Abundant Metal Oxide Nanomaterials for Solar Fuel Generation

Investigations of Earth- Abundant Metal Oxide

Nanomaterials for Solar Fuel Generation

by

Peter Mirtchev

A thesis submitted in conformity with the requirements

for the degree of Doctor of Philosophy

Department of Chemistry

University of Toronto

© Copyright by Peter Mirtchev 2015

ii

Investigations of Earth- Abundant Metal Oxide Nanomaterials

for Solar Fuel Generation

Peter Mirtchev

Doctor of Philosophy

Department of Chemistry

University of Toronto

2015

Abstract

Developing renewable energy technologies to mitigate anthropogenic climate

change is one of the biggest challenges facing humanity in the 21st century. Making

use of the Sun’s bountiful energy is society’s best hope for achieving cheap, efficient

renewable power on a global scale. Solar-assisted conversion of abundant

resources such as CO2 and H2O into valuable hydrocarbons is an attractive

proposition in this respect. This work investigates the synthesis and characterization

of nanoparticulate metal oxides based mainly on abundant iron and copper, and

their application as photocatalysts for CO2 reduction and H2 evolution from water.

We begin by presenting a general introduction to artificial photosynthesis and a brief

literature review of progress in the field. The synthesis, characterization, and hybrid

properties of novel Fe2O3/Cu2O hetero-nanocrystals are then described. These

nanoparticles represent one of the few examples of colloidal oxide-oxide hetero-

structured nanocrystals in the literature. In subsequent work, we explore Cu2O

nanocubes as a semiconducting scaffold for the synthesis of multi-component

photocatalytic architectures. This work then led us to study the activity of metallic Cu

on TiO2 as a model H2 evolution system and examine the effects of alcoholic

scavengers on product distribution in water splitting experiments. Finally, we discuss

iii

iron-copper delafossite CuFeO2 which was found to be an active catalyst for the

light-assisted hydrogenation of CO2 to CO. We conclude by summarizing some of

the lessons learned over the course of this work in trying to develop cheap, efficient

artificial photosynthesis catalysts, and attempt to provide useful guidelines that may

aid future researchers in this pursuit.

iv

Acknowledgements

Firstly I would like to thank my supervisor Geoff Ozin for giving me the

opportunity to work on a project in renewable energy, an area that I’ve always been

passionate about. Geoff’s efforts to secure funding ensuring the continued success

of this project were incredible and we are all indebted to his dedication. I would like

to thank the other members of my committee, past and present, Bob Morris, Greg

Scholes, Gilbert Walker and particularly Doug Perovic for opening the door to great

collaborations that resulted in helpful data. All of our academic collaborators

especially Dr. Frank Osterloh, Dr. Stephen Pennycook, Dr. Chuck Mims, Dr.

Zhenghong Lu and the rest of the solar fuels team deserve a mention for making this

research possible. Enormous thanks to our lab manager Sue Mamiche Afara, who

keeps our lab running smoothly and Dr. Navid Soheilnia who is the most

knowledgeable materials chemist that I have ever met. I’d like to acknowledge the

numerous people who have taught me all I know about the techniques that I’ve

learned to use over the course of the last 5 years: Kristine Liao for her XPS/UPS

knowledge, Veronika Hoepfner for gas chromatography, Paul, O’Brien and Amit

Sandhel for their help in sample testing, Eric Henderson for letting me help out on

the CQD project, and Neil Coombs and Ilya Gourevich at the microscopy facility for

their help in getting nanoparticle images on multiple occasions. And thanks to Anna

Liza Villavelez and the other staff members of the Graduate Office who were always

extremely helpful.

To my friends who had the pleasure of getting to know me, both inside and

outside grad school, I know who you are without having to list names. I hope you

enjoyed our time together as much I did and I hope our roads do not separate too far

while we get to where we’re going.

To my family who I can always rely on, thank you for raising me the way you

did and giving me every opportunity to do what I wanted to do. I wouldn’t have had it

any other way.

v

Table of Contents

List of Tables…………………………………………………….....................................viii

List of Figures……………………………......................................................................ix

List of Abbreviations……………………………………................................................xv

Chapter 1- Introduction to Solar Fuels

1.1 Scientific Motivation…………………………………………………………………1

1.2 Basics of Artificial Photosynthesis ………….………………………….………....2

1.3 Literature Overview ………….……………………………………..……………..11

1.4 Focus of This Thesis………….…………………………………………………...22

1.5 References…………………………………………………………………………24

Chapter 2 – Synthesis of Fe2O3/Cu2O Hetero-Structured Nanocrystals

2.1 Abstract……………………………………………………………………………..34

2.2 Introduction to Hetero-structured Nanocrystals…………………............. ……35

2.3 Results and Discussion……………………………………………………………42

vi

2.4 Conclusions………………………………………………………………………...55

2.5 Experimental………………………………………………………………………..55

2.6 References……………………………………………………………..................58

Chapter 3 – Electronic Properties and Applications of Fe2O3/Cu2O HNCs

3.1 Abstract……………………………………………………………........................64

3.2 Photoelectron Spectroscopy……………………………………………………...65

3.3 Gas Phase CO2 Reduction…………………………………..............................71

3.4 Ligand Removal & Dye Degradation……………………..................................76

3.5 Conclusions………………………………………..............................................84

3.6 Experimental…………………………………….................................................85

3.7 References……………………………………………........................................88

Chapter 4 - Investigations of Cu2O Nanocubes as Semiconducting

Scaffolds for Photocatalytic H2 Evolution and CO2 Reduction

4.1 Abstract……………………………………………………………………………..90

4.2 Introduction………………………………………………………..…....................91

4.3 Results and Discussion ……………………………………...............................92

vii

4.4 Conclusions…………………………………………........................................111

4.5 Experimental……………………………………………………………..............112

4.6 References…………………………………………………………....................116

Chapter 5 – Light-Assisted Hydrogenation of CO2 to CO Using a Mixed Metal

Oxide Delafossite, CuFeO2

5.1 Abstract…………………………………………………………..…....................120

5.2 Introduction……………………………………………………………………......121

5.3 Results and Discussion…………………………………………………............122

5.4 Conclusions……………………………………………………………………….136

5.5 Experimental……………………………………………………………..............137

5.6 References……………………………………………………………………......141

Chapter 6 – Conclusions & Future Outlook

6.1 Concluding Remarks……………………………………....…………...............143

6.2 Future Outlook for Solar Fuels……………………….....................................147

6.3 References………………………………………………………………………..150

viii

List of Tables

Table 1.1 - Reactions of interest in artificial photosynthesis and their

thermodynamic potentials………………………………………………………………….5

Table 1.2 – Summary of notable semiconductor-based CO2 reduction

systems …………………………………………………………………………………….13

Table 3.1 - Binding energy (eV) of Fe 2p core-level lines in γ-Fe2O3

nanocrystals, γ-Fe2O3/Cu2O HNCs, and commercial iron oxide nanopowders…….65

Table 4.1 - Rates of H2 evolution from water with Cu2O nanocubes and

related materials as the photocatalyst…………………………………………………101

Table 4.2 Rates of H2 evolution from water using P25/Cu (10%) mixture and

its separate components………………………………………………………………...107

ix

List of Figures

Figure 1.1 – US energy consumption by type…………………………………...1

Figure 1.2 – The process of artificial photosynthesis…………………………...2

Figure 1.3 – Light absorption, charge migration, and surface reactions in a

heterogeneous photocatalyst………………………………………………………………3

Figure 1.4 – Formation of an electron-hole pair in a semiconductor upon

excitation with light……………………………………………………………………….....4

Figure 1.5 – Positions of VB/CB energies with respect to the redox potentials

of surface molecules………………………………………………………………………..5

Figure 1.6 – Positions of the VB and CB potentials of various semiconductors

at pH=1 relative to the redox potentials of CO2 reduction to different products……..6

Figure 1.7 – Formation of an electron-hole pair and various recombination

pathways inside a semiconductor ………………………………………………………...7

Figure 1.8 – A semiconductor-metal junction where the metal acts as an

electron sink and reducing site for adsorbed reactants…………………………………8

Figure 1.9 – The role of sacrificial reagents in scavenging the majority charge

carriers shown for the water splitting process……………………………………………9

Figure 2.1 – Various morphologies of reported colloidal HNCs including

core/shell, dimer, trimer, and oligomer architectures…………………………………..36

Figure 2.2 – Illustration of the FM, SK, and VW modes for growth of a

secondary material onto a seed nanocrystal …………………………………………..38

Figure 2.3 – Schematic representation of charge carrier confinement regimes

in semiconductor hetero-nanocrystals…………………………………………………..41

Figure 2.4 – Schematic Illustration of a Type II hetero-nanostructure with an

electron rich domain for CO2 reduction and hole rich domain for H2O

oxidation…………………………………………………………………………………....41

Figure 2.5 – Crystal structures and lattice constants of the three components

of the heterostructured nanocrystals…………………………………………………….42

x

Figure 2.6 – 1H NMR spectrum of Fe(oleate)3; inset – IR spectrum of

Fe(oleate)3 ..………………………………………………………………………………..43

Figure 2.7 – TGA scans of Fe(oleate)3 and Cu(I)acetate showing the initial

decomposition temperatures of the precursors..……………………………………….44

Figure 2.8 – a) Representative TEM image and particle size distribution of

isolated Cu2O and b) γ-Fe2O3 nanocrystals..…………………………………………..44

Figure 2.9 – a, b) Low-resolution TEM images of HNCs and physical mixture

of γ-Fe2O3 and Cu2O showing the absence of any ordering into hetero-architectures

c) Particle size distribution of the γ- Fe2O3, Cu, and Cu2O domains in as-synthesized

HNCs dimers and oligomers……………………………………………………………...46

Figure 2.10 – a, b) HRTEM images of as synthesized HNCs c) EDX line scan

across dimer particle showing the Fe-rich and Cu-rich domains (d) PXRD patterns of

Cu2O, γ-Fe2O3, and γ-Fe2O3/Cu2O HNCs as thin films e) Raman spectrum of as-

synthesized γ-Fe2O3 nanocrystals showing the prominent A1g phonon mode at 701

cm-1 indicative of γ-Fe2O3…………………………………………………………………47

Figure 2.11 – a, b) HRTEM images and STEM/EELS maps of the γ-

Fe2O3/Cu2O nanocrystals c-e) STEM-EELS elemental map of Cu, Fe, and O

domains showing the compositional distribution over a larger

area………………………………………………………………………………………….49

Figure 2.12 – UV-VIS optical absorbance spectra of a) Cu2O nanocrystals b)

Cu2O excitonic absorption after 24 hour exposure to air and c) γ-Fe2O3, Cu2O, and γ-

Fe2O3/Cu2O HNCs…………………………………………………………………………50

Figure 2.13 – Percent distribution of isolated, dimer, and oligomer particles

as a function of reaction, time, temperature, and stoichiometry. Inset: Reaction yield

under optimal conditions of 15min, 150°C, and 1 mmol Cu(I) acetate precursor…..51

Figure 2.14 – a, b) Low resolution TEM images of Fe2O3/Cu2O HNCs after

size selective precipitation c) Particle size distribution following size-selective

precipitation of the HNCs…………………………………………………………………52

Figure 2.15 – a) Atomic resolution Z-contrast image of a typical HNC particle

b-d) STEM/EELS elemental maps obtained from Fe L2,3 (blue), Cu L2,3 (red), and O

K edges (green)……………………………………………………………………………53

xi

Figure 2.16 – a,b) HRTEM images of enlarged Cu@Cu2O/Fe2O3 HNCs c-f)

STEM/EELS elemental maps of Cu@Cu2O/Fe2O3 HNCs showing the core-shell

nature of the larger Cu domain and the smaller iron oxide domain………………….54

Figure 3.1 – XPS survey spectra of a) γ-Fe2O3 nanocrystals b) Cu2O

particles and c) γ-Fe2O3/Cu2O …………………………………………………………..66

Figure 3.2 – a) XPS spectra of Fe 2p core-level lines of commercial γ-Fe2O3

and Fe3O4 powders, and the as-synthesized γ-Fe2O3 nanocrystals b) XPS spectrum

of the Fe2p3/2 region of the HNCs and isolated γ-Fe2O3 nanocrystals c) XPS

spectrum of the HNCs O1s region with peak fitting i) O signal from γ-Fe2O3 ii) O

signal from Cu2O iii) O signal from carboxylate ligand d) XPS spectrum of the HNCs

Cu 2p region ……………………………………………………………………………….67

Figure 3.3 – a, b) The secondary electron cut-off region of the γ-Fe2O3/Cu2O

HNCs, their pure components, and the physical mixture of the isolated nanocrystals

c) Valence band edge photoemission spectra of HNCs and their

components………………………………………………………………………………...68

Figure 3.4 – Band energy diagram showing the valence and conduction band

edges and Fermi levels of the HNCs and their constituents. ………………………...69

Figure 3.5 – Optical absorption spectra of a) pure Cu2O b) pure γ-Fe2O3 and

c) γ-Fe2O3/Cu2O HNCs manipulated using the Tauc relation (Ref 58) to determine

their optical bandgaps …………………………………………………………………….70

Figure 3.6 – Scheme of the photocatalytic reactor design and GC/MS product

detection setup …………………………………………………………………………….72

Figure 3.7 – Rate of CO (top) and CH4 (bottom) production from CO2

hydrogenation at different temperatures ………………………………………………..73

Figure 3.8 – Surface photovoltage spectra of as-synthesized a) Cu2O NCs b)

Fe2O3 NCs c) Fe2O3/Cu2O HNCs d) separate components plus HNCs plus

trioctylamine ligand film …………………………………………………………………..74

Figure 3.9 – XPS spectra of the C1s and Fe2p regions before and after

removal of carbon contamination by exposure to X-ray beam for a period of 20

minutes ……………………………………………………………………………………..76

Figure 3.10 – FTIR spectra (left) and PXRD patterns (right) of HNCs before

and after heat treatment at 450°C for 24 hours in air …………………………………77

xii

Figure 3.11 – TEM images of a) HNCs following ligand removal by

calcination at 450°C and b) HNCs following ligand exchange with NOBF4 scale bars-

100nm ………………………………………………………………………………………78

Figure 3.12 – XPS spectra of a) C1s region b) Fe 2p region c) Cu 2p region

and d) UPS spectra of the secondary electron cut-off region of HNC films at a

distance of 5 cm from the UV source in the ZONE cleaner ………………………….79

Figure 3.13 – XPS spectra of a) C1s region b) Fe 2p region c) Cu 2p region

and d) UPS spectra of the secondary electron cut-off region of HNC films at a

distance of <1cm from the UV source in the ZONE cleaner …………………………81

Figure 3.14 – FTIR spectra (left) and TEM images (right) of γ-Fe2O3 NCs

before and after ligand exchange with NOBF4………………………………………….82

Figure 3.15 – a) Extent of MB photocatalytic degradation over various

catalysts as determined by monitoring the main absorption peak at ~ 667nm. b) UV-

Vis spectra of the MB aqueous solution at various intervals over the γ-Fe2O3/Cu2O

photocatalyst……………………………………………………………………………….83

Figure 4.1 – Scheme of Cu2O nanocube formation using sodium citrate as

chelating agent to control particle size…………………………………………………..92

Figure 4.2 – a,b) Dark and bright field TEM images of as synthesized Cu2O

nanocubes with an average size of 82 nm c,d) Dark and bright field SEM images of

Cu2O nanocubes on a TEM grid and glassy carbon electrode respectively………...93

Figure 4.3 – a) Particle size distribution of Cu2O synthesized using 0.75 and

0.25 equivalents of citrate b) PXRD pattern c) UV-VIS absorption spectrum and d)

FTIR spectrum of as-synthesized Cu2O nanocubes…………………………………..94

Figure 4.4 – top) PXRD patterns of Cu2O nanocubes heated at various

temperatures for 1 hour in air bottom) PXRD patterns of Cu2O cubes at increasing

times under 200°C in air…………………………………………………………………..95

Figure 4.5 – a-c) TEM images of as-synthesized Cu2O@TiO2 nanocubes

(scale bar 500 nm) d-f) Ti, Cu, and O elemental signals from EDX linescan of

Cu2O@TiO2 in Panel B……………………………………………………………………96

Figure 4.6 – top) PXRD pattern of as synthesized Cu2O@TiO2 nanocubes

showing absence of TiO2 reflections bottom) PXRD pattern of Cu2O@TiO2

nanocubes following heat treatment at 200°C for 3 days……………………………..97

xiii

Figure 4.7 – a) PXRD pattern of Cu2O@TiO2 composites after calcination at

400°C for 30 minutes under Ar b,c) TEM images of the composites before and after

heat treatment (scale bars 100 and 300 nm respectively)…………………………….98

Figure 4.8 – TEM images of Cu2O nanocubes decorated with various noble

metals a) Au (3%) b) Pt (10%) c) Coated with TiO2 and d) Pd (10%)……………...100

Figure 4.9 – a,c) Photographs of the aqueous phase photocatalytic testing

setup and reactor b,d) H2 evolution as a function of time from Cu2O@TiO2

nanocubes………………………………………………………………………………...102

Figure 4.10 – top) PXRD pattern of Cu2O-Pt@TiO2 particles before testing

bottom) PXRD pattern of Cu2O-Pt@TiO2 particles after testing; (inset) – photograph

of the catalyst suspension showing presence of metallic Cu………………………..103

Figure 4.11 – top) TEM image and PXRD pattern of commercial bulk Cu2O

powder used as a testing reference bottom) TEM image and PXRD pattern of

commercial Cu2O “nanospheres………………………………………………………..105

Figure 4.12 – a) Rates of CO and CH4 over multiple runs from bare Cu2O

nanocubes and Cu2O@TiO2 composites b) CH4 production rates from Cu2O-

Pt(3%)@TiO2 composite (no CO detected) c) CO and CH4 evolution rates on the

Cu2O-Pt(3%) sample over multiple runs……………………………………………….106

Figure 4.13 – a,c) H2 evolution rates as a function of Cu loading in Cu/TiO2

catalysts and H2 evolution profile with time over the P25/Cu (10%) catalyst b,d)

PXRD patterns of P25/Cu (10%) samples before and after testing………………...109

Figure 4.14 – Hydrocarbon and H2 evolution rates over various model

catalysts using different hole scavengers……………………………………………..110

Figure 5.1 – PXRD patterns of as synthesized delafossite CuFeO2 (top) and

CeFe2O4 (bottom)……………………………………………………………………….123

Figure 5.2 – a) PXRD pattern of the mixed 1:1 molar ratio iron/copper

precursor before heating at 900°C a,c,d) SEM images of CuFeO2 particles following

solid state reaction. Panel e) shows a CuFeO2 particle synthesized without the use

of NaOH b) SEM image of a CuFe2O4 particle after solid state reaction and

grinding……………………………………………………………………………………124

Figure 5.3 – a,c) EDX line-scans of CuFeO2 and CuFe2O4 particles showing

the presence of the expected elements (Scale bars 1μm and 2 μm respectively) c,d)

Quantitative signal intensities of the elements detected by EDX confirming expected

stoichiometric ratios of the metals……………………………………………………...125

xiv

Figure 5.4 – a) XPS survey spectrum of CuFeO2 b) PXRD pattern of CuFeO2

following ball-milling to reduce particle size c) PXRD diffraction patterns of CuFeO2

following prolonged heating at 450°C in air d) Diffuse reflectance spectra of CuFeO2

and CuFe2O4 ……………………………………………………………………………..126

Figure 5.5 – a) PXRD patterns of CuFeO2 powder after various amounts of

time at 900°C under Ar b) PXRD patterns of CuFe2O4 powder after calcination at

1000°C in air for various amounts of time ……………………………………………128

Figure 5.6 – TGA-DSC curves of CuFeO2 top) and CuFe2O4 bottom) at a

heating rate of 10°C/min. Arrows indicate temperatures where we performed PXRD

analysis …………………………………………………………………………………...129

Figure 5.7 – PXRD spectra of top) 1:1 molar ratio precursor for CuFeO2 prior

to heating middle) 1:1 molar ratio precursor after heating to 250°C and 550°C for a

period of time corresponding to the synthesis of the material bottom) 1:2 molar ratio

precursor (CuFe2O4) after heating at 450°C for amount of time corresponding to the

synthesis…………………………………………………………………………………..131

Figure 5.8 – a) CO production rate over CuFeO2 over a period of 90 hours

intense illumination b) GC-MS chromatogram of the CO peak showing the presence

of 13CO amongst the products c) 12CO and 13CO signals under alternating low and

high illumination intensities d) CO production rates over CuFeO2 and pure Fe3O4,

Fe3O4/Cu, and pure Cu control sample. Each bar corresponds to a testing run with

an average length of 7 hours……………………………………………………………133

Figure 5.9 – a) Top down SEM image of the CuFeO2 film prior to illumination

b) Cross-sectional SEM image of the CuFeO2 film used to determine approximate

thickness c) PXRD diffraction pattern of the CuFeO2 catalyst after testing showing

the presence of newly formed Fe3O4 and Cu phases (Inset – digital photograph of

discolouration following illumination……………………………………………………134

Figure 6.1 – A general schematic of solar fuels production describing many of

the proposed methods of converting CO2 to fuels using solar energy. The

approximate temperature requirements are color-coded, red = high, yellow =

ambient…………………………………………………………………………………….148

xv

List of Acronyms and Abbreviations

ALD – Atomic Layer Deposition

AM 1.5 – Air Mass 1.5 Solar Spectrum

AMU – Atomic Mass Unit

CB – Conduction Band

CVD – Chemical Vapor Deposition

DSC – Differential Scanning Calorimetry

EDX – Energy Dispersive X-Ray Spectroscopy

EELS – Electron Energy Loss Spectroscopy

FID – Flame Ionization Detector

FTIR – Fourier Transform Infrared

FTO – Fluorine-Doped Tin Oxide

HNC – Hetero-nanocrystal

HOMO – Highest Occupied Molecular Orbital

HRTEM – High Resolution Transmission Electron Microscopy

LUMO – Lowest Unoccupied Molecular Orbital

MB – Methylene Blue

NCs/NPs – Nanocrystals/Nanoparticles

NHE – Normal Hydrogen Electrode

PEC – Photoelectrochemical

PXRD – Powder X-ray Diffraction

RWGS – Reverse Water Gas Shift

SEM – Scanning Electron Microscopy

STEM – Scanning Transmission Electron Microscopy

TEM – Transmission Electron Microscopy

xvi

TGA – Thermogravimetric Analysis

TON – Turnover Number

UPS – Ultraviolet Photoelectron Spectroscopy

UV – Ultraviolet

VB – Valence Band

XPS – X-ray Photoelectron Spectroscopy

1

Chapter 1 – Introduction to Solar Fuels

1.1 Scientific Motivation

Humanity’s reliance on fossil fuels to produce energy and the associated

emissions of greenhouse gases which may have unpredictable effects on Earth’s

climate are arguably the greatest challenges facing society in the 21st century. With

global population expected to reach 10 billion by 2050, it is estimated that we will need

30 – 40 terawatts (TW) of power to maintain our current way of life compared to 17 TW

today.1,2 Currently, close to 80% of our energy needs are derived from the burning of

fossil fuels in the form of oil, coal, and natural gas, Figure 1.1. Technologies based on

renewable resources such as photovoltaics, wind turbines, and biomass conversion

account for only ~ 3% of our total energy production. It is clear that if we want to avoid

damaging changes to Earth’s climate and catastrophic pollution, we must devote

significant efforts to developing cheap, efficient renewable power.

Figure 1.1 US Energy Consumption by Type (Source: NPR,

http://tinyurl.com/pzve2kw, Accessed 20/4/2015)

2

Of the available renewable resources, solar energy stands out in terms of its

abundance. The total solar power irradiating the Earth’s surface is approximately ~ 105

TW per year, meaning that only 0.1% of this resource would be enough to sustainably

meet our needs, provided it can be harvested, converted, and stored.3 However, to date

it remains a significant challenge to efficiently capture and make use of solar power and

to do so in a cost competitive manner in relation to fossil fuels. As chemists, we possess

the necessary tools to develop the materials and processes to make this a reality.

1.2 Basics of Artificial Photosynthesis

As mentioned above, the 3 critical challenges related to utilizing solar power are

energy harvesting, conversion, and storage. Photovoltaic cells, which convert sunlight

into electricity, effectively address only the first two of these issues. Energy storage

continues to be problem in light of the inherent variability of solar irradiation caused by

the time of day and unpredictable weather fluctuations. A possible solution is offered by

the process of photocatalytic solar fuel generation whereby solar energy drives

chemical processes thereby directly converting abundant raw materials such as water

and carbon dioxide into hydrogen or basic hydrocarbons such as formaldehyde,

methanol, and methane, see Figure 1.2.3–5

Figure 1.2 The process of artificial photosynthesis (Reprinted with permission

from Ref (8) Copyright (2012) Royal Society of Chemistry)

3

The direct conversion of solar energy into chemical energy stored in bonds -

termed artificial photosynthesis – is an attractive proposition that allows the issues

associated with electricity storage to be bypassed.6–9 There are typically 3 main

processes that need to take place in an effective photocatalyst for artificial

photosynthesis as shown in Figure 1.3:

1. Light absorption

2. Charge separation and migration,

3. Surface reaction of the photo-generated charges with adsorbed reactant

species at a catalytic centre

In the following, we will discuss how these basic principles affect the

development of materials for efficient solar fuel production.

Figure 1.3 Light absorption, charge migration, and surface reactions in a

heterogeneous photocatalyst (Reprinted with permission from Ref (30) Copyright (2014)

John Wiley and Sons)

Semiconductors are attractive materials for light harvesting. They possess a

band gap: a lack of electronic states between a valence band (VB) filled with ground

state electrons and a conduction band (CB) that is empty of electrons at T=0K.10 The

“size” or energy separation between the top of the VB and the bottom of CB determines

the wavelength of light that a semiconductor can absorb. When light with sufficient

energy shines on a semiconductor, an electron is excited from the VB to the CB, leaving

4

“an electron hole” in the valence band, Figure 1.4. A hole is a quasi-particle that

essentially constitutes the absence of an electron and behaves as a positive charge.

Figure 1.4 Formation of an electron-hole pair in a semiconductor upon excitation

with light (Reprinted with permission from Ref (10) Copyright (2013) John Wiley and

Sons)

Since only photons with incident energy greater than the bandgap can be

absorbed, an overlap between the solar spectrum and the semiconductor’s absorption

profile is crucial for maximizing total light absorption. However, it is not simply a matter

of choosing materials with a narrow bandgap; the positions of the VB and CB edges

with respect to the redox potentials of the reactant adsorbates must also be taken into

account. As an example, the water splitting reaction H2O → H2 + ½ O2 is

thermodynamically unfavourable with ΔG = 237 KJ/mol under standard conditions.11

This corresponds to a ΔE = -1.23 V, meaning that a photocatalyst must have a bandgap

of at least 1.23 eV to drive this reaction. Some other common CO2 reduction reactions

and their minimum required potentials are listed in Table 1.1.12 For reduction reactions,

the CB edge must lie above the lowest unoccupied molecular orbital (LUMO) of the

adsorbed species. In other words the CB potential has to be more negative than the

reduction potential of the acceptor. On the other hand, the transfer of a positive hole to

a hole acceptor molecule requires that the top of the VB lie at a more positive potential

with respect to the highest occupied molecular orbital (HOMO) of the adsorbate. Figure

1.5 provides a helpful visual explanation of these guidelines. These considerations

create a necessary compromise between the opposing requirements of light absorption

and thermodynamic driving force.

5

Table 1.1 Reactions of Interest in Artificial Photosynthesis and Their

Thermodynamic Potentials

Common Reactions of Interest Required Potential, ΔE

H2O → H2 + ½ O2 ΔE = 1.23 V

CO2 + H2O → HCOOH + ½ O2 ΔE = 1.40 V

CO2 + H2O → HCHO + O2 ΔE = 1.34 V

CO2 + 2H2O → CH3OH + 3/2 O2 ΔE = 1.21 V

CO2 + 2H2O → CH4 + 2 O2 ΔE = 1.06 V

On one hand, a narrow band gap is advantageous as it maximizes solar

spectrum absorption. However as discussed above, the VB and CB edges have to span

the reduction and oxidation potentials of the desired reactions necessitating large

bandgaps. There are also overpotentials associated with the kinetic barriers of these

reactions that necessitate even larger driving forces. The general consensus is that for

a single domain photocatalyst, the optimal bandgap is in the range of 2.0 to 2.5 eV as

this allows absorption of some visible light and enough redox driving force for the half-

reactions of CO2 reduction and H2O oxidation.

Figure 1.5 Positions of VB/CB energies with respect to the redox potentials of

surface molecules (Reprinted with permission from Ref (10) Copyright (2013) John

Wiley and Sons)

6

Figure 1.6 shows band edge positions of several commonly studied

semiconductors superimposed onto the reduction potentials of CO2 to different products

versus the normal hydrogen electrode (NHE). From the figure, it becomes apparent that

very few semiconductors possess conduction band edges at potentials negative enough

to reduce to CO2 to the CO2- radical by direct electron injection. However, proton

assisted multi-electron reactions can be thermodynamically driven by a number of

materials to give products such as formic acid and carbon monoxide (HCOOH & CO,

2e-), formaldehyde (HCHO, 4e-), methanol (CH3OH, 6e-), and methane (CH4, 8e-).

