摘要:

J . Chem. SOC., Faraday Trans. I , 1984,80, 1415-1424 Oscillatory Phenomena Application of the D’ Alba-Di Lorenzol Model to the Bray-Liebhafsky2 System and Others Derived from it BY FRANCESCA D’ALBA* Istituto di Ingegneria Chimica, Wale delle Scienze, 90100 Palerrno, Italy AND SERGIO DI LORENZO Istituto Tecnico Commerciale ‘ Bortolo Belotti’, Via Azzano, 24100 Bergamo, Italy Received 13th June, 1983 The model proposed by D’ Alba and Di Lorenzo, which explains oscillations by supersaturation and phase exchange, has been applied to the Bray-Liebhafsky system (hydrogen peroxide, potassium iodate and sulphuric acid), the Briggs-Rauscher system (hydrogen peroxide, potassium iodate, sulphuric acid, manganous sulphate and malonic acid) and the arsenite- iodate-chlori te sys tem. The mechanisms of these reactions and the roles of the phase exchanges have been elucidated.The role of the catalyst and of some of the organic substrates is discussed, along with an interpretation of some of the experimental results. Studies concerning oscillating chemical systems are of increasing interest because they may provide experimentally tractable models of some endogenous rhythms in living systems. Only three different non-biological oscillating systems, the Bray- Liebhafsky,2 the Belousov-Zhabotinskii3 and the Morgan4 systems, have been identified, others (e.g. the Briggs-Rauscher5 system) being derived from them. There have been many attempts to develop a mode16-18 capable of explaining the behaviour of these systems using only the kinetics of the reactions. OSCILLATORY MODEL BASED UPON SUPERSATURATION AND PHASE EXCHANGE We do not agree with the application of these kinetic models to the Bray-Liebhafsky, Belousov-Zhabotinskii and Morgan systems and have presented1 a new theoretical model based upon phase exchange and pulsating supersaturation of the solution. We consider a stirred batch reactor.An irreversible chemical reaction takes place (1) inside the solution with kinetics given by (2) B is a gas at the temperature and pressure of the reaction and hence phase exchange B(so1.) -+ B(gas) (3) -dl\r,/dt(p.e. 1,l) = K2S[B] (4) A(so1.) -+ B(so1.) - d[A]/dt = d[B]/dt(chem.) = &[A]. occurs on going from the solution to the gas above it with kinetics given by if the gas is renewed continuously. 14151416 OSCILLATORY PHENOMENA It is possible to obtain the function B(t) using Vd[B]/dt(overall) = Vd[B]/dt(chem.) + dNB/dt(p.e.1,l) ( 5 ) with the boundary conditions if t = to, then [A] = [A], and [B] = [B],. This function increases up to the value oft, obtained at d[B]/dt(overall) = 0, which gives [B],,,, the maximum value of [B], on substituting it in eqn (5). If [BImax is lower than the value of saturation [B],,,., then [B],ax is obtained and [B] decreases. If [B]m,x > [B]sat., [B],ax cannot be obtained and the concentration of B increases until it reaches [B],,,.. The thermodynamics would prescribe that the system creates a new phase of pure B(gas) as bubbles inside the solution. However, the saturated solution is not in equilibrium with the bubbles and they cannot be nucleated. Hence the concentration of B increases up to a value [Bib allowing nucleation, with kinetics given by - The balance is Vd[B]/dt(overall) = Vd[B]/dt(chem.) + dNB/dt(p.e.1,1) + dNB/dt(p.e. 2). (7) We may have d[B]/dt(overall) $0. If d[B]/dt(overall) < 0, the concentration of B becomes less than that allowing nucleation and we have no new nucleation. However, the presence of the bubbles alters the system, because they can vary in mass and volume; in fact the solution becomes supersaturated with respect to the bubbles. Phase exchange then takes place from the solution to the bubbles. With some simplifications1 we have -dNB/dt(p.e. 1,2) = K3sb V[n,-Cf(f-?b)] ([BI-[B],) (8) and Vd[B]/dt(overall) = Vd[B]/dt(chem.)+ d%/dt(P.e. 1,1) +dN~/dt(P.e.1,2)- (9) When the bubbles have left the solution the phenomenon is repeated and the physical system is governed by the mathematical system This gives oscillations in the concentration of B. We have demonstratedl that this model is capable of explaining the oscillations in the Belodsov-Zhabotinskii, Bray-Liebhafsky and Morgan systems. The influence of the rate of stirring is definitive and unmistakable evidence for discriminating between our modell and that of Noyes, who also agrees about the Morgan system, contrary to his previous work.20 APPLICATION OF THE MODEL TO THE BRAY-LIEBHAFSKY REACTION PRELIMINARY REMARKS In this work we detail the application of our model to the Bray-Liebhafsky system and to the systems derived from it. We explain the mechanisms of the reactions and the role of the catalyst and some of the organic substrates.Since our model freed the kinetics of oscillating reactions from the constraints imposed by models such as the ' Oregonator ',14 it is no longer possible to invoke these constraints to justify some unlikely steps and to neglect other more plausible steps.F. D'ALBA AND S. DI LOREN20 141 7 SETTING UP THE KINETIC MODEL Bray2? 21 has shown that the potassium iodate, hydrogen peroxide and sulphuric acid system has oscillating behaviour at 333 K in a range of concentrations of sulphuric acid. He considered the reactions 5H20, + I, + 2HI0, + 4H20 H,O, -+ H,O + 40, (10) (1 1) (12) 5H,O, + 2HI0, -+ 5 0 , + I, + 6H,O and demonstrated the possibility of a catalytic role for the redox couple 10;/1, in the decomposition of hydrogen peroxide.The oscillations can be explained by the following mechanism. Both reactions (10) and (1 1) are overall reactions that may be derived by bimolecular steps. Reaction (1 1) is derived from the following steps, some demonstrated for a Briggs-Rauscher subsystem by Furrow and Noyes,, by their tests with phenol. Steps A 1 0 , + H+ f HIO, (13) HId, + H+ f +IO, + H,O (14) '-10, + H20, 3 + I 0 + H,O + 0, (15) + I 0 + H20 f HIO, + H+ (16) HIO, f 1 0 , + H+ (17) + I 0 + H,O, -+ I+ + H,O + 0, (18) I+ + H,O f HIO+H+ (19) HI0 f 1 0 - + H+ (20) I++H,O, -+ I-+2H++02 (21) I+ +I- f I, (22) +IO, + I- f I++ 1 0 , (23) + I 0 + I- $ I++ 1 0 - . (24) We do not consider other components because they are not necessary, although this may be incorrect.At the same time, there occur steps giving reaction (10) (though we have no evidence about them) which can be hindered by the removal of iodine2T 21 as soon as it is formed. Steps A are divided into steps B with hydrogen peroxide, considered also by Peard and Cullis22 (who demonstrated that it is possible to obtain the experimental kinetic law, by applying the steady-state method to steps B) and not confuted by N O Y ~ S , , ~ ~ 25 and steps C with iodide. Steps B with hydrogen peroxide [+IO, + H,O, -+ + I 0 + H,O + O,] x 2 [ + I 0 + H,O, + +I + H,O + O,] x 2 +I + H,O, + I- + 2H+ + 0, I+ + I- $ I, 47 F A R 1141 8 OSCILLATORY PHENOMENA which give 2+10,+ 5H202 -+ I,+ 4H20 4- 5 0 2 + 2H+. This gives reaction (1 1) on the addition of two reactions (14).Steps C with iodide +IO, + I- e I+ + 10, + I 0 + I- e I++ IO- (23) (24) (22) Steps C do not alter the stoichiometries of reaction (11). The kinetics and equilibrium parameters of these steps are not available, but we can say that steps C are faster than steps B since iodate reduction by iodide is faster than iodate reduction by hydrogen peroxide. has a dissociation constant2, Hence reaction (- 19) has an equilibrium constant This allows us to assume I+ to be the prevailing form of hypoiodous acid. Reaction (22) has an equilibrium constant,' in aqueous medium K,, = (1.2 x 10-l1)-l. Reaction (22) is faster than both reactions (23) and (24) if reaction (22) is far from its equilibrium, whilst both reactions (23) and (24) become faster than reaction (22) when reaction (22) is reaching its equilibrium and reactions (23) and (24) are not reaching theirs.I+ + I- s I,. HI0 e If +OH- (26) (27) (28) The reaction K,, = 3.2 x 10-l'. K( - 19) = [I+]/[HIO] [H+] = 3.2 x lo4. DERIVING GOVERNING EQUATIONS Equations expressing the kinetics of phase exchange in this system are more complicated than those in the theoretical model. We must consider phase exchange of oxygen from the solution to the gas above it, expressed by (29) where dN,,/dt(p.e. 1,l) is the molar discharge of oxygen from the solution to the gas above it, Kl is the coefficient of phase exchange, expressed by the penetration theory, [O,] is the concentration of oxygen inside the solution, [O,], is the concentration of oxygen inside the solution which would be in equilibrium with that in the gas above it, and S is the surface area available for phase exchange.We must next consider phase exchange of oxygen caused by continuous nucleation of some bubbles, as noted by Bray., This is expressed by -dNo2/dt(p.e* 191) = K1([0,1- [OJe) S where po,,mol is the molar density of oxygen inside a bubble, V, is the volume of a bubble, Vis the volume of the solution, no is the specific nucleation (number of bubbles generated per unit volume) per unit time when the homogeneous concentration of oxygen inside the solution has a value that allows equilibrium with bubbles at nucleation and pO2[0,] is the probability of nucleation expressed as the product of the concentration of oxygen and the coefficient of probability at that concentration. We must also consider phase exchange of oxygen from the solution to the bubbles. This is expressed by (31) -dNo2/dt(~-e* 1 9 3 ) = K3 %([02l- [02le, b) noh ~ ~ 0 , [ 0 2 1 / 2 ~F.D’ALBA AND s. DI LORENZO 1419 where K3 is the coefficient of phase exchange expressed by the penetration theory, [O,] is the concentration of oxygen inside the solution, [O,],.b is the concentration of oxygen equilibrating with that in the bubbles and nohpoz[02]/2v is the average number of bubbles per unit volume, where v is the average ascensional velocity of the bubbles and h is the thickness of the solution. Phase exchange of iodine is represented by several terms. We consider phase exchange of iodine as gas from the solution to the gas above it, expressed by - where the symbols are analogous to those of eqn (29).We next consider phase exchange of iodine as gas from the solution to the bubbles of oxygen, expressed by We also consider phase exchange due to the nucleation of crystals of solid iodine, where the symbols are analogous to those of eqn (30) and [I2], is the concentration of iodine allowing generalizated nucleation. We finally consider phase exchange of solid iodine from the solution to the crystals; this exchange is pulsating since crystals are never present inside the solution. Hence we can writel - dNIZ/dt(p.e* 2,4) = K5 S,([I21 - [I21~) v{ncp[lzlc - a(t - lh)) (35) where K5 is the coefficient of phase exchange, S, is the surface area of a crystal, a is the number of crystals per unit volume that have gone out of the solution per unit time and th is the time of nucleation of the crystals.QUALITATIVE EXPLANATION OF OSCILLATIONS To obtain concentration against time functions we would have to write the balance with respect to oxygen and iodine and the chemical rates of every component that can be derived easily from steps B and C . A computer would be required to resolve the system obtained, putting values in place of the symbols, but this is not possible as most of them are unknown. Hence we attempt to give a qualitative interpretation. At the beginning we have steps with hydrogen peroxide, which are slow but produce iodide by reaction (22). At this time, we have no nucleation of bubbles of oxygen, because reactions (18) and (21) are slow owing to the low concentration of both + I 0 and I+.This is the first part of the induction period, as observed by Bray.2 We then have the steps with iodide. At the start, when the concentration of iodine is low and reaction (22) is far from equilibrium, it is faster than reactions (23) and (24). At this time only reactions (15), (18), (21) and (22) take place. This is the second part of the induction period and there is as yet no nucleation of oxygen bubbles or iodine crystals. Reaction (22) then approaches equilibrium. Reactions (23) and (24) become faster than it and hinder reactions (15) and (18), whilst reaction (21) may become faster. At this time production of oxygen and iodine is very slow, but [+I01 and [I+] are increasing because of reactions (23) and (24). This is the third part of the induction period.The iodine concentration reaches the value [I2], and a generalized nucleation of crystals takes place, suddenly lowering the concentration of iodine. Reaction (22) again becomes faster than reactions (23) and (24), which cannot take place because the iodide concentration is too low to sustain reactions (22), (23) and (24) at the same time. Hence reactions (1 5), (1 8) and (2 1) can take place and are faster than they were 47-21420 OSCILLATORY PHENOMENA in the induction period because of the high values of [+I01 and [I+]. The rate of oxygen production allows the concentration to reach a value where nucleation can take place. The induction period of the iodine oscillations ends at the first nucleation, but the induction period observed by Bray ends shortly after it, because he observed the evolution of bubbles of oxygen.This mechanism explains why Bray2 and Peard22 observed an increase in oxygen evolution shortly after the concentration of iodine reaches its maximum. Peard2, thought that iodine initiates a process producing oxygen, whereas our model shows that the decrease of iodine concentration hinders some processes inhibiting oxygen production. The decrease of iodine concentration allows reaction (22) to occur and the phenomenon starts again. INFLUENCE OF X- (X = C1, Br OR F) ON THE OSCILLATIONS CONSIDERED The influence of halogenides on the oscillations is good evidence for the validity of our model. Degn28 found that increasing concentrations of chloride and bromide ions decrease both the maximum concentration of iodine at the beginning of every oscillation and the induction time, whilst an increasing concentration of fluoride ion increases both the maximum concentration of iodine at the beginning of every oscillation and the induction time.We believe that the hydrogen peroxide oxidizes chloride to chlorine and bromide to bromine: 2C1-+ H,O, + 2H+ + C1, + 2H,O 2Br-+ H,O, + 2H+ -+ Br, + 2H,O. (36) (37) Both the nucleation of gas bubbles of chlorine and liquid bubbles of bromine are faster than that of solid crystals of iodine, since several steps are required to produce iodine. Oscillations in the concentration of chlorine or bromine take place according to our model and a pulsating phase exchange of iodine occurs from the solution to the bubbles of chlorine or bromine, allowing oscillations in the concentration of iodine.This does not allow the iodine concentration to increase to the value allowing its nucleation and hence the concentration of iodine, at the beginning of every induced oscillation, is lower. The induction time decreases because the nucleation of both bromine and chlorine are faster than that of iodine. Fluoride cannot be oxidized by hydrogen peroxide and cannot give analogous phenomena. It remains unchanged inside the solution and we may have I, + F- + 12F- (38) or other absorption phenomena, which hinder the formation crystallization nuclei because of electrostatic repulsion, and we have a higher supersaturation for nucleating crystals. APPLICATION OF THE MODEL OF THE BRIGGS-RAUSCHER5 SYSTEM The Briggs-Rauscher5 system, consisting of hydrogen peroxide + malonic acid + potassium iodate + manganous sulphate and sulphuric or perchloric acid, is derived from the Bray system.The latter oscillates at 333 K because of d[I,]/dt(chem.) is not capable of leading to [I2], at room temperature, but it is possible to obtain a value of d[I,]/dt(chem.) capable of leading to [I2], by using a catalyst. Manganous ion is not capable of producing oscillation^^^ without malonic a~id,~O-~* which must therefore be present if there are to be oscillations.F. D’ALBA AND s. DI LORENZO 1421 THE NOYES AND FURROW MODEL At first Furrow and NoyeP proposed a model, similar to that proposed by Noyes,ls to explain oscillations in the Belousov-Zhabotinskii system, in accordance with De Kepper31 and C~oke.~,-,~ The results of some of their tests23 obliged them to propose the following scheme:30 The overall reaction 10; + 2H,O, + RH + H+ -+ RI + 20, + 3H,O 10; + 2H,O, + H+ + HI0 + 20, +2H,O HI0 + RH -, RI + H,O.(39) (40) (41) where R is an organic radical, is generated by the process Step (42) is generated by the reaction H+ HIO+I-eI,+H,O RH e enol (43) (44) I, +en01 t RI + I- + H+. Process (41) is generated by two chains. The non-radical chain is H+ 10; +I- t 10; + HI0 H+ HIO, + I- s 2HI0 (45) 2 x (HI0 + H,O, + I- + 0, + H+ + H,O). (47) The radical chain is H+ 2 x (10; + HIO, e 2 * 1 0 , + H,O) 4 x [ 1 0 , + Mn2+ + H,O t HIO, + Mn(OH),+] (49) (50) (51) (52) Furrow and Noyes thought that reactions (41) and (44) decreased the sum H[IO]+2[12]+[I-], but these reactions may serve to increase the value of [I-], which passes through a critical value allowing the system to switch between the radical and the non-radical chains.4 x [Mn(OH),+ + H,O -+ Mn2+ + H,O + HOO .] 2 x [2H00 - -+ H,O, + O,] 2HI0, -+ 10; + HI0 + H+. We do not agree with this scheme. EXPERIMENTAL RESULTS AND DISCUSSION We recorded the potential difference of the galvanic chain Pt(so1.) I Agar-Agar + KNO, I KCl,,,,., I Hg,Cl, I Hg, Pt from the beginning to the end of the oscillations. Fig. 1 shows the curves for the maxima and minima of the potential, referred to the standard hydrogen electrode, as functions of time for different systems. Considering that the range of potentials,1422 OSCILLATORY PHENOMENA 0.90 0.80 0 5 tlmin 10 Fig. 1. Curves of the maxima and minima of the potential, referred to the standard hydrogen electrode, as functions of time.[KIO,] = 7.02 x mol drn-,, [H,O,] = 9.9 x 10-1 mol dm-3, [MnSO,] = 3.47 x lo-, mol dm-3, [H2S0,] = 0.692 x mol drn-,, 0, [malonic acid] = 2.62 x mol dm-3 and ., [malonic acid] = 10.47 x lo-, mol dm-3. referred to the standard hydrogen electrode, measured by us goes from 0.76 to 0.92 V and considering the initial concentrations- of reagents, we can relate the measured potential to the redox couple 12/1-. The maxima of the potential are almost steady, but the minima are not, demonstrating that the system does not have cyclical limiting behaviour, as the Noyes theory requires. The maxima are almost steady because they are related to [I2Ic, which is steady within the probability limit, whilst the minima have no specific significance. INFLUENCE OF ORGANIC SUBSTRATES The tests of Furrow and Noyes with organic Tests with the following system: [IO;] = 0.025 mol drn-,, [H+] = 0.1 rnol dm-,, [H,O,] = 0.1 rnol dm-,, [Mn2+] = 0.002 mol dm-3 and [C6H,0H] = 1 x low4 mol dm-, exhibited an extreme induction period such that over 4 h only 1 cm3 of oxygen was evolved.Then, in a few minutes, the rate of oxygen production increased by two orders of magnitude. A naphth-2-01 concentration of 1 x mol dm-3 causes a 10 min induction period. Their tests with trans-CH,-CH=CH-COOH showed a change in the stoichio- metry from reaction (1 1) to are very interesting. Mn2+ 10; + 2H,O, + H+ + C4H60, --* C,H,O,I + 20, + 2H,O (53) with the production of an iodohydrin.The tests with the aromatics showed the importance of +IO,, which reacts with them instead of hydrogen peroxide, preventing oxygen evolution. The tests with crotonic acid showed the importance of I+, which is the prevailing form of hypoiodous acid under these reaction conditions. A test with acrylamide is fundamental to a challenge of the Furrow and Noyes mechanism. They observed that acrylamide reacts with I+ and does not react with radicals, although it is a specific radical trap. Thus no radical is present in this system.F. D’ALBA AND s. DI LORENZO 1423 INFLUENCE OF MALONIC ACID Because of the absence of radicals, oxidation of malonic acid by iodate must occur Malonic acid and the manganous ion can give a complex, although we do not know by a two-electron process.how many malonates are bonded: An analogous complex with cerium was considered by Iwo and no ye^^^ for the Belousov-Zhabotinskii system. The complex can react with anionic iodine compounds in the following way: 10; + 10; + coz + Mn I L O J I I I I (55) [+IO], [I+] and [I-] increase and d[I,]/dt(chem.) increases, leading to a value of [I2Ic capable of causing nucleation. COMMENTS ON THE AsO;/IO;/ClO; SYSTEM De Kepper et al.36 have demonstrated that the arsenite, iodate and chlorite system shows oscillating behaviour. They considered the following reactions : 3H,As03 + 10; + 3H3As0, + I- 4H’ + C10; + 41- +2H,O + C1- + 21, (56) (57) and thought that these reactions explained the oscillations, whilst the arsenite and iodate subsystem does not oscillate. We think that reaction (57) is necessary to obtain d[I,]/dt(chem.) capable of leading to [I2Ic.F. D’Alba and S. Di Lorenzo, J . Chem. SOC., Faraday Trans. I , 1983,79, 39. W. C. Bray, J . Am. Chem. SOC., 1921, 43, 1262. B. P. Belousov, J . Res. Radiat. Med., 1959, 1958, 145. J. S. Morgan, J . Chem. SOC., 1916, 109; 274. T. S. Briggs and W. C. Rauscher, J. Chem. Educ., 1973,50, 496. A. Lokta, J. Phys. Chem., 1910, 14, 271.1424 OSCILLATORY PHENOMENA A. Lokta, Proc. Natl Acad. Sci. USA, 1920, 6,410. A. Lokta, J. Am. Chem. SOC., 1920,42, 1395. A. M. Turing, Philos. Trans. R. SOC. London, Ser. B, 1952, 37, 237. lo I. Prigogine and R. Lefever, J. Chem. Phys., 1968,48,795. l1 B. B. Edelstein, J. Theor. Biol., 1970, 29, 57. l2 C. Vidal, C. R.Acad. Sci., Ser. C, 1972, 275, 523. l3 C. Vidal, C. R. Acad. Sci., Ser. C, 1972, 271, 1713. l4 R. J. Field and R. M. Noyes, J. Chem. Phys., 1974,60, 1877. l5 P. Hanusse and A. Pacault, Proc. 25th Znt. Meeting Societi. de Chimie Physique, 1974 (Elsevier, Amsterdam, 1975), p. 50. J. Boissonade, J. Chem. Phys., 1976, 73, 540. (Wiley-Interscience, New York, 1971). l7 P. GlandsdorfF and I. Prigogine, Thermodynamic Theory of Structure, Stability and Fluctuations l8 R. J. Field, E. Koros and R. M. Noyes, J. Am. Chem. SOC., 1972,94, 8649. l9 P. G. Bowers and R. M. Noyes, J. Am. Chem. SOC., 1983, 105, 2572. 2o K. Showalter and R. M. Noyes, J. Am. Chem. SOC., 1978, 100, 1042. 21 W. C. Bray and H. A. Liebhafsky, J. Am. Chem. SOC., 1931,53, 58. 22 M. G. Peard and C. F. Cullis, Trans. Faraday SOC., 1951,47, 616. 23 S. D. Furrow and R. M. Noyes, J. Am. Chem. SOC., 1982, 104,42. 24 K. R. Sharma and R. M. Noyes, J. Am, Chem. SOC., 1976,98,4345. 25 D. Edelson and R. M. Noyes, J. Phys. Chem., 1979,83, 213. 26 A. F. Clifford, Inorganic Chemistry of Qualitatioe Analysis (Prentice-Hall, Englewood Cliffs, N. J., 27 R. P. Bell and E. Gelles, J. Chem. SOC., 1951, 2734. 28 H. Degn, Acta Chem. Scand., 1967, 21, 1057. 2g S. D. Furrow and R. M. Noyes, J. Am. Chem. SOC., 1982, 104, 38. 30 S. D. Furrow and R. M. Noyes, J. Am. Chem. SOC., 1982, 104, 45. 31 P. De Kepper and I. R. Epstein, J. Am. Chem. SOC., 1982, 104,49. 32 D. 0. Cooke, Inorg. Chim. Acta, 1979, 37, 259. 33 D. 0. Cooke, Znt. J. Chem. Kinet., 1980, 12, 671. 34 D. 0. Cooke, Znt. J. Chem. Kinet., 1980, 12, 683. 35 J. J. Iwo and R. M. Noyes, J. Am. Chem. SOC., 1975,97, 5422. 36 P. De Kepper, I. R. Epstein and K. Kustin, J. Am. Chem. SOC., 1981, 103, 2133. 1961). (PAPER 3/990)

