On the Synergism between Cu and Ni for Photocatalytic Hydrogen Production and their Potential as Substitutes of Noble Metals

Article abstract of DOI:10.1002/cctc.201600697

A series of Cu(OH)₂–Ni(OH)₂/P25 photocatalysts was prepared by co-deposition–precipitation (total metal loading ≈1 wt %) and their performance was evaluated for H₂ production. Among this series, the 0.8 Cu(OH)₂–0.2 Ni(OH)₂/P25 photocatalyst demonstrated very high H₂ production rates in 20 vol % ethanol/water and 5 vol % glycerol/water mixtures (10 and 22 mmol h^?1 g^?1, respectively). Detailed analyses based on reaction kinetics, photoluminescence, X-ray photoelectron spectroscopy (XPS), and charge carrier scavenging suggest that both working catalysts are composed of Cu and Ni metals in their active phases. Cu⁰ is produced directly by the transfer of electrons from the conduction band of TiO₂ to surface Cu(OH)₂ nanoclusters, whereas Ni⁰ is formed indirectly through a process of gradual dissolution of Ni(OH)₂ to yield aqueous Ni²⁺ owing to the acidic environment of the medium, followed by Ni²⁺ reduction by electrons from the TiO₂ conduction band. The high rates of H₂ production that match those obtained with noble metals can be explained owing to a considerably less negative ΔG^o of Cu oxide formation when compared with that of Ni oxide formation and higher work function of Ni than that of Cu.

Full text of DOI:10.1002/cctc.201600697

DOI: 10.1002/cctc.201600697
Full Papers
On the Synergism between Cu and Ni for Photocatalytic
Hydrogen Production and their Potential as Substitutes of
Noble Metals
Imran Majeed,[a] Muhammad Amtiaz Nadeem,*[b] Ejaz Hussain,[a] Geoffrey I. N. Waterhouse,[c]
Amin Badshah,[a] Azhar Iqbal,[a] Muhammad Arif Nadeem,*[a] and Hicham Idriss*[b]
A series of Cu(OH)2–Ni(OH)2/P25 photocatalysts was prepared
by co-deposition–precipitation (total metal loading ꢁ1 wt%)
and their performance was evaluated for H2 production.
Among this series, the 0.8Cu(OH)2–0.2Ni(OH)2/P25 photocata-
lyst demonstrated very high H2 production rates in 20 vol%
ethanol/water and 5 vol% glycerol/water mixtures (10 and
22 mmolhꢀ1 gꢀ1, respectively). Detailed analyses based on reac-
tion kinetics, photoluminescence, X-ray photoelectron spec-
troscopy (XPS), and charge carrier scavenging suggest that
both working catalysts are composed of Cu and Ni metals in
their active phases. Cu0 is produced directly by the transfer of
electrons from the conduction band of TiO2 to surface Cu(OH)2
nanoclusters, whereas Ni0 is formed indirectly through a pro-
cess of gradual dissolution of Ni(OH)2 to yield aqueous Ni2+
owing to the acidic environment of the medium, followed by
Ni2+ reduction by electrons from the TiO2 conduction band.
The high rates of H2 production that match those obtained
with noble metals can be explained owing to a considerably
less negative DGo of Cu oxide formation when compared with
that of Ni oxide formation[1] and higher work function of Ni
than that of Cu.[2]
Introduction
Hydrogen is widely regarded as the logical energy carrier to re-
place fossil fuels for electricity generation and transportation.
The development of the so-called “hydrogen economy” re-
quires new technologies for H2 generation, of which photoca-
talytic water splitting or alcohol photoreforming by using
direct sunlight are viewed as two of the most promising path-
ways. Many photocatalysts have been reported that are capa-
ble of producing hydrogen from water in the presence of sacri-
ficial agents, with M/TiO2 systems (M=Pd, Au, or Pt) generally
showing the best and most stable performances.[3] Photocatal-
ysis relies on capturing the energy of incident photons with
E>Eg, through excitation of electrons from the valence band
of a semiconductor into its conduction band. Electrons and
holes thus produced then drive oxidation and reduction reac-
tions on the semiconductor surface. TiO2 is an ideal support for
photocatalytic hydrogen production from water and alcohols
because of its high chemical stability and resistance to photo-
corrosion in aqueous media, as well as its low cost and well-
matched energy bands with water redox levels.[4] The addition
of noble metals such as platinum, gold, and palladium to TiO2
is needed for photocatalytic H2 production.[3,5] However, noble
metals are expensive, with low natural abundance, hence they
are not especially practical for the design and development of
industrial photocatalysts for hydrogen production. The identifi-
cation of alternative low-cost co-catalysts that enhance the
photocatalytic activity of TiO2 for hydrogen production is a -
priority.
Cu/TiO2 and Ni/TiO2 systems are particularly promising in
this regard and represent cost-effective and efficient photoca-
talyst systems for solar hydrogen production.[6] Yu and co-
workers deposited Ni(OH)2 nanoclusters on TiO2 by a simple
precipitation method and observed a hydrogen production
rate of 3.0 mmolgꢀ1 hꢀ1 in 25 vol% aqueous methanol under
UV excitation.[7] Hydrogen evolution was attributed to the
more positive redox potential of the Ni2+/Ni couple compared
with the conduction band of TiO2, which serves as the active
sites for the reduction of H+ to H2. The same group also fabri-
cated Cu(OH)2/TiO2 photocatalysts by using the same method,
and reported a hydrogen production rate of 3.4 mmolgꢀ1 hꢀ1
in a 0.09m ethylene glycol solution in water under UV irradia-
tion. The enhanced activity was attributed to the transfer of
electrons from the TiO2 conduction band to Cu(OH)2, which
then reduced H+ to H2. However, it is not clear if Cu(OH)2
[a] I. Majeed, E. Hussain, Prof. Dr. A. Badshah, Prof. Dr. A. Iqbal,
Prof. Dr. M. A. Nadeem
Catalysis and Nanomaterials Lab 27
Department of Chemistry
Quaid-i-Azam University
Islamabad 45320 (Pakistan)
E-mail: arifnadeem4u@gmail.com
[b] Dr. M. A. Nadeem, Prof. Dr. H. Idriss
SABIC-Corporate Research and Development (CRD) at KAUST
Thuwal 23955 (Saudi Arabia)
E-mail: nadeemmi@sabic.com
idrissh@sabic.com
[c] Dr. G. I. N. Waterhouse
School of Chemical Sciences
University of Auckland
Private Bag 92019, Auckland (New Zealand).
