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Enhanced photocatalysis on titanium oxide passivated III-V semiconductors
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Enhanced photocatalysis on titanium oxide passivated III-V semiconductors
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Content
ENHANCED PHOTOCATALYSIS ON TITANIUM OXIDE PASSIV ATED III-V
SEMICONDUCTORS
By
Jing Qiu
_________________________________________________________________
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(MATERIALS SCIENCE)
December 2015
Copyright 2015 Jing Qiu
ii
Dedication
Dedicated to my beloved parents and my husband Dr. Wangxue Deng.
iii
Acknowledgments
First and foremost, I would like to express my sincere gratitude to my academic
advisor Prof. Stephen B. Cronin for his invaluable help and guidance. I sincerely appreciate
his encouragement, hard work and dedication. Without the time he spent on me and funding
he invested in me, I could not make the achievement.
My sincere thanks also go to my dissertation committees Prof. Edward Goo, Prof.
Jongseung Yoon and Prof. Wei Wu for their invaluable suggestions, as well as Prof. Andrew
Armani for serving on my qualifying exam committee.
I extend my gratitude to our collaborators Dr. Matthew Mecklenburg in the Center
for Electron Microscopy and Microanalysis in USC and Dr. Mingyuan Ge for helpful
suggestions on difficulties in research project. Also, I want to thank Dr. Kian Kaviani for
offering me a teaching assistant position for his class. The 4 years I spent teaching for the
class were memorable. The professors and staff in the Electrical Engineering department
were extremely helpful. I appreciate all the work Dr. Donghai Zhu did and continues to do
for the Powell clean room at USC. Kim Reid and Shanna Mitchell were extremely helpful
in their administrative support and management of funding.
Nobody finishes a dissertation without fellow researchers. Cronin Group members
at B113, I have to thank every single one of you, I could not have done all this work here
without you. Chia-Chi Chang, Chun-Chung Chen, Guangtong Zeng, Bingya Hou, Haotian,
iv
Shi, Jesse Theiss, Mehmet Aykol, Mohammed Amer, Rohan Dhall, Shermin Arab, Shun-
Wen Chang, Wei-Hsuan "Wayne" Hung, Wenbo Hou, Zhen Li, Nirakar Poudel, Ioannis
Chatzakis and Zuwei Liu, thank you all for making this journey fun. I really enjoyed
working with you.
Finally I take pride in dedicating the current PhD thesis work to my beloved parents
and my husband Wangxue Deng who have given me endless love and encouragement. I
would not have all these achievements today without you.
1
Contents
Dedication ........................................................................................................................... ii
Acknowledgments.............................................................................................................. iii
List of Tables ....................................................................................................................... 4
List of Figures ..................................................................................................................... 5
Abstract ............................................................................................................................. 14
Chapter 1: Introduction to photocatalysis on semiconductor surfaces ............................. 18
1.1 Fundamentals ...................................................................................................... 18
1.2 Applications ........................................................................................................ 20
1.2.1 Water splitting .......................................................................................... 20
1.2.2 CO2 reduction........................................................................................... 24
1.3 Factors limiting the photocatalytic efficiency ..................................................... 28
1.3.1 Mismatching the solar spectrum .............................................................. 28
1.3.2 Photocorrosion of small band gap materials ............................................ 29
1.3.3 Light absorption/carrier diffusion length problem ................................... 31
Chapter 2: Plasmon-enhanced water splitting on TiO2-passivated GaP photocatalysts
........................................................................................................................................... 32
2.1 Introducton .......................................................................................................... 32
2.2 Experimental procedure ...................................................................................... 34
2.3 Results and discussion ........................................................................................ 36
2
2.4 Conclusion .......................................................................................................... 44
Chapter 3: Enhanced Surface Stable Solar-Driven H2 evolution on TiO2-passivated
p-GaAs .............................................................................................................................. 46
3.1 Introducton .......................................................................................................... 46
3.2 Experimental procedure ...................................................................................... 48
3.3 Results and discussion ........................................................................................ 52
3.4 Conclusion .......................................................................................................... 61
Chapter 4: Photocatalytic conversion of CO2 to hydrocarbon fuels on GaP via
plasmon enhanced absorption ........................................................................................... 62
4.1 Introducton .......................................................................................................... 62
4.2 Experimental procedure ...................................................................................... 63
4.3 Results and discussion ........................................................................................ 65
4.4 Conclusion .......................................................................................................... 75
Chapter 5: Artificial photosynthesis on TiO2-passivated InP nanopillars ......................... 76
5.1 Introducton .......................................................................................................... 76
5.2 Experimental procedure ...................................................................................... 76
5.3 Results and discussion ........................................................................................ 80
5.4 Conclusion .......................................................................................................... 90
Chapter 6: A microscopic study of atomic layer deposition of TiO2 on GaAs and its
photocatalytic application ................................................................................................. 91
3
6.1 Introducton .......................................................................................................... 91
6.2 Experimental procedure ...................................................................................... 92
6.3 Results and discussion ........................................................................................ 94
6.4 Conclusion ........................................................................................................ 103
Chapter 7: Future work ................................................................................................... 105
7.1 III-V semiconductor nanowires for CO2 reduction ........................................... 105
7.2 Experimental procedure .................................................................................... 106
Bibliography ................................................................................................................... 109
4
List of Tables
Table 3.1 H2 yields and Faradaic efficiencies of the 12, 24, and 48-hour reactions on
GaAs with 3nm TiO2 under 532nm illumination in a 0.5M H2SO4 solution
Table 4.1 Shift of onset overpotential for samples with different thicknesses of TiO2
Table 5.1 Faraday efficiencies of methanol of 12-hour reaction on InP/TiO2/Cu
sample under 532nm illumination in a CO2 saturated 0.5M KCl solution
Table 5.2 Faraday efficiencies of hydrogen and methanol of 12-hour reaction on
InP/TiO2/Cu sample under 532nm illumination in a CO2 saturated 0.5M KCl solution
Table 7.1 The resulting depths of GaAs holes with different ratio of H2SO4, H2O2 and
H2O in 3min etching
.....66
…………...….88
....89
………………………………………………………………....….108
…….….......60
5
List of Figures
Figure 1.1 Energy band diagram of various semiconductors, plotted together with the
redox potentials of some chemical reaction
Figure 1.2 Schematic diagram of a photoelectrochemical cell for photocatalytic
water splitting
Figure 1.3 Energy band diagram of various semiconductors, plotted together with the
redox potentials of water splitting
Figure 1.4 Schematic diagram of a photoelectrochemical cell for photocatalytic CO2
reduction
Figure 1.5 Energy band diagram of various semiconductors, plotted together with the
redox potentials of CO2 reduction
Figure 1.6 Solar spectrum and absorption spectrum of TiO2
Figure 1.7 Theoretical photocurrent densities for semiconductors under one-sun
illumination. The blue curve is the integrated photon flux at different cutoff energies,
based on the AM1.5G solar spectrum (ASTM G173-03 reference). The theoretical
photocurrent is calculated assuming that all incident photons above the
semiconductor band gap contribute to the photocurrent. The arrows at the bottom
indicate the regions of ultraviolet (UV) (below 400nm), visible (400−750 nm), and
infrared (IR) spectra
Figure 1.8 Absorption length / diffusion length mismatch.illustrated in GaP
...……………………………………......….19
…………………………………………………………………………....21
…………………………………………………..…...24
………………………………………………………………………………....26
…………………………………………………….....27
…………………………...29
……………………………………………………………………..30
…………...31
6
Figure 2.1 Schematic diagram of sample geometry for GaP photocatalysts with TiO2
and Au nanoparticles
Figure 2.2 (a) Photocatalytic current-potential curves measured for GaP
photocatalysts with various thicknesses of TiO2 under 1W/cm
2
532nm illumination
in a 0.5M Na2SO4 pH=7 solution. (b) Photocatalytic current plotted as a function of
potential for thicker TiO2 layers. (c) Relative decrease of the overpotential required
to initiate this reaction and (d) calculated built-in voltage plotted as a function of TiO2
thickness.Energy band diagrams for (e) thin and (f) thick TiO2 layers
Figure 2.3 Photocurrent plotted as a function of voltage for GaP photocatalysts with
various thicknesses of TiO2 with Au nanoparticles under 1W/cm
2
532nm illumination
in a 0.5M Na2SO4 solution
Figure 2.4 Electric field distributions calculated using the finite difference time
domain method (a) in the plane of the Au nanoparticles and (c,d) in the perpendicular
direction across the GaP/TiO2/Au/electrolyte interface. (b) Calculated electric field
enhancement factor plotted as a function of TiO2 thickness using Eq. 1
Figure 2.5 Two terminal photocurrent density (absolute value) plotted as a function
of the applied overpotential for various GaP/TiO2/Au nanoparticle photocatalysts
measured in a pH=0, 0.5M H2SO4 solution under 1W/cm
2
532nm illumination
………………………………………………………………….....35
……………….....37
………………………………………………………………..39
……………......39
……......42
7
Figure 2.6 (a) Time dependence of the photocurrent density of bare GaP illuminated
with 1W/cm
2
532nm light in a 0.5M Na2SO4 solution at an applied overpotential of -
0.7V . (b) Optical microscope image, (c) atomic force microscope image, and (d)
surface topography of the GaP surface after the 5 hour reaction
Figure 2.7 (a) Time dependence of the photocurrent density of TiO2passivated GaP
illuminated with 1W/cm
2
532nm light in a 0.5M Na2SO4 solution at an applied
overpotential of -0.7V . (b) Optical microscope image, (c) atomic force microscope
image, and (d) surface topography of the GaP/TiO2 surface after the 12 hour reaction.
Figure 3.1 (a) Schematic diagram of sample geometry. (b) TEM Image of 3nm TiO2
on GaAs. Note that for TEM sample preparation and imaging purposes, the surface
of the substrate is coated with a thick layer of Pt
Figure 3.2 TEM Image of 3nm native oxide on GaAs. Note that for TEM sample
preparation and imaging purposes, the surface of the substrate is coated with a thick
layer of Pt
Figure 3.3 (a) Ti L edge map and (b) O K edge map from electron energy-loss
spectroscopy. (c) Profile of Ti L edge (blue background) and O K edge map (red line).
(d) STEM image of 3nm TiO2 on GaAs sample. Note that for STEM sample
preparation and imaging purposes, the surface of the substrate and TiO2 is coated with
a thick layer of e-beam deposited Pt. (The microscope used was a JEOL JEM-2100F
at 200kV high tension)
…...…………………...43
....44
………………………………………...48
……………………………………………………………………………......49
…………………………………………………………………...51
8
Figure 3.4 (a) Photocatalytic current-potential curves and (b) corresponding photon-
energy conversion efficiency calculated by using Equation 1 for GaAs photocatalysts
with various thicknesses of TiO2 under an AM1.5 G illumination in a 0.5M H2SO4
pH=0 solution
Figure 3.5 Photocurrent plotted as a function of voltage for GaAs photocatalysts with
10nm and 15nm TiO2 under 532nm illumination (910 mW/cm
2
) in a 0.5M H2SO4
solution
Figure 3.6 (a) Ti 2p and (b) O 1s core level XPS spectra of various thickness of TiO2
on GaAs
Figure 3.7 Dark current plotted as a function of voltage for GaAs photocatalysts with
and without TiO2 thin layer in a 0.5M H2SO4 solution
Figure 3.8 Ti 2p core level XPS spectra of 3nm TiO2 on GaAs after H2 evolution
reaction
Figure 3.9 Photoluminescence spectra of GaAs photocatalysts with various
thicknesses of TiO2 under 532nm excitation
Figure 3.10 Time dependence of the photocurrent density of (a) bare GaAs and 3nm
TiO2 passivated GaAs illuminated for 12 hours and (b) 3nm TiO2 passivated GaAs
illuminated for 48 hours with 532nm light at an applied overpotential of 0.1V vs.
RHE in a 0.5M H2SO4 solution
…………………………………………………………………………....52
…………………………………………………………………………………..53
……………………………………………………………………………….....54
……………………………….....56
..............................................................................................................................56
………………………………………….....57
……………………………………………………….....58
9
Figure 3.11 (a) Gas chromatograph (GC) data taken after 12, 24, and 48-hour
reactions on GaAs with 3nm TiO2 under 532nm illumination in a 0.5M H2SO4
solution. (b) H2 yield plotted as a function of time
Figure 4.1 Schematic diagram of sample geometry
Figure 4.2 (a) Optical microscope image, (b) atomic force microscope image, and (c)
surface topography of bare GaP surface after 8h reaction at -0.5V overpotential. (d)
Optical microscope image, (e) atomic force microscope image, and (f) surface
topography of 5nm TiO2 on GaP surface after 8h reaction
Figure 4.3 (a) Photocatalytic current-potential curves of GaP photocatalysts with
different TiO2 thicknesses in a 0.5M NaCl, 10mM pyridine solution under 532nm
wavelength laser illumination. (b) Decrease of overpotential plotted as a function of
TiO2 thickness on GaP. (c) Calculated built-in voltage plotted as a function of TiO2
thickness. (d) NMR spectra showing methanol production using bare GaP and 5nm
TiO2-passiavated GaP photocatalysts at an overpotential of -0.50V
Figure 4.4 Gas chromatograph (GC) data taken after 8 hour illumination (532nm) of
bare and 5nm TiO2 passivated-GaP in 0.5M NaCl, 10mM pyridine solution at an
overpotential of -0.5V (vs NHE). The GC data is plotted together with a calibration
standard consisting of 10
-4
M methanol in aqueous solution. Based on this data, the
5nm TiO2 passivated-GaP was found to have a Faradaic efficiency of 55%
…………...………………………....59
……………………………………..65
……………………………...65
………………….....67
…………....70
10
Figure 4.5
13
C-NMR spectrum showing methanol peak taken after 8 hour
illumination (532nm) of 5nm TiO2 passivated-GaP at an overpotential of -0.5V (vs
NHE) in 0.5M NaCl, with 10mM pyridine solution
Figure 4.6
1
H-NMR spectra showing methanol peaks taken after 8 hour illumination
(532nm) of 5nm TiO2 passivated-GaP at an overpotential of -0.5V (vs NHE) in 0.5M
NaCl, with 10mM pyridine solution and without 10mM pyridine
Figure 4.7
(a) Optical microscope image, (b) atomic force microscope image, and (c)
surface topography of TiO2 passivated GaP surface after 8h reaction in 0.5M NaCl
without pyridine at -0.5V overpotential
Figure 4.8 Energy band alignment of GaP and TiO2 together with the relevant redox
potentials of CO2
Figure 5.1 (a) Schematic diagram of TiO2-passivated InP nanopillars with Cu
cocatalyst nanoparticles. (b-e) SEM and TEM images of InP nanopillar array with
TiO2 deposition layer and Cu nanoparticles. The high resolution TEM image in (e)
resolves the crystal lattice of the Cu nanoparticles
Figure 5.2 (a) Cu 2p for CuO and Cu, (b) O 1s core level XPS spectra of air-oxide
Cu nanoparticles
Figure 5.3 Time dependence of the photocurrent density of InP with 3nm TiO2 and
Cu nanoparticles illuminated with 532nm light at an applied overpotential of -0.6V
vs. NHE
……………………………………..71
……………………....71
………………………………………………....72
………………………………………………………………………....73
……………………………………...78
………………………………………………………………………....79
……………………………………………………………………………….....80
11
Figure 5.4 (a) Photocatalytic current-potential curves where the dashed line indicates
the potential applied during the methanol test. (b) Log plot of photocatalytic current-
potential curves. (c) Methanol peak in NMR spectra and (d) Faraday efficiencies of
methanol production for InP nanopillars with and without 3nm TiO2 under 532nm
illumination in a CO2 saturated 0.5M KCl solution
Figure 5.5 Ti 2p level XPS spectra of TiO2 on InP, which shows the presence of Ti
3+
states
Figure 5.6 PW-DFT calculated structure for anatase TiO2 with O vacancies (a) before
CO2 adsorption and (b) after CO2 adsorption and relaxation
Figure 5.7 PW-DFT result of neutral CO2 adsorbed to the stoichiometric anatase 101
surface with a binding energy of -0.48 eV
Figure 5.8 (a) Photocatalytic current-potential curves and (b) Faraday efficiencies of
methanol production for samples of bare InP, InP/Cu, and InP/TiO2/Cu under 532nm
illumination in a CO2 saturated 0.5M KCl solution for a12-hour reaction
Figure 5.9 Gas chromatograph (GC) data taken after 12-hour reaction on
InP/TiO2/Cu sample under 532nm illumination in a CO2 saturated 0.5M KCl solution.
Figure 6.1 TEM Image of 15nm anatase TiO2 crystalline on GaAs
…………………………………......82
…………………………………………………………………………………......82
…………………………....84
…………………………………………….....84
……………...87
....89
………………….....95
12
Figure 6.2 The high resolution TEM image of TiO2 films deposited with (a) 25 cycles,
(b) 75 cycles, and (c) 500 cycles of atomic layer deposition on GaAs. Note that for
TEM sample preparation and imaging purposes, the surface of the substrate is coated
with a thick layer of Pt. (d-e) EELS spatial profiles of Ti L edge (green line) and O K
edge map (red line) for the 25, 75, and 500 cycle TiO2 on GaAs samples, respectively.
Figure 6.3 Photocatalytic water splitting current-potential curves for samples of (a)
GaAs photocatalysts with various thicknesses of TiO2 and (b) GaAs with 10nm TiO2
before and after annealing under 532nm illumination in a 0.5M H2SO4 pH=0 solution.
