Langmuir 1987, 3, 303-304
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Photoassisted Catalytic Dissociation of H 2 0 and Reduction of N2 to NH3 on Partially Reduced Fe203 Mahmoud M. Khader,? Norman N. Lichtin,' Gerard H. Vurens, Miquel Salmeron, and Gabor A. Somorjai* Materials & Molecular Research Division, Lawrence Berkeley Laboratory, Department of Chemistry, University of California, Berkeley, California 94720 Received November 4, 1986. I n Final Form: January 22, 1987
H2 and NH, were produced when a suspension of Fe203catalyst particles or sintered pellets in water were illuminated with light of energy greater than 2.3 eV. Catalytic action was demonstrated by detecting Hz yields many times the stoichiometric equivalent of the measured (by X-ray diffraction, Mossbauer spectroscopy, and oxidimetry) Fe(I1) content of the catalyst, up to 80X for Hz and 20X for NH3. The initial rate of formation of H2was 40 ,umol/h.g of catalyst and of NH3was 10 pmo1lh.g of catalyst. The formation labeled water. of O2 was proved by mass spectrometric observation using We report here the catalytic photoassisted splitting of water into H2 and O2 and the photoassisted reduction of N2with water to NHB. The catalyst consists of a mixture of approximately 90% a-Fe203..and10% Fe,04 as determined by X-ray diffraction. Mossbauer experiments also detect the presence of Fe,04 in similar amounts. Other phases of the iron oxide, like y-Fe203could also be present in small amounts. a-Fe203is an n-type semiconductor with a band gap of about 2.3 eV.' It can absorb at least 40% of the solar flux. It has, accordingly, been investigated as a light-absorbing electrode for photoassisted electrolysis of water2-I2and as a catalyst in the photoassisted reduction of N2to ammonia by water', or by aqueous organic s01utions.l~ The photoassisted production of H2 from H 2 0 using Mg-doped Fe2025J6and of NH, from N2 using Fe-doped Ti027918has already been reported. Because of the poor yields, the catalytic nature of these processes could not be assessed. The iron oxide catalyst that we report here is much more active than the iron-containing samples that were previously utilized. The catalytic nature of the reactions could thus be demonstrated. We used the catalyst both in the form of powder and as sintered pellets. It can be prepared in powder form by exposing a-Fe203(Alfa, 99.9%) for 8 h at 450 "C and 1atm to a flowing mixture of 70% H2and 30% H 2 0 vapor and then heating it in O2or air a t 450 "C for 10 min. One-half inch diameter sintered pellets of iron oxide were prepared by pressing the powder to 7000 Kg/cm2 and then heating it in air for 20 h at 1150 OC. The pellets could be activated by reacting them first with 70% H2-30% H 2 0 a t 500 O C and 1 atm for 10 min and then with O2 or air a t 450 "C for 10 min. The composition of the catalyst was estimated from its X-ray powder diffraction pattern by comparing the ratio of the integrated intensities of the (220) spinel and (012) corundum reflections with known standards of spinel and a-Fe203. This analysis indicated that the material contains more than 90% Fe203 and less than 10% Fe304. The Mossbauer spectrum also indicated that the major component is a-Fe203 ( H = 515 kOe, e = 0.12 mm/s, IS = 0.60 mm/s with respect to SNP). Fea04was detectable but constituted less than 10% of the catalyst. Oxidimetric analysis of solutions of the dissolved catalyst in aqueous HC1 with KMn0419indicated that 3-5 atom % of the iron in the catalyst is Fe(I1). BET measurement showed a On leave from the Department of Chemistry, Boston University.