Supporting information and the ORCID identification number(s) for the
author(s) of this article can be found under http://dx.doi.org/10.1002/
cctc.201600697.
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acted as the cathodic site or Cu(OH)2 was reduced to Cu2O or
Cu0 under reaction conditions, which then acted as the site for
H2 evolution. In previous work, some of us have evaluated the
hydrogen production activities of a series of Ni/TiO2 photocata-
lysts in ethanol/water mixtures.[2a] The photocatalytic activity of
Ni/TiO2 photocatalysts was found to be highly dependent on
the Ni loading, with a loading of 0.38 wt% being optimal, af-
fording a rate of 24.3 mmolgꢀ1 hꢀ1 in 95 vol% aqueous ethanol
under UV irradiation. The high activity of the Ni/TiO2 photoca-
talyst was rationalized in terms of the high work function of
metallic Ni and the excellent Schottky junction formed at the
TiO2 surface. In the same work, 0.63 wt% NiO/TiO2 was tested.
A long induction period was observed before a constant hy-
drogen production rate was obtained, reaching about one
fourth of that of Ni/TiO2. This later was simply prepared by in
situ hydrogen production. We have explained this induction
period as being due to in situ reduction of a fraction of NiO by
hydrogen, produced in the initial stage of the reaction, to Ni0.
The Fermi level of Ni0 is below that of the conduction band of
TiO2 anatase. In another work, we have previously studied the
CuO/TiO2 system and found that activity was dependent on
the nominal CuO loading, with 1.25 wt% CuO being optimal
trate is converted to ammonia or nitrite over Pt/TiO2 or Cu/
TiO2, respectively.[15] Similarly, highly selective ammonia synthe-
sis from nitrate can be observed over Pd–Cu/TiO2 photocata-
lysts compared with their monometallic counterparts.[16]
Herein, we report the successful synthesis of a series of
Cu(OH)2–Ni(OH)2/TiO2 photocatalysts by a simple precipitation
method, and the subsequent application of these photocata-
lysts for hydrogen production in different alcohol/water sys-
tems under UV excitation. An oxidized Ni atom in contact with
a Cu atom may become reduced owing to a considerably less
negative DG8 of Cu oxide formation (ꢀ129 kJmolꢀ1) compared
with that of Ni oxide formation (ꢀ430 kJmolꢀ1).[1] This would
then increase hydrogen production because the work function
of Ni is higher than that of Cu.[2] Of particular interest were the
chemical and morphological evolution of the co-catalysts
under the reaction conditions, the synergistic effects of Cu and
Ni co-deposition for H2 evolution, and the effect of the Cu/Ni
ratio on H2 production rates. Results are expected to guide the
future development of efficient, low-cost semiconductor
photocatalysts for solar H2 production.
(H2 production rate=20.3 mmolgꢀ1 hꢀ1 in 80:20 EtOH/H2O).[8] Results and Discussion
Highly dispersed sub-monolayer CuII species on TiO2 surfaces,
Materials characterization
rather than supported CuO nanoparticles, were proposed as
the active site for hydrogen production. In another recent
work, some of us have investigated Au/Cu2O–TiO2 and found
that Cu+ reduction to Cu0 is behind the main catalytic activity.
The redox potentials of Cu+ and Cu2+ are within the band gap
of TiO2. These latter results further confirm the results seen for
the CuO/TiO2 system.[9]
Powder XRD patterns for all Cu(OH)2–Ni(OH)2/P25 photocata-
lysts are shown in Figure S1 in the Supporting Information,
and are dominated by peaks resulting from anatase and rutile.
No clear change in the anatase and rutile diffraction peak posi-
tions is seen, suggesting that the metal hydroxide species are
finely dispersed over the support. The low loadings would also
make identification of Ni(OH)2 and Cu(OH)2 difficult. The ana-
tase and rutile crystallite sizes determined from the powder
XRD data by using the Scherrer equation were 28 nm and
36 nm, respectively.
The enhancement in photocatalytic activity regarding Ni-
containing TiO2 photocatalyst systems is, however, more com-
plex. It is postulated that Ni(OH)2 is converted into NiO after
prolonged exposure to irradiation in the oxidizing environment
provided by the reaction mixture.[10] NiO cannot trap photo-
generated electrons because the conduction band of NiO is
more negative than that of TiO2. There is some controversy in
the literature about the valence and conduction band energies
of NiO.[11] Values in the ranges of ꢀ0.3 to ꢀ1.0 eV are reported
for the conduction band, and 2.4 to 4.3 eV for the valence
band (with respect to the standard hydrogen electrode
(SHE)).[12] Theoretical calculations indicate that H–H recombina-
tion on NiO (100) is exothermic, with an activation barrier of
Figure 1(a) shows the UV/Vis diffuse reflectance spectra
(DRS) for the different mono- and bimetallic photocatalysts,
whilst Figure 1(b) compares fresh and used 0.8Cu(OH)2–
0.2Ni(OH)2/TiO2. All catalysts showed intense absorption below
400 nm, owing to the P25 support (Eg =3.15 eV). All metal-con-
taining photocatalysts showed absorption in the visible region
owing to overlap between the TiO2 and metal hydroxide com-
plexes. The Ni3d and Cu3d states can mix with the TiO2 con-
duction band at the interface region and thus decrease the
band gap energy between the Ti3d and O2p states of TiO2.[17]
The absorption spectra of fresh bimetallic photocatalysts show
a broad band centered above 600 nm and a small absorption
shoulder at 450 nm. The absorption shoulder at 450 nm is
partly due to the interfacial charge transfer (IFCT) from TiO2
conduction band to M2+ cations (where M=Cu and Ni).[15,18]
The enhanced absorption above 600 nm can be assigned to
Ni2+ d–d (600–800 nm) and Cu2+ d–d (700-800 nm) transitions.