Figure 6.4 Photocatalytic CO2 reduction current-potential curves for samples of (a)
GaAs photocatalysts with various thicknesses of TiO2 and (b) GaAs with 10nm TiO2
before and after annealing under 532nm illumination in a 0.02M [EMIM]BF4 non-
aqueous CO2 saturated electrolyte
Figure 6.5 Gas chromatograph (GC) data taken for reactions on GaAs with 3nm TiO2
under 532nm illumination (a) in a 0.5M H2SO4 solution and (b) in a 0.02M
[EMIM]BF4 non-aqueous CO2 saturated electrolyte
Figure 6.6 H2O adsorption on stoichiometric anatase with an adsorption energy of
-1.26 eV
Figure 6.7 (a) PW-DFT calculated structure for anatase TiO2 with O vacancies before
adsorption, (b) after H2O adsorption and relaxation and (c) after CO2 adsorption and
relaxation
....96
....97
…………………………………………………….....98
………………………………….....98
………………………………………………………………………………...102
.........................................................................................................................103
13
Figure 7.1 SEM images of high aspect ratio GaAs nanopillars produced from a 600
nm wide square Au mesh pattern in H2SO4 and KMnO4 solution at 40-45
o
C. (a) 30
o
tilted view at low magnification, (b) 30
o
tilted view at high magnification, (c) cross-
sectional showing the highly vertical nanopillar array
Figure 7.2 (a) Mask design with hole diameter of 300nm and center to center spacing
900nm. (b) Etching depth vs. etching time for GaAs with H2SO4, H2O2 and H2O
(1:8:8)
………………………………....106
…………………………………………………………………………………...107
14
Abstract
In photocatalysis, light absorbed by semiconductors creates electron-hole pairs
which can be used to drive electrochemical redox reactions. Since the first demonstration
of photocatalytic water splitting using TiO2 in 1972, the study of photocatalysis has been
of worldwide interest due to its potential applications in solar fuel generation either through
water splitting or CO2 reduction. For photocatalytic materials, III-V semiconductors such
as GaP, InP, and GaAs are promising candidates with theoretical maximum photocurrent
densities of 9mA/cm
2
, 35mA/cm
2
, and 32mA/cm
2
, respectively. These are significantly
larger than that of the mostly widely studied photocatalytic material, TiO2 (1.1mA/cm
2
).
However, photocatlytic corrosion of these III-V semiconductors prevents them from being
utilized as photocatalysts. We have developed a technology of passivating the surface of
these III-V semiconductors with a thin layer of TiO2, which protects them from corrosion.
In addition, this thin layer of TiO2 provides substantial enhancement in the overall
photoconversion efficiency.
The dissertation will start with a brief introduction of the fundamentals of
photocatalysis on semiconductor surfaces, followed by the basic mechanisms of its
applications to water splitting and CO2 reduction. After that, the factors limiting the
photocatalytic conversion efficiency are introduced. In the following chapters, we will
report the achievements we have made both on the enhanced photocatalytic water splitting
and CO2 reduction processes.
15
In Chapter 2, we demonstrate that a thin layer of n-type TiO2 using atomic layer
deposition (ALD) prevents corrosion of p-type GaP, as evidenced by atomic force
microscopy and photoelectrochemical measurements. In addition, the TiO2 passivation
layer provides an enhancement in photoconversion efficiency through the formation of a
charge separating pn-region. Plasmonic Au nanoparticles deposited on top of the TiO2-
passivated GaP further increases the photoconversion efficiency through local field
enhancement. Finite difference time domain (FDTD) simulations of the electric field
profiles in this photocatalytic heterostructure corroborate the experimental results.
In Chapter 3, we demonstrate that thin (1-5nm) films of TiO2 deposited by ALD on
planar GaAs provide electrochemical stability and substantial improvements in the
efficiency of photocatalytic water splitting. The native oxide of GaAs is removed during
the ALD process and this TiO2 passivation layer produces a shift in the onset potential by
+0.4V and enhances the photocurrent by 32-fold over bare GaAs (at 0V vs. RHE), resulting
in a peak photoconversion efficiency of 1.5% under AM1.5 G illumination. The possible
enhancement mechanism by Ti
3+
active sites provided by ALD TiO2 is discussed.
In Chapter 4, photocatalytic CO2 reduction with water to produce methanol is
demonstrated using TiO2-passivated GaP photocathodes under 532nm wavelength
illumination. In addition to providing a stable photocatalytic surface, the TiO2-passivation
provides substantial enhancement in the photoconversion efficiency. Two possible
mechanisms of enhancement are discussed. One is the passivation of surface states, which
16
cause non-radiative carrier recombination, and the other is pn-junction formation from the
n-type TiO2/p-type GaP junction, which creates a built-in field that assists in the separation
of photogenerated electron-hole pairs, further reducing recombination.
In Chapter 5, photocatalytic CO2 reduction with water to produce methanol is
demonstrated using TiO2-passivated InP nanopillar photocathodes under 532nm
wavelength illumination. In addition to providing a stable photocatalytic surface, the TiO2-
passivation layer provides substantial enhancement in the photoconversion efficiency
through the introduction of O vacancies associated with the non-stoichiometric growth of
TiO2 by atomic layer deposition. Plane wave-density functional theory (PW-DFT)
calculations confirm the role of oxygen vacancies in the TiO2 surface, which serve as
catalytically active sites in the CO2 reduction process. PW-DFT shows that CO2 binds
stably to these oxygen vacancies and CO2 gains an electron (-0.897e) spontaneously from
the TiO2 support. This calculation indicates that the O vacancies provide active sites for
CO2 absorption, and no overpotential is required to form the CO2
-
intermediate.
In Chapter 6, we report a microscopic study of GaAs/TiO2 heterojunctions using
cross-sectional high resolution transmission electron microscopy and electron energy loss
spectroscopy maps. The photocatalytic performance of these heterostructures shows a very
strong dependence on the thickness of the TiO2 over the range of 0-15nm. Thinner TiO2
films (<10nm) are amorphous and show enhanced catalytic performance with respect to
bare GaAs. Thicker TiO2 films (15nm) are crystalline and have poor charge transfer due to
17
their insulating nature, while thinner amorphous TiO2 films are conducting. PW-DFT
calculations show that water molecules and CO2 molecules bind stably to defective TiO2,
which can further improve the photocatalytic charge transfer process.
Finally, in Chapter 7, future work related to photocatalysis using TiO2-passivated
nanowires of III-V semiconductors is discussed.
18
Chapter 1: Introduction to photocatalysis on semiconductor
surfaces
1.1 Fundamentals
From the energy-band theory of crystalline materials, electrons may occupy the
bands of allowed energy states that are separated by bands of forbidden energies. At
absolute zero Kelvin, electrons occupy the lowest energy states. For some crystalline
materials, all states in the lower band (the valence band) will be full, and all states in the
upper band (the conduction band) will be empty. The width of the forbidden energy band
between the top-most valence band and the bottom-most conduction band is the band gap
energy (Eg).
As we know, in the wave-particle duality principle, light waves can be treated as
particles, which are referred to as photons. The energy of a photon is given by E=hv, where
h is Planck’s constant and v is the light frequency. We can relate the wavelength and the
energy by λ=1240/E (nm). When a photon collides with a valence electron in a solid
(usually a semiconductor), if the energy of the photon is equal to or larger than the band
gap energy of the semiconductor, the electron absorb enough energy to be excited into the
conduction band. Such a process generates electron-hole pairs. These generated electron-
hole pairs in the semiconductor can transfer their charger and energy to the reactants in an
electrolyte solution, and cause the photo-assisted decomposition or synthesis of chemical
19
compounds.
In Figure 1.1, we plot the band gap and band edge positions of some
semiconductors with respect to various redox potentials. This data is derived from flat band
potential determinations via capacity measurements. The standard potentials for these
redox couples indicate the thermodynamic limitations for the photoreactions that can be
carried out with the charge carriers. For example, if a reduction of the species in the
electrolyte is to be performed, the conduction band position of the semiconductor must be
positioned above the relevant reduction level. For the oxidation of the species, the valence
band position must be under the oxidation potential of that species.
It has been suggested that at semiconductor electrodes, the charge transfer rates
between photo-generated carriers in semiconductors and the solution species depend on the
correlation of energy levels between the semiconductor and the redox agents in the solution.
Figure 1.1 Energy band diagram of various semiconductors, plotted together with the redox
potentials of some chemical reactions.
-8
-7
-6
-5
-4
-3
InP
H
2
O/H
2
0V
H
2
O/O
2
0.82 V
CO
2
/CO
2
-
-1.90 V
1.7 eV
anatase rutile
2.8 eV
3.2 eV
1.34 eV
2.25 eV
1.1 eV
2.8 eV
2.3 eV
2.4 eV
1.6 eV
3.2 eV
TiO
2
TiO
2
CuO
Cu
2
O
Fe
2
O
3
WO
3
Si
GaP
CdSe
SrTiO
3
PbO
3.0 eV
Vacuum (eV)
NHE (Volts)
CO
2
/CH
3
COOH-0.61V
CO
2
/CH
3
OH -0.38 V
CO
2
/CH
4
-0.244 V
3
2
1
0
-1
20
If the conduction band position of the semiconductor is not positioned above the relevant
redox level, the photo-excited electrons do not have the ability to drive the reductive half-
reaction. In this situation, an external bias voltage must be applied in order to drive the
reaction, which is commonly referred to as applying an overpotential.
1.2 Applications
1.2.1 Water splitting
One of the most important applications of semiconductor photocatalysis in solar
energy conversion is photocatalytic water splitting, that is, the production of hydrogen (H2)
and oxygen (O2) from water by the direct conversion of solar energy. With hydrocarbon
fuels and other nonrenewable fuels becoming increasingly depleted and expensive,
hydrogen fuel presents a promising alternative, which burns cleanly without environmental
pollution, such as emission of greenhouse gases.
Photocatalytic water splitting has been of great interest since the early 1970s, after
the first demonstration by Fujishima and Honda under ultraviolet radiation
3
. This reaction
can be carried out in a photoelectrochemical cell, as shown in Figure 1.2. When the
semiconductor photoelectode is in contact with an electrolyte solution, a space charge layer
is formed inside the semiconductor electrode in order to establish the thermodynamic
equilibration with the ions in solution. When the semiconductor photoelectrode is
irradiated with the light above the band gap, electron-hole pairs are generated and separated
21
in the space charge layer. The photo-generated electrons move through the external circuit
and hydrogen is formed at the surface of the working electrode, while the photo-generated
holes move toward the interface of the counter electrode and the electrolyte and oxygen is
formed
4
.
Figure 1.2 Schematic diagram of a photoelectrochemical cell for photocatalytic water splitting.
Photocathode Pt
Gas out
to GC
A
V
Argon
in
light
ref.
H
2
O
O
2
H
+
e
+
e
-
H
2
H
+
Potentialstat
22
The standard reduction potentials (Ered) of hydrogen evolution and the standard
oxidation potentials (Eox) of oxygen evolution are as follows:
136
Reduction: 4H2O + 4e
-
2H2(g) + 4OH
-
Ered=-0.83V (1.1)
Oxidation: 2H2O + 4h
+
O2(g) + 4H
+
Eox=1.23V (1.2)
Standard reduction-oxidation potentials are measured under standard conditions of 25 ℃,
1M concentration for each ion participating in the reaction, a partial pressure of 1 atm for
each gas that is part of the reaction, and metals in their pure state.
136
These standard potentials given in Equation 1.1 and 1.2, assume an anode chamber
with 1M H
+
and a cathode chamber with 1M OH
-
. However, the actual reduction potentials
are related to the actual concentrations of H
+
and OH
-
ions, or pH. The actual reduction-
oxidation potentials at a specific pH can be calculated using the Nernst Equation 1.3
E=E
o
(-RT/nF ·ln(Cox/Cred) ) (1.3)
where E is the actual reduction-oxidation potential, E
o
is the standard reduction potential,
R is the ideal gas constant, T is the absolute temperature, F is the Faraday constant, n is the
number of electrons needed in the reaction, and Cox and Cred are concentrations of oxidation
and reduction species, respectively. For water splitting reaction, the Nernst Equation can
23
be simplified into the following relation, since only H
+
and OH
-
ions are involved in the
reaction:
E=E
o
± 0.059pH (1.4)
Where “+” is used for the reduction half reaction and “-” is used for oxidation half reaction.
Therefore, the actual reduction potential can be adjusted by pH. For example, the actual
reduction potential of hydrogen evolution is -0.41V vs. NHE and the oxidation potential of
oxygen evolution is 0.82V vs. NHE in pure water, where [H
+
] = [OH
-
] = 10
-7
M. Figure 1.3
shows the reduction-oxidation potentials at various pH values. At pH=0, the reduction
potential needed to drive the reduction of hydrogen from water is 0.0V vs. NHE, and
oxidation potential of oxygen from water is 1.23V vs. NHE. At pH=14, the reduction
potential needed for hydrogen generation is -0.83V vs. NHE and for oxygen formation is
0.4V vs. NHE by Equation 1.4.
By comparing these redox potentials with the energies of the valence and
conduction bands of several semiconductors, as indicated in Figure 1.3, we can see that a
few semiconductors have both the conduction band energy lying higher than the hydrogen
evolution potential and the valence band lying below the oxygen evolution potential, such
as GaP, InP, and CdSe, which can theoretically drive photocatalytic water splitting by light
without an applied overpotential. For those materials with band positions that mismatch
24
the redox potentials of water splitting, an external applied potential is needed in order to
initiate the reaction. To date, most reports of hydrogen evolution from photocatalytic water
splitting is achieved under an additional externally applied voltage.
1.2.2 CO2 reduction
Another important application of semiconductor photocatalysis in the area of solar
energy conversion is the photocatalytic reduction of carbon dioxide with water to form
hydrocarbon fuels. It is well-established that higher atmospheric carbon dioxide (CO2)
levels result in an atmosphere that better retains heat. The significant rise in atmospheric
Figure 1.3 Energy band diagram of various semiconductors, plotted together with the redox
potentials of water splitting.
-8
-7
-6
-5
-4
-3
pH=7
pH=14
pH=0
GaP
pH=7
pH=14
pH=0
1.7 eV
anatase rutile
H
2
O/O
2
2.8 eV
3.2 eV
1.34 eV
2.25 eV
1.1 eV
2.8 eV
2.3 eV
2.4 eV
1.6 eV
3.2 eV
TiO
2
TiO
2
CuO
Cu
2
O
Fe
2
O
3
WO
3
Si
InP
CdSe
SrTiO
3
PbO
3.0 eV
Vacuum (eV)
NHE (Volts)
H
2
O/H
2
3
2
1
0
-1
25
CO2 levels resulting from the combustion of hydrocarbon fuels in the past several decades,
causes the global warming. A solar energy based technology to recycle CO2 into readily
transportable hydrocarbon fuels would be a great help to reduce atmospheric CO2 levels
and to partly fulfill energy demands within the present hydrocarbon based fuel
infrastructure
5
.
One of the earliest reports of photoelectrochemical reduction of CO2 was published
by Halmann in 1987
6
. An electrochemical cell was used, comprised of a single crystal p-
type GaP cathode, carbon anode, and a buffered aqueous electrolyte through which CO2
was bubbled, as illustrated in Figure 1.4. When the GaP crystal was illuminated (mercury
lamp) and a voltage bias applied, current was detected, with the electrolyte solution
showing the presence of formic acid, formaldehyde, and methanol. Many research groups
5
have since investigated the use of different compound semiconductors to achieve higher
visible light catalytic activities in this CO2 reductive reaction system. The mechanism is
almost the same as for water splitting. When the semiconductor photoelectrode is
illuminated under the light above the band gap, the generated electrons move to the surface
of the photocathode to convert CO2 into hydrocarbon with water.
26
The reduction potentials for the reduction half reactions in an aqueous solution of
pH 7 are as follows, assuming the carbon dioxide is reduced directly to give products of
interest:
137
CO2(g) + 8H
+
+ 8e
-
CH4(g) + 2H2O E
o
=-0.24V (1.5)
CO2(g) + 4H
+
+ 4e
-
HCHO(aq) + H2O E
o
=-0.48V (1.6)
CO2(g) + 6H
+
+ 6e
-
CH3OH(aq) +H2O E
o
=-0.38 V (1.7)
CO2(g ) +2H
+
+ 2e
-
HCOOH (aq) E
o
=-0.61V (1.8)
CO2(g) + 2H
+
+ 2e
-
CO(g) +H2O E
o
=-0.52V (1.9)
2CO2(g) + 2H
+
+ 2e
-
H2C2O4 (aq) E
o
=-0.9V (1.10)
Figure 1.4 Schematic diagram of a photoelectrochemical cell for photocatalytic CO 2 reduction.
Photocathode Pt
Gas out
to GC
A
V
CO
2
in
light
ref.
H
2
O
O
2
H
+
e
+
e
-
H
+
Potentialstat
CO
2
CH
3
OH
H
2
O
27
Where the standard redox potentials are given in V olts versus NHE at pH 7.0, and (g) and
(aq) denote the gaseous state and aqueous solution, respectively.
The relationship of the positions of the conduction bands of some semiconductors
and the redox potentials related with CO2 reduction are shown in Figure 1.5. When the
conduction band edges of the semiconductor electrodes are more negative than the
reduction potentials of CO2, the electrons are energetically favorable to transfer from the
conduction band of the semiconductor to the CO2 molecule. It is believed that in direct CO2
reduction, the first electron transfers to CO2 as follows:
Figure 1.5 Energy band diagram of various semiconductors, plotted together with the redox
potentials of CO 2 reduction.