ISSN:0300-9599    

DOI:10.1039/F19848001415

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I , 1984,80, 1425-1434 Order Effects in the Excess Thermodynamic Properties of Benzene + Alkane Mixtures BY RAMON G. RUBIO, CARLOS MENDUIGA AND MATEO DIAZ hfi~ Departamento de Quimica Fisica, Facultad de Quimica, Universidad Complutense, Madrid 3, Spain AND JUAN A. R. RENUNCIO* Catedra de Quimica General, Facultad de Quimica, Universidad de Oviedo, Oviedo, Spain Received 1st July, 1983 Mixtures of benzene and an alkane have been analysed using the method of Patterson. Order contributions in pure benzene, pure n-alkane and benzene + n-alkane mixtures and steric- hindrance effects in some benzene + branched-alkane mixtures are able to explain qualitatively the excess thermodynamic properties of these systems. Order contributions are studied in terms of the model of Heintz and Lichtenthaler.It is shown how the value of XI, obtained from H E data may be predicted from the value of XI, obtained from GE data for a system without order contributions. Solutions containing n-alkanes have been studied frequently. Their thermodynamic properties are considered to be the sum of several contributions: (a) combinatorial, (b) energetic, (c) free volume, ( d ) molecular order and (e) steric hindrance.1*2 Contributions (a)-(c) have usually been calculated using the Prigogine-Flory- Patterson (PFP) model. The thermodynamic behaviour of alkane solutions containing cyclohexane (CC,), n-hexane (nC,) or 2,2-dimethylbutane (22DMB) has been explained successfully in terms of these contributions. However, this is not the case for solutions consisting of an n-alkane and 1 -chloronaphthalene, 1,2,3-trichlorobenzene or cyclic ether^.^ For some of these solutions the value of the heat of mixing (HE) decreases as the number of carbon atoms increases.Recently, Rodriguez and Patterson4 and Grolier et 0 1 . ~ have studied the heat capacities of benzene+alkane solutions. They concluded that the five contributions (a)-(e) are not able to explain the large negative value of the excess heat capacity (CpE) and the experimental value of dC,E/dT for the benzene + 2,2,4-trimethylpentane (224TMP) system, which are different from those obtained for mixtures of the same alkane with cC,, nC, or 22DMB. On the other hand, HE values for some benzene + branched-alkane (br-alkane) mixtures are higher than those obtained for mixtures consisting of benzene and the corresponding normal alkane.It has been suggested that these changes in HE are caused by the different interaction energies of the benzene-methyl and benzene- methylene pairs.4 However, this would lead to high HE values for any benzene+ br-alkane mixture. Experimental HE and excess volume ( VE) data for benzene + n-alkane systems have been obtained in our laboratory.' In order to complete our study of these systems we have recently evaluated their excess Gibbs energies (GE).* In this paper we illustrate the importance of order effects on the thermodynamic behaviour of benzene + alkane mixtures. 1425Table 1. Coefficients of eqn (1) (in cm3 mol-l or J mol-l) and X,, parameter (in J c ~ n - ~ ) for benzene + n-alkane mixtures system A2 A3 A4 4 4 XI, ref.benzene + A0 A , n-hexane n-octane n-decane n-dodecane n-tetradecane n- hexadecane n-hexane n-octane n-decane n-dodecane n-tetradecane n-hexadecane n-octadecane n -pen t ane n-hexane n-heptane n-octane n-decane n-dodecane n-tetradecane n-pentadecane n-hexadecane 22DMBa 23DMBb 3MHe 24DMP" 223TMB" 224TMP' 1.5798 2.9009 3.4192 3.761 8 4.0889 4.4639 I A88 2.78 12 3.2507 3.5624 3.7430 3.8682 3.8740 0.6769 0.6208 0.571 1 0.5667 0.5866 0.548 1 0.4466 0.3042 0.1047 0.1261 0.7974 0.7096 0.6123 0.8084 0.61 57 0.6600 0.3585 0.8569 1.2367 I .7166 2.1908 2.2675 0.08 19 0.5734 1.0024 1.3861 1.6863 1.9749 2.0709 0.0309 0.1306 0.1721 -0.1519 0.0800 -0.1279 -0.0695 - 0.1740 0.1442 0.1735 0.0902 0.1221 0.1702 0.161 1 0.1724 0.0843 - 0.4360 - 0.7334 0.4260 - 0.1473 0.2056 0.3944 0.5499 0.8522 1.1014 1.2287 0.0295 0.0159 0.0241 0.0422 - - - - 0.1751 0.1769 0.0542 -0.01 18 0.0805 0.068 1 0.0444 - VE (298.15 K) - - I - - - VE (323.15 K) 0.0357 0.1934 0.2 186 0.7668 0.9230 0.3564 0.8773 GE (298.15 K) - - - - - - - - 0.1506 - - - - - - - 31.43 30.10 29.60 3 1.03 33.18 35.98 29.23 23.95 21.31 22.17 23.03 23.03 23.48 27.6 25.1 25.5 26.1 27.3 28.9 27.9 23.3 23.5 22.6 36.7 32.7 26.5 34.2 27.0 32.5 7c 7c 7c 7c 7c 7c 7c 7c 7c 7c 7c 7c 7c 9a 9a 9a 9c 9a 8 8 8 8 8 9a 9a 9a 9a 9a 9an-hexane n-heptane n-octane n-decane n-dodecane n-tetradecane n -pen t adecane n-hexadecane n-heptadecane 224TMP 0.5133 0.4532 0.4808 0.3599 0.2341 0.1633 0.070 1 0.0292 0.5190 .O.1338 224TMP 0.4238 n-pentane n-hexane n-heptane n-octane n-nonane n-decane n-undecane n-dodecane n-tridecane n-tetradecdne n-pen tadecane n-hexadecane n- heptadecane n-hexane n-octane n-decane n-dodecane n-tetradecane n-hexadecane n-octadecane 1.3836 I .4334 1.5090 1 S659 1.6112 1.6644 1.7118 1.7783 1.8223 1.9113 1.9443 2.0285 2.0828 1.2085 1.3462 1.4561 1 S258 1.5903 1.6669 1.6868 0.1043 0.1480 0.0847 -0.1407 - - 0.0784 0.1 141 0.1673 0.1432 0.3369 0.3103 0.3371 0.4669 0.5360 0.6122 0.7053 0.8279 0.9826 0.9032 0.9743 0.9678 1.1421 0.21 14 0.3244 0.5238 0.6382 0.7164 0.796 1 0.8429 - 0.0382 0.01 53 0.026 1 0.0484 - - - - - 0.1628 0.0416 0.0406 0.0733 0.0907 0.1329 0.1 154 0.3986 0.3639 0.278 1 0.2730 0.3326 0.5371 0.66 i 3 0.4408 0.0475 0.097 1 0.2051 0.3171 0.4753 0.320 1 - 0.0086 - 0.0083 G” (323.15 K) - - - - - - - - - - GE (348.15 K) H E (298.15 K) - 0.0037 - 0.0762 -0.0152 0.0049 0.1231 0.3693 0.0866 -0.0017 0.0400 0.3308 0.2686 0.4144 0.2060 0.022 1 0.0903 0.0966 0.2182 0.2516 0.2623 0.4107 H E (323.15 K) - - - - - - - - - - - - 0.0506 0.2239 0.2909 0.4993 - -0.3992 - 0.1387 0.1610 0.2757 0.1829 - - 0.4449 - 0.0094 -0.01 16 - 0.0833 0.8353 0.4759 - - 24.2 - - 23.3 - - 25.7 -0.7537 - 24.4 -0.7624 0.1633 22.7 - 1.0956 0.4954 23.1 -0.8404 0.1834 21.5 - 21.4 21 .o - 30.1 - - 0.6226 - - - - 28.4 - 40.9 40.4 40.4 40.8 41.1 41.5 42.0 43.0 43.5 45.1 45.5 46.9 47.9 - 36.1 - 37.4 38.9 - 39.6 40.3 - - 41.5 - - 41.4 - - - - - - - 9a 9a 9a 8 8 8 8 8 9b 9a 9a 7a 9d 7a 9e 9f 7a 7a 9f 7a 7a 7a 7a 7a 7a 7a 7a 7a 7a 7a 10 a 2,2-Dimethylbutane, 2,3-dimethylbutane, 3-methylhexane, 2,4-dimethylpentane, 2,2,3-trimethylbutane, f 2,2,4-trimethylpentane.1428 ORDER EFFECTS IN BENZENE +ALKANE MIXTURES RESULTS AND DISCUSSION For the sake of uniformity, all the available data were fitted to a Pade approximant of the form n x Ai(2~- P / X ( 1 -x) = i-om (1) I + 2 Bj(2~- l)j j-1 where XE represents any of the excess parameters (V*, HE/RT or GE/RT), x is the mole fraction of benzene and Ai and Bi are adjustable parameters.Table 1 summarizes the parameter values. GE data for systems containing alkanes from n-pentane (nC,) to n-octane (nC,) and n-heptadecane (nC,,) were taken from the literat~re.~ HE data for br-alkane systems were estimated from GE data as a function of temperature, except for the system with 224TMP, for which the experimental data were those of Lundberg.lo Fig.1 shows the equimolar values of GE, HE and the entropic term of GE (TSE) at 298.15 K plotted against the alkane order parameter, JZZ, taken from the work of Clement and Bothorel." To avoid confusion, fig. 1 does not show the TSE values for br-alkane systems. Note that the HE curve is almost a straight line. Plots of the HE, GE and TSE data for n-alkanes+cC,, nC, or 22DMB show similar behaviour.12 The plots for TSE and GE in fig. 1 are similar to that for C z in ref. (12). The molecular-order contribution in n-alkanes may explain the variation of HE and TSE in the benzene + n-alkane mixtures. For the systems containing br-alkanes it is possible to distinguish between two different sets of HE values at JZ2 = 0. Some br-alkanes have values above that corresponding to the normal isomer, while other br-alkanes have values which are lower than those of the normal isomer.This unusual behaviour indicates that there is a different interaction energy between the methyl or methylene groups and the aromatic ring.4 Nevertheless, several experimental facts were not taken into account in ref. (4). Depolarized Rayleigh scattering experiments on solutions of benzene and quasi-isotropic m01ecules~~~ l4 indicate that the benzene molecules align themselves perpendicular to each other. This ordering agrees with the results of X-ray diffraction and with theoretical and molecular dynamics studies of liquid benzene.15 HE data for n-alkane mixtures with benzene, toluene or ethylbenzene also provide an indication of the existence of a molecular-order contribution in benzene.16 The maximum value of HE in mixtures of a particular n-alkane (nC,,) and these three solvents is almost the same for the mixtures containing toluene or ethylbenzene and 400 J m o t 1 higher for those containing benzene.The similar values of HE observed for solutions containing toluene or ethylbenzene are expected since the interaction energies of the methyl and methylene groups and the aromatic rings are of similar magnitude and the values of the reduced volumes, 6, of both hydrocarbons are close. Depolarized Rayleigh scattering shows that no molecular order contributions are present in toluene or ethylbenzene. Molecular-order contributions in benzene explain the behaviour described above.The same conclusion is reached after analysing the HE data of trans-decalin +benzene and trans-decalin + toluene mixtures." Depolarized Rayleigh scattering results have been obtained for benzene + n-alkane mixtures and lead to the conclusion that the highest symmetry axes of benzene and the n-alkane tend to be perpendic~1ar.l~ This order correlation has been used successfully to interpret the m* bands of aromatic hydrocarbons when they are mixed with n-alkanes. l8R. G. RUBIO, c. MENDUIGA, M. DIAZ PEGA AND J. A. R. RENUNCIO 1429 1300 1100 9 00 d I - 700 c, w' % 500 300 100 I I I I I 0.0 0.3 0.6 0.9 1.2 1.5 J22 Fig. 1. Equimolar excess magnitudes (0, HE; 0, GE; 0, TSE) plotted against the alkane order parameter. Full lines, n-alkanes; 1, 2,2-dimethylbutane (22DMB); 2, 2-methylpentane (2MP); 3, 2,4-dimethylpentane (24DMP) ; 4, 2,2,4-trimethylpentane (224TMP); 5, 3-methylhexane (3MH); 6, 2,2,3-trimethylbutane (223TMB); 7, 2,3dimethylbutane (23DMB).Molecular-order contributions in pure benzene explain the high values of H E and TSE of its solutions when compared with toluene solutions. Note that free-volume differences are larger for the n-alkane + benzene mixtures than for n-alkane + toluene mixtures. Molecular-order contributions also explain the very different values of the interaction parameters that are obtained from H E and GE data for benzene + 22DMB mixtures, which contrast with the results for cC, solutions where almost the same values are obtained when no molecular-order contributions are present (for1430 1200 1000 * L 800 ; cr \ v, 0 II 3 600 4 00 I I I ORDER EFFECTS IN BENZENE + ALKANE MIXTURES TS 24DMP 223TMB Fig.2. Equimolar magnitudes (HE, GE or TSE) plotted against the molar volume of the br-alkane for hexane and heptane isomers + benzene mixtures. Symbols for the br-alkanes as in fig. 1. cc6 + 22DMB, X12 = 5.1 J ~ m - ~ obtained from GE data and Xlz = 5.4 J C M - ~ obtained from HE data). If we consider that molecular-order effects have a large amount of enthalpic-entropic compensation,19 values of the X12 parameter obtained from HE data are expected to be higher than those calculated from GE data for benzene+ br-alkane mixtures. Order in benzene may also explain the high VE values of benzene + n-alkane mixtures. These values are nearly twice the corresponding values for toluene + n-alkane mixtures.20 Neither free volume nor p* (characteristic pressure in the PFP model) differences between benzene and toluene can explain such a difference in the values of VE.It is known that the disruption of order which occurs during the mixing process produces a positive contribution to VE, thus explaining the large VE values of the benzene systems.21 Finally, molecular-order contributions in benzene would explain the large negative values of C: in benzene + br-alkane mixtures which cannot be explained by the free-volume contribution alone. The order contribution to dC,/dTis positive,22 so the disruption of benzene order would be compatible with the positive values of dC,E/dT suggested for the benzene + 224TMP m i x t ~ r e .~ Order correlations between benzene and the n-alkanes would be equivalent to a small alteration of benzene order during mixing, thus explaining why H E and TSE values for some br-alkanes are larger than those for the corresponding normal isomers. Fig. 2 shows the equimolar H E and GE values plotted against the molar volume of hexane and heptane isomers. Barbe and Patterson1 found similar correlations for alkane+cC,, nC, and 22DMB. Small molar volumes are indicative of high stericR. G . RUBIO, c. MENDUI~A, M. DIAZ PEGA AND J. A. R. RENUNCIO 143 1 51- 1 2 3 4 5 6 HEv/ J .cm -3 Fig. 3. TSE order contribution plotted against H E order contribution per unit of core volume at segment fraction 0.5: A, br-alkane systems; 0, n-alkane systems at 298.15 K; 0, n-alkane systems at 323.15 K. hindrance in the alkane.Steric-hindrance effects lead to large negative contributions to H E and 7'SE.l This effect explains why the H E and TSE values for some br-alkane systems are smaller than those observed for mixtures containing either the normal isomer or a br-alkane of low steric hindrance. GE values for the n-alkane systems are smaller than those for the br-alkane systems. However, previous resultsga obtained for systems with steric-hindrance effects indicate that GE values are smaller for these systems than for those without steric-hindrance effects. On the other hand, GE values decrease as the number of carbon atoms increases in the n-alkane systems. This behaviour is different from that of the n-alkane+cC, and may be due to the compensating effects between all the contributions discussed above.In order to isolate the molecular-order and steric-hindrance contributions from the free-volume contribution, we have used the PFP model. Even though the quadrupolar moment of benzene leads to a non-spherical-symmetry force field around the benzene molecule, we chose the PFP model because it has been widely used in previous calculations, thus allowing a direct comparison of our results. Calculations carried out with non-random models24 or lattice-type models25 lead to similar results. Since the molecular-order contribution decreases exponentially with increasing temperature and there is a high degree of enthalpic-entropic compensation in GE data, it is possible to estimate the X g interaction parameter (without molecular-order or steric-hindrance contributions) from GE data for the benzene + 224TMP at 348.15 K (the boiling points of benzene and 224TMP are 353 and 363 K, respectively). A value of 28.4 0.2 J cmP3 was obtained for X g .Croncher and Patterson26 estimated that the ability of 22DMB and 224TMP to destroy the molecular order is similar, and con- sequently both alkanes may be used to estimate X g with similar results. Fig. 3 shows the order contribution to H E plotted against the order contribution1432 ORDER EFFECTS IN BENZENE + ALKANE MIXTURES I I 1 I I 5 7 9 11 13 15 17 19 n Fig. 4. XI, parameter plotted against the number of carbon atoms for n-alkane+ benzene mixtures obtained from: 0, HE data at 298.15 K; +, Cp” data at 310.15 K; 0, HE data at 298.15 K; ---, calculated according to the Heintz and Lichtenthaler model.to TSE per unit of core volume, at segment fraction 0.5. Note that even though the order contributions arise from different effects for each system, all the points follow the same pattern. This result agrees with that obtained previously for other systems.22 Patterson and Barbe2’ proposed an empirical dependence of the order contribution on temperature, which led to the following relationship where the subscript o refers to the order contribution and z is the characteristic temperature at which the ordered substance undergoes the order-disorder transition. A plot of H,E against TS,E is a straight line. For benzene + br-alkane mixtures we obtain a value of 102 K for z. This value is higher than the values of 90 and 70 K12 obtained for n-alkane systems.Bendler2* proposed a fluctuation model, from which the following expression is derived = - l + T * T / GOE TS,E (3) where the meaning of T* is similar to that of z. For the benzene + 224TMP mixture at 298.15 K we have obtained the values - 0.97 for the zero intercept and 667 K for T*. This value of T* is unrealistic as the critical temperature of benzene is 562 K.29 Heintz and Lichtenthaler30 introduced an order contribution to the PFP model.R. G. RUBIO, c . MENDU~~A, M. DIAZ P E ~ A AND J. A. R. RENUNCIO 1433 According to this model, the X12 parameter is given by T 10 x,, = xg+ Y- T - T, (4) where X g is the interaction parameter in the original PFP formulation, T, is the order- disorder transition temperature, and Y is given by where si, v" and v* have the same meanings as in the original PFP model and N* is related to the number of correlated segments.We have calculated Y and T, for benzene from the benzene+224TMP data using eqn (5) and we obtained values of 280 K for T, and of 0.26 J ~ m - ~ for Y. The T, value may be identified with the melting point of benzene (278 K) within the precision of the model. Eqn (4) can be easily extended to systems containing two ordered components. Assuming for X g the value of 28.4 J ~ m - ~ obtained above, and for Y and To of benzene the values 0.26 J and 280 K, respectively, for Y of the n-alkanes the value 3.5 J ~ m - ~ and for T, of the n-alkanes the values given by Heintz and Li~htenthaler,~~ we have calculated the values of A',, for benzene + n-alkane mixtures at 298.15 and 323.15 K.Fig. 4 shows the values of X12 obtained by this method together with the X12 values obtained from HE data. It also shows the X12 values obtained from C? data at 3 10.15 K4 The agreement is satisfactory, particularly at 323.15 K. The discrepancy for nC16 is caused by the fact that the T, value reported for nC16 is smaller than that for nC,,. There is no explanation for this behaviour in ref. (30). The addition of an order contribution from benzene-alkane pairs would lead to better agreement between experimental and calculated X12 values for hydrocarbons with eight or more carbon atoms, but an unrealistic positive contribution would be necessary for lighter hydrocarbons. An expression for the residual Gibbs energy in which the order contribution affects only to the enthalpic term can be derived from the present model.Therefore, the predicted TSE values for n-alkane systems are smaller than the experimental values, and consequently the GE values are larger than those of table 1. The volume dependence of the interaction energy assumed in ref. (30) does not introduce any change in Flory's original equation of state, and the calculated values of VE are smaller than the experimental values. We conclude that the Heintz and Lichtenthaler model is able to explain the order of magnitude of the differences in the X12 values obtained from HE and GE data. However, the model is not able to reproduce the experimental values of GE and VE. On the other hand, the calculations of Patterson et al.seem to explain qualitatively the behaviour of benzene + alkane mixtures without the necessity of assuming different interaction energies for methyl and methylene groups. M. Barbe and D. Patterson, J . Solution Chem., 1980, 9, 753. A. Heintz and R. Lichtenthaler, Angew. Chem., Int. Ed. Engl., 1982, 21, 184. (a) J. P. E. Grolier, A. Inglese, A. H. Roux and E. Wilhelm, Ber. Bunsenges. Phys. Chem., 1981, 85, 768; (b) E. Wilhelm, Ber. Bunsenges. Phys. Chem., 1977, 81, 1150; (c) A. Inglese, E. Wilhelm, J. P. E. Grolier and H. V. Kehiaian, J. Chem. Thermodyn., 1980, 12, 217. A. T. Rodriguez and D. Patterson, J . Chem. SOC., Faraday Trans. 2, 1982,78, 917. J. P. E. Grolier, A. Faradjzeideh and H. V. Kehiaian, Thermochim.Acta, 1982, 53, 157. P. Tancrede, P. Bothorei, P. St. Romain and D. Patterson, J. Chem. SOC., Faraday Trans. 2, 1977,73, 15.1434 ORDER EFFECTS IN BENZENE + ALKANE MIXTURES (a) M. Diaz Peiia and C. Menduiiia, J. Chem. Thermodyn., 1974, 6, 387; 1974, 6, 1097; (b) M. Diaz Peiia and J. Nuiiez, An. Quim., 1974,70,678; J. Chem. Thermodyn., 1975,71,456; (c) M. Diaz Peiia, G. Tardajos, R. L. de Arenosa and C. Menduiiia, J. Chem. Thermodyn., 1979, 11, 951. R. G. Rubio, J. A. R. Renuncio and M. Diaz Peiia, J. Chem. Thermodyn., 1982,14,983; Thermochim. Acta, 1982,56, 199; J . Solution Chem., 1982, 11, 823; Int. J. Thermophys., 1982,3, 325; Thermochim. Acta, 1983, 65, 69. (a) E. W. Funk and J. M. Prausnitz, Ind. Eng. Chem., 1970, 62, 8; (6) M. Ratzsch and G. Krahn, J.Polym. Sci., Polym. Symp., 1973,42,1001; (c) K. R. Harris and J. P. Dunlop, J. Chem. Thermodyn., 1970, 2, 805; ( d ) H. K. D. Jones and B. C. Y. Lu, J. Chem. Eng. Data, 1966, 11, 488; (e) H. W. Schnaible, Ph.D. Thesis (Purdue University, Indiana, 1955); cf) H. K. D. Jones, D. P. L. Poon, R. F. Lama and B. C. Y. Lu, Can. J. Chem. Eng., 1967,45,22. lo G. W. Lundberg, J. Chem. Eng. Data, 1964, 9, 193. l 1 C. Clement and P. Bothorel, J. Chim. Phys., 1964, 61, 878. l2 S. N. Bhattacharyya and D. Patterson, J. Solution Chem., 1980, 9, 247. l3 C. Clement, J. Chim. Phys., 1978, 75, 747. l4 (a) D. R. Bauer, J. I. Brauman and R. Pecora, J. Chem. Phys., 1975,63,53; Annu. Rev. Phys. Chem., 1976,27,443; (b) F. L. Swinton, in Chemical Thermodynamics (The Chemical Society, London, 1978), vol.2. l5 H. Versmold, in Organic Liquids: Structure, Dynamicsand Chemical Properties, ed. A. D. Buckingham, E. Lippert and S. Bratos (Wiley, New York, 1978); B. J. Berne and R. Pecora, Dynamic Light Scattering with Applications to Chemistry, Biology and Physics (Wiley, New York, 1976). I. Fujihara, M. Kobayashi and S. Murakami, J. Chem. Thermodyn., 1983, 15, 1 . Mantione and P. Claverie, Chem. Phys. Lett., 1974, 27, 515. l6 R. L. de Arenosa, Ph.D. Thesis (Universidad Complutense, Madrid, 1976). l8 M. Lamotte and J. Joussott-Dubien, Chem. Phys., 1973,2,245; M . Lamotte, J. Joussott-Dubien, M. J. l9 M. Barbe and D. Patterson, J. Phys. Chem., 1978, 82, 40. 2o M. Caceres Alonso, J. L. Poveda Vilches, R. G. Sanchez-Pajares and J. Nuiiez Delgado, J. Chem. 21 G. Delmas, P. St. Romain and P. Purves, J. Chem. Soc., Faraday Trans. 1, 1975, 71, 1181. 22 S. N. Bhattacharyya and D. Patterson, J. Phys. Chem., 1979, 63, 2979. 23 I. P. C. Li, B. C.-Y. Lu and E. C. Chen, J. Chem. Eng. Data, 1973,18,305; T. Katayama, E. K. Sung and E. N. Lightfoot, AIChE J., 1965,11,924; J. D. Gomez-Ibaiiez, J. Shieh and E. M. Thorteinson, J. Phys. Chem., 1966, 70, 1998; J. D. Gomez-Ibaiiez and J. Shieh, J. Phys. Chem., 1965, 69, 1660. 24 C. Panayiotou and J. H. Vera, Fluid Phase Equilibria, 1980,5,55; M . D. Donohue and J. M. Prausnitz, AIChE J., 1978, 24, 849. 25 I. Sanchez and R. H. Lacombe, J. Phys. Chem., 1976,80, 2352. 26 M. D. Croucher and D. Patterson, J. Chem. Soc., Faraday Trans. 2, 1974, 70, 1479. 27 D. Patterson and M. Barbe, J. Phys. Chem., 1976, 80, 2435. 28 J. T. Bendler, Macromolecules, 1977, 10, 162. 29 R. R. Dreisbach, Physical Properties of Chemical Compounds (American Chemical Society, 30 A. Heintz and R. Lichtenthaler, Ber. Bunsenges. Phys. Chem., 1980, 84, 890. Thermodyn., 1983, 15, 913. Washington, D.C., 1959), vol. 2. (PAPER 3/1135)

