but in contrast to the first approach, amount of dioxygenin a headspace remains unknown.

In the present thesis, concentration of dissolved (in the electrolyte)dioxygen was determined using fluorescence-based sensor. All oxy-gen evolution measurements were conducted with OxySense GenIII 325i non-invasive oxygen analyzer system. The photoelectrodeswere irradiated from the backside (through the ITO or FTO glass)

3.1 instrumental 43

by a 150 W Xenon lamp (LOT Oriel) equipped with a KG-3 (Schott)heat-absorbing filter. To ensure only visible light irradiation appropri-ate cut-off filters were used. The photocurrent characterization wasdone using Gamry 600 Reference potentiostat (results presented inSection 4.2, Section 5.2, Section 5.3.2, Chapter 6, Chapter 7, and Chap-ter 8) or PalmSense potentiostat (results presented in Section 4.3, Sec-tion 5.3.1, and Section 5.3.3).

Three types of custom-build photoelectrochemical cells were used:

• two-compartment cell for measurements in three-electrode setup(Section 3.1.2.2),

• upgraded version of the first cell (Section 3.1.2.3),

• one-compartment cell for measurements in two-electrode setup(Section 3.1.2.4).

In following paragraphs, design of vessels and methods employed fordetection of dissolved dioxygen produced during photoelectrochemi-cal water splitting are explained in details.

3.1.2.2 First generation cell for DO measurements in three-electrode setup

The key experiments for research presented in this thesis were con-ducted in the photoelectrochemical cell shown in Figure 3.2. As can

Figure 3.2: First generation photoelectrochemical cell for DO measurementsin three-electrode setup during operation.


be seen, there are four inlets, two for electrodes: counter (CE) and ref-erence (RE), and two for Ar circulation system. Moreover, there is one

a)

b)

Figure 3.3: a) A 3-D front and b) a 2-D cross-sectional view of the pho-toelectrochemical cell for DO measurements in three-electrodesetup. Design and construction with help from Mr. Armin Lind-ner (RUB, Feinmechanikwerkstatt).

hole for the photoelectrode, which is pressed against an O-ring leav-ing an irradiated area of 0.5 cm2. The volume of a electrolyte in the an-odic compartment was 5.0 cm3; the oxygen sensitive film (OxyDot®)was attached inside on a transparent glass piece. Dissolved oxygenconcentration was determined by measuring the fluorescence gener-ated in oxygen sensitive film by sensor mounted outside the cell (Fig-

3.1 instrumental 45

ure 3.3). The aluminum foil was used to minimize the temperatureincrease during irradiation.

Argon inlets were closed with ø 8 mm turn-over flange stoppers(Saint-Gobain Performance Plastics), then needles (Braun Sterican 0.90× 40 mm, Emasa Hypodermic 0.90 × 120 mm) were inserted insideto purge electrolyte with argon (Air Liquide N46 99.996 vol %, O2

< 6 ppmv). Before turning on the light, needles were moved up andonly the headspace was purge with argon. Setup was left for 45 minto ensure that no oxygen is leaking from outside. The atmosphereabove electrolyte (in headspace) was dynamic, that is argon was cir-culating continuously also during irradiation.

Results presented in Section 5.2, Section 4.2, and Chapter 7 were ob-tained form measurements conducted in first generation cell withoutstirring.

3.1.2.3 Second generation cell for DO measurements in three-electrodesetup

In contrast to the reactor described in the previous Section, in the up-graded version the atmosphere in the headspace was static, in otherwords, it was not purged with argon during the measurement. Addi-

Figure 3.4: A cross-sectional view of argon inlet/outlet system in first (a)and second (b) generation of photoelectrochemical cell for DOmeasurements in three-electrode setup.

tional extension for gas circulation system was used in order to im-prove detection capability (Figure 3.4). It consisted of threaded plastictube and two ø 8 mm PTFE/red rubber septa. Prior to measurement,space between septa was purged with argon, followed by insertion of


needles and bubbling of electrolyte. Before irradiation needles werewithdrawn and argon was circulating only in the space between septa,while the headspace filled with argon remained closed. Setup was leftfor 45 min to ensure that no oxygen is leaking from outside. The vol-ume of electrolyte in the anodic compartment was 5.7 cm3.

Results presented in Section 4.3, Section 5.3, and Chapter 6 wereobtained from measurements conducted in this cell under stirringconditions.

3.1.2.4 Photoelectrochemical cell for DO measurements in two-electrodesetup

Dissolved oxygen evolution measurements presented in Chapter 8were done in a two-electrode setup with platinum wire as a counterelectrode in the one-compartment cell (Figure 3.5). In a glass vial(VWR International) oxygen sensitive film (OxyDot®; yellow circle)was attached. The vial was sealed using a rubber septum (turn-overflange stopper, ø 15.9 mm) and argon (Air Liquide N46 99.996 vol. %,O2 < 6 ppmv) was used for purging. After 1 h of bubbling of theelectrolyte prior to the measurement, argon was stopped and only at-mosphere above the solution was purged. Before turning on the light,setup was left for 45 min to ensure that no oxygen is leaking fromoutside. Total volume of the vial was 15 ml and electrolyte volumewas 7 ml. Entire area of the photoelectrode was irradiated.

Figure 3.5: Experimental setup for dissolved oxygen measurements in a one-compartment cell in the two-electrode setup.

3.2 characterization 47

3.2 characterization

uv-vis diffuse reflectance spectra were recorded on samples(25 mg) diluted in BaSO4 (0.5 g) using a Harrick praying mantis dif-fuse reflectance accessory mounted in a Perkin Elmer Lambda 650UV−Vis spectrophotometer. BaSO4 was used as the reference. TheKubelka-Munk function F(R∞) was calculated as

F(R∞) =(1− R∞)2

2R∞(3.3)

where R∞ is diffuse reflectance of the sample relative to the reflectanceof a standard.

scanning electron microscopy (sem) and energy-disper-sive x-ray spectroscopy (edx) investigations were performedusing FEI ESEM Dual Beam™ Quanta 3D FEG with EDAX GenesisXM2i system.

uv-vis spectra were recorded using Agilent Technologies Cary60 UV-Vis Spectrophotometer.

x-ray diffraction spectra (xrd) were recorded using BrukerD8 Advance Diffractometer with Mo–K(α) radiation source operatingat 50 KV and 40 mA. Samples were measured in glass capillary tubes(ø 0.5 mm) step 0.01º and in the range of 2θ = 5–35º.

3.3 synthesis

In this Section the general way of photoanode preparation is de-scribed. The consecutive processing or modifications of the protocoldescribed below, e. g. implementation of co-catalyst, are detailed inthe experimental part of each Section in Part III of this thesis.

The conducting ITO-glass substrate (Präzision Glas & Optik, sheetresistance of ∼ 10 Ω sq−1) or FTO-glass substrate (Sigma Aldrich,sheet resistance of ∼ 7 Ω sq−1) was first cut into 2.5 × 1.5 cm pieces,then subsequently degreased by sonicating in acetone and immersingin boiling NaOH (0.1 M), then rinsed with demineralized water, andfinally dried in air stream.

A suspension containing 200 mg of TiO2 (Hombikat UV 100, Sacht-leben, Duiburg, Germany, specific surface area Brunauer-Emmett-Teller(BET) ∼ 300 m2 g-1, crystallite size < 10 nm) in 1 cm3 of ethanol wassonicated for 15 min and then smeared onto ITO- or FTO-coated glassby doctor blading using a scotch tape as frame and spacer. The elec-trodes were then dried at 100 °C for 15 min, covered with a pieceof the normal glass (Thermo Scientific, Menzel-Gläser) pressed for3 min at a pressure of 200 kg cm-2 (IR Hydraluic Press, Perkin-Emer),and calcined in air at 450 °C for 30 min in order to sinter the parti-cles and ensure good electrical contact (Carbolite MTF 12/38/40 tubefurnace). This procedure yields ∼ 2.5 µm thick layer of TiO2 withexcellent mechanical stability.


The TiO2 electrodes were modified with polyheptazine by placinginto a Schlenk tube connected via an adapter with a round bottomflask containing 1 g of urea (J.T Beaker, ACS) and heating in a muffleoven for 30 min at 425 °C (Carbolite ELF 11/23 box furnace).

The photoanodes prepared by above described manufacturing pro-tocol are denoted hereafter as TiO2-PH.

All water solution utilized for both the synthetic purpose and mea-surements were prepared using distillated, demineralized water cleanedby Siemens Ultra Pure Water Systems (ultra clear; conductivity: 0.055µS/cm, TOC content: < 1 ppb, Endotoxins: < 0.001 EU/m).

Part III

R E S U LT S A N D D I S C U S S I O N : H Y B R I DP H O T O A N O D E S L O A D E D W I T H D I F F E R E N T

C O - C ATA LY S T S

4I R I D I U M - B A S E D C O - C ATA LY S T S

4.1 introduction

Iridium-based materials have been known for a few decades to be ex-cellent electrocatalysts for water oxidation. [211–213] Comparative stud-ies of Mills et al. [214] and Trasatti [215] identified IrO2 to be (togetherwith RuO2) the best catalyst out of the group of several metal ox-ides. Iridium oxide-based catalysts still belong to the most studiedones because of their ability to work at very low overpotentials (aslittle as 0.25 V) and outstanding robustness against anodic corro-sion. [216–218] Therefore, they can be found in dimensionally stable an-odes [219] and in industrial water electrolyzers based on solid polymerelectrolytes. [220]

Water photooxidation with iridium-oxide catalyst was first demon-strated by Harriman et al. [157] who used colloidal IrOx · n H2O pho-tosynthesized with ruthenium dye. The method of colloidal iridiumoxide preparation significantly affects the catalytic properties, includ-ing both activity and stability. Different sensitizers were examinedand the size of colloidal particles was optimized in order to achievethe highest possible photooxidation rates. [221–223] It is also importantto mention once again, that iridium oxide nanoparticles can be linkedvia sensitizers (e.g. Ru dyes) to semiconductor particles (TiO2) toserve as a photoanode in dye-sensitized photoelectrochemical cells(see Introduction and Section 2.4.5). [201,203] Interestingly, colloidal nano-particles of iridium oxide were reported to release dioxygen fromaqueous solutions containing electron scavengers (S2O8

2- or Ag+) un-der both UV and visible light irradiation even without sensitization.This unanticipated result was accounted for very small (< 2 nm) sizeof nanoparticles which facilitates oxidation using a material with veryshort life-time of holes. [224]

Due to the very high catalytic activity, iridium oxides were usedquite frequently to promote water oxidation in both photocatalyticand photoelectrocatalytic systems. For example, 2 wt.% loading ofTi-based oxysulfide (Sm2Ti2S2O) photocatalyst with IrOx co-catalystresulted in 5-fold increase in photocatalytic activity. [225] Similarly, co-catalyst loading was demonstrated to increase stability and photo-catalytic activity of the lanthanum titanium oxynitride. [226] The thinfilms of LaTiOxNy were also examined as photoanodes to use in pho-toelectrochemical cells. The iridium oxide co-catalyst was introducedusing the colloidal deposition method, that is by simply soaking ofthe photoanode in the colloidal solution of iridium oxide nanopar-

51

52 iridium-based co-catalysts

ticles, which results in their adsorption onto the photocatalyst sur-faces. [227] The colloidal solution of IrOx was prepared by alkaline hy-drolysis of iridium complexes like [IrCl6]2-. The same method wasused in order to load tantalum oxynitride (TaON) and tantalum ni-tride (Ta3N5). [228,229] The presence of the co-catalyst improves stabil-ity by lowering the photoelectrocatalyst self-oxidation rate, increasesphotocurrents by factor of 4, and enables water photooxidation evenat potentials as low as 0.6 V vs. Pt. In the another example iridiumoxide was loaded by the colloidal deposition method on strontiumniobium oxynitride (SrNbO2N) coated on FTO glass. The photoanodeprepared in this way produced dioxygen from water solution underthe irradiation with visible light at 1.0 V vs. RHE with one order ofmagnitude higher rate than unloaded SrNbO2N. [230]

Iridium oxide co-catalyst was also coupled to hematite photoan-ode (α-Fe2O3). The shift of the photocurrent onset from + 1.0 V to+ 0.8 V vs. RHE and improved photocurrents were observed after thedeposition of the co-catalyst. Importantly, in this case iridium oxideco-catalyst was deposited by electrophoresis which was found to besuperior to the colloidal deposition method. Nevertheless, it was re-ported that iridium oxide particles do not adhere firmly and detachfrom the surfaces of hematite. [231]

In general, iridium-based oxides have a great capability to be im-plemented as co-catalysts in photoelectrochemical devices due to thevery high activity. However, the development of more robust attach-ment methods ensuring better coupling and uniform distribution ofthe co-catalyst on the absorber surface is still needed.

Next Sections present results of investigations of TiO2-PH photoan-odes with iridium-based co-catalyst loaded by the reported in litera-ture colloidal deposition method (IrOx/CD) and the photodepositionmethod developed in our group (IrOx/PD).

4.2 colloidal deposition of iridium oxide

4.2.1 Experimental

Iridium oxide nanoparticles (denoted as IrOx) were synthesized us-ing the method described by Maeda et al. [230] Briefly, 0.0080 g ofsodium hexachloroiridate(IV) hexahydrate, Na2IrCl6· 6 H2O was dis-solved in 50 mL H2O, and the pH of the solution was adjusted to 12with sodium hydroxide (NaOH). The solution was heated at 80 °Cfor 30 min, and then cooled in an ice-water bath. Afterwards, pHwas adjusted carefully to 9 with nitric acid (HNO3) solution. Furtherheating at 80 °C for 30 min resulted in deep blue solution containingcolloidal IrOx. During the colloidal deposition, the TiO2-PH photo-electrodes (prepared like in Section 3.3 on ITO glass) were immersedinto the diluted (∼ 30 %) solution of IrOx nanoparticles overnight and

4.2 colloidal deposition of iridium oxide 53

then dried in air. Photoanodes prepared in this way are denoted asTiO2-PH + IrOx/CD.

Dissolved oxygen measurements were conducted according to Sec-tion 3.1.2.2 without stirring. In order to detect hydrogen peroxideH2O2, horseradish peroxidase-based photometric test (BioVision, Hy-drogen Peroxide Kit) was used.

4.2.2 Results and Discussion

In case of pristine TiO2, loading with IrOx nanoparticles led to a de-crease of IPCE values, most probably due to partial blocking of UVlight absorption by the co-catalyst (Figure 4.1 c). In case of TiO2-PHa similar decrease has been observed in the UV range, while in thevisible (λ > 400 nm) the presence of a co-catalyst increased the IPCEvalues slightly (Figure 4.1 d). At 450 nm the IPCE value was 3 %.

Figure 4.1: IPCE values at different wavelengths measured at 0.5 V vs.Ag/AgCl (1.12 vs. RHE) for TiO2 and TiO2-PH photoelectrodesin a phosphate buffer (pH 7) before and after deposition of IrOxnanoparticles as a co-catalyst.

Notably, when comparing directly the photoresponse of IrOx-loadedTiO2-PH with IrOx-loaded pristine TiO2, a significant enhancement inthe visible range is observed (Figure 4.2 a) that correlates well with


Figure 4.2: a) Photoaction spectrum measured in a phosphate buffer (0.1 M;pH 7) at 0.5 V vs. Ag/AgCl (1.12 vs. RHE) for TiO2-PH andpristine TiO2 photoelectrodes loaded with IrO2 nanoparticles. b)Oxygen evolution and c) photocurrent at 0.5 V vs. Ag/AgCl dur-ing prolonged irradiation by polychromatic visible light (cut-offfilter λ > 420 nm).

the shift of the optical absorption shown in Figure 1.3. Even more im-portantly, irradiation of IrOx-loaded TiO2-PH photoelectrodes withvisible light (λ > 420 nm) led to evolution of dioxygen (Figure 4.2b). As expected, under identical conditions, the pristine TiO2 pho-

4.3 photodeposition of iridium oxide 55

toelectrode modified with IrOx nanoparticles shows only negligiblephotocurrents and does not show any oxygen evolution.

We found out that the presence of the oxygen evolving co-catalystis absolutely crucial in order to observe dioxygen as product of wa-ter photooxidation at TiO2-PH electrodes. TiO2-PH photoelectrodeswithout a co-catalyst exhibited no dioxygen evolution (Figure 4.2b) though photocurrents were relatively high (Figure 4.2 c). Whensearching for possible intermediates of the oxidation reaction we didnot detect any hydrogen peroxide using a standard horseradish per-oxidase-based photometric test. It cannot therefore be ruled out thatin the absence of the co-catalyst the anodic photocurrents lead also tooxidation of polyheptazine or some side-products left from the mod-ification procedure. However, it is apparent that coupling of TiO2-PH with the co-catalyst improved the stability of the photocurrentresponse under prolonged irradiation (Figure 4.2 c). The faradaic ef-ficiency of the current-to-dioxygen conversion was calculated to beca. 19 %, which suggests that the coupling of the IrOx nanoparticlesto the polyheptazine layer is far from perfect. It should be noted thatthe bad adhesion of IrOx nanoparticles to other photoactive materialslike, e.g. Fe2O3, is known from the literature. [231] At any rate, a moreefficient coupling of holes to the catalytic sites for water oxidationcan be expected to avoid excessive accumulation of holes in the poly-heptazine layer, which should improve the conversion efficiency andstability of the hybrid photoelectrodes.