Sections 1.4 and 2.2 will further discuss the selection of suitable semiconductors for the

work performed in this thesis.

Figure 1.6 Positions of the VB and CB potentials of various semiconductors at

pH =7 relative to the redox potentials of CO2 reduction to different products (Reprinted

with permission from Ref (10) Copyright (2013) John Wiley and Sons)

Following light absorption and formation of charge carriers in the form of a

negative electron and a positive hole, the processes of charge separation and migration

to the surface become crucial. Many competing recombination pathways exist and

present a critical factor in limiting the efficiencies of photocatalysts. Naturally, the

electron and hole have a Coulombic attraction due to their opposite charges. The

electrostatically bound electron-hole pair is termed an exciton. Notable, exciton

recombination can reach up to 90% within 10ns following excitation, thereby limiting the

7

number of free charge carriers in the material even prior to charge diffusion.11,13 The

typical distance travelled by charges is called the diffusion length, and it depends on the

material. Limiting the size of the photocatalyst to the nanoscale and introducing porosity

are useful strategies in minimizing the distance travelled by the carriers to reach the

surface. As a consequence of the shorter distances travelled the photogenerated

charges have a higher chance to reach the surface without recombining. Charge

recombination can occur inside the bulk of the semiconductor (volume recombination)

or once the charges reach the surface (surface recombination), Figure 1.7. It is typically

caused by crystalline defects in the bulk, which inhibit the migration of charges to the

surface. Amorphous materials and those with grain boundaries have therefore

traditionally been considered therefore a poor choice for photocatalysis. As a result,

well-defined nanocrystalline materials have received a lot of attention with a view on

overcoming these challenges.14,15

Figure 1.7 Formation of an electron-hole pair and various recombination

pathways inside a semiconductor (Reprinted with permission from Ref (15) Copyright

(2010) American Chemical Society)

The final stage of the artificial photosynthetic process is the redox chemistry at

the surface of the catalyst. This is heavily dependent on the available catalytic sites and

8

surface species present on the catalyst, and is perhaps the least clearly understood.

The surface reactions usually require a co-catalyst, typically a noble metal nanoparticle

or organometallic complex to be tightly interfaced with the semiconductor lowering the

kinetic barriers associated with CO2 activation. The metal acts as an electron sink,

accepting the photogenerated electrons which reach the semiconductor surface. For

this to happen, the metal’s Fermi level must be at a less negative potential than the

semiconductor CB, as illustrated in Figure 1.8. The metal’s Fermi level is then shifted

slightly upward to more negative potentials making the composite more reductive.16 This

forms a classic Schottky barrier at the nanoscale where the electrons accumulate in the

metal and are used to do reduction chemistry while the holes remain on the

semiconducting material.17,18 Charges migrating through the bulk of the catalyst and

reaching its surface are expected to participate in redox chemistry with the surface

adsorbates. However, photogenerated charges can also react with the catalyst itself,

performing unwanted redox chemistry that leads to its decomposition in a process

known as photocorrosion.3 Photocorrosion can either be reductive or oxidative

depending on the material and the redox potentials of those reactions. For example,

CdS is susceptible to oxidative photocorrosion by holes due to oxidation of its S2-

accompanied by leaching of Cd2+, Equation 1.1.19

CdS + 2h+ → Cd2+ + S (1.1)

Figure 1.8 A semiconductor-metal junction where the metal acts as an electron

sink and reducing site for adsorbed reactants (Reprinted with permission from Ref (10)

Copyright (2013) John Wiley and Sons)

9

Photocorrosion of the sulfide anion is common in metal sulfide semiconductors

although it has been known to happen in some oxides as well, namely ZnO and Cu2O.19

Preventing photocorrosion has been addressed by employing core-shell structures

where the active catalyst is coated by a stable inert oxide or by making use of bandgap

engineering and co-catalysts as will be discussed in Chapters 2,3, and 4.20,21 Attaching

co-catalysts to a light absorbing component such as a semiconductor also allows one to

test materials for activity in either of the two half reactions of interest, namely CO2

reduction or H2O oxidation by employing sacrificial reagents.19 For example, the setup

shown in Figure 1.9a would be useful in evaluating reduction products if a hole

scavenger such as MeOH or Na2S2O3 is sacrificially oxidized by the holes leftover in the

semiconductor. Similarly, an oxidizing agent such as a metal ion can be used to mop

up photogenerated electrons and allow the evaluation of O2 evolution activity from H2O,

as shown in Figure 1.9b for water splitting.19 It should be noted that even if a catalyst is

active for these half reactions, it may not be active for the overall process due to

recombination or corrosion issues as discussed above. In addition, employing sacrificial

reagents renders the process less “green” as the scavengers are typically obtained from

non-renewable sources. However, utilising sacrificial reagents remains a useful tool in

stabilizing catalysts that may be susceptible to photocorrosion and determining whether

a material is active for the reductive/oxidative reaction that it was intended for.

Figure 1.9 The role of sacrificial reagents in scavenging the majority charge

carriers shown for the water splitting process (Reprinted with permission from Ref (19)

Copyright (2009) Royal Society of Chemistry)

10

Free CO2 molecules are chemically inert with a linear geometry and D∞h

symmetry.10 Adsorption onto a semiconductor or metallic surface offers a way to

activate the CO2 molecule for reduction. The simplest model involves one-electron

transfer from the photoexcited catalyst to the LUMO of CO2 to give a CO2•- radical

species.22,23 Upon the acceptance of the electron, CO2 undergoes a change from linear

to bent geometry (C2v symmetry) and repulsion exists between the lone pairs on the

oxygen atoms and the unpaired electron on carbon.10 Therefore one-electron reduction

of CO2 to CO2•- radical is extremely unfavourable with a chemical potential of -1.9 V

versus NHE, Figure 1.6.24 As a result, very few, if any, semiconductors possess a CB

edge with a potential reductive enough to initiate this reaction. Proton-assisted multiple

electron reductions are significantly less demanding, comparable to the proton reduction

to H2 which is a one electron process, see Equations 1.2-1.7 below (E0 values versus

NHE at pH = 7).10

(1.2) CO2 + 2H+ + 2e- → HCOOH E0redox = -0.61 V

(1.3) CO2 + 2H+ + 2e- → CO + H2O E0redox = -0.53 V

(1.4) CO2 + 4H+ + 4e- → HCHO + H2O E0redox = -0.48 V

(1.5) 2H+ + 2e- → HCOOH E0redox = -0.41 V

(1.6) CO2 + 6H+ + 6e- → CH3OH + H2O E0redox = -0.38 V

(1.7) CO2 + 8H+ + 8e- → CH4 + 2H2O E0redox = -0.24 V

Multiple semiconductors are able to provide potentials that are reductive enough

to drive reactions like those listed above. However, there has been little experimental

evidence of such processes taking place due to the low likelihood of multiple protons

and electrons coming together in concerted fashion. The reactions are likely to proceed

through a series of one-electron transfer steps with the first electron transfer to CO2

thought to be the limiting step.25,26 Unfortunately, research in elucidating the mechanism

of such proton-coupled electron transfer reactions for CO2 reduction is still in its infancy

due to the complexities involved in identifying transient intermediates. From equations

1.2-1.7, it also becomes apparent that there are multiple gaseous and liquid-phase

11

products that can be formed depending on the number of protons and electrons taking

part in the process. Therefore, product selectivity and unwanted side reactions such as

proton reduction to H2 are further issues that must be addressed when trying to design

systems capable of reducing CO2 to hydrocarbons.

The demanding requirements outlined in this section explain why efficiencies of

artificial photosynthetic processes and in particular CO2 reduction have remained low;

finding materials with broad light absorption, efficient charge transport, and active

catalytic surfaces is extremely challenging. Section 1.3 provides a brief progress review

of the solar fuels field from its inception and growth to the current state-of-the-art

developments, with an emphasis on CO2 reduction.

1.3 Literature Overview

The list of materials that have been explored for CO2 reduction is extensive and

includes most metal oxides, metal sulfides, and to a lesser degree the metal nitrides

and phosphides. Many of the photocatalytically active compounds discussed in the

section were initially developed to study water splitting but have recently found use in

CO2 reduction because the structural and property requirements between the two

reactions are very similar. In general, minimizing recombination, high surface areas, and

low cost/toxicity are always desirable characteristics regardless of the class of

materials. Fujishima and Honda were among the first to report an artificial

photosynthetic process in 1972 when they used TiO2 to split water into H2 and O2 under

UV illumination.27 Halmann was the first to report reduction of CO2 in 1978 using a

single crystal GaP cathode, a carbon anode and a CO2 containing aqueous buffer

solution to which a voltage was applied.28 Shortly after, Inoue and Fujishima

investigated the use of semiconductor powders for CO2 reduction to hydrocarbons in

1979.29 Common semiconductors including TiO2, ZnO, CdS, WO3, and SiC were

illuminated by an ultraviolet Xe lamp while suspended in a saturated aqueous CO2

solution. Small amounts of formic acid, formaldehyde, methanol, and methane were

detected with the authors correlating the amount of product to the redox potential of the

CB edges of the various semiconductors. These early breakthroughs gave impetus to

12

the field and over the past two decades, the number of publications related to

photocatalytic CO2 conversion has skyrocketed.

Metal oxides are attractive mainly due to their chemical stability; most oxides are

easy to synthesize, air-stable, and possess large band gaps that help prevent

photodecomposition processes. Oxides with the metal in d0 (Ti4+, Nb5+, W6+)

configuration have especially been closely studied with TiO2 in particular being perhaps

the most commonly used photocatalyst material. Metals in the filled d10 configuration

(In3+, Ga3+, Sn4+) have also been found to be active.30 As opposed to d0 oxides, d10

oxides possess a conduction band mainly composed of hydridized s and p orbitals.

Mixing of these orbitals leads to a large dispersion in k-space and therefore a high

mobility of photogenerated charge carriers.10 The tops of the VB edge in d0 oxides

exhibit mainly O 2p character and are situated at very positive potentials (+ 2-3 V vs

NHE).31 Therefore, their bandgaps are usually too large to properly utilise the solar

spectrum if the CB edge is to be sufficiently negative to reduce CO2. Nevertheless, TiO2

has been the most common semiconductor used in photocatalysis for a number of

reasons. It is an inexpensive, widely abundant material and has been studied

extensively in terms of its absorption, charge transport and surface chemistry.15,32,33 Its

CB and VB levels are positioned such that the CO2 reduction and water oxidation half

reactions are thermodynamically feasible and it is stable to being degraded by

photocorrosion reactions.34 The standard TiO2 material used in photocatalysis is

referred to as P25, an 80%/20% mixture of the anatase and rutile polymorphs.

Electronic synergism between the two phases was proposed as the reason for the

enhanced activity of P25 compared to phase-pure TiO2 polymorphs.35 The main

drawback is its relatively large indirect, bandgap of 3.0 eV which means that TiO2 can

effectively only absorb approximately 5-10% of the solar spectrum, which severely limits

the numbers of photogenerated carriers and therefore its photocatalytic performance.10

Plenty of approaches have been developed to overcome TiO2’s poor absorption

including doping, sensitization, and hetero-structuring. Some notable results based on

TiO2 and other semiconductor photocatalysts for CO2 reduction are presented here and

summarized in Table 1.2. Section 4.3 presents our work on the H2 evolution activity of

Cu nanoparticle decorated P25 TiO2 composites.

13

Table 1.2 Summary of notable semiconductor-based CO2 reduction systems

Catalyst Reaction

Medium

Light

Source

Products Reference

Zn-doped p-type

GaP single crystal

Aqueous CO2

buffered solution

Hg lamp <

365 nm

Formic acid (major),

formaldehyde,

MeOH

Halmann,

ref 29

SiC, TiO2, GaP,

ZnO, CdS, WO3

powders

Semiconducting

powder dispersed

in CO2 aqueous

solution

500 W Xe

lamp with

various

filters

MeOH (major),

formaldehyde, formic

acid, methane

Inoue, ref

30

Anatase TiO2,

230nm diameter

Powder in

supercritical CO2

990 W Xe

lamp

Formic acid Kaneco,

ref 37

Cu-doped ZnO, Li-

doped TiO2

supported on MgO,

Al2O3, and SiO2

CO2 saturated

solution from

KHCO3 (pH 7.5)

250 W Hg

lamp

Acetone (major),

EtOH, MeOH,

formaldehyde, formic

acid, methane

Subrahma

nyam, ref

44

W18O49 nanowires CO2+H2O gas-

phase system

300 W Xe

lamp,

< 420 nm

Methane (major),

EtOH, acetone

Xi, Ref 67

Bi2WO6 hollow

microspheres

CO2 saturated

aqueous solution

300 W Xe

lamp,

< 420 nm

Methanol Cheng, ref

68

Zn2GeO4

nanoribbons

CO2+H2O gas-

phase system

300 w Xe

lamp

Methane Liu, ref 72

ZnS, CdS powders CO2 saturated

aqueous solution

150 W Hg

lamp <

290 nm

Formic acid Kisch, ref

80

CoPi anode,

NiMoZn cathode,

R. Eutropha

CO2 saturated

phosphate buffer

with R. Eutropha

N/A Isopropanol (major)

acetone, pyruvate

Nocera, ref

84

14

Table 1.2 cont`d

Cu2ZnSnS4

modified with Ru-

polymer

CO2 (aq) in

photoelectro-

chemical cell

Xe lamp,

< 400nm

filter

Formate (major), CO Arai, ref 86

p-GaP, pyridinium

electron shuttle

CO2 (aq) in

photoelectro-

chemical cell

200 W Xe,

arc lamp,

cutoff

filters

Methanol Bocarsly,

ref 88

InP/Ru cathode,

TiO2 anode

CO2 + H2O in

photoelectro-

chemical cell

Xe lamp,

< 400 nm

filter

Formate Sato, ref

94

TiO2-xNx hollow

nanocubes

CO2 + H2O gas

phase system

300W Xe

lamp with

AM 1.5

filter

Methane (major),

ethane, propane,

butane

Schaak,

ref 98

C3N4/Ru

multicomponent

structures

CO2 +

triethanolamine in

polar organic

solvents

400 W Hg

lamp, <

400 nm

Formic acid (major),

H2, CO

Maeda, ref

100

The early work on TiO2 artificial photosynthesis was mainly proof-of-concept

experiments utilising deep UV irradiation and bulk or microcrystalline catalysts with low

surface area and correspondingly low product evolution rates. The most typical setup is

a suspension of TiO2 particles in a CO2 saturated aqueous or alcoholic environment.36–

39 Rates in these reports are usually on the order of μmol/g cat/hour but can be

improved by increasing light intensity or CO2 pressure.40 Since the solubility of CO2 in

aqueous solutions is rather low which can lead to competitive H2 evolution, some

researchers have explored the option of using supercritical CO2 as the reaction

medium.37,41 In these studies, no gaseous products were detected although formic acid

was found in the aqueous phase with acidic media found to increase its production.

TiO2 particles embedded in SiO2 matrices have also been looked at with a view on

15

studying the effect of the reaction solvent.42,43 The dispersed TiO2/SiO2 composites

were found to give formate and CO at higher rates than bulk TiO2, with the product ratio

determined by the interaction of intermediates with the type of solvent. The detection of

C2 or C3 products is rare but has been previously reported using TiO2 photocatalysts.

Subrahmanyam et. al. looked at various metal oxides supported on SiO2, Al2O3, or MgO

and saw acetone and ethanol being formed as the major products with methane and

ethane formed in smaller amounts.44 CO2 reduction occurred preferentially on the basic

oxide supported systems although the acidic oxides were more selective for C2

products. The body of work on modifying TiO2 with various metallic co-catalysts is

immense. As mentioned in Section 1.2, the generally favourable effect of metal loading

on TiO2 is ascribed to the ability of metals to act as an electron sink, reducing charge

recombination in the semiconductor. Copper loading on TiO2 has been studied on

multiple occasions and the consensus has been that Cu particles are selective catalysts

for CH3OH production.45–49 Decoration of TiO2 with Pd nanoparticles has also been

shown to lead to formate and CO with prolonged irradiation leading to deactivation of

the catalyst due to oxidation of Pd to PdO.50,51 Modifications of TiO2 with noble metals

have also been investigated since their plasmonic excitations can help extend the

semiconductor’s absorption.52–54 Doping with up to 5 wt% Ag particles led to an increase

in the formation rates of CH3OH and CH4 compared to bare TiO2.53 Sensitization of TiO2

with organic dyes has also been widely studied since the pioneering work of Gratzel on

dye-sensitized solar cells.55 Visible light is absorbed by the dye whose LUMO level is

located a more negative potential than the TiO2 conduction band edge and the

photogenerated electron is injected into TiO2. The dyes used are organometallic

ruthenium or cobalt complexes.56,57 Enhancement of total reduction products using this

approach was seen by Liu et.al and Woolerton et.al. who attributed it to reduced

electron-hole recombination.58,59 Plenty of work has been carried out on coating TiO2 on

various surfaces and supports such as optical fibers, mesoporous materials, and glass

wool. These materials are more suitable to solid-gas interactions where the morphology

of the catalyst can have a significant effect on the rates. Optical fibers increase the path

length of light which comes in contact with the catalyst; TiO2 is usually coated onto the

fibers from solution by dip-coating. Wu et.al. deposited 50 nm thick films of Cu/TiO2

16

inside the walls of glass optical fibers and detected CH3OH production at μmol/g/h rates

demonstrating the utility of this setup.60 TiO2 coated onto glass wool and glass pellets in

conventional batch reactors have also proven effective catalytic architectures.36,61 The

large surface areas and porosities of mesoporous materials such as zeolites and metal-

organic frameworks have proved useful in the production of solar fuels. Substituting Ti

for some of the elements of the zeolites framework by ion-exchange results in a large

number of isolated Ti catalytic sites that were shown to be active in the presence of CO2

and H2/H2O.62–64 More sophisticated systems have recently begun to be reported using

multiple components. For example, Wang et.al. synthesized CdSe sensitized TiO2

particles decorated with Pt co-catalysts for visible light CO2 reduction.65 Electron

injected from CdSe into TiO2 migrated onto the catalytic Pt centers giving CH3OH, CO

and H2. Such architectures are beneficial as they allow each component to be optimized

separately and its effect on catalytic performance monitored.

The above has been a very brief summary of advances made using TiO2

photocatalysts for CO2 reduction. Despite being the most common material for this

application, TiO2 still has issues related to its limited absorption which are keeping

product formation rates lower than what is required for commercial purposes.

Nevertheless, it provides a convenient platform for studying the effects of co-catalysts,

sacrificial reagents, and reaction conditions as it is known to be active under standard

testing conditions. Other metal oxides have also been explored due to the afore-

mentioned limitations of TiO2. ZnO catalysis is probably the second most common

material although its large bandgap of 3.4 eV means that it too is only active under UV

light.66 Since the main motivation for exploring oxides besides TiO2 is their activity

under visible light, the following will mainly focus on narrow bandgap metal oxides.

Tungsten oxide is one of the aforementioned d0 metal oxides with a slightly narrower

bandgap of 2.7 eV. Ultrathin WO3-x nanowires containing oxygen vacancies were

synthesized by a solution phase route and found to photoreduce CO2 to CH4 under

visible light in the presence of water.67 The number of oxygen vacancies could be

controlled by oxidation with H2O2 and it was found that the number of vacancies would

decrease with increasing W oxidation state. Bismuth tungsten oxide microspheres,

Bi2WO6, were found to be more active than bulk Bi2WO6 producing methanol from CO2

17

at a rate of 16.3 μmol g-1 h-1.68 The microsphere morphology offered a large number of

active surface sites and permeability to allow transfer of reactants in and products out of

its porous structure. Notably, no co-catalysts were used to obtain these rates which is a

significant accomplishment especially under visible light irradiation. Catalysts including

niobium, especially the Nb5+ oxidation state have seen increased use in recent years.

Indium niobate, InNbO4, decorated with nickel or cobalt oxides was found to be

moderately active due to hetero-junctions formed with Ni0 and NiO on the surface.69

NaNbO3 loaded with Pt was investigated for reduction of CO2 to CH4 to compare the

effect of nanowires versus the bulk material. 70,71 Both forms were inactive in the

absence of photo-deposited Pt, however the nanowire sample produced significantly

more methane than the bulk. The increase was attributed to good crystallinity, large

surface area, and improved directional charge transfer along the length of the wires.

Another active nanowire material was recently reported by Zhou and co-workers.

Zn2GeO4 nanoribbons were prepared by a solvothermal approach with the addition of 1

wt% RuO2 and Pt as oxidative and reductive co-catalysts respectively.72 The product

formation rate was over an order of magnitude higher in the hetero-junction samples

compared to the bulk. InTaO4 has been explored extensively for water splitting but it

was not until recently that it was tested for CO2 reduction in the presence of water. Pan

et.al. reported a NiO/InTaO4 system capable of converting CO2 to CH3OH.69 InTaO4

was synthesized by a high temperature reaction between In2O3 and Ta2O5 with nickel

deposited by a solution-calcination method to give the most active hybrid catalysts.

Gallium oxide in the alpha polymorph has been shown to have good affinity towards

CO2 due to its surface OH groups.73 Its CB potential is also sufficiently reductive and so

Park et.al. prepared mesoporous gallium oxide and evaluated its activity for CH4

production compared to a commercial sample. A fivefold improvement was noted which

the authors correlated to better CO2 adsorption due to the larger surface area of the

porous sample. Cu2O is another very attractive material with an optimal bandgap which

has been studied for water splitting.74 Unfortunately it exhibits poor charge diffusion

properties and photocorrosion in aqueous solution.75 One of the few reports on Cu2O

used in CO2 reduction was published by Li and co-workers in 2011.76 SiC particles

decorated with Cu2O produced methanol from CO2 under visible light illumination at

18

higher rates than the separate components although the synergistic effect was not well

understood. Oxides such as the ones discussed above are typically simple to

synthesize as high surface area nanostructures and are chemically stable under the

reaction conditions of artificial photosynthesis. Combined with the ongoing

improvements in their catalytic performance, oxides may be the most promising material

for finding solutions to the CO2 reduction challenge on a global scale.

Sulfides such as CdS and ZnS have historically been widely employed as

photocatalysts due to their narrower bandgaps and well-studied uses as colored

pigments. Their VB edges consist of S 3p orbitals which are shifted to more negative

potentials than the corresponding oxides resulting in broader light absorption. As stated

earlier, perhaps their greatest drawback is the tendency of the sulfide anion to undergo

oxidative photocorrosion, which can be ameliorated by using sacrificial reducing agents.

Zinc sulfide is a wide band semiconductor with an energy gap of 3.6 eV and a CB edge

situated at -1.8 V vs NHE at pH 7.10 It has been shown to be capable of reducing CO2 to

formate and carbon monoxide in respectable photonic quantum yields. Henglein et. al.

looked at colloidal ZnS particles in the presence of alcohol scavengers and found that

the photogenerated charge carriers reacted with the various scavengers on a faster

time scale than charge recombination.77 Small 2-5 nm ZnS crystallites also showed

activity for formate production and competitive H2 evolution from CO2 saturated

aqueous solutions.78 Activity was attributed to low densities of surface defects and the

authors found that addition of Na2S as a suppressor of sulfur vacancies improved

product formation rates. A contrasting result was reported by Fujiwara et. al. who found

that sulfur vacancies cause an increase in activity when using CdS nanocrystals as

catalysts.79 The addition of excess Cd2+ created S vacancies on the particle surface due

to adsorption of extra Cd2+ cations. The S vacancies then provided sites for CO2

coordination and subsequent reduction to CO. ZnS particles have also been loaded

onto large surface area SiO2 supports, with the best results being obtained for a 13% by

weight sample.80 Addition of a Pt co-catalyst resulted in production of formaldehyde and

methanol as opposed to the 2e- product formate in the absence of Pt.80 The cooperative

interaction of both sulfides has also been investigated by Kisch et. al. who loaded CdS

particles onto ZnS supports.81 A 5 % CdS loading was found to increase activity 16-fold

19

compared to unmodified ZnS with the improvement attributed to higher charge

separation efficiency in the coupled semiconductor system. CdS, which has a

significantly narrower bandgap (2.6 eV) than ZnS and a more positive CB edge (-0.9 V

vs NHE), was used to reduce CO2 to CO under visible light irradiation using

triethylamine as hole scavenger. With recent advances in nano metal chalcogenide

synthetic techniques, some more sophisticated systems have been explored. For

example, CdS particles were dispersed on the inorganic clay montmorillonite and the

resulting composites showed activity for CH4 and CO production from CO2 dissolved in

NaOH(aq).82 More recently, Armstrong et.al. functionalized the surface of CdS

nanocrystals with an enzyme – carbon monoxide dehydrogenase – and observed CO

production from CO2 under visible light illumination.83 It was found that larger CdS

particles showed no activity due to grain boundaries acting as recombination centers

while the nature of the sacrificial electron donors also had an effect on the rates. In

2014, Nocera and co-workers reported a hybrid bioelectrochemical device where H2

and O2 generated from water splitting using cobalt phosphate and NiMoZn alloy as the

electrodes were combined with CO2 and converted into hydrocarbons by the genetically

engineered bacterium R. Eutropha.84 Other sulfides have received some attention

including MnS which was found to reduce bicarbonate to formate, acetate, and

propionate in solution with a quantum efficiency of 4.2%.85 The authors proposed that

MnS and related minerals may have been responsible for pre-biotic syntheses of carbon

based biomolecules. Bismuth sulfide, Bi2S3 possesses a very narrow bandgap of 1.3 eV

that renders it an attractive solar absorber. When used in conjunction with CdS,

Bi2S3/CdS composites exhibited activity for methanol production, with the highest rates

obtained for 15% Bi2S3 loading. Complex metal sulfides such as CuxAgyInzZnkSm where

the bandgap can be tuned based on the stoichiometric ratio of the various components

have also been explored.86 Decoration of the sulfide with catalytic amounts of RuO2

produced methanol although H2 was employed as the reducing agent instead of H2O. A

similar system was reported by Arai et. al. who used a Cu2ZnSnS4 (CZTS)

semiconductor as the light absorber modified with a Ru-based polymer.87 Electron

transfer from CZTS to the catalyst initiated CO2 reduction with preference over H2O

reduction, however an external bias was used to prevent photo-oxidation.87 As

20

evidenced by the work above, metal sulfides may still play an important part of artificial

photosynthetic systems going forward, particularly in the role of broad solar absorbers

decorated with active metallic co-catalysts.

Metal nitrides and phosphides have also been studied for CO2 reduction

applications, although to a lesser extent than the oxides or sulfides. Nitrides and

phosphides suffer from some of the same issues associated with metal sulfides

including oxygen and moisture sensitivity, and susceptibility to photocorrosion.

Nevertheless, gallium phosphide (GaP) is a very attractive material due to its optimal

bandgap of 2.3eV and highly reducing conduction band edge.88 Some of the earliest

works in the field such as those by Halmann and Inoue looked at GaP electrodes and

found them to be more active than TiO2 or ZnO electrodes under illumination.28,29 More

recently, Barton et.al. used GaP as an electrode in a photoelectrochemical cell for

CH3OH production from CO2 employing pyridine as a homogeneous co-catalyst.89,90

The selectivity of the process for CH3OH was remarkably high, nearing 100%, a factor

which was ascribed to pyridine/pyridinium acting as a one-electron shuttle to

sequentially transfer six electrons from GaP to CO2. Yang and co-workers have recently

reported the synthesis of large scale GaP nanowire arrays by a low cost VLS approach

that could offer an alternative to expensive single crystals of GaP.91,92

Indium phosphide (InP) is another material with a favourable bandgap that was

explored in early work by Canfield et.al.93 Sato and co-workers recently employed an

InP electrode modified with a molecular Ru complex to reduce CO2 to formate in an

electrochemical cell.94 The same authors then expanded this system by coupling the

InP/Ru cathode to a TiO2 based anode to perform the overall process of CO2 reduction

and H2O oxidation.95 Formate production was achieved at a conversion efficiency of

0.03% without an external bias, and while the efficiency is quite low, the authors

confirmed via isotope experiments that CO2 and H2O were indeed the carbon and

proton sources.95

Metal nitrides are most commonly applied as solid solutions in the form of oxide-

nitride alloys such as (Zn1+xGe)(N2Ox) and (Ga1-xZnx)(N1-xOx).96–98 The nitrogen serves

to introduce states inside the bandgap of the wider gap oxide materials thereby

narrowing it and improving light absorption. Oxynitrides such as TiO2-xNx and TaON

21

have also received attention and shown that they are capable of reducing CO2 to

methane and formate respectively.99,100 In the first case, the titanium oxynitride was

prepared from Cu3N nanocube templates that were oxidized to CuO, and the released

nitrogen was incorporated into a TiO2 coating to give the oxynitride. The resulting CuO-

TiO2-xNx catalysts produced methane under solar irradiation at competitive rates even

without noble metal co-catalysts. TaON was used as the semiconductor scaffold in a Z-

scheme type arrangement with organometallic Ru-complexes acting as sensitizer and

reduction catalyst.99 Another interesting material that has garnered a lot of attention

recently is the polymeric carbon nitride, C3N4.101,102 Although not strictly a metal nitride,

its visible light absorption and graphitic structure have attracted researchers, as has the

earth-abundance of its constituent elements. The graphitic C3N4 materials are usually

prepared by high temperature pyrolysis of organic compounds such as urea, melamine

or cyanamide.103–105 The graphitic structure enhances charge migration across the

planar C3N4 sheets and facilitates deposition of co-catalysts that can enhance rates

even further.106

This section has presented a brief overview of the historical and state-of-the-art

advances in artificial photosynthesis and specifically CO2 reduction. The latest focus in

the field is on developing low-cost materials with novel architectures to minimize charge

recombination and protect against degradation. Current state of the art efficiencies are

not high enough to make the process commercially viable, however the field is receiving

increasing global interest and the current pace of development bodes well for significant

breakthroughs in the near future. A more thorough outlook on the future of the field and

opportunities for advancement is given in Chapter 6. Finally, it should be noted that

some of the results presented in this section should be viewed with an abundance of

caution. The detection of carbon containing products such as CO, CH3OH, or CH4 has

historically been taken as sufficient evidence for CO2 reduction. However, isotope

labelling studies with 13CO2 have shown that in some cases carbon contamination

residual to the photocatalyst surface is the source of the products detected and not the

desired 13CO2.107,108 Care must be taken to eliminate residual carbon contamination

during the preparation of the catalysts and isotope labelling experiments are

indispensable in determining the carbon source of the detected products.