ISSN:0300-9599    

DOI:10.1039/F19848001425

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1984,80, 1435-1446 Catalysis by Amorphous Metal Alloys Part 1.-Hydrogenation of Olefins over Amorphous Ni-P and Ni-B Alloys BY SATOHIRO YOSHIDA,* HIROMI YAMASHITA, TAKUZO FUNABIKI* AND TEIJIRO YONEZAWA Department of Hydrocarbon Chemistry, Faculty of Engineering, Kyoto University, Kyoto, Japan Received 12th July, 1983 Amorphous Ni-P (Ni81P19) and Ni-B (Nie2BS8) alloys have been prepared by the rapid quenching method. The untreated alloys are inactive for the hydrogenation of olefins, but successive pretreatments of the alloys with the dilute HNO,, oxygen and hydrogen bring about greater catalytic activity than that possessed by the crystalline alloys. Measurements of the ESCA spectra of the surfaces of the alloys have indicated that the alloys become active by the partial oxidation of the nickel and phosphorus or boron on the surface.Buta-1,3-diene is hydrogenated to give isomeric butenes of constant composition, and butenes isomerise rapidly to give a similar composition to that obtained on the hydrogenation of buta-1,3-diene. The hydrogenation of buta-1,3-diene is first and zeroth order with respect to the partial pressures of hydrogen and buta-l,3-diene, respectively. These results are explained by assuming that the hydrogen and the olefins are not competing for adsorption on the same sites and that hydrogen is activated on electron-deficient nickel atoms formed by electron transfer from nickel to the oxidised species. Much work has been done on the catalytic reactions over crystalline surfaces of metal alloys, but little has been done on the catalytic reactions over amorphous alloys whose surfaces are considered to be homogene~us.~-* It has been reported that such catalysts are formed in an amorphous state when the reduction of Ni-P and Ni-B alloys, known to be hydrogenation catalysts, is performed under mild condition^,^? but the homogeneity of the state and its role in catalysis are uncertain.It is very probable that the catalytic behaviour of amorphous alloys is very different from that of crystalline alloys, because the surface electronic state and the chemical composition may be different. We have previously reported’ that amorphous Ni-P and Ni-B alloys prepared by the rapid-quenching method are highly active for the hydrogenation of olefins in comparison with crystalline Ni-P and Ni-B alloys, and that the pretreatment conditions are very important.We report here the catalytic activity of the amorphous alloys for the hydrogenation of olefins and the relationship between the catalytic activity and the changes on the surface caused by the pretreatment, studied using ESCA (electron spectroscopy for chemical analysis). EXPERIMENT PREPARATION OF CATALYSTS Amorphous Ni-P (NiE1Pl9) and Ni-B (Nie2B3,J alloys were prepared by the rapid-quenching method8 using a single steel roll as reported previ~usly.~ Ni (99.5 wt %), Ni-P (P, 14 wt %) and Ni-B (B, 15 wt %) were used in the preparation and amorphous alloys in the shape of ribbons ca. 5 mm wide and 10-20 pm thick were obtained. An X-ray diffraction analysis (X.r.d.) was used to verify the amorphous state of the alloys2 before and after the reactions.The surface areas of the alloys were measured by the Brunauer-Emmett-Teller (B.E.T.) method using krypton physisorption at 77 Kg 14351436 CATALYSIS BY AMORPHOUS METAL ALLOYS PRETREATMENT OF CATALYSTS AND HYDROGENATION OF OLEFINS Hydrogen was purified by a hydrogen diffusion purifier (Japan Pure Hydrogen Co., LS-OOB) and commercial oxygen and olefins were used without further purification. The alloys were treated with HNO, (1.5 and 6 mol dm-3 HNO, for Ni-B and Ni-P, respectively) washed thoroughly with water and dried under air for 24 h. Cu. 0.5 g of the alloy was placed in a glass vessel which was connected to a conventional closed circulation system (total volume 210 cm3).The alloy was heated under circulating oxygen (6.7 kPa) and evacuated for 0.5 h. The alloy was then heated at 573 K under circulating hydrogen (13.3 kPa) and evacuated for 1 h. Hydrogen (pH2 = 19.3 kPa) and olefin (polefin = 7.3 kPa) were introduced to the reaction vessel and the reaction took placed under the circulating gases. The reaction was followed by measuring the pressure change and analysing the composition of the product by g.l.c., using a 3 m column of dimethylsulpholane on C22. Initial rates were estimated from the initial pressure change. The pretreatments and reactions were performed under different conditions (temperature and time) depending on the purpose of the experiments; the conditions are given in the captions of the tables and figures.The conditions were selected so as to determine the initial rates as accurately as possible. For example, the conditions for the Ni-P amorphous alloys shown in the tables were chosen so as to be able to determine the rate of highly reactive ethene from the change in the levels of a mercury manometer, and the conditions in fig. 3 were chosen so as to be able to determine the slow rates under low hydrogen pressures. The pretreatment conditions are given in the form (O,, 373 K, 2 h; H,, 573 K, 6 h), which denotes heating under oxygen at 373 K for 2 h, followed by heating under hydrogen at 573 K for 6 h. MEASUREMENTS OF ESCA SPECTRA The effects of the pretreatment and the reaction on the surface structure were studied by measuring the ESCA spectra (Shimadzu ESCA-750) using Mg radiation (10 kV, 30 mA). Amorphous and crystalline alloys were placed in the ESCA analyser chamber and spectra were measured either without sputtering or after sputtering the surface with Ar+ to depths of ca.5 and 25 nm (sputtering time 1 and 5 min, respectively). All binding-energy values (accuracy kO.3 eV) were referred to the value of the contaminant carbon (CIS = 285.0 eV) for convenience. RESULTS SURFACE AREA AND CRYSTALLISATION OF ALLOYS Differential thermal analysis indicated that the amorphous Ni-P and Ni-B alloys started to crystallise at 643 and 783 K, respectively. The crystallised alloys exhibited the X-ray diffraction patterns of crystalline Ni, Ni,,P,, Ni2B, B20, etc. However, X.r.d. indicated that both amorphous alloys continued to be amorphous even after treatment with hydrogen at 573 K for 10 h.The B.E.T. surface areas of the amorphous alloys were 0.12 and 0.094 m2 g-l for Ni-P and Ni-B, respectively, and those of the crystalline alloys were 0.12 and 0.1 1 m2 g-l for Ni-P and Ni-B, respectively. This result indicates that the surface areas of the amorphous and crystalline Ni-P and Ni-B alloys are very similar. HYDROGENATION OF OLEFINS The untreated alloys showed no activity for the hydrogenation of olefins even after hydrogen pretreatment of the alloys at 573 K for 6 h, but the alloys became active after successive treatments with dilute HNO,, oxygen and hydrogen. The three pretreatments were necessary to generate the catalytic activity. Fig. 1 and 2 show the effects of the pretreatments with oxygen and hydrogen, respectively, on the catalytic activity for hydrogenation of buta-1,3-diene.As shown in fig. 1, the catalytic activity is dependent on the temperature of the oxygen pretreatment, and oxygen treatment at a high temperature is necessary for activation of the alloys. In the case of the Ni-P alloy, abrupt enhancement of the activity was observed at 513 K, but the activityS. YOSHIDA, M. YAMASHITA, T. FUNABIKI AND T. YONEZAWA 1437 0 373 473 I 573 1.3 x cd 4 PI I z b) 1.2 2 0 x U .- .* U a W u .- 0.1 2 2 0 temperature of the oxygen pretreatment/K Fig. 1. Effect of the oxygen pretreatment on the catalytic activity for hydrogenation of buta-1,3-diene: 0, amorphous Ni-P; A, amorphous Ni-B. Pretreatments: (O,, 1 h; H,, 573 K, 2 h); reaction temperature: 373 K.Relative activities are the initial rates of hydrogenation normalised to that of amorphous Ni-P pretreated with 0, at 493 K. I .o 0 2 4 6 8 10 time/h Fig. 2. Effect of the hydrogen pretreatment times on the catalytic activity of hydrogenation of buta-1,3-diene: 0, amorphous Ni-P; A, amorphous Ni-B. Pretreatments: Ni-P (O,, 573,2 h; H,, 573 K) and Ni-B (O,, 473 K, 1 h; H,, 573 K); reaction temperature: 473 K (Ni-P) and 423 K (Ni-B). Relative activities are the initial rates normalised to those of Ni-P and Ni-B pretreated by hydrogen for 2 and 6 h, respectively.1438 CATALYSIS BY AMORPHOUS METAL ALLOYS Table 1. Product distribution in the hydrogenation of buta-1,3-diene over Ni-P and Ni-B alloysa products (mol %) conversion cis: trans ratio catalyst (mol %)B butane but- 1 -ene cis-but-2-ene trans-but-2-ene of but-2-ene Ni-P, amor.* 73 1.5 57.7 14.9 25.9 0.58 Ni-P, crystc 77 1.4 59.9 12.7 26.0 0.49 Ni-B, amor.d 73 4.1 52.8 13.2 29.9 0.49 Ni"" 75 12 53 5 31 0.16 Ni,Pe 75 3 32 10 55 0.18 Raney Nif 55-75 66 7 6 21 0.29 a Initial pressure: pH2 = 19.3 kPa andpbutadiene = 7.3 kPa; reaction temperature: 373 K (Ni-P) and 423 K (Ni-B).* Amorphous Ni-P (0.404 g) (02, 523 K, 2 h; H,, 573 K, 2 h). Crystalline Ni-P obtained by heating the above catalyst at 723 K for 0.5 h in uacuo (O,, 523 K, 2 h; H,, 573 K, 2 h). Amorphous Ni-B (0.322 g) (O,, 423 K, 1 h; H,, 573 K, 3 h). eData were estimated from the figure in ref. (6). Ni": alumina-supported nickel phosphate catalyst containing 20 wt % Ni,(PO,), was calcinated at 873 K and reduced at 673 K; Ni,P: the above catalyst reduced at 823 K. f Ref.(1 1). B Based on buta-1,3-diene introduced into the reactor. 10 0 Q) 5 Y 3 P a s 0 0- 0-0- 0.5 1 .o 1.5 1 .o lli E + c, P (u 0 0.5 2 E 0 time/h Fig, 3. Reaction of but-1-ene over amorphous Ni-P: A, but-1-ene; 0, cis-but-2-ene; A, trans-but-2-ene; 0, butane; 0, cis: trans ratio of but-2-enes. Initial pressure: pH2 = 19.3 kPa and pbutene = 7.3 kPa; pretreatment: (O,, 523 K, 2 h; H,, 523 K, 2 h); reaction temperature: 393 K. decreased on treatment at higher temperatures. The activity of the Ni-B alloy increased gradually and reached a maximum at 473 K. The decrease in the activity on treatment at higher temperatures was not significant. Fig. 2 shows the effect of different hydrogen pretreatment times at 573 K.The catalytic activity reaches a maximum after 2 h for Ni-P and after 6 h for Ni-B, and then decreased on prolonged treatment.S. YOSHIDA, M. YAMASHITA, T. FUNABIKI AND T. YONEZAWA 1439 Table 2. Product distribution in the hydrogenation of isoprene over Ni-P and Ni-B alloysa products (mol %) conversion 2-methyl- 2-methyl- 3-methyl- 2-methyl- catalyst (mol %I" butane but-1 -ene but-1-ene but-2-ene Ni-P, amor.b 56 2.7 42.4 13.6 51.3 Ni-B, amor.C 71 2.3 44.9 22.2 30.6 Raney Nid 2 41 16 41 a Initial pressure: pH2 = 19.3 kPa and pisoprene = 7.3 kPa; reaction temperature: 403 K Ref. (17). " Based (Ni-P) and 423 K (Ni-B). b* See footnotes b and din table 1, respectively. on added isoprene introduced to the reactor. Table 3. Initial rates in the hydrogenation of olefins over Ni-P and Ni-B alloysa initial rate/kPa min-' catalyst ethene propene cis-but-2-ene buta- 1,3-diene isoprene Ni-P, amor.b 13.89 3.62 0.37 0.23 0.05 Ni-B, amor.d 0.75 0.57 0.12 0.58 0.3 1 Ni-P, cryst.c 8.79 2.47 0.30 0.08 0.0 1 Ni-B, cryst." c 0.01 - - - < 0.01 a Initial pressure: pH = 19.3 kPa andp,,,,,, = 7.3 kPa; reaction temperature: 373 K (Ni-P) Amorphous Ni-B was and 423 K (Ni-B).b-d See footnotes b-d in table 1, respectively. crystallised by heating at 793 K for 2 h in uacuo. Table 1 shows the product composition for the hydrogenation of buta-1,3-diene catalysed by the amorphous alloys and, for comparison, by other nickel catalysts. Amorphous Ni-P and Ni-B alloys gave similar selectivities and crystallisation did not affect the selectivity. Amorphous alloys formed butenes until the level of conversion reached 95 %, showing that no competitive hydrogenation of butenes occurred in the presence of buta- 1,3-diene.The composition was constant throughout the reaction. Formation of but-1-ene as a main product was similar to that for a nickel catalyst prepared from nickel phosphate on alumina by reduction with hydrogen at 673 K, but different from that for a catalyst reduced at 923 K 6 ~ lo and from that for Raney nicke1.l' The cis: trans ratio for but-2-ene was ca. 0.5 for the amorphous and crystalline alloys, unlike other catalysts. In the absence of buta-1,3-diene, the butenes iso- merised rapidly to give a composition similar to that obtained on the hydrogenation of buta- 1,3-diene.Fig. 3 shows the reaction of but- 1 -ene and indicates that isomerisation proceeds more rapidly than hydrogenation. But-2-enes reacted in a similar way. Isoprene was also hydrogenated, yielding methylbutenes with aconstant composition throughout the reaction; the hydrogenation of methylbutenes did not proceed until the conversion level of isoprene had reached 95%. As shown in table 2, the composition of the products was slightly different for the two amorphous alloys, i.e. 2-methylbut- 1 -ene > 2-methylbut-2-ene > 3-methylbut- 1 -ene (Ni-P) and 2-methyl- but-2-ene > 2-methylbut- 1 -ene > 3-methylbut- 1 -ene (Ni-B). The selectivity was not greatly different from that obtained with Raney nickel.1440 CATALYSIS BY AMORPHOUS METAL ALLOYS I I 6.7 13.3 20.0 PH, Or PC,H6 lkPa 0 Fig.4. Dependence of the initial rates of hydrogenation of buta-1,3-diene on the partial pressures of hydrogen and buta-1,3-diene. Dependence on pC4Hs at !HZ = 19.3 kPa: A, Ni-P; A, Ni-B: Dependence on p H z at pCpHs = 7.3 kPa: a, Ni-P; 0, Ni-B. Pretreatments: Ni-P (02, 523 K, 2 h; 573 K, 2 h) and Ni-B (02, 373 K, 1 h; 573 K, 6 h); hydrogenation temperature: 373 K. Relative rates are the initial rates normalised to those under pHz = 19.3 kPa and pCaHs = 7.3 kPa. Table 3 shows the initial rates of hydrogenation of olefins and diolefins over the amorphous and crystalline Ni-P and Ni-B alloys. The crystalline alloys were less active than the amorphous alloys, and crystalline Ni-B exhibited very little activity under our conditions.The reactivity decreased with increasing methyl substitution, the effect of the methyl substituent being more marked for Ni-P than for Ni-B. A kinetic study of the reaction of buta-1,3-diene was carried out by varying the partial pressures of buta- 1,3-diene and hydrogen. As shown in fig. 4 the relative initial rates can be used to show the dependence on the partial pressures as some variations in the activity were observed on repeated runs. Thus, when the reaction was performed repeatedly without pretreatment with oxygen and hydrogen after each run, the initial rate under the same partial pressures was 80-90% for each run in the first 2 or 3 runs and then became constant. When this catalyst was held at room temperature overnight after evacuation, the rate was 30-50% of that of the fresh catalyst, but the activity was restored to that of the fresh catalyst by pretreatment with oxygen and hydrogen.To correct the variation in the activity caused by repeated use of the same catalyst reaction with 19.3 kPa hydrogen and 7.3 kPa buta-1,3-diene was performed as a standard after each run. In spite of the variation of the magnitude of the standard rate, the ratio of the rates was almost reproducible and the results in fig. 4 indicate that the rates are first and zeroth order with respect to hydrogen and buta-1,3-diene, respectively.S . YOSHIDA, M. YAMASHITA, T. FUNABIKI AND T. YONEZAWA 1441 L -- -- binding energy/eV 860 850 I90 180 860 850 190 180 Fig. 5. ESCA spectra of amorphous Ni-B. Successive treatments: (a) untreated, (b) HNO,, (c) 0, at 373 K for 1 h, (d) H, at 573 K for 6 h, (e) hydrogenation of buta-1,3-diene at 373 K, v) crystallisation at 873 K for 2 h in vacuo.Sputtering: ca. 5 nm; scale of peak intensity ( x , ~ ~ ) : x ( N ~ ~ ~ ~ , ~ ) = (a) 4000, (b) 10000, (c) 4000, ( d ) 2000, (e) 2000, cf) 400; x(B,,) = (a-f) 200. ESCA SPECTRA The changes in the surface states of the amorphous alloys following pretreatment, hydrogenation and crystallisation were studied by ESCA. Fig. 5 shows the spectra of the amorphous Ni-B alloy measured after Ar+ ion sputtering of the surfaces to a depth of ca. 5 nm. The spectrum of the untreated alloy [fig. 5 (a)] shows a band from nickel (2p3,& and two bands from boron (Is). The binding energy of the nickel (852.95 eV) is greater than that of pure nickel (852.2 eV)12 and those of boron (188.20 and 192.20 eV) are also greater than that of pure boron (187.8 eV).12 The latter band, which may correspond to boron bound to oxygen,13 was observed even after sputtering the surface to a depth of ca.25 nm. The spectrum after pretreatment with dilute HNO, [fig. 5(b)] shows a nickel band (852.85 eV) and a boron band (188.10 eV), but the boron oxide band has disappeared. When the spectrum was measured without sputtering of the surface the boron oxide and the oxidised nickel bands were observed at 191.95 and 855.80 eV, respectively. The spectrum after pretreatment with oxygen [fig. 5(c)] shows bands from nickel (853.00 eV), nickel oxide (856.05 eV), boron (1 88.15 eV) and boron oxide (192.20 eV). However, when the spectrum was measured without sputtering, the nickel band was very small.When the oxidised alloy was treated with hydrogen, two bands from nickel (853.15 eV) and nickel oxide (856.60 eV) were observed with little change in the1442 CATALYSIS BY AMORPHOUS METAL ALLOYS L L - - 860 8 50 140 130 binding energy/eV Fig. 6. ESCA spectra of amorphous Ni-P. Pretreatment: (02, 513 K, 1 h; H,, 573, 2 h); sputtering: (a) ca. 25 nm, (b) ca. 5 nm, (c) 0 nm; scale of peak intensity ( x , ~ ~ ) : x ( N ~ ~ ~ ~ , ~ ) = (a) 1000, (b) and (c) = 2000; x(P2J = (a-c) 200. relative intensity. The boron band disappeared and only the boron oxide band (192.80 eV) was observed [fig. 5(d)]. However, when the spectrum was measured without sputtering, a small nickel band was observed, indicating reduction of nickel oxide.Fig. 5 (e), which was measured after hydrogenation of buta- 173-diene, showed the decrease in the intensity of the nickel oxide band (856.80 eV) compared with that of nickel (853.05 eV). The boron band (192.80 eV) remained unaltered. When the alloy was crystallised by pretreatment under severe conditions (873 K, 2 h) in vacuo, the nickel band disappeared and the boron band appeared at higher energy (193.75 eV) [fig. 5 d f ) l . Similar spectral changes were observed with the Ni-P alloy, and bands from nickel (853.35 eV), nickel oxide (856.85 eV), phosphine (129.60 eV) and phosphine oxide (1 33.70 eV) were detected. However, the reactivities of Ni-B and Ni-P with oxygen and hydrogen are different. For example, the nickel band was observed after oxidation even when the spectrum was measured without sputtering, and the nickel oxide band was not observed in the spectrum measured after sputtering to a depth of ca.5 nm. Fig. 6 shows the spectra of the alloy pretreated successively with HNO,, oxygen and hydrogen. The ratio of the peak heights of nickel oxide to nickel decreased withS. YOSHIDA, M. YAMASHITA, T. FUNABIKI AND T. YONEZAWA 1443 increasing time of hydrogen treatment (2.13,O h; 0.69,2 h; 0.35,4 h) but the relative intensities of the peaks of phosphine oxide and phosphine were nearly constant. DISCUSSION We expected that untreated amorphous Ni-P and Ni-B alloys would exhibit catalytic activity for the hydrogenation of olefins, but preliminary experiments indicated that the alloys were inactive, irrespective of the reaction conditions.It was assumed that the surface of the alloys might not be clean as the amorphous films were prepared by rapidly cooling the melted mother alloys in air. However, pretreatment with dilute HNO,, in order to clean the surfaces, or with hydrogen, in order to reduce the oxidised species which might be present, were ineffective. On the contrary, we have found that pretreatment with oxygen at fairly high temperatures is essential for catalytic activity and that successive treatments with acid, oxygen and hydrogen are very effective. This indicates that modifications of the amorphous surface by oxidation and reduction are very important. The results in fig. 1 indicate that the effects of the temperature of pretreatment with oxygen are different for the two alloys.Since the effect appears at a lower temperature for Ni-B than for Ni-P, the oxidation must proceed more readily with Ni-B than Ni-P. Ni-B showed nearly constant activity in the range 473-623 K, but the activity is lower than that of Ni-P. On the other hand, the activity of Ni-P is very sensitive to the pretreatment temperature, and this alloy exhibits very high activity in a very narrow range of temperature. Over-oxidation brings about a large reduction in the activity. The oxygen treatment must form oxygen adducts of nickel and metalloids (phosphine and boron), but the oxidised alloys are not active unless reduced by hydrogen. Fig. 2 shows that the reduction proceeds more readily with Ni-P than with Ni-B, and the activities of both alloys decrease on over-reduction, indicating the importance of the partial reduction.The above results suggest that the catalytic activity of these amorphous alloys requires the presence of both oxidised and reduced species on the surface and that it is dependent on the composition of these two species. This is clearly shown by ESCA. The spectrum of the untreated Ni-B alloy indicates the formation of boron oxide on the surface during the preparation of the alloy [fig. 5(a)]. The binding energies of nickel and boron are shifted from those of pure nickel and boron, indicating an interaction between nickel and boron. The positive shift for boron indicates electron transfer from boron to nickel. On the other hand, the spectrum of the Ni-P alloy shows a negative shift of the phosphorus band compared with pure red phosphorus, indicating electron transfer from nickel to phosphorus.Fig. 5(b) indicates that the metalloid oxides of the untreated alloy are removed by the acid treatment. Since the spectra measured without sputtering shows the bands of oxides of nickel and metalloids, after the acid treatment the surface must be very sensitive to oxygen. Fig. 5(c) indicates that Ni-B is oxidised to form oxides of nickel and boron not only on the surface but also inside the alloy. The nickel species on the surface is mainly nickel oxide because the nickel band is negligible in the spectrum obtained without sputtering. On the other hand, the spectra of Ni-P indicate the presence of both of nickel and nickel oxide on the surface and the absence of nickel oxide after sputtering.These results indicate the more facile oxidation of Ni-B than of Ni-P, which is consistent with the results in fig. 1. The spectra of the alloys measured after hydrogen treatment indicate that hydrogen reduces nickel oxide but not metalloid oxides. The reduction of nickel in Ni-P1444 CATALYSIS BY AMORPHOUS METAL ALLOYS proceeds more readily than that of nickel in Ni-B, which is consistent with the results in fig. 2. The more facile oxidation and the less facile reduction of nickel in Ni-B than in Ni-P seem to be related to the different mode of electron transfer between nickel and the metalloids. Further oxidation of the metalloids proceeds during the reduction and shifts in the binding energies of the oxides of nickel and the metalloids are observed.This is because oxygen bound to nickel is removed on the reduction of the Ni-B alloy with hydrogen and the strength of the boron-oxygen bond increases.12 The results shown by the ESCA spectra indicate that the pretreatments modify the amorphous alloys to form nickel, nickel oxide, metalloid and metalloid oxides, not only on the surface but also inside the alloy. The composition of these species affects the catalytic activity, because it determines the number and electronic state of active nickel species on the surface. The reduced nickel adjacent to oxides of nickel or the metalloids must be more electron deficient than that without oxides as neighbours. Thus, it is important to know the composition which gives the maximum activity and the optimum pretreatment conditions which bring about this composition.The spectrum measured after the hydrogenation of olefin indicates the partial reduction of nickel oxide to nickel, as shown in fig. 5(e). It is probable that the reduction takes place during the hydrogenation, and the change in the composition of species on the surface may result in the observed variation of catalytic activity in repeated hydrogenations. A change in the surface composition also occurs when the alloy is treated under severe conditions so as to cause crystallisation. As shown in fig. 6(f), nickel is transferred from the surface to the interior of the alloy, resulting in a low level of reduced nickel on the surface. The positive shift of the binding energy of boron oxide suggests the formation of stable boron-oxygen compounds such as B203,13 which was detected by X.r.d.The low catalytic activities of the crystalline alloys are caused by these changes in surface composition and structure, as the surface areas are essentially the same for the amorphous and crystalline alloys. A structural change in the amorphous alloy starts at a lower temperature than that estimated by differential thermal analysis and alloys pretreated at high temperatures may undergo structural changes which cannot be detected by X.r.d. The drop of the catalytic activity of Ni-P on treatment at > 533 K under oxygen (fig. 1) seems to be correlated with this type of structural change because the crystallisation temperature of Ni-P is lower than that of Ni-B, but we have not obtained any information about the structures of amorphous alloys in precrystallisation states.The presence of different types of active species is also suggested by the kinetic studies. The rate equation for the hydrogenation of buta- 173-diene is r = kPAZ P! (9 where r, k, p H , and p s denote rate, rate constant and the partial pressures of hydrogen and diene, respectively. This result is well explained by the following reaction sequence: K , Hz 2Hads K3 Hads + Sads f SHads (3) k Hads + SHads SH2 (4)S. YOSHIDA, M. YAMASHITA, T. FUNABIKI AND T. YONEZAWA 1445 assuming that (a) hydrogen and diene do not compete for adsorption on the same site, (b) adsorption of hydrogen is very weak compared with that of diene and (c) desorption of product is very rapid.Assumptions (a) and (b) are suggested by eqn (i)14 and assumption (c) by the result that hydrogenation of diene proceeds without hydrogenation of olefin in the presence of diene. The latter result is explained by the stronger adsorption of diene than olefin, which results in the rapid replacement of adsorbed olefin with diene. With NH andN, denoting the numbers of sites for adsorption of hydrogen and diene, respectively, we obtain ~ NH dK P H , + d K I P H z iHadsl = Then, the rate of hydrogenation of olefins is represented by (ii) (iii) Assumptions (a) and (b), i.e. K2ps 9 1 9 K 1 p H 2 , will lead to which is consistent with eqn (i). Eqn (v) indicates that the rate is affected not only by the partial pressure of hydrogen, but also the amount and electronic states of the active sites. Although we have performed kinetic studies only for the hydrogenation of buta-1,3-diene, it is very probable that eqn (v) is applicable to other olefins.The dependence of the rate on the reactant as shown in table 3 is ascribed not only to the different values of k but also to K3 and Ns, which are affected by the electronic and steric factors of the olefins. It seems reasonable to assume that hydrogen is activated by adsorbing on the electron-deficient nickel atoms and the olefins by adsorbing on the electron-rich nickel atoms. It has been reported that in the hydrogenation of styrene on Ni-B and Ni-P catalysts13 the increased (decreased) electron density on nickel caused by electron transfer from boron to nickel (nickel to phosphorus) weakens (strengthens) the adsorption strength of the reactants, resulting in the reduction (promotion) of self-poisoning effects by the adsorbed reactants.However, stable n-olefin complexes are usually formed with electron-rich low-valent metal complexes, and the stability of the complex is increased (decreased) by electron- withdrawing (donating) substituents on the 01efins.l~ The electronic effects are consistent with the substituent effects in table 3, supporting the explanation of reactivity given by eqn (v) rather than by the self-poising effect. The composition of the butenes produced by the hydrogenation of buta-1,3-diene (but-1-ene > trans-but-2-ene > cis-but-2-ene) and the cis: trans ratio of the but-2-enes (ca. 0.5) suggest that the mechanism involves a number of reversible steps which permit conformational interconversion of the di-n-adsorbed buta- 1,3-diene. l6 The mechanism involves interconversions among half-hydrogenated adsorbed species such as Q- and z-butenyl and n-methylallyl species.The interconversion is very rapid since the composition of the butenes is constant throughout the reaction. Interconversion of these intermediates is involved in the isomerisation of the butenes, because the composition of the butenes obtained on the hydrogenation of buta-l,3-diene is similar1446 CATALYSIS BY AMORPHOUS METAL ALLOYS to that obtained on the isomerisation of the butenes. This indicates that the isomerisation proceeds by the abstraction-addition mechanism via adsorbed n-allylic intermediates rather than the addition-abstraction mechanism via adsorbed a-butyl species.This is consistent with the result that isomerisation of the butenes which does not necessarily require the activation of molecular hydrogen proceeds more rapidly than hydrogenation. Kinetic studies of the isomerisation of the butenes have not been performed, but the rate equation may be different from eqn (v) even if the same intermediates are involved in both the hydrogenation of buta-l,3-diene and the isomerisation of the butenes. In conclusion, amorphous Ni-P and Ni-B alloys prepared by the rapid quenching method become active for the hydrogenation and isomerisation of olefins after pretreatment with acid, oxygen and hydrogen. The pretreatments modify the amor- phous alloys by forming oxides of nickel and metalloids on the surface and inside of alloys.The oxides may be able to form electron-deficient nickel species, and the catalytic activity is dependent on the amount of the nickel species in the different electronic states on the surface. Structural changes which are not detected by X.r.d. may occur in the pretreatment procedures, but the relationship between the catalytic activity and the structural changes has not been clarified in the present study. G. V. Smith, W. E. Brower, M. S. Matyjaszcyk and T. L. Pettit, in Proc. 7th In?. Cong. Catalysis, ed. T. Seiyama and K. Tanabe (Elsevier, Amsterdam, 1981), part 1, p. 335. A. Yokoyama, H. Komiyama, H. Inoue, T. Masumoto and H. M. Kimura, J. Catal., 1981,68, 355. A. Yokoyama, H. Komiyama, H. Inoue, T. Masumoto and H. M. Kimura, Scr. Metall., 1981, 15, 365. A. Yokoyama, H. Komiyama, H. Inoue, T. Masumoto and H. M. Kimura, J. Chem. SOC. Jpn, 1982, 2, 199. R. C. Wade, D. G. Holar, A. N. Hughes and B. C. Hui, Catal. Rev., 1976, 14, 21 1. F. Nozaki and R. Adachi, J. Catal., 1975, 40, 166. 'I S. Yoshida, H. Yamashita, T. Funabiki and T. Yonezawa, J. Chem. SOC., Chem. Commun., 1982,964. * Proc. 3rd In?. Con$ Rapidly Quenched Metals, ed. B. Canter (The Chameleon Press, London, 1978). lo F. Nozaki, T. Kitoh and T. Sodesawa, J. Catal., 1980, 62, 286. R. A. Beebe, J. B. Beckwith and J. M. Honig, J. Am. Chem. SOC., 1945,67, 1554. W. G. Young, R. L. Meier, J. Vinograd, H. Billinger, L. Kaplan and S . L. Linden, J. Am. Chem. SOC., 1947,69, 2046. V. V. Nemoshkalenko, A. I. Kharlamov, T. I. Serebryakova and V. G. Aleshin, Kinet. Katal., 1978, 19, 1567. l3 Y. Okamoto, Y. Nitta, T. Imanaka and S. Teranishi, J. Chem. SOC., Faraday Trans. I , 1979,75,2027. l4 K. Soga, H. Imamura and S. Ikeda, J. Phys. Chem., 1977,81, 1762. l5 E. 0. Fischer and H. Werner, Metal 7r Complexes (Elsevier, Amsterdam, 1966), vol. 1 ; M. Herberhold, Metal 7r Complexes (Elsevier, Amsterdam, 1972), vol. 2. l6 J. J. Phillipson, P. B. Wells and G. R. Wilson, J. Chem. SOC. A, 1969, 1351; P. B. Wells and A. J. Bates, J. Chem. SOC. A, 1968, 3064; B. J. Joice, J. J. Rooney, R. B. Wells and G. R. Wilson, Discuss. Faraday SOC., 1966,41, 223; G. Webb, Catal., 1978, 2, 145. G. C. Bond, Catalysis by Metals (Academic Press, London, 1962), p. 307. (PAPER 3/1205)