4.3 photodeposition of iridium oxide

Photodeposition method is frequently employed in order to depositcatalytically active species onto semiconductor nanostructures. Forexample, it was used to deposit various precious metals [232–234] orPbS quantum dots [235] on the TiO2 surface. Photodeposition methodhas many advantages, for instance: more controllable size and distri-bution of particles, in-situ formation of particles at semiconductor’spreferential active sites, or more efficient charge transfer to semicon-ductor, which is ensured by better contact between components.

Photodeposition method was also used to deposit cobalt or nickeloxygen evolving co-catalysts on many of semiconducting materials(see Chapter 5 and Chapter 6). In this Section, it was used to depositiridium oxide metal clusters, acting as a oxygen evolving co-catalyst,onto TiO2-PH photoanode.

4.3.1 Experimental

Photodeposition method (TiO2-PH + IrOx/PD): the hexahydroxyiri-date(IV) solution was prepared drawing on the method of Mallouket al. [236] 45 mg of Na2IrCl6 · 6 H2O was dissolved in 50 ml of water,


giving brownish coloration of solution. The pH was set to 12.1 usingaqueous NaOH. The solution was heated to 70 °C (during heating thesolution lost color) and then cooled down immediately in an ice bath(resulting in colorless or slightly bluish solution). The pH of the coldsolution was carefully adjusted to 8.5 with HNO3. This procedure re-sulted in the 1.5 mM colorless or slightly yellowish solution, whichwas diluted to 0.45 mM and used for the photoelectrochemical de-position at TiO2-PH photoelectrodes directly after the synthesis. Thedeposition potential was set to 0.5 V vs. Ag/AgCl (1.21 V vs. RHE)and the photoanode was irradiated for 1200 s with monochromaticvisible light (λ = 420 nm).

Colloidal deposition method (TiO2-PH + IrOx/CD): analogically toSection 4.2.1.

Dissolved oxygen measurements were conducted according to Sec-tion 3.1.2.3 with stirring.


Figure 4.3 shows photocurrent generation during irradiation of theTiO2-PH hybrid photoanode in contact with the electrolyte solutioncontaining hexahydroxoirradiate(IV) anions. [Ir(OH)6]2- is an interme-

Figure 4.3: Photocurrent transient recorded at a TiO2-PH photoelectrodeduring photoelectrochemical deposition of IrOx nanoparticlesfrom an [Ir(OH)6]2– solution. The electrode was irradiated withmonochromatic visible light (λ = 420 nm) at 0.5 V vs. Ag/AgCl(pH 8.5). The optical images show the darkening of the hybridphotoelectrode due to formation of IrOx nanoparticles. Scanningelectron micrographs show the surface of the hybrid photoelec-trode before (left) and after (right) the photodeposition process.

diate complex during formation of colloidal IrOx nanoparticles. Elec-trodeposition from this precursor was reported to yield stable and


very active films of IrOx nanoparticles which exhibit very low over-potential (0.2 V) and Tafel slope of b = 39 mV per decade. [236] It washypothesized that the mechanism of IrOx electrocatalyst formationis based on acidic condensation of [Ir(OH)6]2- due to the proton for-mation during the oxidation of H2O molecules. [236] In line with thissuggestion, one can expect that the change of local pH during waterphotooxidation at the irradiated photoanode facilitates the acid con-densation mechanism. It can be assumed that the initial step of thereaction involves oxidation of Ir(IV) to Ir(V) by the photogeneratedholes.

The irradiation with the monochromatic visible light (λ = 420 nm)of the TiO2-PH photoanode led to the dark brown colorization, dueto the formation of the IrOx oxygen evolving co-catalyst, as can beseen in optical images in Figure 4.3. The quantity of the co-catalystmust be fine-tuned in order to avoid blocking of the light absorption.Significantly, the photodeposition method offers, apart from the as-surance that the co-catalyst is formed preferentially at the sites withthe high concentration of the photogenerated holes, also possibilityof controlling its amount by varying the photodeposition time andpotential, the concentration of [Ir(OH)6]2- or the light wavelength.

Figure 4.4: EDX analysis of TiO2-PH + IrOx/PD photoelectrode.

The morphology of TiO2-PH photoelectrodes with photodepositedIrOx was examined with SEM (Figure 4.3). Investigations revealedthat surface of TiO2-PH is covered with IrOx particles whose sizevaries between 20–100 nm, as can be clearly seen by comparison withthe SEM image of TiO2-PH before photodeposition. EDX analysis(Figure 4.4) done on the representative spot1 of the photoanode rev-

1 Representative spot means that EDX spectrum was measured on the bigger (not onthe one shown in SEM micrographs) spot with an average distribution of co-catalystparticles.


Figure 4.5: Photoaction spectra (IPCE vs. wavelength) recorded in a phos-phate buffer (0.1 M; pH 7) at 1.12 V vs. RHE at TiO2-PH hybridphotoelectrodes with IrOx co-catalysts deposited by colloidal de-position (IrOx/CD) and photodeposition (IrOx/PD).

eled that Ir content amounts to 3.71 wt.% (after correcting for Si, In,Ca amount).

In order to asses performance and effectiveness of the photode-position method for implementation of the iridium-oxide based co-catalyst to the photoanode, a series of photoelectrochemical measure-ments was performed. In these experiments TiO2-PH with IrOx loadedby the colloidal deposition method was chosen as a benchmark. Fig-ure 4.5 compares IPCE values at different wavelengths, recorded atthe photoanode loaded with iridium oxide co-catalyst by the colloidaldeposition method (TiO2-PH + IrOx/CD) and the photodepositionmethod (TiO2-PH + IrOx/PD). It can be seen from the graph that interms of photocurrents generated under monochromatic light irradi-ation both methods give similar results, especially in the visible lightrange. Nevertheless, TiO2-PH + IrOx/PD photoanode is capable ofutilizing the photogenerated holes in order to produce dioxygen withmuch higher yields. This is illustrated by Figure 4.6 which plots oxy-gen evolution and photocurrent measurements for TiO2-PH loadedwith IrOx by both methods, recorded in phosphate buffer pH = 7,under the polychromatic visible light (λ > 420 nm) irradiation. TheTiO2-PH + IrOx/PD produced considerably higher dioxygen concen-trations and evolution is observed during the entire irradiation time.In the case of the TiO2-PH + IrOx/CD, after an initial rise, the oxygenconcentration quickly reaches the plateau, followed by a continuousdecrease. The maximum concentration for TiO2-PH with IrOx appliedby the colloidal deposition method is 2.5 times lower than for the pho-toelectrochemical deposition. Moreover, taking into account the factthat photocurrents are comparable (similarly to IPCE presented in


Figure 4.6: Oxygen evolution (a) and photocurrents (b) during prolongedirradiation by polychromatic visible light (cut-off filter 420 nm)in a stirred phosphate buffer solution (0.1 M; pH 7) at 1.12 V vs.RHE.

Figure 4.5), it is clear that the photodeposition method yields muchhigher faradaic efficiencies. These observations clearly indicate thebetter ability of TiO2-PH photoanode with photodeposited IrOx tophotooxidize water, most likely due to the more efficient couplingbetween IrOx/PD co-catalyst and TiO2-PH absorber. This is not sur-prising as colloidal deposition is based simply on soaking of a pho-toanode in a colloidal solution and adsorption of IrOx, while in caseof photodeposition, IrOx is formed photoelectrochemically exactly atsurface sites where photoholes are predominantly generated.

Figure 4.7 plots photocurrent transients recorded during cathodicpotential sweep between + 1.5 V and – 0.4 V vs. RHE. As can be seen,TiO2-PH + IrOx/PD exhibits a “dark” negative peak between + 0.6


Figure 4.7: Photocurrent measured at TiO2-PH photoelectrodes with IrOx co-catalysts deposited by colloidal deposition (IrOx/CD) and pho-todeposition (IrOx/PD) under intermittent monochromatic (λ =420 nm) irradiation (5 s light, 5 s dark) during a cathodic poten-tial sweep (5 mV s-1) in a deaerated phosphate buffer (pH 7).

and + 0.2 V, which is ascribed to the electrochemical reduction (IrIV toIrIII) of IrOx present on the ITO glass. Such peak is not observed forTiO2-PH + IrOx/CD, because in the case of the colloidal depositionthe amount of IrOx located on the ITO glass is expected to be muchsmaller. Another interesting point is the change of photocurrent signfrom positive to negative in this range of potentials. The reason forthis is photoelectrochemical reduction of Ir(IV) to Ir(III) in iridiumoxide co-catalyst present on TiO2-PH particles. More precisely, photo-generated electron reduces iridium oxide instead of being transferredto the conductive substrate (ITO/FTO). At the same time photogen-erated holes are transfered to the conductive substrate, giving riseto negative photocurrent. The reduction of photodeposited iridiumoxide co-catalyst is an important limitation which is, however, not ob-served in case of e.g. photodeposited cobalt-based oxygen evolvingco-catalyst (Co-Pi, Figure 5.6 a). Above the potential of + 0.6 V pho-tocurrents for TiO2-PH + IrOx/PD and TiO2-PH + IrOx/CD are com-parable. However, the clearly less pronounced spike-shape of pho-tocurrents for TiO2-PH + IrOx/PD indicates lower surface recombina-tion rates.

This can be further illustrated by data in Figure 4.8, which com-pares photocurrent transients recorded under potentiostatic conditions,under the monochromatic visible light irradiation. Two issues arenoteworthy. Firstly, as expected after analysis of Figure 4.6 b) pho-tocurrents are approximately two times higher when IrOx co-catalystis present. Second aspect worth explanation is the shape of transients.Both TiO2-PH and TiO2-PH + IrOx/PD transients exhibit the spike-


Figure 4.8: Photocurrent transients recorded at TiO2-PH hybrid photoelec-trodes without and with photodeposited IrOx co-catalyst underirradiation with monochromatic visible light of different wave-lengths in a Na2SO4 (0.1 M; pH 7) at 1.12 V vs. RHE.

like shape, that is a rapid exponential decay of photocurrent imme-diately after switching on and a cathodic current “overshoot” afterswitching off the light. The spike-like shape of photocurrent tran-sients is indicative of surface recombination processes going on. [237,238]

After the initial rise of photocurrent upon switching-on the light,a rapid decay is observed, which can be ascribed to accumulation ofphotogenerated holes in the surface polyheptazine layer. The recom-bination of accumulated photoholes continues even after the light isswitched off, giving a rise to the cathodic current “overshoot”. Signif-icantly, the most striking observation to emerge from comparison inFigure 4.8 is considerably less pronounced spike-like shape (initial de-cay and “overshoot” ) of photocurrents observed for photoelectrodewith the photodeposited co-catalyst (TiO2-PH + IrOx/PD) than forone without a co-catalyst (TiO2-PH). Hence, it can be concluded thatthe presence of IrOx/PD co-catalyst suppresses the excessive accu-mulation of photogenerated holes in polyheptazine layer in this waydecreasing surface recombination rate.

All measurements so far (and most in this dissertation) were donein the same electrolyte: phosphate buffer, pH = 7. This approach isconvenient because it allows for direct comparisons, but at the sametime the influence of electrolyte composition on photoelectrocatalyticproperties remains unknown. The previous literature studies haveshown that buffer anions and pH play a very important role in theOER at Co-based oxygen evolving catalyst (Co-Pi, Section 5.2), in-cluding fundamental changes in the catalytic mechanism and activ-ity. [239,240] On this account, it seems interesting to investigate influ-ence of different electrolytes and pH on photoelectrocatalytic activityof TiO2-PH + IrOx/PD photoanodes.


Figure 4.9: a) Oxygen evolution and b) photocurrents during prolonged ir-radiation by polychromatic visible light (cut-off filter 420 nm) at1.05 V vs. RHE in different electrolytes (pH 5.9) under stirringconditions.

In Figure 4.9 the effect of solution anion on amount of evolveddioxygen and photocurrent is presented. Three different electrolyteswith same molar concentration equal to 0.1 M, and pH equal to 5.9,were used: phosphate buffer, sodium sulfate and lithium perchlo-rate. Although photocurrent is almost same for all three electrolytes,amounts of evolved dioxygen are significantly different. The lowestamount of evolved dioxygen was measured for phosphate buffer,medium for LiClO4, and the highest for Na2SO4 electrolyte. Thesefindings are rather interesting since phosphate buffer is frequentlyutilized in water oxidation research because of perfect buffering prop-erties at neutral pH and resistance against oxidation. Nevertheless,


they seem to be consistent with reports, which found that in compar-ison to HSO4

– or ClO4–, H2PO4

2- anions reduce activity of iridiumoxide in OER reaction. [241] A possible explanation for these resultsmay be specific adsorption of phosphate anions on the iridium oxidesurface. [241] This is in accord with the study demonstrating that phos-phate buffer has beneficial influence on OER rate due to buffering ef-fect only at concentrations below 5 mM. [242] At higher concentrationsthe phosphate buffer has detrimental influence on OER at iridium ox-ides probably due to the coordination to metal active sites, in otherwords, phosphate anions, especially HPO4

2-, can bind more stronglyto high-valent iridium intermediates formed during a catalytic cyclethan neutral water molecules, deactivating thus the catalyst and in-ducing its corrosion. [242]

Apart from the anion effect, another major factor is the pH of theelectrolyte. To asses impact of pH on photoelectrochemical activityof TiO2-PH + IrOx/PD, a series of measurements was conducted inNa2SO4 solution adjusted to selected pH value by addition of H2SO4

or NaOH. Results of these investigations are shown in Figure 4.10.Photocurrent and dissolved oxygen concentration were measured inelectrolytes with six different pH values: 2, 4, 6, 8, 10, 12. Changes inwater oxidation potential with respect to pH were taken into accountby adjusting the external bias in all six measurements to a constantvalue of Eappl = 1.12 V vs. RHE. Local pH changes can be assumedto not play a significant role because of stirring of electrolyte duringmeasurements and relatively low photocurrent densities.

Figure 4.10 c) compares all results: the amount of evolved dioxy-gen, the charge passed (current integrated over the time), and thecalculated faradaic efficiency. These results are revealing in severalways. To begin with, the strong influence of pH on photoelectrocat-alytic activity is observed, indicating fundamental importance of elec-trolyte properties on water photooxidation using TiO2-PH + IrOx/PDsystem. In the investigated pH range, three regions can be distin-guished. In acidic conditions the performance of the photoanode wasfound to be optimal as evidenced by high photocurrents, amounts ofevolved dioxygen, and corresponding faradaic efficiency. Maximumvalue of FE = 44 % was found for pH = 6. Further analysis shows thatin the range of pH from 6 to 10 both photocurrent and amount ofevolved dioxygen decrease remarkably. However, the faradaic effi-ciency, which can be interpreted as an estimation of quality of cou-pling between co-catalyst and the photogenerated in the absorberholes, still preserves high values, and photocurrents are stable. There-fore, the decrease observed in the range of pH from 6 to 10 cannotbe attributed to instability of the system but rather to enhanced re-combination. It can be due to the lower activity of iridium oxide co-catalyst. This explanation is corroborated by studies in the field of“dark” electrocatalysis of OER which have shown that in unbuffered


Figure 4.10: a) Oxygen evolution and b) photocurrents during prolonged ir-radiation by polychromatic visible light (cut-off filter 420 nm)in a stirred Na2SO4 (0.1 M) solution with pH adjusted to dif-ferent values. The bias was adjusted in all cases to 1.12 V vs.RHE. c) Oxygen concentration in the solution (a), total electriccharge passed ( ), and calculated faradaic efficiency (`) after1 hour irradiation at different pH values.

4.4 conclusion 65

solutions overpotentials for water oxidation at IrO2 considerably in-crease in neutral and weakly alkaline conditions, as compared withoverpotentials in acidic conditions. [218,242,243] Consequently, the lowerwater oxidation reaction rate at IrOx co-catalyst in the range of pHfrom 6 to 10 results in stronger accumulation of photoholes in poly-heptazine layer what leads to the increased surfaces recombination.In the third range, that is above pH = 10, the detrimental effect ofalkaline conditions can be clearly seen. At pH = 12, rather high butvery unstable photocurrents can be observed, while the amount ofevolved dioxygen is small, pointing to a very low faradaic efficiency.The same photocurrent behavior was found for TiO2-PH without co-catalyst, indicating that in alkaline media TiO2-PH photoanode isprone to photocorrosion. However, observed instability can be alsoattributed to problems with the iridium oxide co-catalyst. It was re-ported in literature, that the stability of iridium oxide at pH higherthan 11 is limited by formation of highly oxidized species – iridates(IrO4

2-). [243–245] Consequently, the observed deactivation of TiO2-PH+ IrOx/PD in water photooxidation in pH above 10 is most proba-bly due to both instability of IrOx co-catalyst and photocorrosion ofTiO2-PH absorber.

