run 31run 14run 9run 4run 3run 2run 1
0100200300400 7001000 1800
t / h
Amounts of evolved H2 and O2 / µmol
Cu2O as a photocatalyst for overall water splitting under visible light
Michikazu Hara,aTakeshi Kondo,aMutsuko Komoda,aSigeru Ikeda,aKiyoaki Shinohara,bAkira Tanaka,b
Junko N. Kondoaand Kazunari Domen*a
aResearch Laboratory of Resources Utilization, Tokyo Institute of Technology, 4259 Nagatsuta, Midori-ku, Yokohama 226, Japan
bNikon Corp., 1-10-1 Azamizodai, Sagamihara 228, Japan
Photocatalytic decomposition of water into H2and O2on
Cu2O under visible light irradiation is investigated; the
photocatalytic water splitting on Cu2O powder proceeds
without any noticeable decrease in the activity for more than
So far, many photocatalysts have been reported to decompose
water into H2and O2under UV light irradiation.1–5From the
view point of solar energy conversion, however, a photocatalyst
which works under visible light irradiation (>400 nm) is
indispensable, but such a photocatalyst has not yet been found.
In this report, we introduce Cu2O, a well-known p-type
semiconductor, which acts as a photocatalyst for overall water
splitting under visible light irradiation (@600 nm).
The solid state physics of Cu2O, which abundantly exists as
cuprite in nature, has been extensively investigated for a long
time since Cu2O is a simple metal oxide semiconductor with a
small band gap energy. As shown in an energy correlation
between the band gap model of Cu2O and the redox potentials
of relevant electrode reactions in an aqueous solution at pH 7,6
the conduction and valence band edges of Cu2O, which are
separated by a band gap energy of 2.0–2.2 eV,7,8,9seem to be
available for reduction and oxidation of water, respectively.
Therefore, Cu2O is, in principle, capable of decomposing water
into H2and O2under visible light excitation. However, as yet,
such a photochemical reaction has not been accomplished on
any Cu2O electrode since they undergo photodegradation in
aqueous solution.9In fact, visible light irradiation (470 nm) of
a cathode-polarized Cu2O single crystal electrode in aqueous
solution resulted in reduction of Cu2O to metallic Cu.10For this
reason, overall water splitting on Cu2O photocatalysts has not
been investigated despite the band structure available for the
reaction. In this study, we confirmed the photocatalytic overall
water splitting on Cu2O powder under visible light irradia-
Cu2O powder prepared by the hydrolysis of CuCl was used in
this study. CuCl was hydrolyzed by adding 1 maqueous Na3PO4
(40 cm3) to a 5 m aqueous NaCl solution containing 0.04 mol of
CuCl (400 cm3) with vigorous stirring under an Ar flow. A
yellow precipitate was produced by the hydrolysis which was
washed with distilled water (200 cm3) 5–7 times followed by
decantation under vacuum and drying in vacuo. Cu2O powder
was obtained by heating the yellow precipitate at 673 K for 24
h in vacuo, followed by boiling in water under an Ar atmosphere
to remove unreacted CuCl from Cu2O. The particle size and
surface area of Cu2O were estimated to be 0.3–0.5 mm and 6 m2
g21, respectively. Only the XRD pattern due to Cu2O was seen
with no evidence for other diffraction patterns such as for CuO,
metallic Cu or other impurities. The XP spectra of Cu 2p and the
Cu LMM Auger spectra indicated that the surface of Cu2O was
composed of CuI.11,12The band gap energy of Cu2O was
estimated at ca. 2.0 eV (l ca. 620 nm) by UV–VIS
The photodecomposition of water was carried out in a Pyrex
cell with 0.5 g of Cu2O and 200 cm3of distilled water, which
was vigorously magnetically stirred. The cell was irradiated at
room temperature from one side with visible light (l >460 nm)
from a 300 W Xe lamp with a cut-off filter. A closed gas
circulation and evacuation system (300 cm3) made of Pyrex
glass was connected to the reaction cell, and evolved gases were
directly transferred to a gas chromatograph to avoid any
contamination from air.