ISSN:0300-9599    

DOI:10.1039/F19848001435

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I , 1984, 80, 1447-1456 Kinetics of the Decomposition of Hydrogen Peroxide Catalysed by Copper and Nickel Ferrites BY ANTHONY I. ONUCHUKWU Physical and Industrial Chemistry Research Laboratory, Department of Chemistry, Bayero University, P.M.B. 301 1 , Kano, Nigeria Received 13th July, 1983 The kinetic activities of the heterogeneous decomposition of hydrogen peroxide by metal-iron spinel oxides, M,Fe,-,O, (M = Cu or Ni), have been investigated with a view to defining the effects of composition and microstructure on their catalytic activity. The entire composition range, 0 < x < 3, was prepared and characterized. Although the series of nickel catalysts was found to possess a lower aggregate diameter than the copper series, the decomposition activity of the former was surprisingly much less than the latter.This poor performance of the nickel series is explained in terms not of a consideration of microstructural defects but rather of the restricted redox couple represented by Mn/Mn-l in the electronic composition of the catalysts and, possibly, the absence of a more active ion (Mn) at high compositions on the octahedral lattice sites which may initiate the cyclic electron-transfer process on the catalyst surface, as proposed by Abel and Cota et al. The value of x corresponding to maximum activity was experimentally evaluated for each catalyst series by determining the highest specific rate constant having minimum values of the Arrhenius pre-exponential factor, rather than minimum values of the activation energy.The production of oxygen gas has for many years relied on the electrolysis of water. Compression in steel cylinders is possibly the only means of storage of the hydrolysis products. The choice of a suitable peroxide (H,O,) decomposition catalyst as an alternative for oxygen-gas production has proved difficult because the alternatives tried, silver oxide, platinum and palladium blacks, are expensive and therefore unattractive.lP3 Similarly, cheap corrosion-resistant catalysts such as MnO,, C0,03 and Fe,O, are not suitable because of the poor activity of these catalysts towards peroxide decomposition. 4-6 Cobalt-iron spinel oxides of the general formula Co,Fe3_,O4, where x, the composition variable, can take values between 0 and 3, were reported by Cota et al. to possess activity for peroxide decomposition in alkaline media.5 Goldstein et al.prepared the entire composition range of the cobalt-iron spinel oxides by hydroxide and oxalate routes and concluded that the cobalt catalysts produced by these techniques possessed similar structures and ~ ~ m p ~ ~ i t i o n . ~ - ~ Goldstein and Sat0 et al. lo reported high surface areas for catalysts produced via the hydroxide route. Thus the order of activity of the series of catalysts from the two routes was explained in terms of intrinsic factors, e.g. electronic composition and surface morphology of the catalyst p o ~ d e r s . ~ ~ The high activity of the cobalt-iron spinel oxide system towards peroxide decomposition was explained by a redox-couple mechanism in which the presence of Corl ions at the octahedral lattice sites of the cobalt spinel oxide structure initiated a cyclic electron-transfer proces~.~~ ' 7 Prior to the postulation of a redox mechanism, Erdey et al.stated that in peroxide reactions with oxidants containing oxygen the peroxide bond in the peroxide molecule is never dissociated, as was proved by isotopic 14471448 HETEROGENEOUS DECOMPOSITION OF HYDROGEN PEROXIDE measurements.12 These authors proposed a mechanism with a scheme involving an interaction between the peroxide molecule and the perhydroxide ion, HO;. l2 Recently13 other authors have suggested that peroxide decomposition in an alka- line medium proceeds via two mechanisms : the radical-chain mechanism, initiated by the ions of heavy metals (e.g.Cu, Fe or Mn), and the non-radical route, oia unspecified intermediates. However, in the preparation of hydrogen peroxide by the reduction of oxygen in 5 mol dm-3 KOH (to provide sufficient conductance) it was found that traces of any of the heavy-metal ions decomposed the perhydroxide ion in situ, through the redox-couple Interest in the kinetics of peroxide decomposition in dilute solutions of alkali hydroxides stems from its use in oxygen production, bleaching cellulose and textile materials and the electrolytic reduction of ~ x y g e n . ~ ~ lo* 1 5 9 l6 According to Burki et aL2 and Weiss4 peroxide decomposition in alkali-metal hydroxide solutions is a first-order reaction with respect to the peroxide, the decomposition rate increasing with hydroxide concentration.Cota et al. and Parravano also reported first-order kinetics for a series of cobalt-iron spinel The authorlg? 2o has recently reproduced the first-order reaction for cobalt spinel oxides and established that maximum activity in the series occurs at x = 1.5. The present paper describes an investigation of the peroxide decomposition reaction with copper-iron and nickel-iron spinel oxide catalysts of general formulae Cu,Fe,-,O, and Ni, Fe3-, O,, respectively. This study takes cognizance of the redox-couple mechanism in both catalyst systems represented by CuI1/Cu1 and Ni1I1/NiI1. The structural and electronic composition, as well as the surface morphology, of the entire catalyst series have been considered. A comparison of the performance of these spinel oxides with the previously reported series of cobalt catalysts revealed that the copper series is of comparable activity to the cobalt series.However, copper-iron spinel oxide is cheaper than cobalt-iron oxide. The decomposition activity of nickel-iron oxide series is found to be much lower than copper, despite its lower aggregate diameter. The maximum value of x was evaluated in each case by comparing the specific rate constant, activation energy and pre- exponential factor of the two catalyst systems. Thus the activity order in each series was established. The difference in the efficacy of these catalyst series towards peroxide decomposition was found to be influenced by their electronic composition and diffusion effects arising from surface morphology.13+ l4 EXPERIMENTAL PREPARATION OF M, Fe,-, 0, CATALYSTS Two different metal-iron spinel oxides, namely copper-iron and nickel-iron, were investigated. The entire composition range between x = 0 and 3, differing in microstructure and composition, was synthesized by hydroxide coprecipitation following the method of Sat0 er uE.,Io this technique being preferred for preparing high-surface-area catalyst series. Details are given in ref. (2), (6) and (21H23). Oxides with values of x of 0, 0.5, 1.0, 2.0, 2.5 and 3.0 were prepared. CHARACTERIZATION OF THE SPINEL OXIDE SERIES To ensure that the spinel or any other phase had been formed, the oxide samples were subjected to the Debye-Scherrer X-ray analysis technique. The work was carried out on a PlOlO X-ray generator using appropriate filters, Lindemann glass tubes of 0.5 mm internal diameter and an analytical camera of 1 1.46 cm length.The values obtained for the d-plane spacings and their relative intensities of reflection were compared with those from ASTM literature in order to determine the major phases of spinel present. Scanning electron microscopy (SEM) and EDAX analyses were conducted on samples in theA. I. ONUCHUKWU 1449 Cu and Ni spinel oxide series in order to determine particle-size and shape distributions as well as the relative proportions of cations in the spinel series. An Electron Optics model JEMIOOB microscope with an EDAX attachment was used in this investigation. The microstructures of the catalyst series powders were characterized by visual studies using a high-resolution ( x 100) optical microscope and the Coulter-counter particle-sizing technique.An electronic Coulter counter was used by ultrasonically dispersing the powder particle in an isotonic aqueous electrolyte and evaluating the mean aggregate size of the whole oxide powder series. The Brunauer, Emmett and Teller (B.E.T.),* specific surface areas of the powder series were measured by nitrogen adsorption at 80 K, after degassing of the catalyst at 298 K for 1 h.21,24 DETERMINATION OF KINETIC ACTIVITY OF THE PEROXIDE DECOMPOSITION REACTION The kinetic activity of the spinel oxide catalysts for H,O, decomposition was evaluated by the rate of production of oxygen gas in the liquid phase. A constant catalyst weight (50 mg) was injected into a thermostatted reaction vessel containing 5 cm3 of 0.4 mol dm-3 H,O, in 5 mol dm-3 KOH as a diluent for each oxide specimen.The mixture was stirred vigorously using a magnetic stirrer. The alkali and peroxide were AnalaR and Hopkins & Williams reagents, respectively. Both solutions were standardized immediately before use, KOH using standard HC1 ampoules and H,O, using KMnO, solution. Oxygen gas evolved from the reaction vessel according to the equation ks 2HOi + 0, + -OH (9 where the peroxide was present in alkaline solution as the perhydroxide ion, HO;.ls The rate of oxygen evolution was monitored using the gasometric assembly of Cota et aL5 This technique has been used extensively in previous work and described in detail e l ~ e w h e r e . l ~ * ~ ~ $ ~ ~ The timedependent volume, &, of evolved oxygen was monitored at 30 s intervals in all cases studied.The maximum recorded, V,,,, was also indicative that the reaction was complete. Values of the composition variable, x, and the Arrhenius parameter^^^ were studied within the temperature range 300-3 12 K. In a stability study the catalyst retrieved from the reaction vessel was stored in 5moldm-3 KOH in air, and the time taken to attain V,,, was monitored periodically for 72 days. RESULTS X-ray characterization of the oxide series revealed that using deionized distilled water to prepare the solutions produced impurity-free spinel oxides. The effect of x on the particle-size and cation distribution of each system determined by SEM is displayed in tables 1 and 2 for the two spinal systems.Fig. 1 shows the effect of varying the composition on the specific surface area and mean aggregate diameter (d/pm) determined by the B.E.T. and Coulter-counter techniques, respectively. In fig. 2 the volume ( 5 ) of 0, is plotted against the time ( t / s ) of evolution for 50 mg of each of the two catalysts for x = 1.5. These plots were corrected for self- decomposition of the peroxide, i.e. the 0,-evolution rate at 298 K is equal to 1.3 x lop4 cm3 s-1.26-28 The Cu and Ni spinel systems investigated were found to follow a rate law which was first order with respect to peroxide in 5 mol dmP3. The first-order rate constant defined for reaction (i) is (ii) where V, is the initial volume of 0, gas produced by self-decomposition, [HO;], is the initial concentration of the peroxide (0.4 H,O,)mol dm-3 and k, is the rate constant.In order to enhance reproducibility, some authors5* '9 lo, 2o have expressed the peroxide decomposition rate constants per unit mass of catalyst, i.e. ks/s-l g-l. A Commodor microcomputer model PET2001-16NB5 was used to obtain the rate 48 F A R 11450 HETEROGENEOUS DECOMPOSITION OF HYDROGEN PEROXIDE Table 1. Properties of the copper-iron oxide series B.E.T. specific spinel cation valency distributions surface X composition in the spinel structurea area/m2 g-' 0 Cu,Fe,O, (Fe3+)tet[Fe2+Fe3+]oct 0;- 135 Oe5 Cu0.5Fe2.504 (Fe3+)tet[C~~5Fe$=Fe2+loct 0;- 92 1.0 Cu,FE,O, (Fe;+)tet[Cu+Fe3+]oct 0;- 86 2.0 Cu,Fe,O, ( C U + ) ~ ~ ~ [ C U ~ + F ~ ~ + ] ~ ~ ~ 0;- 64 2*5 CU2.5Fe0.504 (Cu+)tet[Cu~;Fei3oct 0;- 66 3.0 Cu,Fe,O, (CU+)tet[CU2+CU+]oct 0:- 78 (inverse) (inverse) Cu1.5Fe1.504 (Cu~5Fe~~)tet[Cu~5Cu$~Fe3+loct 0;- 74 (normal) (normal) ~ ~~ a tet = Tetrahedral site, oct = octahedral site.Table 2. Properties of nickel-iron oxide series B.E.T. specific spinel cation and valency distributions surface X composition in the spinel structurea area/m2 g-l 0 Ni,Fe,O, (Fe3+)tet[Fe2+Fe3+]oct 0;- 135 (inverse) (normal) (inverse) 0.5 Ni,,,Fe,.,O, (Fe3+),,,[Nif:',Fe$;Fe3+],,, 0;- 120 1.0 Ni,Fe,O, (Fe3+)tet[Ni2+Fe3+]oct 0;- 86 1.5 Nil.5Fel.,0, (Ni~?Jtet[Ni~~Nii~Fe3+],,, 0;- 84 2.0 Ni,Fe,O, (Ni2+)tet[Ni3+Fe3+]oct 0;- 81 2.5 Ni,.,Fe,.,O, (Ni2+)tet[Ni3+Fe3+]oct 0;- 83 3.0 Ni,Fe,O, (Ni2+)tet[Ni3+Ni2+]oct 0;- 68 (normal) a tet = Tetrahedral site, oct = octahedral site.constant (k,) for various values of x at temperatures between 300 and 312 K. The reaction rate constants were independent of initial peroxide concentration and a first-order rate plot was linear for over 95% of the total reaction period. Also, the first-order rate constants were directly proportional to the steady mass in the reaction mixture for the weight range 20-100 mg considered. The effect of x on k, is shown for two temperatures (300 and 308 K) in fig. 3. The maximum error check in the programmed solutions of k, values is _+ 1.5% and reproducibility was good. Fig. 4 depicts the influence of the diffusion effect on the Arrhenius activation energy (E,) for each oxide in the two catalyst systems. The activation entropy of decompositionA. 1. ONUCHUKWU 1451 120 3 I w N E .cd cd h I I 0 1 .o 2.0 3.0 X Fig. 1. B.E.T. surface area (triangles) and mean aggregate diameter, d/pm (circles), for Cu, (closed symbols) and Ni, (open symbols) plotted as functions of x. 0 200 LOO 600 t Is Fig. 2. Volume, 5, of oxygen evolution plotted against t / s for Cu, (A, A) and Ni, (0, 0) systems at 300 K. (---) Unstirred and (---) stirred systems. in 5 mol dm-3 for the two series was evaluated by a programmed solution ot'the Eyring equation : (iii) where AH$ and AS$ are the activation enthalpy and entropy changes, respectively, and the other symbols have their usual meanings. For identical values of x (1.5) and under vigorous stirring the Cu and Ni catalysts gave values of 56.3 and 29.7 kJ mol-l, respectively, for the enthalpy of activation.A plot of the effect of diffusion on the variation of entropy of activation with x is shown in fig. 5. The changes in activation enthalpy and entropy evaluated in this study were subject to an experimental error 48-21452 HETEROGENEOUS DECOMPOSITION OF HYDROGEN PEROXIDE 3f 2E 20 " I M I --. - Atm 12 . 4 0 i 1 .o 2.0 3 .O X Fig. 3. Variations of k,/s-l g-l with x for Cu, (m, 0) and Ni, (a) at 300 K and Cu, (0, 0) and Ni, (A, (>) at 308 K. (-) Unstirred and (---) stirred systems. of & 1.5 kJ mol-l. It is also important to note that a stability test on the catalyst for a period of 72 days showed no deterioration in the peroxide decomposition reaction. DISCUSSION Of the various methods reported for the preparation of spinel oxides, the copre- cipitation method of Cota et al.was adopted in this study because the technique produces higher-surface-area oxide powders. Characterization of these high-surface- area powders by X-ray diffraction and EDAX revealed that the spinel oxides prepared were impurity free. Unlike several studies5* 9 v 28 two different catalyst systems (Cu and Ni spinel oxides) have been produced, with differences in composition and microstructure. Thus the significant objective of this work vis u vis the peroxide decomposition reaction was to arrive at an intrinsic order of catalyst activities that is far more dependent on catalyst composition than microstructural differences. In tables 1 and 2 we show the composition (cation distribution) and surface morphology, and specific surface area (in m2 g-l) of the two oxide systems.Again, the variation of mean aggregate diameter and specific surface area with catalyst composition (fig. 1) indicates that the nickel oxide catalysts have much smaller aggregate diameters, seen by the high specific surface area recorded. One possible approach to ascertainingA. I. ONUCHUKWU 1453 \ \ I I I 0 1 .o 2 .o 3.0 Fig. 4. Activation energy as a function of x for Cu, (A, A) and Ni, (m, 0). (-) Unstirred and (---) stirred systems. X the intrinsic activity is to make use of a kinetic parameter which is effectively independent of the catalyst microstructure for each of the two oxide series. A suitable choice of such a parameter would be the activation energy or entropy for the peroxide decomposition reaction.These kinetic parameters are dependent on catalyst composition rather than surface 28 A plot of 0, volume ( q ) against the time of evolution (fig. 2) revealed that the activity of the copper catalyst superseded that of the nickel catalyst under similar experimental conditions. In contrast to the findings of Spalek et aZ.,lS the decomposition of peroxide in 5 mol dm-a KOH obeys the first-order rate equation (ii) with respect to the total peroxide content. This finding is consistent with first-order behaviour for peroxide decomposition in alkaline media reported by several author~.~-~$ ' 9 2o It is, however, interesting to note that despite the observed differences in performance (fig. 2), the variation of the rate constant, k,, with composition (fig. 3) determined at two different temperatures (300 and 308 K) is more interesting. First, an increase in temperature enhanced the activity of the two oxide systems.This increase is more marked for the Cu series than the Ni series (fig. 3). Secondly, although an increase in composition enhanced the activity of Cu catalysts to a maximum at x = 2.0 the nickel catalyst series deteriorated in performance for an identical increase in catalyst1454 HETEROGENEOUS DECOMPOSITION OF HYDROGEN PEROXIDE '*r 0 1 .o 2.0 3.0 A Fig. 5. Effect of stirring (---) on the variation of activation entropy with x in the Cu, (A, A) and Ni, (0, 0) systems. composition. These differences in performance, particularly for the Ni series despite its microstructural advantage (low aggregate particle size), suggest that compositional effects operate within the catalyst series. Hence the poor activity of the Ni catalyst series suggests that an increase in composition inhibits the decomposition reaction ; this is because it is possible that the redox-couple mechanism postulated in earlier studies is hindered, as the formation of nickel as Nil* is more favourable than in a higher oxidation state on the iron oxide support.lV 2 t 4 9 l1 Although the Cu series is favoured by an increase in x, the fall in activity for x > 2 is surprising. However, at x > 2 the Cu series catalysts have a lower mean aggregate diameter, reflected in the high surface area as shown in fig.1. An explanation for this observed anomally may be found by considering diffusion-control effects in the heterogeneous peroxide decomposition reaction.It was observed during the course of the experiment that although 0, evolution produced pronounced turbulence, the high reaction rate with the Cu series was dependent on vigorous mechanical stirring, (fig. 2 and 3). According to Goldstein et al.,9928 who studied the decomposition of peroxide by cobalt spinel oxide, when reaction turbulence occurs near the surface of a solid dispersed in a liquid, a thin liquid layer adheres to the solid surface. Furthermore, they reported28 that for peroxide decomposition characterized by a three-phase [liquid- solid-gas (oxygen)] heterogeneous reaction, the ingress of fresh reactant (H20,) through such a layer to the solid (catalyst) surface to sustain the reaction becomes restricted. This phenomenon would be more serious for solids with cracks and micro- pores where gas bubbles could be occluded.Hence, partial mitigation by vigorous stirring of the reaction indicates that diffusion effects may possibly control the overall decomposition kinetics, especially for high-surface-area solids observed for Cu(x > 2)A. I. ONUCHUKWU 1455 and the entire Ni series. Although an increase in temperature enhanced 0, evolution (fig. 3), the values of the activation energy (fig. 4 and tables 1 and 2) were not consistent with this diffusion-control phenomenon. Despite the vigorous mechanical stirring adopted during the reaction, effects of diffusion control could not be mitigated completely. It may be that the diffusion effect is significant in this study because an increase in the mass of the catalyst provides more particles in which diffusion layers could be set lo, 2o This explains the dependence of the rate constant on the catalyst mass (fig.3), suggesting that diffusion kinetics would be dependent on composition. In order to verify this reasoning, the activation energies (fig. 4) and the entropy (fig. 5), which depend more on composition, were evaluated. The maximum error is of the order of 1.5 kJ mol-l. There is in each case a pronounced minimum at x = 2 and x =2.5 for the Cu and Ni series, respectively. The increase in E: beyond this valley in each system is consistent with a lower mean aggregate diameter, which promotes diffusion control of the catalyst activity, leading to an increase in activation energy (fig. 4).However, vigorous stirring of the reaction has minimized the defusion effect. The restricted redox-couple mechanism (NiI1/Ni1I1) and the increase in composition with minimum surface area could account for the poor activity of the Ni spinel system. Hence, the low activation energy recorded for the Ni oxide series is suspect, because steric hinderance accompanying the attainment of Ei is not reflected in the activation entropy, as shown in fig. 5. The values of the AS$ in the peroxide-catalysed decomposition could be related to the more easily available redox couple CuI1/Cul or the entropy of adsorption of the peroxide on the catalyst. This postulation is reflected in the lower ASS values recorded for the Cu series than for the Ni system.This finding is consistent with the increase in A S (at x > 2.0 and x > 2.5 for the Cu and Ni catalysts), for which a decrease in surface morphology could be responsible for the increase in steric hinderance. Thus the variation of entropy with composition (fig. 5 ) is considered to be significant in the overall assessment of the catalyst systems. Nevertheless, the activities of the two catalyst series may be compared on the basis that the more effective catalyst for peroxide decomposition possessed a lower entropy without a corresponding low activation energy. In this work an attempt has been made to establish the activity order of the two catalyst systems considered using the composition factor. Thus for the two temperatures studied, 300 and 308 K, the order of activity for the copper spinel oxide series is 3.0 < 0.5 < 1 .O < 1.5 < 2.0 < 2.5, while the nickel-containing series maintain the order 0.5 > 1.0 > 1.5 w 3.0 > 2.0 w 2.5.Also, catalysts stored in 5 mol dmP3 KOH for a test period of 72 days showed no deterioration in stability. In conclusion, this work has established the significance of catalyst composition towards the enhancement of catalyst activity in the peroxide decomposition reaction. However, it is equally important to match the microstructural factors with composition in order to achieve optimum efficiency for metal-iron spinel oxides. I thank the technical staff of the Physics Department of The City University, London, for the use of X-ray and electron micrographic facilities.I also thank Mr P. B. Mshelia for preliminary work, and the Research and Higher Degrees Committee, Bayero University, Kano for financial support. G. Tammann, Z . Phys. Chem., 1889,4,441. F. Burki and F. Schaaf, Helu. Chim. Acta, 1921, 4, 418. W. Vielstich, Z. Phys. Chem., 1958, 15, 409. J. Weiss, Trans. Faraday SOC., 1935, 31, 1547.1456 HETEROGENEOUS DECOMPOSITION OF HYDROGEN PEROXIDE H. M. Cota, J. Katan, M. Chim and F. J. Schoenweis, Nature (London), 1964, 203, 1281. W. C. Schnumb, C. N. Satter-Field and R. N. Wentworth, H202 (Reinhold, New York, 1955). J. R. Goldstein and A. C. C. Tseung, J. Muter. Sci., 1972, 7, 1383. A. C. C. Tseung and J. R. Goldstein, J. Phys. Chem., 1972, 76, 3646. J. R. Goldstein, Ph.D. Thesis (The City University, London, 1970). lo T. Sato, M. Sugihara and W. Saito, Rev. Electron Commun. Lab., 11, 1963, 26. l1 G. Blasse, Phillips Res. Rep., 1963, 18, 383. l2 L. Erdey, Acta Chim. Acad. Sci. Hung., 1953, 3, 95. l3 Kh. Raskina, F. I. Sadov and G. A. Bogdanov, Zh. Prisk. Khim., 1970,43,447; 1449; 1977,50,724. l4 K. Balogh, J. Balej and 0. Spalek, Chem. Zvesti, 1976, 30, 385; 61 1. l5 0. Spalek, J. Balej and K. Balogh, Collect. Czech. Chem. Commun., 1970, 42; 312; 394. l6 0. Spalek, J. Balej and I. Paseka, J. Chem. SOC., Faraday Trans. I , 1982, 78, 2349. l7 E. Abel, Monatsh. Chem., 1952,83,422. G. Parravano, Proc. Int. Congr. Catal. (1970), vol. 1, p. 157. A. I. Onuchukwu and A. C. C. Tseung, U S . Patent Application 4,335,754, June, 1982; U.K. Patent Application 2,064,587, November, 1980. 2o A. I. Onuchukwu, J. Catal., submitted for publication. 21 A. I. Onuchukwu, Electrochim. Acta, 1982, 27, 529. 22 A. I. Onuchukwu, J. Electrochem. SOC., 1983, 130, 1077. 23 A. I. Onuchukwu and P. B. Mshelia, J. Chem. Educ., in press. 24 P. H. Emmett, Measurement of the Surface Area of Solid Powders (Reinhold, New York, 1954), 25 M. J. Pilling, Reaction Kinetics (Oxford University Press, Oxford, 1975), vol. 1, p. 10. 26 K. B. Keating and A. G. Romer, J. Phys. Chem., 1965,69, 3658. 27 M. Ardon, Oxygen, Elementary Forms and Hydrogen Peroxide (Benjamin, New York, 1965), p. 84. 28 J. R. Goldstein and A. C. C. Tseung, J. Catal., 1974, 32, 459. V O ~ . 2, pp. 31-74. (PAPER 3/1210)