4.4 conclusion

In conclusion, the photoelectrochemical properties of nanocrystallinephotoelectrodes based on TiO2-PH hybrids loaded with a iridium ox-ide co-catalyst were investigated in order to evaluate their ability tophotooxidize water. The oxygen evolution under visible-light (λ > 420nm) irradiation on hybrid electrodes loaded with IrOx nanoparticles,acting as oxygen evolution co-catalysts, was demonstrated. Impor-tantly, the presence of a co-catalyst was found to be crucial to observedioxygen as a product – TiO2-PH photoanodes without IrOx nanopar-ticles do not evolve dioxygen upon irradiation.

In the first part of this Chapter (Section 4.2) iridium oxide nanopar-ticles were loaded using the colloidal deposition method. The inves-tigations have shown that the quality of the coupling between thecharges produced in the light absorber (inherently a one-electron/holeprocess) and the co-catalyst allowing for chemical transformation alongmultielectron pathways is of fundamental importance and should beimproved. Therefore, the photodeposition method of loading iridiumoxide cluster onto the photoanode surface was developed as an al-ternative to the colloidal deposition method. In the photodepositionmethod the co-catalyst is formed photoelectrochemically in-situ onthe surface of the photoanode in contact with electrolyte containing[Ir(OH)6]2- during irradiation with the visible light. The better per-formance of TiO2-PH with photodeposited IrOx than loaded with col-loidal deposition – 2.5 times higher dioxygen evolution yield – was as-


cribed to much more effective coupling between the photogeneratedholes and the co-catalyst. Moreover, the presence of the co-catalystwas found to decrease surface recombination rate due to faster in-jection of photogenerated holes to electrolyte which minimized theiraccumulation.

The influence of an anion type and pH on photocatalytic propertiesof TiO2-PH + IrOx/PD photoanodes was investigated. It was foundthat photoanodes exhibit highest activity and stability in Na2SO4 atpH = 6. The observed strong dependence of activity on electrolyteproperties indicates that measurement conditions must be cautiouslyselected in order to assure compatibility with all essential compo-nents of photoanode. Consequently, a photoelectrochemical cell op-timization must include parameters related not only to an absorberand a co-catalyst but to an electrolyte as well.

Obviously, co-catalysts based on cheaper and more abundant ma-terials than IrOx should be employed, and research presented in thenext Sections goes in this direction.

5C O B A LT- B A S E D C O - C ATA LY S T S

5.1 introduction

The use of IrOx as a co-catalyst, does not represent a viable strategydue to the high price of iridium. However, since the seminal work ofChen et al. [246] and Harriman et al. [157] it has been known that othertransition metal oxides, based on more abundant elements, for exam-ple on cobalt, can also act as effective catalysts for oxygen evolution.During last 30 years, a large and still growing body of literature hasinvestigated lots of both homo- and heterogeneous catalysts based onthis element. [247,248] Numerous studies attempting to discover a cata-lyst based on Co have found that especially Co3O4 and related cobaltoxo/hydroxo species containing CoII and CoIII are particularly activetowards water oxidation. Some of them are collated in Table 5.1.

catalyst pH remarks ref

elec

troo

xida

tion Co3O4 14

cubic nanoparticles with different

diameter[249]

SWNTs/Co3O4 7nanoparticles supported on carbon

nanotubes[250]

CoOx (Co-Pi) 7oxo/hydroxo clusters; precursor:

Co2+ / phosphate buffer[251]

CoOx 8oxo/hydroxo clusters; precursor:

[Co4(H2O)2(PW9O34)2]10-[252]

Co2O3 NP 7nanoparticles on carbon felt,

obtained by photoreduction of CoCl2[253]

phot

ooxi

dati

on Co/M2P 7small (10 - 60 nm) spherical

nanoparticles; Ru(bpy)33+/ S2O82-

[254]

SBA-15/Co3O4 6nanostructured clusters in the silica

matrix; Ru(bpy)33+/ S2O82-

[255,256]

Co3O4 5 Ru(bpy)33+/ S2O82- [157]

NiCo2O3 5 Ru(bpy)33+/ S2O82- [157]

SiO2/Co(OH)2 9hydroxide clusters on silica

Ru(bpy)33+/ S2O82-

[257]

Table 5.1: Selected Co-based water oxidation catalytic systems.

67

68 cobalt-based co-catalysts

Importantly, Co3O4 can act as a co-catalyst after coupling with pho-toanode material. It was recently demonstrated that Co3O4-decoratedFe2O3 nanorods, synthesized by hydrothermal method, exhibit in-creased activity in water photooxidation. [258] Addition of Co3O4 tothe hematite photoanode resulted in higher photocurrents (1.2 mAcm-2 at 1.23 V vs. RHE, whereas for bare Fe2O3 nanorods 0.72 mAcm-2), improvement in IPCE (doubled in range form 400 to 550 nm)and increase in amount of evolved dioxygen. This study producedfindings which are in line with another results showing that Co3O4

introduced to W-doped BiVO4 photoanodes also increases photocur-rent values and additionally shifts the photocurrent onset potentialto more negative values. [259]

An important breakthrough in electrocatalysis of water oxidationwas done by Nocera et al., who have recently developed a new typeof oxygen-evolving catalyst based on CoOx clusters that can be elec-trocatalytically deposited on the electrode surfaces from phosphate-buffered solutions containing Co2+ ions (“Co-Pi”). Notably, very highcatalytic activity was demonstrated: current density of 1 mA cm-2 atlow overpotential of η = 0.4 V and O2 formation with faradaic effi-ciency close to 100 %. [251]

The catalytic film was formed by oxidation of CoII to CoIII. The lat-ter one is sparingly soluble in the phosphate solution, which resultsin amorphous CoIII oxide/hydroxide participate. Precise structureof the catalyst is still subject of intensive research, nevertheless, oc-currence of edge-sharing CoO6 octahedra, with protonated/deproto-nated µ-O bridging atoms is very likely. [260–264] Phosphate ions mightbe present as terminal, but not bridging ligands. This structural motifcomprises characteristic, possibly distorted, cubane-like units, exhibit-ing therefore, to a certain extend, similarity to the CaMn4O5 center inPSII. [260–266]

Detailed electrochemical, [239] X-ray absorption near-edge structure(XANES), [260] and EPR [267] investigations have suggested that the wa-ter oxidation catalytic action of Co-Pi is associated with oxidationfrom CoIII to CoIV species, which leads to the formation of high-valentCoIV-oxo intermediates.

Discovery and development of Co-Pi electrocatalyst was a signifi-cant step towards meeting the requirements set for system capableto use energy input to generate oxygen and hydrogen. This uniquecatalyst is of particular interest, because:

1. forms in-situ under benign conditions on e.g. ITO or FTO; [240,251,266,268,269]

2. it is manufacturable in controllable way; can be easily depositedfrom proton-accepting buffers: phosphate (Co-Pi), borate (Co-Bi), methylphosphonate (Co-MePi), and fluoride; [240,251,266,268–270]

3. it consists of inexpensive, earth-abundant materials; [269]

5.1 introduction 69

4. operates with high activity under mild conditions – room tem-perature, neutral pH,1 and atmospheric pressure; [240,266,269]

5. its structure mimics the OEC in PS II; [271]

6. has self-healing mechanism allowing to continuous work overlong time; [266,272]

7. works also in impure, waste, and salt natural water; [240,268,269,271,273]

8. is selective – no Cl– oxidation to chlorine was observed; [269]

9. is a basis on which the so-called “artificial leaf” was built (thework of Nocera et al. [74,75] – see p. 15);

10. can be coupled to different charge-separating materials.

Much of the recent research on photoanodes for water oxidationtended to focus on the last point, that is on interfacing Co-Pi electro-catalyst with charge-separating materials (semiconductors), thereforeit will be discussed in more details.

It has been shown that cobalt-based oxygen evolving co-catalystcan enhance the photooxidation kinetics at many different semicon-ductors including: α-Fe2O3, [193,274–276] WO3, [207], ZnO, [277,278] TiO2, [279]

TaON, [204] BiVO4, [82,83,259,280–283] SiO2-BiVO4, [284] and Si-based pn-jun-ctions. [75,285,286] Several studies investigating dynamics of photogen-erated charge carriers in α-Fe2O3 photoanodes with deposited Co-Pi have been carried out. Transient absorption spectroscopy revealedthat addition of Co-Pi can increase the lifetime of photogeneratedholes by more than three orders of magnitude. In other words, ithas been demonstrated that formation of the inorganic heterojunc-tion between the co-catalyst and α-Fe2O3 results in retardation of re-combination. Moreover, the second major finding was that the pho-tocurrent onset was shifted to more negative potentials by ca. 0.2 Vfor α-Fe2O3/Co-Pi in comparison to bare α-Fe2O3 photoanode. [193] Inthe case of W-doped BiVO4 with photodeposited Co-Pi the shift ofthe photocurrent onset even by ca. 0.4 V was demonstrated. [83] Thesephenomenological effects are obvious, but difficulties arise when at-tempts are made to understand the microscopic role of co-catalystslike Co-Pi. [287] Traditional view assumes that photogenerated holesare rapidly extracted from VB and accumulated by a co-catalyst, therebyreducing the electron-hole surface recombination. Consequently, OERoccurs on the co-catalyst surface which can mediate four-hole trans-fer to water molecules more efficiently than the semiconductor itself.This view was confirmed for α-Fe2O3/Co-Pi by transients photocur-rent and impedance spectroscopy. [288] On the other hand, there is anexperimental evidence suggesting that in the α-Fe2O3/Co-Pi compos-ite photoanode, the Co-Pi “co-catalyst” does not catalyze water oxida-tion, but rather increases the band bending in α-Fe2O3 and improves

1 In acidic condition (pH < 3) H2O2 is the main product of oxidation. [240]


thus the charge separation. In this view, the role of Co-Pi is limitedto increasing the probability of the photogenerated hole to reach thesurface of α-Fe2O3 on which OER takes place. [289] Therefore, furtherresearch regarding coupling of Co-Pi with semiconducting materialsand clarification the of the role of the co-catalyst would be a greathelp in development of efficient photoanodes. [287]

In the next Sections results of investigation of TiO2-PH hybrid pho-toanodes with different cobalt-based co-catalysts will be presented.

5.2 photodeposition of cobalt phosphate

5.2.1 Experimental

0.1 M phosphate buffer (Pi) pH = 7 was prepared by dissolving 5.84 g(42.29 mmol) of NaH2PO4 ·H2O and 10.27 g (57.70 mmol) of Na2HPO4

· 2 H2O in 1 L of water. The photodeposition of Co-Pi was performedin phosphate-buffered solution containing 0.3 mmol dm-3 Co2+. Tothe photoelectrochemical cell containing 40 mL of Pi buffer, 0.5 mLof Co2+ solution, prepared by dissolution of 0.0707 g (0.243 mmol)of Co(NO3)2 · 6 H2O in 10 mL of water, was added directly beforemeasurement.

Prior to photodeposition, the TiO2-PH photoelectrodes were pre-pared on ITO glass, according to procedure described in Section 3.3.Dissolved oxygen measurements were conducted without stirring ac-cording to method described in Section 3.1.2.2.


Preliminary investigations of the Co-Pi co-catalyst deposition on theTiO2-PH photoelectrodes have shown that formation of Co-Pi leadsto brownish coloration of the TiO2-PH electrode (Figure 5.1), and thatthe amount of Co-Pi deposited must be fine-tuned to avoid blockingof light absorption by the co-catalyst.

Figure 5.1: TiO2-PH photoanodes after (on the right) and before (on the left)photodeposition of Co-Pi.

As already noted above, the Co-Pi co-catalyst was loaded on theTiO2-PH photoelectrodes photoelectrochemically to ensure that Co-Pi

5.2 photodeposition of cobalt phosphate 71

Figure 5.2: Photocurrent transients recorded at a TiO2-PH photoelectrodeunder irradiation with monochromatic visible light (λ = 450 nm)at 0.4 V vs. Ag/AgCl in a phosphate buffer (0.1 M; pH 7) without(a) and with (b) addition of Co2+ (0.3 mmol dm-3).

was formed preferentially at sites with high concentration of photo-generated holes, and to be able to control the progress of depositionby observing the photocurrent changes during photodeposition.

During irradiation with monochromatic visible light (λ = 450 nm)in the absence of cobalt ions (Figure 5.2), the TiO2-PH photoanodeshows a spike-like behavior typical of photocurrent generation pro-cesses associated with surface recombination. [237,238] After the initial

Figure 5.3: EDX analysis of a TiO2-PH electrode loaded with Co-Pi. EDXwas measured on four different spots.

rise after switching on the light, a rapid decrease of photocurrent isobserved due to accumulation of photogenerated holes in the sur-


face polyheptazine layer, which increases the rate of their recombi-nation with electrons in the conduction band of TiO2. [44,45,290] WhenCo2+ ions are present in the buffered solution a similar behavior isobserved in the initial phase. However, after several minutes the pho-tocurrent starts to increase and the photoelectrode turns brownish.The rise of photocurrent can be ascribed to formation of Co-Pi atthe surface of TiO2-PH, which leads to enhanced kinetics of wateroxidation. After about 10 min the photocurrent leveled off and thephotodeposition was interrupted. It is noted that during prolonged ir-radiation, particularly in solutions with higher concentrations of Co2+

ions than used in the present case (0.3 mM), the photocurrent tendedto decrease again, presumably due to the build-up of larger amountsof Co-Pi. The energy-dispersive X-ray spectroscopy (EDX) elementalanalysis performed on four different spots revealed a homogeneousdistribution of a relatively small amount of Co-Pi with the cobalt con-tent of approximately 0.9 wt. % (Figure 5.3).

Figure 5.4: Oxygen evolution measured in a phosphate buffer (pH 7) at 0.5 Vvs. Ag/AgCl (1.12 V vs. RHE) during prolonged irradiation bypolychromatic visible light (λ > 420 nm). The irradiated area was0.5 cm2, and the volume of the anode compartment was 5 mL.

Notably, irradiation of a Co-Pi-loaded TiO2-PH photoelectrode withvisible light (λ > 420 nm) led to evolution of dioxygen (Figure 5.4).TiO2-PH photoelectrodes without a co-catalyst exhibited no oxygenevolution, as well as pristine TiO2 photoelectrodes modified with Co-Pi by irradiation with UV light (λ = 350 nm) under otherwise identicalconditions.

As compared to a pristine TiO2 electrode modified with Co-Pi (Fig-ure 5.5 a), the photocurrent transients recorded on Co-Pi loaded TiO2-PH electrodes at different wavelengths (Figure 5.5 b) still reveal a morepronounced spike-like behavior. This indicates that the kinetics of pri-mary recombination processes occurring in TiO2-PH is still signifi-

5.2 photodeposition of cobalt phosphate 73

Figure 5.5: Photocurrent transients (without correction for the changes in ir-radiation intensity) under intermittent irradiation (5 s light/10 sdark) at different wavelengths recorded in a phosphate buffer(pH 7) at 0.5 V vs. Ag/AgCl (1.12 V vs. RHE) on TiO2 (a) andTiO2–PH (b) photoelectrodes loaded with Co-Pi. c) The corre-sponding photoaction spectra (IPCE versus λ plots).

cantly faster than the Co-Pi-assisted hole transfer to water molecules.Nevertheless, a significant portion of photogenerated holes is ableto escape recombination, as evidenced not only by the visible-light-driven evolution of oxygen but also by relatively significant valuesof incident photon-to-current efficiency (IPCE) in the visible region(Figure 5.5 c). The photoaction spectra of Co-Pi-loaded TiO2 and TiO2-PH correlate well with their respective optical absorption properties


shown in Figure 1.3. This proves that the interfacial charge-transferabsorption in TiO2-PH hybrids leads to efficient injection of electronsinto the conduction band of TiO2 whereby the holes are generatedin the polyheptazine layer and can take part in the multielectron-transfer reactions catalyzed by Co-Pi.

Figure 5.6: Photocurrent measured at a) TiO2–PH photoelectrodes with andwithout Co-Pi and b) pristine TiO2 under intermittent monochro-matic (λ = 450 nm) irradiation (5 s light, 5 s dark) during a ca-thodic potential sweep (5 mV s-1) in a deaerated phosphate buffer(pH 7).