The increased absorption around 450–600 nm for the used
photocatalyst in Figure 1(b) is likely due to the formation of
Cu+ oxide species during photoreactions.[29]
only 40 kJmolꢀ1 [13]
High hydrogen production rates observed
.
on Ni(OH)2-containing catalysts can be rationalized in terms of
the ability of NiO to act as an adsorbed atomic hydrogen re-
combination site after H+ has accepted an electron directly
from the conduction band of TiO2.[10] Bimetallic heterogeneous
catalysts demonstrate excellent performance in many reac-
tions, offering increased activity and selectivity to the desired
products. Bimetallic nanoparticles deposited on TiO2 can create
surface electronic states and active sites different to those real-
ized in monometallic systems owing to synergistic physical
and chemical interactions amongst the different atoms and
phases.[14] For example, Pt–Cu/TiO2 photocatalysts can selec-
tively reduce nitrate to N2 under UV irradiation. In contrast, ni-
Figure S2 in the Supporting Information shows digital pho-
tographs of the as-prepared photocatalysts under UV irradia-
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Figure 2. TEM images of fresh 1.0Cu(OH)2/P25 photocatalyst.
difficult to discern metal hydroxide nanoparticles on the sur-
face of the TiO2 by using TEM owing to the similarity in the
atomic number of Ni, Cu, and Ti, among other factors (Fig-
ure 2a and b). It has been reported that hydroxide nanoparti-
cles on TiO2 support could only be identified with any certainty
at a metal loadings of greater than 5 wt% M(OH)2 (M=Ni,
Cu).[7,21] It is likely that the hydroxide nanoparticles are present
in a very small size (1–3 nm) as seen in the magnified TEM
image only (Figure 2b).[22] The presence of small metal hydrox-
ide particles suggest a strong interaction with the P25 support.
Energy-dispersive X-ray (EDX) analysis of the same photocata-
lyst indicates the presence of both Cu and Ni species (Fig-
ure S3 in the Supporting Information).
Figure 3 presents the Cu2p, Ni2p, Ti2p, C1s, and O1s XPS
spectra of fresh 0.8Cu(OH)2–0.2Ni(OH)2/P25 photocatalyst
whereas the Cu2p and Ni2p spectra of all fresh photocatalysts
are given in Figure S4 in the Supporting Information. Nitrogen
was the only other element detected with an overall atomic
percentage of less than 0.5%. The X-ray photoelectron spec-
troscopy (XPS) of the Cu2p region for different Cu oxidation
Figure 1. (a) UV/Vis spectra of P25, “as prepared” mono- and bimetallic pho-
tocatalysts as indicated; (b) UV/Vis of as prepared and used 0.8Cu(OH)2–
0.2Ni(OH)2/TiO2.
tion (20 mg of each of photocatalyst suspended in 20 vol%
ethanol/water mixtures with total volumes of 25 mL). For the
as-prepared photocatalysts, the light blue-green color in-
creased with an increase in Cu(OH)2 loading owing to the pres-
ence of octahedrally coordinated Cu2+ ions.[20] The grey color
of calcined 0.8Cu(OH)2–0.2Ni(OH)2/P25 catalyst indicates the
formation of metal oxides.[8] Under photoreaction conditions,
the color change increases gradually with the increase of the
amount of Cu(OH)2 whereas no change is observed for the
Ni(OH)2-containing catalysts. The latter observation might be
due to the fact that Ni2+ in Ni(OH)2 cannot be reduced under
the given conditions. Upon exposure to air, the change in
Cu(OH)2/P25 color reverts back to the original color, which is
most likely due to re-oxidation of Cu0.
TEM images for 1.0Cu(OH)2/P25 photocatalyst are shown in
Figure 2 at different magnifications. P25 is known to be com-
posed of small nearly spherical anatase crystallites of average
size 20–30 nm, and larger rutile crystallites with average size
40–60 nm, both of which are evident in the TEM images. It is
Figure 3. XPS results of fresh bimetallic 0.8Cu(OH)2–0.2Ni(OH)2/P25 photo-
catalyst.
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states has been studied thoroughly.[23] Typical XPS Cu2p3/2
values for Cu0 and Cu+ are observed between 932.3–932.5 eV.
making Cu0 and Cu+ difficult to distinguish, whereas that of
Cu2+ is 1 eV higher at about 933.6 eV. The XPS of Cu2p of
Cu2+ has characteristic satellites at about 6 eV above the Cu2p
peaks. It can be noted that copper is detected in its reduced
state (Cu0/Cu+) whereas nickel is detected as Ni(OH)2 in 0.98
and 0.19 at%, respectively (Table 1). Ni2p3/2 and Ni2p1/2 peaks
are observed at 855.6 eV and 873.2 eV, respectively. The bind-
ing energy position, absence of multiple splittings, and addi-
tional “shake-up” satellite features at approximately 6 eV above
the Ni2p peaks indicate that Ni is present as Ni(OH)2.
intensity was observed to increase with an increase in loading.
To understand the electron transfer rate, time-resolved PL was
conducted. We were anticipating a decrease in the PL lifetimes
of 1.0Cu(OH)2/P25 compared with P25 as a result of electron
transfer from the photo-excited P25 to Cu(OH)2. However, in-
stead we observed the same PL kinetics behavior for all three
samples. This strongly suggests that the electron transfer
occurs much faster than the instrument response function
ꢁ500 ps of our time-correlated single photon set-up. The mea-
sured PL kinetics was best fitted by the built-in single expo-
nential model (t=Aeꢀkt, where t, A, k, and t are the time con-
stant, amplitude in counts, first-order rate constant, and time,
respectively) suggesting a time constant of ꢁ1.24ꢂ0.006 ns.