ISSN:0300-9599    

DOI:10.1039/F19848001447

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984, 80, 1457-1465 Catalytic Activity of Dealuminated Y and HZSM-5 Zeolites Measured by the Temperature-programmed Desorption of Small Amounts of Preadsorbed Methanol and by the Low-pressure Flow Reaction of Methanol BY JANA NOVAKOVA,* LUDMILA KUBELKOVA, KAREL HABERSBERGER AND ZDENEK DOLEJSEK J. Heyrovsky Institute of Physical Chemistry and Electrochemistry, Czechoslovak Academy of Sciences, Machova 7, 121 38 Prague 2, Czechoslovakia Received 13th July, 1983 The activity of HZSM-5, NH,Y and Y zeolites, dealuminated using SiCl,, in methanol transformation has been compared and correlated with the acid properties and the Si: A1 ratio of the respective zeolites. The temperature-programmed desorption of small amounts of methanol preadsorbed at ambient temperature together with measurements in a flow reactor under low pressure allowed the primary reaction steps to be observed.HZSM-5 and Y zeolites which had been moderately dealuminated showed the highest activity; however, the product distribution was different above 630 K : HZSM-5 gave predominantly methane and formaldehyde (aromatics appearing at higher temperatures), while C3-6 olefins and aromatics were formed above the Y zeolite. HY and highly dealuminated Y zeolite with Si/Al 3 20 were less active than moderately dealuminated Y and HZSM-5 zeolites. The high activity of the latter zeolites is caused by the action of strong proton-donor sites and electron-acceptor centres (extralattice aluminium in dealuminated zeolites). The difference between HZSM-5 and dealuminated Y zeolites is caused by their different structures, which lead to different reaction mechanisms.For comparison, the temperature-programmed desorption of preadsorbed ethanol was also studied. This resulted predominantly in ethene formation via intramolecular dehydration of ethanol. The unusual catalytic properties of HZSM-5 zeolites have been widely discussed over the last decade: their high Si: A1 ratio, the strong acid sites and a structure which includes a special channel pore system have been assumed to be the main causes of their selectivity and catalytic stabi1ity.l Comparative studies of ZSM, mordenites, faujasites and other types of zeolites have been published, the differences in their behaviour being explained by some of the features mentioned In the present paper, we have tried to compare the methanol interaction with HZSM-5 and with dealuminated Y zeolites.Y zeolites with Si : A1 ranging from 2.5 to 43 were prepared by Beyer's method using SiC1,. It was shown in a previous paper' that this thermochemical treatment not only changes the Si:Al ratio in the lattice but also results in the appearance of extralattice A1 species, both of which influence the acidity of the zeolite. Thus it was possible to study the catalytic behaviour of HY zeolites covering a certain range of Si:Al ratio and acidity together with the behaviour of HZSM-5 whose Si: A1 ratio and acidity lay within the same range. Comparison was also made between hydrothermally treated Y zeolites containing oxidic extralattice A1 species in cavities and AlHY containing cationic Al.In some cases, the interaction of ethanol was also investigated in order to study the behaviour of ethene, which is its main p r o d ~ c t . ~ Ethene is often assumed to be the first intermediate with a C-C bond in methanol conversion [see ref. (10) and references therein]. 14571458 CATALYTIC ACTIVITY OF DEALUMINATED ZEOLITES Temperature-programmed desorption (t.p.d.) of small amounts of methanol or ethanol, preadsorbed at room temperature (the number of alcohol molecules being one order of magnitude lower than the number of acid sites) was investigated mass spectrometrically. The catalytic conversion of methanol on our zeolites was performed in a flow reactor at low pressures (1 Pa) with mass-spectrometric analysis of the products. EXPERIMENTAL The parent NaY zeolite (Vurup, Bratislava, Czechoslovakia) was dealuminated by Beyer's method using SiC1, [for details see ref.(7)]; after the reaction and careful washing it was exchanged with NH,NO, in the same way as the parent zeolite, further referred to as HY. Dealuminated zeolites were denoted Y-A, Y-B, Y-C and Y-E in order of decreasing Si: A1 ratio. Sample Y-A, with the highest level of dealumination, was kindly provided by Dr H. K. Beyer. The zeolite Al,,H,,Y (denoted AlHY) was prepared from NH,Y by ion exchange with 0.1 mol dm-, Al(NO,), at pH 4. The hydrothermally stabilized Y (denoted SY) was obtained from NH,Y by steaming at 1030 K for 3 h [details of both preparations are given in ref. (1 1)) All the zeolites exhibited good crystallinity, as is evident from the infrared spectra, X-ray diffraction patterns and sorption capacities.ZSM-5 zeolite was synthesized according to ref. (12) and converted into the acid form using 0.5 mol dm-, HNO,. Prior to the measurements, the zeolites were dehydrated and deammonized at 670 K for 18 h in uucuo (lo-, Pa). Dealuminated Y zeolites were heated before this treatment in an oxygen flow at 800 K for 5 h and then rehydrated. The characteristics of the zeolites used are listed in table I. Si: Altotal was determined using classical chemical analysis and Si : Allattice from mid-infrared spectra.'. l3 The Ar sorption capacities were measured in the usual way." The number of strong Bronsted and Lewis acid sites was calculated from the heights of infrared bands of pyridinium ions (1545 cm-l) and pyridine complexes with aluminium electron-acceptor centres (1455 cm-l), both these complexes being retained in the zeolites after pyridine desorption at 530 K.' The acid strength was com- pared using t.p.d.of ammonia (see fig. 1):3 x 10ls molecules of NH, were adsorbed at room temperature on 0.01 g of the zeolite for 30 min; then NH, was pumped off through the vacuum system of a mass spectrometer (modified MI 1305, U.S.S.R.). The temperature dependence of the NH, peaks was recorded, the heating rate being 6 K min-l. From fig. 1 it follows that the maximum amount of ammonia evolved with HY occurs at a much lower temperature than the second maximum observed with HZSM-5 and dealuminated Y. The interaction of zeolites with methanol (or ethanol) was studied by t.p.d.The alcohols were of A.R. grade and purified by repeating freezing and pumping. Alcoh,ol vapour (10 pmol) was adsorbed at room temperature on 0.1 g of zeolite; after 1 h of adsorption the desorbate was analysed in the same way as for ammonia. The maximum pressure of released gases did not exceed 10-1 Pa. The catalytic reaction of methanol under low pressure (1 Pa) was investigated with HZSM-5, HY, Y-C and Y-B at 293, 510, 630 and 670 K. The flow rate was 3 x lo-, dm3 h-l. In these measurements, 0.05 g of zeolite was placed directly inside the mass spectrometer in the reactor previously used for the study of radical reactions. l4 RESULTS T.P.D. MEASUREMENTS Typical t.p.d. curves for the products of the methanol-zeolite interaction are shown in fig.2. The products appear in the following order: methanol at 300-600 K, dimethyl ether (and a small amount of methyl chloride at slightly higher temperatures, not shown in fig. 2) at 400-600 K, aromatics (up to Clo) and C3-5 olefins from Y zeolites and aromatics (especially xylenes) from HZSM-5 at 600-700 K. The release of aromatics from HZSM-5 is preceeded by the evolution of methane and formaldehydeJ. NOVLKOVL, L. KUBELKOVA, K. HABERSBERGER AND Z. DOLEJSEK 1459 Table 1. Characteristics of zeolites useda C.4J B/ 1 020 L/ 1020 zeolite Si : Altotal Si : Allattice mmol g-' 8-l g-' HY Y -E Y -c Y -B Y -A HZSM-5 AlHY SY 2.5 2.5 3.4 5.2 5.7 10.8 - 19.9 43.5 17.5 - - 2.0 2.9 2.9 9.4 10.9 9.2 9.2 9.4 9.4 5.0 9.0 8.6 7.9 0.15 4.2 2.2 3.8 2.0 0.9 0.75 0.2 0.25 2.4 0.6 3.3 2.6 0.6 1.4 a C,, = Ar sorption capacity; B = number of Bronsted sites; L = number of Lewis sites.r-- I 0 - Y-A 0, ot- 1 TIK Fig. 1. T.p.d. of ammonia, preadsorbed on HY, dealuminated Y and HZSM-5.1460 CATALYTIC ACTIVITY OF DEALUMINATED ZEOLITES HZSM-5 I' Y .- Ej 1 AIHY TIK Fig. 2. T.p.d. of methanol and its interaction products with HY, Y-C, Y-A, AlHY, SY and , CH,OH; B, (CH,),O; IDI, CH,+HCHO; 113, C=C; .,aromatics. in the approximate ratio 1 : 1. The latter two compounds are observed at the same temperature as the aromatics over highly dealuminated Y-A and HZSM-11 (the latter not being dealt with in this paper). Note that the most active zeolites are HZSM-5 and moderately dealuminated Y zeolite Y-C as well as AlHY and SY.HY and highly dealuminated Y-A are less active. In fig. 3, the maxima of released products are plotted against the molar fraction of aluminium. Fig. 3 includes all the dealuminated Y zeolites, HY and HZSM-5. Desorbates of both methanol and ethanol are shown. The moderately dealuminated zeolites Y-E and Y-C are the most active as far as the formation of aromatics and C3-5 olefins from methanol and ethene from ethanol is concerned. With increasing dealumination and for HY the amount of these products is decreased; HY gives the largest amount of diethyl ether from ethanol; Y-A has a yield of ether similar to that of Y-B and Y-C while the amount of other products formed from ethanol is very low. The formation of dimethyl ether from methanol does not depend strongly on the A1 content.HZSM-5 differs from Y-B in spite of possessing almost the same molar fraction of aluminium and has a total activity [fig. 3(a)] comparable to that of moderately dealuminated zeolites. Fig. 3 also shows the formation of methyl and ethyl chloride from methanol and ethanol, respectively, from traces of chlorine left in the dealuminated zeolites. This amount of chlorine is so small that it cannot be detected by any other method; only in Y-E was a very small amount of C1 observed, by X.P.S.,J. NOVAKOVA, L. KUBELKOVA, K. HABERSBERGER AND Z. DoLEJSEK 1461 Y-A Y-B Y-E HY 01 -*--- 0.1 0.2 AI/Si + A1 Fig. 3. Dependence of amount of methanol, ethanol and their interaction products on the aluminium mole fraction (Al:Si+Al): x , HY and dealuminated Y; e, HZSM-5.(a) Methanol and ethanol, (b) methyl chloride from methanol, (c) ethyl chloride x 10 from methanol, (d) dimethyl ether from methanol, (e) diethyl ether x 10 from ethanol, U, CSp5 olefins from methanol, (g) aromatics from methanol, (h) ethylene from ethanol and (i) aromatics from ethanol. in the surface layers of the zeolite. HY and HZSM-5 also give negligible methyl and ethyl chloride formation, probably because of the chlorine impurities introduced into these zeolites during their preparation. From methanol, the maximum amount of methyl chloride is found with Y-E, which has, relatively, highest chlorine content. The release of ethyl chloride shows the same dependence as the formation of ethyl ether. LOW-PRESSURE FLOW MEASUREMENTS The time dependence of the amount of the individual components in the reactor output for the HY, Y-C, Y-B and HZSM-5 zeolites at various reaction temperatures is shown in fig.4. No selective reaction products were found under these conditions with HY: it exhibits minimum absorption of methanol at room temperature, at higher1462 CATALYTIC ACTIVITY OF DEALUMINATED ZEOLITES HY Y -c Y-B HZSM-5 tlmin Fig. 4. Time and temperature dependence of methanol conversion in a low-pressure flow reactor over HY, Y-C, Y-B and HZSM-5. (-) Ethanol, (----) dimethyl ether, ( x - x - x -) aromatics, (. . . .) olefins, (----) methane and formaldehyde and (Q-@-@-) carbon dioxide; (a) 295, (b) 510, (c) 630 and ( d ) 670 K. temperatures only dimethyl ether is formed and at 670 K an appreciable amount of carbon dioxide appears.Moderately dealuminated Y-C absorbs all the methanol input within the studied time interval at room temperature and the composition of the reaction products at higher temperatures is the same as in the t.p.d. experiments. Y-B, dealuminated to a greater extent, behaves as for the t.p.d. experiments, exhibiting lower activity than Y-C. HZSM-5 also gives the same products at the relevant temperatures as for the t.p.d. experiments, except that the formation of formaldehyde was greater than that of methane. After increasing the temperature to 770 K (not shown in fig. 4) aromatics appear as the reaction products. DISCUSSION Amongst the numerous mechanisms for methanol conversion on zeolites which have been suggested in the literature is the assumption of the principal action of strong Bronsted acid sites [e.g.ref. (2), ( 5 ) and (1 5)] supported by the presence of Lewis acid sites.ls, l7 Though the amount of information on methanol conversion keeps growing, the mechanism of this reaction, especially formation of the first C-C bond, has still not been elucidated. It is also not clear whether the basic steps are the same for all zeolites or whether the properties of the relevant zeolite affect the interaction from its very beginning. It is difficult to establish the reaction mechanism because of a number of sequential and parallel reactions which are more or less pronounced under various experimental conditions and which give rise to a broad range of products. The experimental conditions used in our investigation allowed us to study the reactions at very low concentrations of both surface species and gas molecules.The path of the latter between the catalyst and the analyser is so short that interaction between the molecules themselves is almost eliminated. Using this technique we can approach the primary steps of the reaction. The first interesting result is the coincidence of the product composition for the t.p.d.J. NOVAKOVA, L. KUBELKOVA, K. HABERSBERGER AND z. DOLEJSEK 1463 experiments and the low-pressure flow experiments at the relevant temperatures. This might be explained by the participation of the same surface species, the reactivity and composition of which are strongly temperature dependent. Similarly, this temperature dependence also holds for the function of the active sites.Thus, it follows that dimethyl ether need not be an intermediate in the reaction of methanol to hydrocarbons: (i) it does not appear in the gas phase above 600K and (ii) its formation at lower temperatures does not follow the activity of zeolites under study as far as the conversion to higher hydrocarbons is concerned. The formation of dimethyl ether is a well known property of acid catalysts1* and it seems that neither strong acidity of the Bronsted type nor a large number of these centres is needed for this reaction. HY, releasing almost no selective products during the t.p.d. of methanol, exhibits at 670 K in the flow experiments increased consumption of methanol. However, only the products of deep oxidation appear in the gas phase.Carbon dioxide was found at higher temperatures in the reaction with methanol for all the zeolites, but the amount was much lower compared with HY. Recently, the use of silica-alumina as a catalyst in methanol oxidation has been described.lg In our case only methanol (without added oxygen) interacts with the zeolites and this carbon dioxide formation might take place via the incorporation of oxygen from the zeolite or by the decomposition of methanol molecules. Experiments to ascertain the stoichiometry of this reaction (hydrogen and water release) are in progress. Both t.p.d. and flow experiments showed the maximum consumption of methanol with moderately dealuminated Y zeolites. In the flow experiments, the appearance of not only methanol at room temperature but also other products at higher temperatures is connected with an induction period which is especially pronounced for Y-C.In this case olefins and aromatics are apparently secondary reaction products formed via oligomerization, cyclization, hydrogen transfer etc. ; no assumptions about the primary steps of the methanol-zeolite interaction can be made on the basis of these experiments. A not very pronounced induction period is found with HZSM-5, which is the only zeolite studied here which releases simple C, products (methane and formaldehyde) at a relatively low temperature. Methane is assumed to be a primary product from methanol and not a compound accompanying coking as supposed previ~usly.~ Methane was mentioned as a product in the very low conversion of methanol over HZSM-5 in ref.(20). Note that the absence of methane from the products was considered to be evidence against the participation of methyl cations in the first C-C bond formation [ref. (20) and (21)]. In our opinion, methane formation might also support the assumption of Zatorski and Krzyzanowski22 on the radical mechanism : some papers have recently appeared concerning the formation of radical species on the surface of HZSM-5 during the interaction with hydrocarbon^.^^ Formaldehyde was assumed to be an intermediate in methanol conversion in ref. (5). The equimolecular desorption of methane and formaldehyde during the experiments indicates random attack of acid sites on the different bonds of the methanol molecule (or its fragments) or the simultaneous action of different zeolite sites (Bronsted or Lewis).Note that the methane : formaldehyde ratio changes towards higher formaldehyde yield in the low-pressure flow experiments. A more detailed study of the surface species under such low coverages will be needed to explain these findings. In the t.p.d. of methanol, aromatics are released from HZSM-5 at a slightly higher temperature than with dealuminated Y, so it is not surprising that in the flow experiments at 670 K almost no aromatics appeared. It seems therefore that aromatics are more strongly bound in HZSM-5 and that diffusion limitations may play an important role. Although HZSM-5 has a similar Si:Al ratio as dealuminated Y-B zeolite, its activity is1464 CATALYTIC ACTIVITY OF DEALUMINATED ZEOLITES considerably higher; it reaches the activity of moderately dealuminated Y zeolites.Nevertheless, from the different composition of the products from methanol conversion the pathway of the reaction is different. The properties of HZSM-5, including the acidity, Si:Al ratio and structure, are responsible for its high activity as well as for the reaction mechanism, which is probably different from that on Y zeolites. The high activity and selectivity of moderately dealuminated zeolites is caused by the presence of an appreciable number of strong proton-donor sites together with electron-accepting sites originating from extralattice aluminium species, whose form- ation in Y zeolites treated by SIC& was proved in ref. (7) and explained by the hydrolysis of aluminium chloride complexes remaining in the zeolite cavities after substitution of lattice A1 with silicon from SiCl,: This aluminium helps to compensate the lattice charge as cationic or oxidic species and it exhibits strong electron-accepting properties.The influence of the Lewis acid sites on the catalytic activity was confirmed by the behaviour of AlHY and SY zeolites. In the former case, aluminium was introduced into the zeolite cavities by cation exchange while in the latter case the extralattice aluminium species were formed by delocalization of lattice aluminium during the hydrothermal process. The reaction of methanol with traces of chlorine left in the zeolite proceeds regardless of all the properties of the studied zeolites and is most pronounced with Y-E, which contains a relative maximum of chlorine.The low temperature of the methyl chloride release shows that abstraction of the methyl group from methanol can proceed, in these cases, quite readily. The presence of traces of chlorine has no influence on the type and relative amounts of the products. In the flow experiments chlorine in the form of methyl chloride was removed in a very short time while the course of formation of the remaining products with temperature and time was unchanged. The t.p.d. of ethanol with all the dealuminated Y zeolites, as well as with HY and HZSM-5, was investigated for comparison. The amount of unreacted ethanol was almost the same as that of methanol, the main product of the reaction being ethene, with the highest yield over moderately dealuminated Y and HZSM-5.With less active zeolites the formation of diethyl ether was also observed. The formation of ethyl chloride showed the same dependence on the molar fraction of A1 in the zeolites as the release of diethyl ether. The yields of aromatics were appreciably lower than from methanol, the highest amount being formed on HZSM-5. From the reaction studies at lower pumping speed (i.e. with a longer contact time of the ethene with the zeolite) and from the interaction of ethene alone it seems that the formation of aromatics from ethanol proceeds uiu ethylene oligomarization and cyclization. As almost no ethene was found under our conditions in the transformation of methanol, we assume in accordance with ref. (1 5) that ethylene need not be a primary product in methanol conversion.CONCLUSIONS The t.p.d. of small amounts of alcohols and the catalytic flow experiments at low pressure are evidently controlled by the most active sites, so a deeper insight into the first steps of the interaction of zeolites with these reactants may be obtained. The correlation between the acid properties, the Si : A1 ratio and the activity for methanol transformation over zeolites of different structural types reveals the importance of these features for the reaction in question. The most active Y zeolites are thoseJ. N O V ~ K O V ~ , L. KUBELKOV~, K. HABERSBERGER AND z. DOLEJSEK 1465 moderately dealuminated (Si : Allattice = 5 ) containing an appreciable amount of extra- lattice aluminium with electron-acceptor properties and a high number of strong proton-donor sites.HZSM-5 (Si:Al = 17.5) with a similar number of strong proton- donor sites and a lesser amount of Lewis sites exhibits comparable activity, although with a different product distribution. This difference is most pronounced at lower interaction temperatures, where simple C , compounds (methane and formaldehyde) appear over HZSM-5 as the primary products. Dealuminated Y zeolites yield C3-5 olefins and aromatics directly, apparently as secondary reaction products. Their release is accompanied by a long induction period. From the temperature dependence of the formation of dimethyl ether during the conversion of methanol it may be assumed that dimethyl ether is not an intermediate in this process.Similarly, the same assumption concerning ethene may be made when the interactions of methanol and ethanol with zeolites are compared. P. B. Weisz, Proc. 7th Znt. Congr. Catal., ed. T. Seiyama and K. Tanabe (Kodansha, Tokyo and Elsevier, Amsterdam, 198 l), p. A3. 2 J. C. Vkdrine, A. Auroux, V. Bolis, P. Dejaifve, C. Naccache, P. Wierzchowski, E. G. Derouane, J. B. Nagy, J. P. Gilson, J. H. C. van Hooff, J. P. van den Berg and J. P. Wolthuizen, J . Catal., 1979, 59, 248. E. G. Derouane, Catalysis by Zeolites, ed. B. Imelik (Elsevier, Amsterdam, 1980), p. 5. H. Bremer, W. Reschetilowski, Do Guy Son and P. Wendlandt, 2. Chem., 1981, 21, 77. B. E. Langer, Appl. Catal., 1982, 2, 289. P. A. Jacobs, J. A. Martens, J. Weitkamp and H. K. Beyer, Faraday Discuss. Chem. SOC., 1982, 72, 353. ’ L. Kubelkova, V. Seidl, J. Novakova, S. Bednai-ova and P. Jiru, J . Chem. SOC., Faraday Trans. 1 , 1984,80, 1367. H. K. Beyer and J. Belenykaya, Catalysis by Zeolites, ed. B. Imelik (Elsevier, Amsterdam, 1980), p. 203. E. G. Derouane, J. B. Nagy, P. Dejaifve, J. H. C. van Hooff, B. P. Spekman, J. C. Vedrine and C. Naccache, J . Catal., 1978, 53, 40. lo J. P. van den Berg, J. P. Wolthuizen, A. D. H. Clague, G. R. Hays, R. Huis and J. H. C. van Hooff, J. Catal., 1983, 80, 130. l1 B. Wichterlova, J. Novakova, L. Kubelkova and P. Ji1-6, Proc. 5th Znt. Con$ Zeolites, ed. L. V. C. Rees (Heyden, London, 1980), p. 373. l2 R. J. Argauer and G. R. Landolt, U S . Patent 3702886, 1972; K. Habersberger and V. Seidl, Ropa Uhlie, 1981, 23, 207. l 3 E. M. Flanigen, ACS Monogr., 1976, 171, 80. l4 Z. DolejSek and J. Novakova, J. Catal., 1975, 37, 540. l5 R. M. Dessau and R. B. LaPierre, J. Catal., 1982, 78, 136. l7 J. P. van den Berg, Ph.D. Thesis (Technical University of Eindhoven, 1981). l8 P. A. Jacobs, Carboniogenic Activity of Zeolites (Elsevier, Amsterdam, 1977). lB L. Cairati and F. Trifiro, J. Catal., 1983, 80, 25. 21 C. D. Chang and A. J. Silvestri, J. Catal., 1977, 47, 249. 2 2 W. Zatorski and S. Krzyzanowski, Acta Phys. Chem., 1978, 24, 347. 23 S. Shih, J . Catal., 1983, 79, 390. 24 C. D. Chang, Catal. Rev., 1983, 25, 1 . J. Novakova, L. Kubelkova, Z. DolejSek and P. Jirb, Collect. Czech. Chem. Commun., 1979,44,3341. W. 0. Haag, R. M. Lago and P. G. Rodenwald, J . Mol. Catal., 1982, 37, 161. (PAPER 3/1213)