-8
-7
-6
-5
-4
-3
InP
H
2
O/O
2
0.82 V
CO
2
/CO
2
-
-1.90 V
1.7 eV
anatase rutile
2.8 eV
3.2 eV
1.34 eV
2.25 eV
1.1 eV
2.8 eV
2.3 eV
2.4 eV
1.6 eV
3.2 eV
TiO
2
TiO
2
CuO
Cu
2
O
Fe
2
O
3
WO
3
Si
GaP
CdSe
SrTiO
3
PbO
3.0 eV
Vacuum (eV)
NHE (Volts)
CO
2
/CH
3
COOH-0.61V
CO
2
/CH
3
OH -0.38 V
CO
2
/CH
4
-0.244 V
3
2
1
0
-1
28
CO2 + e- CO2
-
(1.11)
This is intermediate step requires a potential of -1.9V versus NHE in order to initiate charge
transfer. Unfortunately, few semiconductors have their conduction band edge lying above
this potential. Thus, an applied potential is usually needed to enable CO2 reduction reaction
to occur.
1.3 Factors limiting the photocatalytic efficiency
1.3.1 Mismatching the solar spectrum
The absorption of sunlight and its conversion to a storable fuel motivates the
investigation about photocatlysis. However, semiconductor candidates with relative large
band gaps (TiO2, WO3, SrTiO3 etc.) can not be used as photoelectrodes for efficient
photocatalytic reactions due to the inherent mismatch between the absorption spectra of
semiconductors and the solar spectrum.
For example, TiO2, as a self-cleaning catalyst, does not suffer from the corrosion
problems associated with photovoltaic cells, and thus it has been the most studied
photocatalyst to date. However, it does not absorb light in the visible region of the
electromagnetic spectrum. Figure 1.6 shows this problem graphically with the absorption
spectrum of TiO2 superimposed over the solar spectrum (AM1.5).
138
Because of TiO2’s
short wavelength cutoff, there are very few solar photons (~4%) that can be used to drive
29
this photocatalyst. Several attempts have been made previously to extend the cutoff
wavelength of this catalyst, including doping
7, 8
, and defect creation
9
. While these efforts
have resulted in slight improvements in the absorption in the visible range, leaving a
majority of the solar spectrum unable to drive this photocatalyst
7, 8, 10
.
1.3.2 Photocorrosion of small band gap materials
Taking into account overpotentials and other losses, the most desirable band gaps
for photoelectrodes are between 1.1 and 1.7eV for optimized efficiency from theoretical
calculations.
11, 12
Further, with examining several typical semiconductors for photocatalytic
applications, the choice of available materials with band gaps ranging between 1.5 and
2.0eV is limited.
1
Figure 1.7 shows the theoretical photocurrent density under one-sun
Figure 1.6 Solar spectrum and absorption spectrum of TiO 2.
30
illumination. According to this figure, Si (44mA/cm
2
) and III-V semicondutors, such as
GaP (9mA/cm
2
), InP (35mA/cm
2
) and GaAs (32mA/cm
2
), are better candidates than TiO2
(1~2mA/cm
2
). Unfortunately, most semiconductors in the range between 1.5 and 2.0eV
corrode readily when incorporated in an aqueous photoelectrochemical cell.
13, 14
Such
corrosion is not expected in practical use for photocatalytic reaction.
Figure 1.7 Theoretical photocurrent densities for semiconductors under one-sun illumination. The
blue curve is the integrated photon flux at different cutoff energies, based on the AM1.5G solar
spectrum (ASTM G173-03 reference). The theoretical photocurrent is calculated assuming that all
incident photons above the semiconductor band gap contribute to the photocurrent. The arrows at
the bottom indicate the regions of ultraviolet (UV) (below 400nm), visible (400−750 nm), and
infrared (IR) spectra.
1
31
1.3.3 Light absorption/carrier diffusion length problem
Another important problem in photocatalysts is the inherent mismatch between the
light absorption length and the minority carrier diffusion length. As an example, Figure 1.8
illustrates this problem schematically for p-type GaP, where about 100μm of GaP film is
needed to absorb 63% of the incident light at λ=532nm
15
. Thus, a majority of the photo-
generated carriers are produced within a distance of 100μm from the
semiconductor/electrolyte interface. However, the minority carriers (electrons) recombine
rapidly with the majority carriers (holes) over a length scale of 100nm. Because of this
short minority carrier diffusion length, most of the photo-generated carriers (~98%)
recombine before reaching the semiconductor-electrolyte interface and, therefore, do not
contribute to the photocatalysis.
15, 16
32
Chapter 2: Plasmon-enhanced water splitting on TiO2-passivated
GaP photocatalysts
2.1 Introducton
In photocatalysis, the energy of photons can be utilized to drive many important
chemical reactions, including H2 production
17
, CO2 reduction
18-20
, and water purification
21,
22
. Unfortunately, the efficiencies of most photocatalytic processes are far too low for
practical large scale applications. The direct conversion of solar-to-chemical energy has
several advantages over solar-to-electric energy conversion, most notably, the ability to
store large amounts of energy (~GW) in chemical bonds that can later be released without
producing harmful byproducts. As the cost of direct solar-to-electrical energy becomes
competitive with fossil fuels, there will be a need to store large amounts of the solar energy
for use during nights, cloudy days, and winter months. Therefore, our energy infrastructure
has much to gain by finding more efficient ways to enhance these photocatalytic processes.
Over the past few years, several research groups have demonstrated a new method for
improving the efficiency of photocatalytic processes by exploiting the strong plasmon
resonance of small metal nanoparticles
23-26
. While these studies have clearly shown proof-
of-principle of this enhancement mechanism, the overall photoconversion efficiencies are
still very low
24
. Torimoto et al. also demonstrated enhanced photocatalytic water splitting
by depositing CdS nanoparticles on SiO2-coated Au nanoparticles
27
. However, these
33
plasmonic-photocatalytic complexes demonstrated enhancement factors less than 2. Liu et
al.and Ingram et al. observed enhanced photocatalytic water splitting under visible
illumination by depositing plasmonic metal nanoparticles on top of anatase TiO2
27-32
. While
these previous works reported enhancement factors of approximately 10-fold with the
incorporation of plasmonic nanoparticles, the overall photoconversion efficiencies were
still quite low in the visible wavelength range because of TiO2’s large band gap (Eg=3.2eV).
As a result, these proof-of-principle studies relied on short-lived, sub-band gap defect states
for optical absorption in the visible wavelength range.
GaP(Eg=2.25eV) has a substantially smaller band gap than TiO2, can absorb more
than 18% of the solar spectrum, and is better matched to the plasmon resonance energy of
Au nanoparticles. While GaP absorbs a more extended range of the solar spectrum than
TiO2, its main advantage as a photocatalyst is its relatively high conduction band energy,
which exceeds that of TiO2 by more than 1eV . However, GaP is more expensive than TiO2
and is not oxidatively stable over a wide range of pH. As a result, a vast majority of previous
photocatalytic studies have been carried out on TiO2. In the work presented here, we
investigate the photocatalytic stability of TiO2-passivated GaP using atomic force
microscopy (AFM), photoelectrochemistry, and optical microscopy. The photocatalytic
efficiency is studied systematically as a function of TiO2 layer thickness. Further catalytic
enhancement is explored with the addition of plasmonic gold nanoparticles, and is studied
as a function of TiO2 layer thickness. The effects of plasmonic enhancementare
34
distinguished from the natural catalytic properties of Au by evaluating similar
photocatalyticTiO2/GaP structures with catalytic, non-plasmonic metals (i.e., Pt) instead of
Au.
2.2 Experimental procedure
In the work presented here, Zn doped p-type (100) oriented GaP with a dopant
concentration of 2x10
18
cm
-3
was used as a photocatalyst for water splitting. Atomic layer
deposition (ALD) of TiO2 was performed at 250
o
C on the p-GaP wafers with TiCl4 as the
titanium source and water vapor as the oxygen source. The carrier gas during the deposition
was argon with a flow rate of 20sccm. The rate of deposition was about 0.4Å per cycle. A
500nm thick aluminum film was evaporated on the back of the p-GaP to form an Ohmic
contact. We then evaporated a gold film with a nominal thickness of 5nm on the top surface
of the TiO2. This thin gold film is known to form island-like growth that is strongly
plasmonic and serves as a good substrate for surface enhanced Raman spectroscopy
(SERS)
33, 34
and photocatalytic enhancement
28, 35
. Samples were prepared with and without
the plasmonic gold nanoparticles. The schematic diagram of sample geometry is shown in
the schematic diagrams of Figure 2.1. The aluminum contact was then connected to the
external circuitry with a copper wire and coated with epoxy cement to insulate it from the
electrolytic solution. The photocatalytic reaction rates of two sets of samples were
measured in 0.5 M Na2SO4 and 0.5M H2SO4 solutions using a three-terminal potentiostat
35
with the prepared samples, a Ag/AgCl electrode, and a graphite electrode functioning as
the working, reference, and counter electrodes, respectively.
Finite difference time domain (FDTD) simulations were carried out using a cell of
size 1000 nm x 800 nm x 500 nm. A grid spacing of 2Å is used in the volume of 500 nm x
500 nm x 40 nm around the film and 10 nm elsewhere. A 0.002 fs temporal grid is used
with a total of 100 000 time steps. In the simulation, the samples are irradiated with a
planewave source with a Gaussian pulse containing a spectrum of wavelengths ranging
from 300 nm to 800 nm. Perfectly matched layers (PML) boundary conditions are used
with 25 layers. The dielectric function of Au is based on the optical constants given by
Palik and Ghosh.
139
Figure 2.1 Schematic diagram of sample geometry for GaP photocatalysts with TiO 2 and Au
nanoparticles.
36
2.3 Results and discussion
Figure 2.2a shows the photocurrent-voltage curves for GaP passivated with various
thicknesses of TiO2 measured in a pH=7 solution of 0.5M Na2SO4 under 532nm
illumination. Bare GaP (blue curve) has an onset of photocurrent at a potential of
approximately -0.66V . For TiO2 passivated GaP, we see a clear shift in the overpotential
that scales linearly with the thickness of the TiO2, as shown in Figure 2.2c. For 10nm TiO2
(pink curve), this overpotential/onset potential is shifted by approximately 0.46V . This shift
is attributed to the formation of a pn-junction, since the TiO2 is n-type doped due to oxygen
vacancies
363531314343
. While TiO2 does not absorb light at 532nm, the pn-junction formed
with the GaP enables separation of photogenerated charge in the actively absorbing GaP.
Figure 2.2d shows the built-in potential for the junction calculated using the relation
=
(
)
(
)
, with a doping concentration of Na=5x10
18
cm
3
. Here,
is the
depletion width of the GaP-TiO2 junction, which is a function of the TiO2 thickness layer.
This calculation shows a similar trend to the experimentally observed shift in the
overpotential. Typically, p-type GaP acts as a photocathode due to the direction of band
bending at its interface with water, promoting photoelectron flow to the water. For n-type
semiconductors, the bending is usually upward, resulting in photohole flow to the water. It
is, therefore, somewhat surprising that when n-type TiO2 up to 10 nm thick is coated on
top of p-GaP, the electrode still functions as a photocathode. This is most likely due to the
finite thickness of the TiO2 layer, which is fully depleted and does not provide enough
37
Figure 2.2 Photocatalytic current-potential curves measured for GaP photocatalysts with various
thicknesses of TiO 2 under 1W/cm
2
532nm illumination in a 0.5M Na 2SO 4 pH=7 solution. (c)
Relative decrease of the overpotential required to initiate this reaction and (d) calculated built-in
voltage plotted as a function of TiO 2 thickness.Energy band diagrams for (e) thin and (f) thick TiO 2
layers. (g) Photocatalytic current plotted as a function of potential for thicker TiO 2 layers.
(e)
-1.5 -1.0 -0.5 0.0
-3.5
-3.0
-2.5
-2.0
-1.5
-1.0
-0.5
0.0
Dark
Bare GaP
GaP w/1nm TiO
2
GaP w/3nm TiO
2
GaP w/5nm TiO
2
GaP w/10nm TiO
2
Potential vs. Ag/AgCl (V)
Current Density (mA/cm
2
)
RHE=0
-1.5 -1.0 -0.5 0.0
-1.4
-1.2
-1.0
-0.8
-0.6
-0.4
-0.2
0.0
Dark
Bare GaP
GaP w/15nm TiO2
GaP w/20nm TiO2
GaP w/25nm TiO2
Potential vs. Ag/AgCl (V)
Current density (mA/cm
2
)
0 1 2 3 4 5 6 7 8 9 10
0.0
0.1
0.2
0.3
0.4
0.5
Thickness of TiO
2
(nm)
Decrease of Overpotential (V)
1 2 3 4 5 6 7 8 9 10
0.0
0.1
0.2
0.3
0.4
0.5
Built-in Voltage (eV)
TiO
2
Thickness (nm)
(a) (b)
(c)
(d)
(f)
38
surface charge to invert the material from p- to n-type, as illustrated in Figure 2.2e. When
the TiO2 thickness is further increased, enough donor impurities to bend the band upward,
hindering the electron flow to the water, as shown in Figure 2.2f. This result was
corroborated experimentally, when the thickness of TiO2 is increased beyond 15nm, we
observe no photocurrent, as shown in Figure 2.2b.
Figure 2.3 shows the photocatalytic I-V characteristics for TiO2 passivated GaP with
and without gold nanoparticles. Au nanoparticles deposited on GaP without TiO2
passivation (purple curve) improve the photocatalytic reaction rate by a slight shift of the
I-V curve to a lower overpotential with respect to bare GaP by approximately 0.2V .
Depositing Au nanoparticles on top of TiO2 passivated GaP results in a further
improvement in the I-V characteristics. For 0.5nm thick TiO2, however, significant
enhancement is observed in the photocatalytic reaction rate. For this dataset, we observe
both a downshift of the overpotential (by approximately -0.58V) and an increase in the
photocurrent (i.e., increased photoinduced charge) due to plasmonic field enhancement. At
V=-0.7V , the photocurrent is enhanced by a factor of 4X with respect to bare GaP, after
accounting for the shift in the overpotential. The photocatalytic properties of samples with
0.5nm TiO2 are substantially better than those with 1nm TiO2. This anomalous behavior of
0.5nm thick TiO2 was observed in several other samples consistently, and is the result of a
tradeoff between pn-junction formation and coupling of the localized plasmonic fields of
the Au nanoparticles to the actively absorbing GaP layer, as discussed below.
39
-1.5 -1.0 -0.5 0.0
-4.0
-3.5
-3.0
-2.5
-2.0
-1.5
-1.0
-0.5
0.0
Dark
Bare GaP
GaP w/ 5nm Au
GaP w/ 0.5nmTiO
2
and 5nm Au
GaP w/ 1nmTiO
2
and 5nm Au
GaP w/ 3nmTiO
2
and 5nm Au
Potential vs. Ag/AgCl (V)
Current Density (mA/cm
2
)
Figure 2.3 Photocurrent plotted as a function of voltage for GaP photocatalysts with various
thicknesses of TiO 2 with Au nanoparticles under 1W/cm
2
532nm illumination in a 0.5M Na 2SO 4
solution.
Figure 2.4 Electric field distributions calculated using the finite difference time domain method (a)
in the plane of the Au nanoparticles and (c,d) in the perpendicular direction across the
GaP/TiO 2/Au/electrolyte interface. (b) Calculated electric field enhancement factor plotted as a
function of TiO 2 thickness using Eq. 1.
40
Figure 2.4 shows the results of a finite difference time domain (FDTD) simulation
performed on a 5nm gold nanoparticle (nano island) film. The electric field distribution in
the plane of the gold island film is shown in Figure 2.4a. Here, localized hot spots can be
seen in regions between nearly touching nanoparticles/nanoislands separated by
approximately 2-3nm. This phenomenon has been studied in detail previously
28, 35, 37, 38
.
Figure 2.4c shows the electric field distribution in the perpendicular direction, across the
GaP/TiO2/Au/electrolyte interface at one of the hot spots between two nearly touching Au
nanoparticles with a TiO2 thickness of 0.5nm. Here, the electric field intensity can reach
1000X the incident electric field. These plasmonic nanoparticles couple light very
effectively from the far field to the near field at the GaP/TiO2 pn-junction. This is
advantageous for photocatalysis for two reasons. First, there is an increased electron-hole
pair generation rate in close proximity to the electrolyte interface and charge separating
region, enabling a larger fraction of the photoinduced charge to diffuse to the catalytic
surface and contribute to catalysis. We can calculate this fraction by integrating E
2
over the
volume of the catalyst from the GaP surface to one minority carrier diffusion length (100nm)
below this surface, as described by Eq 1.
(2.1)
41
In the denominator, the electric field intensity without the TiO2 and Au nanoparticles, Eo,
is integrated over the same volume as in the numerator. The second enhancement
mechanism arises from the increased light intensity at the pn-junction, which produces a
larger open circuit voltage
=
ln
+ 1 , and this further reduces the overpotential.
Here, the plasmon-enhanced electric fields create a larger photocurrent IL, which in turn
produces a larger Voc. In addition to electric field enhancement, these plasmonic
nanoparticles may produce hot electrons that can drive catalytic processes at a lower
applied overpotential
25
. Figure 2.4d shows the electric field distribution for a GaP/TiO2/Au
nanoparticle structure with a 3nm thick TiO2 film. Here, the localized plasmonic fields do
not extend into the actively absorbing GaP layer. Figure 2.4b shows the integrated electric
field enhancement factor calculated using Eq. 2.1 plotted as a function of TiO2 thickness.
In this Figure, a maximum enhancement of 42% can be seen at 0.5nm TiO2 thickness. The
sharp drop off of the EF as the TiO2 thickness increases is a direct result of the highly
localized nature of the plasmon-enhanced fields. The 42% electric field enhancement factor
is significantly smaller than those obtained in our previous studies of plasmon-enhanced
TiO2 (without GaP)
28
. This is a direct result of GaP’s relatively long minority carrier
diffusion length (100nm), which results in a larger integrated volume in Eq 2.1.