It should be noted that a moderate positive bias is needed to en-hance charge separation and observe photocurrents. Figure 5.6 a)shows that under visible-light irradiation (λ = 450 nm) the photocur-rent onset potential is at + 0.1 V vs. RHE for TiO2-PH electrodesboth with and without Co-Pi. This is practically the same photocur-rent onset potential as in case of pristine TiO2 photoelectrodes un-der UV-light illumination, which confirms that the potential of photo-generated electrons in TiO2-PH hybrids coincides with the potentialof the conduction band edge in TiO2 (Figure 5.6 b). Moreover, two

5.3 mixtures of metal oxides 75

important points are noteworthy when discussing the potential de-pendence of the photocurrent in Figure 5.6 a). First, as compared topristine TiO2 (Figure 5.6 b), at TiO2-PH a relatively more pronouncedrecombination is observed in the vicinity of the photocurrent onsetpotential. Second, the situation is significantly improved after the de-position of Co-Pi. This clearly demonstrates that introduction of effi-cient catalytic sites for water oxidation into the polyheptazine layercan, at least partially, reduce the undesirable accumulation of holesand alleviate the need for positive electric bias. It should be also notedthat reasonable photocurrents are still observed in the potential rangefrom + 0.6 to + 0.8 V vs. RHE where, for example, the α-Fe2O3 pho-toanodes show no photocurrent response. [96,193,198]

5.3 mixtures of metal oxides

The successful loading of TiO2-PH photoanode with cobalt oxide/hy-droxide clusters (Co-Pi) motivated us to investigate other Co-basedco-catalysts: spinel-type cobalt(II,III) oxide (Co3O4) and nickel cobaltite(NiCo2O4), and ilmenite-type cobalt titanate (CoTiO3). These metaloxides were not introduced by the photodeposition method, but werefirst synthesized, then mixed with TiO2, and after preparing the elec-trodes modified with polyheptazine.

5.3.1 Cobalt(II,III) Oxide (Co3O4)

5.3.1.1 Experimental

Photoelectrodes with cobalt(II,III) oxide as a co-catalyst were manu-factured using two different methods:

method 1 : In the first stage, cobalt(II, III) oxide powders were pre-pared using four different protocols reported in literature:

a. Hydrothermal method: 0.50 g (4.2 mmol) of Co(CH3COO)2 · 4 H2Owas dissolved in 25.0 mL of water, and 1.9 ml (29 mmol) of33 % ammonia solution (NH3 · H2O) was added under strongstirring. The mixture was stirred for 10 min to form a homo-geneous dark brown-gray suspension. After this, the suspen-sion was transferred into a 50.0 ml autoclave, sealed, and main-tained at 250 °C for 3 h, followed by natural cooling to roomtemperature. In the next step product was washed with watervia centrifugation–redispersion (3 × 5 min, 3300 rpm) and laterdried at 60 °C for 4 h. According to literature report, resultingblack powder contained Co3O4 nanoparticles with size of ap-prox. 40 nm. [291]

b. Hydrothermal method: likewise point A, but autoclave was main-tained at 70 °C for 3 h. According to literature report, resulting


black powder contained Co3O4 nanoparticles with size of ap-prox. 20 nm. [291]

c. Sol/gel method: 2.32 g (8.0 mmol) of Co(NO3)2 · 6 H2O and0.02 g (0.1 mmol) of citric acid (acting as a surfactant) was dis-solved in 50 mL of ethanol. Then, the solution was heated for2 h at 50 °C in the first step and for 2 h at 85 °C in the secondstep to allow hydrolysis and olation of Co2+. The precipitatewas dried at 80 °C and heated for 3 h at 500 °C in air, givingblack Co3O4 powder with particles size reported in literature tobe in range 70–520 nm with maximum for 120 nm. [292]

d. Co-precipitation method: to 50 mL of 0.75 M cobalt(II) nitrateaqueous solution, excess of sodium carbonate (Na2CO3) wasadded dropwise under vigorous stirring, resulting in dark bluecobalt(II) carbonate (CoCO3) precipitate. Subsequently, the sus-pension was stirred for 2 h, filtered using Büchner funnel (MN615 d 60 Macherey-Nagel filter), and dried at 80 °C for 3 h. Suc-cessive heating in air for 3 h at 500 °C resulted in black Co3O4

oxide powder, with particle size reported in literature to be inrange of 250–800 nm with maximum for 325 nm. [292]

In the second stage of method 1, anatase TiO2 (Hombikat) and oneof Co3O4 powders synthesized using methods A, B, C or D weremixed together. Typically, 15 wt. % solid solution of Co3O4 in TiO2

was prepared (265 mg of TiO2 and 45 mg of respective Co3O4; 300 mgin total) and mixed (slightly ground) in a mortar. 200 mg of mixed ox-ides powder was diluted in 1 mL of ethanol and sonicated for 15 min,smeared on ITO glass and dried at 100 ºC for 15 min. Afterwards,photoelectrodes were covered with a piece of aluminum foil (the mattside to the catalyst surface), then with a piece of normal glass, andpressed for three minutes at a pressure of 150 kg cm-2. Finally, elec-trodes were calcined at 450 ºC for 30 min and modified with polyhep-tazine (30 min, 425 ºC).

method 2 : 2 g of anatase TiO2 (Hombikat) and 100, 200 or 500 mgof Co(NO3)2 · 6 H2O was dissolved in 50 mL of water, fol-lowed by dropwise addition under strong stirring of 5, 10 or25 mL of 0.4 M sodium carbonate Na2CO3, respectively. Themixtures were stirred overnight, filtered, and dried at 80 °C for3 h. This procedure yielded gray powder which was used for thephotoanodes preparation according to the standard protocol de-scribed in Section 3.3; in brief: 200 mg of respective gray powderwas added to 1 mL of ethanol, sonicated for 15 min, smearedon ITO glass, dried at 100 °C for 15 min, pressed, calcined at450 °C for 30 min, and finally modified with polyheptazine at425 °C for 30 min.


Dissolved oxygen measurements were conducted according to Sec-tion 3.1.2.3 with stirring.

5.3.1.2 Results and discussion

Cobalt(II,III) oxide is a well-known electrocatalyst, [249,250] photocat-alyst, [157,255,256] and co-catalyst [258,259] for water photooxidation. InFigure 5.7 results of dissolved dioxygen concentration measurementsduring irradiation of TiO2/Co3O4-PH with the visible light (λ > 420nm) are presented. Four different sets of samples were prepared us-

Figure 5.7: Dissolved oxygen concentration measurements recorded onTiO2/Co3O4-PH photoanodes manufactured with different sortsof Co3O4 in agreement with method 1 protocol A, B, C, D(see Section 5.3.1.1) during irradiation with polychromatic visi-ble light (λ > 420 nm) at 0.5 V vs Ag/AgCl (1.12 vs. RHE) in0.1 M Pi buffer pH = 7. For protocol A three samples with 5, 15,25 % concentration of Co3O4 were measured, while for protocolB, C, and D only with 15 %.


Figure 5.8: Photocurrent density recorded on TiO2/Co3O4-PH manufac-tured in agreement with method 1, protocol A, B, C, D (see Sec-tion 5.3.1.1) with different sort of Co3O4 during irradiation withpolychromatic visible light (λ > 420 nm) at 0.5 V vs. Ag/AgCl(1.12 vs. RHE) in 0.1 M Pi buffer pH = 7. For protocol A threesamples with 5, 15, 25 % concentration of Co3O4 were measured,while for protocol B, C, and D only with 15 %.

ing method 1. For the first set of samples (Figure 5.7 A) measure-ments were done with different concentration in order to find optimalamount of Co3O4. Analogically to photodeposition of Co-Pi or IrOx,amount of the co-catalyst must be well-tuned in order to avoid light-


shielding effects. In the case of in-situ photodeposition experiments,it was achieved by varying time, potential, concentration of metal ionsin the solution, and wavelength, whereas in case of metal oxides mix-tures the weight of added co-catalyst has to be optimized. It is in linewith previous studies, which have noted importance of amount of theco-catalyst: Fe2O3 + Co3O4 and graphitic carbon nitrite + Co3O4 sys-tems exhibit optimal activity when 5 wt. % and 3 wt. %, respectively,of the co-catalyst is present in the mixture. [258,293]

Figure 5.9: a) SEM micrographs of TiO2/Co3O4-PH photoanode preparedaccording to method 1 d). b) EDX analysis of same photoanodedone on the representative spot (see footnote on p. 57).

Turning now to the experimental results shown in Figure 5.7 A)– dissolved oxygen measurements of three samples with concentra-tions equal to 5, 15, and 25 %, illustrate the above-described trade-offbetween electrocatalysis and light attenuation by a co-catalyst. Thesample containing 5 % of Co3O4, regardless of high photocurrent (Fig-ure 5.8 A), does not show any dioxygen evolution. If concentrationof Co3O4 is raised to 15 %, at the almost same photocurrent values,dioxygen evolution can be evidently observed. However, if even moreof the co-catalyst is present in the mixture, photocurrent and amountof evolved dioxygen drops remarkably, because too high amount ofCo3O4 blocks light absorption by TiO2-PH. It can be illustrated brieflyby the color of TiO2/Co3O4 photoanodes after modification with PH:5 % content of Co3O4 hold typical yellow color, 15 % content makeselectrode much darker, and eventually photoanode with 25 % of the


co-catalyst content is almost completely black. Consequently, for thenext sets of samples (B, C, D) only concentration of 15 % was investi-gated as the optimal one.

Figure 5.10: a) Photocurrent density and b) dissolved oxygen evolutionconcentration measured on TiO2/Co3O4-PH manufactured asstated in method 2 (Section 5.3.1.1) with different amount ofCo(NO3)2 · 6 H2O in starting mixture, during irradiation withpolychromatic visible light (λ > 420 nm) at 0.5 V vs Ag/AgCl(1.12 V vs. RHE) in 0.1 phosphate buffer pH = 7.

Cobalt(II,II) oxide powders used to manufacture four sets of sam-ples (A, B, C, D), whose properties were investigated in Figure 5.7and Figure 5.8 are believed to differ mostly by size of particles. Ac-cording to literature reports, samples A should consist of nanocubes40 nm size, sample B of nanoparticles 20 nm size, samples C of muchbigger particles, which should be 120 nm in size, and eventually sam-ple D 325 nm in size. The highest values of both photocurrent (0.15mA cm-2) and amount of dioxygen (40 µmol dm-3) produced duringirradiation were measured for set of samples denoted as D. It is some-


what surprising because this result differs from the previous researchdone in the field, which showed that the activity of Co3O4 increaseswith decreasing size of nanoparticles. [208,255,258,293,294]

The most active sample D prepared by method 1(Figure 5.7 D)has not only bigger particles than another powders, but they are alsopoorly distributed among TiO2 forming rather big islands (over 5 µm)as can be seen in Figure 5.9 a). On this account, method 2 wasdeveloped, which is similar to method 1, but precipitating agent(Na2CO3) was added under strong stirring directly to suspension ofTiO2 containing Co2+ ions. Such procedure results in much better dis-tribution of Co3O4 in TiO2. However, as can be seen in Figure 5.10both photocurrent and amounts of dioxygen for all samples preparedusing method 2 are much lower than for sample D prepared usingmethod 1. It can be therefore concluded, that for TiO2/Co3O4-PHphotoanodes rather big particles (300–800 nm) of Co3O4 which aregrouped in µm-sized domains are preferred.

5.3.2 Nickel Cobaltite (NiCo2O4)


Nickel cobaltite (NiCo2O4) was synthesized using a co-precipitationmethod reported in the literature. [295,296] In short, 50 mL of 2.0 MNaOH, was quickly added to 150 mL of the aqueous solution contain-ing 0.02 mol of Co(NO3)2 and 0.01 mol of Ni(NO3)2 and then stronglystirred for 30 min. Afterwards, suspension was filtered using a Büch-ner funnel and thoroughly washed by distilled water. The participatewas dried at 80 °C overnight, followed by grinding in a mortar. Thedark green precursor powder was calcined at 300 °C for 3h yieldingblack NiCo2O4 powder. Both greenish “precursor” (before calcina-tion) and black “actual” NiCo2O3 (after calcination) powders wereused for photoelectrodes preparation.

The photoelectrodes were manufactured on ITO glass, accordingto the protocol developed for mixed oxides. Firstly, a commercialanatase TiO2 powder (Hombikat) was mixed with different amountsof NiCo2O4 “precursor” or “actual” NiCo2O4 powder. Samples withthree weight ratios were prepared: 10, 20 and 30 wt. % for both theprecursor of NiCo2O4 and the actual NiCo2O4 powder. TiO2 andNiCo2O4 powders were ground together in a mortar, and then 200 mgof mixture was weighted and dissolved in 1 mL of ethanol. Subse-quently, the suspension was sonicated for 15 min, smeared on theITO glass and dried at 100 ºC for 15 min. In the following step, pho-toelectrodes were covered (the matt side to the catalyst surface) witha piece of aluminum foil, and a normal glass and pressed for threeminutes at the pressure of 150 kg cm-2. Lastly, the electrodes were cal-


cined at 450 ºC for 30 min and modified with polyheptazine (30 min,425 ºC).

Dissolved oxygen measurements were conducted according to de-scription in Section 3.1.2 without stirring.


Nickel cobaltite is a mixed valence oxide with spinel structure. It canbe obtained by the co-precipitation method from solutions contain-ing stoichiometric ratio of Ni2+ and Co2+ ions. NiCo2O4 was investi-

Figure 5.11: a) Photocurrent and b) dissolved dioxygen concentration mea-surements under irradiation with polychromatic visible light(λ > 420 nm) at 0.5 V vs. Ag/AgCl (1.12 V vs. RHE) in a phos-phate buffer (0.1 M; pH 7) for photoanodes manufactured bymixing TiO2 with three different (10, 20, 30 wt. %) concentra-tion of NiCo2O4 “precursor” and subsequent modification withpolyheptazine.


gated as a electrocatalyst for water oxidation, but was also reportedto oxidize water under illumination if photosynthesizer and electronacceptor are present in the mixture. [157]

Figure 5.11 presents photocurrents and dissolved oxygen measure-ments results for TiO2/NiCo2O4 precursor modified with polyhep-tazine. The precursor was brucite-like mixture of Co(OH)2 and Ni(OH)2, [297]

which transforms after heat treatment to NiCo2O4. Three weight ra-tios: 10, 20, and 30 % of precursor in the mixture with TiO2 were inves-tigated. The set of samples in Figure 5.11 was denoted as “precursor”because the precursor, that is mixture of Co(OH)2 and Ni(OH)2, wasdirectly mixed with TiO2, sonicated, smeared on ITO glass, pressed,and then calcined.

Figure 5.12: EDX elemental analysis of TiO2/NiCo2O4 (20 % precursor)modified with PH. Analysis was done on the representativespot (see footnote on p. 57) of the photoanode.

The transformation of the precursor to NiCo2O4 occurs just in thelast step, during heating of electrodes in tube furnace at 450 °C. Thehighest and most stable photocurrents, validated by highest amountsof O2 generated during visible light illumination, were exhibited bythe sample with 20 % concentration of precursor. EDX elemental anal-ysis (Figure 5.12) performed on sample prepared in such way re-vealed almost stoichiometric amounts of Ni and Co, equal to 3.40 %and 6.35 %, respectively (after correcting for Si, In, and Ca content).

The second set of samples was manufactured by mixing “actual”NiCo2O4 with TiO2 followed by standard procedure. It means thatNiCo2O4 was first separately synthesized by heating the precursorat 300 °C for 3 h, and then mixed with TiO2. To asses activity inwater photooxidation of photoanodes prepared in this way, a seriesof photocurrent and dissolve oxygen measurements was conducted(Figure 5.13). Although values of photocurrents and amounts of gen-


Figure 5.13: a) Photocurrent and b) dissolved oxygen concentration mea-surements recorded under irradiation with polychromatic vis-ible light (λ > 420 nm) of hybrid photoanodes manufactured bymixing TiO2 with three different (10, 20, 30 wt. %) concentrationof “actual” NiCo2O4 and subsequent modification with poly-heptazine. Photoelectrodes were biased at 0.5 V vs. Ag/AgCl(1.12 V vs. RHE); phosphate buffer (0.1 M; pH 7) was used asan electrolyte.

erated dioxygen achieved for the best concentration of 30 % are signif-icant ( ∼ 0.08 mA cm-2 and 26 µmol dm-3, respectively), they are lowerthan those for optimal 20 % of “precursor” sample (∼ 0.10 mA cm-2

and 31 µmol dm-3, respectively). Superior behavior of “precursor”samples may be explained by the transformation to final NiCo2O4

during heating together with TiO2. This should result in the more ef-ficient TiO2/NiCo2O4 junction and the better electric contact with theITO substrate due to stronger sintering between particles of both com-ponents. Consequently, TiO2/NiCo2O4 20 % precursor samples were


Figure 5.14: A) Photocurrent and B) dissolved oxygen concentration mea-sured under prolongate irradiation with polychromatic visiblelight (a, b, c - different cut-off filters) of hybrid photoanode man-ufactured by mixing TiO2 with 20 wt. % of NiCo2O4 “precursor”and subsequent modification with polyheptazine. The unmodi-fied TiO2/NiCo2O4 20 wt. % photoanode irradiated with poly-chromatic visible light (d) λ > 420) was measured as a bench-mark. External bias was equal to 0.5 V vs. Ag/AgCl (1.12 Vvs. RHE); measurements were done in phosphate buffer (0.1 M;pH 7).

chosen in order to investigate long-term stability under 5 h illumi-nation with different ranges of the polychromatic light (Figure 5.14).Three cut-off filters were used: a) λ > 400 nm, b) λ > 420 nm, andc) λ > 455 nm to provide irradiation with different fragments of thevisible light. These results are interesting in several aspects. To begin


with, apparently operational time of TiO2/NiCo2O4-PH photoanodesis limited to ca. 3 h. After this time, the photoelectrocatalytic systemdeactivates. Furthermore, a significant drop (50 % between the secondand the third hour of irradiation) of the photocurrent (Figure 5.14 A)can be observed for samples irradiated with cut-off filters 400 nmand 420 nm. Interestingly, such drop of photocurrent is not observedin the case of irradiation with c) λ > 455 nm light. This differenceis most probably due to the higher accumulation of the photoholes,which can damage polyheptazine layer or its bond to TiO2 what re-sult in observed instability during irradiation with the broader spec-trum of the light: λ > 400 nm and λ > 420 nm. In the case of c) λ >455 nm, photocurrent is stable, because photoanode can better matchsmaller incoming light flux to water oxidation reaction and/or thehigh-energy UV photons which can destroy the TiO2-PH bond aremissing. In other words, in the long-time perspective, the NiCo2O4

co-catalyst is not efficient enough to keep up hole transfer from ab-sorber to water molecules with higher light flux (cut-off filters 400and 420 nm) or it just destroys PH layer. Finally, as expected, unmod-ified TiO2/NiCo2O4 benchmark sample does not exhibit any activityunder irradiation with visible light (λ > 420 nm).