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1.4 Focus of This Thesis

Having given an overview of artificial photosynthesis and a brief literature review,

this section will provide a summary description of the work contained in this thesis,

especially focusing on the motivations and over-arching ideas tying the different

chapters together.

The main topic of this thesis is the synthesis, characterization, and investigation

of the photocatalytic properties of earth-abundant metal oxide nanomaterials for solar

fuel production. Much of our motivation for focusing on low cost, abundant materials and

processes was based on the fact that the artificial photosynthesis challenge is global in

nature and any significant breakthroughs would need to make sense from an economic

point of view. We acknowledge that there are alternatives to this approach; developing

high-cost, high-efficiency renewable technologies for suitable applications is another

viable option. To give an example, organic solar cells are a massive topic of research

because the solubility of the active layers allows the cells to be solution-processed by

spin or spray coating at low costs.109 Typical laboratory organic photovoltaics are in the

range of 5-10% efficient at converting solar energy into electricity.110 In comparison,

triple junction solar cells based on III-IV semiconductors have achieved efficiencies of

over 44% under concentrated sunlight.111 These technologies are prohibitively

expensive for large scale implementation, but can find specific applications where cost

is not the primary concern. Similar trade-offs will be encountered in the field of artificial

photosynthesis as improving efficiencies clash with the higher costs of noble metal

catalysts or reaction conditions i.e. high temperatures, pressures, solar concentration

etc. The majority of the work in this thesis is based on materials that currently perform

poorly for solar fuel production but have the potential to be interesting if certain

fundamental challenges are solved. A further distinguishing characteristic of this work is

our focus on gas-phase heterogeneous CO2 reduction as opposed to catalyst

suspensions in CO2 saturated aqueous solutions. Working in the gas-phase allows us to

use high reactant pressures and circumvent the low solubility are correspondingly low

conversion rates of CO2 in water. We also believe that gas-solid heterogeneous

processes will ultimately prove to be safer and easier to implement when applied at

23

industrial scales. We therefore focus on the hydrogenation of CO2 to hydrocarbons such

as methane or methanol in the form of the well-known Sabatier and Reverse-Water-

Gas-Shift (RWGS) reactions, Equations 1.8 and 1.9:

CO2 + 4 H2 → CH4 + 2 H2O (1.8)

CO2 + H2 CO + H2O (1.9)

These reactions are currently driven thermochemically at temperatures on the

order of a few hundred degrees Celsius. Here we attempt to lower the energy demands

of these processes by using sunlight to lower the required reaction temperatures.

Occasionally we do explore aqueous H2 evolution experiments as a preliminary test to

gauge the activity of potential catalysts since H2 evolution imposes many of the same

material property requirements as CO2 reduction.

Generally, each of the following chapters begins with a sub-intro section

providing the reader with background and motivation behind the research. We then

report the synthetic preparation of the materials, followed by structural characterization

and any applicable catalytic performance testing for CO2 reduction or H2 evolution.

Chapters 2 and 3 detail our work on Fe2O3/Cu2O heteronanocrystals (HNCs). In

Chapter 2, we introduce the concept of HNCs and explain how hetero-structuring could

be useful for solar fuel production. The synthesis of the separate component

nanocrystals is presented first, followed by the newly-developed synthesis of the hetero-

nanocrystals. Chapter 3 exhibits the XPS/UPS surface studies of the HNCs, along with

gas-phase CO2 reduction results, ligand removal and liquid phase dye degradation

experiments. All of Chapter 2 and parts of Chapter 3 were published in J. Mater. Chem.

A, 2014, 2, 8525-8533. The following is a list of experimental contributions related to

those chapters:

Nanocrystal synthesis, characterization, and data analysis – Peter

Mirtchev with assistance from Elizabeth Jaluague

Photoelectron spectroscopy – Kristine Liao

STEM/EELS imaging – Qiao Qiao

Surface photovoltage spectroscopy – Zongkai Wu

24

Raman spectroscopy – Peter Mirtchev with assistance from Yao Tian

PI’s – Maria Varela, Kenneth Burch, Stephen Pennycook, Doug Perovic,

and Geoffrey Ozin

Chapter 4 is based on unpublished work on Cu2O nanocubes as a

semiconducting platform for deposition of metal co-catalysts and subsequent H2

evolution experiments from solution with sacrificial reducing agents. Chapter 5 presents

unpublished work on developing mixed iron & copper oxide delafossites (CuFeO2) and

spinels (CuFe2O4) for light-assisted CO2 hydrogenation. Experimental contributions

related to these two chapters are as follows:

Nanomaterial synthesis, characterization and data analysis – Peter

Mirtchev

Photoelectron spectroscopy – Kristine Liao

Gas-phase CO2 hydrogenation and diffuse reflectance – Paul O’Brien

Aqueous phase hole scavenger experiments – Peter Mirtchev with

assistance from Veronika Hoepfner

Ball milling – Navid Soheilnia

Chapter 6 serves as a concluding summary of the work and a discussion of

future directions in developing materials for artificial photosynthesis.

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34

Chapter 2 – Synthesis of Fe2O3/Cu2O Hetero-Structured Nanocrystals

(Reproduced in part with permission from J. Mater. Chem A, 2014, 2, 8525-8533

Copyright 2014 Royal Society of Chemistry)

2.1 Abstract

This chapter describes the synthesis of γ-Fe2O3/Cu2O HNCs using a solution-

phase seeded-growth approach. It begins with an introduction to colloidal hetero-

nanocrystals (HNCs) and their nucleation and growth. The motivations behind the

selection of Cu2O and Fe2O3 are then explained as is the expected bandgap

engineering of a type II heterojunction which was hypothesized to improve their

eventual photocatalytic performance. The synthesis of the separate component particles

is outlined followed by the synthesis of the HNCs. γ-Fe2O3 nanocrystals were used as

seeds for the nucleation of metallic Cu followed by oxidation of the Cu domain to Cu2O

upon exposure to air. Structural characterization in the form of HRTEM, STEM/EELS,

UV-VIS, TGA, PXRD, and Raman data is presented. The iron oxide component was

found to likely be γ-Fe2O3 as opposed to Fe3O4 based on its Raman signature. A max

HNC yield of 72% was achieved by reducing particle growth time at a lower temperature

with respect to the individual component particles. Additional HRTEM images and

STEM/EELS mapping results are presented to further elucidate the mechanism of HNC

formation.

35

2.2 Introduction to HNCs

Colloidal inorganic nanocrystals are at the cutting edge of synthetic advances in

nanochemistry due to the high degree of control with which their sizes, shapes, crystal

structures, and surface functionalities can be engineered. Their unique size-dependent

properties make them a fascinating subject for study from a fundamental point of view

and their versatility makes them amenable to being used in numerous technological

applications including optoelectronics, catalysis, energy conversion, and biomedicine.1–4

Wet chemistry approaches in particular have been very successful in producing a

variety of semiconductor, oxide, and metal nanocrystals (NC’s) for the aforementioned

applications.5 Judicious selection of precursors, solvents, ligands, and reaction time and

temperature allows the thermodynamic and kinetic considerations controlling

nanocrystal growth to be controlled resulting in particles with sub-nm level

reproducibility. However, the continued requirement for materials with improved

performance and multiple functionalities has recently resulted in increased interest in

colloidal hetero-nanocrystals, HNCs, which contain two or more distinct materials in the

same nanoparticle.6–9 These inorganic domains are joined through chemical bonding

interfaces and exhibit heterojunctions without the use of organic molecule linking agents

or capping molecules. Figure 2.1 shows some of the different morphologies of HNCs

that have so far been reported in literature. These include most possible combinations

of semiconductors, metals, and insulators in core/shell, dimer, trimer, and branched

architectures. This increased level of complexity allows HNCs to exhibit multi-

dimensional functionalities beyond the limitations imposed by the structural and

compositional options of mono-component nanocrystals. This opens up many scenarios

where HNCs may be applied making use of their hybrid optical, electronic, magnetic,

and chemical properties. Semiconductor-metallic hybrids have already been shown to

improve photocatalytic performance by reducing charge recombination.10,11 Magnetic-

fluorescent particles have seen growing interest as recoverable bio-imaging labels.12,13

Coating photoluminescent particles with insulating oxides in a core-shell configuration

has been shown to significantly improve luminescence quantum yields.14,15

36

Figure 2.1 Various morphologies of reported colloidal HNCs including core/shell,

dimer, trimer, and oligomer architectures. (Reprinted with permission from Ref (6)


Research on HNCs has expanded tremendously over the previous decade as

evidenced by reviews focusing on different classes of HNC materials, configurations,

and applications.16–18 The synthesis of colloidal HNCs requires an even higher degree

of synthetic ingenuity than that of single component particles. The

thermodynamic/kinetic parameters that typically control nanocrystal growth are here

complicated by the interplay of facet-specific reactivity, interfacial strain and atomic

diffusion or exchange.9 Simple one-component colloidal NCs in general are formed

upon decomposition of molecular precursors containing the desired elements in a liquid

solution of a high boiling point solvent and in the presence of stabilizing capping ligands.

Once the reaction is initiated at a certain temperature, the molecular precursors

decompose to form active ‘monomer’ species, which are the basic building blocks of

nanocrystals. The monomers aggregate to form small clusters that are

thermodynamically unstable and dissolve out of solution, unless they reach a critical

radius and overcome a critical free energy barrier thereby becoming thermodynamically

stable.19 Experimental parameters, such as temperature, heating rate, and the relative

concentration of precursors to stabilizers affects the growth rate of the resulting

37

nanocrystals. Mechanistic insights have suggested that the key to obtaining

monodisperse, reproducible NCs is a separation of the nucleation and growth

processes.20,21 This is typically achieved by promoting a short burst of nucleation via

quick precursor injection at elevated temperature inducing supersaturation of the

monomers. This induces their nucleation into stable seeds followed by a slow diffusion-

controlled growth into larger nanocrystals. In the case of HNCs, two or more different

materials are involved in the nucleation process creating an inorganic interface at the

junction. This is typically achieved by the seeded growth approach where fully-formed

nanocrystals of a primary material are present in solution and act as nucleation points.

The precursors of a secondary material are then introduced into the reaction mixture

and preferentially nucleate onto the existing seeds which is thermodynamically favoured

as opposed to homogeneous nucleation in solution.9 This approach is based on the fact

that the energy barrier ΔGhet, that has to be overcome for heterogeneous nucleation is

lower than the barrier for homogeneous nucleation ΔGhom for the corresponding

material.9 The growth of the second material onto the seed is determined by the Gibbs

free energy surface function, ΔGs, similar to the well-developed theories of hetero-

epitaxial thin film growth onto crystallographically oriented substrates. ΔGs is

determined by the surface energies of the respective materials (γ1 and γ2) and the

interfacial strain energy (γ1,2), Equation 2.1.9

ΔGs = γ1 – γ2 + γ1,2 (2.1)

The first two terms are heavily affected by the binding of surfactants, capping

ligands and coordinating solvents, while the interfacial strain energy is a function of

crystallographic lattice mismatch between the two materials. If the second material

exhibits lower energy surfaces γ1 > γ2, and if the two substances have a good lattice

match such that the γ1,2 term is small, then ΔGs > 0 and the growth will take place in a

layer by layer fashion resulting in uniform coverage and a core-shell type configuration.

This regime is called the FM or Franck – van der Merwe mode and is illustrated in the

top panel of Figure 2.2.

38

Figure 2.2 Illustration of the FM Mode, SK Mode, and VW mode for growth of a

secondary material onto a seed nanocrystal (Reprinted with permission from Ref (9)

Copyright (2010) Elsevier)

If the secondary material exposes higher energy surfaces or if there is significant

lattice mismatch between the two domains, then the secondary material will deposit in

island-like formations as opposed to layer-by-layer in order to minimize its contact with

the surface of the seed. This growth mode is called the Volmer-Weber or VW mode and

results in non-core shell architectures such as the dimers and trimers shown in Figure

2.1. A third regime called the Stranski-Krastinov or SK mode can occur when the

secondary material initially deposits in a layer-by-layer fashion but switches to island

growth when the layers exceed a certain critical thickness and the interfacial strain term

becomes dominant.9

39

The majority of HNCs reported in literature are based on transition metal

chalcogenides because their syntheses are well studied and generally reproducible.

Many reports detail chalcogenides combined with other chalcogenides (CdSe/CdS,

CdSe/ZnS, CdS/EuS)2,22,23, metals (CdSe/Au)24 and oxides (CdS/Fe2O3)25–27 in a variety

of architectures. Functional devices ranging from solar cells28 to photocatalytic reactors

have also been demonstrated.29–31 However, metal chalcogenides suffer from chemical

stability issues and are susceptible to oxidation, photocorrosion32, and thermal

degradation which limits their performance for certain applications. Metal oxides are

typically more resistant to these problems but comparatively fewer papers exist on

metal oxide heterojunctions synthesized by colloidal chemistry methods.33–36 Cao and

co-workers synthesized UO2/In2O3 heterodimers by high temperature solution-phase

annealing of UO2 and In2O3 seeds and suggested that epitaxial growth preferentially

occurs at crystal facets where the first atomic monolayer has the strongest affinity for

the seed nanocrystal.37 Different methods to prepare ZnO/FexOy heterodimers have also

been reported and their magnetic/luminescent properties were investigated.38,39 Cozzoli

et al. synthesized binary γ-Fe2O3/TiO2 HNCs by heterogeneous nucleation of iron oxide

onto the longitudinal facets of anatase TiO2 nanorods in a ternary surfactant mixture

and described their mechanism of formation.40,41 Tremel and co-workers recently

reported the synthesis of core/shell and heterodimer Cu@Fe3O4 nanoparticles using

Cu(II) acetate and Fe(CO)5 as organometallic reagents.42 However, information about

the reaction yield, the extent of oxidation of the Cu component, and the electronic

properties of the particles was not provided.

Here we report for the first time, the synthesis of γ-Fe2O3/Cu2O HNCs by a

solution-phase seeded growth approach and investigate their structure and composition

by HRTEM, STEM/EELS, PXRD, and Raman spectroscopy. γ-Fe2O3 is an n-type

semiconductor with the inverse spinel structure of magnetite (Fe3O4) but with Fe (II)

vacancies in the octahedral sites. The α-Fe2O3 polymorph, hematite, has received a lot

of attention as a photoanode in PEC cells due to having a favorable bandgap (2.5 eV),

excellent stability, elemental abundance and a suitable VB edge position for H2O

oxidation.43–47 In its own right, Cu2O is an extremely promising material as a

photocathode as it is one of the only naturally occurring p-type metal oxides.48 It has a

40

bandgap of 2.2 eV which allows for visible light absorption. It is non-toxic, widely

abundant, cheap, and has a CB edge that is 0.7 eV negative than the hydrogen

evolution potential.49,50 As a result, it has garnered interest for water splitting and to a

lesser extent CO2 reduction.51,52 Despite these promising characteristics, Fe2O3 and

Cu2O have significant limitations that have prevented them from being used in efficient

artificial photosynthetic systems so far. Fe2O3 has poor absorptivity near the bandgap

(α-1 ~ 0.12 μm at λ = 550nm), low charge mobility, short excited carrier lifetime (~10 ps),

and short minority carrier diffusion length (2-4nm).53,54 Cu2O also suffers from short

minority carrier diffusion lengths and poor stability under illumination due to photo-

corrosion in electrolyte solution.55,56 As a result, most of the photogenerated charge

carriers undergo recombination and do not contribute to photocatalytic surface

reactions. This is where nano-structuring can be beneficial; reducing the size of the

oxide domains reduces the distance that charges have to migrate to reach the surface

and makes it comparable to the diffusion lengths which are on the order of a few

nanometers. A further advantage can be realized by utilizing the ability of HNCs to

create novel electronic properties. Depending on the energy offsets between the VB

and CB edges of the two materials, “bandgap engineering” can be used to direct the

separation and localization of charge carriers following photoexcitation.6 Three

disparate regimes can be identified as a result of the band offsets: Type me, Type II,

and Type I1/2 (or quasi Type II), see Figure 2.3. In the Type I regime the bandgap of one

semiconductor lies entirely within the bandgap of the other. Following excitation, the

electron and hole are confined by the energy offsets into the same part of the HNC. This

leads to an increased overlap of electron-hole wavefunctions and a corresponding

increase in the probability of recombination. This regime is undesirable for

photocatalytic applications, although it is useful in situations where radiative

recombination is required such as light emitting diodes. In the Type II regime, the

staggered band configuration results in the electron and the hole being localized on

different domains of the HNC. This spatial separation reduces their wavefunction

overlap and should suppress unwanted charge recombination. This should result in

improved photocatalytic performance as more charge carriers are able to reach the

surface and react with adsorbates.

41

Figure 2.3 Schematic representation of charge carrier confinement regimes in

semiconductor hetero-nanocrystals (Reprinted with permission from Ref (6) Copyright

(2011) Royal Society of Chemistry)

The Type I1/2 or quasi Type II configuration occurs when one of the carriers is

confined while the other is delocalized over both domains because of a negligible

energy barrier offset in either the VB or CB edges. By confining electrons and holes in

separate domains of the HNCs, we aimed to create a novel photocatalytic architecture

with oxidizing and reducing sites incorporated in the same nanoparticle. Figure 2.4

schematically illustrates this idea. At the same time, the preparation of HNCs made of

Fe2O3 and Cu2O has not been reported and would represent a synthetic advance in the

field.

Figure 2.4 Schematic Illustration of a Type II hetero-nanostructure with an

electron rich domain for CO2 reduction and hole rich domain for H2O oxidation

(Reprinted with permission from Adv. Mater. 2014, 26, 4607–4626. Copyright (2014)

John Wiley and Sons)

42

2.3 Results and Discussion

As mentioned above, the compatibility of the crystal structures and lattice

constants of HNC components has a big effect on the resulting morphology with well

lattice-matched systems typically resulting in a core-shell configuration, and larger

mismatches resulting in dimer, and dumbbell-shaped particles due to the interfacial

strain energy. Figure 2.5 shows the crystal structures of γ-Fe2O3, Cu2O and metallic Cu

which is initially nucleated on the Fe2O3 seeds and then oxidized to Cu2O as will be

explained below. All 3 compounds crystallize in a cubic unit cell with lattice constants of

8.34 Å, 4.27 Å, and 3.61 Å respectively. In designing a synthesis for novel

heterostructures, the choice of which material should be used as the seed crystal can

be crucial. In the case of γ-Fe2O3/Cu2O we decided to employ the iron oxide

nanocrystals as seeds due to their reproducible synthetic protocols and their improved

stability under high temperature relative to Cu and Cu2O, which are susceptible to

oxidation and subsequent aggregation.

Figure 2.5 Crystal structures and lattice constants of the three components of

the heterostructured nanocrystals

Iron oxide nanocrystal seeds were prepared by a thermal decomposition of

Fe(oleate)3 in 1-octadecene at 320°C according to a modified literature procedure.57

Fe(oleate)3 was synthesized from FeCl3·6H2O and sodium oleate in a biphasic reaction

mixture and its structure confirmed by 1H NMR and IR spectroscopy, Figure 2.6. No

43

significant impurities were found in the precursor. The product was a waxy red liquid

that was used for Fe2O3 nanocrystal synthesis without additional purification.

Figure 2.6 1H NMR spectrum of Fe(oleate)3; inset – IR spectrum of Fe(oleate)3

Thermogravimetric analysis (TGA) of Fe(oleate)3 indicated a loss of the first

oleate ligand at 200°C followed by loss of the remaining two at 320°C in agreement with

published work, Figure 2.7.58 It has been suggested, that metal oleates decompose by

CO2 elimination to give thermal free radicals which can recombine, decompose into

smaller clusters, or react with other metal carboxylate species to propagate the

decomposition reaction.58 This leads to the formation of small metal oxide nuclei

alongside other byproducts such as H2, and CO although the exact reaction route and

stoichiometry has not been clearly determined. The temperature difference between the

nucleation (250°C – 300°C) and growth (320°C) is likely the reason for the good

monodispersity achieved in this reaction. The resulting iron oxide nanocrystals are

mainly spherical with an average size of 12.6 ± 1.5 nm although some cubic particles

are observed likely due to annealing the nanocrystals for an extended time at the

growth temperature of 320°C.

44

Figure 2.7 TGA scans of Fe(Oleate)3 and Cu(I)acetate showing the initial

decomposition temperatures of the precursors

Figure 2.8 a) Representative TEM image and particle size distribution of

isolated Cu2O nanocrystals b) Representative TEM image and particle size distribution

of isolated γ-Fe2O3 nanocrystals

45

The Cu@Cu2O particles were synthesized according to a modified literature

procedure by reduction of Cu(I) acetate to metallic Cu with trioctylamine at 270°C

followed by post-synthetic oxidation to Cu2O upon exposure to air.59 Spherical particles

with an average size of 13.8 ± 2.6 nm were produced as seen in Figure 2.8b. Initially,

the reaction mixture is a dark burgundy color indicating the presence of metallic copper.

Upon exposure to the atmosphere during workup, the mixture begins to turn dark green

as Cu oxidizes to Cu2O. The oxidation process takes several hours to go to completion

and doesn’t seem to induce any changes to the morphology of the initial particles,

although an increase in size of 1-2 nm can be observed due to oxygen diffusion into the

lattice. During HNC formation, the copper component is introduced by lowering the

injection temperature of the Cu(I) acetate precursor solution to 150°C and allowing for a

period of growth ranging from 15 to 60 minutes. It is interesting to note that nucleation of

metallic Cu is observed at the significantly lower temperature of 150°C in the HNC

synthesis as compared to 250°C when the bare Cu@Cu2O were made without pre-

existing γ-Fe2O3 seeds in solution. This is consistent with heterogeneous nucleation

being more thermodynamically favorable than homogeneous nucleation and suggests

that γ-Fe2O3/Cu2O HNCs form by a seeded growth mechanism. TGA scans of the Cu(I)

acetate precursor (see Figure 2.7) confirm its decomposition temperature in the range

of 120-160°C consistent with the above observation. Attempts to carry out the Cu(I)

acetate reduction on pre-formed γ-Fe2O3 seeds did not result in HNC formation

indicating that Cu nucleation only occurs in-situ. This suggests that the presence of an

oxygen deficient iron oxide phase may be responsible for chemically reducing the Cu+1

precursor to Cu0 with simultaneous oxidation of iron sites to Fe3+ upon exposure to air.

Figure 2.9a,b shows representative low-resolution TEM images of the resulting HNCs.

Three distinct particle morphologies are observed including isolated seeds, γ-

Fe2O3/Cu2O heterodimers, and higher oligomers consisting mainly of Cu2O/γ-

Fe2O3/Cu2O hetero-trimers. The γ-Fe2O3 component of the HNCs has an average

diameter of 11.6 ± 1.4 nm in agreement with the isolated iron oxide particles. In

contrast, the Cu2O components are significantly smaller with average diameters of 8.2 ±

1.9 nm and 7.1 ± 1.3 nm in the dimers and oligomers respectively as compared to 13.8

± 2.6 nm in the isolated particles. Control experiments were performed to confirm that

46

the HNC architecture was not a result of post-synthetic assembly, see Figure 2.9b. A

physical mixture of Cu2O and γ-Fe2O3 particles did not exhibit spontaneous ordering

into hetero-nanocrystals suggesting that the observed dimer and oligomer architectures

are a result of seeded growth of Cu on γ-Fe2O3 in solution.

Figure 2.9 a, b) Low-res TEM images of HNCs and physical mixture of γ-Fe2O3

and Cu2O showing the absence of any ordering into hetero-architectures c) Particle size

distribution of the γ- Fe2O3, Cu, and Cu2O domains in as-synthesized HNCs dimers and

oligomers

Figure 2.10a,b shows HRTEM images of the as-synthesized HNCs offering a

better view of the dimer, and trimer morphologies. To confirm that the particles consist

of distinct iron and copper-containing domains we performed energy-dispersive X-ray

analysis (EDX) on the HNCs as shown in Figure 2.10c. The resulting spectra are

overlaid with the TEM image of the examined particle, which shows the Fe and Cu

signals corresponding to separate iron (left) and copper oxide (right) domains. The

47

powder X-ray diffraction patterns of the pure iron oxide and copper oxide particles and

the HNCs are shown in Figure 2.10d.

Figure 2.10 a) Representative bright-field TEM image of as synthesized HNCs

(scale bar is 20 nm) b) HRTEM image showing isolated, dimer, and trimer morphologies

(scale bar is 5nm) c) EDX line scan across dimer particle showing the Fe-rich and Cu-

rich domains (the spectra are shifted up for clarity, scale bar 5 nm). d) PXRD patterns of

Cu2O, γ-Fe2O3, and γ-Fe2O3/Cu2O HNCs as thin films on a Si wafer e) Raman spectrum

of as-synthesized γ-Fe2O3 nanocrystals showing the prominent A1g phonon mode at 701

cm-1 indicative of γ-Fe2O3

The prominent Cu2O (111) reflection and the absence of metallic Cu reflections

indicate fully oxidized Cu2O particles. The pure iron oxide particles exhibit reflections

assigned to an inverse spinel structure, either Fe3O4 or γ-Fe2O3. This is typically the

case in solution phase syntheses which are limited by the boiling points of common

48

solvents. Formation of the thermodynamically stable α-Fe2O3 requires temperatures

above 400°C. The Fe3O4 and γ-Fe2O3 polymorphs cannot be distinguished by PXRD

because they have very similar crystal structures, with the only difference being the

presence of Fe(II) vacancies in the octahedral site of the gamma phase. Therefore,

Raman spectroscopy was used to distinguish between these phases as has been

reported previously.26 Pure Fe3O4 exhibits a A1g phonon mode at ~ 670 cm-1 which

broadens and shifts to ~700 cm-1 as the sample is oxidized to γ-Fe2O3.26,60 The Raman

spectrum of the pure iron oxide nanocrystals is shown in Figure 1e with the A1g mode

present at 701 cm-1 suggesting that the iron oxide component is likely γ-Fe2O3. The

PXRD spectrum of the HNCs in Figure 1c exhibits broadened reflections corresponding

to γ-Fe2O3. The Cu2O reflections are not apparent likely due to their lower abundance,

smaller size and a low signal-to-noise ratio caused by X-ray fluorescence from the γ-

Fe2O3 component.

Figure 2.11 shows a high resolution TEM image of a single γ-Fe2O3/Cu2O dimer.

The lattice spacing of the two domains were measured to be 0.247 Å and 0.294 Å

corresponding to the (111) and (220) planes of Cu2O and γ-Fe2O3 respectively. Figure

2.11b-e show electron energy loss spectroscopy (EELS) elemental maps tracing the

compositional distribution of a single dimer and over a larger area. The Fe and O

signals overlap perfectly giving the location of the γ- Fe2O3 seeds. The Cu-containing

domains are clearly visible adjacent to the predominant γ-Fe2O3 nanocrystals. The O

content of the Cu2O domains is also confirmed by the blue regions in Figure 2.11c,

confirming their oxidation from metallic Cu to Cu2O. The optical properties of the HNCs

are presented in Figure 2.12. Initially the dominant feature in the optical spectrum is a

Cu d-d transition at 572 nm which rapidly decreases in intensity as a result of oxidation

in air, Figure 2.12a.61 After allowing oxidation to proceed for 24 hours, the appearance

of a new feature at 620 nm is observed, Figure 2.12b. This spectral signature is

attributed to the excitonic transition of Cu2O in the range of 2.0-2.2 eV.62 The UV-VIS

spectra of pure γ-Fe2O3, the fully oxidized Cu2O component and the HNCs are shown in

Figure 2.12c. The HNC spectrum shows features of both constituent spectra including

the Cu2O transition and the tail into the UV characteristic of γ-Fe2O3.