Figure 2.5 shows the two terminal photocurrent densities plotted as a function of
the applied overpotential for various GaP/TiO2/Au nanoparticle photocatalysts, measured
with a Pt counter electrode in a pH=0, 0.5M H2SO4 solution illuminated at 532nm. A large
42
reduction in the overpotential from approximately -0.8 to -0.2V can be seen for the
GaP/TiO2/Au structure with 0.5nm TiO2 thickness. Again, the plasmon-enhanced
photocatalyst with 0.5nm TiO2 is substantially better than those with thicker TiO2 films and
TiO2 without gold nanoparticles.
Perhaps the most important aspect of these TiO2 passivated GaP photocatalysts is
their photochemical stability in the electrolytic solution. Figures 2.6 and 2.7 show the time
dependence of the photocatalytic current and surface roughness without and with TiO2
passivation, respectively. In Figure 2.6a, the photocurrent density is plotted as a function
of time for bare GaP illuminated at 532nm in a 0.5M Na2SO4 solution with an applied
overpotential of -0.7V for 5 hours. An exponential decay can be seen with a time constant
of 0.45 hours indicating significant corrosion of the surface. An abrupt drop in the
-1.4 -1.2 -1.0 -0.8 -0.6 -0.4 -0.2 0.0
0.01
0.1
Dark
Bare GaP
GaP w/0.5nmTiO
2
GaP w/0.5nmTiO
2
&Au
GaP w/10nmTiO
2
GaP w/10nmTiO
2
&Au
Applied Overpotential (V)
Current Density (mA/cm
2
)
Figure 2.5 Two terminal photocurrent density (absolute value) plotted as a function of the applied
overpotential for various GaP/TiO 2/Au nanoparticle photocatalysts measured in a pH=0, 0.5M H 2SO 4
solution under 1W/cm
2
532nm illumination.
43
photocurrent occurs at 4.8 hours when the device failed. Figures 2.6b and 2.6c show optical
microscope and atomic force microscope images of the GaP surface after 5 hours of
illumination. Figure 2.6d shows a plot of the surface topography obtained along the dashed
white line indicated in Figure 2.6c, showing an RMS roughness of ±143nm. In contrast,
the photocurrent density of TiO2 passivated GaP is stable for 12 hours, as plotted in Figure
2.7a. The optical microscope image (Figure 2.7b) and atomic force microscope image
(Figure 2.7c) exhibit no evidence of surface corrosion or damage after 12 hours, with an
RMS surface roughness of ±1.0nm (Figure 2.7d).
Figure 2.6 (a) Time dependence of the photocurrent density of bare GaP illuminated with 1W/cm
2
532nm light in a 0.5M Na 2SO 4 solution at an applied overpotential of -0.7V. (b) Optical microscope
image, (c) atomic force microscope image, and (d) surface topography of the GaP surface after the 5
hour reaction.
44
2.4 Conclusion
In conclusion, plasmon-enhanced photocatalytic water splitting is observed on
TiO2passivatedGaP. The TiO2 passivation layer prevents corrosion of the GaP surface,
making it stable in pH=0 solution. In addition to preventing corrosion, the TiO2 passivation
layer provides enhancement in the photoconversion efficiency through the formation of a
Figure 2.7 (a) Time dependence of the photocurrent density of TiO 2passivated GaP illuminated with
1W/cm
2
532nm light in a 0.5M Na 2SO 4 solution at an applied overpotential of -0.7V. (b) Optical
microscope image, (c) atomic force microscope image, and (d) surface topography of the GaP/TiO 2
surface after the 12 hour reaction.
45
charge separating pn-region, which decreases carrier recombination and lowers the
overpotential required to initiate this reaction. Plasmonic Au nanoparticles deposited on
top of TiO2 passivated GaP further improve the photocatalytic process through plasmonic
field enhancement. These two enhancement mechanisms result in an optimum thickness of
the TiO2 layer of 0.5nm. Electromagnetic simulations performed using the finite difference
time domain (FDTD) method indicate that a 0.5nm film of TiO2 enables significant
coupling of the localized plasmonic fields of the Au nanoparticles with the actively
absorbing GaP layer. This general approach of passivating narrower band gap
semiconductors with TiO2will enable more efficient photocatalysts to be developed.
46
Chapter 3: Enhanced Surface Stable Solar-Driven H2 evolution on
TiO2-passivated p-GaAs
3.1 Introducton
Photocatalytic water splitting is of great interest for its potential to store solar
energy in the form of chemical bonds that can later be released without producing harmful
byproducts. In the initial demonstration of photocatalytic water splitting in 1972,
Fujishima and Honda used TiO2 under ultraviolet irradiation.
39
Because of TiO2’s wide
band gap (Eg=3.2eV), however, very few solar photons (∼4%) can be used to drive this
photocatalyst, and thus it will never provide efficient solar energy conversion on its own.
Despite this inherent limitation, the low cost, easy synthesis and chemical robustness of
TiO2 have made it by far the most extensively studied photocatalytic material.
7-9, 40
The III-
V compound semiconductors, such as GaAs and InP, are better candidate materials for solar
utilization with band gaps of 1.42eV and 1.34eV, respectively. These band gap energies
nearly match the Shockley-Queisser optimum band gap for solar energy conversion.
41
While the Shockley-Queisser limit is typically applied to electrical solar cells, its results
are general and also apply to photocatalysts.
42
The long-standing record high applied-bias
photon to current efficiency of 13.3% was obtained on p-type InP photocathodes covered
with Pt catalysts.
43
More recently, Ali Javey’s group reported 14% efficient H2 evolution
using InP nanopillars with a Ru catalyst.
44
While GaAs, a more widely used and more
47
readily available material, has a band gap and electron affinity that is similar to InP, it has
not been reported to provide high efficiency solar-to-hydrogen conversion. One of the main
reasons for this is the extremely high surface recombination velocity (10
6
cm/s) in GaAs,
which is 1 to 2 orders of magnitude higher than most other III-V compound semiconductors.
The rapid electron-hole pair recombination associated with this high surface recombination
velocity lowers the overall efficiency of photocatalytic energy conversion.
45-47
Another
major problem preventing GaAs from being utilized as a photocatalyst is that its surface is
not photochemically stable, like other III-V semiconductors.
14
The protective role of TiO2
for photoanodes has been discussed in the works of Hu et al. and Chen et al.
48, 49
Our
previous work demonstrated that a very thin layer of TiO2 can not only make GaP
photochemically stable, but also enhance photoconversion performance of GaP.
50, 51
More
recently, Lin et al. showed that photocatalytic H2 evolution by InP can be improved by
depositing TiO2 films with various ALD precursors, including TDMAT (i.e., [(CH3)2N]4Ti)
and titanium isopropoxide (i.e., C12H28O4Ti).
52
Although this phenomenon has been
discussed in these previous works, here, we provide direct evidence for the specific
mechanisms of catalytic enhancement through TEM, XPS, and PL spectroscopy.
In this work, we deposit very thin layers of TiO2 (1 to 5nm) on GaAs by atomic
layer deposition using TiCl4 as the Ti precursor, and observe drastic enhancement in the
photocatalytic efficiency compared to bare GaAs. We systematically investigate the
photocatalytic performance as a function of TiO2 thickness using gas chromatography and
48
by measuring their photo-I-V characteristics as a function of reference potential. These
photocatalytic surfaces are further studied using photoluminescence (PL) spectroscopy,
high resolution transmission electron microscopy (HRTEM), electron energy-loss
spectroscopy (EELS), and X-ray photoemission spectroscopy (XPS). To evaluate the
photochemical stability of the TiO2-passivated GaAs surfaces, we monitor the
photocatalytic yield over 48 hours under continuous illumination during the H2 evolution
reaction.
3.2 Experimental procedure
Figure 3.1 (a) Schematic diagram of sample geometry. (b) TEM Image of 3nm TiO 2 on GaAs. Note
that for TEM sample preparation and imaging purposes, the surface of the substrate is coated with
a thick layer of Pt.
(a)
(b)
49
We use p-type (111) oriented GaAs substrates with a Zn dopant concentration of
6x10
16
cm
-3
(obtained from University Wafer, Inc.) as the photocathode for H2 evolution.
Atomic layer deposition (ALD) of TiO2 was performed at 250
o
C on the p-GaAs wafers
using TiCl4 as the titanium source and water vapor as the oxygen source. The carrier gas
during the deposition was argon with a flow rate of 20sccm, and TiCl4 is always used for
the first half-cycle. The thicknesses of TiO2 were measured by ellipsometry. Ohmic back
contacts were made to the p-GaAs by evaporating 1nm thick Ti followed by 50nm of Au.
The Ti-Au film was then connected to the external circuitry with a copper wire and coated
with epoxy cement to insulate it from the electrolytic solution, as illustrated in Figure 3.1a.
Figure 3.1b shows a high resolution transmission electron microscope (HRTEM) image of
3nm amorphous TiO2 on GaAs below a layer of electron beam deposited Pt (the Pt was
Figure 3.2 TEM Image of 3nm native oxide on GaAs. Note that for TEM sample preparation and
imaging purposes, the surface of the substrate is coated with a thick layer of Pt.
50
deposited as part of the cross section sample preparation using standard focused ion beam
sample preparation techniques). This high resolution TEM image confirms the thickness
of the TiO2 film, which was independently measured by ellipsometry. An HRTEM image
of the GaAs surface prior to TiO2 deposition is also shown in Figure 3.2, which shows a
3nm native oxide on GaAs. A scanning TEM (STEM) image and spatial maps of the Ti and
O species obtained from EELS spectra are shown in Figure 3.3. Based on these spatial
maps, it is evident that the native oxide of the GaAs has been removed during the ALD
process due to the Cl
-
ions from the TiCl4 precursor, consistent with previous reports in the
literature.
53
If there were a native oxide present, the oxygen EELS signal would increase
before the titanium signal. Instead, the O and Ti EELS signals increase simultaneously,
corresponding to the TiO2 interface. As a surface passivation layer, amorphous TiO2 is far
better than crystalline TiO2, mainly because crystalline TiO2 is insulating and, therefore,
impedes charge transfer to the ions in solution, whereas amorphous TiO2 is conducting.
48,
54
We measured the photocatalytic reaction rates of two sets of samples in a pH=0 solution
of 0.5 M H2SO4 using a three-terminal potentiostat with the prepared samples, a Ag/AgCl
electrode, and a graphite electrode functioning as the working, reference, and counter
electrodes, respectively. A 910 mW/cm
2
532nm laser and AM1.5 G illumination solar
simulator (100 mW/cm
2
) were used for illumination. The area of the exposed electrode
surface is about 0.2 cm
2
. H2 evolution was verified and quantified by gas chromatography.
Photoluminescence spectra were collected using a 100X objective lens, a 1800 l/mm
51
grating, and a silicon CCD detector. A 532nm continuous-wave laser was used to excite the
samples, and spectra were collected over the 750 nm to 1000 nm wavelength range. Low
to moderate power excitation (10
2
-10
5
mW/cm
2
) was used to avoid optical heating.
Figure 3.3 (a) Ti L edge map and (b) O K edge map from electron energy-loss spectroscopy. (c) Profile of Ti
L edge (blue background) and O K edge map (red line). (d) STEM image of 3nm TiO2 on GaAs sample. Note
that for STEM sample preparation and imaging purposes, the surface of the substrate and TiO2 is coated with
a thick layer of e-beam deposited Pt. (The microscope used was a JEOL JEM-2100F at 200kV high tension. )
52
3.3 Results and discussion
Figure 3.4a shows the photocurrent-voltage curves of p-GaAs photocathodes with
various thicknesses of TiO2 under an AM1.5 G illumination solar simulator, plotted
together with bare GaAs. The bare GaAs (blue curve) exhibits an onset of photocurrent at
a potential of approximately -0.05V vs. RHE. The TiO2-passivated GaAs shows a clear
shift in the onset potential. The sample with a nominal thickness of 1nm TiO2 exhibits the
most prominent shift, lowering the overpotential (i.e., increasing the onset potential) by
approximately 0.4V (red curve) with a 32-fold enhancement of photocurrent over bare
GaAs (at 0V vs RHE). It is interesting that such a large shift in the onset potential is
obtained from such a thin layer of material. Nevertheless, these results were reproduced
consistently in several different sets of samples. We observe no improvement over bare
GaAs for TiO2 thicknesses above 10nm, as shown in Figure 3.5. The hydrogen evolution
Figure 3.4 (a) Photocatalytic current-potential curves and (b) corresponding photon-energy
conversion efficiency calculated by using Equation 1 for GaAs photocatalysts with various
thicknesses of TiO 2 under an AM1.5 G illumination in a 0.5M H 2SO 4 pH=0 solution.
-0.3 -0.2 -0.1 0.0 0.1 0.2 0.3 0.4 0.5 0.6
-20
-15
-10
-5
0
Dark
Bare GaAs
GaAs w/1nm TiO
2
GaAs w/3nm TiO
2
GaAs w/5nm TiO
2
Potential vs. RHE (V)
Current Density (mA/cm
2
)
0.0 0.1 0.2 0.3 0.4 0.5
0.0
0.5
1.0
1.5
Potential vs. RHE (V)
Bare GaAs
GaAs w/1nm TiO
2
GaAs w/3nm TiO
2
GaAs w/5nm TiO
2
Efficiency (%)
(a) (b)
53
efficiency (η) was evaluated by computing the applied-bias photon to current efficiency of
the photocatalytic reaction according to the following equation
55
=
(
)∗
100% (3.1)
where Ein is the applied reference potential vs. RHE, Eout is the voltage of the redox couple
(H
+
/H2) vs. RHE (which is 0V in this case), I is the photocurrent density (mA/cm
2
)
observed, and P0 is the incident optical power density (100 mW/cm
2
). This equation gives
the electrode efficiency for the H2 evolution half-reaction.
56, 57
The maximum applied-bias
photon to current efficiency for p-GaAs with a 1nm TiO2 passivation layer reaches
approximately 1.5% at an applied potential of +0.2 V vs. RHE, as shown in Figure 3.4b. It
should be noted that these results were obtained from planar substrates without any co-
catayst particles. With further modification, such as using a nanopillar structuring and/or
decorating with Ru or Pt co-catalyst particles, the H2 production efficiency can be further
improved.
-0.8 -0.6 -0.4 -0.2 0.0 0.2 0.4 0.6
-100
-80
-60
-40
-20
0
Dark
Bare GaAs
1nmTiO
2
GaAs
10nm TiO
2
GaAs
15nm TiO
2
GaAs
Potential vs. RHE (V)
Current Density (mA/cm
2
)
Figure 3.5 Photocurrent plotted as a function of voltage for GaAs photocatalysts with 10nm and 15nm
TiO 2 under 532nm illumination (910 mW/cm
2
) in a 0.5M H 2SO 4 solution.
54
Figure 3.6 shows X-ray photoemission spectroscopy (XPS) of these TiO2 deposited
GaAs samples. Figure 3.6a shows the core level binding energies of Ti 2p1/2 and Ti 2p3/2 in
the TiO2 film at 464.7 and 459.0 eV , respectively. Additional XPS peaks with lower binding
energy of 457.1 eV (red fit curve) are assigned to Ti
3+
. The area ratios of Ti
3+
to Ti 2p1/2 are
0.15, 0.09, and 0.061 for the 1nm, 3nm, and 5nm TiO2 films, respectively, indicating a
larger Ti
3+
state density in thinner TiO2 films. Figure 3.6b shows the O 1s core level peaks
for samples with 1nm, 3nm, and 5nm thick TiO2 deposited on GaAs. Since the native oxide
Figure 3.6 (a) Ti 2p and (b) O 1s core level XPS spectra of various thickness of TiO 2 on GaAs.
4 6 8 4 6 6 4 6 4 4 6 2 4 6 0 4 5 8 4 5 6 4 5 4
GaAs w/ 1nm TiO
2
Ti
3+
GaAs w/ 3nm TiO
2
GaAs w/ 5nm TiO
2
Experimatal
Fitted
Ti
3+
Ti
3+
Binding Energy (eV)
5 3 6 5 3 4 5 3 2 5 3 0 5 2 8
Oa
Ob
Experimatal
Fitted
GaAs w/ 1nm TiO
2
GaAs w/ 5nm TiO
2
GaAs w/ 3nm TiO
2
Binding Energy (eV)
(a) (b)
55
of the GaAs is removed during the ALD process, as is evident from the HRTEM data (see
Figure S1 of the supporting information), both O peaks in these XPS spectra are believed
to originate from the TiO2 thin layers. The O 1s core level peaks are fitted with two
symmetric Gaussian curves denoted as Oa and Ob. The Oa peak is ascribed to oxygen
atoms of stoichiometric TiO2 while the Ob peak corresponds to oxygen vacancies.
58
The
area ratios of Ob to Oa are 2.4, 1.6, and 1.4 for the 1nm, 3nm, and 5nm TiO2 films
respectively, indicating that the ratio of Ob to Oa decreases as the thickness increases. This
implies that there are higher densities of oxygen vacancies in thinner TiO2 films, which is
consistent with the results of the Ti
3+
XPS peaks. It is widely accepted that the reduction
of H2O to form H2 occurs predominantly at the Ti
3+
-O vacancy sites due to the reaction of
adsorbed H
+
ions.
59, 60
Thus, a higher concentration of O vacancies correspond to more
active sites, resulting in a higher hydrogen generation efficiency. The dark I-V
characteristics of TiO2-passivated GaAs (Figure 3.7) also shows a large shift in the onset
potential indicating that the catalytically active sites are primarily responsible for the
enhanced H2 evolution reaction. XPS data of GaAs with 3nm TiO2 passivation after a 12
hour reaction (see Figure 3.8) shows the presence of Ti
3+
peaks after the reaction.