* Department of Chemistry, Boston University.
surface area of 1.85 m2/g of the powdered catalyst. To measure the hydrogen yield, the reaction was performed in a closed loop with either N2 or Ar circulating through a suspension of 30 mg of catalyst in 30 mL of deionized water in a thermostated Pyrex cell (temperature 29-30 "C).The suspension was illuminated with 20 mW of light from a 150-W Xe lamp. The gas stream was periodically leaked into a UHV chamber where it was analyzed with a calibrated quadrupole mass spectrometer. In its pure form neither a-Fe203nor the reduced form of the oxide, Fe304,is active as photocatalyst. Assuming that all the catalyst particles were active during the experiment, the initial rate of hydrogen formation was 40 ,umol/h-g of catalyst. To determine the yield of ammonia, N2 gas was bubbled through the aqueous slurry of catalyst and illuminated by 20 mW of light. The gas stream was subsequently passed through a trap containing dilute aqueous HC1 at 0-4 OC. Aliquots of solution from both the cell and the trap were analyzed spectrophotometrically by the trichloramine method.20 The initial rate of ammonia (1) Strenlow, W. H.; Cook, E. L. J.Phys. Chem. Ref. Data 1973,2,163.
(2) Wilhelm, S. M.; Yun, K. S.; Ballenger, L. W. Hackermann, N. J. Electrochem. Soc. 1979, 126, 419. (3) Curran, J. S.; Gissler, W. J. Electrochem. SOC. 1979, 126, 1983. 1979,126, 1893. (4) Fredlein, R. A.; Bard, A. J. J.Electrochem. SOC. (5) Kennedy, J. H.; Frese, K. W. J. Electrochem. SOC.1978,125,723. (6) Shinar, R.; Kennedy, J. H. Sol. Energy Mater. 1982, 6, 323. (7) Launay, J. C.; Horowitz, G. J. Cryst. Growth 1982,57, 118. (8) Pollert, E.; Hejtmanek, J.; Doumerc, J. P.; Claverie, J.; Hagenmuller, P. J. Phys. Chem. Solids 1983, 44, 273. (9) Redon, A.M.; Vigeron, J.; Heindl, R.; Sella, C.; Martin, C.; Dalbera, J. P. Sol. Cells 1981, 3 , 179. (10) 1981.129. . . McCann.. J. F.:. Badwall. S. P. S. J.Electrochem. SOC. . . 551. (11) Hardee, K. L.; Bard, A. J. J. Electrochem. SOC. 1986, 123, 1024. (12) Giordano, N.; Passalacqua, E.; Antonucci, V.; Bart, J. C. J. Znt. J. Hydrogen Energy 1983,8, 763. (13) Lichtin, N. N.; Vijayakumar, K. M. "Abstracts of the 163rd Meeting of Electrochemical Society", May 8-13, 1983; No. 782. (14) Lichtin, N. N.; Vijayakumar, K. M.; Khader, M. M. Proceedings of the June 1985 meeting of the International Solar Energy Society, Montreal, Canada. (15) Leygraf, C.; Hendewerk, M.; Somorjai, G. A. J. Phys. Chem. 1982, 86, 4484. (16) Turner, J. E.; Hendewerk, M.; Somorjai, G. A. Chem. Phys. Lett. 1984, 105, 581. (17) Schrauzer, J. N.; Guth, T. D. J. Am. Chem. Soc. 1977,99, 7189. (18) Augugliaro, V.; Lauricella, A.; Rizzuti, L.; Schiavello, M.; Sclafani, A. "Hydrogen Energy Progress", 3rd World Hydrogen Energy Conference, Tokyo, June 1980; Vol. 2. (19) Hall, W. T. Analytical Chemistry; Wiley: New York, 1955; Vol. 2, p 543. (20) Zitomer, F.; Lambert, J. Anal. Chem. 1962, 34, 1738.