Table 1. XPS analysis of fresh bimetallic 0.8Cu(OH)2–0.2Ni(OH)2/P25
photocatalyst.
Photoreactions
Chemical
composition
Core
level
Peak
position [eV]
FWHM
[eV]
at%
H2 production
Photocatalytic H2 production was evaluated by using either
ethanol or glycerol as sacrificial agents. Degussa P25 showed
negligible activity because of the rapid recombination of con-
duction band electrons and valence band holes and as a result
TiO2
TiO2
Ti2p3/2
Ti2p1/2
O1s
458.7
464.4
530
1.0
1.9
1.1
2.1
1.4
1.3
1.3
1.7
2.1
2.1
2.5
2.5
33.8
46.4
19.6
Metal oxide
Metal hydroxide
Carbon (adv.)
CꢀOꢀC (adv.)
CꢀO=C (adv.)
Cu/Cu2O
532.1
284.8
286.4
288.6
932.4
952.2
855.6
873.2
873.2
C1s
C
of the absence of H recombination centers for hydrogen evo-
lution.[10,24] Figure 5a shows the effect of the precipitation
Cu2p3/2
Cu2p1/2
Ni2p3/2
Ni2p1/2
Ni2p1/2
0.39
0.10
Ni(OH)2
To gain insight regarding the electron transfer dynamics,
steady-state and time-resolved photoluminescence (PL) of the
P25, 1.0Cu(OH)2/P25 and 1.0Ni(OH)2/P25 samples were con-
ducted and the results are presented in Figure 4a and b. Fig-
Figure 5. (a) Effect of NaOH concentration on H2 production with
0.5Cu(OH)2–0.5Ni(OH)2/P25 photocatalysts. (b) Effect of metal loading
amount on H2 production in 20 vol% ethanol/water mixtures by using
xCu(OH)2–yNi(OH)2/P25 photocatalysts where x=y.
agent (NaOH) concentration, during catalyst preparation, on
the hydrogen production rate when using 0.5Cu(OH)2–
0.5Ni(OH)2/P25 catalysts. The small observed dependence of
the hydrogen production rate on NaOH concentration might
be due to the better dispersion or agglomeration of metal hy-
droxides particles on the P25 surface. The effect of metal load-
ing concentration was studied by increasing the overall load-
ing amount from 0.5 to 3 wt% (Figure 5b).
Figure 4. (a) Steady-state PL and (b) time-resolved PL spectra following
excitation at 305 nm.
ure 4a demonstrates that 1.0Cu(OH)2/P25 exhibits PL approxi-
mately 2 times lower in intensity compared with P25. This is
a clear indication of photo-excited electron transfer from the
conduction band of the P25 to the Cu(OH)2. Conversely, P25
and 1.0NiOH)2/P25 exhibit similar intensities. The presence of
Ni may also passivate some surface defects, thus resulting in
a slight increase in the PL intensity. To confirm this, PL meas-
urements as a function of Ni(OH)2 loading were performed. PL
The optimum hydrogen production rate was obtained for
a nominal loading amount of 1 wt%. A slight decrease in activ-
ity at 1.5 wt% metal loading was observed and after that
a sharp decrease in activity was observed from 2 to 3% load-
ings. The decrease in photocatalytic activity with increasing
metal loading has been observed by many other groups, how-
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ever, no clear explanation is yet known.[8,25] Upon increasing of
metal particle coverage, a larger fraction of the semiconductor
surface may become unavailable for light adsorption. It can
also result in increased surface defects at the metal–semicon-
ductor interface, leading to an increase in electron–hole re-
combination centers. For further study, the overall nominal
metal loading was kept at 1 wt% and 0.5m NaOH solution was
used as the precipitating agent.
Comparison of hydrogen production in 20 vol% ethanol/
water mixture with various samples is shown in Figure 6. Hy-
drogen production experiments in different ethanol to water
volume ratios with the 0.8Cu(OH)2–0.2Ni(OH)2/P25 catalyst
Figure 7. Comparison of the photocatalytic activities of 0.8Cu(OH)2–
0.2Ni(OH)2/P25, 0.2Cu(OH)2–0.8Ni(OH)2/P25, 0.5Cu(OH)2–0.5Ni(OH)2/P25,
1.0Cu(OH)2/P25, and 1.0Ni(OH)2/P25 samples for the photocatalytic H2
production in 5 vol% glycerol/water mixtures.
larger numbers of a-H atoms, and hydroxyl groups deliver
higher hydrogen production rates.[26] The oxidation potential,
number of a-H atoms, and hydroxyl groups are 0.004 V, 5, and
3 for glycerol whereas they are 0.084 V, 2, and 1 for ethanol,
respectively.[26] In alcohol/water mixtures, the photoreaction is
initiated by the reaction of an alcohol with a hole in the TiO2
valence band to form a-hydroxyl radicals. The formation of a-
hydroxyl radicals have been identified previously on TiO2 by
using EPR and has been proposed by many others.[27] Lower
hydrogen production rates were observed for glycerol concen-
trations greater than 30 vol% (data not shown) owing to an in-
crease in reaction mixture viscosity. Inspection of the hydrogen
production with time indicates an induction period in both
cases. Two particular points are, however, worth mentioning.
First, Ni(OH)2 catalysts had the longer induction period in both
cases, also Ni-containing catalysts showed longer induction pe-
riods, which increased with concentration. Second, glycerol
and water mixtures showed much shorter induction periods
compared with that of ethanol and water mixtures. In previous
work,[26] we have shown a relationship between hydrogen pro-
duction rates and the exponential of redox potential, by fol-
lowing Marcus theory.[28] The rates for glycerol were about
twice as fast compared with those for ethanol reactions. The
results reported in this study on Cu and Ni are thus not too
different, indicating that the main drive for the reaction (at
least on TiO2) is the difference in the oxidation potential of the
donor molecules and the valence band of the semiconductor.
Table 2 presents a compilation of results from other work to
compare with the ones presented here. It can be noted that
better hydrogen production rates are obtained with the bimet-
allic composition used in this study compared with mono-
metallic catalysts reported in literature.