56
-0.8 -0.6 -0.4 -0.2 0.0 0.2
-1.5
-1.0
-0.5
0.0
Bare GaAs
1nmTiO
2
GaAs
Potential vs. RHE (V)
Current Density (mA/cm
2
)
Figure 3.7 Dark current plotted as a function of voltage for GaAs photocatalysts with and without
TiO 2 thin layer in a 0.5M H 2SO 4 solution.
468 466 464 462 460 458 456 454
Experimatal data
Fitting data
Intensity (a.u.)
Binding Energy (eV)
GaAs w/ 3nm TiO
2
Ti
3+
Figure 3.8 Ti 2p core level XPS spectra of 3nm TiO 2 on GaAs after H 2 evolution reaction.
57
In order to further explore the role of this TiO2 (or TiO2-x) surface layer, we also
measured the photoluminescence (PL) spectra of these GaAs samples with various
thicknesses of TiO2. These spectra are plotted in Figure 3.9. Here, we find that the sample
with the highest PL efficiency (i.e., bare GaAs) has the lowest photocatalytic efficiency,
and vice versa, the sample with the lowest PL efficiency (1nm TiO2) has the highest
photocatalytic efficiency. Initially, this was somewhat surprising, since materials with
strong photoluminescence efficiencies typically make good solar cells and photocatalysts.
However, it is apparent that the catalytically active surface states also cause strong electron-
hole recombination, which limits the PL efficiency. However, the benefit that these surface
states provide by lowering the potential barrier of the reaction and promoting charge
transfer outweighs their detriment associated with charge recombination. For the sample
with nominally 1nm TiO2, we observe a 5-fold reduction in the photoluminescence
750 800 850 900 950 1000
0
1x10
4
2x10
4
3x10
4
4x10
4
5x10
4
Worst photocatalyst
Bare GaAs
GaAs w/1nm TiO
2
GaAs w/3nm TiO
2
GaAs w/5nm TiO
2
Wavelength (nm)
PL Intensity (Counts)
Best photocatalyst
Figure 3.9 Photoluminescence spectra of GaAs photocatalysts with various thicknesses of TiO 2
under 532nm excitation.
58
intensity (efficiency), indicating a high density of surface states, which act as non-radiative
recombination centers, thus lowering the PL efficiency. GaAs with slightly thicker (3nm
and 5nm) TiO2 exhibits higher PL intensities than 1 nm TiO2, by reducing the effects of
these surface states. It is tempting to attribute the reduction in PL intensity to the built-in
field at the n-TiO2/p-GaAs junction. Here, the band bending will sweep photogenerated
electrons out to the ions in solution, preventing them from recombining (radiatively) with
the holes in the semiconductor. However, this is unlikely for a 1nm film of TiO2, which
will have a negligible effect on the charge separation fields compared to the semiconductor-
liquid fields.
Figure 3.10 Time dependence of the photocurrent density of (a) bare GaAs and 3nm TiO 2 passivated
GaAs illuminated for 12 hours and (b) 3nm TiO 2 passivated GaAs illuminated for 48 hours with
532nm light at an applied overpotential of 0.1V vs. RHE in a 0.5M H 2SO 4 solution.
0 2 4 6 8 10 12
1
10
100
Bare GaAs
1nm TiO
2
on GaAs
ICurrent DensityI (mA)
Time (hours)
0 10 20 30 40 50
0
20
40
60
80
100
120
ICurrent DensityI (mA)
Time (hours)
(a)
(b)
59
From thermodynamic considerations alone, GaAs (either p- or n-type) is known
to be unstable under both anodic and cathodic conditions.
61
In our previous work of
H2 evolution on TiO2-passivated GaP
50
, atomic force microscopy (AFM) as well as the
photo-I-V characteristics showed that thin layers of TiO2 protect the underlying GaP from
corrosion in the strongly acid electrolytic solution. We observe similar findings for TiO2-
passivated GaAs. Figure 3.10a plots the time dependent photocurrent of bare GaAs and
TiO2-passivated GaAs, which shows that bare GaAs corrodes readily while TiO2-
passivated GaAs is stable. We performed measurements extending over many hours
monitoring H2 evolution produced by GaAs coated with 3nm TiO2 under 532nm
wavelength illumination with an applied potential of +0.1V (vs. RHE) in a 0.5M H2SO4
solution. The yields of H2 were measured by gas chromatography after 12, 24, and 48 hour
Figure 3.11 (a) Gas chromatograph (GC) data taken after 12, 24, and 48-hour reactions on GaAs
with 3nm TiO 2 under 532nm illumination in a 0.5M H 2SO 4 solution. (b) H 2 yield plotted as a
function of time.
2.0 2.1 2.2 2.3
GC Signal (a.u.)
12hr
24hr
48hr
H
2
Air
Retention Time (min)
0 10 20 30 40 50
0
50
100
150
H
2
Yields ( mol)
Reaction Time (hours)
(a)
(b)
60
illumination times and are plotted in Figure 3.11a. The H2 peak appears at a GC retention
time of 2.1 min, and the integrated area increases linearly with the illumination time. After
calibration, the yields of H2 in µmol was plotted as a function of time, as shown in Figure
3.11b. The linear relationship of the H2 yield versus time indicates that the TiO2-passivated
GaAs is photochemically stable in this strongly acidic electrolyte (pH=0). Also, the time
dependence of the photocatalytic current taken over a 48-hour period is plotted in Figure
3.10b, which shows that this photocatalyst is stable over 48 hours of illumination. The
Faraday efficiencies of the 12, 24 and 48 hour reactions are around 92%,as summarized in
Table 3.1. We also measured the Faraday efficiency of bare GaAs to be only 6%, which is
15X lower than the TiO2-passivated Faraday efficiency. This is most likely due to the
photoelectrochemically-driven corrosion of the GaAs crystal structure that is circumvented
by the TiO2 layer.
Table 3.1 H 2 yields and Faradaic efficiencies of the 12, 24, and 48-hour reactions on GaAs with
3nm TiO 2 under 532nm illumination in a 0.5M H2SO4 solution.
61
3.4 Conclusion
In conclusion, we observe enhanced photocatalytic H2 evolution on TiO2-
passivated GaAs. The TiO2 layer hosts active surface states on the GaAs, which increase
the recombination rate of photogenerated electron-hole pairs. The benefit of these active
sites in lowering the potential barrier for this reaction outweigh the unfavorable effects of
charge recombination, resulting in a net enhancement in the photocurrent density. The
highest efficiency is obtained from GaAs with a 1nm TiO2 passivation layer under AM1.5
G illumination, which reaches 1.5% at an applied voltage of approximately +0.2V vs. RHE.
We observe no improvement over bare GaAs for TiO2 thicknesses above 10nm, where the
insulating nature of the TiO2 eventually outweighs its benefits. The TiO2 passivation layer
prevents photocorrosion of the GaAs surface, providing a viable, long-term stable
photocatalyst. Moreover, the Faraday efficiency of TiO2-passivate GaAs reaches around
92% that is 15X higher than the one of bare GaAs which is only 6%.
62
Chapter 4: Photocatalytic conversion of CO2 to hydrocarbon fuels
on GaP via plasmon enhanced absorption
4.1 Introducton
The photoelectrochemical reduction of CO2 is an exciting reaction system with the
ability to convert an abundant greenhouse gas to combustible hydrocarbon fuels using
sunlight. The direct conversion of solar-to-chemical energy has several advantages over
solar-to-electric energy conversion, most notably, the ability to store large amounts of
energy (~GW) in chemical bonds that can later be released in a carbon neutral cycle.
62-72
Many attempts have been made to reduce CO2 by 2e
-
to various species such as CO and
formic acid, as reported in previous literature.
73-80
Few researchers have achieved further
reduction to CH3OH or CH4.
81-83
Methanol is an attractive product with a relatively high
energy density, which can be easily integrated into the existing liquid fuel technologies.
68,
84
However, the photocatalytic reduction of CO2 with H2O to methanol requires 6 electrons
and many intermediate species, some of which have extremely high energy barriers.
85
The
most likely first step in this multi-electron reaction is the one electron reduction to the CO2
-
intermediate,
86
which lies 1.7eV above the conduction band of TiO2 and 1.2eV above GaP.
The mechanism for electrochemical CO2 reduction was first proposed by Bockris et al.
87-
89
The high overpotential required for this reaction was attributed to the formation of the
CO2
-
intermediate, which consequently converts to CO via the general process CO2 + e
-
63
CO2
-
, CO2
-
+2H
+
+ e
-
CO + H2O.
64, 70, 81, 90
In 1978, Hallman’s group first reported CO2
reduction on p-GaP under 365nm illumination with an applied overpotential of -1.4V (vs
SCE).
73
Fujishima and Honda demonstrated photoelectrocatalytic reduction of CO2 to
formaldehyde and methanol by irradiating TiO2 and GaP with the UV light at an
overpotential of -1.5V (vs SCE).
20, 74
Canfield later reported CO2 reduction to methanol on
p-InP with an overpotential of -1.3V (vs SCE).
80
More recently, Bocarsly’s group
demonstrated pyridinium-catalyzed CO2 reduction on GaP photocathodes with
overpotentials between -0.7V and -0.2V (vs SCE) under UV light.
81
Despite these
interesting prior results, the stability of these materials against photocorrosion has not been
addressed.
In the work presented here, we investigate the photocatalytic performance and
stability of TiO2-passivated p-GaP using atomic force microscopy (AFM),
photoelectrochemistry, and optical microscopy. The photocatalytic efficiency is studied
systematically as a function of TiO2 layer thickness using a three-terminal potentiostat. The
products are detected using NMR spectroscopy and gas chromatography, systematically as
a function of applied overpotential.
4.2 Experimental procedure
Zn doped p-type (100) oriented GaP with a dopant concentration of 2x10
18
cm
-3
was used as the photocatalyst for CO2 reduction with an active area of 0.5cm×1cm. Atomic
64
layer deposition (ALD) of anatase TiO2 was performed at 250ºC on the p-GaP wafers with
TiCl4 as the titanium source and water vapor as the oxygen source. The antase crystal phase
was verified by Raman spectroscopy. Using ellipsometry, we established that 100 cycles
of ALD produces a 4nm thick TiO2 film and 1000 cycles produces a 40nm film. A Ga-In
eutectic film was painted on the back of the p-GaP to form an Ohmic contact. The Ga-In
contact was then connected to the external circuitry with a copper wire and coated with
epoxy cement to insulate it from the electrolytic solution, as illustrated in Figure 4.1. While
this planar geometry is not ideal for high efficiency photoconversion, it enables us to study
surface stability. A three-terminal potentiostat was used with the prepared semiconductor
samples as the working electrode, a Ag/AgCl electrode as the reference electrode, and a Pt
electrode functioning as a counter electrode. The photocatalytic reaction rates of two sets
of samples were measured in a 2ml solution of 0.5 M NaCl, with and without 10mM
pyridine, while continuously bubbling CO2 through the solution. 0.5M NaCl was chosen
as the electrolyte solution because of its high conductance, and its ability to stabilize the
intermediate states involved in the CO2 reduction.
91
In this setup, we analyze the products
evolving at the working electrode, instead of at the counter electrode. It is likely that
oxygen is also produced in the reaction.
65
4.3 Results and discussion
Figure 4.1 Schematic diagram of sample geometry.
Figure 4.2 (a) Optical microscope image, (b) atomic force microscope image, and (c) surface
topography of bare GaP surface after 8h reaction at -0.5V overpotential. (d) Optical microscope
image, (e) atomic force microscope image, and (f) surface topography of 5nm TiO 2 on GaP surface
after 8h reaction.
66
While photocatalysis on GaP (and other III-V compound semiconductors) has been
demonstrated previously,
80, 81, 92
this material corrodes rapidly under photo-electrochemical
conditions and is significantly degraded after just 30 minutes of illumination. In order to
make GaP photochemically stable, we passivated the surface using a thin film of TiO2
deposited by ALD. Figures 4.2a and 4.2b show optical microscope and atomic force
microscope images of the bare GaP surface after 8 hours of illumination. Figure 4.2c shows
a plot of the surface topography obtained along the dashed white line in Figure 4.2b,
showing an RMS roughness of ±54nm, which indicates that substantial photocorrosion has
taken place and that this will not serve as a viable photocatalyst. In contrast, the
photocurrent density of TiO2-passivated GaP is stable for 8 hours. The optical microscope
image (Figure 4.2d) and atomic force microscope image (Figure 4.2e) exhibit no evidence
of surface corrosion or damage after 8 hours, with an RMS roughness of ±1nm (Figure
4.2f), indicating that this is a long term, stable photocatalyst. Here, the TiO2 significantly
improves the photo-stablilty of the GaP surface, however, more extensive time-dependent
studies are needed in order to establish the extent of this long term stability.
Catalysts Shift of onset overpotential
compared to Bare GaP (V)
Bare GaP 0
1nm TiO
2
@GaP 0.13
3nm TiO
2
@GaP 0.31
5nm TiO
2
@GaP 0.40
10nm TiO
2
@GaP 0.52
Table 4.1 Shift of onset overpotential for samples with different thicknesses of TiO 2.
67
In addition to providing a stable photocatalytic surface, the TiO2 passivation layer
results in an increase in the photoconversion efficiency. Figure 4.3a shows the
photocurrent-voltage curves for GaP passivated with various thicknesses of TiO2 measured
in a 0.5M NaCl, 10mM pyridine solution under 532nm illumination. During these
measurements, CO2 is continuously bubbled through the solution. Bare GaP (green curve)
(d) (c)
3.5 3.4 3.3 3.2 3.1 3.0
Bare GaP
5nm TiO
2
passivated GaP
ppm
Methanol
-1.0 -0.5 0.0
-1.4
-1.2
-1.0
-0.8
-0.6
-0.4
-0.2
0.0
Dark
Bare GaP
GaP w/1nm TiO
2
GaP w/3nm TiO
2
GaP w/5nm TiO
2
GaP w/10nm TiO
2
Potential vs. NHE (V)
Current Density (mA/cm
2
)
(a)
(b)
0 1 2 3 4 5 6 7 8 9 10
0.0
0.1
0.2
0.3
0.4
0.5
0.6
Thickness of TiO
2
(nm)
Decrease of Overpotential (V)
0.5V
1 2 3 4 5 6 7 8 9 10
0.0
0.1
0.2
0.3
0.4
0.5
Built-in Voltage (eV)
TiO
2
Thickness (nm)
Figure 4.3 (a) Photocatalytic current-potential curves of GaP photocatalysts with different TiO 2
thicknesses in a 0.5M NaCl, 10mM pyridine solution under 532nm wavelength laser illumination.
(b) Decrease of overpotential plotted as a function of TiO 2 thickness on GaP. (c) Calculated built-in
voltage plotted as a function of TiO 2 thickness. (d) NMR spectra showing methanol production using
bare GaP and 5nm TiO 2-passiavated GaP photocatalysts at an overpotential of -0.50V.
68
has an onset of photocurrent at a potential of approximately -0.15V (vs NHE). For TiO2-
passivated GaP, we see a clear shift in the overpotential required to drive this reaction with
increasing thickness of the TiO2, as plotted in Figure 4.3b. Table 4.1 lists the shift of onset
overpotential of samples with different thicknesses of TiO2. For example, the onset
potential for 10nm TiO2 (red curve) is shifted by 0.5V with the respect to bare GaP. This
shift is attributed to the passivation of surface states that cause non-radiative recombination
and the formation of a pn-junction, which is created because the ALD-deposited TiO2 tends
to be n-type doped due to oxygen vacancies
36, 93
. Figure 4.3c shows the built-in potential
for the junction calculated using
=
(
)
(
)
, assuming a doping
concentration of Na=5x10
18
cm
3
.
94
Here,
is the depletion width of the GaP-TiO2
junction, which is a function of the TiO2 layer thickness. This simple calculation predicts
values similar to the experimentally observed shift in the overpotential plotted in Figure
4.3b. Beyond 10nm, however, the photocurrent decreases rapidly with increasing TiO2
thickness due to band bending at the n-type TiO2/electrolyte interface, which blocks
electrons. No enhancement is observed for TiO2 thicknesses above 10nm, due to the
insulating nature of the TiO2, which eventually outweighs the benefits of passivation.
While, TiO2 does not absorb light at 532nm, the pn-junction formed with the GaP enables
separation of the photo-generated charge in the actively absorbing GaP. Figure 4.3d shows
the NMR spectra taken after 8 hours of illumination with an overpotential of -0.50V vs
NHE for GaP with and without TiO2 passivation. This data shows a clear peak
69
corresponding to methanol, as reported previously by Barton et al.
81
Gas chromatography
FID data has also been used to verify the production of methanol, as shown in Figure 4.4.
Based on this GC FID data for the 5nm thick TiO2 sample, we calculated that 4.9 (±0.02)
µmol of CH3OH are produced during an 8 hour reaction consuming 5.2 Coulombs of
charge. Dividing by this ratio by the stoichiometric factor of 6, yields a Faradaic efficiency
of 55%. Also, according to the GC TCD data of the same experiment, H2 is produced with
a Faradaic efficiency of 30%. The photo-conversion efficiency, however, can be
significantly less than this due to non-radiative recombination, which is unknown,
particularly for this planar sample geometry. As a control experiment, the same reaction
was run under the same electrochemical conditions of -0.50V vs NHE without laser
illumination, which resulted in no measureable current and no detectable methanol in the
NMR spectra. In order to rule out other sources of carbon in this reaction, we used
isotopically labeled
13
CO2 as the carbon source in this reaction and observed
13
CH3OH in
the
13
C-NMR spectrum shown in Figure 4.5
95
. In addition, we repeated the experiment,
purging with Ar instead of CO2, and found no production of hydrocarbons. Therefore, we
are confident that CO2 is the only carbon source in this reaction. Previously, it was reported
that a pyridine catalyst is required to drive this reaction on GaP. The pyridinium radical
serves as a one-electron charge-transfer mediator, which is capable of efficiently
transferring all six electrons to reduce CO2 to methanol, thereby circumventing the high
energy barrier of the one-electron reduction of CO2 mentioned above.
