Decomposition of Hydrogen Peroxide on Supported Metal Oxides (34) (35) (36) (37) (38) (39) (40) (41) (42) (43) (44) (45)
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The Journal of Physical Chemistry, Vol. 82, No. 13, 1978
M. Martinaud and G. Nouchi, J . Chim. Phys., 70, 450 (1973). R. Calas and R. Lalande, Buli. SOC.Chim. Fr., 763, 766, 770 (1959). E. A. Chandross, J . Chem. Phys., 43, 4175 (1965). M. Irie, T. Kamijo, M. Aikawa, T. Takemura, K. Hayashi, and H. Baba, J. Phys. Chem., 81, 1571 (1977). D. M. Hanson, J . Chem. Phys., 52, 3409 (1970). K. E. Mauser, H. Port, and H. C. Wolf, Chem. Phys., 1, 74 (1973). S. Leach, J. Phys. C3, 28, 134 (1967). J. Ferguson and W. G. Schneider, J . Chem. Phys., 28, 761 (1958). V . Zanker and J. Preuss, Z. Angew. Phys., 27, 363 (1969). A. Bree and L. E. Lyons, J . Chem. SOC.,2662 (1956). M. S. Brodin and S. V. Marisova, Opt. Spectrosc., 10, 242 (1961); 19, 132 (1965). Another species had already been defined: the aggregate.34 This species when it exists corresponds to a third new spectrum and is defined by exclusion of oligomer and crystallite. Its existence is not as general as that of oligomer and crystallite since it is observed with only two solutes: 9,10-(CI),A3* and 9-Cn-A. It appears on heating in a narrow concentration range as an intermediate for the evolution of the oligomer to the crystallite, but with a spectrum which is unambiguously different from those of the oligomer and of the crystallite. I t could seem that ambiguity arises if, for example, a crystallike spectrum appears at low concentrations and low temperatures. The above given definitions imply that it is simultaneously due to crystallite (crystallike spectrum) and oligomers (low temperatures and low
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concentrations). How to break this difficulty?. Of course this species is a crystallite. Referring to the general behavior of the aggregated species we know that the oligomer always exists at lower temperatures and lower concentrations than for the crystallite; these are conditions of low solute diffusion rate. Thus the oligomer must be searched over experimental conditions which diminish the solute diffusion rate. For instance this problem occurred with tetracene and was solved by using a very fast cooling rate.31 M. Martinaud and F. Dupuy, Mol. Specfrosc. Dense Phases, Proc. Eur. Congr. Mol. Specfrosc., 12fh, 1975, 433 (1976). M. Itoh, T. Mimura, H. Usui, and T. Okamoto, J . Am. Chem. Soc., 95, 4388 (1973). M. Itoh and T. Mimura, Chem. Phys. Lett., 24, 551 (1974). C. Tanford, "Physical Chemistry of Macromolecules", Wiley, New York, N.Y., 1961. G. A. Von Salis and H. Labhart, J . Phys. Chem., 72, 752 (1968). H. Greenspan and E. Fischer, J . Phys. Chem., 69, 2466 (1965). Note that with the two solutes (9,10-(Cl)2-A,9-Cn-A) with which the so-called aggregate is observed, region I ,does not exist and in region I1 sequence C is replaced by sequence F: oligomer aggregate small crystallite monomer. The behavior in the other regions is unchanged. We arrive at the same conclusion with other solutes. In region I, and region I 1 the small crystallite is not formed on heating if, by modifying the preceding cooling branch the formation of the oligomer is prevented.