Figure 6. Comparison of the photocatalytic activities of 0.8Cu(OH)2–
0.2Ni(OH)2/P25, 1.0Cu(OH)2/P25, 0.5Cu(OH)2–0.5Ni(OH)2/P25, 0.2Cu(OH)2–
0.8Ni(OH)2/P25, 1.0Ni(OH)2/P25, and 0.8Cu(OH)2–0.2Ni(OH)2/P25 (calcined)
samples for the photocatalytic H2 production in 20 vol% ethanol/water
mixture.
were also performed. An increase in rates with increasing etha-
nol proportion, was noticed (Figure S5 in the Supporting Infor-
mation). Further investigations were performed in 20 vol% eth-
anol/water mixtures while keeping the ethanol concentration
to a minimum to afford reliable measurements owing to the
amount of hydrogen produced with our experimental set up.
Catalyst 1.0Ni(OH)2/TiO2 was the least active with a hydrogen
production rate two times lower than 1.0Cu(OH)2/TiO2. Hydro-
gen production rates increase with increasing Cu wt% ratio in
bimetallic catalysts and 0.8Cu(OH)2–0.2Ni(OH)2/TiO2 was found
to be the most active. However, the 0.8Cu(OH)2–0.2Ni(OH)2/
TiO2 catalyst losses its activity considerably if calcined at
3008C. As the Brunauer–Emmett–Teller (BET) surface area of
TiO2 P25 did not change upon calcination at 3008C, the deacti-
vation could be linked to the formation of larger particles of
oxides of Cu and Ni, which are more difficult to reduce and
therefore do not provide the needed Schottky barrier for
electron transfer to occur.
The activity and stability of 0.8Cu(OH)2–0.2Ni(OH)2/P25 cata-
lysts was investigated in 20 vol% ethanol/water mixtures. As
can be seen from Figure 8, stable photocatalytic hydrogen pro-
duction was observed over the duration of the experiments
(16 h). The slight decrease in rate with time (up to 15% for the
In comparison with ethanol/water mixtures, glycerol/water
mixtures showed higher rates of hydrogen production
(Figure 7). In general, alcohols with lower oxidation potential,
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Table 2. Comparison of hydrogen production rates over Cu- and Ni-containing catalysts reported in the literature with those of the present study.
Photocatalyst (metal loading method)
Sacrificial reagents
[vol%]
Irradiation source
H2 production
Ref.
[mmolgꢀ1 hꢀ1
]
1 wt% Cu–Ni/P25
1 wt% Cu–Ni/P25 (solvothermal)
1.25 wt% CuO/TiO2 (complex precipitation)
29 mol% Cu(OH)2/P25 (precipitation)
0.5 mol% Ni(OH)2/P25 (precipitation-deposition)
1.3 wt% CuO/P25 (impregnation)
5% glycerol/water
100 W lamp (UV)
300 W lamp (UV)
100 W lamp (UV)
3 W LEDs (UV)
3 W LEDs (UV)
3 W LEDs (UV)
400 W lamp (UV)
200 W lamp (UV)
208 W (UV) lights
100 W lamp (UV)
3 W UV-LEDs (UV)
300 W lamp (UV)
400 W lamp (UV)
450 W (UV/Vis)
2.4 W lamp (UV)
250 W (UV/Vis)
8 W (UV-B)
22.32
13.5
20.3
3.42
3.056
2.06
20
Present study
[29]
[8]
[21]
[7]
[30]
[31]
[32]
[6c]
[2a]
[7]
[6e]
[33]
[34]
[35]
[36]
[37]
[38]
[39]
[40]
[41]
[42]
[27]
37% methanol/water
80% ethanol/water
0.09m ethylene glycol/water
25% methanol/water
0.1m glycerol/water
10% methanol/water
20% ethanol/water
30% methanol/water
95% ethanol/water
25% of methanol/water
10% methanol/water
25% methanol/water
5% ethanol/water
50% methanol/water
6% methanol/water
50% ethanol/water
50% ethanol/water
25% glycerol/water
80% ethanol/water
16% glycerol/water
10% glycerol/water
10% glycerol/water
8 at%Cu/P25 (chemical reduction)
1 wt% Cu2O/P25 (impregnation)
3.48
3.2
1.2 wt% Cu/TiO2 anatase (incipient wetness)
1.25 wt% Ni/P25 (complex precipitation)
0.5 mol% Ni(OH)2/P25 (precipitation deposition)
1.5 wt% NiO/TiO2-anatase (sol–gel)
0.8 wt% Au/TiO2-anatase (photodeposition)
4 wt% Au/P25 (photodeposition)
1 wt% Au/TiO2ꢀ anatase (photodeposition)
1 wt% Au/TiO2-anatase (flame spray pyrolysis)
1 wt% Pd/TiO2-anatase+rutile nanofibers (impregnation)
0.5 wt% Au–0.5 wt% Pt/TiO2-anatase (impregnation)
0.25 wt% Au–0.75 wt% Pd/P25 (sol-immobilization)
1.5 wt% Au/P25 (deposition precipitation with urea)
2 wt% NiO/TiO2-anatase+rutile (impregnation)
1 wt% Pt/P25 (deposition precipitation with urea)
1.5 wt% Au/P25 (deposition precipitation with urea)
20.7
3.056
0.162
1.54
6.12
8.4
8
16.2
8
19.6
32.2
1.23
27.1
27.9
125 W (UV lamp)
LED (365 nm)
100 W (365 nm)
500 W (UV)
200 W (UV)
100 W (365 nm)
Electron/hole transfer
Table 3 presents Rhodamine B dye degradation over P25,
1.0Ni(OH)2/P25, 1.0Cu(OH)2/P25, and 0.8Cu(OH)2–0.2Ni(OH)2/
P25 catalysts in the absence or presence of either a hole
(Na2EDTA) or an electron scavenger (benzoquinone).[45] The ob-
jective here is to further probe the electron/hole transfer pro-
cess in the presence of Cu and Ni hydroxides. The absorption
spectra were collected in 400–600 nm range during the moni-
toring process. To test the activity, we have measured the time
needed for dye degradation, specifically, the time needed to
decompose 80% of the dye. It was noted that the time re-
quired to degrade the dye is longer in the case of metal-
Figure 8. Different successive runs for the hydrogen production experiments
with 0.8Cu(OH)2–0.2Ni(OH)2/P25 catalysts by using 20 vol% ethanol/water
mixtures.