49

Figure 2.11 a) HRTEM image of the γ-Fe2O3/Cu2O nanocrystals showing the

Cu2O (111) and γ-Fe2O3 (220) lattice planes b) HR-STEM image and STEM-EELS

elemental map obtained from the Cu L2,3 (red), Fe L2,3 (blue), and O K (green) edges,

showing the compositional distribution of a single heterodimer. Data acquired in an

aberration corrected Nion UltraSTEM100 operated at 100 kV c-e) STEM-EELS

elemental map of Cu, Fe, and O domains showing the compositional distribution over a

larger area. The color of the image is proportional to signal intensity with red indicating

the strongest signal and blue the weakest

50

Figure 2.12 UV-VIS optical absorbance spectra of a) Cu2O nanocrystals

showing initial Cu d-d transition up to 10 min after exposure to air b) Cu2O excitonic

absorption after 24 hour exposure to air and c) γ-Fe2O3, Cu2O, and γ-Fe2O3/Cu2O HNCs

We then examined the effect of varying the reaction parameters in an attempt to

maximize the yield of HNCs with respect to isolated particles. Figure 3 shows the

percentage particle distribution as a function of varying the reaction time, stoichiometry,

and Cu precursor injection temperature. In determining the yield of HNCs we consider

the proportion of dimers and higher oligomers in the total particle count. Increasing the

growth time at a fixed temperature of 150°C leads to a decrease in the yield of HNCs

from a total of ~65% to just over 30%. This is likely caused by thermal de-attachment of

the Cu component from the seeds upon prolonged heating. The injection temperature of

the Cu(I) acetate precursor also has an effect on HNC yield as seen in the right side of

Figure 4. The HNC yield decreases on increasing the injection temperature from 100°C

to 150°C, and 200°C. This effect is consistent with previous knowledge in the field of

hetero-nanocrystal synthesis; heterogeneous nucleation is facilitated by the seed

surface thereby lowering the activation energy for nucleation. However, using an

injection temperature of 100°C resulted in some colloidally unstable byproduct of bulk

51

Cu2O which had to be removed before the yield was determined. When the reaction

was repeated at 150°C, no bulk byproduct was observed and therefore 150°C was

identified as the optimal injection temperature despite giving a slightly lower HNC yield

than the 100°C reaction.

Figure 2.13 Percent distribution of isolated, dimer, and oligomer particles as a

function of reaction, time, temperature, and stoichiometry. Inset: Reaction yield under

optimal conditions of 15min, 150°C, and 1 mmol Cu(I) acetate precursor

Varying the amount of Cu(I) acetate precursor in the range of 1,2, and 4 mmol

while holding the growth time (15min), temperature (150°C) and amount of Fe(oleate)3

(2mmol) constant also had an effect on the yield. We found that using 1 mmol of Cu(I)

acetate (2:1 Fe:Cu molar ratio) results in higher HNC yield than when a 1:1 Fe:Cu ratio

was employed. This is likely caused by incomplete decomposition of Fe(oleate)3, as

evidenced by unreacted precursor that had to be removed by centrifugation post-

synthesis. Increasing the Cu(I) acetate amount to 4 mmol resulted exclusively in

52

isolated γ-Fe2O3 particles and bulk Cu2O precipitate. Under the optimal reaction

conditions identified in Figure 2.13, an HNC yield of 72% was achieved, consisting of

approximately 51% heterodimers and 21% trimers and higher oligomers. Upon size-

selective precipitation with ethanol, the trimers and larger oligomers can essentially be

removed from solution leading to fractions that are enriched in dimers with a

corresponding amount of remaining isolated γ-Fe2O3 particles, Figure 2.14.

Figure 2.14 a, b) Low resolution TEM images of Fe2O3/Cu2O HNCs after size

selective precipitation removing the majority of trimers and higher oligomers c) Particle

size distribution following size-selective precipitation of the HNCs

Further HRTEM and STEM/EELS analysis of the dimer-enriched samples is

shown in Figure 2.15. The atomic resolution z-contrast image of a typical HNC in

53

Figure 2.15a shows the crystallinity of the system although the interface at the junction

is not epitaxial. Interestingly Figure 2.15c shows the presence of a thin Cu shell around

the γ-Fe2O3 domain, likely from residual Cu. This suggests that Cu nucleates non-

epitaxially coating the entire Fe2O3 seed and then coalesces and crystallizes into a

distinct particle upon thermal annealing in solution. The overlap between the Cu and O

signals suggests that partial or complete oxidation to Cu2O has occurred during workup.

Figure 2.15 – a) Atomic resolution Z-contrast image of a typical HNC particle

b-d) STEM/EELS elemental maps obtained from Fe L2,3 (blue), Cu L2,3 (red), and O K

edges (green)

54

By varying the reaction conditions as discussed above, we can increase the size

of the Cu domains thereby preventing their complete oxidation to Cu2O and resulting in

Cu@Cu2O/Fe2O3 particles. By doubling the amount of Cu(I) acetate precursor injected,

larger polycrystalline Cu particles in the range of 12-17nm are formed compared to the

typical size of 6-10 nm under the optimized reaction conditions, Figure 2.16. The larger

size seems to prevent complete oxidation of the Cu core. Under these conditions, the

Fe2O3 seeds are significantly smaller than the overgrowing Cu domain although the

junction is still preserved. However, a thin shell of Cu around the Fe2O3 was not

observed as in Figure 2.15. This demonstrates some of the synthetic control available in

this reaction; however particles synthesized under the optimized conditions were used

in the applications described in the following chapter.

Figure 2.16 – a,b) HRTEM images of enlarged Cu@Cu2O/Fe2O3 HNCs c-f)

STEM/EELS elemental maps of Cu@Cu2O/Fe2O3 HNCs showing the core-shell nature

of the larger Cu domain and the smaller iron oxide domain

55

2.4 Conclusions

We successfully synthesized colloidal γ-Fe2O3/Cu2O hetero-nanocrystals by

thermal decomposition of Cu(I) acetate on γ-Fe2O3 at 150°C leading to the nucleation of

a metallic Cu domain followed by its conversion to Cu2O upon exposure to air. The

oxidation process could be followed to completion by monitoring the disappearance of

the Cu d-d transition and the growth of the Cu2O exciton peak by UV-VIS spectroscopy.

We achieved a yield of 72% HNCs with approximately 50% dimers and 21% higher

oligomers. Size selective precipitation was found to remove some of the oligomers

resulting in solutions enriched in dimers. HRTEM, EDX, and STEM/EELS mapping

confirmed that the particles are comprised of joined but distinct γ-Fe2O3 and Cu2O

domains without alloying or mixed phases. PXRD was inconclusive, but Raman

spectroscopy suggested that the iron oxide domain was likely γ-Fe2O3 as opposed to

Fe3O4. STEM/EELS showed residual copper around the iron oxide seeds indicating that

the particles likely form in a core-shell configuration during growth and segregate into

dimers and trimers upon annealing.

2.5 Experimental Details

Chemicals. Iron trichloride (FeCl3·6H2O, 97%), sodium oleate (82% fatty acid

basis) Cu(I) acetate 97%, trioctylamine (TOA, 98%), 1-octadecene (ODE, tech. 90%),

oleic acid (OA, tech. 90%), iron (II) oxide nanopowder, iron (II,III) oxide nanopowder,

and anhydrous organic solvents were purchased from Sigma Aldrich and used without

further purification.

Synthesis Procedures. All synthetic manipulations were done using standard

airless techniques. Fe-oleate was synthesized according to a previously reported

procedure.57 Iron trichloride (4.32 g, 16.0 mmol) and sodium oleate (14.6 g, 48 mmol),

were dissolved in 110 mL of 4:3:7 ethanol:water:hexane mixture and refluxed at 70°C

under Ar for 4 hours. The solution was cooled to room temperature; the upper organic

layer was separated and washed twice with 20 mL distilled water. Solvent was removed

on a rotary evaporator giving a viscous deep red liquid. The product was dried at 70°C

in a vacuum oven for 48 hours, resulting in Fe(oleate)3 in the form of a waxy red solid.

56

γ-Fe2O3 Nanocrystals – γ-Fe2O3 nanocrystals used as seeds in the synthesis of

HNCs were prepared according to a modified literature procedure.57 2g Fe(oleate)3, 14

mL 1-octadecene, and 0.35 mL oleic acid were added to a 100 mL 2-neck round bottom

flask and heated to ~60°C for 10 minutes to solubilize Fe(oleate)3. Using a spherical

heating mantle, the reaction mixture was heated to 320°C over the course of 20 minutes

at an average rate of 13°C per minute, under Ar flow. The temperature was maintained

at 320°C for 30 minutes, after which the reaction mixture was cooled to room

temperature and 5 mL toluene added. The nanocrystals were precipitated by addition of

excess EtOH, centrifuged at 7800 rpm for 20 minutes, and redispersed in heptane. Two

redispersion/precipitation cycles were performed to remove excess Fe(oleate)3.

Cu@Cu2O Nanocrystals – Cu@Cu2O core-shell nanocrystals were prepared

according to a modified literature procedure.63 0.245g Cu(I) acetate, 7.5 mL

trioctylamine, and 2 mL oleic acid were added to a 50 mL 3-neck round bottom flask in a

N2-filled glovebox. The flask was connected to a Schlenk line and degassed under

vacuum at 60°C for at least 30 minutes. The flask was then filled with Ar, heated to

180°C at a rate of 12°C per minute, and kept at that temperature for 45 minutes. The

reaction mixture was then heated to 270°C at 10°C resulting in a color change to deep

burgundy indicative of nucleation of elemental Cu. Particle growth was continued at

270°C for 45 minutes. The mixture was then cooled to room temperature, 5 mL of

toluene was added and the particles were precipitated by centrifugation in excess EtOH

at 7800 rpm for 20 minutes, followed by redispersion in heptane. Complete oxidation to

Cu2O was observed within 48 hours upon storage in air.

γ-Fe2O3/Cu2O HNCs – The γ-Fe2O3 seeds were synthesized as described above

but the particles were not isolated. Instead, following 30 minutes at 320°C, the reaction

mixture was cooled to 150°C, at which point a degassed, 60°C solution of 0.123 g Cu(I)

acetate, 7.5 mL trioctylamine, and 2 mL oleic acid was rapidly injected via metal

syringe resulting in a temperature decrease to ~ 120°C. The mixture was then rapidly

heated back to 150°C and particle growth was continued for 15-60 minutes. After the

completion of the growth period, the flask was cooled to room temperature, 5 mL

toluene was added, and the HNCs were precipitated by centrifugation in excess EtOH

57

at 7800 rpm for 20 minutes, followed by redispersion in heptane. Three

redispersion/precipitation cycles were performed to remove unreacted starting material

and free ligand.

General Characterization - TGA curves were acquired on a TA Instruments

Q500 thermogravimetric analyzer at a constant ramp rate of 5°C under N2 atmosphere.

UV-VIS absorption spectra were recorded on a Perkin Elmer Lambda 900 UV/VIS/NIR

spectrophotometer in dilute heptane solutions. NMR spectra of the precursors were

acquired on a Varian Mercury 400 MHz spectrometer in CDCl3.

Electron Microscopy - Low resolution TEM images were acquired on a Hitachi

H-7000 conventional TEM operating at 100kV. HRTEM images and STEM EELS

spectra were acquired on a Hitachi H-3300 ETEM, a JEOL JEM 2010 operating at

200kV and a Nion UltraSTEM100 operated at 100 kV and equipped with a Gatan Enfina

spectrometer. Sample preparation involved dropping a dilute nanocrystal solution on a

carbon coated Ni TEM grid. EDX analysis was performed on a Hitachi S-5200 SEM

operating in TEM mode using an Oxford Inca detector. Particle size and yield

determination was done manually using the free ImageJ software on a minimum of 200

particles. STEM HAADF images were obtained from an aberration corrected Nion

UltraSTEM200 dedicated STEM operating at 200 kV. STEM EEL spectra were obtained

from an aberration corrected Nion UltraSTEM100 dedicated STEM operating at 100 kV.

Powder X-Ray Diffraction - Powder X-ray diffraction patterns were recorded on

Siemens D5000 and Bruker D2 Phaser instruments using CuKα line as excitation

source. Samples were prepared by drop-casting a concentrated nanocrystal solution

onto Si(100) substrates to give films of at least 1 micrometer in thickness.

Raman Spectroscopy - Raman spectra were measured in backscattering

configuration utilizing a 532nm solid-state laser, Tornado Hyperflux U1 spectrometer,

and a cooled CCD detector. The spectral resolution was 5 cm-1 and the beam size on

the sample was 10 microns. The laser power was 0.5 mW to avoid laser induced

transition of γ-Fe2O3 to alpha-Fe2O3. Raman analysis samples were prepared by drop-

58

casting concentrated nanocrystal solutions onto a silica glass substrate to give a film of

at least 1 micrometer in thickness.

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64

Chapter 3 – Electronic Properties and Applications of Fe2O3/Cu2O

Heterostructured Nanocrystals

(Reproduced in part with permission from J. Mater. Chem A, 2014, 2, 8525-8533

Copyright 2014 Royal Society of Chemistry)

3.1 Abstract

This chapter includes further characterization of Fe2O3/Cu2O nanocrystals and

their first use as photocatalysts for CO2 reduction. Photoelectron spectroscopy data is

used to determine the absolute VB and CB energies of the particles and compare them

to known standards confirming the desired type II band alignment. The oxidation states

and chemical environment of the surface elements are determined and found to agree

with the results of Chapter 2. Surface photovoltage experiments of the as-synthesized

nanocrystals suggested that the organic ligand shell was preventing reactions with the

desired adsorbates from occurring at the surface and giving false positives in gas-phase

CO2 reduction testing. We then investigated methods to remove the organic capping

molecules including solution-phase ligand exchange, thermal treatment, and ultraviolet

photolysis. We examined the effects of the UV photolysis process on the band energies

of our materials by XPS/UPS. When used for organic dye degradation in aqueous

solution, the HNC performed more effectively than their separate components, though

activity for gas-phase CO2 conversion was not observed.

65

3.2 Photoelectron Spectroscopy

In order to fully determine the chemical and electronic properties of γ-

Fe2O3/Cu2O HNCs we performed an X-ray/ultraviolet photoelectron spectroscopy (PES)

study of our materials. PES is a useful, non-destructive tool for studying the chemical

and electronic structure of nanocrystalline samples as it can provide information about

elemental composition, oxidation state, and density of states near the Fermi level.1 The

presence of both Cu and Fe in the HNCs was confirmed by the characteristic Fe and Cu

2p doublets in the XPS survey scans, see Figure 3.1. We further employed XPS to

ascertain that the iron oxide component of our HNCs was indeed γ-Fe2O3 as

determined by Raman spectroscopy. XPS has previously been used to differentiate

between Fe2O3 and Fe3O4 based on the difference in binding energy between Fe2+ and

Fe3+.2,3 Figure 3.2 shows the Fe 2p core-level peaks of the as-synthesized pure iron

oxide nanocrystals and commercial γ-Fe2O3 and Fe3O4 nanopowders analyzed under

the same conditions. The binding energies of the Fe 2p1/2 and Fe 2p3/2 peaks are

summarized in Table 3.1 which indicates that the nanocrystalline iron oxide lines closely

correspond to the commercial γ-Fe2O3 powder thereby confirming our assignment.

Table 3.1 Binding energy (eV) of Fe 2p core-level lines in γ-Fe2O3 nanocrystals,

γ-Fe2O3/Cu2O HNCs, and commercial iron oxide nanopowders

Sample Fe 2p3/2 Energy Fe 2p1/2 Energy

γ-Fe2O3 Nanocrystals 710.2 eV 723.9 eV

γ-Fe2O3/Cu2O HNCs 709.4 eV 723.0 eV

Commercial γ-Fe2O3

Nanopowder

710.6 eV 724.4 eV

Commercial Fe3O4

Nanopowder

710.9 eV 724.8 eV

66

Figure 3.1 XPS survey spectra of a) γ-Fe2O3 nanocrystals b) Cu2O particles and

c) γ-Fe2O3/Cu2O

The Fe 2p3/2 region of the HNCs is shown in Figure 3.2b. The main peak is

shifted to a lower binding energy of 709.4 eV, as a result of the presence of the Cu2O

domain. A small shoulder at 706.9 eV is also present which may be assigned to Fe0

likely at the interfacial junction region. Figure 3.2c shows the O1s region of the HNCs.

The peak can be de-convoluted into three component peaks corresponding to oxygen in

Fe2O3 and Cu2O environments, and the carboxylic group of surface oleate ligands, in

agreement with database values. Finally, the Cu 2p region of the HNCs is shown in

Figure 3.2d. The Cu 2p3/2 and 2p1/2 lines are present at 932.9 and 952.7 eV

respectively, consistent with Cu+ literature values.4 In contrast to O’Brien et. al. we did

not observe any satellite peaks in the range of 934-940 eV that would indicate a CuO

layer on the Cu2O surface.4

67

Figure 3.2 a) XPS spectra of Fe 2p core-level lines of commercial γ-Fe2O3 and

Fe3O4 powders, and the as-synthesized γ-Fe2O3 nanocrystals b) XPS spectrum of the

Fe2p3/2 region of the HNCs and isolated γ-Fe2O3 nanocrystals c) XPS spectrum of the

HNCs O1s region with peak fitting i) O signal from γ-Fe2O3 ii) O signal from Cu2O iii) O

signal from carboxylate ligand d) XPS spectrum of the HNCs Cu 2p region

The electronic structure of hetero-nanocrystals is of particular interest with

respect to using these materials in photocatalysis or photovoltaics. Using ultraviolet

photoelectron spectroscopy we probed the density of states near the Fermi level, which

allows extraction of the Fermi energy and valence band (VB) maximum. Figure 3.3

shows the secondary electron cutoff peak of our materials, which allows the

determination of the work function and Fermi level. The HNC sample’s valence band

maximum (VBM) is found between the two isolated components suggesting electronic

68

contact between the two domains. A physical mixture of the two isolated nanocrystals

was also evaluated. In such a mixture there is no electronic contact between the two

components and UPS will only detect electrons with the lowest kinetic energy

corresponding to the γ-Fe2O3 component in our system. Figure 3.3 shows that the

secondary electron cutoff edge of the physical mixture spectrum overlaps that of the γ-

Fe2O3. The valence photoemission spectra in Figure 3.3 allow determination of the VBM

energy with respect to the Fermi level. The densities of states of the HNCs originate

mostly from the γ-Fe2O3 component. The small shoulder at 1.5 eV in the spectra of γ-

Fe2O3, the HNCs, and the physical mixture corresponds to a Fe2+ satellite peak caused

by reduction of Fe3+ by the beam. The Cu2O valence band spectrum indicates that the

valence band electron density extends all the way to Fermi level which is consistent with

a fully occupied d band.1

Figure 3.3 a, b) The secondary electron cut-off region of the γ-Fe2O3/Cu2O

HNCs, their pure components, and the physical mixture of the isolated nanocrystals c)

Valence band edge photoemission spectra of HNCs and their components

69

Having determined the position of the valence band maxima of the HNCs and

their constituents by UPS, we can construct an electronic band energy diagram as a

step towards understanding the charge carrier behavior in our system. Figure 3.4 shows

the Fermi levels, and conduction and valence band energies of the HNCs, their

individual components, and the corresponding commercial samples. The positions of

the conduction band (CB) minima were calculated by adding the bandgap as

determined by UV-VIS spectroscopy, see Figure 3.5, to the valence band maxima found

by UPS.5

Figure 3.4 Band energy diagram showing the valence and conduction band

edges and Fermi levels of the HNCs and their constituents. The commercial samples

of copper and iron oxides were evaluated under the same conditions for comparison

70

Figure 3.5 Optical absorption spectra of a) pure Cu2O b) pure γ-Fe2O3 and c) γ-

Fe2O3/Cu2O HNCs manipulated using the Tauc relation (Ref 58) to determine their

optical bandgaps

The Fermi level and valence band maximum of the HNCs is found to be between

those of γ-Fe2O3 and Cu2O pointing to contributions from both components. In

agreement with the literature on the bulk materials, we found that γ-Fe2O3 and Cu2O

nanocrystals are intrinsically n-doped and p-doped, likely from anion and cation

vacancies, respectively. Considering the staggered type II band alignment, a

photoexcited electron in the conduction band of Cu2O would relax to the conduction

band of γ-Fe2O3, promoting its separation from the hole in the Cu2O valence band.

Alternatively a two-photon Z-scheme could also be observed where initial excitation of

γ-Fe2O3 and relaxation into the VB of Cu2O is followed by absorption of a secondary

photon and promotion to its CB analogous to the process in Photosystems I and II.

71

3.3 Gas-phase CO2 Reduction

The overall reaction between CO2 and H2O is highly endergonic, with the majority

of the energy used in the kinetically more difficult water oxidation reaction. Therefore,

catalysts that simultaneously drive CO2 reduction and H2O oxidation without an external

bias are rare and operate at very low efficiencies.6 Vast improvements in H2O splitting

systems have opened the possibilities for a sustainable and economically competitive

supply of H2 as a reactant gas. The hydrogenation of CO2 with H2 is thermodynamically

favourable relative to the reaction between H2O and CO2. Thus, H2 generated

separately via solar-driven water splitting can be used in the subsequent photocatalytic

reduction of CO2 to maximize the potential of the harvested sunlight. CO2 reduction

photocatalysts that operate in an H2 environment at moderate temperatures provide

valuable insights into CO2 reduction mechanisms and provide opportunities for the

discovery of active cathode components in a scalable artificial photosynthetic process.

The majority of CO2 reduction photocatalysts reported in the literature operate at room

temperature or 80 °C for aqueous and gas phase reactions, respectively. Here we

evaluate γ-Fe2O3/Cu2O HNCs for photo-assisted CO2 reduction using H2 as the

reducing agent in a batch process. CO2 hydrogenation to methane has been well known

since the beginning of the 20th century. It was first discovered by Paul Sabatier and

involves the reaction of CO2 and H2 at high temperatures (300-400°C) and pressures

over nickel or ruthenium heterogeneous catalysts dispersed on solid supports, Equation

3.1.7 The reaction is exothermic, with an enthalpy of ΔH = -165 KJ/mol. The process

has been proposed for use onboard spacecraft to produce water for human use from

waste H2 generated by water electrolysis and CO2 emitted by respiration.

CO2 + 4 H2 → CH4 + 2 H2O (3.1)

The reverse water gas shift reaction can also occur under similar conditions,

Equation 3.2. The methane and carbon monoxide generated by these reactions can be

fed into existing natural gas infrastructure or added to Fischer Tropsch processes

respectively. Here we attempt to lower the required temperatures for generating these

products from CO2 by partially driving these processes with solar energy.

72

CO2 + H2 CO + H2O (3.2)

Films of γ-Fe2O3/Cu2O were drop-cast from hexane solutions onto glass

substrates, dried in vacuo and introduced into the reactor. Figure 3.6 shows a

schematic diagram of the multi-reactor and the spectral output of the metal halide lamp

used in this study.

Figure 3.6 Scheme of the reactor design and GC/MS setup

Samples were irradiated for 16 hours following which products were manually

sampled from the reactor. The full details of the testing procedure are given in the

Experimental section. The presence of CO and CH4 was detected at nmol/hr rates, as

shown in Figure 3.7. The empty reactor and reactor+sample were first tested at room

temperature providing us with background rates of hydrocarbon production. Under

illumination and in the absence of additional heating, the sample heated up to 67°C and

product concentrations approximately doubled. Even with an empty reactor, the product

rates were higher under illumination relative to in the dark. However, no increase was

detected in the presence of the catalyst relative to the empty reactor under light. The

detection of hydrocarbons in the absence of a catalyst pointed to a background level of

73

carbon contamination likely arising from rubber tubing, O-rings, or sample

contamination from handling. Under the combined effects of heating and illumination

such that the thermocouple in contact with the sample read 150°C, the rates increased

several fold. In the case of CO, the highest rate of 30 nmol/hr was obtained for the

empty reactor under illumination indicating that the catalyst had no effect. In the case of

CH4, a rate of 2 nmol/hr was reached for the sample under illumination which was

twofold higher than the catalyst in dark or the empty reactor. However, we suspected

that the CH4 produced originated from degradation of the organic capping ligands under

the combined effects of heat and light.

Figure 3.7 – Rate of CO (top) and CH4 (bottom) production from CO2

hydrogenation at different temperatures

74

We hypothesized that the organic ligand shell on the surface likely prevents

adsorption of reactant gases and prevents charges from reacting at the particle surface

as well as providing false positives in hydrocarbon detection. To study whether the

photoactivity of the samples was inhibited by the ligands, we employed surface

photovoltage spectroscopy (SPS). SPS is a contactless technique that probes contact

potential difference changes (ΔCPD) in thin films upon excitation with light.8 SPS

spectra allow photochemical charge separation to be monitored concurrently with the

sample’s light absorption. The dominant ΔCPD signal is expected to correlate with

excitation of electrons from the valence to the conductions band under illumination with

light energy comparable to the film’s bandgap.

Figure 3.8 Surface photovoltage spectra of as-synthesized a) Cu2O NCs b)

Fe2O3 NCs c) Fe2O3/Cu2O HNCs d) separate components plus HNCs plus trioctylamine

ligand film

75

Comparisons of SPS and UV/VIS absorption spectra for pure Cu2O, pure Fe2O3

NPs, and Fe2O3/Cu2O HNCs are shown in Figure 3.8a-c. The onset energy of the SPS

spectra and UV/VIS absorption do not correspond to each other. This suggests the

photovoltage signal in the SPS spectrum is not caused by excitation of the

nanoparticles. All three samples produced similar SPS spectra, Figure 3.8d. We

suspect that the SPS signal is from the excitation of FTO film itself, and that the organic

ligand trioctylamine can accept the holes from the FTO therefore magnifying the

response of FTO. This was confirmed in the SPS spectrum of trioctylamine on FTO in

Figure 3.8d. This spectrum reproduces all features of the nanocrystal spectra. In

summary, the SPS experiment did not observe photochemical charge transfer in the

nanoparticle samples. The likely reason is that electron or hole transfer from the NPs to

the substrate is blocked by the organic surfactants, which form an insulating layer.

Having observed the deleterious effects of the organic capping layer on the

photoelectrochemical properties of the particles, we performed additional XPS

measurements looking at C1s region which would indicate the presence of surface

carbon species, Figure 3.9a. The prominent peak at a binding energy of 286 eV is due

to carbon in sp3 configuration and the shoulder at 289 eV likely indicates the presence

of carbonyls. Surface carbon in XPS can be sputtered away by exposing a region of the

sample film to continuous X-ray irradiation which is energetic enough to break carbon-

carbon bonds. Over a period of 20 minutes, we sputtered away approximately 3nm of

carbon and the resulting XPS spectrum is shown in Figure 3.9c. A clear twofold

reduction in the signal intensity is observed as is the disappearance of the shoulder at

289 eV. With increased cleaning time, it is expected that the surface carbon can be

largely removed from the surface. We also looked at the effect of the carbon shell on

the other elemental peaks present in the sample. The Fe2p region before and after

sputtering is shown in Figure 3.9b,d. After cleaning, the signal to noise ratio is

significantly improved and identification of the fine structure of the Fe peaks is much

easier. Unfortunately, removing carbon contamination in this way is not practical as it is

time consuming and only works on exceedingly small areas of the sample. We therefore

set out to investigate other methods of removing surface carbon ligands.

76

Figure 3.9 XPS spectra of the C1s and Fe2p regions before and after removal of

carbon contamination by exposure to X-ray beam for a period of 20 minutes

3.4 Ligand Removal & Dye Degradation

Our first attempt to remove the ligands inherent in the particles synthesis was

based on a heat treatment of the nanocrystal film at elevated temperature. Such

calcination is straightforward to do as it simply pyrolyzes any surface carbon species

although leftover carbon may still be present as a “carbon black” residue. Another

concern is that the prolonged heating in air may alter the crystal structures of the

nanoparticles. Alloying of the separate HNC domains was also a concern at higher

temperatures, and therefore we selected 450°C as the calcination temperature, which

should be high enough to burn off surface carbon but not high enough to destroy the

hetero-structures. As-synthesized nanoparticles in heptane were drop-cast on Si

substrates and calcined at 450°C for 24 hours. Surface carbon content was monitored

by FTIR. The as-synthesized particles are capped with a mixture of oleic acid and

77

trioctylamine as evidenced by the dominant C-H stretch signal in the FTIR spectrum in

Figure 3.10. Following heat treatment, the FTIR is essentially featureless which

indicates complete removal of any carbon species from the surface of the film. The

PXRD patterns of the film before and after heating are shown in the right panel of Figure

3.10. The particles do not seem to have undergone a phase change despite the

prolonged heating at elevated temperature. The reflections of the iron oxide phase are

still dominant and the signal-to-noise ratio has improved due to the removal of the

amorphous carbon shell.