ISSN:0300-9599    

DOI:10.1039/F19848001457

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1984, 80, 1467-1477 Physical Adsorption of Gases on Heterogeneous Surfaces Series Expansion of Isotherms using Central Moments of the Adsorption Energy Distribution BY JAMES A. O’BRIEN AND ALAN L. MYERS* Chemical Engineering Department, University of Pennsylvania, Philadelphia, Pennsylvania 19104, U.S.A. Received 14th July, 1983 An alternative approach to the use of the usual integral equation for heterogeneous adsorption is developed. It provides a means of obtaining an explicit result for arbitrary local isotherms and distributions of energy of adsorption, The approximation is compared with exact numerical integrations, using the Langmuir local isotherm and a normal energy distribution. The series fails to converge for energy distributions that are very wide.Truncation of the expansion after the first-order term yields an equation which is capable of fitting experimental isotherm data. Although the existence of surface heterogeneity has long been recegnized, much of the prcgress in the field has been made in the last 15 years. There are some comprehensive review articles available, of which the most recent is ref. (1). Other reviews of the problem include those by Jaroniec2 and Zolandz and Myers;3 the latter deals primarily with single-component gas adsorption on heterogeneous solids. The basis for most approaches has been to account somehow for a heterogeneous surface as a collection of homogeneous surfaces and to sum the resultant effects. It is obvious that this becomes quite complex as heterogeneity becomes more pronounced.However, the very complexity of the problem leads to simplification in that hetero- geneity may be taken into account by the methods of probability theory. It is generally postulated that a heterogeneous surface is composed of many energetically homogeneous, non-interacting patches, usually referred to as homotattic patches. Furthermore, each patch is considered to be large enough to assume the existence of a so-called local isotherm for that patch and, at the same time, small enough for the surface energies to be characterized by a continuous probability density function. These adsorption energies depend in general on both adsorbate and adsorbent. If the probability of a patch having an energy between E and E+ dE i s f ( E ) dE, then the key equation for heterogeneous adsorption is This is simply the expected value of O,(T,p, E), where E is distributed in the manner specified byf(E) over a domain of energies, D.Eqn (1) is the starting point for most treatments of heterogeneity, but it has been employed in different ways by various authors. Several, including S ~ P S , ~ ~ Misraa3 and Toth,8 assumed that O,( T , p , E ) was given by the Langrnuir isotherm, and they derived the function, f(E), explicitly for certain isotherms O(T,p) by use of the Stieltjes 14671468 ADSORPTION OF GASES ON HETEROGENEOUS SURFACES transform. A similar method was used by Jaroniec and Sokolo~ski,~ who derivedf(E) by means of the Laplace transform, assuming that OL(T,p,E) was given by the Jovanoviclo isotherm and O( T, p ) was the Dubinin-Radushkevich equation.A different type of approach was taken by Myers and 0u,l1 who assumed that OL(T,p, E ) was a step function (the ‘condensation approximation’) and that the energies followed a beta distribution. Eqn ( I ) was then integrated analytically to obtain a new isotherm which is capable of fitting experimental data for heterogeneous adsorption. Still another method of using eqn (1) has been to pick the form of O,( T , p , E ) (for example Langmuir) and, using experimental data for O( T,p), to calculate f(E) numerically. Probably the best-known algorithms for such calculations are CAEDMON by Ross and Morrison12 and HILDA by House and Jay~0ck.l~ Considerable care is necessary in attempting to perform these calculations, as the problem is ill-posed and there are complications associated with numerical stability if the data contain random pertur- bations (which will be the case with all experimental data).In addition, it has been shown by House and c o ~ o r k e r s ~ ~ ~ l5 that the function,f(E), obtained depends strongly on the form of the local isotherm that is chosen. In this paper we present an approach different from those outlined above. We reformulate the right-hand side of eqn (1) as an infinite series whose terms involve (a) derivatives of OL(T,p, E ) with respect to E, evaluated at the mean of the energy distribution, and (b) the central moments of the distribution. This has an important advantage over the previously mentioned methods since the integral in eqn (1) can only be analytically evaluated for a very limited class of combinations of OL(T,p,E) andf(E), beyond which one must resort to numerical integration.With the approach presented here nb integration is necessary. In addition, if the infinite series converges rapidly one can evaluate O( T, p ) with knowledge of only the first few central moments of the energy distribution. This can also provide information on how the shape of the distribution affects the generated isotherm O(T,p), since the central moments of a distribution are related to its shape. DEVELOPMENT OF EXPANSION PROCEDURE Consider a general case of eqn (1) where we take the expectation of a function of a random variable, g(x), with respect to its probability density function, f ( x ) , valid on the domain D of the random variable, x: Suppose that the mean offlx) is p and expand g(x) in a Taylor series about this point.This requires that g(x) have an infinite number of derivatives with respect to x at this point and also that it be continuous there. Thus where, for compactness, dng g y x ) = -(x) dxn g q x ) = g(x).J. A. O’BRIEN AND A. L. MYERS 1469 Inserting eqn (3) into eqn (2) and interchanging the order of integration (note that the subtleties and implications of this operation for convergence have been overlooked for the present), we obtain We now recognize that the quantity in square brackets in eqn (4) is the nth central moment off@), which is normally denoted by pn = E [ ( X - ~ ) ~ ] = (x-p)nf(x)dx. ( 5 ) D Note that ,u is not one of the p n : it is, however, the symbol most commonly employed to denote the mean of the distribution.Inserting the definition eqn (5) into eqn (4) All that now remains is to transcribe eqn (6) to the usual notation of heterogeneous adsorption. From here on it is convenient to work not with the energy, E, but with the dimensionless quantity E-Emin Z T RT * (7) All other energy variables referred to later are also made dimensionless by dividing through by RT. We also note that the local adsorption isotherm e,(T,p, z ) must possess the properties demanded of g(x) in the above treatment. The result is: Eqn (8) is the basis for the remainder of this work. Note that we have replaced the problem of integrating eqn (1) by that of differentiation and summation. Nothing has been said, up to this point, about the convergence properties of eqn (8).However, if we regard eqn (8) as a perturbation expansion on OL(T,p,p) it seems intuitive that convergence may be expected when the moments p n are sufficiently small. In general, by definition, Po = 1, Pl = o Now, p2 is more commonly denoted by 02, the variance of the distribution. In addition one can define the parameterslG a,=? P a, G1” o3 ’ 0, (9) where a3 is the ‘coefficient of skewness’ and is a measure of the asymmetry off(z), and a4 is the ‘kurtosis’ of f(z) and is proportional to the ‘peakedness’ of the distribution. Thus both the normal and the uniform distribution have a3 = 0. For the uniform distribution a, = 1.8 and for the normal distribution a, = 3. Eqn (8’) may be rewritten as1470 ADSORPTION OF GASES ON HETEROGENEOUS SURFACES where the second, third and fourth terms may be identified with the effects of spread, asymmetry and peakedness of the distribution, respectively.We note from eqn (10) that if the series is converging rapidly the effect of the width of the distribution is more pronounced than that of its detailed shape. IMPLICATIONS FOR HENRY’S CONSTANT The above derivation has an interesting implication for the effect of heterogeneity on the Henry’s constant (or surface second virial coefficient), defined as de e p+odP p+oP H E lim - = lim -. Consider eqn (6), where we can insert HF)(p) for g ( n ) ( p ) and H for Y. HL(z) is the Henry’s constant for the local isotherm O,( T,p, z) on a patch of energy z. Performing this substitution and noticing that since for all local isotherms, H(Ln)(p) = H&) (for all n) HL(z) cc ez we arrive at This is a general result, applying to adsorption on a surface of arbitrary adsorption- energy distribution.By picking various distribution functions, JTz), with their associ- ated central moments, we can now obtain an interesting insight into the dependence of H on heterogeneity. For a uniform distribution of energies with mean p and variance 02, it can be shown that Similarly, for a Gaussian distribution of energies, given by it can be shown that H = H,(p)exp - (3 where o2 and p have the same meanings as in eqn (14). In the derivation of eqn (16) it is assumed that the probability of negative energies can be neglected, a physically realistic situation. This last result is essentially the same as that derived by Pierotti and Th0mas.l’ Both eqn (14) and eqn (16) illustrate the point that the variance or spread of the energy distribution strongly influences the Henry’s constant of the heterogeneous isotherm, with a functionality that is exponential in character. It can also be seen that modest values of o will produce isotherms which have high values of slope at zero coverage, an observation which is typical in experimental data for adsorption on heterogeneous surfaces.In addition, these equations demonstrate that H does not depend merely on the energy of the highest-energy patch on the surface.J. A. O'BRIEN AND A. L. MYERS 1471 TESTING OF THE SERIES EXPANSION: A MODEL STUDY The energy distribution employed here is the Gaussian form given in eqn (1 5), under the condition that negative energies can be neglected.For such a density function it can be shown15 that ( n - l)(n-3) ...( 5)(3) (l)on,neven pn=( 0, n odd. The disappearance of all central moments of odd order is a direct result of the symmetry of the density function about the mean. All that now remains is to choose a local isotherm. A preliminary investigation of the Hill-de Boer equation for a two-dimensional van der Waals fluid has been undertaken because it accounts for adsorbate-adsorbate interactions. This equation may be written p = 1 -eL ~,e-Z(*)exp[(*)--~,~~] I -eL where the significance of Kl and K , are well understood.'* However, the differentiations required for substitution into eqn (8') quickly become unmanageable owing to the fact that eqn (18) is implicit in 8,.Therefore we did not carry the expansion beyond the second derivative in eqn (1 0). Another suitable choice, although physically less realistic than eqn (18) at least in terms of its neglect of intermolecular interactions, is the Langmuir isotherm: Note that temperature is implicitly contained in z owing to the method of non- dimensionalizing energy variables in use here. Unlike eqn (18), eqn (19) does not exhibit so-called cooperative effects, but it has, nevertheless, been frequently used as a local isotherm in studies of heterogeneity. It also avoids the problem of distinguishing between the patchwise and random models of adsorption heterogeneity. There is a further, more pragmatic, reason for its use here, and that is to avoid problems associated with taking an arbitrary number of derivatives.Rewriting eqn (19) in the form where we have defined x by it can be shown that the general form of the nth derivative is (20) x = CpeZ (21) X &(x) = - x+ 1 n A recursion relation may be derived which gives the coefficients an+,, in terms of the an, as follows: - an+,,, - -%.,I an+,, n+l - an, n = 1. a,,,, i+l = (n - i+ l)a,, - ( i + l)an, i+l (i = 1,2, . . . , n- 1) -1472 ADSORPTION OF GASES ON HETEROGENEOUS SURFACES The procedure may be initiated by noting that so that al, = 1 . By recursion, as many z-derivatives of the Langmuir equation as desired may be calculated. We note that the expansion will consist of terms involving the product of a central moment of the distribution and a rational function of p.Eqn (20)-(22c) together with eqn (17) may now be easily combined in an algorithm for generating approximations to any desired order, which we implemented as a short program in BASIC computer language. For purposes of illustration, the first few terms of the expansion, eqn (S'), are presented using the Langmuir local isotherm with x evaluated at the mean energy z = p. For an arbitrary density function This technique may also be applied to the Jovanoviclo equation or, rewritten in terms of x from eqn (21), &(x) = 1 -eP5. The general form of the nth derivative here is d"8 n L = e-5 an,ixi dzn i - 1 I I I I I I I I 1 Z Fig. 1. Normal distribution of energies used to calculate adsorption isotherms.Mean energy p = 10. (1) o/p = 0.075, (2) o/p = 0.125.J. A. O’BRIEN AND A. L. MYERS 0.8 I 1 I I I 0.7 - 0.6 - 0.5 - 1473 3 P Fig. 2. Exact numerical integration of eqn (1) using Langmuir local isotherm and distributions shown in fig. 1 . (1) o/p = 0, (2) g/p = 0.075, (3) g/p = 0.125. and we have, once again, a set of recursion relationships - an+,, 1 - an, 1 = 1 an+l,i+l = (i+ l)an,i+l-an,i %+I, n+1 - -an, n - (i = 1, . . . , n - l ) - Once again, the procedure is started by noting that so that al, = 1. To evaluate the effectiveness of the expansion exact results were necessary. These were obtained by numerical integration of eqn (1) using eqn (1 5) and (19) asf(z) and 8,( T,p, z), respectively. The algorithm employed was that of Romberg integration.19 It was necessary to choose specific values of the constants p, C and B before proceeding.We set p = 10, which corresponds to an energy of 24.94 kJ mol-1 at 300 K. This is the correct order of magnitude for a typical heat of adsorption. The value of C = e-l0 was chosen so that 8 = 0.5 atp = 1 .O. Since the product (Cp) is dimensionless, the units of p are insignificant. Finally, the parameter 0 was varied to investigate the effect of heterogeneity. RESULTS OF MODEL STUDY The density function,flz), is plotted on fig. 1 for values of B studied. Fig. 2 presents the exact, numerically integrated isotherms for each B. The corresponding results,1474 ADSORPTION OF GASES ON HETEROGENEOUS SURFACES 0 0.5 1.0 1.5 2.0 2 5 3.0 P Fig. 3. Convergence of eqn (8’): percentage error as a function of pressure and number of terms in series.Labels 0, 1, 2 and 3 refer to the number of correction terms added to O,(T,p,p). Table 1. Parameters of eqn (28) for CO, adsorption data20 m RTP RTa C T/K /mmol g-l /kJ mol-l /kJ mol-l /kPa-l A 212.7 1 1.005 19.371 1.8 6.933 x 0.0188 260.2 10.500 18.848 2.1035 6.933 x 0.006 301.4 9.2 197 19.003 2.2153 6.933 x 0.0084 obtained using the approximation eqn (S’), are presented in terms of their percentage deviation from the exact results: ) x 100. exact - approx. Fig. 3 gives a plot of 6 against p for approximations with varying numbers df terms, using G = 0.75. It is apparent that the series gives an oscillating convergence. The maximum deviation after four correction terms is 0.05%. After only one correction term, the deviation is < 1.5%.It should be mentioned in this connection that the addition of m correction terms, when using the Gaussian distribution, involves terms up to the 2mth derivative owing to the fact that every second term in the expansion, eqn (S’), disappears because of eqn (17). A similar calculation of 6 against p for 0 = 1.25 yielded oscillatory, increasing deviations after four correction terms. If the series eventually converges for CT greater than unity, too many terms are necessary for practical use. Although no rigorous mathematical analysis of the convergence properties of eqn (8’) has been undertaken here, it has been shown that rapid convergence may be expected for values of G < 1. Recalling that the dimensionless value of CT under discussion is obtained by dividing it by RT, the series converges for ratios of o/p < 0.1,J.A. O’BRIEN AND A. L. MYERS 1475 10.0 8.0 7 6.0 00 c.l E 3 4.0 2.0 Fig. 4(a) and (b). For legend see page 1476. or at room temperature for values of CT < ca. 2 kJ mol-l. A similar study using the uniform distribution fails to converge eventually, but is convergent for values of o / p > 0.1. Thus it would appear that the Gaussian distribution gives a ‘worst-case’ estimate for convergence. These results might be phrased in the alternative form that the expansion procedure is convergent for ‘moderately’ heterogeneous surfaces. For such surfaces the method produces an approximate (although slightly unwieldy) version of eqn (I) which is explicit and, at least in principle, may be used with an arbitrary local isotherm and distribution of energies.APPLICATION TO EXPERIMENTAL DATA Although the series expansion, eqn (S’), has not converged after one correction term, truncation at this point should give an equation which can be used to correlate experimental isotherm data. Thus the first-order approximation, using the Langmuir1476 ADSORPTION OF GASES ON HETEROGENEOUS SURFACES P m a Fig. 4. Comparison of eqn (28) with CO, adsorption data for various temperatures. (a) 212.7, (b) 260.2 and (c) 301.4 K. local isotherm, has been fitted to the experimental data of Reich et aL20 for the adsorption of CO, on an activated carbon at three temperatures. This equation is written as where x is given, as before, by eqn (21). The resulting parameters, as well as the objective function for the fit are given in table 1 .Fig. 4(a)-(c) show the comparison between the equation and the data points, and the agreement is seen to be good. Thus this modification of the Langmuir equation has endowed it with enough heterogeneous character to describe adsorption on an activated carbon. These results indicate the possibility of modifying any of the usual homogeneous isotherm equations to account for heterogeneous surfaces. Even implicit isotherms, such as eqn (18), may be used in this way since it is possible, although tedious, to calculate d26/dz2, the only derivative necessary in this case. However, only the Langmuir case has been investigated in this work.J. A. O’BRIEN AND A. L. MYERS 1477 NOMENCLATURE an, i C D E f H HL rn n n c a l ~ nexptl P R T X Z a3 a4 6 A e 8, P P n 0 coefficients in eqn (22a) and (26a) constact in Langmuir isotherm, eqn (19) domain of energy distribution energy of adsorption probability density function for energy Henry’s constant, eqn (1 1) Henry’s constant on local patch adsorption saturation capacity number of moles adsorbed per gram of adsorbent calculated value of n from eqn (28) experimental value of n pressure gas constant temperature dimensionless group, eqn (21) dimensionless energy of adsorption coefficient of skewness off(z), eqn (9) kurtosis off(z), eqn (9) percentage deviation, eqn (24) objective function minimized in data fit, eqn (29) fractional coverage, entire surface fractional coverage, local patch mean of energy distribution nth central moment of energy distribution standard deviation of energy distribution This work was supported by National Science Foundation Grant CPE-8 1 17 188.REFERENCES W. A. House Adsorption on Heterogeneous Surfaces, in Colloid Science (Spec. Period. Rep., The Royal Society of Chemistry, London, 1983), vol. 4. M. Jaroniec, Adv. Coll. Interface Sci., 1983, 18, 149. R. R. Zolandz and A. L. Myers, Prog. Filtration Separation Sci., 1979, 1, I . R. Sips, J . Chem. Phys., 1948, 16, 490. R. Sips, J. Chem. Phys., 1950, 18, 1024. D. N. Misra, Sur- Sci., 1969, 18, 367. J. Toth, W. Rudzinski, A. Waksmundzki, M. Jaroniec and S . Sokolowski, Acta Chim., 1974, 82, 1 I . M. Jaroniec and S. Sokolowski, Colloid Polym. Sci., 1977, 255, 374. ’ D. N. Misra, J. Chem. Phys., 1970, 52, 5499. lo D. S. Jovanovic, Kolloid Z., 1969, 235, 1203. l1 A. L. Myers and D. Y. Ou, AIChE Symp. Ser., 1983, 230, 79. l2 S. Ross and I. D. Morrison, Surf. Sci., 1975, 52, 103. l3 W. A. House and M. J. Jaycock, Colloid Polym. Sci., 1978, 256, 52. l4 E. W. Sidebottom, W. A. House and M. J. Jaycock, J . Chem. SOC., Faraday Trans. I , 1976,72,2709. l6 W. A. House and M. J. Jaycock, J. Chem. SOC., Faraday Trans. 1, 1977,73, 942. l6 V. K. Rohatgi, An Introduction to Probability Theory and Mathematical Statistics (Wiley, New York, l7 R. A. Pierotti and H. E. Thomas, J. Chem. SOC., Faraday Trans. I, 1974,70, 1725. l8 S. Ross and J. P. Olivier, On Physical Adorption (Interscience, New York, 1964), pp. 17, 170 and l9 B. Carnahan, H. A. Luther and J. 0. Wilkes, Applied Numerical Methods (Wiley, New York, 1969), 2o R. Reich, W. T. Ziegler and K. A. Rogers, Ind. Eng. Chem., Process Des. Dev., 1980, 19, 336. 1976), pp. 93 and 221. 180. p. 90. (PAPER 3/ 121 8)

