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J. Phys. Chem. B 2004, 108, 19076-19077
COMMENTS Comment on “A Study of Mechano-Catalysis for Overall Water Splitting” David S. Ross* US Geological SurVey, 345 Middlefield Rd., Menlo Park, California 94025 ReceiVed: May 6, 2004
Recently Hara et al. reported on the splitting of water to H2 and O2 at nominally ambient temperatures in experiments in which metal oxides were mechanically ground against the bottom of a Pyrex reaction vessel.1 This fascinating observation was noted for Co3O4, NiO, Cu2O, and Fe3O4, whereas an additional 44 oxides, including CoO, CuO, and Fe2O3, were inert.2 Typically for the active oxides, Hara et al. collected tens of Torr of gas over 20-30 h intervals, during which time the gas production rate slowed. The rate recovered after evacuation of the vessel, however, suggesting that the process is reversible and therefore under thermodynamic control. Other features of the study included a nominally stoichiometric product mixture, that is, H2 ) 2O2. An X-ray diffraction study of the recovered oxide in experiments with Cu2O showed the production of small quantities of Cu[0] but no CuO. Work with Cu[0] itself yielded solely H2 for tens of hours, after which time O2 was produced in increasing quantities over longer periods. In seeking to explain their findings, the authors suggested cycles such as
Cu2O a Cu + 1/2O2
(1)
2Cu + H2O a Cu2O + H2
(2)
/4Co3O4 a 3/4Co + 1/2O2
(3)
/4Co + H2O a 1/4Co3O4 + H2
(4)
and 1
3
at 25 °C, while recognizing that, because they are highly endoergic, these cycles cannot be easily reconciled with their findings. In addition, an experiment with 18O-tagged water yielded O2 with an 18O/16O ratio identical to that of the water, a result they acknowledged to be inconsistent with the participation of eq +1.9 They conceded that the rationale they put forward, a catalytic process at ambient temperatures employing the cycles and based on the mechanical energy input, was speculative, and that further attention to the reaction mechanism was necessary. A closer look at the findings is presented here and reveals that low temperatures and a route involving the oxide chemistry * Corresponding author. E-mail:
[email protected]. Retired, current address: 149 Walter Hays Dr., Palo Alto, CA 94303.
Figure 1. Equilibrium hydrogen pressures (s) developed from eq 5, and copper-copper oxide stability fields developed from eqs 2 and 8. The stability boundaries are also shown (- - -).
Figure 2. H2/(2O2) ratios for copper and cobalt oxides developed from eqs 1-4 and 6 and 7. The thermochemical values for the cobalt cases were provided for temperatures up to 1500 °C5 and were smoothly extrapolated to higher temperatures for this Figure; CoO/Co (s), Co3O4/ Co (- - -), and Cu2O/Cu (‚ ‚ ‚).
directly do not apply. A reinterpretation of the data begins with consideration of the overall governing process
H2O a H2 + 1/2O2
(5)
the thermodynamics of which provides the limits to the ultimate quantities of the gases formed at a given set of conditions, reaction rates aside. The solid curve in Figure 1 presents the pressure of H2 in Torr in equilibrium with water as a function of temperature, and it is observed that, at ambient temperatures, the maximum possible pressure is well below those reported. It is apparent from the Figure that the tens of Torr described by Hara et al. are only likely at temperatures no lower than about 1500 °C. An examination of the role of the metal oxides is presented in Figure 2, developed from the thermochemistry of eqs 1-4.3-5 The Figure presents the pressure ratios H2/2O2 for the active
10.1021/jp040336l CCC: $27.50 © 2004 American Chemical Society Published on Web 11/11/2004
Comments
J. Phys. Chem. B, Vol. 108, No. 49, 2004 19077
copper and cobalt oxides, and the nonactive CoO, represented by eqs 6 and 7, is included for comparison.