92
However, we
70
observe the same methanol peak in our NMR spectra without pyridine in solution (see
Figure 4.6), indicating that this catalyst is not, in fact, required to drive this reaction at low
overpotentials. While CH3OH products are observed without pyridine, the yield is one third
that of the system with pyridine, indicating that the pyridine, in fact, helps lower the energy
barriers of the reaction by forming an inner-sphere-type electron transfer system.
92, 96
.
Atomic force microscopy shows that the GaP/TiO2 is photochemically stable without
pyridine, as shown in Figure 4.7.
4.4 4.6 4.8 5.0 5.2 5.4 5.6 5.8 6.0
Bare-gap
10
-4
methanol
TiO
2
passivated p-GaP
GC Signal (a.u.)
Retention Time (min)
methanol peak
Figure 4.4 Gas chromatograph (GC) data taken after 8 hour illumination (532nm) of bare and 5nm
TiO 2 passivated-GaP in 0.5M NaCl, 10mM pyridine solution at an overpotential of -0.5V (vs NHE).
The GC data is plotted together with a calibration standard consisting of 10
-4
M methanol in aqueous
solution. Based on this data, the 5nm TiO 2 passivated-GaP was found to have a Faradaic efficiency
of 55%.
71
52 51 50 49 48
ppm
5nm TiO
2
w/ GaP
13
CH
3
OH, 13-C peak
Figure 4.5
13
C-NMR spectrum showing methanol peak taken after 8 hour illumination (532nm) of
5nm TiO 2 passivated-GaP at an overpotential of -0.5V (vs NHE) in 0.5M NaCl, with 10mM pyridine
solution.
3.5 3.4 3.3 3.2 3.1 3.0
With pyridine
Without pyridine
ppm
Methanol
Figure 4.6
1
H-NMR spectra showing methanol peaks taken after 8 hour illumination (532nm) of
5nm TiO 2 passivated-GaP at an overpotential of -0.5V (vs NHE) in 0.5M NaCl, with 10mM pyridine
solution and without 10mM pyridine.
72
In order to understand the mechanism of this reaction, we must consider the
energetics of the electrons in this photocatalytic structure. The conduction bands of GaP
and TiO2 lie slightly above NHE at -0.7V and -0.2V vs NHE, respectively, as plotted in
Figure 4.8. This leaves an energy barrier of at least 1.2V for the electrons to overcome in
the reduction of CO2. The -0.5V externally applied overpotential (Vext) accounts for part of
this required energy, and the photovoltage produced at the internal pn-junction and/or the
liquid-semiconductor junction (VPV) can easily account for the remaining -0.7V . From the
flat band voltage, we can obtain the open circuit voltage, as follows: (Voc)max=|Vfb-Vredox|,
where Vfb is flat-band potential and Vredox is the potential of the redox couple
97
. From Mott-
Schottky measurements, we obtained a flat-band potential of 0.4 V versus NHE, which is
consistent with previous values from literature
81, 98
. Using Vredox
(Ferricyanide/Ferroyanide)=0.36V , we obtain an open circuit voltage of Voc=0.76V , which
is large enough to cover the remaining -0.7V depicted in Figure 4.8. This photovoltage is
Figure 4.7
(a) Optical microscope image, (b) atomic force microscope image, and (c) surface
topography of TiO 2 passivated GaP surface after 8h reaction in 0.5M NaCl without pyridine at -
0.5V overpotential.
73
reasonable considering GaP’s relatively large band gap of 2.25eV .
These estimations assume that the electrons traverse the TiO2 layer ballistically, and
do not equilibrate to the TiO2 conduction band edge. In the diffusive case, an additional
0.5V would be required. We believe this is one of the reasons why the TiO2 layer must be
made very thin. Several aspects of these results and their underlying mechanism are quite
surprising. First, we had initially thought that the TiO2 layer, which is insulating and has a
low conduction band energy, would lower the overall photocatalytic efficiency (i.e.,
e
-
-1.9V CO
2
-
/CO
2
p-GaP n-TiO
2
-0.7V
0V (NHE)
GaP
Conduction
Band
E
g
=2.25eV
e
+
V
ext
=-0.5V
V
pv
=-0.7V
Electrolyte
Figure 4.8 Energy band alignment of GaP and TiO 2 together with the relevant redox potentials of CO 2.
74
photocurrent), but would, at least, provide a stable, viable catalyst. Much to our surprise,
the TiO2 layer actually improved the overall photo-conversion efficiency. The reasons for
this are three-fold: 1.) The TiO2 reduces non-radiative recombination of the photo-excited
electron-hole pairs. 2.) The electrons traverse the TiO2 ballistically and, therefore, do not
relax to the conduction band edge. 3.) The formation of a pn-junction provides an additional
photovoltage required to drive the reaction.
It is important to note, however, that the CO2
-
reduction potential of -1.9V vs NHE is
calculated from simple thermodynamic considerations for isolated CO2
-
species, and does
not include the effects of the solution or catalytic surface. As a result, the energetics of the
actual CO2
-
intermediates can be quite different due to the presence of the aqueous solution
and/or the catalytic surface. In a mechanism proposed by Anpo et al., the CO2
-
intermediate
is strongly bound to a proposed Ti
3+
active site (oxygen vacancy) on the TiO2 surface, thus
lowering its energy
99
. Another strategy for lowering the reaction barrier is stabilizing the
CO2
-
intermediate, which was recently demonstrated using an ionic liquid electrolyte co-
catalyst where the cation forms a complex with the anionic intermediate.
100
Two-electron
processes have also been proposed by Tananka et al., which would circumvent this first
intermediate step altogether
101
. While several mechanisms have been proposed in the
literature, further spectroscopic studies are needed in order to verify the catalytic reaction
pathway.
75
4.4 Conclusion
In conclusion, we report photocatalytic CO2 reduction on TiO2-passivated GaP. The
TiO2 passivation layer successfully stabilizes the GaP surface in solution, preventing it
from photocorrosion. In addition, the TiO2 passivation layer provides enhancement in the
photoconversion efficiency through the passivation of surface states and the formation of
a charge separating pn-region, which reduces carrier recombination and lowers the
overpotential required to initiate this reaction by approximately 0.5V . This general
approach of passivating narrower band gap semiconductors with TiO2 will enable more
efficient photocatalysts to be developed and a broad range of materials to be considered for
photocatalysis that make more efficient use of the solar spectrum. We also observe CH3OH
evolution with and without pyridine catalyst, indicating that this catalysts is not, in fact,
required to drive this reaction at low overpotentials.
76
Chapter 5: Artificial photosynthesis on TiO2-passivated InP
nanopillars
5.1 Introducton
Here, we utilize TiO2-passivated nanotextured InP photo-cathodes to explore
aqueous CO2 reduction to methanol under 532nm illumination as a function of applied
potential. The selectivity of methanol production is compared for InP samples prepared
with and without Cu nanoparticles and TiO2-passivation layers. The photocatalytic surface
is characterized by high resolution transmission electron microscopy (HRTEM), in order
to provide a detailed picture of the nanoparticle/TiO2/InP interface. Plane wave density
functional theory (PW-DFT) calculations are carried out to explore the role of surface
binding of reactants and intermediate species to the TiO2 surface.
5.2 Experimental procedure
A schematic diagram of the sample geometry is illustrated in Figure 5.1a. Zn doped
p-type (100) oriented InP nanopillars (NPLs) with a dopant concentration of 5x10
17
cm
-3
were used as the photocatalyst for CO2 reduction. The InP NPLs are around 80nm in
diameter and approximately 400 to 600 nm tall, and have an average pitch of approximately
50 nm (Figure 5.1b). Briefly, InP bulk wafers are treated in a reactive ion O2 plasma
77
treatment followed by a two-minute wet-etching step in HCl/H3PO4 (3:1) to remove the
surface-damaged layers and contaminants.
56
Atomic layer deposition (ALD) of TiO2 was
performed at 250
o
C on the p-InP wafers with TiCl4 as the titanium source and water vapor
as the oxygen source. The average rate of deposition is approximately 0.44Å per cycle, as
calibrated by ellipsometry. We then evaporate copper with a nominal thickness of 0.5nm
on the top surface of the TiO2. A Zn-Au film was evaporated on the back of the p-InP to
form an Ohmic contact. The Zn-Au contact was then connected to the external circuitry
with a copper wire and coated with epoxy cement to insulate it from the electrolytic
solution. Figures 5.1c–5.1e show transmission electron microscope (TEM) images of the
0.5nm Cu on the top of the InP NPLs. These figures indicate that, instead of forming a
uniform continuous film, the evaporated 0.5nm Cu forms crystal nanoparticles with
diameters around 20nm. In Figure 5.1e, a thin amorphous layer of CuO can be seen on the
surface of the Cu nanoparticles, which is formed in air. The XPS spectra shown in Figure
5.2. confirms the existence of Cu 2p and O 1s peaks corresponding to CuO. The thickness
of the deposited TiO2 is around 3nm, as shown in Figure 5.1d.
78
(c) (b)
(e) (d)
(a)
Figure 5.1 (a) Schematic diagram of TiO 2-passivated InP nanopillars with Cu cocatalyst
nanoparticles. (b-e) SEM and TEM images of InP nanopillar array with TiO 2 deposition layer and
Cu nanoparticles. The high resolution TEM image in (e) resolves the crystal lattice of the Cu
nanoparticles.
79
A three-terminal potentiostat was used with the prepared semiconductor samples as
the working electrode, a Ag/AgCl electrode as the reference electrode, and a Pt electrode
functioning as the counter electrode. The photocatalytic reaction rates of two sets of
samples were measured in a 0.5 M KCl solution, while continuously bubbling CO2 through
the solution. The products are detected using NMR spectroscopy and gas chromatography
(GC). While photocatalytic CO2 reduction on InP and other III-V compound
semiconductors has been demonstrated previously,
81, 102
these materials corrode rapidly
under photo-electrochemical conditions and are significantly degraded after just 30
minutes of illumination.
93, 105
Several research groups have shown that by depositing thin
films of TiO2 on these unstable semiconductors, they can be protected from corrosion.
48, 49,
56, 93, 105
In order to make InP photochemically stable, we passivated the surface using a
536 534 532 530 528 526 524
Intensity (a.u.)
O
Binding Energy (eV)
980 970 960 950 940 930
Intensity (a.u.)
Experimental Data
Fitting Data
Cu Peak
CuO Peak
Binding Energy (eV)
Figure 5.2 (a) Cu 2p for CuO and Cu, (b) O 1s core level XPS spectra of air-oxide Cu nanoparticles.
80
3nm thin film of TiO2 deposited by ALD, as illustrated schematically in Figure 5.1a. Under
532nm illumination, the photocurrent of InP nanopillars with TiO2-passivation is stable for
at least 12 hours, as shown in Figure 5.3.
5.3 Results and discussion
Figure 5.4a shows the photocurrent-voltage curves measured in a 0.5M KCl
solution under 532nm illumination for bare InP nanopillars and InP nanopillars passivated
with TiO2. With TiO2-passivation, the photocurrent is substantially increased for all applied
potentials. The bare InP nanopillar sample (blue curve) has an onset of photocurrent at a
potential of approximately 0.03V (vs NHE), as indicated in Figure 5.4b. For TiO2-
passivated InP (red curve), we observe a clear shift in the onset potential by about 0.1V in
0 2 4 6 8 10 12
0
1
2
3
4
5
6
7
8
Time (hours)
Current Density (mA/cm
2
)
Figure 5.3 Time dependence of the photocurrent density of InP with 3nm TiO 2 and Cu nanoparticles
illuminated with 532nm light at an applied overpotential of -0.6V vs. NHE.
81
these photo-I-V characteristics. We attribute this, in part, to a pn-junction formed between
the TiO2, which is n-type due to oxygen vacancies, and the p-type InP.
36, 93
This pn-junction
creates a built-in potential that assists in the separation of photogenerated electron-hole
pairs and results in a shift of the onset potential of this reaction. Figure 5.4c shows the
methanol peaks observed in the NMR spectra of bare InP nanopillars and TiO2-passivated
InP nanopillars after 12-hour reactions at an applied potential of -0.6V vs. NHE under
532nm illumination. If we consider the one electron reduction of CO2 to CO2
-
as the first
step on this reaction, this is 1.3V below the E
o
(CO2/CO2
-
)=-1.9eV standard redox potential.
Figure 5.4d shows the Faraday efficiencies of methanol (i.e., the selectivity of methanol)
for these two types of samples. This figure indicates that the TiO2-passivation layer not
only enhances the overall photon-conversion efficiency, but also increases the selectivity
of methanol from H2 and other hydrocarbons from 0.85% to 4.8%. We attribute the
increased methanol selectivity to oxygen vacancies that are inherent to these thin TiO2 films,
which provide the catalytically active sites for CO2 reduction.
106-108
Peaks corresponding
to Ti
3+
states (i.e., O-vacancies) are observed in the XPS spectra taken from TiO2 deposited
on InP, as shown in Figure 5.5.
82
-0.15 -0.10 -0.05 0.00 0.05 0.10 0.15
0.1
1
Bare InP
InP w/TiO
2
Potential vs. NHE (V)
Absolute Current Density (mA/cm
2
)
(b)
0
1
2
3
4
5
6
4.79%
InP with TiO
2
Faraday Efficency of Methanol (%)
Bare InP
0.846%
(d)
3.50 3.45 3.40 3.35 3.30
Bare InP
InP w/TiO
2
ppm
Methanol
(c)
-1.0 -0.8 -0.6 -0.4 -0.2 0.0 0.2 0.4
-30
-25
-20
-15
-10
-5
0
Dark
Bare InP
InP w/TiO
2
Potential vs. NHE(V)
Current Density (mA/cm
2
)
(a)
Figure 5.4 (a) Photocatalytic current-potential curves where the dashed line indicates the potential
applied during the methanol test. (b) Log plot of photocatalytic current-potential curves. (c)
Methanol peak in NMR spectra and (d) Faraday efficiencies of methanol production for InP
nanopillars with and without 3nm TiO 2 under 532nm illumination in a CO 2 saturated 0.5M KCl
solution.
468 466 464 462 460 458 456 454
Intensity (a.u.)
Binding Energy (eV)
Ti
3+
Experimatal data
Fitting data
Figure 5.5 Ti 2p level XPS spectra of TiO 2 on InP, which shows the presence of Ti
3+
states.
83
In order to verify the role of these oxygen vacancies in the photocatalytic reduction
of CO2, all plane wave density functional theory (PW-DFT) calculations were performed
with the Quantum Espresso package using the most recently available ultrasoft
pseudopotentials with scalar relativistic corrections,
109-112
and spin-unrestricted
calculations were done employing the Perdew-Burke-Ernzerhof (PBE) functional.
113
Clean,
stoichiometric anatase and defective anatase with an oxygen vacancy were both
investigated in this study.
In all plane wave density functional theory (PW-DFT) calculations, large kinetic
energy cutoffs of 435.2 eV and 4.352 eV were applied to the wave functions and charge
density, respectively. A 1x1x1 Monkhorst-Pack k-point grid was used for all calculations.
The anatase slab was modeled as a 3x3 cell with lattice constants of a = 3.7845 Å, c =
9.5143 Å from experimental crystallographic data.
114
The appropriate cuts were made to
construct the most stable and dominant facet (>94%) of the anatase crystal, the 101
surface,
115
as a 16 TiO2 cell, the optimal size for molecular deposition of small molecules
such as water.
116
A vacuum gap of about 9 Å (a factor of 2.5 greater than the height of the
cell) was used. A vacuum gap of about 11 Å (a factor of 2.5 greater than the height of the
cell) was used. The effect of nonlocal forces such as dispersion forces was accounted for
using Grimme’s method,
4
which correctly predicted the thermodynamic stability of the
phases of TiO2 (rutile > brookite > anatase).
5
Differences in adsorption energies and Bader
charges between our calculations and Sorescu, et al.’s study may be accounted for by our
84
different methods (their dispersion method was developed by Tkatchenko and Scheffler),
cell sizes (doubled in the x and y-axes), and U (their U = 3.5 eV, our U = 3.6 eV). We
attributed our higher energies to interaction across the periodic cells.
6
CO2
CO2ˉ
(b)
(a)
Figure 5.6 PW-DFT calculated structure for anatase TiO 2 with O vacancies (a) before CO 2
adsorption and (b) after CO 2 adsorption and relaxation.
Figure 5.7 PW-DFT result of neutral CO 2 adsorbed to the stoichiometric anatase 101 surface with
a binding energy of -0.48 eV.
85
The DFT+U approach was adopted in order to recover these highly localized states
with a self-consistently computed U term of 3.6 eV applied to Ti atoms, well within Finazzi
et. al.’s suggested range of 3-4 eV for anatase.
117, 118
Although one study found the oxygen
vacancy to be present deep in slabs of varying size,
119
others were unable to determine the
most stable vacancy structure.