ISSN:0300-9599    

DOI:10.1039/F19848001467

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I , 1984,80, 1479-1489 Paramagnetic Metal and Oxygen Species Observed with [Co,(CO),,l and [Rh,(CO),,I Carbonyl Clusters Pyrolysed on ?-Alumina and Zirconia Supports BY TIZIANA BERINGHELLI, ANTONELLA GERVASINI, FRANCA MORAZZONI* AND DONATELLA STRUMOLO Dipartimento di Chimica Inorganica e Metallorganica, Universita di Milano, Via Venezian 21, 20133 Milano, Italia AND SECONDO MARTINENGO Centro di Studio per la Sintesi e la Struttura dei Composti dei Metalli di Transizione nei bassi stati di ossidazione, Via Venezian 21, 20133 Milano, Italia AND LUCIANO ZANDERIGHI Dipartimento di Chimica Fisica ed Elettrochimica, Universita di Milano, Via Venezian 21, 20133 Milano, Italia Received 26th July, 1983 The pyrolysis of y-Al,O,- and Zr0,-supported [Co,(CO),,] and [Rh,(CO),,] gives highly dispersed metal systems whose electronic properties have been studied by e.s.r.spectroscopy. Paramagnetic formal Rh" species (g,, = 2.21, gl = 2.11) were observed on Rh/y-Al,O, samples. The pattern of the hyperfine structure suggests that metal-metal interactions are present and that paramagnetic transition-metal carbonyl clusters can be assumed as model compounds. No paramagnetic species involving the metal were observed for either y-Al,O,- or Zr0,-supported cobalt samples. This behaviour is consistent with a higher acidic character of y-Al,O, with respect to ZrO,. On contact with 0, and CO the following paramagnetic species were observed: O;-A13+ (g,, = 2.034, g , = 2.00) over Co/y-Al,O, samples, O;-Rh"' (8, = 2.09, g, = 2.03, g , = 2.00) over Rh/y-Al,O, samples, O;-RhlI1 and O;-Zr4+ (8, = 2.03, g , = 2.00) over Rh/ZrO, samples and CO-RhII (g = 2.034) over Rh/y-Al,O, samples.The 0, and CO bond strengths are higher for the y-Al,O,-supported samples than for the Zr0,-supported samples and are tentatively related to the catalytic activity of these metal systems. In spite of the technological importance of supported transition-metal systems, there have been few extensive physico-chemical investigations on the nature of the active metal species. Specifically, whenever it was observed that the activity and selectivity depended on the properties of the support, the presence of a strong metal-support interaction was postu1ated.l Because of the dearth of experimental data it is still difficult to assess the relevance of such an effect.It is well known that supported mononuclear transition-metal complexes can give highly dispersed metal catalysts2 and that molecular spectroscopy and magnetic techniques can be helpful in characterizing the electronic structure of the surface metal centres. However, the properties of these systems are different from those of conventional metal catalysts. In recent years another technique, the decomposition of supported metal carbonyl cluster^,^ has been tested for the preparation of highly dispersed metal catalysts. It is reasonable to expect that the activity of such systems is more analogous to that of 14791480 PARAMAGNETIC METAL AND OXYGEN SPECIES conventional metal catalysts and that the electronic configuration of the small metal aggregates is more amenable to investigation by molecular physical methods. These reasons, together with the increasing interest being shown in catalytic processes assisted by transition-metal clusters4 and the lack of extensive physico-chemical characterization of such catalysts, led us to consider the possibility that: (i) the paramagnetic metal species which are possibly present on the surface of pyrolysed supported-metal carbonyl clusters could be investigated by e.s.r.spectroscopy and (ii) new paramagnetic species could be generated by reaction with oxygen and/or carbon monoxide, which are used to probe the electronic properties of both paramagnetic and diamagnetic surface species. The e.s.r. work reported in the present paper was carried out on systems differing in either the supported transition-metal carbonyl cluster or the support metal oxide in order to investigate the different roles of the transition metal and of the metal-support interaction in the generation of the paramagnetic species.The interpretation of the e.s.r. spectra was based on previous work on paramagnetic transition-metal car- bony1 cluster^.^? The reaction with oxygen was considered because the paramagnetic 0, molecule can easily interact, depending on the electronic properties of the surface transition-metal centres, with either paramagnetic or diamagnetic species and may often produce new paramagnetic oxygenated products. Moreover, the use of 0,, which usually acts as a z-acceptor system, could be diagnostic of the ability of the metal to transfer electrons from its dn orbitals to the z* molecular orbitals of the acceptor molecule.Several catalytic processes probably require this mechanism and the catalytic activity could thus depend on the amount of electron back-donation. The interpretation of the e.s.r. spectra of the oxygenated product was based on extensive investigations of the oxygen paramagnetic species generated by supported transition- metal ions' and by undiluted or supported transition-metal complexe~.~~~ Finally, the reaction of the surface paramagnetic metal species with CO was tested, in view of the possible use of these systems as catalysts for the CO+H, reaction.1° EXPERIMENTAL PREPARATION OF [Co,(CO),,]/y-Al,O,, ZrO, AND [Rh,(CO),,]/y-Al,O,, ZrO, Solutions of [Co,(CO),,] and [Rh,(CO),,] in anhydrous deaerated n-pentane were added dropwise, under an inert atmosphere, to deaerated y-Al,O, or ZrO, suspended in the same solvent.The carbonyl solutions lost their colour on chemisorption of the cluster on the support. When the addition was complete the mixture was allowed to equilibrate for 2 h while being stirred; the solid was then filtered off. Samples were dried in vacuo at room temperature for 2 h and stored under a N, atmosphere. Contact with air was avoided as it could affect the formation of the paramagnetic species after decarbonylation. For example, if [Rh,(CO),,]/y- A1,0, was allowed to oxidize in air until the Rh(CO), fragment species was formedll and subsequently decarbonylated (see later), no paramagnetic species were observed.The percentages of the supported compound, expressed as g of carbonyl metal cluster per 100 g of the support were: [Co,(CO),,]/y-Al,O,, 2.9 %; [Co,(CO),,]/Zr02, 2.5 %; [Rh,(CO),,]/y- Al,O,, 3 % ; [Rh,(CO),,]/ZrO,, 2.5 % . The i.r. spectrum of [Rh,(CO),,]/y-Al,O, shows two absorptions attributable to CO vibrations of [Rh,(CO),,], as already observed by Ichikawa.', [Rh,(CO),,] and [Co,(CO),,] were prepared as in ref. (13). y-Al,O, was from Akzo Kemie, CK 300, 50-150 mesh, thermally pretreated for 1 h at 200 "C, 1 h at 300 "C, 1 h at 400 "C and 6 h at 550 "C under an inert atmosphere (surface area 200 m2 ggl). ZrO, was from Strem, containing 98% ZrO, and 2% y-Al,O,, 50-150 mesh, thermally pretreated at 500 "C for 7.5 h under an inert atmosphere (surface area 70 m2 g-l).VACUUM TREATMENT OF THE SAMPLES [Co4(C0),,]/y-A1,0,, ZrO, and [Rh,(CO),,]/y-A120,, ZrO, were treated in vacuo ( lo-, Pa) for 2 h at 180 "C (cobalt cluster) and 250 "C (rhodium cluster). The decarbonylation progressT. BERINGHELLI et al. 1481 was monitored by i.r. spectroscopy. Treatments in uucuo were performed in a small flask connected to an e.s.r. tube (internal diameter 3 mm). Contacts with gas at controlled pressures were carried out on a standard gas-vacuum line. The e.s.r. spectra were recorded on a Varian E-109 spectrometer equipped with an automatic Varian temperature control. COMPUTATIONAL METHOD FOR THE SPIN CONCENTRATION The spin concentration was determined by double integration of the area under the resonance lines, taking as reference area that of the Varian weak pitch, which contains 1013 spin per cm of length.The sensitive region of the e.s.r. cavity was 1 cm. The apparent density of samples was 1 g cm-, for y-Al,O,-supported samples and 3.5 g cmP3 for Zr0,-supported samples. The broad shape of some resonances and the overlap between others limited the accuracy of the quantitative evaluation and thus only the orders of magnitude of the spin concentration can be safely assumed. RESULTS DETECTION AND CHARACTERIZATION OF PARAMAGNETIC SPECIES SIJPPORT SYSTEMS y-Al,O, was treated in vacuo ( Pa) at 180 and 250 OC, then contacted with argon (26.6 kPa). The only observed resonances were those due to the high-spin tetragonally distorted Fe3+ centres ( g = 4.25). The three-line low intensity signals at g x 2 could be due to very small quantities of low-spin rhombic Fe3+.14 Control experiments on thermally treated y-Al,O, contacted with either 0, or CO, under the same conditions as the samples, showed no changes in the support resonance lines, thus ruling out interaction of these gases with the paramagnetic centres of the support.ZrO, was treated in vacuo ( lo-, Pa) at 180 and 250 "C, and contacted with argon and then with 0, and CO under the same conditions as the samples. The support, after thermal treatment, showed the resonance lines of a radical species and were not affected by contact with the gases. As the observed ZrO, signal could interfere with the signals of the metal-containing samples, their spectra were corrected accordingly. DECARBONYLATED SYSTEMS Co/y-Al,O,, under an argon atmosphere (26.6 kPa), gave symmetrical resonance lines, whose linewidth (12 G) and position ( g = 2.00) suggest the presence of a radical species.The resonance lines were not observed before decarbonylation of [Co,(CO),,]/y-Al,O,. It is difficult to draw any conclusion as to the nature of this paramagnetic species, although the cobalt valence electrons are not responsible for this signal. On the other hand, because of the low intensity of these lines it was decided not to investigate the nature of this radical species further. Rh/y-Al,O, under an argon atmosphere (26.6 kPa) gave strong asymmetrical lines, whose position (gll = 2.21, gl = 2.1 1) and width (fig. 1) suggest the location of an unpaired electron on Rh centres.The lines are strongly temperature dependent; they decrease in intensity and broaden at room temperature, as expected for 4d transition- metal-ion resonances, for which the spin-lattice relaxation time controls the linewidth. The presence of hyperfine structure in the perpendicular region of the resonances indicates interaction with more than one Rh(l = 1/2) nucleus. Although the spectra are easily reproducible, the number of lines is difficult to define. Measurements of the hyperfine coupling constant are not warranted as the spectral pattern depends on the pyrolysis temperature (see fig. 1). The multi-line nature of the spectra seems to arise from a contribution of several polynuclear paramagnetic species with very similar g values rather than from the hyperfine coupling of a single species.The ground state of the paramagnetic polynuclear species can be assigned by assuming a conventional paramagnetic centre6 and it requires a molecular orbital comprised mainly of the same 49 FAR 11482 PARAMAGNETIC METAL AND OXYGEN SPECIES 4 \r 200 G Fig. 1. X-band e x . spectra recorded under an argon atmosphere (26.6 kPa) at - 150 "C: (a) y-Al,O, ; (6) Rh/y-Al,O, from [Rh,(CO),,]/y-Al,O, pyrolysis at I50 "C ; (c) Rh/y-Al,O, from [Rh,(CO),,]/y-Al,O, pyrolysis at 250 "C. atomic orbitals suggested for the conventional centre by the crystal-field analysis of the g tensor. The values of the g-tensor components of a mononuclear Rh paramagnetic species with gI1 > gl 2 2 could be diagnostic of two different electronic configurations, d7 RhT1 with the unpaired electron in the dz2--.2/2 orbital or d9 Rho with the unpaired electron in the same orbital, corresponding to compressed or elongated tetragonal symmetry, respectively.Contrasting reports on the assignment of the Rh oxidation state based on e.s.r. data have been p~b1ished.l~ In the present case the clustered Rh atoms could be represented magnetically by a d9 or a d7 centre; both are in agreement with an oxidative interaction between Rh and y-Al,O, as the number of interacting Rh atoms is not known. A definite conclusion on the electronic configuration of the Rh species will be drawn from an analysis of the interaction with oxygen.T. BERINGHELLI et al. 1483 Fig. 2. X-band e.s.r. spectra, recorded at - 150 "C: (a) ZrO,; (b) Rh/ZrO, recorded under an 0, (10 Pa) atmosphere; (c) Rh/y-Al,O, recorded under an 0, (10 Pa) atmosphere; ( d ) Rh/ZrO, recorded under an 0, (10 Pa) and CO (10 Pa) atmosphere; ( e ) Co/y-Al,O, recorded under an 0, (10 Pa) atmosphere.Co/ZrO, and Rh/ZrO, were treated analogously to the y-Al,O,-supported samples and did not give resonance lines other than those due to the support. REACTION WITH OXYGEN Co/y-Al,O, was contacted with 0, (26.6 kPa) at room temperature for 5 min. The e.s.r. spectrum shows strong new resonance lines having a different shape and higher intensity than the radical species under an argon atmosphere. The intensity and resolution of the new signals increase at low pressure (lo-, Pa) and the magnetic anisotropy becomes evident (fig. 2).The spectrum is obtained either at low (- 150 "C) or at room temperature. This behaviour, together with the values of the g-tensor 49-21484 PARAMAGNETIC METAL AND OXYGEN SPECIES components (gll = 2.034, gl = 2.00) suggests that the resonances can be attributed to 0; fixed on A13+ c e n t r e ~ . ~ * l ~ The bonding interaction between 0; and A13+ cannot be removed by vacuum or vacuum and thermal treatment (1 80 "C, 1 0-3 Pa), in agreement with the usual behaviour of the A13+-O; species.17 Rh/y-Al,O, was contacted with 0, (26.6 kPa) at room temperature for 5 min and did not reveal any resonance lines; the lines seen under the argon atmosphere are no longer noticeable. The sample, evacuated to 10 Pa, showed new resonance lines (fig. 2) whose intensity and resolution are strongly dependent on the temperature.At - 150 "C g-tensor anisotropy is evident (gl = 2.09, g , = 2.03, g , = 2.00). If we consider the rhombic anisotropy of the g tensor, the value of its principal components and the relative intensity of the e.s.r. lines, satisfactory agreement is found with the g-tensor properties of the monomeric superoxo compound [Rh(en),C1(O,)]+ls and of dimeric superoxo-bridged compounds such as [(Rh(en),Cl),(p and O,)I3+ l8 [(Rh(bpy),C1),(p-0,)]3+.1B The values of the g-tensor components suggest that the unpaired electron occupies a 0, n* orbital, with a small contribution from the Rh d . orbitals. However, the amount of the actual charge on 0, cannot be defined from the e.s.r. data and the location of the unpaired electron on the 0, n* orbitals does not necessarily mean that one negative charge is transferred to 0,.For simplicity, whenever we refer to the RhIL1-0; formulation, we mean that the negative charge is localized mainly on 0,. The anisotropic character of the g tensor implies that the 0-0 bond is bent with respect to the Rh-0 bond direction, but we cannot decide between monomeric and dimeric superoxo species. The suggestion of a superoxo- bridged species does not disagree with the hypothesis that Rh-Rh interactions are not fully removed by the supporting processes. The formation of 0x0-bridged rhodium clusters was also tentatively suggested for Rh/y-Al,O, samples obtained from pyrolysis of [Rh6(CO)l,]/y-Al,0,.20 The interaction between 0, and Rh/y-Al,O, is not removed by vacuum treatment ( Pa); vacuum and thermal treatment (250 "C) remove 0, from the coordination, but new e.s.r.lines, different from those of the unoxygenated precursor, appear. These lines are probably related to the irreversible oxidation of Rh and therefore the usefulness of 0, as a probe for the electronic state of the surface metal centres or as model of catalytically important electron-acceptor systems is doubtful. Irreversibly modified systems were not studied. RhIII-0; formulation requires that the spectrum observed after pyrolysis is that of a RhI1 species and that the oxygen adduct is formed by a one-electron-transfer process, as invoked for d7 centres.,' Co/ZrO,, after contact with O,, did not give any resonance lines in addition to those observed for 21-0,.Rh/ZrO, was contacted with 0, (26.6 kPa) for 5 min at room temperature. Comparison with the spectra of ZrO,, analogously treated, did not reveal any new paramagnetic species. On evacuation at 10 Pa two different new paramagnetic species appeared at - 150 "C (fig. 2): one was identical to that observed on interaction of 0, with Rh/y-Al,O, and the other is almost temperature independent and could be easily separated from the first by recording the spectrum at room temperature. The values of the g-tensor components (g,, = 2.03, gl = 2.00) suggest that the latter signal could be attributed to 0; fixed on Zr4+ centres;,, as the assignment of the perpendicular component is doubtful when strong ZrO, signals are present, other evidences on the nature of this resonance will be given in the following section.Both 0; signals disappear on vacuum treatment Pa) at room temperature for 24 h, thus ruling out the possibility that the RhI"-O; species found on the Rh/ZrO, sample is due to the low content of y-Al,O, in ZrO,.T. BERINGHELLI et ai. L 1485 Fig. 3. X-band e.s.r. spectrum of Rh/y-Al,O, recorded under a CO (10 Pa) atmosphere. REACTION WITH co Co/y-Al,O,, ZrO, and Rh/ZrO, were contacted with CO (26.6 kPa) at room temperature. No new resonance lines appeared. Rh/y-Al,O, was treated with CO (26.6-10 kPa) and revealed strong modification of the e.s.r. spectrum with respect to that observed under argon (fig. 3). Signals attributed to formal Rhl* polynuclear species disappear. The new resonance lines, whose intensity does not depend on the CO pressure, are because of their width still indicative of the location of the unpaired electron on the Rh centre.However, their small paramagnetic anisotropy and the displacement of the g-tensor value from that of the free electron suggest that the extent of electron location on Rh decreases on interaction with the CO molecule. In the absence of hyperfine structure, no direct proof of the nuclearity of the CO adduct is available. Thermal treatment (100 "C) in uucuo (lo-, Pa) removes the Rh-CO interaction. REACTION WITH co AND 0, On successive contact with CO and O,(26.6 kPa), and further evacuation to 10 Pa, Co/y-Al,O, reveals the lines of 0; fixed on A13+, whose stability properties have been already described, and Co/ZrO, does not show any other resonances beside those of the support.Rh/y-Al,03 reveals the lines of 0; fixed on A13+, in addition those of the Rh-CO interaction product. Neither Rh-CO or A13+-O; interactions could be removed by vacuum treatment. Thermal treatment (100 "C) in vacuu Pa) removes only the Rh-CO interaction. Rh/ZrO, reveals only the species attributed to 0; fixed on Zr4+ (fig. 2). Following vacuum treatment ( Pa) for 4 h the intensity of these lines decreases and the resonances due to 0; fixed on Rh become evident. After 24 h under vacuum at room temperature all the reduced forms of oxygen are removed from the ZrO, surface. The formation of 0, reduced species is a reversible process. The behaviour observed following contact with CO and then with 0, can be interpreted as follows : on y-Al,O,-supported samples the CO molecule successfully competes with oxygen for the interaction with Rh, so that the only available coordination sites for 0; are provided by the support and the adduct can be formulated as A13+-O;. If we invoke the same competitive process in the case of Zr0,-supported samples we can conclude that the 0, coordinates on the Zr4+ centres, thus confirming that these e.s.r.signals are attributable to Zr4+-O; species. Vacuum1486 PARAMAGNETIC METAL AND OXYGEN SPECIES treatment switches 0; from Zr4+ to Rh, probably as a consequence of the partial removal of CO from Rh. DISCUSSION The values of the magnetic tensors of the species are collected in table 1, together with the number of paramagnetic centres derived from the e.s.r.analysis. The assignment of the g-tensor components was based on the relative intensity of the resonances. We have made the following observations. (i) The number of paramagnetic centres on Rh/y-Al,O, is much lower than the number of supported Rh, units (ca. 0.02%). This can be explained by supposing that besides the Rh, units interacting with y-Al,O,, higher aggregates also exist. It is also likely that not all the A13+ centres have the same acidic strength, as found for thermally pretreated y-Al,O,, and their electron-attracting power is not sufficient to produce electron-deficient paramagnetic centres. (ii) 0, interaction with Rh/y-Al,O, involves an amount of Rh centres slightly less than or comparable to the number of paramagnetic Rh centres observed before contact with oxygen. This behaviour seems to be in keeping with that observed when the species interacting with 0, are ionic and monon~clear.~~ (iii) The presence of paramagnetic metal centres before oxygenation is not a necessary condition for electron transfer to 0,.Indeed the reduction of 0, over Rh/ZrO, samples proceeds in the absence of a detectable number of paramagnetic metal centres, producing the same number of oxygenated centres as found over Rh/y-Al,O, samples; moreover, 0, reduction over Co/y-Al,O, produces more paramagnetic centres than in the starting material, confirming that no relationship exists between the initial and final paramagnetic species. On the other hand, only paramagnetic centres can generate paramagnetic CO adduc t s. The above results suggest that the formation of paramagnetic species by decar- bonylation of supported clusters depends on either the support and/or the metal properties.Indeed the presence of detectable paramagnetic centres for only y- Al,O,-supported samples is consistent with the greater acidic character of y-Al,O, compared with ZrO,. The larger number of paramagnetic centres found for Rh samples with respect to Co samples (whatever the nature of the cobalt paramagnetic species) is due to better overlap between the A13+ p orbitals and the Rh d, orbitals, the overlap being caused by a larger expansion of the Rh orbital lobes. The resonance lines found for Rh/y-Al,O, samples suggest that the Rh-Rh interaction in the supported Rh cluster is not totally removed by the metal-support interaction; therefore the metal-metal interaction is more likely to occur for the e.s.r.-inactive Zr0,-supported samples, where the metal-support interaction is less than on y-Al,O,. Note that if the support does not remove the metal-metal interaction, the vacancy in the valence electrons generated by the interaction of 7-Al,O, with rhodium should be distributed among several metal centres, as usually observed in stable paramagnetic metal carbonyl clusters where hyperfine structure is e ~ i d e n t .~ * , ~ The oxidizing effect of y-Al,O, on a metal centre could be transferred to another atom, which may not be in contact with the oxide surface. The possibility that a perturbation generated on a metal centre, for example after a change in the ligand molecule, is transferred to other metal centres in a cluster is well known in the chemistry of metal carbonyl interaction with the support, in our metal systems, could result in such an effect.The results of the e.s.r. investigation concerning the chemisorption of 0, and COTable 1. E.s.r. dataa for paramagnetic centres for supported metalsb no. of 0, co paramagnetic pressure pressure centres paramagnetic /spin g-l speciesC sample /Pa /Pa g,, g x x or yy g y y or xx notes c04(c0)1 2/YmA12'3 0 10 0 10 10 Rh4(C0)12/y-A1203 0 0 10 Rh4(C0)12/Zr02 0 10 0 10 0 0 10 10 0 0 10 10 0 0 10 10 2.00 2.034 2.034 2.2 1 2.09 - 2.034 2.034 { 2.034 - - 2.03 2.00 2.00 2.00 2.1 1 2.03 - 2.034 2.00 2.034 2.03 2.00 2.00 - - 2.00 2.00 2.00 2.1 1 2.00 - 2.034 2.00 2.034 - 2.00 2*oo 1 - 2.00 0.28 x 1014 radical 0.56 x 1015 ~ 1 3 + - 0 ; 1.12 x 1015 ~ 1 3 + - 0 ; 0.46 x 10'' RhlI1- 0, - - 0.56 x 10l6 RhII 0.28 x 10l6 RhII-CO 1 0.42 x 1015 ~ 1 3 + - 0 ; 1.38 x 1015 RhII-CO - stable at 180 "C and lop3 Pa stable at 180 "C and lop3 Pa multinuclear species thermally removed at 250 "C, - c3 % lop3 Pa with irreversible z sample modification p Pa z z % # thermally removed at 100 "C, stable at 100 "C, Pa removed by vacuum treatment at Pa, reversible process on vacuum treatment ( Pa) Rh"I-0, is formed as well as - - zr4+-0; a From spectra at on y-Al,03 = 0.32 x - 150 "C.1020 g-l. Rh4 units supported on y-Al,03 = 0.24 x 1020 g-l; Rh, units supported on ZrO, = 0.17 x 1020 g-l; CO, units supported The meaning of the oxidation state indicated for the paramagnetic species is detailed in the text.1488 PARAMAGNETIC METAL AND OXYGEN SPECIES molecules can be interpreted as follows.Oxygen is reduced by the supported metal clusters with the reaction pathway and products depending on the electronic structure of the metal and/or the metal-support interaction. y-Al,O,-supported Co is not rich enough in d, electrons to stabilize bonding with oxygen. It is well known that bonding between oxygen and transition-metal ions is partially a d, - n* bond interaction; in the presence of A13+ centres it is likely that electron transfer from the cobalt d, orbitals to the A13+p orbitals successfully competes with that to the 0, n* orbitals. The y-Al,O, surface acts as an electron-transfer medium for the electrons which reduce the 0, fixed on AP+ to 0;.By a different reaction pathway 0, could interact with cobalt and after reduction could migrate over A13+; however, the formation of O;, in the presence of CO, is difficult to explain. Reduction of 0, does not take place over Co/ZrO, because of the lower ability of this support to accept electrons from the cobalt. In the case of Rh/y-Al,O, and Rh/ZrO, the metal has more expanded d, orbitals than cobalt, thus allowing better overlap with the oxygen n* orbitals. Bonding with oxygen is partially covalent, as suggested by the anisotropy and the temperature dependence of the e.s.r. resonances; moreover, it is very likely that the internuclear axis of 0; is bent with respect to the Rh-0 axis as the interaction with 0, should depend strongly on the expansion of the d, orbital lobes.On the other hand, the reduction of 0, by Rh/y-Al,O, after contact with CO can be explained only by electron transfer from rhodium ' via y-Al,O, ' as the Rh centres are presumably bound to CO. With ZrO, supports both Rh and Zr centres are involved in the fixation of O;, but it is improbable that 0, can accept electrons from rhodium 'via support' because of the low electron-accepting properties of ZrO,. Most likely the superoxide anion is produced on Rh centres, then partially migrates towards Zr centres. Note that the strength of Rh-0, bonding is lower on Zr0,- than on y-Al,O,-supported samples, so that by simple vacuum treatment 0; jumps between Rh and Zr and is then removed from the coordination. The same conclusion can be reached for the CO interaction; indeed the vacuum treatment removes CO from the Rh/ZrO, samples, while only thermal treatment destabilizes the Rh-CO bonding in y-Al,O,-supported samples.The higher strength of RhIII-0; and RhII-CO bonds in y-Al,O,-supported samples is due to the higher acidic character of Rh centres for y-Al,O,-supported samples. This situation is favourable to a CJ interaction between Rh and the chemisorbed molecules. CONCLUSIONS One of the most important problems related to the physico-chemical investigation of supported-metal systems is concerned with knowledge of the oxidation state of the metal. There seems to be evidence that the selectivity of such systems depends on metal centres having a positive oxidation state.l'~,~ By dispersing metal centres over supports of different acidity it is possible to obtain the metal in an oxidized form following interaction with support.Conventional catalysts, obtained by reduction of supported metal ions, show, in general, oxidized centres whose amount is limited to those interacting with the support. The formation of large metal particles hinders the interaction and lowers the reactivity of the metal centres towards the molecules involved in the catalytic reactions. On the basis of our results it seems reasonable that, when the metal precursor is a pyrolysed cluster, the metal on the surface is not a mononuclear system and the valence positive charge generated by support interaction can reasonably be distributed among all the interacting metal centres. Thus by means of metal clusters we can increase the availability of active metal centres and have more chance than withT.BERINGHELLI et al. 1489 mononuclear metal complexes to affect the reactivity through the choice of the support. Finally, although the number of paramagnetic centres found on the surface is small compared with the total number of metal atoms, they can still act as catalytic sites in some catalytic reactions, as has been shown in the case of Ag-supported systems.17 Samples of Rh/y-Al,O, and Rh/ZrO,, obtained from pyrolysis of [Rh,(CO),,], have been tested as catalysts for the CO + H, reaction27 and show that: (i) y-Al,O,-supported samples produce methane and higher hydrocarbons and (ii) Zr0,-supported samples selectively form oxygenated products.Our results on the e.s.r.-active species show that the interaction of n* acceptor molecules such as CO and 0, with Rh is stronger on y-Al,O, than on ZrO,. As a consequence, on y-Al,O, the stronger metal-CO interaction weakens the C-0 bond and leads to hydrocarbon formation; on ZrO, the weaker interaction cannot lead to dissociation of the C-0 bond and oxygenated products are obtained. We thank M.P.I. and the Italian C.N.R. (Progetto Finalizzato Chimica Fine e Secondaria) for financial support. S. J. Tauster, S. C. Fung and R. L. Garten, J. Am. Chem. SOC., 1978, 100, 170. R. H. Grubbs, Chemtech., 1977, 7, 512. B. C. Gates and J. Lieto, Chem. Technol., 1980, 195; 1980, 248. B. C. Gates, in Chemistry and Chemical Engineering of Catalytic Processes, ed. R. Prins and G.C. A. Schuit (Sijthoff & Noordhoff, Amsterdam, 1980), p. 427. G. Longoni and F. Morazzoni, J. Chem. SOC., Dalton Trans., 1981, 1735. T. Beringhelli, F. Morazzoni and D. Strumolo, J. Organomet. Chem., 1982, 236, 109. J. H. Lunsford, Catal. Rev., 1973, 8, 135. (a) F. Campadelli, F. Cariati, P. Carniti, F. Morazzoni and V. Ragaini, J. Catal., 1976, 43, 167; (b) G. Mercati, F. Morazzoni, M. Barzaghi, P. Carniti, V. Ragaini and F. Campadelli, J. Chem. SOC., Faraday Trans. I , 1979, 75, 1857. M. Barzaghi, T. Beringhelli and F. Morazzoni, J. Mol. Catal., 1982, 14, 357. A. Ceriotti, S. Martinengo, L. Zanderighi, G. Tonelli and A. Iannibello, Atti del ZZZ Congress0 Nazionale di Catalisi, Torino, Italy, 1982, p. 26; A. Ceriotti, S. Martinengo, L. Zanderighi, G. Tonelli, A. Iannibello and A. Girelli, J. Chem. SOC., Faraday Trans. I, 1984, 80, 1605. l 1 G. C. Smith, T. P. Chojnacki, S. R. Dagupta, K. Ivatake and K. L. Watter, Znorg. Chem., 1975, 14, 1429; A. K. Smith, F. Hughues, A. Theolier, J. M. Basset, R. Ugo, G. M. Zanderighi, J. L. Bilhou, V. Bilhou-Bougnol and W. F. Graydon, Znorg. Chem., 1979, 18, 3104. l 2 M. Ichikawa and K. Shikakura, in New Horizons in Catalysis, ed. T. Seiyama and K. Tanabe (Elsevier, Amsterdam, 1980), p. 925. I3 (a) S. Martinengo, G. Giordano and P. Chini, Znorg. Synth., 1980,20,209; (b) R. Ercoli, P. Chini and M. Massi-Mauri, Chim. Ind. (Milan), 1959, 41, 132. l4 J. Caspary, D. A. Lanzo and C. Niziak, Biochemistry, 1981, 20, 3868. l5 M. D. Sastry, K. Savitri and B. D. Joshi, J. Chem. Phys., 1980,73, 5568. l6 A. A. Gezalov, G. M. Zhabrova, V. V. Nikisha, G. B. Pariiskii and K. N. Spiridonov, Kinet. Catal., l 7 R. B. Clarkson and A. C. Cirillo Jr, J . Catal., 1974, 33, 392. l8 J. B. Raynor, R. D. Gillard and J. D. Pedrosa de Jesus, J. Chem. SOC., Dalton Trans., 1982, 1 165. Is H. Caldararu, K. De Armond and K. Hanck, Inorg. Chem., 1978, 17, 2030. 2o K. L. Watters, R. F. Howe, T. P. Chojnacki, Chia-Min Fu, R. L. Schneider and Ning-Bew Wong, J. 21 F. Basolo, B. M. Hoffman and J. A. Ibers, Acc. Chem. Res., 1975, 8, 384; B. B. Wayland and 22 M. Setaka, S. Fukuzawa, Y. Kirino and T. Kwan, Chem. Pharm. Bull., 1968, 16, 1240. 23 J. R. Katzer, G. C. A. Schuit and J. H. C. Van Hooff, J. Catal., 1979, 59, 278. 24 C . E. Strouse and L. F. Dahl, Discuss. Faraday Soc., 1969, 47, 93. 25 K. J. Karel and J. R. Norton, J. Am. Chem. SOC., 1974, 96, 6812. 26 S. Sivasankar, A. V. Ramaswamy and P. Ratnasamy, J. Catal., 1977, 46, 420. 27 M. Ichikawa, J. Chem. SOC., Chem. Commun., 1978, 566. 1968, 9, 462. Catal., 1980, 66, 424. A. R. Newman, J. Am. Chem. SOC., 1979, 101, 6472. (PAPER 3/1295)