5.3.3 Cobalt Titanate (CoTiO3)


Ilmenite-type cobalt titanate (CoTiO3) was synthesized according tothe procedure reported in literature. [298] Briefly, 7.1 g (0.025 mol) oftitanium isopropoxide Ti(OPr)4 and 6.3 g (0.025 mol) of cobalt acetatetetrahydrate Co(CH3COO)2 · 4 H2O (1:1 molar ratio) were dissolvedin 40 mL of 2-methoxyethanol. The solution was stirred overnightto produce deep purple, sol-like precursor. Afterwards, the solutionwas heated in a porcelain bowl for one hour at 150 °C and thenpost-heated for two hours at 550 °C yielding dark violet CoTiO3

powder. X-ray diffraction spectrum of the synthesized powder is pre-sented in Figure 5.15. It is consistent with those of literature stud-ies, [298] and clearly shows that the powder is mostly composed of theCoTiO3 phase. The peaks labeled with form pattern characteris-tic to CoTiO3, nevertheless, two minor peaks, labeled with c can beascribe to either TiO2 or Co3O4 phases.

The photoelectrodes were prepared on the FTO glass, using thestandard preparation protocol for mixed oxides, described previouslyfor Co3O4 and NiCo2O4. Initially, a commercial anatase TiO2 (Hom-bikat) was mixed with different amounts of CoTiO3. For example, toprepare samples with 7.5, 10.0 and 12.5 wt. % of CoTiO3, 23, 30 and38 mg of CoTiO3 were weighted and mixed with 277, 270 and 262 mgof TiO2 (300 mg in total), respectively. Subsequently, the two pow-


Figure 5.15: X-ray diffraction (XRD) pattern for cobalt titanate (CoTiO3). Thepeaks labeled with form pattern characteristic to CoTiO3; thepeaks labeled with c can be ascribe to either TiO2 or Co3O4phases.

ders were ground in a mortar and then 200 mg of the mixture wasweighted and dissolved in 1 mL of ethanol. Later, the suspension wassonicated for 15 min, smeared on the FTO glass, and dried at 100 ºCfor 15 min. In the next step, photoelectrodes were covered (the mattside to the catalyst surface) with a piece of aluminum foil and a nor-mal glass piece, and pressed for three minutes at the pressure of 150kg cm-2. Finally, photoelectrodes were calcined at 450 ºC for 30 minand modified with polyheptazine (30 min, 425 ºC).

Dissolved oxygen measurements were conducted with stirring ac-cording to description in Section 3.1.2.3.


Figure 5.16 compares photocurrents and dissolve dioxygen measure-ments results for photoanodes having of 5, 10 or 20 % weight contentof CoTiO3 in the mixture with TiO2 after modification with polyhep-tazine. These initial measurements reveled that 20 % of CoTiO3 inthe mixture is too high because it blocks the light absorption, as ev-idenced by photocurrents below 50 µA cm-2 and a raise of dioxygenconcentration by only 4 µmol dm-3 during 1 h irradiation with thevisible light (λ > 420 nm). On the other hand, sample with 5 % of wt.content of CoTiO3 in the mixture exhibited 2 times higher photocur-rents and dioxygen concentration raised by 13 µmol dm-3. However,it is still lower than optimal amount of co-catalyst, i. e. 10 %, whichexhibited a raise of dioxygen concentration by 16 µmol dm-3. In the


Figure 5.16: Initial measurements of a) photocurrent and b) dissolved dioxy-gen evolved during irradiation with polychromatic visible light(λ > 420 nm) of TiO2/CoTiO3-PH photoelectrodes (three differ-ent wt. % of CoTiO3) biased at 0.5 V vs. Ag/AgCl (1.12 V vs.RHE) and immersed in phosphate buffer (0.1 M; pH 7).

next measurements additional samples containing: 7.5 % and 12.5 %of CoTiO3 were measured. Results are shown in Figure 5.17. It can beinferred, that the best ratio of CoTiO3 is 12.5 %, because such photoan-ode released the highest amount of dioxygen (rise by 19 µmol dm-3),whereas photocurrents were quite similar for all three samples. Ad-ditionally, two samples which were not modified with polyheptazine(TiO2/CoTiO3) were measured for comparison. As expected, they didnot exhibit any activity under visible light irradiation (λ > 420 nm).

Best capability of samples containing 12.5 % of CoTiO3 was con-firmed by IPCE measurements shown in Figure 5.18. Three samplesfor every concentration were measured in order to calculate the stan-dard deviation and draw the error bars. The IPCE values at 400 nm


Figure 5.17: a) Photocurrent and b) dissolved oxygen concentrationrecorded at TiO2/CoTiO3-PH and TiO2/CoTiO3 photoelec-trodes (three different content (wt. %) of CoTiO3 – 7.5 %, 10.0 %,12.5 % ) under irradiation with polychromatic visible light (λ >420 nm) at 0.5 V vs. Ag/AgCl in a phosphate buffer (0.1 M;pH 7).

for TiO2/CoTiO3-PH photoanodes with 12.5 % weight ratio of the co-catalyst were equal to approx. 4 %, hence, they are close to those ofTiO2-PH + Co-Pi (Figure 5.5). Photoanodes with lower content of theco-catalyst showed smaller IPCE values: 3.2 % and 2.7 % for photoan-odes containing 10.0 % and 7.5 % of CoTiO3, respectively. Similarlyto previously described materials, modification with polyheptazineis obligatory to observe response in the visible light range – IPCEvalues at 400 nm for all TiO2/CoTiO3 samples were close to zero.


Figure 5.18: Photoaction spectra (IPCE vs. λ) of a TiO2/CoTiO3-PH andTiO2/CoTiO3 photoelectrodes with three different content(wt. %) of CoTiO3 measured at 0.5 V vs. Ag/AgCl (1.12 V vs.RHE) in a phosphate buffer (0.1 M; pH 7).

Interestingly, the TiO2/CoTiO3-PH 12.5 % photoanode, the best inthe short-time (≤ 1 h) perspective is not at all best in the long-timeexperiments (> 2 h). It can be seen in Figure 5.19 which comparesTiO2/CoTiO3-PH photoanodes containing 12.5 % and 10.0 % of theco-catalyst in terms of photocurrent and amount of evolved dioxy-


Figure 5.19: a) Photocurrent and b) dissolved oxygen concentrationrecorded at TiO2/CoTiO3-PH (10.0 and 12.5 wt. % of CoTiO3)and TiO2/CoTiO3 (10.0 wt. % of CoTiO3) during prolonged ir-radiation (5 h) with polychromatic visible light (λ > 420 nm) at0.5 V vs. Ag/AgCl in a phosphate buffer (0.1 M; pH 7).

gen during prolonged (5 h) irradiation with the visible light (λ >420 nm). It is apparent from this graph, that the photoanode with10.0 % of CoTiO3 was much more stable in the long-time perspective(> 2 h). Three more issues are noteworthy. Firstly, the concentrationof dioxygen measured at TiO2/CoTiO3-PH photoanode with 10.0 %of the co-catalyst reached level of almost 70 µmol dm-3. It is a ratherhigh value – two times higher than the one for NiCo2O4 (∼ 35 µmoldm-3), and only approximately three times lower than the standardconcentration of dissolved O2 in water remaining in equilibrium withair (∼ 200 µmol dm-3). Secondly, as compared with TiO2/NiCo2O4-PH 20 % “precursor”, operational time for TiO2/CoTiO3-PH 10.0 % islonger. It can be seen that generation of dioxygen can be observedeven after 4 h of irradiation, albeit dioxygen generation rate (slop of


the plot) decreases. Thirdly, as expected, the unmodified TiO2/CoTiO3

(10.0 %) photoanode does not show any O2 evolution. However, thisresult confirms, that the experimental setup is very tight and even insuch long period of time there is no leakage of atmospheric dioxygen.

Figure 5.20 presents SEM micrographs of TiO2/CoTiO3-PH (10.0 %)photoanode. Islands of CoTiO3 incorporated to TiO2 are clearly visi-ble. EDX elemental analysis done on the representative spot (see foot-note on p. 57) revealed the presence of Co in amount of 3.34 % (aftercorrection for Si, In, Ca).

Figure 5.20: a) SEM micrographs and b) EDX analysis of TiO2/CoTiO3-PH(10 %) photoanode.

Figure 5.21 compares photocurrent transients measured during ca-thodic potential sweep at (a, b) TiO2/CoTiO3-PH and c) TiO2/CoTiO3.This results are significant in at least two major respects. Firstly, thephotocurrent onset for samples modified with polyheptazine (Fig-ure 5.21 a) is shifted to positive potentials and exhibits sensitivity toapplied potentials in contrast to unmodified samples (Figure 5.21 c).This behavior originates probably from enhanced primary recombi-nation processes (back electron transfer after electron excitation formPH to TiO2; Figure 1.2) unless a sufficient electric bias is applied. Sec-ondly, it can be clearly seen in Figure 5.21 b), that the spike-shapecharacter of transients decreases with increasing amount of the co-catalyst. Assuming, that the spike-shape character of photocurrentsarises from the enhanced surface recombination, it could be conceiv-ably hypothesized that higher amount of the co-catalyst results infaster extraction of the photogenerated holes from the polyheptazine


Figure 5.21: Photocurrents measured at (a, b) TiO2/CoTiO3-PH and c)TiO2/CoTiO3 photoelectrodes with different content (wt. %) ofCoTiO3, recorded under intermittent monochromatic irradia-tion (5 s light, 5 s dark) during a cathodic potential sweep(5 mV s-1) in a deaerated phosphate buffer (0.1 M; pH 7).

layer and their faster injection to the electrolyte. On the other hand,photocurrents for highest amount of CoTiO3 (12.5 %) are lower than


for medium one (10.0 %) probably because of blockage of the lightabsorption.

5.4 conclusion

In the first part of this Chapter we have demonstrated that TiO2-PHhybrid photoelectrodes can be loaded with small amounts of a cobaltoxide-based oxygen-evolving co-catalyst (Co-Pi) by irradiation withvisible light. The thus fabricated photoelectrodes, based solely onearth-abundant elements, exhibit visible (λ > 420 nm) light drivencomplete photooxidation of water to oxygen under a relatively mod-erate external electric bias.

In the second part of this Chapter we have demonstrated, that Co-based co-catalysts can be introduced also by mixing cobalt-containingoxides with pristine TiO2 and subsequent modification with polyhep-tazine. In this way three mixtures of the following oxides: Co3O4,NiCo2O4 and CoTiO3 with TiO2 were prepared. Amount of co-catalystswere optimized in order to ensure the maximum light absorption andhighest possible reaction rates at the same time. The most importantfinding to emerge from this study was that such mixtures of oxidesoxidize water with high rates under photoelectrochemical conditionssimilarly to photoelectrodes with photodeposited cobalt oxide-basedoxygen-evolving co-catalyst (Co-Pi).

Taken together, results presented in this Chapter prove ability ofTiO2-PH photoanodes, successfully loaded with cheap and earth-abun-dant Co-based co-catalysts, to drive water photooxidation. Apart fromits crucial role in the complete oxidation of water to oxygen, the co-catalyst was found to reduce significantly the recombination of pho-togenerated charges, particularly at low bias potentials, by diminish-ing partially the accumulation of holes in the surface polyheptazinelayer. This is highly important in view of reducing the need for ex-ternal electric bias at hybrid photoanodes for water splitting, whichcan apparently be achieved if further and more effective ways of in-troducing catalytic sites for water oxidation into the surface structureof the hybrid photoanodes are developed.

6N I C K E L - B A S E D C O - C ATA LY S T S

6.1 introduction

A considerable amount of literature has been published on water de-composition using powder photocatalysts, e. g. Ta2O5, [299,300] K4Nb6O17, [301]

Sr2Ta2O7, [302] La:NaTaO3, [303] and SrTiO3[304–307] with NiOx (0 < x <1)

as a co-catalyst. In these studies, formation of core/shell (Ni/NiO)structure, which act as water reduction center (while water oxidationoccurs on photocatalyst surface) was assumed. [308] This concept hasbeen recently challenged by studies revising role of NiOx in photocat-alytic process. [309] It has been demonstrated on the basis of surfacephotovoltage, photoelectrochemical, and H2/O2 evolution measure-ments, that NiOx is not involved in water reduction. It is more likely,that vast amount of Ni(0) and NiO particles are distributed separatelyon the SrTiO3 surface and, even more importantly, only Ni(0) reducesprotons, whereas NiO acts as a water oxidation catalyst.

Another important finding in this context was that an amorphousnickel oxide/hydroxide, electrodeposited on ITO glass form aqueousborate buffer (Bi) solution (pH = 9.2) containing Ni2+ cations, can cat-alyze oxygen evolution reaction. [310] Nickel borate (Ni-Bi) films re-sembles cobalt-phospate (Co-Pi; see Section 5.1) in terms of prepara-tion route, possibility of precise thickness control, high catalytic ac-tivity, long-term stability, mild operation conditions, and even atomicstructure. [273] Nickel centers in as-deposited Ni-Bi electrocatalyst ex-ist mainly as Ni(III). Subsequent activation of catalyst by anodiza-tion1 results in almost 3 orders of magnitude improvement in cat-alytic activity, attributed to an increase in number of bis-oxo/hydroxobridged Ni(IV) centers (average oxidation state ∼ 3.6) and change inshort-range structure, while long-range order remains amorphous, asmentioned above. [311] The Bi solution keeps bulk and local pH con-stant during O2 evolution and borate anion itself plays crucial rolein catalytic mechanism acting as a proton acceptor, similarly to phos-phate in case of Co-Pi. [312] Above all, the nickel-borate (Ni-Bi) thinelectrocatalyst films exhibit excellent activity of 1 mA cm-2 at overpo-tential of 0.42 V and very low Tafel slope of 58 mV per decade. [310]

Interestingly, nickel borate was recently successfully coupled toBiVO4 electrode. [313] The results of the study indicate that active layerof Ni-Bi can be photodeposited on the surface of BiVO4 from 0.1 mMBi buffer containing 1 mM Ni2+ under AM 1.5-light (400 mW cm-2) ir-

1 By application of 1.1 V, with correction to the ohmic loss, for 2.5 h in stirred 1.0 Mborate buffer pH = 9.2 [311]

95

96 nickel-based co-catalysts

radiation, without application of external potential. Significantly, pho-tocurrent generation observed for BiVO4/Ni-Bi photoanode was in-creased by factor of ca. 4 (from ca. 0.3 to 1.2 mA cm-2 at 1.23 V vs.RHE) in comparison to pristine BiVO4. Furthermore, photoelectro-chemical water oxidation under AM 1.5 light (4 sun) irradiation wasdemonstrated with almost 100 % efficiency and photocurrent of ca.1.8 mA cm-2 at 1.23 V vs. RHE, proving ability of Ni-Bi to trap photo-generated holes and drive their transfer to water molecules. Moreover,crucial role of borate (good proton acceptor) was confirmed by com-parison with measurements in nitrate (poor acceptor) which showeddiminished activity. [313]

For these reasons, and because of the fact that nickel is cheaper andmore abundant than cobalt, Ni-Bi is definitely an attractive co-catalystto be coupled with TiO2-PH photoelectrodes.

6.2 experimental

0.1 M potassium borate buffer electrolyte pH = 9.0 (Bi solution) wasprepared by dissolving 6.18 g of boric acid (H3BO3) and 3.00 g ofpotassium hydroxide in 1 dm3 of water.