36
Our DFT+U calculations identified the oxygen vacancy to
be at the surface, specifically, the two-fold coordinated bridging oxygen (see Figure 5.6),
as energetically favorable compared to other sites. Adsorption energies were calculated by
subtracting the two components (molecule and surface) from the adsorbed system:
Eads=E[surf+molecule]-E[surf]-E[molecule]
where the molecule was CO2, or CO2
-
and the surface, stoichiometric anatase or defective
anatase with a surface oxygen vacancy. A thorough, manual search was performed to
determine the global minimum of adsorbed molecules to the anatase support. Adsorption
energies were favorable between neutral CO2 adsorbed to stoichiometric (-0.48 eV) and
defective anatase (-0.94 eV). The global minima found reproduced the geometries of
Sorescu, et al.’s PW-DFT study of the adsorption of CO2 on anatase.
34
The difference in
adsorption energies to Sorescu, et al.’s results may be attributed to our smaller cell and
computational parameter. In our study, we investigated alternative roles of the anatase
support such as stabilization of CO
2-
intermediate. However, the adsorption of CO2
-
to
stoichiometric and defective anatase resulted in repulsive, unstable systems, requiring
thermodynamically unfavorable energies of 4.39 eV and 2.57 eV , respectively, to form. The
86
chemical bonding analysis obtained using the Bader charge localization scheme is shown
in Figure 5.7.
120-122
In Figure 5.6b, the linear CO2 molecule becomes bent upon adsorption
to the defective anatase support, its C effectively filling the bridging oxygen vacancy.
Moreover, CO2 gains an electron (-0.897e) spontaneously from the TiO2 support. This
calculation indicates that the O vacancies provide active sites for CO2 absorption, and no
overpotential is required to form the CO2
-
intermediate. In fact, 4 of the 8 minima found by
Sorescu, et. al. formed the CO
2-
intermediate.
34
These CO2
-
intermediates then react with
H2O to form methanol. Therefore, this thin TiO2 film can not only protect InP from
photocorrosion, but also increase the methanol yields without applying high overpotentials.
In order to further improve the selectivity of methanol with respect to H2 evolution
in aqueous solution, we deposited copper nanoparticles on the TiO2-passivated InP
nanopillars. Cu and its oxide are known catalysts for CO2 reduction by lowering the energy
barriers of intermediate states in the reaction
123, 124
and have been intensively studied in the
electrochemical reduction of CO2.
125, 126
Figure 5.8a shows the photo-I-V characteristics of
InP nanopillars with copper nanoparticles (green curve) and TiO2-passivated InP
nanopillars with copper (purple line). From this Figure, we find that the addition of Cu
nanoparticles to the bare InP NPLs (green curve) does not change the photo-I-V
characteristics. However, it does increase the Faraday efficiency of methanol from 0.85%
to 2.8%, indicating that copper is an active catalyst for CO2 reduction to methanol. By
deposition of copper nanoparticles on the TiO2-passivated InP nanopillars, we successfully
87
increase the selectivity to 8.7%, as shown in Figure 5.8b. This improvement in the
selectivity likely results from the interface of the copper nanoparticles and the TiO2-
passivation layer. For Cu as a metal catalyst (and likely CuO), there have been several
previous works reporting that it has a moderate hydrogen overvoltage and weak CO2
adsorption characteristics, and that it can facilitate the reaction of CO with H2 to generate
hydrocarbons, aldehydes, and alcohols as major products.
125, 127
We believe it is the oxygen
vacancies in the TiO2 and Cu together, which promote the selectivity of methanol over H2
production, however a detailed reaction mechanism is still under investigation.
-1.0 -0.8 -0.6 -0.4 -0.2 0.0 0.2 0.4
-35
-30
-25
-20
-15
-10
-5
0
Dark
Bare InP
InP w/Cu
InP w/ TiO
2
&Cu
Potential vs. NHE (V)
Current density (mA/cm
2
)
(a)
(b)
0
2
4
6
8
InP with Cu InP with TiO
2
&Cu
Faraday Efficency of Methanol (%)
Bare InP
Figure 5.8 (a) Photocatalytic current-potential curves and (b) Faraday efficiencies of methanol
production for samples of bare InP, InP/Cu, and InP/TiO 2/Cu under 532nm illumination in a CO 2
saturated 0.5M KCl solution for a12-hour reaction.
88
The previous work of Lee et al. showed that these nanotextured InP photo-cathodes
exhibit about a 40% enhancement in the photoconversion efficiency compared to planar
surfaces in photocatalytic water splitting.
56
In the work presented here, we also observed
an enhancement by a factor of 2 in total conversion efficiency when comparing InP
nanopillars to planar InP. The Faraday efficiencies of methanol production, however, are
around 1% for both planar and nanopillar InP, and, thus, the selectivity of methanol with
respect to H2 is almost the same, as listed in Table 5.1 in the Supporting Document. By
depositing Cu nanoparticles on TiO2-passivated InP nanopillars, the Faraday efficiency
increases to 8.7%. The H2 yields produced by these samples were measured by GC after
12 hours of illumination with 532nm wavelength light under an applied potential of -0.6V
vs. NHE, as shown in Figure 5.9. Table 5.2 lists the H2 yields from these GC measurements,
methanol yields from NMR spectroscopy, and the total integrated charge running through
the device. For 3nm TiO2-passivated InP nanopillars, the Faraday efficiency is 83.6% in
-0.6V vs.NHE
12hr Reaction
Methanol (mol) Q(C) Faraday Efficiency
Planar 2.76E-8 1.414 1.13%
Nanopillar 5.0E-8 3.39 0.864%
Table 5.1. Faraday efficiencies of methanol of 12-hour reaction on InP/TiO 2/Cu sample under
532nm illumination in a CO 2 saturated 0.5M KCl solution.
89
total and 76.9% for H2 evolution. This indicates that, in aqueous solution, simultaneous H2
evolution is an inevitable competing reaction with a lower energy barrier than CO2
reduction. In order to suppress the H2 competing reaction, a non-aqueous solution should
be used in future studies.
1.6 1.8 2.0 2.2 2.4 2.6 2.8 3.0 3.2
H
2
CO
2
Re tention Time (min)
N2
Figure 5.9 Gas chromatograph (GC) data taken after 12-hour reaction on InP/TiO 2/Cu sample under
532nm illumination in a CO 2 saturated 0.5M KCl solution.
3nm TiO2 & Cu H2 CH3OH H2 & CH3OH
Faraday Efficiency 76.9% 8.7%
85.6%
Table 5.2. Faraday efficiencies of hydrogen and methanol of 12-hour reaction on InP/TiO 2/Cu
sample under 532nm illumination in a CO 2 saturated 0.5M KCl solution.
90
5.4 Conclusion
In summary, photocatalytic CO2 reduction with water to produce methanol is
observed using TiO2-passivated InP nanopillar photocathodes under 532nm wavelength
illumination. In addition to providing a stable photocatalytic surface, the TiO2-passivation
provides substantial enhancement in the photoconversion efficiency through the
introduction of O vacancies associated with the non-stoichiometric growth of TiO2 by
atomic layer deposition. The role of these oxygen vacancies as catalytically active sites in
the photocatalytic reduction of CO2 is established by PW-DFT calculations, which indicate
that CO2 binds stably to these oxygen vacancies and gains an electron (-0.897e)
spontaneously from the TiO2 support. Therefore, no externally applied overpotential is
required to form the CO2
-
intermediate, which can subsequently react with H2O to form
methanol. Copper nanoparticles deposited on the TiO2 act as a cocatalyst, which further
improves the selectivity and yield of methanol production by up to 8 fold giving a Faraday
efficiency of 8.7%.
91
Chapter 6: A microscopic study of atomic layer deposition of TiO2
on GaAs and its photocatalytic application
6.1 Introducton
Since Fujishima and Honda initially demonstrated photocatalytic water splitting
using TiO2 in 1972,
3
the study of photocatalysis has received worldwide interest due to its
potential applications in solar fuel generation either through water splitting or CO2
reduction. While TiO2 is a widely studied material, its theoretical photocurrent density
under one-sun illumination is only about 1-2mA/cm
2
. Narrower band gap materials like Si,
InP, and GaAs are more promising candidates, since their theoretical photocurrent densities
are 44mA/cm
2
, 35mA/cm
2
, and 32mA/cm
2
, respectively.
1
However, one of the primary
problems, which prevents these narrower band gap materials from being utilized as
photocatalysts, is that their surfaces are not photochemically stable.
14
Therefore,
passivating their surfaces to protect them from photocorrosion without sacrificing their
photoconversion efficiency will enable more efficient photocatalytic systems to be
developed. Recently, several groups have demonstrated that a very thin layer of TiO2 can
be used to make these narrower band gap materials stable. Chen et al. first showed that
TiO2 deposited by atomic layer deposition (ALD) can stabilize silicon photoanodes for
water oxidation.
49
Hu et al. also reported that amorphous TiO2 coatings stabilize Si, GaAs,
and GaP photoanodes for efficient water oxidation.
48
For the study of photocathodes, it is
92
also reported that TiO2 can be used to protect p-InP from corrosion and provide efficient
solar-driven H2 production.
56
In contrast, our group’s previous work showed that a thin
layer of TiO2 can not only make GaP photocathodes photochemically stable, but also
provide substantial enhancement in the photoconversion efficiency for both water splitting
and CO2 reduction.
93, 105
More recently, Lin et al. has also discussed similar enhancement
in photocatalytic H2 evolution using InP using TiO2 films deposited with various ALD
precursors, including TDMAT (i.e., [(CH3)2N]4Ti) and titanium isopropoxide (i.e.,
C12H28O4Ti).
52
While these previous works are very encouraging, a detailed microscopic
picture of these TiO2-passivated heterostructures is lacking.
In this report, we provide a detailed study of TiO2 layer growth on GaAs by atomic
layer deposition using high resolution transmission electron microscopy (HRTEM) and
electron energy-loss spectroscopy (EELS), as well as plane wave density functional theory
(PW-DFT) calculations, which are carried out to provide a detailed picture of the
GaAs/TiO2 interface and explore the role of surface binding of reactants and intermediate
species to the TiO2 surface.
6.2 Experimental procedure
In this work, p-type (111) oriented GaAs substrates with a Zn dopant concentration
of 6×10
16
cm
-3
(obtained from University Wafer, Inc.) are used as the photocathode. Atomic
layer deposition of TiO2 was performed at 250
o
C on the p-GaAs wafers using TiCl4 as the
93
titanium source and water vapor as the oxygen source. TiCl4 is used for the first half-cycle
and argon was the carrier gas with a flow rate of 20sccm during deposition. High resolution
transmission electron microscope images of TiO2 on GaAs were taken with a JEOL JEM-
2100F TEM equipped with a Gatan Quantum SE GIF quantum energy filter. A layer of Pt
was deposited on the surfaces of samples as part of the cross section sample preparation
using standard focused ion beam sample preparation techniques. The photocatalytic water
splitting measurements were carried out in a pH=0 solution of 0.5 M H2SO4 using a three-
terminal potentiostat with the prepared samples, a Ag/AgCl electrode, and a graphite
electrode functioning as the working, reference, and counter electrodes, respectively. The
area of the exposed electrode surface is about 0.1 cm
2
. A 532nm continuous-wave laser
was used to excite the samples and low to moderate power excitation (910 mW/cm
2
) was
used to avoid optical heating. The non-aqueous solution (for CO2 reduction) was prepared
using acetonitrile (AcN, 99.99%), dimethyl sulfoxide (DMSO, 99.96% D, Sigma Aldrich),
and 1-Ethyl-3-methylimidazolium tetrafluoroborate ([EMIM]BF4, 99.0%, HPLC, Sigma
Aldrich). An anion exchange membrane (Selemion AMV , Anion Exchange Membrane,
AGC Inc.) was used to separate the working and counter electrodes to prevent oxidation of
the reduced CO2 products. This anion exchange membrane only allows negative ions to
transfer through the membrane, which prevents the oxylate intermediate from participating
in the oxygen evolution half-reaction at the counter electrode. A three-terminal potentiostat
was used with the prepared semiconductor samples as the working electrode, a Ag/AgNO3
94
reference electrode, and a Pt wire as the counter electrode. The Ag/AgNO3 electrode was
made of a silver wire immersed in 0.01 M silver nitrate dissolved in 0.1M TEAP/AcN.
Also, the reference electrode was calibrated against a ferrocene/ferrocenium (Fc
+
/Fc) redox
couple to confirm that it gave the right potential with respect to NHE. Before each
measurement, CO2 was purged through the solution on the working electrode side for 30
minutes. H2 evolution and the gaseous products from the CO2 reduction reaction were
analyzed with a gas chromatograph (GC) (Bruker GC-450) equipped with a thermal
conductivity detector (TCD) and a carbon-plot column (Agilent).
6.3 Results and discussion
Figures 6.2a, 6.2b, and 6.2c show HRTEM images of TiO2 films deposited with 25
cycles, 75 cycles, and 500 cycles of atomic layer deposition on GaAs. These films have
nominal thicknesses of 1nm, 3nm, and 15nm, respectively. Figures 6.2d, 6.2e, and 6.2f
show EELS spatial profiles of the Ti and O species for the 25, 75, and 500 cycle TiO2 on
GaAs samples, respectively. In Figure 6.2d, the oxygen signal increases approximately
0.5nm before the Ti signal, indicating the presence of the GaAs native oxide underneath
the TiO2 layer. Therefore, based on these EELS spectra, there is both native GaAs oxide
and amorphous TiO2 on the GaAs surface after the 25 cycle ALD process. It should be
noted that the thickness of GaAs’s native oxide is substantially lower than the 3nm native
oxide observed on bare GaAs (without TiO2 deposition). After 25 ALD cycles, the native
95
oxide of the GaAs was partially removed due to the Cl
-
ions from the TiCl4 precursor.
Ohmic back contacts were made to the p-GaAs by evaporating 5nm of Ti followed by 50nm
of Au. The Ti-Au film was then connected to the external circuitry with a copper wire and
coated with epoxy cement to insulate it from the electrolytic solution. Figure 6.2e shows
the EELS spatial maps of GaAs with 75 cycles of atomic layer deposition of TiO2. Here,
both the O and Ti signals increase together, indicating that the native oxide of GaAs has
been completely removed during the ALD process. With further growth of 500 cycles, the
amorphous TiO2 becomes crystalline as shown in Figure 6.2f. This layer of TiO2 is in the
anatase phase with clear lattice planes along the 101 direction, identified by the inter-plane
distance of 3.5±0.1Å, as indicated in Figure 6.1. Shi’s work also shows anatase phase TiO2
formed by high temperature atomic layer deposition.
128
Figure 6.1 TEM Image of 15nm anatase TiO 2 crystalline on GaAs.
96
Figure 6.3a shows the water splitting photocurrent-voltage curves of p-GaAs
photocathodes with various thicknesses of TiO2 under 910 mW/cm
2
532nm illumination in
a 0.5M H2SO4 electrolyte, plotted together with bare GaAs. The bare GaAs (blue curve)
exhibits an onset of photocurrent at a potential of approximately -0.05V vs. RHE. The
TiO2-passivated GaAs shows a clear shift in the onset potential. The sample with a nominal
thickness of 1nm TiO2 exhibits the most prominent shift, lowering the overpotential (i.e.,
increasing the onset potential) by approximately 180mV (red curve) with a 120X-fold
Figure 6.2 The high resolution TEM image of TiO 2 films deposited with (a) 25 cycles, (b) 75 cycles,
and (c) 500 cycles of atomic layer deposition on GaAs. Note that for TEM sample preparation and
imaging purposes, the surface of the substrate is coated with a thick layer of Pt. (d-e) EELS spatial
profiles of Ti L edge (green line) and O K edge map (red line) for the 25, 75, and 500 cycle TiO 2 on
GaAs samples, respectively.
97
enhancement of photocurrent over bare GaAs at 0V vs RHE. These results were reproduced
consistently in several different sets of samples. The enhancement is attributed to the
oxygen vacancies (Ti
3+
states), which provide the photocatalytically active site for water
splitting, which is consistent with the results from PW-DFT calculations. At the same time,
we observe no improvement when the TiO2 thickness reaches 15nm, as shown by the
orange curve. This is because amorphous TiO2 is far more conductive than crystalline TiO2,
which is insulating and, therefore, impedes charge transfer to the ions in solution. Figure
6.3b shows the photocurrent-voltage curves for 10nm TiO2 on GaAs before and after
annealing. The photocurrent decreases dramatically after annealing, due to the
crystallization of the TiO2 during annealing, which is consistent with the result of the 15nm
crystalline TiO2 film on GaAs. The H2 product was detected using gas chromatography, as
shown in Figure 6.5a.
-0.6 -0.4 -0.2 0.0 0.2 0.4 0.6
-120
-100
-80
-60
-40
-20
0
Dark
Bare GaAs
GaAs w/1nm TiO
2
GaAs w/3nm TiO
2
GaAs w/10nm TiO
2
GaAs w/15nm TiO
2
Potential vs. RHE (V)
Current Density (mA/cm
2
)
(a)
-0.6 -0.4 -0.2 0.0 0.2
-60
-40
-20
0
Dark
GaAs w/10nm TiO
2
before annealing
GaAs w/10nm TiO
2
after annealing
Potential vs. RHE (V)
Current Density (mA/cm
2
)
(b)
Figure 6.3 Photocatalytic water splitting current-potential curves for samples of (a) GaAs
photocatalysts with various thicknesses of TiO 2 and (b) GaAs with 10nm TiO 2 before and after
annealing under 532nm illumination in a 0.5M H 2SO 4 pH=0 solution.