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Formation of Superoxide Ion during the Decomposition of Hydrogen Peroxide on Supported Metal Oxides Nobumasa Kitajima, Shun-ichi Fukuzumi, and Yoshio Ono" Department of Chemical Engineering, Tokyo Institute of Technology, Ookayama, Meguro-ku, Tokyo, Japan (Received January 30, 1978) Publication costs assisted by the Tokyo Institute of Technology
Superoxide ion (02Jis formed in the aqueous phase during decomposition of hydrogen peroxide over metal oxides supported on alumina. The kinetics of the decomposition and superoxide ion formation have been studied. The same kinetic equations can be applied as previously reported for the decomposition over metals supported on alumina. The kinetic results are explained based on Weiss's mechanism. A correlation between the catalytic activities of metal oxides for the decomposition and the oxidation potential of metal cations are accounted for by the mechanism. Introduction A number of mechani~msl-~ have been proposed for the decomposition of hydrogen peroxide over heterogeneous catalysts, most of them based on the classical Haber-Weiss m e ~ h a n i s m ~which - ~ involves free-radical chain reactions. The mechanisms proposed have been based mostly on kinetic investigations and no attempts have been made to obtain direct evidence for the participation of free-radical intermediates, though they have been observed by ESR in the homogeneous decomposition of hydrogen per~xide.~-l~ In the previous paper,15we reported that the superoxide ion (02Jis formed in the aqueous phase during the decomposition of hydrogen peroxide by supported metals and the kinetics of the decomposition was accounted for by the mechanism proposed by Weiss.I This work will show that the superoxide ion is observed also in the decomposition of hydrogen peroxide over supported metal oxides and the kinetics of the decomposition over nickel oxide is studied in detail. The catalytic activities of metal oxides for the decomposition are rationalized with the oxidation potentials of the metal cations. Experimental Section Catalysts. Metal oxides supported on alumina were prepared as follows: alumina from Sumitomo Chemicals 0022-365417812082-1505$01 .OO/O
was sieved into 16-24 mesh and heated in air 550 "C for 5 h. The alumina was then immersed in a 0.5 M aqueous solution of Ni(NO3I2,FeCl,, CoC12, C U ( N O ~ Cr(N03),, )~, or ZnC12, or a 0.1 M solution of H2PtCl6overnight, filtered, dried, and calcined in air a t 600 "C for 5 h. The metal contents of the supported catalysts were determined by atomic absorption spectroscopy. In the cases of ruthenium or silver oxides, alumina was immersed in an aqueous solution of RuC13 and AgNO,. The solution was then dried, and the resulting alumina containing 0.5 wt % of the metal was calcined in air a t 600 and 150 OC, respectively. Reagent. Hydrogen peroxide (30% aqueous solution) was obtained from Mitsubishi Gas Chemical Industry. The concentration was determined by titration with 0.1 M potassium permanganate aqueous solution. Procedure. The apparatus and the procedure were described in detail e1~ewhere.l~ The catalyst (0.5-1.5 g ) was packed in a Pyrex tubing (7 mm i.d.1. The hydrogen peroxide solution was fed a t 275 K from a reservoir with the rate of 7-15 cm3 mi& by a roller pump. The concentration of hydrogen peroxide a t the outlet of the catalytic column was determined by titration with potassium permanganate. The outflow was frozen with liquid nitrogen and its ESR spectrum was measured. The concentration of superoxide ion was determined by comparison of the area under the absorption 0 1978 American
Chemical Society
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The Journal of Physical Chemistry, Vol. 82, No. 13, 1978
N. Kitajima,
S.Fukuzumi, and Y. Ono
TABLE I : Rate Constants for t h e Decomposition of Hydrogen Peroxide with Various Metals or Metal Oxides Supported o n Al,O, Oxidation potential, Noes Catalysts eV k,*b 102k,*b 10-2k,*b 1 2 3 4 5 6 7 8 9 10 11 12 13
Fe ,O ,/AI20 , NiO/Al,03 A d 4z 0 3 Ag ,O / Al,03 Pt / Al, 0, P t 0/A1,0, C O , 0, /A120 3 RuO,/Al,O, Pd/A1,03 Rh / Al, 0 , CuO/Al,O, Cr, 0 / A1,0 , ZnO/Al,O,
-0.56 (2+-3+) -0.49 (2+-4') -0.20 ( 0 - l + ) -0.20 ( 0 - - l + ) -0.15 (0-2") -0.15 (0-2+l -0.14 ( 2 + - 3 + ) -0.10 ( 2 + - 4 + ) -0.07 (0-2+) -0.04 (0-3') 0.08 (1+-2") 1.1(2+-3+) 1 . 2 (0-2+)
2.2 3.5 76 93 66 48 13 13 47 8.2
5.4 p2(Fe); 0 2 > H20 (Ni); O2 > H20 (A1 rubbed only for a short time); H20 > O2 (A1 rubbed repeatedly); H 2 0 = O2 (Cu). The activity of OSEE from iron and nickel after exposure to oxygen at various pressures decreased with increasing pressure, particularly the former being strongly suppressed by the adsorption of a small amount of oxygen. The activity of OSEE from nickel was also enhanced by the interaction of hydrogen. The OSEE from iron rubbed in various gaseous environments was well correlated with the amount of iron wear particles produced during this friction process except for the environment of oxygen and water vapor. The ratio of the OSEE to the amount of the wear particles with the water vapor was much larger than that with the other environments.