Figure 3.10 FTIR spectra (left) and PXRD patterns (right) of HNCs before and

after heat treatment at 450°C for 24hours in air

The heat treatment also had certain drawbacks. The removal of organic

surfactants leads to a loss in solubility and film quality subsequently suffers. The

prolonged heating in the absence of separating ligands also leads to sintering of the

particles and a loss of the HNC architecture. Figure 3.11a shows a TEM image of the

HNC particles after heat treatment. The nanocrystals have aggregated to sizes

exceeding 100 nm and the well-defined nature of the HNC domains has not been

preserved. As a result, thermal calcination of nanoparticle films does not seem to be an

effective method of ligand removal even though it is capable of removing most of the

surface carbon species.

78

Figure 3.11 TEM images of a) HNCs following ligand removal by calcination at

450°C and b) HNCs following ligand exchange with NOBF4 scale bars- 100nm

We then began exploring other methods of removing carbon contamination from

thin films of inorganic nanoparticles. Our experience with electron microscopy led us to

think about existing sample cleaning processes for hydrocarbon elimination from

sample grids in SEM and TEM. These treatments efficiently remove hydrocarbons

under reduced pressure and deep UV irradiation to break carbon-carbon bonds.

Oxygen or air plasma then reacts with the hydrocarbons to give species such as H2O,

CO, CO2 which are pumped away from the sample chamber. We used Hitachi’s ZONE

SEM cleaner and evaluated its effect on surface carbon contamination by XPS and

UPS. We tested the ZONE cleaner’s performance at two different distances from the UV

source and investigated the effect of time on carbon removal. The sample film was

placed as close as possible to the lamp (<0.5cm) or relatively far away (5cm). XPS

spectra were recorded after 5, 10, 20, and 50 total minutes of treatment. The data when

the sample was held relatively far from the lamp is summarized in Figure 3.12.

79

Figure 3.12 XPS spectra of a) C1s region b) Fe 2p region c) Cu 2p region and d)

UPS spectra of the secondary electron cut-off region of HNC films at a distance of 5 cm

from the UV source in the ZONE cleaner

There is a pronounced decrease of approximately 30% in the intensity of the

carbon signal shown in Figure 3.12a after cleaning for a total of 50 minutes. Carbon

removal seems to quite efficient up to 10 minutes but the effect decreases with

prolonged cleaning. Therefore the treatment was not continued beyond 50 minutes. The

shoulder peaks corresponding to carbonates were still present, indicating that those

species were not effectively removed. The Fe and Cu 2p regions in Figure 3.12b,c show

a significant improvement in signal intensity due to carbon removal after only 5 minutes

of ZONE cleaning. Subsequent cycles seem to have no big effect on signal strength.

The secondary electron cutoff region of the UPS spectrum is shown in Figure 3.12d.

The left-hand x-intercept is used to calculate the work function of the material, and so

80

the observed shifts show how easily the work function can be influenced by extraneous

surface species. The shoulder peaks are likely due to substrate effects, but there is a

clear shift to lower binding energies which decreases with increasing treatment time.

We then performed a similar XPS/UPS analysis while holding the sample film as

close as possible (<1cm) to the UV lamp inside the reactor. Carbon removal was a lot

more efficient compared to when the sample was held at 5cm from the lamp, Figure

3.13a. The carbon signal is reduced to roughly 1/3 of its original level after 20 minutes

treatment, and is essentially reduced to zero after 50 minutes. There is a much more

pronounced shift to lower binding energies with increasing treatment time likely caused

by increased sample charging. The Fe and Cu 2p signals in Figure 3.13b,c show a

similar signal enhancement as the 5cm sample even after only 5 minutes of cleaning.

Interestingly, the signal after 50 minutes is significantly lower than expected for both Fe

and Cu. This is possibly associated with a surface rearrangement caused by the

prolonged exposure to UV light or it could be an artifact. The secondary electron cutoff

spectra in Figure 3.13d show a huge variation in the work function of the materials.

Such variations of over 8 eV are unlikely to be realistic even under UV irradiation and

high vacuum conditions. We suspect that severe surface charging and material

degradation occurs when the film is held close to the lamp. Based on these results, 20

minutes was selected as the optimum time for future ZONE cleaning treatments of

carbon-coated nanomaterials. The ZONE treated films were then tested for photo-

assisted CO2 reduction using the conditions described above. We did not see any

increase in the production rates of CO and CH4, which were still barely in the nmol/hr

range. Such low concentration of isotope-labelled products was approaching the

detection limit of our GC/MS instrument, and we decided that it was not worth

performing isotope labelling experiments with 13CO2 as rates that low are likely to be

due to remaining carbon contamination and not genuine CO2 reduction.

81

Figure 3.13 XPS spectra of a) C1s region b) Fe 2p region c) Cu 2p region and d)

UPS spectra of the secondary electron cut-off region of HNC films at a distance of <1cm

from the UV source in the ZONE cleaner

After attempting to remove carbon from thin films of our particles, we switched

our attention to solution-phase ligand exchange protocols which have been growing in

versatility over the last few years.9–11 The majority of these protocols aim to remove

long-chain organic surfactants to improve the electronic properties of colloidal

nanocrystals for applications in photovoltaics and solar energy conversion. This typically

involves introducing much shorter, charged capping molecules which reduce

interparticle spacing and enable increased electronic coupling of neighboring

nanocrystals. Murray et. al. recently developed a ligand exchange methodology

employing nitrosonium tetrafluoroborate (NOBF4) to replace hydrocarbon ligands and

solubilise nanocrystals in polar hydrophilic solvents.11 It was shown that NOBF4 was

effective in stabilizing a number of different materials including Fe3O4, TiO2, and FePt.11

We chose to use this protocol to attempt to exchange oleic acid and trioctylamine on the

surface of HNCs with the charged BF4- ion. Performing the ligand exchange on the

82

isolated γ-Fe2O3 particles was successful; the particles became soluble in

dimethylformamide.

Figure 3.14 FTIR spectra (left) and TEM images (right) of γ-Fe2O3 NCs before

and after ligand exchange with NOBF4

FTIR did not show any residual carbon post-exchange and TEM images of the

resulting polar-soluble particles indicated that the particles maintained their size and

shape although some aggregation into large clusters was observed, Figure 3.14. When

we performed the ligand exchange on HNCs particles, we observed aggregation of the

particles and loss of the HNC architecture, see Figure 3.11b. Solubility in DMF was poor

compared to the isolated iron oxide nanocrystals because of their much larger particle

sizes. Evidently, the Cu2O domains were not stable to NOBF4 since the separate Fe2O3

particles withstood the ligand exchange.

Despite the aggregation observed after ligand exchange process, we elected to

proceed and use the HNC particles in liquid phase dye degradation. Dye degradation

experiments are among the simplest ways to demonstrate photocatalytic activity. An

organic dye is added to an aqueous suspension of the catalyst and its UV-VIS

absorption monitored over time under illumination. The dyes are typically oxidatively

degraded by holes in the semiconductor’s valence band. We predicted that the type II

83

band alignment between γ-Fe2O3 and Cu2O will lead to reduced overlap electron and

hole wavefunctions and extend the lifetime of the carriers. The charge carriers should

then more efficiently carry out surface-based redox reactions with adsorbed reactants

compared to the separate components. To show this experimentally, we evaluated the

photocatalytic degradation of methylene blue (MB) using γ-Fe2O3/Cu2O HNCs, their

separate components, and TiO2 as a reference photocatalyst. The bare Cu2O and

Fe2O3 were also treated with NOBF4 prior to the dye degradation test. The ligand

exchange renders the nanocrystal surface accessible for dye adsorption which was

allowed to proceed for 40 minutes in the dark prior to illumination. The performance of

the various catalysts under illumination is shown in Figure 3.15a.

Figure 3.15 a) Extent of MB photocatalytic degradation over various catalysts as

determined by monitoring the main absorption peak at ~ 667nm. b) UV-Vis spectra of

the MB aqueous solution at various intervals over the γ-Fe2O3/Cu2O photocatalyst

84

The HNCs exhibit a higher rate of MB degradation than their components despite

the loss of the heterostructure. A synergistic effect between the copper and iron oxide

phases is likely occurring leading to improved charge separation. Methylene blue

degradation is thought to proceed through aromatic ring opening hydroxylation driven by

hydroxyl radicals generated upon neutralization of OH- groups by photogenerated

holes.12 This process will occur more efficiently if electron-hole recombination is

suppressed in the γ-Fe2O3/Cu2O composites. However, a control experiment with P25

TiO2 exhibited much improved activity compared to the HNCs despite its inferior visible

light absorption. The lower relative performance of HNCs can be explained by poor

dispersion of the aggregated HNCs in aqueous solution following ligand exchange. In

contrast, P25 has a hydroxylated surface which makes it much more soluble in water.

Figure 3.15b also shows the temporal UV-Vis spectra of the MB solution which exhibit a

decrease in both the aromatic and chromophoric absorbance maxima of MB indicating

that dye is being degraded and not simply decolorized.

3.5 Conclusions

In this chapter, we studied the surface properties of Fe2O3/Cu2O by XPS/UPS,

investigated ligand exchange and removal techniques, and applied the materials to gas

phase CO2 reduction and aqueous methylene blue degradation. The binding energy of

the Fe2p region of the as-synthesized particles confirmed that the iron oxide domain

was likely γ-Fe2O3 as determined by Raman spectroscopy in Chapter 2. UPS allowed

us to construct a band diagram of the components in the system and confirmed the type

II band alignment between Cu2O and Fe2O3. The HNCs were found to have an

intermediate VB/CB edges and work function due to contributions by both domains. The

particles did not exhibit a surface photovoltage response likely due to the insulating

nature of the capping ligands. We investigated UV photolysis, thermal calcination, and

solution-phase ligand exchange as different options for ligand removal and found that

the latter two, though effective, led to aggregation and loss of the HNC architecture. The

UV treatment was less effective with up to 30% carbon still present after 20 minutes.

Complete carbon removal was possible after prolonged UV treatment but the irradiation

seemed to have a questionable effect on the HNC films, inducing huge shifts in the work

85

function which were likely not realistic. Applying the HNCs to photo-assisted CO2

hydrogenation produced nmol/hr rates of CO and CH4 which were attributed to

background carbon contamination from the ligands. The UV-cleaned samples did not

show improved rates however HNCs treated with NOBF4 were found to be more active

than their separate components for aqueous phase methylene blue degradation. A

synergistic effect between the Cu2O and Fe2O3 phases was likely responsible despite

the loss of the HNC architecture.

3.6 Experimental Section

General Characterization - Powder X-ray diffraction (PXRD) was performed on

a Bruker D2-Phaser X-ray diffractometer, using Cu Kα radiation at 30 kV. Fourier

transform infrared spectroscopy (FT-IR) was performed using a Perkin Elmer Spectrum-

One FT-IR fitted with a universal attenuated total reflectance (ATR) sampling accessory

with a diamond coated zinc selenide window. UV-VIS absorbance of the samples was

measured using a Lambda 1050 UV/VIS/NIR spectrometer from Perkin Elmer. Low

resolution TEM images were acquired on a Hitachi H-7000 conventional TEM operating

at 100kV.

Photoelectron Spectroscopy - XPS/UPS spectra were acquired using a PHI

5500 instrument. An Aluminum K-alpha light source with X-ray wavelengths of 1486.7

eV under UHV conditions (< 1 x 10-9 Torr) was used for XPS spectra. Photons with

energy of 21.22 eV generated by helium plasma with a back pressure of 2 x 10-5 Torr

were used for UPS spectra. A beam of Xenon ions with kinetic energy of 3.0 eV was

used to sputter-clean the sample surface of organic ligand prior to analysis. Sputtering

was performed for an average of three minutes corresponding to a sputtering depth of ~

2nm. Samples for XPS/UPS analysis were prepared by drop-casting dilute nanocrystal

solutions on p-doped Si(100) or FTO substrates to give a film of ~50 nm thickness. The

substrates were cleaned of organics by immersion into 3:1 NH4OH:H2O2 solution at

50°C for 12 hours prior to film formation.

Gas Phase Photocatalytic Measurements - Gas-phase photocatalytic rate

measurements were conducted in a custom fabricated 1.5 mL stainless steel batch

86

reactor with a fused silica view port sealed with Viton O-rings. The reactors were

evacuated using an Alcatel dry pump prior to being purged with the reactant gases H2

(99.9995%) and CO2 (99.999%) at a flow rate of 6 mL min-1 and a stoichiometry of

either 4:1 (stoichiometric for Sabatier reaction) or 1:1 (stoichiometric for reverse water

gas shift reaction). During purging, the reactors were sealed once they had been heated

to the desired temperature. The reactor temperatures were controlled by an OMEGA

CN616 6-Zone temperature controller, with a thermocouple placed in contact with the

sample. The pressure inside the reactor during reaction was monitored during the

reaction using an Omega PX309 pressure transducer. Reactors were irradiated with a

1000 W Hortilux Blue metal halide bulb or a Newport 300 W Xe Lamp (at a distance of 4

cm and a light intensity of 2.2 suns) for a period of 16 hours. Product gases were

analyzed by a flame ionization detector (FID) and thermal conductivity detector (TCD)

installed in a SRI-8610 Gas Chromatograph (GC) with a 3’ Mole Sieve 13a and 6‘

Haysep D column. The reactor was held in a custom designed stand. Heating was

supplied from a heated copper block fixed below the fixed catalyst bed. A thermocouple

was in contact with the top of the reactor so that the reactor maintained a constant

temperature of 150 °C. Samples were prepared by drop-casting nanocrystal solutions

onto high resistivity glass slides (2 inch by 2 inch) to give a sample loading in the range

of 20-50 mg followed by drying of the films in vacuo at 60°C overnight.

Surface Photovoltage Measurements - Solutions of Cu2O NPs, pure Fe2O3

NPs, and Cu2O/Fe2O3 hetero-structured NPs, dispersed in hexane, were drop coated

one FTO substrates, and dried in air. The drop-coating was repeated 5 times to achieve

films with sufficient thickness. Trioctylamine (TOA) films were made by coating an FTO

substrate for 6 hours, followed by washing with methanol, and drying in air overnight.

UV/VIS absorption spectra were recorded for the solutions. Surface photovoltage

Spectroscopy (SPS) measurements were conducted using a vibrating gold Kelvin probe

(Delta PHI Besocke) mounted inside a home-built vacuum chamber (<10-4 mbar). Film

Samples were illuminated with monochromatic light from a 150 W Xe lamp filtered

through an Oriel Cornerstone 130 monochromator (1-10 mW·cm-2). The SPV spectra

were corrected for drift effects by subtracting dark scan background.

87

Ligand Exchange and Removal – Solution phase ligand exchange with

nitrosonium tetrafluoroborate was performed according to the published procedure by

Murray et.al.11 Briefly, 5 mL of NC dispersion in heptane (~5 mg/mL) was combined with

5 mL of dichloromethane solution of NOBF4 (0.01 M) at room temperature. The resulting

mixture was shaken gently until the precipitation of NCs was observed, typically within 5

min. After centrifugation to remove the supernatant at 7800 rpm for 5 min, the

precipitated NCs were re-dispersed in acetonitrile. To purify NCs, toluene and heptane

(1:1 by volume) were added to aggregate the NC dispersion. After centrifugation,

acetonitrile was again added to redisperse NCs to form a stable colloidal dispersion.

Ligand removal by thermal treatment was performed on drop-coated nanocrystal

films on glass substrates. A solution of NCs in heptane (~5 mg/mL) was drop cast onto

a glass slides to give a film of several hundred nanometers in thickness. The film was

calcined in an oven at 450°C in air for 24 hours and the powder was scraped off the film

with a spatula in order to do PXRD.

Ligand removal by ZONE cleaning was performed on nanocrystal films drop-cast

from heptane solutions using the ZONE SEM cleaning system from Hitachi. The films

were placed in the ZONE cleaner at distances of 50mm and <10mm from the UV light

source and ZONE cleaning was performed for a duration of 5, 10, 20, or 50 minutes.

The pressure inside the chamber was ~ 10-4 Torr. Analysis of the carbon content by

XPS was performed at the above time intervals and the sample was placed back in the

ZONE cleaner

Photocatalytic Dye Degradation - 25 mg of the various catalysts following

ligand exchange with NH4SCN were added to 100 mL of 2.5 mg/L MB aqueous solution

in a Pyrex round bottom flask. The suspensions were stirred in the dark for 40 min in

order to reach an adsorption-desorption equilibrium prior to exposure to light. A 120W

Xe lamp equipped with an AM 1.5 filter was used as the light source. The reactor was

placed approximately 5 inches from the light source resulting in an average illumination

intensity of 1.3 Suns. The degradation of MB was monitored by taking aliquots from the

reaction suspension at various intervals and following the intensity of the main

absorbance peak at ~667nm. The suspended catalyst was removed by centrifugation

88

(7800 rpm for 10 minutes) followed by filtration of the supernatant through a 0.22 μm

PTFE filter.

3.7 References

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(2) Poulin, S.; Franca, R.; Moreau-Belanger, L.; Sacher, E. Confirmation of X-Ray Photoelectron Spectroscopy Peak Attributions of Nanoparticulate Iron Oxides, Using Symmetric Peak Component Line Shapes. J. Phys. Chem. C 2010, 114, 10711–10718.

(3) Teng, X.; Black, D.; Watkins, N. J.; Gao, Y.; Yang, H. Platinum-Maghemite Core-Shell Nanoparticles Using a Sequential Synthesis. Nano Lett. 2003, 3, 261–264.


(5) Tauc, J.; Grigorovici, R.; Vancu, A. Optical Properties and Electronic Structure of Amorphous Germanium. Phys. Status Solidi 1966, 15, 627–637.


(7) Brooks, K. P.; Hu, J.; Zhu, H.; Kee, R. J. Methanation of Carbon Dioxide by Hydrogen Reduction Using the Sabatier Process in Microchannel Reactors. Chem. Eng. Sci. 2007, 62, 1161–1170.

(8) Zhao, J.; Osterloh, F. E. Photochemical Charge Separation in Nanocrystal Photocatalyst Films: Insights from Surface Photovoltage Spectroscopy. J. Phys. Chem. Lett. 2014, 5, 782–786.

(9) Rosen, E. L.; Buonsanti, R.; Llordes, A.; Sawvel, A. M.; Milliron, D. J.; Helms, B. A. Exceptionally Mild Reactive Stripping of Native Ligands from Nanocrystal Surfaces by Using Meerwein’s Salt. Angew. Chemie - Int. Ed. 2012, 51, 684–689.

(10) Nag, A.; Kovalenko, M. V.; Lee, J. S.; Liu, W.; Spokoyny, B.; Talapin, D. V. Metal-Free Inorganic Ligands for Colloidal Nanocrystals: S2-, HS-, Se2-, HSe-, Te2-, HTe-, TeS3

2-, OH-, and NH2- as Surface Ligands. J. Am. Chem. Soc. 2011, 133, 10612–10620.

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(11) Dong, A.; Ye, X.; Chen, J.; Kang, Y.; Gordon, T.; Kikkawa, J. M.; Murray, C. B. A Generalized Ligand-Exchange Strategy Enabling Sequential Surface Functionalization of Colloidal Nanocrystals. J. Am. Chem. Soc. 2011, 133, 998–1006.

(12) Houas, A. Photocatalytic Degradation Pathway of Methylene Blue in Water. Appl. Catal. B Environ. 2001, 31, 145–157.

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Chapter 4 – Investigations of Cu2O Nanocubes as Semiconducting Scaffolds for

Photocatalytic H2 Evolution and CO2 Reduction

4.1 Abstract

This chapter presents our work on the synthesis, characterization, and

gas/aqueous phase CO2 reduction and H2 evolution activity of Cu2O nanocubes. We

used a hydrophilic approach to reproducibly synthesize Cu2O nanocubes followed by

characterization of their structural and electronic properties by electron microscopy, X-

ray diffraction, and photoelectron spectroscopy. We then developed a solution-phase

coating approach to uniformly deposit TiO2 on the surface of Cu2O to give core-shell

Cu2O@TiO2 composites. Decoration of the Cu2O cubes and Cu2O/TiO2 composites with

noble metal co-catalysts such as Au, Pt, and Pd was then performed by photo-

deposition. The resulting catalysts were tested for light assisted gas-phase CO2

hydrogenation and H2 evolution from water/alcohol mixtures employing a sacrificial hole

scavenger. In the aqueous phase, it was found that bare Cu2O cubes and many of the

synthesized composites were inactive as photocatalysts for these reactions, however

some activity was observed with samples that had apparently undergone reduction of

Cu2O to Cu0 under irradiation. Therefore we then studied the performance of

commercial Cu0 nanopowder dispersed on P25-TiO2 as a model system and found the

optimal Cu loading for H2 evolution. We concluded the aqueous phase study by

examining the effects of the alcoholic hole scavengers (MeOH, EtOH) and metal co-

catalysts (Au, Pt, Cu) on product selectivity using TiO2 as a proven semiconductor

scaffold. We found that light-assisted reforming of the hole scavenger to CO, CH4 and

higher hydrocarbons is responsible for product formation along with H2O reduction. In

the gas phase, we detected possible evidence for light enhancement of product

formation over the Cu2O-based catalysts but the rates of CO and CH4 were too low to

be confirmed with isotope labelling experiments.

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4.2 Introduction

In the previous chapter, we described the challenges caused by the presence of

organic capping ligands on the surfaces of photocatalytic nanomaterials. These

insulating long-chain hydrocarbons hinder electronic coupling between neighboring

nanocrystals, prevent reactant molecules’ access to the nanoparticle surface, and can

decompose under testing conditions to give false positive signals in hydrocarbon

detection. Unfortunately, high temperature solution phase synthetic methods invariably

make use of these molecules to prevent nanocrystal aggregation and provide solubility

in organic solvents. Their complete removal by ligand exchange or calcination can be

challenging as it typically results in undesirable chemical and structural changes to the

nanoparticles’ inorganic cores. As a result, we decided to focus on aqueous synthetic

procedures which rely on charge stabilization of colloids for solution stability. Typically,

hydrophilic synthetic methods have resulted in comparatively poor control of size,

shape, and crystallinity relative to high temperature non-aqueous approaches. Recently,

a number of reproducible aqueous synthetic methods have been reviewed yielding

various nanocrystalline metals and metal oxides.1–7 In this chapter, we will continue to

focus on Cu2O which is one of the most promising cathode materials for solar fuel

applications. Its narrow bandgap, favourable energetics, elemental abundance, and low

cost are some of the reasons why there has been a growing interest in developing

systems based on this material. Numerous papers reporting reproducible syntheses of

Cu2O nanocubes, octahedra, nanowires, and hollow particles have recently been

published.8–13 In terms of shape, nanocubes were chosen for our study because of their

ease of synthesis and ability to be converted into other shapes if needed.9,14 Studies

demonstrating Cu2O’s potential for water splitting, dye-degradation, gas sensing, and as

electrode material in lithium ion batteries have been published.15–19 Most of the

challenges associated with Cu2O revolve around its oxidative and reductive instability in

aqueous solution and poor charge carrier diffusion length.20 The general strategy to deal

with these issues is deposition of a metal co-catalyst that acts an electron sink to

simultaneously promote charge separation and alleviate self-photocorrosion of Cu2O.21–

24 The rest of this chapter will describe our work in using Cu2O nanocubes as a template

for the development of multi-component solar fuel photocatalysts.

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Aqueous Cu2O nanocubes were synthesized from Cu(SO)4·5H2O by a modified

literature procedure.25 Initially, Cu(SO)4·5H2O was converted to the corresponding

hydroxide, Cu(OH)2, by treatment with NaOH. Sodium citrate was added to the

precursor solution to act as a chelating agent for Cu2+ thereby reducing the

concentration of Cu(OH)2 with increasing amount of citrate added. Addition of ascorbic

acid as the reducing agent gave Cu2O nanocubes over the course of a 30 minute

reaction at room temperature. By controlling the amount of citrate added, the size of the

corresponding cubes could be varied in the range of 50-100 nm, see Figure 4.2.

Formation of copper citrate reduces the concentration of Cu(OH)2 crystallites which act

as seeds for the nucleation of Cu2O following reduction with ascorbic acid. Fewer seeds

go on to form larger particles whereas higher Cu(OH)2 concentrations lead to a larger

number of seeds and correspondingly smaller cubes, Figure 4.1.

Figure 4.1 Scheme of Cu2O nanocube formation using sodium citrate as

chelating agent to control particle size (Reprinted with permission from Ref (25)


An average particle size of 82 nm was observed when the molar ratio of citrate to

copper sulfate was maintained at 0.75:1. When the ratio is reduced to 0.25:1, the

increased number of Cu(OH)2 seeds reduces the average particle size to 58 nm.

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Nanocubes were the only morphology observed for the particles by TEM. The

monodispersity of the synthesis is quite high, with a standard deviation of only 5% for

the regular prep and 9% for the reduced citrate synthesis. In the absence of citrate,

small 30-50 nm Cu2O cubes agglomerate into larger structures of over 100 nm in

diameter. By imaging a film of nanocubes on a glassy carbon electrode, we confirmed

that there is significant porosity in the nanocrystal film which we hypothesized would

allow reactant molecules in the aqueous and gas phases to infiltrate and reach the

majority of particle surfaces.

Figure 4.2 a,b) Dark and bright field TEM images of as synthesized Cu2O

nanocubes with an average size of 82 nm (scale bar 100nm) c,d) Dark and bright field

SEM images of Cu2O nanocubes on a TEM grid and glassy carbon electrode

respectively

The powder X-ray diffraction pattern of the nanocubes in Figure 4.3b indicated

that the sample was pure cuprite (Cu2O) without any impurity phases. Using XPS we

found a small amount of Cu2+ present on the surface of the cubes indicating residual

copper sulfate precursor or a thin shell of CuO. Its absence from the PXRD spectrum

94

indicated that it is likely a very thin or amorphous Cu2+ layer. The FTIR spectrum

suggested that the particle surface is generally free of carbon species. The citrate and

ascorbate species present on the surface are not covalently bound and are likely

washed away during purification of the nanocubes. A weak OH peak in the FTIR

indicates surface hydroxides or adsorbed moisture, and there is no evidence of C-H

stretching modes.

Figure 4.3 a) Particle size distribution of Cu2O synthesized using 0.75 and 0.25

equivalents of citrate b) PXRD pattern c) XPS scan of Cu2p region and d) FTIR

spectrum of as-synthesized Cu2O nanocubes

Once we had confirmed the synthesis of the cubes by TEM and PXRD, we

turned to studying their thermal oxidation stability which is a concern for Cu2O as it

contains copper in the unstable +1 oxidation state. We first heated the powdered

nanocubes for 1 hour at temperatures ranging from 100°C – 500°C in air and looked at

the resulting PXRD patterns. The cubes were stable at 100°C and 200°C under these

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conditions as seen in the top of Figure 4.4. After 1 hour at 300°C, the Cu2O cubes

began to partially convert to CuO, and after treatment at 500°C the oxidation was

essentially complete. Based on these results, we selected 200°C as the temperature to

conduct a prolonged heating of the cubes, shown in the bottom of Figure 4.4. At 200°C,

no oxidation was observed after 4 hours of heating in air, which was crucial for

depositing a thin shell of TiO2 and its crystallization, as explained below.

Figure 4.4 top) PXRD patterns of Cu2O nanocubes heated at various

temperatures for 1 hour in air bottom) PXRD patterns of Cu2O cubes at increasing

times under 200°C in air

96

With a view on further improving the stability of our Cu2O nanocubes, we

developed a solution phase approach to coating them with a thin layer of TiO2. Cu2O is

prone to photo-degradation since the redox potentials of Cu1+ reduction to Cu0 and Cu1+

oxidation to Cu2+ in aqueous solution are both found within the energy range of the

Cu2O bandgap. In order to prevent the Cu2O surface from coming in contact with the

surrounding electrolyte, coating it with a protective layer of another metal oxide has

been shown to be a useful strategy. Gratzel and co-workers employed this method to

stabilize thin films of Cu2O with a thin layer of aluminum zinc oxide deposited by ALD.15

Here we used a slow hydrolysis of Ti(OBu)4 in alcoholic solution to lay down a

conformal amorphous TiO2 coating on the pre-formed Cu2O cubes. Figure 4.5 shows

TEM images of the resulting core-shell particles confirming complete surface coverage.

We also performed an EDX analysis which confirmed the presence of a thin TiO2 shell

on the surface of the cubes. Based on the TEM images and EDX signal intensities, we

estimated that the thickness of the TiO2 shell was in the range of 20-40 nm.

Figure 4.5 a-c) TEM images of Cu2O@TiO2 nanocubes (scale bar 500 nm) d-f)

Ti, Cu, and O elemental signals from EDX linescan in Panel B

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Figure 4.6 top) PXRD pattern of as synthesized Cu2O@TiO2 nanocubes

showing absence of TiO2 reflections bottom) PXRD pattern of Cu2O@TiO2 nanocubes

following heat treatment at 200°C for 3 days

Interestingly, when we looked at the PXRD patterns of the resulting core-shell

particles we did not see any reflections corresponding to TiO2, see Figure 4.6 (top). The

thickness of the TiO2 layer as determined by TEM should be enough to give a sufficient

X-ray diffraction signal therefore its absence suggested that the TiO2 layer was

amorphous. An amorphous TiO2 layer might be effective in preventing Cu2O

decomposition but it is likely to adversely affect charge transport in the particles and

subsequently their photocatalytic performance. Therefore we attempted to crystallize

the TiO2 by thermal treatment in air at 200°C, a temperature at which the underlying

Cu2O cubes were found to be stable. We started out by heating for several hours but

did not observe any noticeable change in the PXRD patterns and so we continued the

98

heat treatment up to a period of 3 days. The results are shown in Figure 4.6. We still did

not observe any TiO2 reflections, but a partial oxidation of Cu2O to CuO had taken

place. It was apparent that 200°C was not a sufficiently high temperature for TiO2

crystallization. We then increased the temperature to 400°C and performed the heat

treatment under Ar atmosphere to try and prevent oxidation of the nanocubes to CuO.