Table 3. Results of Rhodamine B dye degradation over different photoca-
talysts. Degradation time is taken as that needed to decompose 80% of
the dye.
fourth run) was due to the deposition of catalyst particles
(clearly seen with the naked eye) on the reactor walls with
time, thus scattering away the incoming radiation. We per-
formed recyclability tests with similar bimetallic loadings on
TiO2 nanofibers, and no decrease in hydrogen production rates
was observed over the same length of time. The effect of pH
on H2 generation with 0.8Cu(OH)2–0.2Ni(OH)2/P25 catalysts
was investigated in a 5 vol% glycerol/water mixture (Figure S6
in the Supporting Information). In general, a decrease in rate
with an increase in pH was observed. We did observe negligi-
ble H2 production in water alone. H2 production in the case of
pure water on metal supported TiO2 is about two orders of
magnitudes smaller compared with that in the presence of the
sacrificial agents.[43]
Photocatalyst
P25
Hole
Superoxide
Degradation
time [min]
scavenger scavenger
–
–
32
1.0Ni(OH)2/P25
1.0Ni(OH)2/P25
1.0Ni(OH)2/P25
–
–
38
38
70
Na2EDTA
benzoquinone
1.0Cu(OH)2/P25
1.0Cu(OH)2/P25
1.0Cu(OH)2/P25
–
–
–
110
105
Na2EDTA
–
benzoquinone 170
0.8Cu(OH)2–0.2Ni(OH)2/P25
0.8Cu(OH)2–0.2Ni(OH)2/P25 Na2EDTA
0.8Cu(OH)2–0.2Ni(OH)2/P25
–
–
–
90
90
–
benzoquinone 135
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loaded catalysts compared with P25 alone in all cases. This in-
dicates that doping TiO2 with these elements decreases its ac-
tivity, in sharp contrast with hydrogen production. More specif-
ically, the dye degradation time is longer over all Cu supported
catalysts compared with 1.0Ni(OH)2/P25 (with almost the same
activity as TiO2 alone). It is to be noted that this last observa-
tion is also opposite to the hydrogen production results. The
reason is most likely competition of conduction band (CB)
electrons. In other words, Cu in Cu(OH)2/P25 has a suitable re-
duction potential, as a result of which there is competition be-
tween (Cu2+ +2eꢀ(CB)!Cu0)—not needed for the dye degra-
ꢀC
dation—and (O2 +eꢀ(CB)!O2 —needed for the dye degrada-
tion reaction. However, in the case of Ni(OH)2/P25, the conduc-
tion band electrons are unable to reduce the Ni in Ni(OH)2,
and thus are therefore exclusively used to form superoxide
radicals that in turn oxidize the dye. To probe further into this,
we have investigated the effect of hole and electron scaveng-
ers on the dye degradation time. In the presence of Na2EDTA,
a hole scavenger, the rate of dye degradation all catalysts was
close to that obtained without any scavenger. However, the
degradation was much slower in the presence of benzoqui-
Figure 10. Schematic illustration of the hydrogen production, charge trans-
fer, separation, and chemical conversions with Ni(OH)2 and Cu(OH)2 nano-
clusters on TiO2 under photoreaction conditions.
none, a superoxide radical scavenger (Table 3). These results
with the decrease in Cu content (Figure S2). 2) The presence of
Cu in its reduced state in all XPS results owing to the excita-
tion of TiO2 electrons by X-rays and their subsequent transfer
to Cu(OH)2 nanoclusters. 3) Two times lower PL intensity in the
case of 1.0Cu(OH)2/P25 compared with P25, indicating photo-
excited electron transfer from the CB of the P25 to the
Cu(OH)2. 4) The approximately three-fold decrease in dye deg-
radation rate on Cu(OH)2 loaded P25 compared with P25 alone
owing to the competition between electron transfer from the
conduction band of TiO2 to Cu(OH)2 and dye molecule, proba-
+
C
also suggest that the role of OH radicals or h is negligible in
this degradation process whereas electrons/superoxide radicals
are the main active species.
The reduction potential of Cu(OH)2/Cu is slightly lower than
the conduction band potential of anatase and rutile TiO2
(Cu(OH)2 +2eꢀ =Cu+2OHꢀ, E8=ꢀ0.224 V; Figure 9). Under
photoreaction conditions, electrons can directly transfer from
the CB of TiO2 to Cu(OH)2 nanoclusters to reduce Cu2+ to Cu0
as shown in Figure 10. This direct conversion can be supported
by four observations in this study. 1) Instant color change on
exposing the catalysts slurry to UV radiation in the presence of
a sacrificial regent. This change in color gradually decreases
ꢀC
bly via superoxide radical (O2 ) formation. This is further sup-
ported by an increase in dye degradation time in the presence
of superoxide scavenger benzoquinone whereas no change in
Figure 9. Schematic illustration of the hydrogen production, charge transfer, separation, and chemical conversions with Ni(OH)2 and Cu(OH)2 nanoclusters on
TiO2.[2a,7,9,21,47]
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dye degradation time was observed in the presence of a hole
scavenger, Na2EDTA (Table 3).[44] In the metallic state, Cu acts as
a co-catalyst for electron transfer to proton for hydrogen
production.[45]
0.2Cu(OH)2–0.8Ni(OH)2/TiO2 is due to the gradual decrease of
the Cu amount. The very small hydrogen production rate ob-
served with 0.8Cu(OH)2–0.2Ni(OH)2/TiO2 calcined at 3008C in
air might be due to conversion of Cu(OH)2 and Ni(OH)2 to their
oxides (CuO and NiO, respectively) and their subsequent
growth into larger particles by sintering. This explanation is
also in line with what has been reported previously where a de-
crease in H2 production rate with an increase of CuO particle
size was observed.[8]
The rate of hydrogen production observed with a similar
loading of Ni was far less than that with Cu (Figures 6 and 7).