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Langmuir 1987, 3, 304-306
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100.0
0.0 00
200 0
400 0
600 0
REACTION TIME (hours) Figure 1. Cumulative amounts of H2and NH3 obtained from
photodissociation of H 2 0 and photoreduction of N2 by 20 mW of light of energy greater than 2.3 eV. The catalyst consists of aqueous suspensions of partially reduced particles of iron oxide. For comparison we indicate the amounts of H2 that would correspond to the stoichiometric oxidation of Fe(I1) by H20 in samples containing 100% and 5% Fe(I1). formation was 10 pmol/h.g of catalyst. We can roughly express the rate in terms of catalyst surface area. For our experimental conditions the result is 100 pmol/h.m2 (which would be equivalent to 0.5 pA/cm2), assuming that only half of the particles are illuminated on one side a t any given time. By using Ar instead of N2 as the circulating gas, ammonia production was less than the detection limit of the trichloramine method (2 X In the case of Ha,
the yield was unaffected by using Ar instead of N2 as the circulating gas. NH3 and H2 and O2 thus appear to be formed in parallel processes. The production of both H2 and NH, was only observed when the catalyst was illuminated with light of wavelength shorter than 540 nm. When the reaction was carried in the dark neither H2 nor NH, was observed. The formation of O2was demonstrated by using a pellet suspended in 3 mL of water labeled with I80 and detecting products of mass 36, 34, and 32 mass spectrometrically. The observed O2 to H2 ratios varied between 0.41 and 0.51 in four different measurements. The oxidation products formed in the NH, reaction are not known at present. The results of these experiments are presented in Figure 1. Our studies demonstrate that the catalyst remains active for production of both NH, and H2 for about 450 h. Assuming that the catalyst contained 5 atom % of Fe(I1) (in the form of Fe304),the yield of H2 obtained in about 450 h would be equivalent to 80 times the stoichiometric reducing capacity of the catalyst while the yield of NH, obtained in about 580 h of illumination would be equivalent to 20 times the stoichiometric reducing capacity of the catalyst. In fact, no significant consumption of Fe(I1) was detected either by X-ray diffraction or by oxidimetric measurements on samples of the used catalyst. It can be concluded that the observed reactions are catalytic. As can be seen in the figure, a significant decay of catalytic activity was observed after several hundred hours of illumination. The reasons for this decay are under investigation. A more detailed account of these and other experiments will be published in a forthcoming paper.
Acknowledgment. This work was supported by the Director, Office of Energy Research, Office of Basic Energy Sciences, Chemical Sciences Division of the U.S. Department of Energy under Contract DE-AC03-76SF00098. M.M.K. thanks the Egyptian Ministry of Higher Education for the award of a support fellowship.
Geometric Formalism of the Thermodynamics of Adsorption at Interfaces between Two Fluid Phases Kinsi Motomura* and Makoto Aratono Department of Chemistry, Faculty of Science, Kyushu University 33, Fukuoka 812, Japan Received J u l y 2, 1986 The adsorption of surface-active solutes at interfaces between two fluid phases has been shown to be analyzed thermodynamically by use of the equations derived by introducing two dividing planes chosen so as to make the excess numbers of moles of solvents of the phases with reference to them zero. The positions of the dividing planes have been related to the variation in the concentrations of the solvents. Finally, this geometric formalism of the thermodynamics of adsorption has been proven to be identical with the Hansen convention. Gibbs made it possible to investigate the adsorption at interfaces between two bulk phases thermodynamically by introducing the dividing surface.' This geometric formalism was extended by Guggenheim and Adam.2 Defay et aL3 and Adamson4 elucidated the connection of the (1) Gibbs, J. W. Collected Works; Dover: New York, 1961; Vol. 1. (2) Guggenheim, E. A.; Adam, N. K. Proc. R. SOC.London, A 1933, 139, 218.
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location of dividing surface with the variation in the concentration of a component across the inhomogeneous region between the bulk phases. In 1962, Hansen5 devised the algebraic formalism which does not rely on any dividing (3) Defay, R.; Prigogine, I.; Bellemans, A. Surface Tension and Adsorption; Everett, D. H., Translator; Longmans: London, 1966; p 21. (4) Adamson, A. W. Physical Chemistry of Surfaces; Wiley-Interscience: New York, 1976. (5) Hansen, R. S. J.Phys. Chem. 1962, 66, 410.
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