Fig. 1 shows several typical time courses of H2 and O2
evolution from Cu2O under visible light irradiation (>460 nm).
The reaction system was evacuated after each run. As shown in
run 1, only O2evolved for 10 h after the beginning of the
reaction, and then the evolution of H2was observed as the
reaction proceeded. The rate of H2 evolution increased
gradually in the subsequent runs. The ratio of the amount of
evolved H2to O2(H2/O2) was 0.8 in run 1 and increased to 1.8
in run 4. After run 4, the ratio was between 2.0 and 2.5. The
reaction proceeds without any noticeable decrease in the
activity for more than 30 runs as shown in Fig. 1. The total
amounts of evolved H2and O2for 1900 h reached 3.8 and 1.9
mmol, respectively, and are comparable to the amount of Cu2O
used (0.5 g, 3.5 mmol). Furthermore, there was no noticeable
difference in pH of the suspension before reaction (pH 7.3) and
after run 31 (pH 7.1). In order to elucidate the origin of the
evolved O2, an experiment using H218O was carried out. In
another small Pyrex cell (50 cm3), 0.1 g of Cu2O (after reaction
for 400 h) suspended in a mixture of H216O (5 cm3) and H218O
(1 cm3) was irradiated with visible light (>460 nm). H2and O2
were stoichiometrically evolved after light irradiation and it was
confirmed by mass spectral analysis that the ratio,
16O2:16O18O:18O2, in the evolved O2species for 24 h was
254:94:13. The result indicates that the atomic ratio, 16O:18O,
in the total amount of evolved O2is 5.0 corresponding to that in
the mixed water. As a result, the evolved O2is attributed to the
water cleavage on Cu2O. The photoresponse on Cu2O was
observed for visible light through a cut-off filter of 600 nm,
while there was no photoresponse at l >650 nm.
Fig. 1 Time courses of H2(open circles) and O2(filled circles) evolution in
Cu2O under visible light (l !460 nm) irradiation. Catalyst: 0.5 g, H2O: 200
cm3. The reaction system was evacuated with light irradiation after each
run. Time courses in runs 5–8, 10–13 and 15–30 are omitted.
Chem. Commun., 1998357
Intensity / a.u.
Kinetic energy / eV
Eb / eV
Cu LMM Auger
Fig. 2 shows the XP spectra of Cu 2p and Cu LMM Auger
peaks of Cu2O before and after reaction for 400 h. There was no
noticeable difference in the XP spectra of Cu2O, indicating that
Cu2O powder was neither reduced nor oxidized after photo-
catalytic reaction. These results are in total contrast to the
observation on Cu2O electrodes and strongly suggest that Cu2O
powder catalytically decomposes water into H2and O2under
visible light irradiation. To the best of our knowledge, such a
reaction on Cu2O photocatalysts has not yet been reported. The
reaction mechanism as well as the reason for the difference
between the electrode and the powder systems are still not clear.
Nevertheless it is inferred that the photocatalytic reaction on a
Cu2O particle in distilled water is clearly different from the
photoelectrochemical reaction on polarized Cu2O electrodes in
an aqueous electrolyte. The quantum efficiency of the photo-
catalytic reaction was estimated at ca. 0.3% between 550 and
One of the characteristic features of the Cu2O photocatalyst is
the excess evolution of O2above the stoichiometry at the early
stage of the reaction (runs 1 and 2). Cu2O is known to absorb a
relatively large amount of oxygen in bulk as well as adsorbing
oxygen as O2or O22on the surface.13,14The excess oxygen on
the surface or in the bulk leads to p-type semiconducting
behaviour and unique oxidation catalysis of Cu2O. The release
of these excess oxygen species from Cu2O by visible light
irradiation may cause the excess evolution of O2above the
stoichiometry at the early stage of the reaction. Another feature
to be noted is the O2pressure dependence of the reaction.
As shown in runs 3 and 4 of Fig. 1, the evolution rates of H2
and O2became slow or stopped when the amount of evolved O2
exceeded ca. 80 mmol which corresponded to 500 Pa of O2in
our system. In all runs after run 5, H2and O2evolved without
any significant decrease in the activity so long as the evolved
gas was evacuated before the pressure of O2reached 500 Pa.