The reduction potential of Ni(OH)2/Ni [Ni(OH)2 +2eꢀ =Ni+
2OHꢀ E8=ꢀ0.76 V] is more negative than the conduction
band potential of TiO2. Thus, the electron transfer from the CB
band of TiO2 to Ni(OH)2 is thermodynamically unfavorable. Al-
though there is evidence from photoluminescence and dye
degradation experiments that the Ni(OH)2 level lies above the
TiO2 CB for the photocatalysts used in this study, stable H2 pro-
duction on monometallic Ni(OH)2 loaded catalysts is observed
during the hydrogen production experiments. Furthermore,
0.8Cu(OH)2–0.2Ni(OH)2/P25 exhibits higher activity to that of
the sum of both 0.2Ni(OH)2/P25 and 1.0Cu(OH)2/P25. These
observations reject the possibility of Ni(OH)2 remaining inactive
under photoreaction conditions. Typically, in a photoreaction
experiment involving ethanol or glycerol as sacrificial agent,
the pH of the reaction mixture drops to approximately 4
within the first hour as a result of their oxidation. There is
a possibility of Ni(OH)2 dissolution in the reaction mixture as
a Ni2+ concentration of 0.1 molLꢀ1 at pH 6 is predicted in
Conclusions
A series of Cu(OH)2–Ni(OH)2/P25 photocatalysts were prepared
by co-deposition–precipitation (total metal loading ꢁ1 wt%),
characterized, and their performance evaluated for H2 produc-
tion in different alcohol/water mixtures under UV excitation.
Among this series, the 0.8Cu(OH)2–0.2Ni(OH)2/P25 photocata-
lyst demonstrated very high hydrogen production rates in
20 vol% ethanol/water and 5 vol% glycerol/water mixtures
(10 mmolhꢀ1 gꢀ1 and 22 mmolhꢀ1 gꢀ1, respectively). Detailed
analyses based on reaction kinetics, photoluminescence, XPS,
and charge carrier scavenging suggest that both working cata-
lysts are composed of Cu and Ni metals in their active phase.
Cu0 is produced directly by the transfer of electrons from the
conduction band of TiO2 to surface Cu(OH)2 nanoclusters,
whilst Ni0 is formed indirectly through a process of gradual dis-
solution of Ni(OH)2 to yield aqueous Ni2+ as a result of the
acidic environment of the medium, followed by Ni2+ reduction
by electrons from the conduction band of the semiconductor.
The high rates of H2 production, which match those obtained
with noble metals, can be explained by the following rationale.
An oxidized Ni atom in contact with a Cu atom may become
reduced owing to the considerably less negative DG8 of Cu
oxide formation (ꢀ129 kJmolꢀ1) compared with that of Ni
oxide formation (ꢀ430 kJmolꢀ1).[1] This would then increase hy-
drogen production because the work function of Ni is higher
than that of Cu.[2] The present work suggests that bimetallic
Cu–Ni catalysts formed on TiO2 are promising alternatives to
noble metals for hydrogen production.[25a,26]
[46]
equilibrium with Ni(OH)2.
To confirm this, we performed photoreactions with the high
concentration of 1% Ni(OH)2/P25 catalyst (100 mg) to allow
Ni2+ detection in the reaction mixture, by using dimethylglyox-
ime (DMG) as a complexing agent. The appearance of a red
color after 3 h clearly indicated the dissolution of Ni2+ in the
alcohol/water mixture. Based on these observations, the stable
hydrogen production, and the appropriate redox potential of
Ni2+/Ni couple (ꢀ0.23 V), we propose that Ni(OH)2 is first dis-
solved into the reaction mixture and Ni2+(aq) thus formed is
then photodeposited on P25 as shown in Figures 9 and 10.
This is also supported by the observation that relatively longer
induction time is needed to observe hydrogen production in
the case of Ni-containing catalysts (Figures 6 and 7).
The considerable hydrogen production rate in the case of
1.0Cu(OH)2/TiO2 was due to its appropriate work function
(5.1 eV). The very high rate of hydrogen production,
10 mmolhꢀ1 gꢀ1, in the ethanol/water mixture observed for
0.8Cu(OH)2–0.2Ni(OH)2/P25 and that of 22 mmolhꢀ1 gꢀ1 in the
glycerol/water mixture observed for 0.5Cu(OH)2–0.5Ni(OH)2/
P25 is a result of the synergistic effect of Cu deposited directly
from Cu(OH)2 precipitates and re-adsorption of Ni2+ cations in
the solution. These Ni2+ ions may selectively be photodeposit-
ed over electron-rich sites on the P25 surface, which are away
from the Cu nanoclusters or on the Cu nanoclusters them-
selves, forming a Cu/Ni alloy (Figure 10). It has been reported
that for an ideal composition of Cu/Ni alloy, a more suitable
Schottky barrier height can be made.[29] The reducibility of Ni is
enhanced in the Cu/Ni alloy owing to the considerably less
negative DG8 of Cu oxide formation (ꢀ129 kJmolꢀ1) compared
with that of Ni oxide formation (ꢀ430 kJmolꢀ1),[1] which is fa-
vorable for water reduction.[29] The decrease in the hydrogen
production rate on going from 0.5Cu(OH)2–0.5Ni(OH)2/P25 to
Experimental Section
Catalyst preparation
All the reagents used were of analytical grade and used without
further purification. Distilled water was used in all experiments.