CoO a Co + 1/2O2
(6)
Co + H2O a H2 + CoO
(7)
It is immediately clear from the Figure that, when the thermochemistry of the cycles is included in the exercise, the product ratios vary enormously with temperature. The ratios approach the stoichiometric balance point of 1.0 only at high temperatures in the range of 1000-1300 °C, but at temperatures nonetheless below the 1500 °C minimum where the product mixtures would be highly oxygen rich. It is notable, moreover, that the stoichiometric ratio is struck by the active Co and the Cu systems at points separated by more than 200 °C. Recognizing the general uniformity of the product ratios around the stoichiometric point reported for the active oxides, we can reason that the fundamentals of the Hara et al. observations must not include the thermodynamics of the cycles. It is also significant that the active Cu and the inactive Co oxides appear to provide a unit ratio at about the same temperature, reflecting no trend distinguishing the nonactive from the active oxides We are led to conclude that the generation of H2 and O2 in the experiments takes place simply through the thermal decomposition of water, that is, eq 5, which of course directly leads to the stoichiometric product ratio, taking place most likely in small regions of transient, high temperatures generated by frictional forces. The reaction is probably catalyzed by the active oxides and is essentially at equilibrium. The hot regions could involve frictionally driven cavitation of the medium where temperatures can reach 5000 °C,6 and the absence of any effect for the inactive oxides must simply be due to their not creating sufficiently high temperatures and/or the fact that they are not catalytic. The observations with Cu and its oxides specifically can be understood with attention to the stability regions in Figure 1 generated from eqs 2 and 8.
Cu2O + H2O a CuO + H2
(8)
The Figure shows that at the high temperatures discussed above the equilibrium pressure of H2 is well above the CuO stability field, and the Cu[II] oxide is therefore not expected. Reactions at 1500 °C are in the Cu[0] region, and thus, the native metal
is formed via eq -2. Its formation directly supports the claim of reaction temperatures no less than about 1500 °C. The reaction is slow relative to eq -5, however, in accord with there being no oxide thermodynamics component in the process. The intriguing observation with Cu[0], initially producing H2, and then later an H2-O2 mixture, is best rationalized through a scheme in which the relatively soft Cu[0] both generates frictional temperatures below 1500 °C and does not catalyze splitting. Thus, eq 5 is not established, but hydrogen is generated via eq +2. Cu2O is then produced, which in turn brings about higher frictional temperatures; eq 5 eventually follows, and both gases are then formed. Overall, the promotion of the splitting by the active oxides must be due to factors ranging from individual catalytic capacities to a diversity of mechanical and frictional properties. Whatever the detailed features of the process, it is notable that water splitting as an economic source for hydrogen is, at present, an active research area. Extensive programs featuring advanced photoelectrochemical and solar thermal conversions have been reported,7 as have studies of biological routes to splitting.8 It would be of considerable interest to add simple friction to the list. Acknowledgment. The author acknowledges very helpful correspondence with Professer Michikazu Hara and Dr. Theodore Mill. References and Notes (1) Hara, M.; Komoda, M.; Hasei, H.; Yashima, M.; Ikeda, S.; Takata, T.; Kondo, J.; Domen, K. J. Phy. Chem. B 2000, 104, 780. (2) Ikeda, S.; Takata, T.; Komoda, M.; Hara, M.; Kondo, J.; Domen, K.; Tanaka, A.; Hosono, H.; Kawazoe, H. Phys. Chem. Chem. Phys. 1999, 1, 4485. (3) Chou, I.-M. In Hydrothermal Experimental Techniques, John Wiley and Sons: New York, 1987; pp 61-99. (4) Stull, D. R.; Prophet, H. JANAF Thermochemical Tables, 2nd ed.; NSRDS-NBS 37, National Bureau of Standards: Washington, D.C., 1971. (5) Chemistry WebBook at http://webbook.nist.gov/chemistry/, NIST Standard Reference Database Number 69, 2003 (6) Suslick, K.; McNamara, W., III; Didenko Y. In Sonochemistry and Sonoluminescence, Kluwer Publishers: Dordrecht, The Netherlands, 1999; pp 191-204. (7) Glatzmaier, G.; Blake, D.; Showalter S. Assessment of Methods for Hydrogen Production Using Concentrated Solar Energy; Report NREL/ TP-570-23629; National Renewable Energy Laboratory: Golden, CO, 1998. (8) Markov, S.; Weaver, P. F.; Seibert, M. In Hydrogen Energy Progress XI, Proceedings of the 11th World Hydrogen Energy Conference, Stuttgart, Germany, 23-28 June 1996; Verziroglu, T. N., Veziroglu, T. N., Winter, C.-J., Baselt, J. P., Kreysa, G., Eds.; Frankfurt, Germany, 1996. (9) The “+” and “-” designations here refer, respectively, to the forward and reverse reactions of the pair.