Under these conditions, we found that a heating time of 30 minutes was sufficient to

crystallize the TiO2 coating to a mixture of anatase and rutile polymorphs, Figure 4.7.

Figure 4.7 a) PXRD pattern of Cu2O@TiO2 composites after calcination at 400°C

for 30 minutes under Ar b,c) TEM images of the composites before and after heat

treatment (scale bars 100 and 300 nm respectively)

PXRD confirmed that the cuprite crystal structure had been largely maintained

even under elevated temperature due to the inert atmosphere. TEM images of the

Cu2O@TiO2 composites taken before and after calcination indicate that the treatment

99

induced partial agglomeration of the particles into larger clusters of an average size in

the range of 200-400 nm. The increased contrast of the image following calcination is

caused by the crystalline TiO2 coating. With the TiO2 layer being no longer transparent

under the electron beam, we were not able to confirm that the underlying cube

morphology was preserved. However, Cu2O was still present by PXRD.

Once we had a developed an effective method to coat the Cu2O nanocubes in a

shell of crystalline TiO2 without changing its crystallinity, we turned our attention to the

deposition of noble metal co-catalysts to act as active sites on the semiconductor

cubes. Plenty of publications exist on the topic of decorating Cu2O nanomaterials with

metals such as Au, Pt, and Pd.21,26–28 The most straight forward approach is simple

galvanic replacement of Cu+ with a cation that is more easily reduced on the

electrochemical series, although this often results in poor size control of the resulting

metal particles. Here we decided to employ photo-deposition by decomposing labile

metal precursors under UV irradiation. The Cu2O nanocubes were first illuminated in

order to excite electrons to the conduction band and improve the efficacy of metal cation

reduction. A solution of the metal ions in ethanol was then injected and irradiated for a

further period of time to complete the deposition. We varied the mass loadings of Pt, Au,

and Pd in the range of 2-10%. The small size and relatively low loadings of the

deposited metals made them difficult to detect by PXRD however the particles were

clearly visible by TEM as seen in Figure 4.8. We were able to successfully deposit small

(2-5 nm) particles of the above metals on bare Cu2O cubes and Cu2O@TiO2

composites. In the case of Cu2O/Metal particles, our motivation was to utilise the metal

domains as reduction centers where photogenerated electrons would funnel and reduce

CO2 or H2O to hydrocarbons or hydrogen. To protect Cu2O from oxidation, we

employed sacrificial reducing agents in the aqueous phase during testing to soak up

photogenerated holes. In the case of Cu2O@TiO2, the titania shell should act as a

barrier to photocorrosion of the Cu2O core, while being thin enough for photogenerated

electrons to tunnel through and reach the Pt sites on the surface. We chose to do an in-

depth study of the Cu2O/Pt system as Pt has been shown to be a very promising

reduction co-catalyst when used in conjunction with other semiconductors such as TiO2,

CdSe and Nb3O7OH.29–32

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Figure 4.8 TEM images of Cu2O nanocubes decorated with various noble

metals a) Au (3%) b) Pt (10%) c) Coated with TiO2 and d) Pd (10%)

We started by testing the materials discussed so far for aqueous phase CO2

reduction and/or H2 evolution from water employing a sacrificial hole scavenger. We

utilised a 300 W Xe lamp and an AM 1.5 filter to simulate solar radiation and selected

MeOH as the hole scavenger. We elected to do preliminary testing on H2 evolution from

water instead of aqueous CO2 reduction because the reaction is kinetically simpler but

imposes many of the same light absorption, stability, and charge transport

requirements. It can therefore be used as a pre-screening tool on potential

photocatalysts. A photograph of the reactor setup is shown in Figure 4.9a and the

101

results are summarized in Table 4.1. Of all the samples tested, only the Cu2O@TiO2

and Cu2O-Pt@TiO

2 calcined samples produced any H2 from H2O/MeOH under AM 1.5

illumination. To clarify, these two samples represent Cu2O nanocubes coated with

crystalline TiO2 and Cu2O cubes decorated with Pt nanoparticles and subsequently

crystalline TiO2 respectively. Some general observations can be made from the data in

Table 4.1. The bare Cu2O cubes on their own are inactive likely due to charge

recombination within the particles or a lack of suitable surface activating sites for H2

evolution. The cubes were stable in aqueous solution under AM 1.5 which could be an

indication that carrier generation and charge transport within the nanocube volume was

the main issue. The stability was somewhat unexpected in view of Cu2O`s well-known

problems regarding photocorrosion although it could be a further indication that

electron-hole recombination in the Cu2O cubes was ultra fast and the charges did not

reach the surface.

Table 4.1 Rates of H2 Evolution from Water with Cu2O Nanocubes and Related

Materials as the Photocatalyst

Sample Rate of H2 Production

Bare Cu2O N/A

Cu2O@TiO

2 amorphous N/A

Cu2O@TiO

2 calcined 70 μmol g

-1 h

-1

Cu2O/Pt N/A

Cu2O/Au N/A

Cu2O/Pd N/A

Cu2O@TiO

2-Pt amorphous N/A

Cu2O@TiO

2-Pt calcined N/A

Cu2O-Pt@TiO

2 amorphous N/A

Cu2O-Pt@TiO

2 calcined 19 μmol g

-1 h

-1

Cu2O/WS

2 N/A

None of the Cu2O/Metal composite samples showed any activity either despite

the presence of the noble metal reducing sites. Again, this points to charge

recombination or insufficient charge generation by the cubes. We also looked at WS2

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nanosheets developed in our lab as a co-catalyst but did not observe any activity.33 The

samples coated with TiO2 were the ones to show observable rates of H2 evolution. We

looked at Cu2O@TiO

2 with an amorphous or crystalline TiO2 coating, and observed an

average H2 production rate of 70 μmol /g h only with the calcined sample. This held true

across all tested samples; none of the amorphous samples showed any activity

confirming the importance of crystallinity to charge transport. For Cu2O@TiO

2, an

induction period was observed as seen in Figure 4.9b, where no H2 was produced in the

first few hours of the reaction, but a linear production rate was observed overnight.

Figure 4.9 a,c) Photographs of the aqueous phase photocatalytic testing setup

and reactor b,d) H2 evolution as a function of time from Cu2O@TiO2 nanocubes

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We studied the Cu2O@TiO2 sample by PXRD before and after testing in Figure

4.10. The presence of Pt could not be confirmed by PXRD as expected, however

reflections attributed to the Cu2O core and TiO2 shell were clearly present. The

presence of metallic Cu after testing was also clearly observed indicating reduction of

Cu+1 to Cu0. It is possible that electrons generated on Cu2O migrated to the TiO2/Cu2O

interface and were unable to tunnel through the TiO2 layer. The accumulated electrons

then likely reduced Cu+1 to the metallic Cu seen in the PXRD pattern.

Figure 4.10 top) PXRD pattern of Cu2O-Pt@TiO2 particles before testing

bottom) PXRD pattern of Cu2O-Pt@TiO2 particles after testing; (inset) – photograph of

the catalyst suspension showing presence of metallic Cu

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We also tested a series of 4 similar samples consisting of combinations of Cu2O,

Pt, and TiO2. The samples labelled Cu2O@TiO2-Pt were made by photo-depositing

platinum particles on Cu2O@TiO2 composites such that the Pt sites were depositing on

TiO2. The remaining two samples, labelled Cu2O-Pt@TiO2, were synthesized by photo-

depositing 3% Pt directly on the Cu2O cubes and then coating the entire assembly in a

thin layer of TiO2. The calcined Cu2O-Pt@TiO

2 sample was the second to show any

activity at an average rate of 19 μmol H2 /g h. The fact that both active samples

exhibited a TiO2 surface could be a sign that its surface chemistry is required to reduce

protons and could explain the lack of activity for Cu2O. No reductive decomposition was

seen in the Cu2O-Pt@TiO2 sample where Pt particles were present between the Cu2O

and TiO2 layers. However, we did not observe any activity for this sample. The active

Cu2O@TiO2 and Cu2O-Pt(3%)@TiO2 samples did not produce any hydrocarbon

products when tested in 0.1 M NaHCO3 as a source of CO2. We then suspected that the

nanocube morphology could be another reason why we were not observing products

from the majority of our catalysts. The cubes are typically terminated with {100} facets

which may have been inactive and so we purchased two commercial forms of Cu2O: a

bulk powder in the micron range and a sample labelled Cu2O “nanospheres” with

particle sizes in the 50-100nm range. Both of these samples exhibited irregular particle

morphologies that expose additional facets not seen in the nanocubes. Figure 4.11

shows the characterization we performed on these commercial samples prior to testing.

Both samples were nearly pure crystalline Cu2O however, the “nanospheres” were quite

poorly defined and did not show any identifiable nanoparticles, but rather a network of

interconnected copper oxide species. Regardless, when tested neither sample

produced H2 under the same testing conditions as above.

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Figure 4.11 top) TEM image and PXRD pattern of commercial bulk Cu2O

powder used as a testing reference bottom) TEM image and PXRD pattern of

commercial Cu2O “nanospheres

Following the aqueous phase tests where we observed possible activity from two

of the samples, we decided to proceed with gas-phase testing. In the gas phase, the

absence of an electrolyte to conduct H+ could eliminate many of the photocorrosion

concerns around Cu2O-based materials. Light assisted CO2 hydrogenation was

performed under the standard conditions described in Chapter 3 and outlined in the

experimental section. We tested bare Cu2O nanocubes as a reference along with Cu2O-

Pt and the active Cu2O@TiO2 and Cu2O-Pt(3%)@TiO2 catalysts. Samples were

prepared by drop-casting nanocrystal ethanolic solutions onto high resistivity quartz

substrates and drying in vacuo. The results are summarized in Figure 4.12. We

detected the presence of CO and CH4 at rates of approximately 1 nmol/hr, similar to the

rates we previously observed from Cu2O/Fe2O3 HNCs. The experiments were repeated

two or three times and showed good reproducibility from batch to batch. We first

compared the bare Cu2O cubes with the core-shell Cu2O@TiO2 particles.

106

Figure 4.12 a) Rates of CO and CH4 over multiple runs from bare Cu2O

nanocubes and Cu2O@TiO2 composites b) CH4 production rates from Cu2O-

Pt(3%)@TiO2 composite (no CO detected) c) CO and CH4 evolution rates on the Cu2O-

Pt(3%) sample over multiple runs

107

The latter consistently produced slightly higher hydrocarbon evolution rates

including an amplifying effect under irradiation in the case of CO. The Cu2O-

Pt(3%)@TiO2 sample only produced methane and no CO as seen in Figure 4.12b with

a twofold rate enhancement under light. Finally, the Cu2O-Pt particles gave both CO

and CH4 at similar nmol/hr rates; once again CO rates were higher under illumination

whereas CH4 production was the same under dark and light conditions. The absence of

a light effect in the case of CH4 suggests that it was being produced from carbon

contamination sources and not CO2 reduction. The light enhancement observed for CO

could be promising however the overall rates were still too low to be confirmed with

13CO2 isotope labelling by GC/MS. Based on this data, we concluded that the Cu2O

nanocube catalysts were not active under these conditions for the gas phase

hydrogenation of CO2. After observing H2 production activity for the Cu2O@TiO2

sample, we hypothesized that it may have been caused by the reduction to Cu0 seen in

Figure 4.10. A Cu/TiO2 system formed in-situ by reduction of Cu2O@TiO2 could be an

effective water reduction catalyst as it contains a metal-semiconductor junction to

separate charges, and a catalytic metal site. We synthesized a model Cu/TiO2 catalyst

by simple physical grinding of commercial P25 and nanoparticulate Cu powders. We

varied the % Cu loading in the range of 0.1% to 50% and also looked at pure Cu

nanopowder and pure P25 particles as controls. Using the aqueous phase testing

setup, we obtained rates of H2 production that were up to 2 orders of magnitude higher

than those obtained using the catalytic systems based on Cu2O nanocubes, Table 4.2.

Table 4.2 Rates of H2 evolution from water using P25/Cu (10%) mixture and its

separate components

Sample Rate of H2 Production

(mmol g-1 h-1) Cu Nanopowder N/A

P25 TiO2 0.23 ± 0.02 P25/Cu (0.1%) 0.30 ± 0.9 P25/Cu (1%) 4.0 ± 0.8

P25/Cu (10%) 3.9 ± 0.9 P25/Cu (50%) 1.6 ± 0.9

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Bare TiO2 produced H2 at a rate of ~230 μmol g-1 h-1 compared to the highest

rate of 70 μmol g-1 h-1 obtained using the Cu2O@TiO2 system. The fact that TiO2 which

only absorbs approximately 4% of the solar spectrum performed better than the Cu2O-

based systems indicates that charge generation and/or surface chemistry was inefficient

in the Cu2O nanocubes. When various amounts of Cu were added, a significant

increase in H2 evolution was observed. Maximum rates on the order of 4.0 mmol g-1 h-1

were obtained for the 1% and 10% Cu loading catalysts, Figure 4.13. Amounts of CH4

and CO in the μmol range were also found which are likely a product of side reactions of

the hole scavenger since no additional carbon source such as NaHCO3 had been

added except the alcoholic scavengers. The effect of scavengers on product distribution

will be discussed below. PXRD patterns taken before and after testing did not show any

structural changes caused by illumination. We also performed a control experiment

where we prepared a physical mixture of Cu nanopowder and Cu2O nanocubes in the

same fashion as the Cu/TiO2 mixture. Once again, this system was inactive despite the

presence of the metal, indicating that the Cu2O cubes were not effective in photo-

generating charge carriers.

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Figure 4.13 a,c) H2 evolution rates of a function of Cu loading in Cu/TiO2

catalysts and H2 evolution profile with time over the P25/Cu (10%) catalyst b,d) PXRD

patterns of P25/Cu (10%) samples before and after testing

We studied the effect of different hole scavengers (MeOH, EtOH etc.) and

different metallic co-catalysts (Au, Pt) on the product distribution over P25 samples. Our

findings are summarized in Figure 4.14. All catalysts were tested under AM 1.5

illumination with 25 % v/v scavenger in deionized water. Control reactions in the dark

and in the absence of a catalyst did not result in any detectable products above the

background. Pure P25 produced mainly hydrogen with small amounts of CO and trace

CH4. This result was in good agreement with the data in Table 4.2 on pure P25. We

then looked at the effect of the alcohol scavenger on the products over a P25/Pt (1%)

catalyst which was prepared in the same way as the P25/Cu samples earlier.

110

Figure 4.14 Hydrocarbon and H2 evolution rates over various model catalysts

using different hole scavengers

In the presence of MeOH, H2 is the predominant product with a production rate of

almost 15 mmol g-1 h-1. Recall that using the P25/Cu (1%) sample, we obtained an

average rate of 4 mmol g-1 h-1. The dramatic increase in H2 production compared to the

bare P25 sample was noted, confirming the beneficial effects of metal semiconductor

junctions on photocatalysis. CH4 is seen as a side product at a rate of ~ 12 μmol g-1 h-1.

If EtOH is used as the scavenger, hydrogen production is relatively unchanged but the

amount of CH4 produced increases drastically to ~260 μmol g-1 h-1. We suspect that this

is due to decarboxylation of acetic acid under illumination; a pathway termed the photo-

Kolbe reaction.34,35 EtOH likely undergoes consecutive 2e- oxidations to acetaldehyde

111

and then acetic acid. Acetic acid then reacts with a further oxidizing equivalent to give

CH4 and CO2 according to Equations 4.1 and 4.2. Ethane was also detected as a

relatively minor product as per Equation 4.3.

(4.1) CH3COOH + h+ → •CH3 +CO2 + H+

(4.2) •CH3 + CH3COOH → •CH2COOH + CH4

(4.3) 2 CH3COO- → CH3CH3 + 2CO2

We then looked at the effect of changing the co-catalyst on product distribution.

We prepared P25/Au (1%) samples by physical grinding of Au and TiO2 nanopowders

as was done for Pt and Cu. We still obtained H2 as a major product albeit at significantly

lower amounts compared to Pt. Interestingly, the most abundant product when using

MeOH as the scavenger was CO which was not detected for P25/Pt. The mixture of CO

and H2 produced by the P25/Au system is referred to as syngas. In the presence of

EtOH, CH4 is produced at the expense of CO likely similar to the pathway using P25/Pt.

We also looked at two inorganic scavengers, sodium iodide (NaI) and sodium

thiosulfate (Na2S2O3) to rule out the presence of carbon. We did not see any detectable

amounts of hydrocarbons or H2 above the background levels of these products. This

data indicates that the identity of the organic scavenger is very important in determining

the products of aqueous “water splitting” experiments. Furthermore, the protons in the

scavenger undergo rapid exchange with protons from H2O so that the products detected

originate from both H2O and the scavenger. Therefore it is hard to decouple genuine

water splitting from MeOH/EtOH reforming and many published results in the field

should be viewed with caution if water splitting is claimed as the sole method of H2

evolution.

4.4 Conclusions

In this chapter we investigated Cu2O nanocubes as a semiconducting scaffold for

the synthesis of multiple component photocatalytic architectures. We first described the

synthesis and provided characterization of the nanocubes using TEM, powder X-ray

diffraction and XPS. The nanocubes were reproducibly synthesized with sizes on the

range of 50-100 nm and were found to contain much lower levels of carbon species on

the surface as compared to the Fe2O3/Cu2O particles in Chapters 3 and 4. We then

112

developed a methodology for coating the nanocubes with a conformal layer of

amorphous TiO2 and crystallizing it via heat treatment under inert conditions. We then

deposited noble metals (Au, Pt, Pd) by photo-decomposition of labile precursors onto

the bare Cu2O cubes and the core-shell Cu2O@TiO2 and tested these materials for H2

evolution from H2O/MeOH solutions under illumination. We found the majority of the

catalysts were inactive with the exception of Cu2O@TiO2 and Cu2O/Pt@TiO2 which

gave modest rates in the μmol g-1 h-1 range. Both of these samples exhibited a

crystalline TiO2 surface and we observed metallic Cu by PXRD in the case of

Cu2O@TiO2 which led us to believe that we were in fact forming Cu/TiO2 hetero-

structures. We then prepared this model system by simple mixing and grinding of its two

components and found that it exhibited H2 production rates that were 2 orders of

magnitude higher than the samples derived from Cu2O nanocubes. Cu loadings of 0.1,

1, 10 and 50% w/w were tested and the 1% w/w found to give the highest rates. Finally,

we studied the effect of the hole scavenger and co-catalyst on a working H2 evolution

scaffold, namely TiO2. We found vastly different amounts of H2, CO, CH4, and higher

hydrocarbons were produced depending on the nature of the scavenger and the metal.

These products likely arise from simultaneous H2O reduction and MeOH/EtOH

reforming and so it is difficult to claim these processes as pure water splitting. We also

tested the active samples for gas phase light-assisted CO2 hydrogenation to CO and

CH4. In the case of CO, we saw some evidence of a catalytic effect in the presence of

light, however the rates were still very low (nmol h-1) and we were not able to confirm

that the hydrocarbons originated from CO2 reduction or from extraneous carbon

contamination.


General Characterization - All chemicals were purchased from Sigma Aldrich

and used directly without further purification. Powder X-ray diffraction (PXRD) was

performed on a Bruker D2-Phaser X-ray diffractometer, using Cu Kα radiation at 30 kV.

Fourier transform infrared spectroscopy (FT-IR) was performed using a Perkin Elmer

Spectrum-One FT-IR fitted with a universal attenuated total reflectance (ATR) sampling

accessory with a diamond coated zinc selenide window. UV-VIS absorbance of the

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samples was measured using a Lambda 1050 UV/VIS/NIR spectrometer from Perkin

Elmer.

Cu2O Nanocubes – Cu2O nanocubes were synthesized according to a modified

literature procedure.25 The Cu2O nanocubes were synthesized from copper sulfate

(CuSO4, 99%), sodium hydroxide (NaOH, 99%), trisodium citrate dihydrate

(C6H5Na3O7．2H2O, 99%), and ascorbic acid (C6H8O6, 99.9%). In a typical procedure, a

round-bottom flask was filled with 300 mL of deionized water and 0.265 g (0.9 mmol) of

trisodium citrate was added with vigorous stirring until dissolved. Then, 1 mL of 1.2 M

CuSO4 solution (1.2 mmol) was rapidly injected using a pipette. After 5 min, 1 mL of 4.8

M NaOH (4.8 mmol) solution was injected into the solution. The clear blue solution

immediately turned turbid blue, indicating the precipitation of Cu(OH)2. After another 5

min, 1 mL of 1.2 M (1.2 mmol) ascorbic acid was injected as reducing agent. The color

of the solution rapidly changed from blue to green to light orange. The solution was

stirred at room temperature for 30 minutes accompanied by a color change to bright

orange at the end. The reaction mixture was centrifuged at 4000 rpm for 2 minutes to

separate the cubes which were then dispersed in 30 mL of deionised water. The

process was repeated 3 more times with the particles dispersed in absolute EtOH after

the final washing. The nanocubes can be stored in solution or as a powder if the EtOH

is evaporated in a drying oven at 60°C.

Cu2O@TiO2 nanocubes – 30 mg of powdered Cu2O nanocubes were dispersed

in 20mL anhydrous EtOH in a 50ml round bottom flask with magnetic stirring. 2.5 mL of

a 0.1 M Ti(OBu)4 solution in EtOH were added dropwise and the mixture was stirred for

2 hours in air at room temperature. 1.2 mL of 1:3 H2O:EtOH mix was added dropwise

and stirred for another 2 hours. The Cu2O@TiO2 particles were centrifuged at 4000 rpm

for 2 minutes and redispersed in 20 mL absolute EtOH.

Deposition of Pt, Au, Pd – Metal nanoparticles of Pt, Au, Pd were deposited

onto Cu2O nanocubes or Cu2O@TiO2 composites by photo-decomposition of H2PtCl6,

HAuCl4, or Pd(acetate)2 under UV irradiation. The loadings were calculated on a mass

percent basis and were in the range of 1-10% metal. Powdered Cu2O or Cu2O@TiO2

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nanocubes were dispersed in ~20 mL absolute EtOH with magnetic stirring in a 100 mL

round bottom flask. The flask’s exterior was wrapped in aluminum foil and a UV finger

light source was used to illuminate the dispersions for one hour in order to activate the

surface and create energetic electrons for subsequent reduction of metal salts.

Appropriate amounts of the noble metal precursors were introduced with a 1 mL syringe

from ethanolic solutions and the UV irradiation was continued for another hour. The

particles were then centrifuged at 4000rpm and redispersed in water or absolute

ethanol.


H-7000 conventional TEM operating at 100kV. SEM Images and EDX spectra were

acquired on a Hitachi S-5200 operating at 30kV using an Oxford Inca detector.

Additional SEM images were acquired on a FEI Quanta FEG 250 Environmental SEM.

Sample preparation involved dropping a dilute nanocrystal solution on a carbon coated

Ni TEM grid.
















115












Haysep D column. The reactor was held in a custom designed stand. Heating was

supplied from a heated copper block fixed below the fixed catalyst bed. A thermocouple

was in contact with the top of the reactor so that the reactor maintained a constant

temperature of 80 °C. Samples were prepared by drop-casting nanocrystal solutions

onto high resistivity glass slides (2 inch by 2 inch) to give a sample loading in the range

of 20-50 mg followed by drying of the films in vacuo at 60°C overnight.

Aqueous Photocatalytic Testing - All photocatalytic tests were performed in a

home built, gas-tight pyrex reactor with a quartz window. The photocatalyst powder (1

gL-1) was dispersed in distilled water and methanol (25 vol%, HPLC grade). Prior to

each run, the reactor was purged with nitrogen for 20 min and the first sample was

taken before the photocatalytic reaction was started. A Newport 120 W or 300 W Xe

lamp was used for solar irradiation and an air mass (AM) 1.5 filter (Newport) was

applied to simulate the direct solar spectrum when the sun is at a zenith angle of 48.28.

To test the long-term stability of the photocatalyst and the reproducibility of the H2

formation, the reactor was purged after a few hours to remove all products and start a

new formation cycle. Gas samples were extracted with a gas-tight syringe, separated by

gas chromatography (GC, Agilent 7820A GC) equipped with a thermal conductivity

detector (TCD). A 1.5 m Molesieve 13 X column (80–100 mesh) was used for the

separation of H2, O2, and N2, and an integrated 1 m Haysep Q column (80–100 mesh)

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was used to first separate CO2 and water vapour to avoid contamination of the

Molesieve column.

4.6 References

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(5) Wang, X.; Zhuang, J.; Peng, Q.; Li, Y. A General Strategy for Nanocrystal Synthesis. Nature, 2005, 437, 121–124.

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(7) Lim, B.; Wang, J.; Camargo, P. H. C.; Jiang, M.; Kim, M. J.; Xia, Y. Facile Synthesis of Bimetallic Nanoplates Consisting of Pd Cores and Pt Shells through Seeded Epitaxial Growth. Nano Lett. 2008, 8, 2535–2540.

(8) Singh, D. P.; Neti, N. R.; Sinha, A. S. K.; Srivastava, O. N. Growth of Different Nanostructures of Cu2O (Nanothreads, Nanowires, and Nanocubes) by Simple Electrolysis Based Oxidation of Copper. J. Phys. Chem. C 2007, 111, 1638–1645.

(9) Siegfried, M. J.; Choi, K.-S. Directing the Architecture of Cuprous Oxide Crystals during Electrochemical Growth. Angew. Chemie, Int. Ed. 2005, 44, 3218–3223.

(10) Siegfried, M. J.; Choi, K.-S. Electrochemical Crystallization of Cuprous Oxide with Systematic Shape Evolution. Adv. Mater. (Weinheim, Ger. 2004, 16, 1743–1746.

(11) Li, H.; Liu, R.; Zhao, R.; Zheng, Y.; Chen, W.; Xu, Z. Morphology Control of Electrodeposited Cu2O Crystals in Aqueous Solutions Using Room Temperature Hydrophilic Ionic Liquids. Cryst. Growth Des. 2006, 6, 2795–2798.

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(12) Wang, W.; Wang, G.; Wang, X.; Zhan, Y.; Liu, Y.; Zheng, C. Synthesis and Characterization of Cu2O Nanowires by a Novel Reduction Route. Adv. Mater. (Weinheim, Ger.) 2002, 14, 67–69.

(13) Chang, Y.; Zeng, H. C. Manipulative Synthesis of Multipod Frameworks for Self-Organization and Self-Amplification of Cu2O Microcrystals. Cryst. Growth Des. 2004, 4, 273–278.

(14) Siegfried, M. J.; Choi, K.-S. Elucidating the Effect of Additives on the Growth and Stability of Cu2O Surfaces via Shape Transformation of Pre-Grown Crystals. J. Am. Chem. Soc. 2006, 128, 10356–10357.

(15) Morales-Guio, C. G.; Tilley, S. D.; Vrubel, H.; Grätzel, M.; Hu, X. Hydrogen Evolution from a Copper(I) Oxide Photocathode Coated with an Amorphous Molybdenum Sulphide Catalyst. Nat. Commun. 2014, 5, 3059.

(16) Hara, M.; Kondo, T.; Komoda, M.; Ikeda, S.; Shinohara, K.; Tanaka, A.; Kondo, J. N.; Domen, K. Cu2O as a Photocatalyst for Overall Water Splitting under Visible Light Irradiation. Chem. Commun. 1998, 357–358.

(17) Zhang, J.; Liu, J.; Peng, Q.; Wang, X.; Li, Y. Nearly Monodisperse Cu2O and CuO Nanospheres: Preparation and Applications for Sensitive Gas Sensors. Chem. Mater. 2006, 18, 867–871.

(18) Xu, H.; Wang, W.; Zhu, W. Shape Evolution and Size-Controllable Synthesis of Cu2O Octahedra and Their Morphology-Dependent Photocatalytic Properties. J. Phys. Chem. B 2006, 110, 13829–13834.

(19) Poizot, P.; Laruelle, S.; Grugeon, S.; Dupont, L.; Tarascon, J. M. Nano-Sized Transition-Metal Oxides as Negative-Electrode Materials for Lithium-Ion Batteries. Nature 2000, 407, 496–499.

(20) Zhang, Z.; Dua, R.; Zhang, L.; Zhu, H.; Zhang, H.; Wang, P. Carbon-Layer-Protected Cuprous Oxide Nanowire Arrays for Efficient Water Reduction. ACS Nano 2013, 7, 1709–1717.

(21) Zahran, E. M.; Bedford, N. M.; Nguyen, M. a.; Chang, Y. J.; Guiton, B. S.; Naik, R. R.; Bachas, L. G.; Knecht, M. R. Light-Activated Tandem Catalysis Driven by Multicomponent Nanomaterials. J. Am. Chem. Soc. 2014, 136, 32–35.