The solution for photodeposition of Ni-Bi co-catalyst was preparedby method describe in literature. [310] In brief, 0.0232 g (0.08 mmol) ofNi(NO3)2 · 6 H2O was dissolved in 40 mL of water (0.2 mM Ni2+) andcombined with 40 mL of 0.1 M Bi buffer solution. Instantaneously af-ter addition of solution containing Ni2+cations, Bi buffer became tur-bid, probably because of Ni(OH)2↓ precipitation. The solution (withtotal volume of V = 80 mL) was filtered through an ashless filter paper(Macherey-Nagel MN640d 125 mm, 0.17 mm thick) and used directlyafter preparation.

TiO2-PH photoanodes were manufactured in the standard way (seeSection 3.3) on ITO glass. Dissolved oxygen measurements were con-ducted according to description in Section 3.1.2 with stirring. All mea-surements were done using borate buffer (Bi) pH = 9.0 as an elec-trolyte.

6.3 photodeposition of nickel borate

Similarly to formation of IrOx or CoOx (Co-Pi) catalytic layers, de-scribed in previous Chapters, Ni-Bi co-catalyst was formed at theTiO2-PH surface by photodeposition method. As mentioned in litera-ture review, catalytic layer of Ni-Bi was also photodeposited on BiVO4

photoanode. In that case, Ni-Bi co-catalyst was photodeposited un-der polychromatic light irradiation under purely photochemical con-ditions (without application of external bias). [313] In contrary, in thisstudy the external bias was used as one of the factors, apart from lightwavelength and time, to control Ni-Bi co-catalyst properties by gov-

6.3 photodeposition of nickel borate 97

Figure 6.1: Photocurrent density during photodeposition of Ni-Bi co-catalyst on a) TiO2-PH photoanode (irradiation with monochro-matic visible light λ = 450 nm) and b) pristine TiO2 photoanode(irradiation with monochromatic UV light λ = 350 nm) from Bielectrolyte (pH = 9.0) containing Ni2+ cations (see Section 6.2).Photoelectrodes were biased at 0.5 V vs. Ag/AgCl (1.24 V vs.RHE). The electrolyte was not stirred.

erning the number of photogenerated holes able to oxidize Ni2+ ionsto form NiOx oxide/hydroxide layer at the photoelectrode’s surface.

Figure 6.2: EDX analysis of TiO2-PH + Ni-Bi.

Analogously to cases described in previous Chapters, irradiation ofthe photoanode in contact with solution containing Ni2+ cations ledto brown colorization due to the Ni-Bi formation. The photocurrentsrecored during photodeposition of Ni-Bi on (a) pristine TiO2 and (b)TiO2-PH with monochromatic (a) visible and (b) UV light are pre-


sented in Figure 6.1. In first two minutes of photodeposition after aninitial raise of photocurrent, fast decay is observed – this spike-likeshape can be attributed to fast surface recombination processes. Be-tween 2–14 min of photodeposition, when Ni-Bi is built-up, decreaseof the photocurrent can be also caused by insufficiently fast transportof Ni2+ from the bulk of the electrolyte to the surface of the pho-toanode, because it was not stirred. The photodeposition at 1.24 Vvs. RHE for 14 min yielded fine layer of Ni-Bi co-catalysts, as canbe seen in SEM picture in Figure 6.1. Nevertheless, further optimiza-tion of conditions during photodeposition (time, potential, stirring ofelectrolyte) can be beneficial.

EDX elemental analysis performed on representative spot (see foot-note on p. 57) confirmed presence of Ni on the photoelectrode surface(Figure 6.2). After correcting for Si, Ca, and In, weight content of Niamounts to ∼ 3.5 %.

Figure 6.3: Photoaction spectra (IPCE vs. λ plots) for (− •−) hybrid pho-toanode and (−−) pristine TiO2 with photodeposited Ni-basedco-catalyst, measured at 0.5 V vs. Ag/AgCl (1.24 V vs. RHE) in0.1 M borate buffer electrolyte pH = 9.0.

Figure 6.3 compares photoaction spectra (IPCE vs. wavelength) forpristine TiO2 and TiO2-PH hybrid photoanode with photodepositedNi-Bi co-catalyst. As expected, TiO2 modified with polyheptazine andloaded with Ni-Bi (TiO2-PH + Ni-Bi) exhibits remarkably better prop-erties in terms of both the values and the shape of the spectrum ascompared to unmodified TiO2 loaded with Ni-Bi. These data mustbe interpreted with caution because these measurements were donein solution with pH = 9 (1.24 vs. RHE), and thus cannot be directlycompared to values measured at pH = 7 (1.12 V vs. RHE) – like thosein Figure 5.5 (TiO2-PH + Co-Pi) or Figure 4.5 (TiO2-PH + IrOx/PD).However, they can be compared to those in Figure 8.1. Accordingly,

6.3 photodeposition of nickel borate 99

if values recorded at the same potential vs. RHE are compared, it canbe clearly seen that IPCE for Ni-Bi are lower by, at least, 1 %.

Figure 6.4: a) Photocurrent and b) dissolved oxygen evolution measured inborate buffer (pH = 9.0) during prolonged irradiation by poly-chromatic visible light (λ > 420 nm) of TiO2-PH and pristine TiO2with photodeposited Ni-based co-catalyst at 0.5 V vs. Ag/AgCl(1.24 V vs.RHE)

Importantly, TiO2-PH hybrid photoanodes with photodeposited Ni-Bi are able to oxidize water under visible light (λ > 420 nm) irradia-tion, as evidenced by dissolved oxygen measurements results shownin Figure 6.4. However, stable, but moderate photocurrent densities(90 µA cm-2) and rather small amounts of evolved dioxygen (13 µmoldm-3 after 1 h) are not satisfactory, especially if higher pH of the elec-trolyte (higher potential vs. RHE) is taken into account. Relativelysmall faradaic efficiency of FE = 16 % (if compared to for example FE


up to 45 % for IrOx/PD) suggests, that the reason for this is ratherimperfect coupling between the co-catalyst and TiO2-PH.

Figure 6.5: Photocurrent measured at a) TiO2-PH and b) TiO2 photoelec-trodes without and with photodeposited Ni-based co-catalystrecorded under intermittent monochromatic light irradiation (5 son, 5 s off, λ = 420 nm for TiO2-PH and λ = 350 nm for TiO2)during cathodic potential sweep (5 mV s-1) in deaerated boratebuffer (Bi) electrolyte solution (pH = 9).

Figure 6.5 shows photocurrents measured under intermittent irra-diation during cathodic potential sweep . There are three interestingpoints in this plots. Firstly, the “dark current” peaks at + 0.2 V vs.RHE in the case of both TiO2-PH + Ni-Bi and TiO2 + Ni-Bi are as-cribed to reduction of NiIII to NiII in the Ni-Bi co-catalyst. Secondly,a comparison of modified (TiO2-PH, TiO2-PH + Ni-Bix) and unmod-

6.4 conclusion 101

ified (TiO2, TiO2 + Ni-Bi) photoelectrodes clearly show that in thecase of modified samples photocurrents are more sensitive to appliedbias and photocurrent onset is shifted by 0.4 V. It is mainly becauseof primary recombination processes which prevail especially at lowvalues of applied bias. Thirdly, a comparison between TiO2-PH andTiO2-PH + Ni-Bi reveals slightly higher photocurrent at potentialsabove 1.2 V vs. RHE. It can be explained by blocking of light absorp-tion by the co-catalyst. [158] Nevertheless, a beneficial influence of theco-catalyst can be seen specifically at lower values of applied bias,where kinetic competition between surface recombination and trans-fer of holes to water molecules plays more important role. The lesspronounced spike-shape decay and higher photocurrents in range be-tween + 0.8 and + 1.2 V vs. RHE (inset in Figure 6.5 a) indicate, thataddition of Ni-Bi co-catalyst improves water oxidation kinetics, pre-venting thereby, at least partially, the hole accumulation and decreas-ing surface recombination rate.

6.4 conclusion

In conclusion, in this Chapter photooxidation of water under vis-ible light irradiation on TiO2-PH hybrid photoanode loaded withnickel-based oxygen evolving co-catalyst was demonstrated. The co-catalyst was introduced to TiO2-PH photoanode by the photodeposi-tion method, which was once again confirmed to be efficient way ofloading of small amounts of metal oxides onto photoactive semicon-ductor material. However, due to the unsatisfactory results (low con-centration of evolved oxygen), further studies with focus on tuningphotodeposition conditions are suggested in order to improve ratherweak coupling between Ni-Bi co-catalyst and TiO2-PH absorber. More-over, additional per-treatment step similar to anodization of electro-catalyst film, which was reported in literature [310] to be necessary toreach optimum activity by Ni-Bi electrocatalyst, may also improveperformance of TiO2-PH + Ni-Bi photoanodes.

7M A N G A N E S E - B A S E D C O - C ATA LY S T S

7.1 introduction

Manganese-based oxygen evolving catalysts have attracted significantinterest motivated mainly by the fact that core of the water oxidationcomplex (CaMn4O5) in PSII consists of this element. In addition, man-ganese is abundant in the earth crust, environmental-friendly, andinexpensive, thus many publications in the past described both ho-mogeneous metal complexes and heterogeneous catalysts based onmanganese. [314,315]

Taking into account the amazing activity of water oxidation com-plex in PSII, it seems interesting to develop its synthetic analogues.This challenging task was undertaken in recent reports, demonstrat-ing successful synthesis of bio-mimetic calcium/manganese oxides:CaMnO3, Ca2Mn3O8

[167,316] and CaMn2O3· x H2O. [168] Latter one ex-hibits photocatalytic activity under illumination with visible lightin the presence of Ru(bpy)3

3+ photosensitizer and S2O82- sacrificial

electron acceptor. Detailed structural investigations by XAS revealedclose structural similarity to natural complex, manifested in charac-teristic edge-sharing cuban units with Ca ion in the corner. [317] Theoxidation state of Mn was found to be close to + 4 (+ 3.8), indicatingthat some of Mn atoms are at + III oxidation state.

The systematic study by Dismukes et al. [318] compared several ox-ide structures containing MnIII and MnIV in order to asses their abil-ity to oxidize water under visible light irradiation, in neutral pHconditions and in the presence of Ru(bpy)3

3+ as a photosensitizerand S2O8

2- as a sacrificial electron acceptor. It has been conclusivelyshown, that catalytic activity decreases in series Mn2O3 > Mn3O4

>> λ-MnO2. Amounts of O2 produced during irradiation of λ-MnO2

were very small, whereas other MnO2 polymorphs (R-, α-, β-, γ-) ex-hibit no light-driven O2 evolution at all. Although these results differfrom some published studies [319] where authors claimed water oxida-tion with α-MnO2 nanowires, they are consistent with those of Frei etal. [208] and Suib et al. [210] The first study (Frei et al. [208]), reports nanos-tructured manganese oxide clusters supported on mesoporous silicaprepared by impregnation with Mn(NO3)2 followed by calcination.It was found that the most active clusters are mostly composed ofMn2O3, while less active of MnO2 phase. In the second study (Suib etal. [210]), both electro- and photocatalytic water oxidation was demon-strated with mixed-valent (MnIV/MnIII) amorphous manganese ox-ide consisting of nanoparticles with high surface area obtained by

103

104 manganese-based co-catalysts

reduction of KMnO4 with oxalic acid. Both studies used Ru(bpy)33+

photosensitizer and S2O82- sacrificial electron acceptor under mild pH

(5–6) conditions.Apparently, the compounds most active towards water oxidation

possess Mn at + III oxidation state, whereas the inactive compoundspossess Mn at + IV oxidation state. It was hypothesized that weaken-ing of Mn-O bond in MnIII(d4) edge sharing MnO6 octahedra due toJahn-Teller effect (distortion) caused higher catalytically activity, be-cause of increased structural flexibility. [318] Indeed, Jahn-Teller effectwas reported to occur also in Ni-Bi catalyst; distorted structure wasalso suggested for Co-Pi (see Section 6.1 and Section 5.1)

Recent developments in the electrocatalysis with highly-active Co-Pi and Ni-Bi have motivated investigation of analogues manganesemetal oxo/hydroxo clusters, that can be electrodeposited on conduc-tive substrates. Zaharieva et al. [160] have screened several electrodepo-sition protocols trying to relate conditions (voltage, solution) to ac-tivity. It was somewhat surprising that the deposition at constant an-odic potential, similar to Co-OEC and Ni-OEC, yielded inactive man-ganese oxides. Accordingly to their study, only voltage-cycling pro-tocols facilitate formation of active catalysts. Potential was cycled inde-ionized water containing Mn2+ ions for 25 min between – 0.75 and+ 2.15 V (each cycle 1 min) giving amorphous Mn oxo-/hydroxo- clus-ters (MnCat) exhibiting moderate current of 0.1 mA cm-2 recorded inphosphate buffer pH = 7 at overpotential of 0.5 V. Subsequent heat-ing of electrodes at relatively low temperature (≤ 120 °C) can lowerTafel slopes due to conductivity improvement arising form partial de-hydration of catalytic film. [320] Superior activity of the electrocatalystdeposited using the cycling potential protocol over those depositedusing constant potential was assigned to hindering of formation ofMnIVO2 which is inactive in water oxidation, as mentioned before.Similarly to the above discussed cases, the active MnOx thin film,electrodeposited using cycling-voltage protocol, contains around 20 %of MnIII (mean oxidation state + 3.8). [160] Therefore this study con-tributed in two ways: a new catalyst was found which exhibited ac-tivity close to that needed for interfacing with photoanode material,and the necessity of usage of cycling protocol supported the idea thatpresence of Mn at oxidation state + III plays a crucial role in catalyticprocess.

In following Sections the photoelectrochemical properties of TiO2-PH hybrid photoanodes loaded with manganese-based co-catalystsare examined in details.

7.2 photodeposition of manganese oxide clusters 105

7.2 photodeposition of manganese oxide clusters

7.2.1 Experimental

Manganese-based co-catalyst was photodeposited on TiO2-PH pho-toanode (fabrication method described in Section 3.3; ITO glass wasused). Procedure was similar to manufacturing protocols reported inliterature for electrodeposition of MnOx on bare ITO, [160] but with si-multaneous visible light irradiation.

The precursor solution, which contained 0.5 mM Mn2+ was ob-tained by dissolving 0.173 g of Mn(CH3COO)2 · 4 H2O (Sigma-Aldrich,≥ 99.0 %) in 1 L of water or 0.1 M MgSO4. Time, potential, solution(deionized water or MgSO4) and wavelength were varied accordingto protocols listed below:

a. Photodeposition with monochromatic visible light λ = 450 nmat constant potential Eappl. = 0.5 V vs. Ag/AgCl for ca. 10 min indeionized water;

b. Photodeposition with monochromatic visible light λ = 450 nmat constant potential Eappl. = 0.5 V vs. Ag/AgCl for ca. 30 min indeionized water;

c. Photodeposition with monochromatic visible light λ = 420 nmat constant potential Eappl. = 0.5 V vs. Ag/AgCl for 20 min in0.1 M MgSO4;

d. Photodeposition with monochromatic visible light λ = 450 nmat potential changed stepwise between + 0.25 and + 0.70 V vs.Ag/AgCl (see Figure 7.1 D) for ca. 25 min in 0.1 M MgSO4.

Dissolved oxygen measurements were conducted according to Sec-tion 3.1.2.2 without stirring.


Manganese-based oxygen evolving co-catalysts were introduced toTiO2-PH by the photodeposition method using both the constant po-tential protocol (similar to those of IrOx, CoOx (Co-Pi) and NiOx (Ni-Bi)) and the alternate potential steps protocol. In Figure 7.1 photocur-rents during photodeposition are presented. Three different protocols,varying the wavelength and the photodeposition time, but at constantpotential (A, B, C), were used. Additional alternate potential protocol(D) was also utilized because it was reported in literature to yieldparticularly active electrocatalysts.

All four methods caused dark brown colorization of the photoan-ode due to the expected MnOx formation. The co-catalyst was me-chanically stable, did not come out of the photoelectrode, could not be


Figure 7.1: Photocurrents during photodeposition of MnOx from solutioncontaining 0.5 Mn2+at constant potential Eappl. = 0.5 V vs.Ag/AgCl (see Section 7.2.1 manufacturing protocols A, B, C) andat alternative potential steps (see Section 7.2.1 protocol D).

washed out, and photoanode retained its color even after prolongedirradiation. Therefore, it was quite surprising that dissolved oxygenmeasurements have shown no oxygen evolution, as presented in Fig-ure 7.2. As can be seen, for all four photodeposition protocols pho-tocurrents are below 80 µAcm-2 and amounts of detected dioxygenare below 3 µmol dm-3. Such small amounts O2 cannot be attributedto water oxidation on the photoanode, but only to small leak of O2

from outside or temperature effects influencing functioning of thefluorescent-based O2 sensor.