Commercially available Degussa P25 was obtained from Evonik In-
dustries, Germany. In a typical synthesis, P25 (500 mg) was added
to 0.5m NaOH (50 mL) and sonicated well to give a homogeneous
slurry. Specific volumes of aqueous solutions of Cu(NO3)2·3H2O and
Ni(NO3)2·6H2O were then added dropwise to the P25 dispersion
with continuous stirring. The resulting dispersions were sonicated
for 10 min and then stirred for a further 2 h. Finally, the Cu(OH)2
and Ni(OH)2 impregnated P25 photocatalysts were collected by
vacuum filtration, washed several times with water, and then dried
in air at 808C for 24 h. The nominal weight percentages of Ni and
Cu in the photocatalysts were 1.0Cu(OH)2, 1.0Ni(OH)2, 0.5Cu(OH)2–
0.5Ni(OH)2, 0.8Cu(OH)2–0.2Ni(OH)2, and 0.2Cu(OH)2–0.8Ni(OH)2,
where the prefixes represent the weight percentage of each metal.
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A 0.8Cu(OH)2–0.2Ni(OH)2/P25 sample was also calcined at 3008C
for 2 h to examine the effect of metal oxide formation on photoca-
talytic activity. The Brunauer–Emmett–Teller (BET) surface area of
TiO2 is approximately 50 m2 gꢀ1 in all the catalysts.
only difference was that the reactor was purged with nitrogen at
a flow rate of 10 mLminꢀ1 for 30 min at 4 h increments to remove
dissolved and headspace gases.
0.05m HCl or 0.05m NH4OH solutions were used to adjust the pH
for experiments performed to see the effect of pH on H2 genera-
tion rate.
Photocatalysts characterization
Powder XRD patterns were collected with a Siemens D5000 Diffrac-
tometer equipped with a curved graphite filter monochromator.
XRD data was collected over the 2q range 20–708 (step size 0.058,
scan rate 28minꢀ1) by using CuKa X-rays (l=1.5418 ꢁ, 40 mA,
40 kV). Anatase and rutile crystallite sizes (L) were determined by
using the Scherrer equation (L=0.9l/bcosq) and line-widths of the
anatase (101) reflection at 2q=25.38 and the rutile (110) reflection
at 2q=27.48. UV/Vis absorbance spectra of the photocatalysts
were collected over the wavelength range 200–900 nm with
a Thermo Fisher Scientific UV/Vis spectrophotometer equipped
with a praying mantis diffuse reflectance accessory. TEM analyses
were performed by using a Philips CM12/STEM Electron Micro-
scope, PW 6030 (120 kV). XPS data was collected by using a Kratos
Axis UltraDLD equipped with a hemispherical electron energy ana-
lyzer and an analysis chamber at base pressure ꢁ1ꢂ10ꢀ9 Torr.
Samples were excited by using monochromatic AlKa X-rays
(1486.69 eV) with the X-ray source operating at 150 W. Samples
were gently pressed into thin pellets of ꢁ0.1 mm thickness for the
analyses. A charge neutralization system was used to alleviate
sample charge build up during analysis. Survey scans were collect-
ed at a pass energy of 80 eV over the binding energy range 1200–
0 eV, whereas core level scans were collected with a pass energy of
20 eV. The spectra were calibrated against the C1s signal at
284.80 eV from adventitious hydrocarbons. Steady-state and time-
resolved photoluminescence (PL) measurements were carried out
by using a Flau Time 300 (FT-300) steady-state and lifetime spec-
trometer, PicoQuant GmbH, Germany. The PL was measured by
using a pulsed LED laser excitation source, PLS-300, centered at
305 nm with full-width half-maximum (FWHM) of ꢁ416 ps and
pulse energy 0.077 pJ. The PL curves were fitted by using Easy Tau
and FluoFit software.
Dye degradation experiments
Rhodamine B photodegradation experiments were conducted in
a 140 mL Pyrex reactor containing aqueous Rhodamine B (40 mL,
12 mgLꢀ1) and photocatalyst (10 mg). Prior to starting an experi-
ment, the reaction mixture was stirred in the dark for one hour to
establish the adsorption–desorption equilibrium of the dye on the
catalyst surface. The same irradiation set up was used as that de-
scribed in the hydrogen production experiment section. At regular
intervals, a small portion of the reaction mixture was collected by
pipette, centrifuged to remove suspended photocatalyst particles,
and its absorbance (400–600 nm) was measured by using a UV
spectrophotometer (lambda 25 PerkinElmer). For radical trap ex-
periments, 5 mL of 50 mmol solutions of Na2EDTA or benzoqui-
none were added directly into the initial reaction mixture. Negligi-
ble Rhodamine B degradation was observed under UV irradiation
in the absence of a photocatalyst.
Acknowledgements
This work was supported by Higher Education Commission (HEC)
of Pakistan (No. 20-2704/NRPU/R&D/HEC/12). The manuscript
was written through contributions of all authors. All authors
have given approval to the final version of the manuscript.
Keywords: metallic copper · nickel deposition · non-noble
metals · P25 · water splitting
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Received: June 19, 2016
Revised: July 15, 2016
Published online on && &&, 0000
&
ChemCatChem 2016, 8, 1 – 11
www.chemcatchem.org
10
ꢀ 2016 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
ÝÝ These are not the final page numbers!

FULL PAPERS
Noble metal substitutes: A series of
Cu(OH)2–Ni(OH)2/P25 photocatalysts
was prepared by co-deposition–precipi-
tation (total metal loading ꢁ1 wt%)
and their performance was evaluated
for H2 production. Among this series,
the 0.8Cu(OH)2–0.2Ni(OH)2/P25 photo-
catalyst demonstrated very high H2 pro-
duction rates in 20 vol% ethanol/water
and 5 vol% glycerol/water mixtures (10
and 22 mmolhꢀ1 gꢀ1, respectively).
I. Majeed, M. A. Nadeem,* E. Hussain,
G. I. N. Waterhouse, A. Badshah, A. Iqbal,
M. A. Nadeem,* H. Idriss*
&& – &&
On the Synergism between Cu and Ni
for Photocatalytic Hydrogen
Production and their Potential as
Substitutes of Noble Metals
ChemCatChem 2016, 8, 1 – 11
www.chemcatchem.org
11
ꢀ 2016 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
&
These are not the final page numbers! ÞÞ