Introduction EEE from new mechanically created surfaces contributes to enhanced activity of the surfaces, which can lead to rapid surface oxidation in the presence of oxygen, or other chemical reactions with environmental species.1#2However, to our knowledge, only a few attempts have been made to establish the use of mechanically induced EEE as a technique in trib~logy.~ The use of a Geiger counter, which exposes the surface under test to the counter gas, makes unequivocal interpretation of the adsorption processes difficult. However, the technique is considered to be useful in providing an empirical method of attack on problems associated with friction and wear. Previous investigations4 have demonstrated the dependence of EEE from ground aluminum powder on the adsorption of gases. It is of considerable interest to investigate EEE not only after grinding but also during mechanical treatment such as sliding friction, when there is intensive development of defects and formation of new surfaces. The present paper is associated with practical metal surfaces interacting with the environments, not with well-defined surfaces. A new procedure was developed for studying the interaction of gases with rubbing surfaces, and a rig, which permitted continuous measurement of EEE during and after friction, was constructed by modifying the Geiger counter. The following tests were conducted: (1)the emission behavior during and after friction, ( 2 ) the behavior of OSEE after exposure to gaseous
environments such as oxygen and water vapor, and (3) the relationship between the OSEE from iron rubbed in various environments and the amount of iron particles produced during this friction process.
Experimental Section (a) Materials. Metal specimens were rolled sheets of iron (purity, >99.7%), nickel (>99.7%), aluminum (> 99.5%), and copper (>99.9%). The flat metal specimen was 30 X 30 X 0.1 mm (Fe, Ni, and Cu) and 30 X 30 X 0.3 mm (Al). These specimens were degreased with benzene solvent and then annealed in vacuo (about 0.13 N/m2) for 1 h a t 955 OC (Fe), for 2 h at 350 "C (Fe, Ni, and Cu) (iron specimens of this type were usually used), and for 1h a t 300 "C (Al) before use. Gaseous environments used were argon (purity, >99.99%), oxygen (>99.95% 1, redistilled water vapor, hydrogen (>99.99% ), ethanol vapor, and an atmosphere under vacuum (the pressure in this case was usually about 0.67 N/m2). The composition of the Geiger-counter gas was a mixture of organic vapor (C2H50H,CH,CN, C6H6,or n-C3H7NH2)(2700 N/m2) and argon (11200 N/m2), as shown in Table I, and Q gas (helium plus 1%isobutane) a t atmospheric pressure. The latter was used as a gas-flow counter gas for some tests with nickel. A counter gas of C2H50H-Arwas usually used for all the metals. Here, it should be noted that the gas compositions of these counter gases and other mixtures such as argon-oxygen mentioned later were obtained by use of a mercury manometer.
0022-3654/78/2082-1509$01 .OO/O 0 1978 American Chemical Society