A possible reason of inert behavior of MnOx as a co-catalyst inphotoelectrochemical water oxidation is low activity of manganeseoxide/hydroxide due to high strength of the M-OHad bond. (see Sec-tion 2.3.3; the activity of MOx decreases in order: Ni > Co > Fe >Mn). This is in line with the comparative study of water oxidationon BiVO4 loaded with CoOx, MnOx, IrOx, and RuOx co-catalysts,which showed that MnOx has the lowest activity, actually lower, than


Figure 7.2: a) Photocurrent and b) dissolved oxygen concentration measuredin 0.1 M phosphate buffer pH = 7 during polychromatic visiblelight (λ > 420 nm) irradiation of TiO2-PH + MnOx photoanodesbiased at 0.5 V vs. Ag/AgCl (1.12 V vs. RHE). Manganese oxidecluster were photodeposited as described in Section 7.2.1.

pristine BiVO4. In other words, loading with MnOx co-catalyst wasdemonstrated to even decrease activity of original material. [282]

Due to the inactivity of TiO2-PH photoanodes with photodepositedMn-based metal oxide clusters in water photooxidation reaction, theywere not investigated in more detail.

7.3 mixtures of metal oxides

7.3.1 Experimental

Apart from the photodeposition method, photoelectrodes based onmixtures of TiO2 and manganese oxides were investigated. The pho-toanodes containing manganese-based oxides were manufactured byseven different methods:


a. The procedure was analogous to one utilized for cobalt oxidesmixed with TiO2, presented in Section 5.3. In short, mixture of1 wt. % of commercially available Mn2O3 (Sigma-Aldrich) withTiO2 was ground in a mortar. After that, 200 mg of mixturewas suspended in 1 mL of ethanol. Prior to smearing on ITOglass and drying at 100 ºC for 15 min, suspension was sonicatedfor 15 min. In the next step, the photoelectrodes were covered(the matt side to the catalyst surface) with a piece of aluminumfoil and a normal glass, and then pressed for three minutes ata pressure of 150 kg cm-2. Lastly, electrodes were calcined at450 ºC for 30 min and modified by placing into a Schlenk tubeconnected via an adapter with a round bottom flask containing1 g of urea and heating in a muffle oven for 30 min at 425 °C.

b. Analogously to point A, but 2.5 wt. % of commercial Mn2O3

was present in the mixture.

c. In the first stage, Mn2O3 was synthesized using hydrothermalmethod reported in literature. [318] Briefly, 500 mg of MnCO3

powder was dissolved in a 0.5 M HCl and heated in an autoclavewith a Teflon liner at 150 ºC for 10 h. Then, the product waswashed with water and dried at 100 °C overnight. The result-ing powder was calcined at 550 °C for 4 h. In the second stage,photoanodes were prepared on ITO glass using procedure anal-ogous to point A with 10 wt. % of as synthesized Mn2O3.

d. Introduction of MnOx co-catalyst to TiO2-PH using impregna-tion method commenced with dipping the photoanode in the0.187 mg mL-1 of Mn(NO3)2 precursor solution. Then, the pho-toanode was quickly taken out and dried at 80 ºC for 10 min.This process was repeated 10 times. Finally, the photoelectrodewas calcined at 250 °C for 1 h. [282]

e. In the first stage CaMn2O4 ·H2O was synthesized using methodreported in the literature. [168] Shortly, 825 mg (3.5 mmol) ofCa(NO3)2 · 4 H2O and 550 mg (5.6 mmol) of MnCl2· 4 H2O wasdissolved in 10 mL of water and stirred for 15 min. Aqueoussolution of KMnO4 (2.4 mmol, 190 mg) and KOH (300 mmol,16.8 g) in 10 mL of water was added dropwise, resulting inthe dark precipitate. The precipitate was collected by centrifu-gation, resuspended in 10 mL of water, transferred to a Teflonlined stainless steel reactor and placed in the oven at 210 °C for72 h. The obtained suspension was filtered and washed withwater before being allowed to dry for one day at 60 °C in air.The dried powder was heated to 400 °C for 10 h in air to obtainbrown powder. In the second stage, photoanodes were preparedon ITO glass using procedure analogous to point A, but with10 wt. % of as synthesized CaMn2O4 · H2O.

7.4 conclusion 109

f. In accordance with point E, but 20 wt. % of CaMn2O4 · H2O wasused in the mixture.

g. In accordance with point E, but 40 wt. % of CaMn2O4 · H2O wasused in the mixture.

Dissolved oxygen measurements were conducted according to Sec-tion 3.1.2.2 without stirring.


In Figure 7.3 results of dioxygen measurements for mixtures of TiO2

and manganese oxides modified with polyheptazine are presented.Mn-based additives used for photoanodes preparation were: A, B –

commercial Mn2O3 nanopowder in different concentrations, C – syn-thetic Mn2O3, reported to have very high activity [318], D – MnOx ob-tained by heat treatment after impregnation in Mn(NO3)2, and E, F, G– CaMn2O4 · H2O bio-mimetic oxides in different concentrations. De-spite relatively high photocurrent for some samples (A, B), none ofthe oxides mixtures generated amounts of dioxygen which may be anevidence of complete water photooxidation to dioxygen. In all casesamounts of detected dissolved O2 were below 5 µmol dm-3. Smallrise of dioxygen concentration before switching on the light (45 min;measurement C and D) can be ascribed to slight leakage of O2 fromoutside.

Accepting the above-described results, Mn-based oxide species in-troduced to TiO2-PH are inactive in water photooxidation under vis-ible light. Actually, it is not easy to find examples of successful cou-pling of Mn-based co-catalyst to photoactive material, in spite of largevolume of published studies on water photooxidation. In reviewingthe literature, only two studies were found. First one demonstratesphotocatalytic (powder photocatalysis) overall water splitting withGaN:ZnO loaded with Rh/Cr2O3 (core/shell) and Mn3O4 nanoparti-cles acting as H2 and O2 evolving co-catalysts, respectively. [321] Sec-ond one investigates WO3 photoanode with improved activity due topresence of Mn-based catalyst whose structure is not exactly known. [322]

However, it is not clear, if the co-catalyst in this study really improvedwater oxidation rate or only acted as protective layer against WO3 dis-solution occurring normally at pH above 4. [322]

7.4 conclusion

In conclusion, TiO2-PH photoanodes loaded with Mn-based oxidesdo not oxidize water under visible light irradiation. Mn-based co-catalysts were implemented to TiO2-PH hybrid material using twodifferent methods: the photodeposition form solutions containing Mn2+

ions or by mixing different MnxOy oxides with TiO2 followed by mod-


Figure 7.3: a) Photocurrents and b) dissolved oxygen concentration recordedin phosphate buffer pH = 7 during irradiation with visible light(λ > 420 nm) of photoanodes consist of manganese oxides mixedwith TiO2 and modified with PH. Samples were prepared accord-ing to protocols described above. Photoelectrodes were biased at0.5 V vs. Ag/AgCl (1.12 V vs RHE).

ification with polyheptazine. Eleven samples, prepared by both meth-ods have been measured. Neither of them yielded dioxygen as a prod-uct of photooxidation. A possible explanation for this might be intrin-sically low electrocatalytic activity of the prepared manganese oxidesin water oxidation. Therefore, after implementation into the photoan-ode material, they act rather as light blocking layer or recombinationcenters.

8E F F E C T O F E X T E R N A L B I A S

8.1 introduction

In previous Chapters, we have observed that complete photooxida-tion of water to dioxygen was feasible only after the TiO2-PH hybridphotoanodes were loaded with an oxygen-evolving co-catalyst likeIrOx nanoparticles or amorphous CoOx clusters (“Co-Pi”). The cou-pling of the holes photogenerated in the polyheptazine layer to anefficient co-catalyst is therefore very important. In this Chapter, theTiO2-PH photoanodes with two different co-catalysts: IrOx nanoparti-cles deposited by the colloidal deposition and Co-Pi deposited by thephotodeposition method, will be compared in terms of their photo-conversion efficiency and stability. In particular the effect of appliedpotential on the efficiency of photocurrent and oxygen generationinvestigate will be investigated in details. Additionally, the behav-ior of TiO2-PH photoanodes under visible light irradiation with thatof pristine TiO2 photoelectrodes under UV light irradiation will becompared. Based on this analysis some key factors limiting the per-formance of TiO2-PH hybrid photoanodes in water splitting will bediscussed.

8.2 experimental

Iridium oxide nanoparticles were synthesized using the method de-scribed by Maeda et al. [230] The colloidal deposition method was de-scribed in Section 4.2.1.

CoOx oxygen-evolving co-catalyst (Co-Pi) was deposited on TiO2-PH photoelectrodes by method described in Section 5.2.1, but fromsolution containing 0.5 mM instead of 0.3 mM Co2+. The control sam-ples using pristine TiO2 electrode (TiO2 + Co-Pi) and photodepositionwavelength of 350 nm were prepared in the same way.

Dissolved oxygen measurements were performed in one compart-ment cell with two-electrode setup, in accordance to procedure de-scribed in Section 3.1.2.4 without stirring.

8.3 results and discussion

Figure 8.1 shows photoaction spectra (IPCE versus λ plots) recordedat TiO2 and TiO2-PH photoelectrodes loaded with either IrOx/CD orCo-Pi. As expected, in contrast to pristine TiO2 whose photoresponseis confined to the UV range (< 400 nm), TiO2-PH photoelectrodes

111

112 effect of external bias

Figure 8.1: Photoaction spectra (IPCE vs. λ plots) measured in a three-electrode setup at different potentials in deaerated phosphatebuffer (pH 7) at pristine (a) TiO2, and (b) TiO2-PH loaded withIrOx/CD, and at (c) TiO2 and (d) TiO2-PH loaded with Co-Pi.

exhibit significant IPCE values down to ∼ 540 nm, corresponding tothe optical absorption edge of TiO2-PH (Figure 1.3). [44] Similar IPCEvalues are observed on photoelectrodes modified with IrOx/CD andCo-Pi. Most interestingly, while at pristine TiO2 photoelectrodes theIPCE values are relatively insensitive to applied potential, at TiO2-PH photoanodes the IPCE values increase continuously with morepositive bias potential. (Figure 8.1).

This is further illustrated in Figure 8.2 showing photocurrent tran-sients recorded during a slow cathodic potential sweep. Here pristineTiO2 photoelectrodes exhibit a typical profile of nanocrystalline pho-toelectrodes in which the photocurrent generation is governed by thekinetics of holes at the TiO2-electrolyte interface and by diffusion ofelectrons through the porous nanocrystalline layer, rather than by thepotential drop within the space-charge layer, as in case of compactsingle- or polycrystalline semiconductors. [195,323]

8.3 results and discussion 113

Figure 8.2: Photocurrent measured at (a) pristine TiO2 and at (b) TiO2-PHphotoelectrodes without and with different co-catalysts recordedunder intermittent monochromatic irradiation (5 s light, 5 s dark;λ = 350 nm for TiO2, λ = 420 nm for TiO2-PH) during a cathodicpotential sweep (5 mV s-1) in deaerated phosphate buffer (pH 7).

Two further points are noteworthy. First, the deposition of Co-Pileads to a slight cathodic shift of the photocurrent onset potentialfor both TiO2 and TiO2-PH, while the colloidal deposition of IrOx

nanoparticles leaves the photocurrent onset practically unchanged.Assuming that at potentials near the conduction band edge of TiO2

(– 0.2 V versus RHE) the photocurrent generation is mainly deter-mined by kinetic competition between the transfer of holes to watermolecules and the electron-hole recombination, we conclude that Co-Pi performs better as a co-catalyst in our electrodes. This is not sur-prising as IrOx nanoparticles are deposited simply by the colloidaldeposition, Co-Pi is formed photoelectrochemically exactly at surface


sites where holes are predominantly generated, which results in bet-ter coupling between holes and the co-catalyst. Second, as comparedwith the UV (λ = 350 nm) light induced photocurrent response atpristine TiO2, the visible (λ = 420 nm) light photocurrents at hybridTiO2-PH electrodes exhibit pronounced sensitivity to applied poten-tial and the photocurrent onset is shifted anodically by ∼ 0.4 V. Thisstrong dependence of photoconversion efficiency on applied externalbias even at potentials relatively far from the potential of the con-duction band edge of TiO2 suggests that in TiO2-PH hybrids severeprimary recombination processes [direct back charge transfer follow-ing the optical charge-transfer absorption (Figure 1.2 c)] occur unlessexternal bias is applied to drive electrons away. Clearly, these recom-bination processes are here much faster than the kinetics of wateroxidation by holes, even in the presence of co-catalysts. Furthermore,it suggests that the coupling between the photogenerated holes andthe Co-Pi co-catalyst is still far from perfect, which leads to accumula-tion of positive charges in the surface environment of TiO2 and mightresult in an anodic shift of its conduction band edge.

Photocurrent and oxygen evolution measurements under polychro-matic visible light irradiation (λ > 420 nm) for TiO2-PH without andwith IrOx/CD and Co-Pi co-catalysts are shown in Figure 8.3. Thesemeasurements are performed in a two-electrode set-up for effectingthe most realistic estimation of the external electric bias needed todrive the water-splitting reaction. At both bias voltages (0.7 V and1.0 V), the results clearly demonstrate the superior behavior of Co-Pi-loaded photoanodes. Both photocurrent and the amount of evolvedoxygen are about doubled as compared with photoanodes loadedwith IrOx/CD. This might seem rather surprising taking into accountthe very similar IPCE values shown in Figure 8.1. This seeming dis-crepancy is explained by the very different time-scale of the corre-sponding measurements, and by the fact that the photocurrent tran-sients recorded at TiO2-PH photoanodes exhibit a “spike-like” behav-ior (a very fast rise of photocurrent followed by decrease), which isa clear fingerprint of surface recombination processes that are go-ing on. [238] In other words, after switching on the light, the holesvery quickly accumulate in the surface polyheptazine layer, whichincreases probability of their recombination with electrons photogen-erated in the conduction band of TiO2. However, when sufficient ex-ternal electric bias is applied to drive electrons away and a suitableco-catalyst is present to catalyze transfer of holes to water molecules,a significant portion of photogenerated holes is able to escape recom-bination and induce water oxidation to dioxygen. The very small dif-ference in IPCE values between photoelectrodes loaded with IrOx/CDand Co-Pi arises from the fact that the IPCE measurements are per-formed under relatively short-pulsed monochromatic irradiation (5 slight, 10 s dark) so that photocurrents are read-off very quickly after


Figure 8.3: a) Photocurrent and b) oxygen evolution measured in a two-electrode setup in phosphate buffer (pH 7) during prolongedirradiation by polychromatic visible light (λ > 420 nm). Work-ing electrode was biased at 0.7 V versus Pt in the first part (till125 min) and at 1.0 versus Pt in the second part of experiment(after 125 min).

the initial rise. Obviously, experiments using prolonged (time-scale∼ hours) irradiation shown in Figure 8.3 give a much more real-istic estimation of photoanode performance since the system is al-lowed to reach steady-state conditions. It should be noted that thebetter performance of Co-Pi-loaded photoanodes as compared withIrOx-loaded ones is apparently not only due to better coupling ofphotogenerated holes with the in situ grown Co-Pi but also due to


bad adhesion of IrOx/CD to the surface of TiO2-PH photoanodes, ashas been reported for other photoactive materials (like e.g. Fe2O3)loaded with IrOx. [231] Interestingly, the TiO2-PH photoelectrode with-

Figure 8.4: a) Photocurrent and b) oxygen evolution measured in a phos-phate buffer (pH 7) during prolonged irradiation by polychro-matic visible light (λ > 420 nm). TiO2-PH electrodes with pho-todeposited Co-Pi were biased at 0.8, 0.9, and 1.0 V versus Pt.The control samples, i.e. pristine TiO2 with photodeposited Co-Pi and TiO2-PH without Co-Pi, were biased at 0.9 V.

out any co-catalyst shows quite significant photocurrents at higherbias (1 V versus Pt), while no oxygen production is observed. Whensearching for possible intermediates of the water oxidation reactionwe failed to detect any hydrogen peroxide (H2O2) using a standard


horseradish peroxidase-based photometric test. It cannot be ruled outthat in the absence of the co-catalyst the anodic photocurrents leadalso to partial oxidation of polyheptazine or some side-products fromthe polyheptazine-modification procedure. The behavior of TiO2-PHphotoelectrodes modified with Co-Pi was investigated in more detailat different bias voltages (Figure 8.4).