98
-1.5 -1.0 -0.5 0.0 0.5
-6
-5
-4
-3
-2
-1
0
Bare GaAs
GaAs w/1nm TiO
2
GaAs w/3nm TiO
2
GaAs w/10nm TiO
2
GaAs w/15nm TiO
2
Potential vs. RHE (V)
Current Density (mA/cm
2
)
(a)
-1.5 -1.0 -0.5 0.0 0.5
-5
-4
-3
-2
-1
0
GaAs w/10nm TiO
2
before annealing
GaAs w/10nm TiO
2
after annealing
Potential vs. RHE (V)
Current Density (mA/cm
2
)
(b)
Figure 6.4 Photocatalytic CO 2 reduction current-potential curves for samples of (a) GaAs
photocatalysts with various thicknesses of TiO 2 and (b) GaAs with 10nm TiO 2 before and after
annealing under 532nm illumination in a 0.02M [EMIM]BF 4 non-aqueous CO 2 saturated electrolyte.
2.0 2.1 2.2 2.3
GC Signal (a.u.)
H
2
N
2
and O
2
Retention Time (min)
(a)
1.90 1.95 2.00
Retention Time (min)
GC Signal (a.u.)
N
2
and O
2
CO
(b)
Figure 6.5 Gas chromatograph (GC) data taken for reactions on GaAs with 3nm TiO 2 under 532nm
illumination (a) in a 0.5M H 2SO 4 solution and (b) in a 0.02M [EMIM]BF 4 non-aqueous CO 2
saturated electrolyte.
99
In addition to water splitting, we also characterized the photocatalytic performance
of these GaAs/TiO2 heterostructures for CO2 reduction to CO. A non-aqueous ionic liquid
electrolyte was used to exclude the effect of H2O splitting, which has a lower onset
potential than CO2 reduction and tends to dominate the Faradaic efficiency. Figure 6.4a
shows the photocurrent-voltage curves for GaAs passivated with various thicknesses of
TiO2 measured in an AcN electrolyte with 0.02M [EMIM]BF4 under 532nm wavelength
illumination. During these measurements, CO2 is continuously bubbled through the
solution. At any given voltage, the TiO2-passivated GaAs samples have a higher
photocurrent than the bare GaAs (blue curve). Also, clear shifts in the onset potential
required to initiate this reaction can also be seen in the I-V characteristics of samples
passivated with TiO2, which is consistent with the predictions of the PW-DFT calculations
discussed below. Again, we observed no improvement when the TiO2 thickness reaches
15nm, as shown in the orange curve, due to its insulating crystalline nature. Figure 6.4b
shows the photocurrent-voltage curves for 10nm TiO2 on GaAs before and after annealing.
The photocurrent decreases dramatically after annealing, which is due to the crystallization
of TiO2 during annealing. The gaseous products from the CO2 reduction reaction were
analyzed using gas chromatography, as shown in Figure 6.5b
It is widely accepted that the reduction of H2O to form H2 occurs predominantly at
the Ti
3+
-O vacancy sites due to the reaction of adsorbed H
+
ions.
59, 60
Thus, a higher
concentration of O vacancies correspond to more active sites, resulting in a higher
100
hydrogen generation efficiency. Similarly, these Ti
3+
-O vacancy sites can also provide the
catalytically active sites for CO2 reduction.
106-108
As was previously done,
129
plane wave
density functional theory (PW-DFT) calculations were performed on a supercell of anatase
composed of 16 TiO2 units exposing the 101 surface. The Quantum Espresso package was
used with the most recently available ultrasoft pseudopotentials with scalar relativistic
corrections
109-112
and spin-unrestricted calculations were done employing the Perdew-
Burke-Ernzerhof (PBE) functional.
113
Kinetic energy cutoffs of 435.2 (4.352) eV were
applied to the wave functions (charge density) with a 1x1x1 Monkhorst-Pack k-point grid
centered at Γ. A DFT+D+U approach was instituted with Grimme’s method
130
for
dispersion forces (+D) and a Hubbard U parameter of 3.6 eV (+U). An explanation for this
approach was provided in our previous study.
129
Clean, stoichiometric anatase and
defective anatase with a surface oxygen vacancy were both investigated in this study.
Adsorption energies were calculated by subtracting energies of the two components (H2O
and surface) from the energy of the adsorbed system:
Eads=E[surf+H2O]-E[surf]-E[H2O]
where the surface was stoichiometric anatase or defective anatase with a surface oxygen
vacancy.
A manual search was performed to determine the global minimum of adsorbed H2O
to the anatase support reproducing the same geometries found in Tilocca et al.’s study on
water on anatase.
131
Adsorption energies were favorable between neutral H2O adsorbed to
101
stoichiometric (-1.26 eV as indicated in Figure 6.6) and defective anatase (-1.50 eV). These
energies differ from those of Tilocca et al. due to possible interactions across our smaller
supercell and the additional consideration of dispersion forces, which lowered the
adsorption to stoichiometric anatase significantly (Tilocca et al. reported -0.74 eV) and to
defective anatase minimally (Tilocca et al. reported -1.48 eV). As in our findings with CO2,
molecules adsorbed to defective anatase tend to fill the surface oxygen vacancy. In Figure
6.7b, the water molecule attempts to both fill the oxygen vacancy with its oxygen (colored
green) and retain a hydrogen-bond to the neighboring surface oxygen on the oxide support.
Moreover, the hydrogen-bond might be considered an “activated” bond with a length of
1.71 Å as compared to 1.89 Å on the stoichiometric support, indicating that the system is
approaching the proton transfer to the O atom of the support. In a molecular dynamics
study, Tilocca et al. estimated the activation barrier from adsorption of H2O to dissociation
of the H2O to two hydroxyls to be ~0.1 eV .
132
Moreover, Fujimori, et al. maintained that
hydroxylation of their MgO support provided the pathway for two mechanisms of
hydrogen evolution: the direct redox process and the water-gas shift reaction in the
presence of CO.
133
This indicates that Ti
3+
-O sites are more energetically favorable for H2O
adsorption, which results in higher H2 evolution efficiencies, consistent with previous
reports in the literature.
60
This calculation combined with our previous finding
129
of
spontaneous CO2 reduction on defective anatase (a charge transfer of 0.897 e from the
support to CO2) indicates that the O vacancies provide catalytically active sites for CO2
102
and H2O absorption, and no overpotential is required to form the CO2
-
intermediate.
Figure 6.6 H 2O adsorption on stoichiometric anatase with an adsorption energy of -1.26 eV.
103
6.4 Conclusion
In conclusion, we have carried out a microscopic study of GaAs photocathodes
passivated with various thicknesses of TiO2 using atomic layer deposition. High resolution
transmission electron microscopy (HRTEM) and electron energy-loss spectroscopy (EELS)
show that the native oxide of GaAs is removed by the TiCl4 precursor during the TiO2
Figure 6.7 (a) PW-DFT calculated structure for anatase TiO 2 with O vacancies before adsorption,
(b) after H 2O adsorption and relaxation and (c) after CO 2 adsorption and relaxation.
104
growth. This removal of the native oxide improves the photocatalytic performance by
facilitating charge transfer from the GaAs semiconductor to the ions in solution. The
photocatalytic performance for both water splitting and CO2 reduction of these
heterostructures show a very strong dependence on the thickness of the TiO2 over the range
of 0-15nm. Thinner TiO2 films are amorphous and show enhanced catalytic performance
with respect to bare GaAs, whereas thicker TiO2 films (15nm) are single crystal and have
poor charge transfer due to the insulating nature of crystalline TiO2. DFT calculations show
that the water and CO2 molecules bind stably to TiO2, which can further improve the
photocatalytic charge transfer process.
105
Chapter 7: Future work
7.1 III-V semiconductor nanowires for CO2 reduction
With amorphous thin TiO2 passivation, planar III-V semiconductors can not only
remain stable during photocatalytic reactions, but also achieve enhanced photoconversion
efficiency due to the existence of Ti
3+
active sites. In order to further enhance
photoconversion efficiency to reach the theoretical photocurrent density, there are two
major issues that should be addressed: (1) the absorption ratio of incident light and (2) the
inherent mismatch between the light absorption length and the minority carrier diffusion
length. The introduction of nanowire morphology could help to improve the performance
of existing photocatalytic materials. First of all, the nanowire morphology provides a large
surface area for light absorbing, which significantly decreases specular reflectance of
semiconductors comparing that of planar ones. The increase of roughness factor (the ratio
between surface area and the projected electrode area) reduces the required overpotential
by decreasing the surface flux of charge carriers.
1
Secondly, the nanowire morphology
helps to alleviate the mismatch between the minority carrier diffusion length and the light
absorption length by controlling the diameters of nanowire arrays. From theoretically
calculations, nanowires with a diameter of 300nm, separation of 600nm, and length of 2μm
are believed to produce near optimum efficiency. Recently, as a photocathode for
photochemical hydrogen production, nanowires have been extensively investigated.
134, 135
106
Compared with water splitting, there is inadequacy of research of nanowire photocathode
for CO2 reduction. Thus, it will be of great interest to study CO2 reduction using nanowires
of III-V semiconductors. As an example, GaAs nanowires are used in the following
research proposal.
7.2 Experimental procedure
It has been well demonstrated that periodic high aspect ratio GaAs nanowires with
widths in the range of 500-1000 nm can be produced my metal-assisted chemical etching
using Au catalyst films patterned with soft lithography by Li’s group.
2
Figure 7.1 shows an
array of vertically aligned GaAs nanowires ~3.5μm in length and ~600nm in diameter.
These nanorwires were produced from an Au mesh patterned with 600nm diameter
openings in a solution of H2SO4 over-saturated with KMnO4 at slightly elevated
temperature for 5min. With slightly tuning of this recipe, the optimum geometry of
nanowire arrays for photocatalysis will be achieved.
Figure 7.1 SEM images of high aspect ratio GaAs nanopillars produced from a 600 nm wide square
Au mesh pattern in H 2SO 4 and KMnO 4 solution at 40-45
o
C. (a) 30
o
tilted view at low magnification,
(b) 30
o
tilted view at high magnification, (c) cross-sectional showing the highly vertical nanopillar
array.
2
107
Alternatively, we can fabricate GaAs nanowires with photolithography and wet
chemical etching. First of all, a mask with hole diameter of 300nm and center to center
spacing 900nm will be designed as shown in Figure 7.2a, following with photolithography
to create one layer of photoresist with desired pattern on GaAs substrate. After that, wet
chemical etchant with certain ratio of H2SO4, H2O2 and H2O will be used to etch GaAs
substrate. With preliminary experiments, the etching rate of H2SO4, H2O2 and H2O (1:8:8)
has been demonstrated to be 0.0839μm/sec, as shown in Figure 7.2b. In order to decrease
the etching rate, different ratios of H2SO4, H2O2 and H2O have been studied. Table 6.1
shows the resulting depths for different ratios in 3min etching time. From this table, the
ratio of 1:8:40 or 1:8:50 may give us a reasonable time to produce desired GaAs nanowires
with length of 2μm.
Figure 7.2 (a) Mask design with hole diameter of 300nm and center to center spacing 900nm. (b)
Etching depth vs. etching time for GaAs with H 2SO 4, H 2O 2 and H 2O (1:8:8).
0 100 200 300 400
0
5
10
15
20
25
30
Depth ( m)
Etching Time (s)
(a) (b)
108
A protection thin layer of ALD TiO2 will be performed at 250
o
C on the GaAs
nanowires using TiCl4 as the titanium source and water vapor as the oxygen source. A three-
terminal potentiostat will be used with the prepared semiconductor samples as the working
electrode, a Ag/AgNO3 reference electrode, and a Pt wire as the counter electrode. The
non-aqueous solution for CO2 reduction will be prepared using acetonitrile, dimethyl
sulfoxide, and 1-ethyl-3-methylimidazolium tetrafluoroborate. An anion exchange
membrane (Selemion AMV , Anion Exchange Membrane, AGC Inc.) will be used to
separate the working and counter electrodes to prevent oxidation of the reduced CO2
products. Before each measurement, CO2 will be purged through the solution on the
working electrode side for 30 minutes. Gaseous products from the CO2 reduction reaction
will be analyzed with a gas chromatograph (GC) (Bruker GC-450) equipped with a thermal
conductivity detector (TCD) and a carbon-plot column (Agilent).
Ratio ) Depth (um)
1:8:10 11.213
1:8:20 7.818
1:8:30 4.174
1:8:40 3.681
1:8:50 3.028
Table 7.1. The resulting depths of GaAs holes with different ratio of H 2SO 4, H 2O 2 and H 2O in 3min
etching.
109
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Abstract (if available)
Abstract
In photocatalysis, light absorbed by semiconductors creates electron-hole pairs which can be used to drive electrochemical redox reactions. Since the first demonstration of photocatalytic water splitting using TiO₂ in 1972, the study of photocatalysis has been of worldwide interest due to its potential applications in solar fuel generation either through water splitting or CO₂ reduction. For photocatalytic materials, III-V semiconductors such as GaP, InP, and GaAs are promising candidates with theoretical maximum photocurrent densities of 9mA/cm², 35mA/cm², and 32mA/cm², respectively. These are significantly larger than that of the mostly widely studied photocatalytic material, TiO₂ (1.1mA/cm²). However, photocatalytic corrosion of these III-V semiconductors prevents them from being utilized as photocatalysts. We have developed a technology of passivating the surface of these III-V semiconductors with a thin layer of TiO₂, which protects them from corrosion. In addition, this thin layer of TiO₂ provides substantial enhancement in the overall photoconversion efficiency. ❧ The dissertation will start with a brief introduction of the fundamentals of photocatalysis on semiconductor surfaces, followed by the basic mechanisms of its applications to water splitting and CO₂ reduction. After that, the factors limiting the photocatalytic conversion efficiency are introduced. In the following chapters, we will report the achievements we have made both on the enhanced photocatalytic water splitting and CO₂ reduction processes. ❧ In Chapter 2, we demonstrate that a thin layer of n-type TiO₂ using atomic layer deposition (ALD) prevents corrosion of p-type GaP, as evidenced by atomic force microscopy and photoelectrochemical measurements. In addition, the TiO₂ passivation layer provides an enhancement in photoconversion efficiency through the formation of a charge separating pn-region. Plasmonic Au nanoparticles deposited on top of the TiO₂-passivated GaP further increases the photoconversion efficiency through local field enhancement. Finite difference time domain (FDTD) simulations of the electric field profiles in this photocatalytic heterostructure corroborate the experimental results. ❧ In Chapter 3, we demonstrate that thin (1-5nm) films of TiO₂ deposited by ALD on planar GaAs provide electrochemical stability and substantial improvements in the efficiency of photocatalytic water splitting. The native oxide of GaAs is removed during the ALD process and this TiO₂ passivation layer produces a shift in the onset potential by +0.4V and enhances the photocurrent by 32-fold over bare GaAs (at 0V vs. RHE), resulting in a peak photoconversion efficiency of 1.5% under AM1.5 G illumination. The possible enhancement mechanism by Ti³⁺ active sites provided by ALD TiO₂ is discussed. ❧ In Chapter 4, photocatalytic CO₂ reduction with water to produce methanol is demonstrated using TiO₂-passivated GaP photocathodes under 532nm wavelength illumination. In addition to providing a stable photocatalytic surface, the TiO₂-passivation provides substantial enhancement in the photoconversion efficiency. Two possible mechanisms of enhancement are discussed. One is the passivation of surface states, which cause non-radiative carrier recombination, and the other is pn-junction formation from the n-type TiO₂/p-type GaP junction, which creates a built-in field that assists in the separation of photogenerated electron-hole pairs, further reducing recombination. ❧ In Chapter 5, photocatalytic CO₂ reduction with water to produce methanol is demonstrated using TiO₂-passivated InP nanopillar photocathodes under 532nm wavelength illumination. In addition to providing a stable photocatalytic surface, the TiO₂-passivation layer provides substantial enhancement in the photoconversion efficiency through the introduction of O vacancies associated with the non-stoichiometric growth of TiO₂ by atomic layer deposition. Plane wave-density functional theory (PW-DFT) calculations confirm the role of oxygen vacancies in the TiO₂ surface, which serve as catalytically active sites in the CO₂ reduction process. PW-DFT shows that CO₂ binds stably to these oxygen vacancies and CO₂ gains an electron (-0.897e) spontaneously from the TiO₂ support. This calculation indicates that the O vacancies provide active sites for CO₂ absorption, and no overpotential is required to form the CO₂⁻ intermediate. ❧ In Chapter 6, we report a microscopic study of GaAs/TiO₂ heterojunctions using cross-sectional high resolution transmission electron microscopy and electron energy loss spectroscopy maps. The photocatalytic performance of these heterostructures shows a very strong dependence on the thickness of the TiO₂ over the range of 0-15nm. Thinner TiO₂ films (<10nm) are amorphous and show enhanced catalytic performance with respect to bare GaAs. Thicker TiO₂ films (15nm) are crystalline and have poor charge transfer due to their insulating nature, while thinner amorphous TiO₂ films are conducting. PW-DFT calculations show that water molecules and CO₂ molecules bind stably to defective TiO₂, which can further improve the photocatalytic charge transfer process. ❧ Finally, in Chapter 7, future work related to photocatalysis using TiO₂-passivated nanowires of III-V semiconductors is discussed.
Linked assets
University of Southern California Dissertations and Theses
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Qiu, Jing
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Core Title
Enhanced photocatalysis on titanium oxide passivated III-V semiconductors
School
Viterbi School of Engineering
Degree
Doctor of Philosophy
Degree Program
Materials Science
Publication Date
10/06/2015
Defense Date
10/06/2015
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University of Southern California
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CO₂ reduction,GaAs,gap,H₂ evolution,InP,OAI-PMH Harvest,photocatalysis,TiO₂
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jingq@usc.edu,qiujing1986@gmail.com
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CO₂ reduction
GaAs
gap
H₂ evolution
InP
photocatalysis
TiO₂