(22) Read, C. G.; Steinmiller, E. M. P.; Choi, K.-S. Atomic Plane-Selective Deposition of Gold Nanoparticles on Metal Oxide Crystals Exploiting Preferential Adsorption of Additives. J. Am. Chem. Soc. 2009, 131, 12040–12041.

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(23) Kuo, C.-H.; Yang, Y.-C.; Gwo, S.-J.; Huang, M. H. Facet-Dependent and Au Nanocrystal-Enhanced Electrical and Photocatalytic Properties of Au-Cu2O Core-Shell Heterostructures. J. Am. Chem. Soc. 2011, 133, 1052–1057.

(24) Ho, J.-Y.; Huang, M. H. Synthesis of Submicrometer-Sized Cu2O Crystals with Morphological Evolution from Cubic to Hexapod Structures and Their Comparative Photocatalytic Activity. J. Phys. Chem. C 2009, 113, 14159–14164.

(25) Chang, I.-C.; Chen, P.-C.; Tsai, M.-C.; Chen, T.-T.; Yang, M.-H.; Chiu, H.-T.; Lee, C.-Y. Large-Scale Synthesis of Uniform Cu2O Nanocubes with Tunable Sizes by in-Situ Nucleation. CrystEngComm 2013, 15, 2363.

(26) Pang, M.; Wang, Q.; Zeng, H. C. Self-Generated Etchant for Synthetic Sculpturing of Cu2O-Au, Cu2O@Au, Au/Cu2O, and 3D-Au Nanostructures. Chemistry 2012, 18, 14605–14609.

(27) Li, Q.; Xu, P.; Zhang, B.; Wu, G.; Zhao, H.; Fu, E.; Wang, H.-L. Self-Supported Pt Nanoclusters via Galvanic Replacement from Cu2O Nanocubes as Efficient Electrocatalysts. Nanoscale 2013, 5, 7397–7402.

(28) Pastor, E.; Pesci, F. M.; Reynal, A.; Handoko, A. D.; Guo, M.; An, X.; Cowan, A. J.; Klug, D. R.; Durrant, J. R.; Tang, J. Interfacial Charge Separation in Cu2O/RuOx as a Visible Light Driven CO2 Reduction Catalyst. Phys. Chem. Chem. Phys. 2014, 16, 5922–5926.

(29) Hmadeh, M.; Hoepfner, V.; Larios, E.; Liao, K.; Jia, J.; Jose-Yacaman, M.; Ozin, G. A. New Hydrogen-Evolution Heteronanostructured Photocatalysts: Pt-Nb3O7 (OH) and Cu-Nb3O7(OH). ChemSusChem 2014, 7, 2104–2109.

(30) Zhang, Z.; Wang, Z.; Cao, S.; Xue, C. Au/Pt Nanoparticle-Decorated TiO2 Nano Fibers with Plasmon- Enhanced Photocatalytic Activities for Solar-to-Fuel Conversion. J. Phys. Chem C, 2013, 117, 25939-25947.

(31) Wang, W. N.; An, W. J.; Ramalingam, B.; Mukherjee, S.; Niedzwiedzki, D. M.; Gangopadhyay, S.; Biswas, P. Size and Structure Matter: Enhanced CO2 Photoreduction Efficiency by Size-Resolved Ultrafine Pt Nanoparticles on TiO2 Single Crystals. J Am Chem Soc 2012, 134, 11276–11281.

(32) Wang, C.; Thompson, R. L.; Baltrus, J.; Matranga, C. Visible Light Photoreduction of CO2 Using CdSe/Pt/TiO2 Heterostructured Catalysts. J. Phys. Chem. Lett. 2010, 1, 48–53.

(33) Mahler, B.; Hoepfner, V.; Liao, K.; Ozin, G. Colloidal Synthesis of 1T-WS2 and 2H-WS2 Nanosheets : Applications for Photocatalytic Hydrogen Evolution. J. Am. Chem. Soc. 2014, 136, 14121-14127.

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(35) Kraeutler, B.; Bard, A. J. Heterogeneous Photocatalytic Synthesis of Methane from Acetic Acid - New Kolbe Reaction Pathway. J. Am. Chem. Soc. 1978, 100, 2239–2240.

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Chapter 5 – Light-Assisted Hydrogenation of CO2 to CO Using a Mixed Metal

Oxide Delafossite, CuFeO2

5.1 Abstract

This chapter details our work on the synthesis, characterization and

photocatalytic testing of mixed copper-iron delafossite, CuFeO2, and spinel, CuFe2O4,

as a continuation of our work on Cu2O and Fe2O3 described in previous chapters. We

synthesize the materials from stoichiometric precursor mixtures of metal salts followed

by high temperature solid state reaction to give the desired phases. We characterize the

resulting mixed metal oxides by SEM, PXRD, and EDX and study their formation by

TGA/DSC and PXRD over the course of the heat treatment. We found that both

precursors are first converted to a mixture of Fe3O4 and CuO, which forms the mixed

oxide phases upon continued heating at temperatures above ~800°C. The optimal

annealing times to give phase-pure materials were identified and both ternary oxides

were found to be stable in air post-synthesis with heating up to 450°C. Lacking solubility

in water, CuFeO2 and CuFe2O4 were found to be inactive for aqueous phase H2

evolution from water. When tested for light-assisted CO2 hydrogenation under intense

illumination (~ 20 Suns), CO was produced at an average rate of at least 0.1 mmol g-1 h-

1 over the CuFeO2 catalyst. Isotope labelling experiments with 13CO2 confirmed that the

products largely originated from CO2 reduction and not adventitious carbon

contamination. PXRD analysis of the active catalyst after testing determined that

CuFeO2 transformed to a mixture of metallic Cu and likely Fe3O4 as a result of the

testing conditions. However, despite the apparent structural change, the CO evolution

activity of CuFeO2 was found to be stable over the course at least 90 hours under

intense illumination. Control experiments with the identified by-product phases, namely

Cu, Fe3O4 and a physical mixture of the two, indicated that iron oxide is likely the active

component. Interestingly, the activity of the Fe3O4 and Fe3O4/Cu samples degraded

after prolonged illumination, in contrast to the CuFeO2. Future work on this system will

involve confirming the active components, increasing catalyst surface area, and

decoration with co-catalysts to study changes in product yield and distribution.

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5.2 Introduction

As a continuation of our progress towards the discovery of earth abundant metal

oxides for solar energy conversion, this chapter details our work on mixed metal oxides

of iron and copper, namely delafossites with general formula A(+1)B(+3)O2 and spinels,

A(+2)B2(+3)O4. Having developed expertise working with the binary oxides Cu2O and

Fe2O3, we were interested in incorporating these metals in a single catalyst for solar fuel

generation. As has been emphasized previously, there is huge interest in developing

efficient photoelectrodes based on metal oxides due to their low cost and

thermodynamic stability.1 Unfortunately most oxides possess large bandgaps that

require ultraviolet excitation and are typically n-type, meaning that that they are suitable

for driving oxidation reactions using holes, the minority charge carriers. Photocathode

reactions such as CO2 reduction and H2 evolution require a p-type electrode, preferably

with significant solar spectrum absorption. Some examples based on III-V

semiconductors such as GaAs and GaP, or WSe2 have been reported; however, their

performance depends on expensive high-vacuum processing and deposition techniques

such as CVD and ALD.2–4 The metal pnictides and chalcogenides are also known to

require high overpotentials and are relatively unstable in aqueous solution under

illumination.5,6 Notably, p-type metal oxides have received little attention and represent

an interesting area of study. A recent report described a p-type spinel CaFe2O4

photoelectrode capable of producing hydrogen from water.7 In CaFe2O4, the valence

band primarily has O 2p character and the conduction band is dominated by Fe 3d

states.8 Based on this example, it appears that oxides with Fe 3d-dominated CB states

are energetic enough to drive CO2 reduction. Relative to oxides with valence bands

dominated by O 2p states, VBs consisting of mainly copper 3d10 character have been

calculated to result in higher energy VB edges and therefore smaller bandgaps.9 With

this in mind, Cu(I) delafossite materials such as CuCrO2, CuAlO2, CuRhO2 have been

investigated and shown to be stable in aqueous environments and under excitation.10–12

The stability under illumination was attributed to the fact that optical excitations in

CuMO2 compounds involve metal-to-metal transitions between the Cu d10 orbitals and

the other metal’s 3d states which should protect the metal-oxide bonds and improve

stability.13 Based on all of the above considerations, the most promising candidate

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among the delafossites is CuFeO2; it is naturally p-type due to Cu vacancies in its

structure and unlike Cu2O, has been reported as stable under reductive conditions.14

Furthermore, it has a bandgap of about 1.5 eV and its conduction band, located at - 0.4

V versus the reversible hydrogen electrode (RHE), is suitably positioned to reduce

water under illumination.15 This chapter will describe our work on the synthesis,

characterization and photocatalytic properties of CuFeO2 and its related spinel

CuFe2O4.


There are two predominant approaches in literature to synthesizing CuFeO2:

electrodeposition, and high temperature solid-state reaction. Solid-state approaches are

amenable to producing much larger quantities of material but require long reaction times

at high temperatures due to slow diffusion in the solid state. Electrodeposition

approaches are less energy-intensive and have the advantages of good thin film

formation which enables electrochemical characterization. On the other hand, the

amount of material deposited is small and many cycles may be required to achieve films

thick enough to maximize photocurrent generation. Both approaches involve annealing

or crystallization under inert gas due to the presence of copper in the +1 oxidation state,

which is unstable in ambient conditions. Here we used a 1:1 molar ratio of iron and

copper nitrate hydrates, Fe(NO3)3·9H2O & Cu(NO3)2·5/2 H2O, as starting materials

followed by addition of excess NaOH to give a mixed Cu/Fe oxide-hydroxide precursor

powder. The powder was then heated at 900°C in a tube furnace under Ar atmosphere

for varying periods of time to give the desired delafossite CuFeO2, Figure 5.1. When a

2:1 ratio of iron to copper was used and we heated the resulting powder in air at

1000°C, the related CuFe2O4 spinel phase could be obtained. Both phases were

obtained largely without any impurities as judged by PXRD, which has a detection limit

of approximately 5% by weight. SEM images of the two materials are presented in

Figure 5.2 and show the irregular shapes and large sizes of the resulting mixed metal

oxides. Delafossite and spinel are both high temperature phases and require extended

heating above ~650°C and 700°C respectively in the case of a solid state reaction.

Prolonged heating in the absence of structure-directing agents tends to lead to large

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particle sizes so the CuFeO2 and CuFe2O4 particles are on the order of a few

micrometers.

Figure 5.1 PXRD patterns of as synthesized delafossite CuFeO2 (top) and

CuFe2O4 (bottom)

The BET surface area of the two materials was identical at 2 m2/g. The particles

lacked a consistent morphology; SEM images show large micron-sized chunks dotted

with smaller domains on the order of several hundred nanometers. The powders are

grinded manually in a mortar and pestle following solid-state synthesis, which likely

accounts for the fragmentation into smaller chunks. The synthesis of CuFeO2 can also

be performed without the use of NaOH, which can introduce sodium impurities on the

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surface as was determined by XPS. By simply mixing the two precursors in EtOH and

evaporating to dryness, a precursor powder can be obtained that gives comparably-

sized particles of CuFeO2 when subjected to heat treatment at 900°C under Ar. An

NaOH-free particle is shown in Figure 5.2d notable for its smoother surface and lack of

fragmentation.

Figure 5.2 a) PXRD pattern of the mixed 1:1 molar ratio iron/copper precursor

before heating at 900°C a,c,d) SEM images of CuFeO2 particles following solid state

reaction. Panel e) shows a CuFeO2 particle synthesized without the use of NaOH b)

SEM image of a CuFe2O4 particle after solid state reaction and grinding

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Apart from PXRD, we also confirmed the identity of the mixed metal oxide

phases by EDX, Figure 5.3. Performing EDX line-scans across a single CuFeO2 or

CuFe2O4 particle gave the stoichiometrically expected 1:1 or 1:2 ratio of Cu to Fe. No

other elements were found above the background signal which was arbitrarily assigned

as titanium.

Figure 5.3 a,c) EDX line-scans of CuFeO2 and CuFe2O4 particles showing the

presence of the expected elements (Scale bars 1μm and 2 μm respectively) c,d)

Quantitative signal intensities of the elements detected by EDX confirming expected

stoichiometric ratios of the metals

XPS analysis of CuFeO2 confirmed the 1:1 Fe:Cu ratio that was found by EDX as

well as the presence of Na on the surface as was discussed above. We then tried to

reduce the particle size and increase the surface area of CuFeO2 by ball milling the

powder since surface area is a crucial consideration in catalysis. A high energy ball mill

was used which combines centrifugal forces with ceramic ZrO2 balls as the grinding

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tool. We were able to obtain a much larger proportion of smaller particles with sizes in

the range of a few hundred nanometers up to 1 micron and the BET surface area of the

material increased to 4 m2.

Figure 5.4 a) XPS survey spectrum of CuFeO2 b) PXRD pattern of CuFeO2

following ball-milling to reduce particle size c) PXRD diffraction patterns of CuFeO2

following prolonged heating at 450°C in air d) Diffuse reflectance spectra of CuFeO2

and CuFe2O4

Unfortunately, undesired abrasion of ZrO2 took place as shown in the PXRD

pattern in Figure 5.4b. The PXRD shows that post ball milling, the material was mainly

composed of ZrO2; CuFeO2 was present as a minor phase. We were unable to prevent

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this by reducing the rotation speed of the ball mill and ratio of ZrO2 to sample, therefore

we ruled out ball milling as a viable method for increasing the surface are of our

materials. Figure 5.4c shows the PXRD pattern of CuFeO2 heated at 450°C in air over

the course of 24 hours. We don’t observe any oxidation or degradation of the material

indicating that once formed CuFeO2 is actually quite stable in air despite the fact that its

synthesis requires inert atmosphere. The optical properties of the two mixed oxides

were determined from the diffuse reflectance spectra shown in Figure 5.4d. The spectra

were fitted with a modified Kubelka Munk function to determine the optical band gap of

each sample.16 For CuFeO2, a bandgap of 0.9-1 eV was calculated based on a direct

forbidden transition. The value of the optical gap of CuFeO2 has been reported as being

in the range of 1.3-1.6 eV based on spectroscopic data and DFT modelling.14,17 In the

case of CuFe2O4, an optical band gap of 1.2-1.4 eV has been reported in literature,

however, we did not observe any transitions in that energy range.18 A featureless

reflections spectrum was seen for CuFe2O4 indicative of pseudo-metallic behaviour and

a small bandgap.

Having successfully synthesized the mixed metal oxides, we then aimed to study

their crystallization as a function of time in an attempt to reduce the reaction time and

temperature required to obtain phase-pure materials. Our first step was to look at the

PXRD patterns of CuFeO2 and CuFe2O4 at different periods of time after calcination at

900°C and 1000°C respectively. For CuFeO2, the heating profile includes a fast

temperature ramp up to 900°C over the course of an hour followed by an extended

period of crystallization at that temperature. The PXRD of the powder following the ramp

up without any crystallization period is shown in Figure 5.5a. It shows that the sample is

already largely composed of delafossite with minor impurity phases consisting of CuO

and likely Fe3O4. For clarity, we have omitted the diffractograms obtained after 5 and 10

hours of calcination which show the progressive decrease of the impurity phase with

increasing time. The sample calcined for 20 hours is essentially phase-pure CuFeO2

without any impurity phase by PXRD. It should be noted that the time required to obtain

phase-pure materials scales with the amount of sample; larger scale reactions require

correspondingly longer reaction times. Performing the same study on CuFe2O4 gave

similar results; the desired spinel phase was overwhelmingly the major component after

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1 hour at 1000°C. CuO and Fe3O4 were once again identified as impurity phases which

decrease and eventually disappear with increasing calcination time.

Figure 5.5 a) PXRD patterns of CuFeO2 powder after various amounts of time at

900°C under Ar b) PXRD patterns of CuFe2O4 powder after calcination at 1000°C in air

for various amounts of time

We also performed a TGA/DSC study of the heat treatment process in order to

gain information about the temperatures associated with precursor conversion to the

desired mixed oxide phases. We set the heating rate of the TGA experiment to

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10°C/min to correspond to the heating rate used in synthesizing the materials and

performed the runs under argon and air for CuFeO2 and CuFe2O4 respectively. The

TGA curve for the spinel showed a continuous mass loss up to 400°C followed by a

gradual increase which could be explained by precursor decomposition and subsequent

reaction with atmospheric O2.

Figure 5.6 TGA-DSC curves of CuFeO2 top) and CuFe2O4 bottom) at a heating

rate of 10°C/min. Arrows indicate temperatures where we performed PXRD analysis

The TGA curve of the 1:1 CuFeO2 precursor shows a steep mass loss followed

by a plateau in the range of 300°C - 800°C and smaller mass loss past 800°C. The DSC

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scan seems to indicate 3 phase transitions based on the local maxima in Figure 5.6. We

hypothesize that the delafossite precursor is first converted to Fe3O4 and CuO around

300°C followed by annealing until 800°C at which point crystallization of CuFeO2 takes

place. The arrows indicate temperatures at which a phase transition was suspected to

have taken place; PXRD was used to probe the composition of the material at these

temperatures in Figure 5.7. A PXRD pattern of the 1:1 molar ratio precursor is shown at

the top of Figure 5.7. It shows that prior to heating the precursor is a mixture of CuO

and amorphous iron oxide/hydroxide components. After heating to 250°C over a period

of 25 minutes and annealing at that temperature for 45 minutes, there is little change

observed in sample composition; broad peaks corresponding to CuO are the only

features in the diffractogram. If heated to 550°C and annealed for 45 minutes,

reflections attributed to Fe3O4 begin to appear in the precursor’s PXRD pattern. This

agrees well with the phase transition observed at this temperature in the DSC scan. A

similar mixture of Fe3O4 and CuO is seen when the 2:1 precursor for CuFe2O4 is heated

to 450°C. Based on this data, we can conclude that the precursors are first converted to

a mixture of Fe3O4 and CuO which reacts to form the desired mixed metal oxide at

temperatures above ~800°C based on the stoichiometric metal ratio in the precursor.

Thermal treatment of the corresponding binary oxides is a typical way to synthesize

ternary mixed oxides and has been reported in the preparation of CuFeO2 from CuO,

Fe2O3 and CuFe2O4.19,20

The as-prepared materials were found to be inactive for H2 evolution from

water/MeOH mixtures. The main issue was likely the complete lack of solubility and low

surface area of both CuFeO2 and CuFe2O4 in water due to an absence of surface

functional groups. We increased catalyst loading from the usual 10 mg up to 1g in an

attempt to increase H2 production but the rates obtained were comparable to control

experiments with pure H2O and H2O/MeOH under illumination but in the absence of a

catalyst. We concluded that these materials are a lot more amenable to being tested as

powdered catalysts in a gas-solid heterogeneous system and explored their activity for

light-assisted CO2 hydrogenation.

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Figure 5.7 PXRD spectra of top) 1:1 molar ratio precursor for CuFeO2 prior to

heating middle) 1:1 molar ratio precursor after heating to 250°C and 550°C for a period

of time corresponding to the synthesis of the material bottom) 1:2 molar ratio precursor

(CuFe2O4) after heating at 450°C for amount of time corresponding to the synthesis

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Gas phase hydrogenation experiments were carried out as described in previous

chapters with the exception that the intensity of irradiation was increased up to 20 Suns.

The reactor was kept at room temperature with the only source of heat being the

photothermal effect by the lamp. Under such high photon fluxes, our group has reported

that both photochemical and thermochemical processes can contribute to catalysis.21 A

thermocouple in contact with the sample determined the temperature at the surface of

the film to be approximately 50°C although the local temperature of catalyst particles

may be much higher. Under these conditions, we obtained significant activity for the

reverse water gas shift (RWGS) hydrogenation of CO2 to CO over the CuFeO2 catalyst,

see Figure 5.8. An average CO production rate of approximately 0.1 mmol g-1 h-1 was

calculated which corresponds to 5 µmol g-1 Sun-1 h-1. These rates compare well with our

groups previously published results on CO2 hydrogenation over In2O3-xOHy nanocrystals

and Ru/Si nanowires.21,22 It should be noted that the calculated values are a

conservative estimate of the actual CO production rates. The amount of CO produced

saturated the FID detector on our GC, restricting the peak area used in the calculations

to only 60-70 % of its full value. In addition, normalizing the rate per unit mass is an

overestimation as only ~1/5 of the catalyst film is under illumination and produces

products, as will be discussed below. The CO evolution activity of the CuFeO2 sample is

reproducible and stable over time; we repeated the testing several times and sampled

the amount of CO produced after each run, Figure 5.8a. The activity remains constant

at an average rate of 0.1 mmol g-1 h-1 over a total of more than 90 hours under intense

illumination with 20 Suns. To confirm the origin of CO production, we performed

experiments using isotope-labelled 13CO2 as the reactant gas and analyzed the

products by GC-MS. Figure 5.8b shows the relative intensities of the 28 atomic mass

unit (AMU) peak of 12CO and the 29 AMU peak corresponding to 13CO. The 13CO peak

dominates and we can calculate that over 70% of the CO signal originates from

reduction of 13CO2 to 13CO. No detectable amounts of 13CO are found in the dark or

when an inactive reference catalyst is tested under the same conditions. Corma et.al.

recently published the photocatalytic hydrogenation of CO2 to CO and CH4 on a number

of supported metal oxide catalysts.23 As part of the study, the authors investigated the

use of commercial nanopowders of Fe2O3 and Fe3O4 which were found to produce CO

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at %CO2 conversions rates of 51.0% and 12.3% respectively. These conversions

correspond to CO production rates of approximately 7 mmol g-1 h-1 for Fe2O3 and 1

mmol g-1 h-1 for Fe3O4. These rates are roughly one order of magnitude higher than the

rates observed with our CuFeO2-derived catalysts. We suspect the significantly lower

surface area of our materials (2 m2/g vs ~100 m2/g) is responsible.

Figure 5.8 a) CO production rate over CuFeO2 over a period of 90 hours intense

illumination b) GC-MS chromatogram of the CO peak showing the presence of 13CO

amongst the products c) 12CO and 13CO signals under alternating low and high

illumination intensities d) CO production rates over CuFeO2 and pure Fe3O4, Fe3O4/Cu,

and pure Cu controls. Each bar corresponds to a run with an average length of 7 hours

At this point, we became interested in the origin of activity of the CuFeO2 films.

We noticed a visible discolouration in the area of the film that had been exposed to light

and performed PXRD analysis on the catalyst after testing to determine if structural

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changes had occurred. Figure 5.9c shows the resulting PXRD pattern and a digital

photograph of the film after illumination.

Figure 5.9 a) Top down SEM image of the CuFeO2 film prior to illumination b)

Cross-sectional SEM image of the CuFeO2 film used to determine approximate

thickness c) PXRD diffraction pattern of the CuFeO2 catalyst after testing showing the

presence of newly formed Fe3O4 and Cu phases (Inset – digital photograph of

discolouration following illumination)

The presence of Fe3O4 or γ-Fe2O3 and Cu phases is evident, likely as a

by-product of CuFeO2 transformation. It should be noted that despite the fact that

135

CuFeO2 still appears to be the dominant phase, the by-product phases were only

located on the illuminated area of the film whereas PXRD samples the film as a whole.

It is therefore likely that conversion to Fe3O4 and Cu is complete under light of this

intensity. This will be verified in future experiments by only looking at the PXRD pattern

of the illuminated spot. Figure 5.9a,b shows SEM images of the catalyst film prior to

testing; no morphological change is seen after the apparent decomposition of CuFeO2.

Despite the apparent degradation of the original CuFeO2 catalyst, CO production

remained stable after continued illumination of the same sample area, Figure 5.8a. A

sample sputtered with 50nm of Ru was also tested but did not show significantly

improved rates compared to CuFeO2. We then varied the intensity of the light source as

shown in Figure 5.8c. Under low illumination (3 Suns), we did not observe any 13CO by

GC-MS; only a low signal attributed to 12CO contamination was detected. We then

illuminated the same sample with light of 20 Sun intensity and observed the typical

amounts of 13CO being evolved. We then decreased the lamp intensity down to 3 Suns

again, and once more did not see any 13CO. This indicated that the sample is not

activated by an initial illumination under high light intensity, but rather that the high light

intensity is necessary for activity. Finally, we hypothesized that the component

responsible for activity was not CuFeO2 but one or both of the decomposition by-

product phases. We prepared films of pure Cu, pure Fe3O4, and a 50 w/w % mixture of

Fe3O4/Cu from the corresponding commercially available nanopowders. Pure Cu was

not active but the Fe3O4 and Fe3O4/Cu samples produced comparable rates of 13CO as

the CuFeO2 films, suggesting that the iron oxide component was the active phase in our

system. However, both control samples sharply decreased in activity upon repeated

testing in contrast to the CuFeO2 catalyst which maintained its performance for over 90

hours, see Figure 5.8d. The Fe3O4 films were prepared from nanoparticle powders in

the 50-100 nm size range and could sinter and aggregate leading to the observed loss

of activity. The cause of the unexpected stability of the CuFeO2-derived materials will be

investigated further in future experiments. The observation that iron oxide is the active

component agrees well with literature reports, where iron oxide is a major component in

the most widely applied water gas shift catalyst (75% Fe2O3, 10% Cr2O3, 0.2% MgO,

balance volatiles).24 This mixture is typically used in high temperature WGS processes

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in the range of 320 - 450°C and so its activity at room temperature under illumination

could be a good way to reduce the energy requirements of the process. In the interest

of determining whether the activity we observed was indeed catalytic or simply

stoichiometric, we performed a quick “back of the envelope calculation” to estimate the

number of active metal sites in our films. In the case of a typical overnight run using the

CuFeO2 catalyst, we obtain a minimum of 1 mmol CO, or approximately 6 x 1020

molecules of product. Using liberal over-estimates for the particle size, film thickness

and area, we calculate that the number of metal atoms in the illuminated sample is

approximately 3.9 x 1020. Keeping in mind that we included Cu atoms in the calculation,

as well as all metal sites in the bulk of the catalyst particles, the likely number of active

sites is much lower. This suggests that the turnover number is greater than 1, which

would be indicative of a catalytic process. Much work remains to be done on this system

beginning with investigating the increased stability of the CuFeO2-derived catalysts. We

would also like to explore new synthetic approaches to reducing particle size and

increasing surface area. We will monitor catalytic performance with various cut-off filters

which block parts of the solar spectrum as well as investigate if hydrocarbons such as

CH4 can be produced with co-catalysts or different reaction conditions.

5.4 Conclusions

This chapter was mainly focused on the chemical and physical properties and

photocatalytic activity of a mixed metal oxide delafossite, CuFeO2, for the photo-

assisted hydrogenation of CO2 to CO. CuFeO2 was synthesized according to a standard

high temperature solid state method from iron and copper nitrate precursors at 900°C

under an Ar atmosphere. Using a combination of PXRD and TGA/DSC we determined

the optimal reaction time and crystallization temperatures required to obtain phase pure

CuFeO2 and its related spinel CuFe2O4. The materials were inactive for H2 evolution

from water due to a lack of solubility, but CuFeO2 produced significant amounts of CO

when tested for the reverse water gas shift reaction under 20 Suns illumination in the

gas phase. The detection of mostly 13CO by GC-MS indicated that CO2 reduction was

genuine and not a result of carbon contamination. The catalytic activity of the CuFeO2

was stable over prolonged irradiation. As a result of the high light intensity, we found

137

that CuFeO2 was decomposing to a mixture of Fe3O4/γ-Fe2O3 and metallic Cu. We

tested these phases as control samples and found that the iron oxide domain was likely

responsible for the activity, however the control films largely deactivated under intense

sunlight in contrast to CuFeO2. Future experiments will attempt to enhance surface

area, determine the origin of the stability, and examine the effect of co-catalysts on the

rate and product distribution. Light-driven catalysis of the water gas shift reaction at

significantly reduced temperatures could be a promising contribution towards reducing

the energy requirements of that process.


General Characterization - All chemicals were purchased from Sigma Aldrich

and used directly without further purification. Powder X-ray diffraction (PXRD) was

performed on a Bruker D2-Phaser X-ray diffractometer, using Cu Kα radiation at 30 kV.

Absorption and reflection measurements were performed using a Lambda 1050

UV/VIS/NIR spectrometer from PerkinElmer equipped with an integrating sphere with a

diameter of 150 mm and a center mount holder. Nitrogen Brunauer-Emmet-Teller (BET)

adsorption isotherms were obtained at 77 K using a Quantachrome Autosorb-1-C.

TGA/DSC curves were acquired on a TA Instruments Q500 thermogravimetric analyzer

at a constant ramp rate of 5°C under Ar or air atmosphere. Ball milling was performed

for a total of 5 hours with a Fritsch Pulverisette 7 planetary micro mill using ZrO2 balls as

the ceramic material at 1000 rpm. The amount of CuFeO2 used was 3g along with 20g

of ZrO2 balls with a diameter of ~1cm each. The sample had to be periodically taken out

and cooled down to room temperature.