At all voltages the photocurrent transients recorded at Co-Pi-modifiedphotoanodes still exhibit a “spike-like” behavior. This means thatthese recombination processes are proceeding faster than the Co-Pi-catalyzed transfer of holes to water molecules. It is noteworthy that,as expected, neither TiO2-PH without a co-catalyst nor pristine TiO2

loaded with Co-Pi show oxygen evolution under visible light irra-diation (λ > 420 nm). Furthermore, both the photocurrent and the

Figure 8.5: (•) An amount of oxygen evolved and (a) electric charge passedat a TiO2-PH + Co-Pi photoelectrode biased at different externalpotentials during 1h irradiation with polychromatic visible light(λ > 420 nm).

amount of evolved oxygen increase linearly with increase of the ap-plied external bias. This can be further illustrated in Figure 8.5 whichshows the linear dependence of oxygen concentration and the over-all passed electric charge (photocurrent integrated over time) on theapplied bias. The minimal external bias needed can be estimated tobe approximately 0.6 V versus Pt. This value coincides well with thephotocurrent onset measured in a three-electrode setup (Figure 8.2),which suggests that the proton reduction reaction proceeds practi-cally without significant overpotential at the platinum electrode. Itshould be also noted that reasonable photocurrents and oxygen evo-lution could be still observed in the bias range (∼ 0.7 V versus Pt)in which, for example, the mostly studied α-Fe2O3 photoanodes typ-ically show no photocurrent response. [96] However, as already notedabove, this bias voltage is significantly higher than that one would


expect taking into account the fact that the potential of the conduc-tion band edge of TiO2 slightly straddles the reduction potential ofprotons (Figure 1.2).

It is important to emphasize that the better performance of TiO2-PH+ Co-Pi photoelectrodes in comparison to TiO2-PH + IrOx/CD photo-electrodes does not mean that cobalt oxide-based oxygen evolving co-catalyst is better than iridium oxide-based one. The superior behav-ior of Co-Pi co-catalyst presented in this Chapter originates mainlyfrom more efficient deposition method (photodeposition). Neverthe-less, TiO2-PH photoelectrodes loaded with iridium oxide-based co-catalyst by the photodeposition method (IrOx/PD; Section 4.3) exhibitcomparable photoelectrochemical properties than (also photodeposited)cobalt oxide-based co-catalyst (Co-Pi; Section 5.2). It confirms onceagain great importance of the method used to interface a co-catalystwith an absorber. Additionally, it should be noted that direct compar-isons between photoelectrodes loaded with various co-catalysts arerather difficult due to the different stability and activity windows(e.g. in function of pH) of different co-catalysts.

8.4 conclusion

In summary, the photoelectrochemical properties of TiO2-PH hybridphotoanodes loaded with IrOx/CD and Co-Pi as oxygen-evolving co-catalysts were investigated, particularly with the focus on the effectof external electric bias on the efficiency of visible light-driven wa-ter photooxidation. As compared with IrOx-loaded photoanodes, theTiO2-PH photoelectrodes modified with Co-Pi co-catalyst exhibitedbetter stability and performance under prolonged irradiation, pre-sumably because the photoelectrochemical deposition precipitationof Co-Pi helped to establish a better contact between the photogen-erated holes and the co-catalyst. Under visible light irradiation, theminimum external electric bias needed to observe complete photooxi-dation of water to dioxygen at Co-Pi modified TiO2-PH photoanodeswas estimated to be ∼ 0.6 V at pH 7. The key factor limiting the photo-conversion efficiency at low bias potentials is the fast primary recom-bination of photogenerated charges. Further improvements can bepossibly achieved if fast back-charge transfer processes are avoidedby engineering the interface between polyheptazine and TiO2, anda more effective coupling of photogenerated holes to a co-catalyst isestablished for avoiding accumulation of positive charges in the sur-face polyheptazine layer.

Part IV

S U M M A RY

9S U M M A RY

The past decades have seen a rapid development of systems capableto transform solar energy to heat or electricity, which was triggered bya growing need for covering the increasing energy demand in cheapand completely sustainable way. However, the development of solar-to-chemical energy conversion systems seems to be more attractive ap-proach because energy can be directly stored in chemical bonds, e.g.of dihydrogen molecules.

One of the most appealing methods for direct conversion of sun-light energy into chemical energy is the use of photoelectrochem-ical cells in which simultaneous oxidation and reduction of waterto hydrogen and oxygen takes place. Although impressive solar-to-hydrogen efficiencies reaching far beyond the 10 % threshold requiredfor commercial implementation have been demonstrated repeatedly,all these high-efficiency cells have been based on very expensive ma-terials and cannot be easily scaled-up. In this context it is important torealize that the major challenge in photoelectrochemical water split-ting devices is the development of low-cost, efficient and stable pho-toanodes. This is particularly because of the complex chemistry in-volved in four-electron oxidation of water to dioxygen, which typi-cally translates into high overvoltages needed to drive the reaction,which, in turn, enhances the risk of oxidative photocorrosion.

Traditionally, the attention in scope of water photooxidation is drawnmostly to pristine or doped metal oxides like TiO2, WO3, and Fe2O3.Whereas these materials have relatively high inherent stability, themajor problems are: too positive position of CB or too large bandgap,and limited flexibility of their structure, which brings about increasedoverpotentials for oxygen evolution. Consequently, significant exter-nal bias application is a prerequisite to drive water photooxidation atpristine metal oxides.

A different approach is to use dye-sensitized photoanodes in whichnanocrystalline layer of a stable wide-bandgap oxide sensitized bya visible light absorbing dye is coupled to the oxygen-evolving cata-lyst. Despite many advantages, like visible light absorption, flexibility,and kinetic charge separation, such materials often suffer because ofinstability of the dye in harsh water oxidation conditions.

An alternative way, which might combine advantages of pristinemetal oxides (stability) and dye-sensitized materials (flexibility, ki-netic charge separation) is represented by photoanodes based on a novelclass of visible-light photoactive hybrid materials: nanocrystalline TiO2

modified at the surface by a thin (< 1–3 nm) layer of polyheptazine

121

122 summary

(TiO2-PH). One of the most attractive features of these inorganic/or-ganic hybrids is the high thermal (up to 550 °C in air) and chemi-cal stability of polyheptazine-type compounds as compared to con-ventional organic dyes. It was shown that polyheptazine and TiO2

form an interfacial charge-transfer complex, which effectively shiftsthe optical absorption edge of the TiO2-polyheptazine hybrids intothe visible (2.3 eV; ∼ 540 nm) as compared to the bandgaps of both ofthe single components – TiO2 (3.2 eV; ∼ 390 nm) and polyheptazine(bandgap of 2.9 eV; ∼ 428 nm). [44] In other words, the direct opticalcharge transfer is expected to produce electrons with a relatively neg-ative potential in the conduction band of TiO2, while the photoholesare located in the surface polyheptazine layer. Those properties makesTiO2-PH an excellent platform for development of photoanodes foruse in photoelectrochemical applications.

Difficulties arise, however, when an attempt is made to use simplyTiO2-PH photoanode to oxidize water. TiO2-PH is a very good visiblelight absorber and charge separating material, but complete photoox-idation of water to dioxygen it feasible only after coupling with addi-tional metal oxide sites (co-catalyst) able to catalyze complicated holetransfer to water molecule.

The present study was designed to determine the effect of loadingof TiO2-PH with various co-catalysts by careful evaluation of evolveddioxygen concentration measurements results obtained by means ofthe home-build photoelectrochemical experimental setup. In particu-lar, this dissertation sought to address the following questions:

• Which type of co-catalysts can efficiently mediate hole transferfrom TiO2-PH hybrid to water molecules?

• How the photooxidation of water depends on the method, bywhich the co-catalyst was introduced to the light-absorbing pho-toanode material?

• What is the influence of conditions, e.g. external bias, pH, typeof an electrolyte at the activity of TiO2-PH loaded with a co-catalyst in water oxidation reaction?

The content of the present thesis can be summarized as follows:

chapter 2 has introductory character and brings basic informationsabout photoelectrochemical cells, design of photoanodes, prin-ciples of water oxidation reaction, and strategies which can beemployed to improve system efficiency.

chapter 3 compiles experimental details about photoelectrochem-ical measurements, design of experimental setup for dissolvedoxygen measurements, instrumental characterization techniques,and synthesis of TiO2-PH photoanodes.

summary 123

chapter 4 begins with concise literature review introducing recentdevelopments regarding iridium-based catalysts in context of wa-ter oxidation, followed by presentation of the author’s investi-gation results. In the first part, nanocrystalline TiO2-polyhept-azine (TiO2-PH) hybrid photoelectrodes were loaded with irid-ium oxide nanoparticles, acting as oxygen evolving co-catalyst,by the colloidal deposition method (TiO2-PH/CD). Significantly,visible-light irradiation (λ = 420 nm) led to dioxygen evolution.In the second part of Chapter 4, iridium oxide co-catalyst wasphotodeposited onto TiO2-PH photoelectrodes from solution con-taining [Ir(OH)6]2- anions. Iridium oxide clusters were formedin-situ under the irradiation of monochromatic visible light. Itwas found that TiO2-PH photoanodes with photodeposited IrOx

yield 2.5 times higher amount of evolved dioxygen than IrOx

loaded using colloidal deposition method mainly due to thebetter coupling to holes photogenerated in TiO2-PH absorber.It was also shown that presence of photodeposited IrOx co-catalyst reduces the surface recombination rate. In addition, thephotodeposition method is very reproducible, highly control-lable, and yields stable co-catalyst layers which allowed to inves-tigate properties of TiO2-PH photoelectrode loaded with IrOx

co-catalyst in the broad range of pH. A strong dependence ofactivity on electrolyte properties (anion, pH) was found. TiO2-PH + IrOx/PD photoanode exhibit highest activity and stabilityin Na2SO4 at pH = 6. These results have shown that measure-ment conditions must be cautiously selected in order to assurecompatibility with all essential components of photoanode.

chapter 5 commences with brief literature review of cobalt-basedmaterials investigated in scope of water photooxidation, fol-lowed by presentation of investigation results. In the first part,a cobalt oxide-based oxygen-evolving co-catalyst (Co-Pi) wasphotodeposited by visible-light irradiation onto TiO2-PH pho-toelectrodes from phosphate-buffered solution containing Co2+.In the second part, cobalt-based co-catalysts were introducedby mixing three different oxides (Co3O4, NiCo2O4, and CoTiO3)with pristine TiO2 and subsequent modification with polyhep-tazine. Both mixtures of oxides and Co-Pi co-catalyst coupleeffectively to photoholes generated in polyheptazine layer ofTiO2-PH photoanode, as evidenced by complete photooxidationof water to oxygen under visible-light (λ > 420 nm) irradiationat moderate bias potentials. In addition, the presence of theco-catalyst again reduces the recombination of photogeneratedcharges, particularly at low bias potentials, which was ascribedto better photooxidation kinetics. This suggests that further im-provements of photoconversion efficiency can be achieved if

124 summary

more effective catalytic sites for water oxidation are introducedto the surface structure of the hybrid photoanodes.

chapter 6 begins with short literature review of nickel-based materi-als investigated in context of water oxidation. The next Sectionpresents investigation results of TiO2-PH loaded with a nickel-based oxygen evolving catalyst (NiOx) photodeposited underthe visible light irradiation from borate-buffered solution con-taining Ni2+ cations. It was found, that the amount of dioxy-gen generated by TiO2-PH + Ni-Bi and the IPCE values arerather moderate, especially if compared with those of cobalt-or iridium-based oxygen evolving co-catalysts created by thephotodeposition method.

chapter 7 starts with concise presentation of recent developmentsregarding managanese-based materials in scope of water oxida-tion. In the following Sections attempts to interfaces Mn-basedoxides with TiO2-PH by two methods: in-situ photodepositionand mixing of metal oxides are presented. None of the sam-ples exhibit capability to photoxodize water, as evidenced byunchanged dissolved dioxygen concentration during irradiationwith visible light. It is probably due to the low, in comparison toiridium, cobalt or nickel oxides, electrocatalytic activity of hereprepared manganese oxides acting in this case rather as a lightblocking layer.

chapter 8 describes investigation of visible (λ > 420 nm) light -drivenphotooxidation of water at TiO2-PH electrodes loaded with twodifferent metal oxide co-catalysts: Co-Pi and IrOx/CD. As com-pared with TiO2-PH photoanodes loaded with iridium oxide bycolloidal deposition (IrOx/CD), photoelectrodes modified withCoOx oxygen-evolving co-catalyst (Co-Pi) deposited by photoas-sisted deposition precipitation method showed both higher pho-tocurrents and more efficient oxygen evolution under prolongedirradiation. The minimum external electric bias needed to ob-serve complete photooxidation of water to dioxygen at TiO2-PHphotoanodes modified with Co-Pi was estimated to be ∼ 0.6 Vat pH 7. The key factor limiting the photoconversion efficiencyat low bias potentials is the fast primary recombination of pho-togenerated charges.

In conclusion, this thesis presents results of investigation of visiblelight active TiO2-PH photoanodes loaded with additional metal ox-ides acting as oxygen evolving centers. It is focused on evolved dioxy-gen measurements and photoelectrochemical characterization.

One of the more significant findings to emerge from this study isthat after loading with iridium, cobalt or nickel oxide species, TiO2-PH photoanode oxidizes water under irradiation with visible light

summary 125

(λ > 420 nm), as evidenced by dissolved dioxygen measurements. Onthe contrary, it was also found, that manganese oxide species arecompletely inactive as an oxygen evolving co-catalyst. The secondmajor finding is that photooxidation efficiency strongly depends ona co-catalyst loading method, as illustrated by higher activity of IrOx

introduced by the photodeposition method than the colloidal deposi-tion method, mostly due to the better coupling with absorber (TiO2-PH). It is important to emphasize that the photodeposition method(particularly investigated for IrOx co-catalyst implementation) is verypromising for the further research. For example, it would be inter-esting to assess the possibility of simultaneous deposition of two ormore co-catalysts on the photoanode surface. The next conclusionwhich can be drawn from the present study is that the presence ofthe co-catalyst reduces accumulation of holes, thereby retarding re-combination of photogenerated charges, particularly at low bias po-tentials. Finally, results of this thesis have also indicated importancethe electrolyte type and its pH for efficiency and stability of hybridphotoanodes.

Taken together, the findings of this thesis add to our understand-ing about methods which can be used to implement catalytic sitesinto semiconducting and hybrid (e.g. dye-sensitized) light absorbingmaterials in order to use photogenerated holes to drive water oxi-dation. Further research aiming to improve electrocatalytic activityof the co-catalyst, the visible-light photoactivity of absorber and cou-pling between those two, would be of great help in development ofstable and efficient photoanodes which are in the center of our effortsdevoted to create new solar-to-chemical energy conversion systems.

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C O N F E R E N C E C O N T R I B U T I O N S

TA L K S :

1. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek "HybridPhotoanodes for Visible Light-driven Water Splitting", 2013 MRSSpring Meeting, San Francisco (USA), April 01-05, 2013.

2. M. Bledowski, L. Wang, A. Ramakrishnan, O. V. Khavryuchenko,R. Beranek "Photoelectrochemical water splitting at hybrid pho-toanodes" ISE-2012, 63rd Annual Meeting of the InternationalSociety of Electrochemistry, Prague (Czech Republic), August19-24, 2012.

P O S T E R S :

1. M. Bledowski, L. Wang, R. Beranek, Visible light-driven watersplitting at hybrid photoanodes, International Conference on"Solar Energy for World Peace", Istanbul (Turkey), August 17-19, 2013.

2. M. Bledowski, L. Wang, R. Beranek, Visible light-driven wa-ter splitting at TiO2-polyheptazine hybrid photoanodes, Inter-national Conference on New Advances in Materials Researchfor Solar Fuels Production (SolarFuel13), Granada (Spain), June12-14 June 2013

3. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek "Visiblelight-driven photooxidation of water at TiO2-polyheptazine hy-brid photoanodes" 2013 MRS Spring Meeting, San Francisco(USA), April 01-05, 2013.

4. L. Wang, M. Bledowski, A. Ramakrishnan, R. Beranek, "Visiblelight-driven water splitting at hybrid photoanodes”, 10th Mate-rials Day, Bochum (Germany), November 08-09, 2012.

5. L. Wang, M. Bledowski, A. Ramakrishnan, R. Beranek "Visiblelight-driven photooxidation of water at hybrid photoanodes"GDCh Photochemie 2012, Potsdam (Germany), October 08-10,2012.

6. L. Wang, M. Bledowski, A. Ramakrishnan, R. Beránek "Photo-electrochemical Water Splitting at Hybrid Photoanodes" GDChElectrochemistry 2012, Munich (Germany), September 17-19, 2012.

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7. L. Wang, M. Bledowski, A. Ramakrishnan, R. Beranek "Hybridmaterials for photoelectrocatalytic solar energy conversion" 15thInternational Congress on Catalysis 2012, Munich (Germany),July 01-06, 2012.

8. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek "HybridMaterials for Photoelectrochemical Water Splitting" The interna-tional CECAM-Workshop "Photo- meets Electrocatalysis: UnitedWe Split (...Water)", Delmenhorst (Germany), October 04-07, 2011

9. M. Bledowski, L. Wang, A. Ramakrishnan, R. Beranek "NovelHybrid Materials for Photoelectrochemical Water Splitting" 5thGerischer Symposium „Photoelectrochemistry: From Fundamen-tals to Solar Applications“, Berlin (Germany), June 22-24, 2011

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Visible light-driven photooxidation of water at hybrid photoanodes

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