These results suggest that O2at more than a certain pressure
(500 Pa) in the reaction system inhibits the overall water
splitting on Cu2O. Such an inhibition might be attributed to the
photoadsorption of oxygen on the Cu2O surface. p-Type
semiconductors are known to photoadsorb O2 under light
irradiation when O2in gas phase exceeds a certain pressure.15
The photoadsorption largely depends on the O2pressure as well
as on the wavelength and intensity of incident light, tem-
perature, etc. Although the dependence of photoadsorption on
O2pressure in a Cu2O–H2O–O2/H2system as in the present
case has not yet been investigated, it is probable that preferential
O2photoadsorption inhibits the overall water splitting on the
Although Cu2O has been regarded as an unstable material for
water decomposition under light irradiation from the results of
photoelectrochemistry, the present study has revealed Cu2O to
be a photocatalyst able to decompose water into H2and O2
under visible light irradiation. The reaction mechanism on
Cu2O is under investigation.
Recently, we have also found that CuFeO2evolves H2and O2
under visible light irradiation, and detailed results will be
reported soon. CuFeO2has a delafossite type layered structure
where the iron oxide layers are connected to each other through
linear –O–CuI–O– bonds.16–18The Cu2O lattice consists of
chains of linear bonds. This suggests that CuIcontaining
materials with linear –O–CuI–O– bonds are available for the
overall water splitting under visible light irradiation. Such CuI
containing materials may become potential candidates for
converting solar energy into H2energy.
Notes and References
* E-mail: firstname.lastname@example.org
1 D. Duonghong, E. Borgarello and M. Graetzel, J. Am. Chem. Soc., 1981,
2 K. Domen, S. Naito, T. Onishi, K. Tamaru and M. Soma, J. Phys.
Chem., 1982, 86, 3657.
3 K. Domen, A. Kudo, A. Shinozaki, A. Tanaka, K. Maruya and T.
Onishi, J. Chem. Soc., Chem. Commun., 1986, 356.
4 Y. Inoue, T. Kubokawa and K. Sato, J. Chem. Soc., Chem. Commun.,
5 K. Sayama and H. Arakawa, J. Chem. Soc., Chem. Commun., 1992,
6 H. Gerischer, J. Electroanal. Chem., 1977, 82, 133.
7 C. Kittel, in Introduction to Solid State Physics, 5th edn., Wiley, New
York, 1976, p. 341.
8 P. W. Baumeister, Phys. Rev., 1961, 121, 359.
9 G. Nagasubramanian, A. S. Gioda and A. J. Bard, J. Electrochemcal
Soc., 1981, 128, 2158.
10 H. Gerischer, Ber. Bunsenges Phys. Chem., 1971, 75, 1237.
11 R. V. Siriwardane and J. A. Poston, Appl. Surf. Sci., 1993, 68, 65.
12 K. Domen, S. Naito, M. Soma, T. Onishi and K. Tamaru, J. Chem. Soc.,
Faraday Trans. 1, 1982, 78, 845.
13 H. D¨ unwald and C. Wagner, Z. Phys. Chem. B, 1933, 40, 197.
14 B. J. Wood, H. Wise and R. S. Yolles, J. Catal., 1969, 15, 355.
15 Th. Wokenstein and IV. Karpenko, J. Appl. Phys., 1962, 33, 460.
16 A. Pabst, Am. Mineral., 1946, 31, 539.
17 R. D. Shannon, D. B. Rogers and C. T. Prewitt, Inorg. Chem., 1971 10,
18 C. Prewitt, R. D. Shanonn and D. B. Rogers, Inorg. Chem., 1971, 10,
Received in Cambridge, UK, 15th October 1997; 7/07440I
Fig. 2 X-Ray photoelectron spectra of Cu 2p and Cu LMM Auger spectra
of Cu2O before (a) and after (b) reaction for 400 h. The binding and kinetic
energies were referenced to the Au 4f7/2level at 83.8 eV.
Chem. Commun., 1998