Synthesis of CuFeO2 & CuFe2O4 – To prepare the precursor for CuFeO2

deposition, 1.392 g (6 mmol) of Cu(NO3)2·5/2 and 2.424 g (6 mmol) of Fe(NO3)3·9H2O

were dissolved in 40 mL of absolute EtOH at room temperature. 1.2 g (30 mmol) of

NaOH were dissolved in a further 40 mL of EtOH with slight heating by heat gun on the

outside of a 50 mL round bottom flask. The warm NaOH solution was added to the

metal precursors and allowed to react under vigorous magnetic stirring for 30 minutes at

room temperature. The mixture was centrifuged at 2000 rpm for 5 minutes and washed

138

with deionized water 3 times and EtOH once. The powder was dried in vacuo for 48

hours at 60°C and grinded up in an agate mortar and pestle. The precursor was placed

in an alumina crucible inside a tube furnace and purged with flowing Ar for 1hour at

room temperature followed by a gradual ramp to 900°C over the course of 1 hour. The

temperature was held at 900°C for varying amounts of time (0-20 hrs) to study the effect

of time on the crystallization. CuFe2O4 was prepared in the same way except 0.696 g (3

mmol) of Cu(NO3)2·5/2 were used and the mixed precursor powder was treated in air at

1000°C in a box furnace. Alternatively, CuFeO2 the use of NaOH by simply dissolving

the desired stoichiometric ratio of starting materials in 50 mL of absolute EtOH and

heating on an oil bath to evaporate the solvent. Subsequent heat treatment gives the

desired delafossite and eliminates the need to wash away sodium hydroxide. CuFe2O4

is not accessible without the use of NaOH.

Catalyst Fabrication - Sample films were prepared for photocatalytic testing by

drop-casting 100 mg of each sample powder – suspended via sonication in absolute

EtOH – onto 1” × 1” binder free borosilicate glass microfiber filters (Whatman, GF/F, 0.7

μm) placed on top of a vacuum filtration funnel held under very weak vacuum. The

borosilicate filters were then dried in vacuo at 60°C for 48 hours. For samples

containing Ru, the CuFeO2 powder was first drop-cast on high resistivity glass

substrates (2 inch by 2 inch). The wafers were initially cleaned with piranha solution

(H2SO4:H2O2 = 3:1 by volume) for 3 h and then rinsed with de-ionized water. Ru was

sputtered onto these samples in a custom-built sputtering system (Kurt J. Lesker Co.)

by radio frequency (RF) magnetron sputtering using a 99.95% pure Ru sputtering target

purchased from Angstrom Sciences, Inc. The base pressure of the sputtering chamber

was pumped down to 1 × 10 −7 Torr before argon was introduced into the chamber at a

flow rate of 20 sccm. The chamber pressure was set to 3 mTorr during the deposition,

which was carried out at room temperature. The forward power was 100 W and the

substrate-to-target distance was 14 cm. The sputtering process was terminated when

10 nm of Ru, as measured from an in-situ thickness monitor (SQM-242 from Sigma),

had been deposited. The CuFeO2/Ru was then scraped off the glass substrate,

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suspended in absolute EtOH (2 mL) and drop-cast onto the borosilicate glass microfiber

filter for testing


H-7000 conventional TEM operating at 100kV. SEM Images and EDX spectra of the

particles were acquired on a Hitachi S-5200 operating at 30kV using an Oxford Inca

detector. Sample preparation involved dropping a dilute nanocrystal solution on a

carbon coated Ni TEM grid. Images of the catalyst films were acquired on an FEI

Quanta FEG250 operating at 20kV. The catalyst film was glued on a standard aluminum

SEM pin stub mount for imaging.












Aqueous Photocatalytic Testing - All photocatalytic tests were performed in a

home built, gas-tight pyrex reactor with a quartz window. The photocatalyst powder (20

gL-1) was dispersed in distilled water and methanol (25 vol%, HPLC grade). Prior to

each run, the reactor was purged with nitrogen for 20 min and the first sample was

taken before the photocatalytic reaction was started. A Newport 120 W or 300 W Xe

lamp was used for solar irradiation and an air mass (AM) 1.5 filter (Newport) was

applied to simulate the direct solar spectrum when the sun is at a zenith angle of 48.28.

To test the long-term stability of the photocatalyst and the reproducibility of the H2

formation, the reactor was purged after a few hours to remove all products and start a

140

new formation cycle. Gas samples were extracted with a gas-tight syringe, separated by

gas chromatography (GC, Agilent 7820A GC) equipped with a thermal conductivity

detector (TCD). A 1.5 m Molesieve 13 X column (80–100 mesh) was used for the

separation of H2, O2, and N2, and an integrated 1 m Haysep Q column (80–100 mesh)

was used to first separate CO2 and water vapour to avoid contamination of the

Molesieve column.
















Haysep D column. The reactor was held in a custom designed stand. The thermocouple

was placed at the front face of the sample and shielded from incident light unless

otherwise specified. Isotope tracing experiments were performed using 13CO2 (99.9

atomic% Sigma Aldrich). Isotope product gases were separated using a 60 m GS-

Carbonplot column and measured using an Agilent 7890A gas chromatographic mass

spectrometer (GC-MS).

141

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(2) McKone, J. R.; Pieterick, A. P.; Gray, H. B.; Lewis, N. S. Hydrogen Evolution from Pt/Ru-Coated P-Type WSe2 Photocathodes. J. Am. Chem. Soc. 2013, 135, 223–231.

(3) Khaselev, O.; Turner, J. a. A Monolithic Photovoltaic-Photoelectrochemical Device for Hydrogen Production via Water Splitting Science 2014, 280, 425-427.

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(5) Aurian-Blajeni, B.; Halmann, M.; Manassen, J. Electrochemical Measurement on the Photoelectrochemical Reduction of Aqueous Carbon Dioxide on p-Gallium Phosphide and p-Gallium Arsenide Semiconductor Electrodes. Sol. Energy Mater. 1983, 8, 425–440.

(6) Preusser, S.; Herlem, M.; Etcheberry, A.; Jaume, J. The Photodissolution of InP. Electrochim. Acta 1992, 37, 289–295.

(7) Ida, S.; Yamada, K.; Matsunaga, T.; Hagiwara, H.; Matsumoto, Y.; Ishihara, T. Preparation of p-Type CaFe2O4 Photocathodes for Producing Hydrogen from Water. J. Am. Chem. Soc. 2010, 132, 17343–17345.

(8) Matsumoto, Y.; Obata, M.; Hombo, J. Photocatalytic Reduction of Carbon Dioxide on p-Type CaFe2O4 Powder. J. Phys. Chem. 1994, 98, 2950–2951.

(9) Galakhov, V. R.; Poteryaev, A. I.; Kurmaev, E. Z.; Anisimov, V. I.; Bartkowski, S.; Neumann, M.; Lu, Z. W.; Klein, B. M.; Zhao, T.-R. Valence-Band Spectra and Electronic Structure of CuFeO2. Phys. Rev. B 1997, 56, 4584–4591.

(10) Saadi, S.; Bouguelia, A.; Trari, M. Photocatalytic Hydrogen Evolution over CuCrO2. Sol. Energy 2006, 80, 272–280.

(11) Gu, J.; Yan, Y.; Krizan, J. W.; Gibson, Q. D.; Detweiler, Z. M.; Cava, R. J.; Bocarsly, A. B. p-Type CuRhO2 as a Self-Healing Photoelectrode for Water Reduction under Visible Light. J. Am. Chem. Soc. 2014, 136, 830–833.

(12) Smith, J. R.; Van Steenkiste, T. H.; Wang, X. G. Thermal Photocatalytic Generation of H2 over CuAlO2 Nanoparticle Catalysts in H2O. Phys. Rev. B - Condens. Matter Mater. Phys. 2009, 79, 2–5.

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(13) Thurston, T. R.; Wilcoxon, J. P. Photooxidation of Organic Chemicals Catalyzed by Nanoscale MoS2. J. Phys. Chem. B 1999, 103, 11–17.

(14) Read, C. G.; Park, Y.; Choi, K.-S. Electrochemical Synthesis of p-Type CuFeO2 Electrodes for Use in a Photoelectrochemical Cell. J. Phys. Chem. Lett. 2012, 3, 1872–1876.

(15) Prévot, M. S.; Guijarro, N.; Sivula, K. Enhancing the Performance of a Robust Sol-Gel-Processed P-Type Delafossite CuFeO2 Photocathode for Solar Water Reduction. ChemSusChem 2015, 8, 1359-1367.

(16) López, R.; Gómez, R. Band-Gap Energy Estimation from Diffuse Reflectance Measurements on Sol-Gel and Commercial TiO2: A Comparative Study. J. Sol-Gel Sci. Technol. 2012, 61, 1–7.

(17) Ong, K. P.; Bai, K.; Blaha, P.; Wu, P. Electronic Structure and Optical Properties of AFeO2 (A = Ag, Cu) within GGA Calculations. Chem. Mater. 2007, 19, 634–640.

(18) Reitz, C.; Suchomski, C.; Haetge, J.; Leichtweiss, T.; Jagličić, Z.; Djerdj, I.; Brezesinski, T. Soft-Templating Synthesis of Mesoporous Magnetic CuFe2O4 Thin Films with Ordered 3D Honeycomb Structure and Partially Inverted Nanocrystalline Spinel Domains. Chem. Commun. 2012, 48, 4471.

(19) Bassaid, S.; Chaib, M.; Omeiri, S.; Bouguelia, a.; Trari, M. Photocatalytic Reduction of Cadmium over CuFeO2 Synthesized by Sol-Gel. J. Photochem. Photobiol. A Chem. 2009, 201, 62–68.

(20) Gu, J.; Wuttig, A.; Krizan, J. W.; Hu, Y.; Detweiler, Z. M.; Cava, R. J.; Bocarsly, A. B. Mg-Doped CuFeO 2 Photocathodes for Photoelectrochemical Reduction of Carbon Dioxide. J. Phys. Chem. C 2013, 117,

(21) O’Brien, P. G.; Sandhel, A.; Wood, T. E.; Jelle, A. a.; Hoch, L. B.; Perovic, D. D.; Mims, C. A.; Ozin, G. A. Photomethanation of Gaseous CO2 over Ru/Silicon Nanowire Catalysts with Visible and Near-Infrared Photons. Adv. Sci. 2014, 1.

(22) Hoch, L. B.; Wood, T. E.; O’Brien, P. G.; Liao, K.; Reyes, L. M.; Mims, C. A.; Ozin, G. A. The Rational Design of a Single-Component Photocatalyst for Gas-Phase CO2 Reduction Using Both UV and Visible Light. Adv. Sci. 2014, 1.

(23) Sastre, F.; Puga, A.V.; Liu, L.; Corma, A.; Garcia, H. J. Am. Chem. Soc. 2014, 136, 6798-6801.

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Chapter 6 – Conclusions and Future Outlook

6.1 Concluding Remarks

In the work presented in this thesis, we explored the potential of various earth

abundant metal oxides to be used as catalysts for light assisted generation of fuels from

CO2 and H2O. Our approach was driven by a focus on low cost materials that currently

lack practical efficiencies but have the potential to be scaled to globally significant

quantities with improved performance. We approached this challenge from a synthetic

materials chemistry perspective based on our expertise in nanomaterials. In doing so,

we made use of a number of synthetic methods to access the desired materials. In

Chapters 2 and 3, we used high temperature, organic-phase, colloidal chemistry

techniques to obtain previously unreported oxide-oxide heterojunction nanocrystals

based on Fe2O3 and Cu2O. We then transitioned to low-temperature, aqueous phase

approaches to prepare Cu2O nanocubes and various multicomponent architectures. In

chapter 5, we prepared mixed metal oxides based on Cu and Fe using classic solid-

state chemistry techniques. Some key findings and advances presented in the thesis

include:

Synthesis of novel Fe2O3/Cu2O hetero-structured nanocrystals and elucidation

of their hybrid properties

Evaluation of low-pressure UV photolysis for organic ligand removal on colloidal

nanocrystals by XPS and comparison to thermal and colloidal ligand-exchange

techniques

Studying the effect of hole scavengers on gas-phase product distribution for

water splitting experiments in the presence of widely used P25 and P25/(Pt, Au)

catalysts

Investigating the light-assisted hydrogenation of CO2 to CO by an iron/copper

delafossite, CuFeO2

Future work will focus on the CuFeO2 system and determining the origin of its

apparent catalytic properties for CO2 hydrogenation. Firstly we would like to increase

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the particle’s surface area by decreasing their average size. Infiltrating the channels of

commercially available periodic mesoporous SiO2 with the CuFeO2 precursors would

confine the growth of the delafossite to within the pores of the SiO2 structure. The

template can then be dissolved away with NaOH(aq) to give the delafossite in

nanoparticulate form. We would then examine the effect of surface area on product

rates. In terms of further gas-phase testing, we would like to examine the effect of light

intensity and wavelength on the observed rates. A series of long-wavelength pass filters

(ex. 400 nm, 500 nm etc) will be used to block out the shorter wavelengths of the visible

spectrum while we monitor activity. Based on the narrow bandgap of CuFeO2, the entire

visible spectrum would be expected to contribute to its photocatalytic properties. In

addition, we have plans to perform a full light intensity study by varying the illumination

from 1 Sun to 25 Sun equivalents in order to establish a minimum illumination threshold.

We will then switch testing from a batch to a flow setup which will enable us to do

continuous product as a function of illumination or temperature and test powders as

opposed to thin films. In the flow reactor, we will be able to heat the catalyst to

temperatures up to 500°C and monitor its performance in the absence of illumination.

This will give us some insights into how much of the activity is driven by a photothermal

effect as opposed to a photocatalytic one. There are plans to vary the reactant gases;

we are equipped to detect the products of water splitting and so it would be interesting

to study if CuFeO2 is able to drive gas-phase overall water splitting. If the sample

transforms into Cu and Fe3O4, the metal component could be a reducing site with

oxidation taking place on the oxide.

At this stage it would be useful to summarize some of the lessons learned over the

course of this work, which may be valuable to fellow researchers in the field of artificial

photosynthesis. One of the main challenges that we encountered was dealing with the

presence of organic capping ligands on the surface of our nanoparticulate catalysts.

Molecules containing long hydrocarbon chains and a polar head group such as

oleylamine, oleic acid, or trioctylphosphine oxide are invariably used to mediate the

growth of nanocrystals in organic solvents and to provide solubility in non-polar media.

Unfortunately, they also insulate adjacent nanocrystals from electronic coupling, prevent

reactants reaching the nanoparticle surface and can give false positive hydrocarbon

145

signals when testing for CO2 reduction products by GC. In Chapter 3, we described the

main techniques for removing hydrocarbon ligands including solution-phase ligand

exchange, thermal calcination, and UV photolysis. For applications that require

nanocrystals to be soluble, chemical removal by ligand exchange is the preferred

approach. Care must be taken in ascertaining that the incoming ligand reacts only with

the organic shell and not the inorganic nanocrystal core as we observed using NOBF4

and the Fe2O3/Cu2O HNCs. For applications where thin films or nanocrystal powders

are desired, physical carbon removal methods are more suitable. Simple thermal

treatment may be effective as long as it does not induce structural changes in the

nanocrystals. UV photolysis was found to be very effective and faster than thermal

treatment in removing carbon contamination and we recommend other researchers to

consider this method as an option. One drawback of this technique is the slow “batch”

throughput which limits the efficiency of the process. Surface rearrangements caused

by prolonged UV irradiation can also occur as evidenced by the UPS data in Chapter 3.

Overall, we would advocate avoiding the use of long-chain hydrocarbons whenever

possible by using aqueous phase synthetic methodologies. High temperature reactions

in organic media afford exquisite control over particle size and shape but this many not

be necessary for many applications where monodispersity is not crucial. We would also

like to state the importance of 13C isotope labelling for determining whether the detected

products of CO2 reduction experiments originate from CO2 or from adventitious carbon

contamination. Many publications omit this experiment but still report active catalysts.1–3

Relevant control experiments such as carrying out the desired reactions in the absence

of light and/or CO2 are certainly useful and support manuscripts’ claims of

photocatalysis. However, using 13CO2 as the reactant gas and detecting 13C-labelled

products by GC-MS remains the most convincing way to demonstrate genuine reduction

of CO2 to hydrocarbons.

The nature of the reaction medium is another consideration that must be carefully

taken into account by researchers. Most publications report liquid-phase systems where

the active catalyst is suspended in CO2-saturated aqueous or organic solutions. This

setup has several drawbacks as the solubility of CO2 in water is quite low and the

presence of the electrolyte can contribute to catalyst decomposition through

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photocorrosion. These challenges can be largely resolved by using light-assisted gas-

solid heterogeneous catalysis. In the gas-phase, high pressures of reactant gases are

easily achieved and catalyst degradation may be less of an issue. We believe that

working in the gas-phase is the better approach when it comes to doing artificial

photosynthesis on industrial scales whereas aqueous setups are better suited to

preliminary testing in laboratory environments. In terms of materials development, we

believe it will be difficult to discover a single “magic bullet” material with the required

light absorption, charge transport, and surface chemistry properties which can drive

overall CO2 reduction or H2O splitting. Multi-component architectures which combine

light absorbing and catalytic domains have been shown to be capable of driving these

processes albeit at very low efficiencies of fractions of a percent.4 An advantage of

these approaches is they allow for separate optimization of each functional domain

which can then be interfaced with other components in the final device. In particular,

combining semiconductor light absorbers with molecular catalysts is a very promising

approach that combines efficient solar absorption with selective CO2 conversion. The

major limiting factor here is still the catalysis itself regardless of whether one uses

metal/semiconductor surfaces, organometallic complexes, or nanocrystals.5–8

Developing efficient, earth-abundant H2 and O2 evolving catalysts along with those

capable of reducing CO2 to hydrocarbons will remain a topic of intense interest in the

coming years.

Furthermore, we have increasingly found that the classical solid state physics way

of thinking is unable to fully describe the complexities of CO2 reduction. Judicious

selection of materials with proper conduction and valence band energies is not sufficient

to ensure CO2 conversion; the surface chemistry of the catalyst is evidently equally

important. We noted evidence of this with Cu2O nanocubes in Chapter 4 which were

completely inactive when it came to doing chemistry with CO2 and H2O despite having

energetically suitable band energies. It is therefore crucial for researchers to develop an

understanding of the surface active sites, kinetics, and reaction mechanisms of potential

nanoparticulate catalysts. As an example, our group recently reported a detailed

spectroscopic and DFT analysis of the surface reaction chemistry responsible for the

light-driven conversion of CO2 to CO by hydroxylated indium oxide nanocrystals.9 We

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proposed a reaction mechanism whereby surface active sites of In2O3-x(OH)y are

composed of a Lewis basic hydroxide adjacent to a Lewis acidic indium, which in the

presence of an oxygen vacancy, assist the adsorption and heterolytic dissociation of H2

which enables the reaction of CO2 to form CO and H2O. The proximal Lewis acid and

base sites suggest a mechanism analogous to molecular Frustrated Lewis pairs, an

exciting discovery for surface chemistry on semiconductors. Elucidating similar

structure-property-activity relationships will be a crucial step for future researchers in the

field of solar fuels.

6.2 Future Outlook for Solar Fuels

In this concluding section of the thesis, we would like to share our views on the

current directions and future development of solar fuel technologies towards

commercialization, once again with an emphasis on CO2 as the primary carbon source.

Several approaches currently being considered including solar-driven photocatalytic,

photo-electrochemical, and electrochemical CO2 conversions, alongside CO2

hydrogenation with solar derived H2, and solar-driven thermochemical processes.10 It is

evident that two principal routes emerge from the myriad CO2 conversion technologies

summarized in Figure 6.1:

1. Catalytic conversion using solar derived H2

2. Direct CO2 reduction with H2O

In the first route, CO2 is hydrogenated by H2 derived from solar energy through

various technologies including photo-electrochemical water splitting, electrolysis, and

thermochemical methods. Water electrolysis is a commercially demonstrated

technology, and has reasonably high system efficiencies. Coupling electrolytic modules

with photovoltaic panels is a near-term solution to producing clean hydrogen with

reasonable solar to H2 efficiencies of approximately 10%.6

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Figure 6.1 A general schematic of solar fuels production describing many of the

proposed methods of converting CO2 to fuels using solar energy. The approximate

temperature requirements are color-coded, red = high, yellow = ambient (Reprinted with

permission from Ref (6) Copyright (2015) Royal Society of Chemistry)

Alternatively solar thermal energy can be used to drive thermochemical cycles

that have the net effect of releasing hydrogen and oxygen.11 These are typically based

on metal oxides such as ZnO or Fe3O4 and operate at temperatures above 1000°C, see

Equations 6.1 and 6.2.10 Such thermochemical cycles can achieve a solar to H2

conversion efficiency of almost 20% however they are complex to engineer and require

large up-front capital investments.12

MOx → MOx-α + α/2O2 (6.1)

MOx-α + αH2O → αH2 + MOx (6.2)

Photoelectrochemical and photocatalytic water splitting approaches have

generated tremendous interest for their ability to generate hydrogen from water using

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only solar irradiation or a small applied potential. Important strides have been made in

recent years towards improving the performance of these technologies but many issues

still remain. Photo-electrochemical cells are only able to achieve competitive H2

evolution rates using expensive multi-junction systems such as Nocera’s artificial leaf.13

Despite using earth abundant electrodes such as CoPi and NiMoZn alloy, the

engineering cost of this system was deemed too high for commercialization. In a similar

vein, photocatalytic systems are attractive for their simplicity but efficiencies lag even

further behind PEC technology. The state of the art catalyst based on a Ga1-xZnxN1-xOx

solid solution with a mixed Rh-Cr oxide co-catalyst only achieves a solar to H2

conversion of 0.2 %.14 The H2 generated by the above methods can then be used to

reduce CO2 to fuels using mature industrial processes as has been discussed earlier. In

particular, the reverse-water–gas-shift reaction (RWGS) can be used to convert CO2

and hydrogen to CO and water as we observed in Chapter 5. CO mixed with hydrogen

in varying proportions produces syngas, which can be used to synthesize a variety of

products, including methanol, dimethylether, or hydrocarbons through Fischer–Tropsch

synthesis. There also exist direct routes for hydrogenating CO2 to products

hydrocarbons including methanol, methane, and formic acid.

Besides these indirect hydrogenation approaches, CO2 can also be directly

converted to fuels using H2O and solar energy electrocatalytic, photo-electrochemical,

thermochemical, or photocatalytic methods. The simplicity of such direct CO2

conversion processes in this category makes them very attractive despite the fact that

these technologies are not as mature as those in the previous section. The major

challenge is twofold; firstly CO2 solubility in the aqueous phase is low which leads to low

reduction rates. Additionally, H2 evolution from water competes with CO2 reduction

because of its similar thermodynamic reduction potential thereby lowering CO2

conversion efficiencies. Nevertheless, important progress has been made in the

electrocatalytic, photo-electrochemical, and photocatalytic conversion of CO2. Much of

the electrocatalytic work has been done on metal surfaces with copper being most

effective at producing hydrocarbons from CO2.15 The major product is usually CO on

Ag, Au, and Pd, formate on Pb, Hg, In, Cd, and Sn, and hydrogen on Ni, Pt, Fe and Ti.6

Semiconductor surfaces have also been investigated with the most common

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photocathodes being Si, p-InP, p-GaAs, and p-CdTe.16 Many semiconductor systems

are designed to operate in organic solvents instead of water to circumvent competitive

H2 evolution. Common solvents include acetonitrile, dimethylformamide, methanol, and

ionic liquids, but the use of such additives renders the overall process less

sustainable.10,16 It quickly becomes evident that the major challenge of CO2 reduction is

catalytic in nature. In particular conversion rates and selectivities lag far behind those of

water splitting prohibiting any economically viable implementation of CO2 conversion

catalysis. At the same time, this means that there is a huge incentive for chemists to

come up with materials solutions to these challenges. A number of well-funded research

centres are currently tackling CO2 reduction including the Joint Centre for Artificial

Photosynthesis (JCAP) in California, the DOE-sponsored Argonne-Northwestern Solar

Energy Research Center, and the SUNCAT Centre for Interface Science and Catalysis

at Stanford Unieversity.6

As society moves towards reducing our dependence on fossil fuels and the

implementation of more sustainable energy resources, we are faced with one of our

biggest technological challenges yet. In order to attract the attention of the general

public and lawmakers, solar fuel technologies must reach a point of commercial

competitiveness with fossil fuels. With a sustained collaborative effort among

researchers from the fields of chemistry, chemical engineering, materials science, and

physics, we are capable of realizing our goal of an energy economy based on

renewable fuels. We fully believe that with continued funding and extensive research

efforts focused on developing cheap, efficient catalyst materials to overcome the

challenges of insufficient activity, poor product selectivity and low stability, the

technology of solar fuels will move towards commercialization in the near future.

6.3 References

(1) Liang, Y. T.; Vijayan, B. K.; Gray, K. A.; Hersam, M. C. Minimizing Graphene Defects Enhances Titania Nanocomposite-Based Photocatalytic Reduction of CO2 for Improved Solar Fuel Production. Nano Lett. 2011, 11, 2865–2870.

(2) Liu, Q.; Zhou, Y.; Kou, J.; Chen, X.; Tian, Z.; Gao, J.; Yan, S.; Zou, Z. High-Yield Synthesis of Ultralong and Ultrathin Zn2GeO4 Nanoribbons toward Improved

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Photocatalytic Reduction of CO2 into Renewable Hydrocarbon Fuel. J. Am. Chem. Soc. 2010, 132, 14385–14387.

(3) Xi, G.; Ouyang, S.; Li, P.; Ye, J.; Ma, Q.; Su, N.; Bai, H.; Wang, C. Ultrathin W18O49 Nanowires with Diameters below 1 nm: Synthesis, Near-Infrared Absorption, Photoluminescence, and Photochemical Reduction of Carbon Dioxide. Angew. Chemie, Int. Ed. 2012, 51, 2395–2399, S2395/1–S2395/9.


(5) Luo, J.; Im, J.; Mayer, M. T.; Schreier, M.; Nazeeruddin, M. K.; Park, N.; Tilley, S. D.; Fan, H. J.; Grätzel, M. Water Photolysis at 12.3% Efficiency via Perovskite Photovoltaics and Earth-Abundant Catalysts. Science 2014, 345, 1593-1596.

(6) Sastre, F.; Puga, A. V; Liu, L.; Corma, A.; García, H. Complete Photocatalytic Reduction of CO2 to Methane by H2 under Solar Light Irradiation. J. Am. Chem. Soc. 2014, 136, 6798–6801.

(7) Manthiram, K.; Beberwyck, B. J.; Alivisatos, A. P. Enhanced Electrochemical Methanation of Carbon Dioxide with a Dispersible Nanoscale Copper Catalyst. J. Am. Chem. Soc. 2014.

(8) Finn, C.; Schnittger, S.; Yellowlees, L. J.; Love, J. B. Molecular Approaches to the Electrochemical Reduction of Carbon Dioxide. Chem. Commun. 2012, 48, 1392.

(9) Ghuman, K. K.; Wood, T. E.; Hoch, L. B.; Mims, C. A.; Ozin, G. A.; Singh, C. V. Illuminating CO 2 Reduction on Frustrated Lewis Pair Surfaces: Investigating the Role of Surface Hydroxides and Oxygen Vacancies on Nanocrystalline In2O3−x (OH)y. Phys. Chem. Chem. Phys. 2015, 17, 14623–14635.

(10) Herron, J. A.; Kim, J.; Upadhye, A. A.; Huber, G. W.; Maravelias, C. T. A General Framework for the Assessment of Solar Fuel Technologies. Energy Environ. Sci. 2015, 8, 126–157.

(11) Steinfeld, A. Solar Thermochemical Production of Hydrogen - A Review. Sol. Energy 2005, 78, 603–615.

(12) Ferreira, L. S.; Trierweiler, J. O. Modeling and Simulation of the Polymeric Nanocapsule Formation Process. IFAC Proc. Vol. 2009, 7, 405–410.

(13) Reece, S. Y.; Hamel, J. a.; Sung, K.; Jarvi, T. D.; Esswein, a. J.; Pijpers, J. J. H.; Nocera, D. G. Wireless Solar Water Splitting Using Silicon-Based Semiconductors and Earth-Abundant Catalysts. Science 2011, 334, 645–648.

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(14) Maeda, K.; Teramura, K.; Lu, D.; Takata, T.; Saito, N.; Inoue, Y.; Domen, K. Photocatalyst Releasing Hydrogen from Water. Nature 2006, 440, 295.

(15) Yano, J.; Yamasaki, S. Pulse-Mode Electrochemical Reduction of Carbon Dioxide Using Copper and Copper Oxide Electrodes for Selective Ethylene Formation. J. Appl. Electrochem. 2008, 38, 1721–1726.

(16) Kumar, B.; Llorente, M.; Froehlich, J.; Dang, T.; Sathrum, A.; Kubiak, C. P. Photochemical and Photoelectrochemical Reduction of CO2. Annu. Rev. Phys. Chem. 2012, 63, 541–569.

Abundant Metal Oxide Nanomaterials for Solar Fuel Generation

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