ISSN:0300-9599    

DOI:10.1039/F19848001479

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1984, 80, 1491-1498 Adsorption of Goethite onto Quartz and Kaolinite B Y MARVIN c . GOLDBERG,* EUGENE R. WEINER~ AND PAUL M. BOYMELT U.S. Geological Survey, P.O. Box 25046 Mail Stop 424, Denver Federal Center, Lakewood, Colorado 80225, U.S.A. Received 27th July, 1983 The adsorption of colloidal goethite onto quartz and kaolinite substrates has been studied as a function of pH and NaCl concentration. Goethite adsorption was measured quantitatively by Fourier-transform infrared spectroscopy. The results indicate that adsorption onto both substrates is due primarily to coulombic forces; however, the pH dependence of adsorption is very different for the two substrates. This is explained by the fact that the surface charge on quartz is entirely pH-dependent, while kaolinite has surface faces which carry a permanent negative charge.Adsorption of goethite on to kaolinite increases markedly with increasing NaCl concentration, while adsorption onto quartz is relatively independent of NaCl concentration. This can be explained by the influence of NaCl concentration upon the development of surface charge on the substrates. A method is described for separating surface-bound goethite from free goethite. Sediments in natural waters can serve as sources or sinks of water-borne chemicals by virtue of adsorption and desorption processes occurring at their surfaces.l? The nature of sediment surfaces may change over time as the sediments acquire adsorbed organic and inorganic materials through interactions with dissolved and colloidal chemical species.The adsorbed species themselves then become important factors influencing further sorption processes and hence the chemical composition of the waters in which they reside. Among the most common sorbed species are the Fe1I1 hydrous oxides. Jenne2 has proposed that sorption to the hydrous oxides of iron and manganese is the principal control of heavy-metal mobility in soils and fresh waters. Crystalline iron(m) oxides such as goethite (a-FeOOH), lepidocrocite v-FeOOH) and haematite (Fe20,) are common constituents of the sediments, occurring as free particles and as sorbed layers on other particles such as clays and quartz.2v3 A gen- erally accepted mechanism for forming sorbed layers of crystalline Fe1I1 is one where ferric ion species [Fe3+, FeOH2+ and Fe(OH)t] and/or amorphous Fe(OH), are adsorbed onto the surface of a sedimentary mineral and then become converted into crystalline ferric layers over a period of time.Follett4 and F ~ r d h a m ~ - ~ have studied this process in detail for substrates of kaolinite, quartz and gibbsite. Follett found that amorphous Fe(OH), in colloidal suspension adsorbs specifically to all three substrates and that its adherence is unaffected by pH changes, prolonged washing, ultrasonic vibration or exposure to high concentrations of NHZ, Ca2+ or A13+. Fordham ex- amined the adsorption of ferric ion species from dilute acidic solutions onto kaolinite. He followed the change of ferric ion species adsorbed on the surface into amorphous Fe(OH), and its subsequent conversion into crystalline lepidocrocite.Although he observed no goethite or haematite after 15 weeks of ageing, these minerals, which are more stable thermodynamically than 1epidocrocitelO and are abundant in the natural environment, presumably could be formed eventually under appropriate conditions. t Permanent address: Department of Chemistry, University of Denver, Denver, Colorado 80208, U.S.A. $ Present address: Carborundum Co., Niagara Falls, New York 14302, U S A . 14911492 ADSORPTION OF GOETHITE ONTO QUARTZ AND KAOLINITE There seem to be no previously published reports concerning the direct adsorption of colloidal crystalline goethite onto sediments as a first step in developing crystalline coatings. The purpose of our study was to determine whether goethite coatings on sediments required precursors of ferric ion species or amorphous Fe(OH),, or whether they might also form by direct sorption of free crystalline goethite particles.The substrates studied were kaolinite and quartz. EXPERIMENTAL Goethite was prepared by the method of Atkinson et aZ.ll using an OH/Fe molar ratio of unity. Quartz, produced by the Ottawa Silica Co. as Flint Shot, was washed with 0.1 mol dm-3 HCl followed by 3 or 4 distilled-water washes to remove adsorbed ions. A well crystallized kaolinite was used as received from the Georgia Kaolinite Co. The point of zero charge (P.z.c.) for goethite was determined by the pH-drift method of Berube and de Bruyn12 to be at pH 7.2. Points of zero charge for quartz and kaolinite were taken from the literature.The P.Z.C. for quartz appears to be in the region of pH 2-313* l4 and that of kaolinite is reported to be somewhere between pH 4.5 and 8.2, depending on the ionic strength and method of meas~rement.~~’ l6 The uncertainties in P.Z.C. values for quartz and kaolinite are discussed further in the Results and Discussion section. A typical experiment consisted of stirring 0.1 g of kaolinite or quartz (particle-size range between 10 and 20 pm) with a distilled-water suspension of goethite (0.2 g per 250 cm3, particle- size range < 1 pm). The pH was adjusted with either 0.01 mol dm-3 HCI or 0.01 mol dm-3 NaOH. NaCl was added to adjust the ionic strength. After stirring for 19 h the final pH was recorded and the solids allowed to settle. During stirring, the pH change was always toward neutrality and the variation was < 1 pH unit.The final pH after stirring was taken as the pH at which sorption occurred. Sorption of goethite onto the substrate produced larger composite particles which were visually distinguishable from the untreated substrate. Electron micrographs of goethite-sorbed particles showed goethite platelets adhering to substrate particles. Goethite-sorbed substrate was separated from free goethite by exploiting their different settling-out times after stirring ceased. Goethite was formed as a distilled-water suspension with an average particle size < 1 pm. A portion of the suspension was tested for the presence of amorphous ferric hydroxide by treatment with acid ammonium oxalate.lS The amount of amorphous hydroxide was < 1 % .Quartz and kaolinite were sieved to obtain a particle-size distribution between 10 and 20 pm. Free goethite remained in suspension after 2 days, whereas the quartz and kaolinite particles settled out in CQ. 30 min (fall distance 10 cm). When a run ended and stirring stopped, the larger goethite-sorbed substrate particles settled out in ca. 15 min. The unsorbed goethite remaining in suspension was removed by siphoning off the suspension. In experiments where varying the solution ionic strength might have caused flocculation of goethite, the salt concentrations were adjusted in the goethite suspensions before adding the substrate material. These suspensions were stirred for 19 h to complete the flocculation process and all the flocculated goethite was allowed to settle out.Further flocculation did not occur. The goethte remaining in suspension was siphoned off and used in subsequent experiments. At the end of the sorption experiment the separated goethite-sorbed substrate particles were repeatedly washed with distilled water adjusted to an ionic strength matching that of the experiment, in order to remove any non-bound goethite. Goethite adsorption to kaolinite occurred at all pH values between 2 and 7. Wash solutions between pH 2 and 10 had no effect on the adherence of goethite to kaolinite, regardless of the pH at which the adsorption occurred. Goethite adsorption onto quartz was detected only in suspensions between pH 5 and 8.5. Wash solutions with pH values outside this range removed goethite from the substrate.After washing, the goethite-sorbed substrate particles were air-dried and examined by infrared spectroscopy to determine the quantity of goethite adsorbed. Infrared spectra were obtained from 1.0 and 0.1 % KBr pellets on a Digilab FTS-20B IR Fourier-transform spectrometer, interfaced with a Nova 2 computer. Computer analysis of the recorded spectra produced relative measurements of the amount of goethite adsorbed onto quartz and kaoliniteM. C. GOLDBERG, E. R. WINER AND P. M. BOYMEL 1493 I I I I I I I I 0 3600 3200 2800 2400 2000 I600 1200 800 wavenum berlcm -' 0 Fig. 1. Infrared spectra from the goethite-kaolinite system: (a) standard spectrum of pure goethite, from one of a series of standardized samples, (b) pure kaolinite, (c) goethite-sorbed kaolinite and ( d ) spectrum of surface-bound goethite resulting from subtraction of (b) from (c).under different conditions of pH and salt concentration. Because the particle-size distribution can influence the i.r. spectral pattern," care was taken to ensure that all calibration and experimental samples had the same particle-size distribution. Because the quartz spectrum was especially sensitive to particle-size distribution, we restricted the quartz particle-size range to < 20pm. Fig. 1 and 2 illustrate the spectral subtraction procedure used to determine the amount of goethite adsorbed to the substrate. Spectra were obtained from pure substrate, goethite-sorbed substrate and standardized samples of pure goethite. The pure-substrate spectrum was computer-subtracted from the spectrum of goethite-sorbed substrate, leaving a sample goethite spectrum.A standard-goethite spectrum was then compared by subtraction with the sample- goethite spectrum to obtain a quantitative measure of the amount of goethite absorbed. The characteristics of the pure-goethite spectrum [fig. l(a)], a broad band centred at 3 160 cm-l [v(OH) from hydroxy groups hydrogen-bonded to adsorbed water molecule^^^] and1494 ADSORPTION OF GOETHITE ONTO QUARTZ AND KAOLINITE 4000 3600 3200 2800 2400 2000 1600 1200 800 400 Fig. 2. Infrared spectra from the goethite-quartz system : (a) pure quartz, (b) goethite-sorbed quartz, ( c ) spectrum of surface-bound goethite resulting from subtraction of (a) from (b). Spectrum (c) is compared by subtraction with a pure-goethite standard spectrum, such as fig.1 (a), to obtain a quantitative measure of surface-bound goethite. wavenum berlcm-' sharp peaks at 900, 800 and 640 cm-I, make goethite easy to recognize in the presence of other spectral features. The absence of the v(0H) bands at 3660 and 3486cmp1, which are characteristic of dry goethite,'*. l9 indicates that the goethite in the measured samples was covered with adsorbed water. With a kaolinite substrate the sharp kaolinite peaks between 3600 and 3750 cm-' [fig. 1 (b)] can serve as a convenient internal reference for quantitative subtraction of kaolinite from goethite-sorbed kaolinite [fig. 1 ( c ) ] . After subtracting kaolinite, the height of the remaining goethite band at 3160 cm-' [fig l(d)], which is relatively free from overlap by kaolinite spectral features but is close in frequency to the kaolinite reference peaks, was compared with a standard-goethite spectra to measure the amount of goethite adsorbed.Reproducibility was satisfactory only if the goethite band closest to the kaolinite reference peaks was used for the measurement, possibly because the influence of particle-size distribution upon band shape, mentioned earlier, is wavelength-dependent. The subtraction procedure is more difficult with a quartz substrate. The broad quartz absorption band around 3460 cm-l overlaps the goethite band at 3 160 cm-l, and structure in the quartz spectrum below 900 cm-I interferes with the goethite bands in that region (see fig. 2). The most satisfactory spectral subtraction procedure proved to be one where the quartz band at 3460 cm-l was subtracted to leave no trace on the side of the goethite band at 3 160 cm-l.The height of the remaining goethite band at 3160 cm-l [fig. 2(c)] was then compared with standard-goethite spectra, as in the goethite-kaolinite system.I 50 60 70 80 1495 Fig. 3. Adsorption isotherm of colloidal goethite onto quartz. There is no measurable dependence on NaCl concentration. RESULTS AND DISCUSSION GOETHITE-QUARTZ The adsorption isotherm of goethite particles onto quartz is shown in fig. 3. Adsorption was observed only between pH 5 and 8.5, with a maximum at pH 6.7. Goethite adsorption onto quartz is expected at pH values between their respective P.Z.C. values of pH 7.2 (goethite) and 2 or 3 (quartz), where goethite and quartz have surface charges of opposite sign. However, the negative surface charge density on quartz develops slowly at first, making the P.Z.C.difficult to determine a~curate1y.l~ It remains quite small (< 10 pC cm-2) until ca. pH 5 or 6, above which the charge development is ac~e1erated.l~ Our spectroscopic measurement technique detected no adsorbed goethite until this pH region was reached. Increasing the ionic strength by adding NaCl had no detectable influence at any pH value. Although increasing the NaCl concentration is known to increase the quartz surface charge density, the magnitude of the increase is sma1114 and any influence on goethite adsorption was below our limit of sensitivity. Washing goethite-coated quartz removed goethite if the wash solution pH was below 5 or above 8.5.The behaviour of the goethite-quartz system is consistent with a non-specific adsorption mechanism and indicates that purely electrostatic attractions between pH-dependent charge sites are responsible for goethite adsorption onto quartz. GOETHITE-KAOLINITE Adsorption isotherms of goethite onto kaolinite are shown in fig. 4. The dependence upon NaCl concentration and pH is very different from the goethite-quartz system and suggests a different surface-adsorption mechanism. For pure kaolinite the1496 0 l 0 - n M k x 008- E 0 .- W .- u L Y N 0.06 5 m \ u .4 3 004- 0 M 002 ADSORPTION OF GOETHITE ONTO QUARTZ AND KAOLINITE - - 0 12 nQ 2.0 3.0 4.0 5.0 ' 60 7.0 PH Fig. 4. Adsorption isotherm of colloidal goethite onto kaolinite, for several NaCl concentrations: 0, 0.001; A, 0.01 and 0, 0.1 mol dm-3.dependence of surface charge upon pH is more complicated than for quartz. Kaolinite plates are believed to carry a permanent small negative charge on their basal planes, owing to isomorphous substitution, and a pH-dependent charge on the plate edges that has a P.Z.C. between pH 4.5 and 8.2, depending on the ionic strength and the measuring technique.14*15 The edge charge is positive at pH values below the P.Z.C. owing to adsorption of hydronium ions to the edges, and is negative above the P.z.c., owing to adsorption of hydroxyl ions. The overall negative charge density is small at low pH, owing to occupation of negative charge sites by H+, but it increases rapidly as the pH is increased14 and appears to go through a maximum around pH 10.5.15 The adsorption of cations, such as Fe3+, FeOH2+ and Fe(OH)i, onto kaolinite will occur only on the basal faces in very acidic solutions but will extend to the plate edges at pH values above the P.Z.C.of the kaolinite edge charge. Colloidal particles such as amorphous Fe(OH), will also be absorbed on the negative kaolinite basal faces at pH values below the P.Z.C. for the colloid, where the colloid particles carry a net positive charge. In the pH region above the P.Z.C. values for both the colloid and kaolinite, adsorption of colloid is negligible because both particles carry net negative charges. The surface charge on goethite is positive at pH values below its P.Z.C.of 7.2, and adsorption onto kaolinite is expected in this region. We were unable to obtain consistent measurements above pH 7. The data suggested that coagulation of the goethite colloid might have been the reason, causing larger particles of pure goethite to settle and be collected along with goethite-sorbed substrate particles. Fig. 4 shows that adsorption onto kaolinite occurs at pH 2 and increases with pH for low NaCl concentrations. It seems likely that the concentration of positive geothite particles is greatest adjacent to the negative basal planes of kaolinite and small nearM. C. GOLDBERG, E. R. WEINER AND P. M. BOYMEL 1497 the positively charged plate edges, resulting in preferential adsorption onto the basal planes.* As the pH increases, the net negative charge density also increases as the amount of surface-bound H+ decreases and more goethite is adsorbed.However, the fact that goethite, which was adsorbed at higher pH values, is not removed when the pH is lowered suggests a bonding mechanism more complex than in the goethite-quartz system. We observed no changes in the goethite spectrum that might be attributed to the adsorption process, although small changes in the broad goethite band would be hard to detect. Nor did we find any new bands from goethite-surface bonds. Although these results are inconclusive, they at least give no support to a specific bonding mechanism. Follett4 reported similar adsorption behaviour for the adsorption of amorphous Fe(OH), onto kaolinite. He proposed an adsorption mechanism whereby positively charged colloid particles are attracted to the permanently negative surface of kaolinite, increasing the net positive charge on the kaolinite.Changes in pH can increase or decrease the additional charge density carried by the colloid, but have no influence on the portion of the colloid neutralized by being in contact with permanent negative sites on the kaolinite surface. Thus, pH changes do not release colloid particles back into suspension. Our observations of goethite adsorption onto kaolinite in acidic suspensions are consistent with Follett’s mechanism. Fig. 4 also shows that adsorption of goethite increases at any pH < 7 if the NaCl concentration is increased, with saturation occurring around 0.10 mol dmP3 NaC1. This behaviour is expected because the negative surface charge density of kaolinite increases rapidly with increasing NaCl concentration,14 and this should increase goethite adsorption.Another possible mechanism, which also may play a role, is easier adsorption of goethite onto Na+-exchanged kaolinite surfaces than onto H+-exchanged surfaces. Na+ is an easily exchangeable cation, while H+ is much more difficult to displace.20 At low pH and low Na+ concentrations, the kaolinite negative surface sites are occupied by tightly bound H+. Increasing the pH or increasing the Na+ concentration allows Na+ to compete more effectively for surface positions. Eventually, at high enough pH and/or Na+ concentration, most of the H+ is displaced from the negative surface sites. The fact that the extent of goethite adsorption follows the displacement of H+ by Na+ could indicate that goethite absorbs more readily onto a kaolinite surface exchanged with Na+ than onto one exchanged with H+.A similar increased affinity for adsorption of Cu2+ onto Na+-exchanged kaolinite surfaces has been observed by McBride.21 CONCLUSIONS Colloidal goethite can adsorb onto the surface of both quartz and kaolinite. The pH dependence of adsorption onto quartz indicates a purely coulombic mechanism governed by the pH dependence of the goethite and quartz surface charges. The pH dependence of goethite adsorption onto kaolinite indicates a different binding mechanism. A possible mechanism is based on the fact that in acid solutions the kaolinite has permanent negative sites. Adsorption of positively charged goethite onto these sites results in a binding force, owing to effective charge neutralization in the region where the particles are in contact, which is not pH-dependent.The influence of NaCl concentration on goethite adsorption was found to follow closely its influence on the substrate surface charge concentration, both adsorption and charge concen- tration increasing with NaCl concentration. The data also allow the possibility that goethite displaces Na+ more readily than H+ from a kaolinite surface.1498 ADSORPTION OF GOETHITE ONTO QUARTZ AND KAOLINITE J. D. Hem, U.S. Geological Survey Water Supply Paper 1473 (U.S. Government Printing Office, Washington, D.C., 1970). E. A. Jenne, in Molybdenum in the Environment, ed. W. R. Chappell and K. K. Peterson (Marcel Dekker, New York, 1977), vol. 2, pp. 425-553. U. Schwertmann and R. M. Taylor, in Minerals in Soil Environments, ed. J. B. Dixon and S. B. Weed (Soil Science Society of America, Madison, Wisconsin, 1977), pp. 145-176. E. A. C. Follett, J. Soil Sci., 1965, 16, 333. A. W. Fordham, J. Soil Res., 1969, 7, 185. A. W. Fordham, Aust. J. Soil Res., 1969, 7, 199. A. W. Fordham, Aust. J. Soil Res., 1970, 8, 107. A. W. Fordham, Aust. J. Soil Res., 1973, 11, 185. A. W. Fordham, Ausr. J. Soil Res., 1973, 11, 197. lo W. L. Lindsay, Chemical Equilibria in Soils (Wiley, New York, 1979), p. 133. l 1 R. J. Atkinson, A. M. Posner and J. P. Quirk, J. Inorg. Nucl. Chem., 1968, 30, 2371. l2 Y. G. Berube and P. L. DeBruyn, J. Colloid Interface Sci., 1968, 27, 314. l3 R. 0. James and T. W. Healy, J. Colloid Interface Sci., 1972, 40, 65. l4 A. C. Riese, Ph.D. Thesis (Colorado School of Mines, Golden, Colorado). lE, J. B. Dixon, in Minerals in Soil Environments, ed. J. B. Dixon and S. B. Weed (Soil Science Society of America; Madison, Wisconsin, 1977), pp. 385-388. E. R. Landa and R. G. Gast, Clays Clay Miner., 1973, 21, 121. l7 J. Hlavay and J. Inczedy, Acta Chem. Acad. Sci. Hung., 1979, 102, 11. IR J. D. Russell, R. L. Parfitt, A. R. Fraser and V. C. Farmer, Nature (London), 1974, 248, 220. 2o R. N. Yong and B. P. Warkentin, Soil Properties and Behaviour (Elsevier, New York, 1975). '' M. B. McBride, Clays Clay Miner., 1978, 26, 101. R. L. Parfitt, J. D. Russell, and V. C. Farmer, J. Chem. Soc., Faraday Trans. I , 1976, 72, 1082. (PAPER 3/1306)

ISSN:0300-9599    

DOI:10.1039/F19848001491

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1984,80, 1499-1506 Laser-powered Homogeneous and Heterogeneous Pyrolysis of 2-Nitropropane BY JOSEF POLA,* MARTA FARKACOVA AND PAVEL KUBAT Institute of Chemical Process Fundamentals, Czechoslovak Academy of Sciences, 165 02 Prague, Czechoslovakia AND ANTON~N TRKA Institute of Organic Chemistry and Biochemistry, Czechoslovak Academy of Sciences, 166 10 Prague, Czechoslovakia Received 22nd August, 1983 The pyrolysis of 2-nitropropane has been studied under both homogeneous conditions (using a continuous-wave CO, laser and sensitized by sulphur hexafluoride) and heterogeneous conditions (pulsed pyrolysis) at temperatures in the range 82&1210 K. A comparison of both decompositions is aided by the determination of spatial temperature profiles and the effective temperature in the laser-driven reaction by using calculations of the temperature distribution, thermocouple techniques and elimination of ethyl acetate as a 'chemical thermometer '.The main process in the laser-driven decomposition is the clean decomposition into propene and nitrous acid occurring via C-N bond fission. In heterogeneous pyrolysis additional products such as 2-nitropropene, acetonitrile and acetone are formed. The high yield of 2-nitropropene at 820 K indicates a unique route to this compound. The gas-phase thermal decomposition of 2-nitropropane as a representative of the nitroalkanes at temperatures from 513 to 1403 K has been studied intensively by static, 1-3 very-low-pressure6 and shock-wave-induced' pyrolysis techniques for both practical and theoretical reasons.839 A detailed study of the product distribution performed only over a rather narrow temperature range (523-61 0 K) revealed2 the reaction to proceed via two competitivelo processes, namely HONO molecular elimination and C-N bond fission to form propylene and nitrous acid, followed by a series of secondary processes initiated by the attack of the decomposition products of nitrous acid on the parent molecule.The secondary processes were, based on a comparison of the rates and product distributions in reactors with different surface-to-volume ratios, inferred partly to be heterogeneous, but the effect of the reaction-vessel wall upon the product distribution has not yet been studied using more sophisticated techniques.The recently developed technique of laser-powered homogeneous pyrolysis1'* l2 completely obviates the effect of hot reactor walls and ensures that the gas-phase decomposition of an organic compound can occur without heterogeneous reaction stages. In the present investigation this technique, along with surface pyrolysis in a micropulse reactor, is applied to the decomposition of 2-nitropropane in order to determine surface effects on the distribution of the decomposition products at temperatures higher than those previously chosen. 14991500 PYROLYSIS OF 2-NITROPROPANE EXPERIMENTAL A continuous-wave CO, laser with NaCl Brewster windows and grating-wavelength tuning, operating at the P(34) line of the OO"1 -+ 10'0 transition (931 cm-l) with 5-8 W output and 0.8 cm beam diameter, was used for the irradiation experiments.The unimolecular decomposition of ethyl acetate (EA) was chosen as the reference reaction. To provide identical thermal conditions for the systems to be compared, i.e. the laser-photosensitized decomposition of 2-nitropropane (2NP) and that of EA, equal quantities of the reactant (2NP or EA) and SF, sensitizer with or without a diluent (helium or argon), ensuring rapid thermalization,13 were separately irradiated using laser power with amplitude stability of & 1.5 % (quartz-stabilized laser cavity) in an identical cylindrical glass (Simax) cell of 10 cm path length and 3.6 cm inner diameter fitted with NaCl windows and one PTFE stopcock. The laser beam was focussed (Ge lens, focal length 25 cm) at the centre of the vertically positioned cell and entered from below.The laser output was measured using a Coherent model 201 power meter, and the laser line used for irradiation was verified with a model 16-A spectrum analyser (Optical Eng. Co.). The same cell filled with a 2NP+SF, mixture equipped with six or less thermocouples (jacket diameter 0.15 mm, VEB Walzwerk Hettstadt) was used for temperature measurements. The thermocouples were placed 12, 38, 63 and 88 mm ( x ) above the entrance window at different distances ( r = 0-18 mm) from the laser beam passing through the cell along its symmetry axis. The points in space at which the thermocouple measurements were made are given as circles in fig. 2 (see later). Exact adjustment of the path of the continuous-wave CO, laser beam was made by matching with a visible He-Ne laser (model HNA 50, Carl Zeiss, Jena).The cells were separately attached to the apparatus as shown in fig. 1. The samples for laser irradiation were prepared by a standard vacuum-line technique and the concentrations of the individual components were checked by measuring the i.r. spectra at 1524 (2NP), 1047 (EA) and 987 (SF,) cm-l on a Perkin-Elmer model 62 1 i.r. spectrometer. The absorptioncoefficients were ascertained by measuring the spectra of the pure samples. The progress of the reaction with gaseous mixtures irradiated at measured intervals was monitored using a sampling valve13 coupled with a Chrom 3 gas chromatograph provided with flame ionization and thermal conductivity detectors and a temperature programming facility.A thermocouple method of measuring the temperature inside the cell enabled the voltage on the individual thermocouples to be recorded at a rate of 10 channels per second by means of a digital data logger microscan (Dynamco Systems, Chertsey) with an accuracy of 0.05 mV. The results were fed onto punched paper tape and evaluated by an EC 1033 computer. The micropulse reactor used was similar to that described in ref. (15) and consisted of a 18 cm x 0.2 cm i d . electrically heated stainless-steel tube, closed at one end by a silicone-rubber septum and at the other end attached to the injection port of the gas chromatograph (Chrom-3). The carrier gas was nitrogen and its flow rate was allowed to settle to obtain a suitable residence time.The temperature profile inside the middle 12 cm of the reactor was steady to within k 3 K. The gas-chromatographic analysis of the decomposition of 2NP and EA was carried out on 25% squalane on Chromosorb (90/105 mesh; carrier gas nitrogen; separation of 2NP, 2-nitropropene, acetonitrile, acetone, propene and methane ; separation of ethyl acetate and ethylene) and on Porapak N (80/ 100 mesh; carrier gas hydrogen; separation of propene, nitrous oxides, nitric oxide and water). For the identification of the decomposition products both the comparison of their retention time with that of a standard compound and the g.c./m.s. system were used. The latter consisted of a Pye-Unicam 104 gas chromatograph equipped with a packed column (3 m x 3 mm id., 10% OVE 17 silicone elastomer on Chromosorb W-DMCS) connected to an AEI MS 902 double-focussing mass spectrometer (Associated Electric Industries, Manchester) by means of a steel capillary and a Watson-Biemann separator.The quantitative analysis was done by peak-area measurements, taking into account detector response factors. Sulphur hexafluoride (Montedison Milano, I.E.C. standard) and ethyl acetate (Lachema, Brno) were commercial products. 2-Nitropropane was obtained by treatment of 2-bromopropane with sodium nitrite, a procedure similar to the preparation of 2-nitro-octane.16 The yield of 2NP (at 65 "C) depended upon the reaction conditions (solvent, reaction time) and was 0% (dimethylformamide, 13 h), 4.5 % (dimethylsulphoxide, 4 h) and 1 1 % (hexamethylphosphotri- amide, 5 h).J.POLA, M. FARKACOVA, P. KUBAT AND A. TRKA 1501 -m I i d Fig. 1. Apparatus used for the laser-driven pyrolysis of 2NP and the measurement of temperature distribution. a, CO, laser; b, He-Ne laser; c, rectangular NaCl window; d, revolving chromium coated mirror; e, Coherent model 201 powermeter; f, reaction cell (fi) or cell for measurement of temperature (f,); g, thermocouples; h, reference cell for i.r. analysis; i, vacuum line; j, digital centre; k, computer; 1, sampling valve; m, gas chromatograph; n, i.r. spectrometer. The computer program developed for the estimation of the temperature throughout the cell in the laser experiments as a function of the cell radius (r) and length (x) was similar to that described in ref. (1 7), where a steady-state temperature distribution, insignificant convection, thermal conductivity being dependent on temperature, and monotonic dependence of laser- beam absorption on pathlength were assumed.In our program the absorption of the laser beam was considered to be controlled for the relatively high concentration of SF, used in our experiments by the Beer-Lambert law. Moreover, except for argon and helium,18 the thermal conductivity was assumed to be independent of temperature (the value at 300 K for SF, diluent systems was equal to that of pure diluent). RESULTS AND DISCUSSION Chemical reactions induced with a C.W. CO, laser are generally assumed to proceed far from the reactor walls and to have a non-uniform temperature distribution. In reviewing the previously described methods for evaluation of thermal and/or kinetic parameters, three approaches must be considered.Convection being unimportant, the rate constant and translational temperature at a given position of a reactor can be obtained1'* l9 through a series of special procedures based on measurements of laser-induced reaction rate and optical absorption which include interferometric measurements of temperature. The temperature can also be measured directly using probes as thermocouples,20 but this may introduce errors in regions of large temperature gradients and inside the beam region. Provided that convection times are rapid relative to the reaction rate a practical procedure can be applied, either treating the reacting system as though it were characterised by a mean effective temperature throughout the reactor volumell or assuming non-isothermal conditions in the form of a parabolic temperature distribution21 in a limited reactor space (the so-called ' hot-zone').In the present investigation the temperature distribution was measured by means of a thermocouple technique. The vertically positioned tube-like cell reactor was1502 PYROLYSIS OF 2-NITROPROPANE 1600 rlmm 0 18 Fig. 2. Temperature distribution inside the cell (r = 18 mm, x = 100 mm) after 3 s irradiation of the 2NP (1.7 kPa)+SF, (0.7 kPa) mixture with the focussed laser beam [P(34) line, 6 W output] . irradiated at its centre from the lower window; in such a case the temperature distribution was found to be symmetrical along the laser beam. The temperature profiles acquired for the 2NP + SF, mixture (fig.2) irradiated with the laser beam at 6 W output after 3 s exposition (a time at which 2NP decomposition had already progressed to a pronounced conversion and a steady temperature state had been achieved) are extrapolated (fig. 2, dashed curve) by consulting the computed temperature profiles for laser-irradiated sulphur hexafluoride and its mixtures with argon or helium (fig. 3). The difference in character between the experimental (fig. 2) and calculated (fig. 3) temperature distributions supports the importance of convection currents neglected in the calculation. Maximum temperatures in the SF,, SF, + Ar and SF, + He systems computed for thermal conductivity independent of temperature decrease in the above order, taking the values 4407, 2578 and 547 K.Lower values of T,,, for SF, + Ar and SF, + He were computed when the thermal conductivity was assumed to be dependent on temperature, and are 139 1 and 520 K, respectively. These latter values are certainly more realistic, T,,, for the SF, + Ar system being reasonably higher than the mean effective temperature for this system obtained from kinetic data on the decomposition of 2NP and EA (table 1). The laser-photosensitized decomposition of 2NP yields propene, nitrogen oxides, water and methane and is strongly influenced by the cleanliness of the cell windows and the configuration (vertical or horizontal) of the irradiated cell. Contamination of the entrance window by dust or dirt, and to a lesser extent a horizontal cell configuration, facilitates the formation of additional compounds, acetone and acetonitrile.Their yield diminished upon careful cleaning of the cell windows and did not increase even when the laser output was increased from 5 to 8 W. By analogy with other examples of laser-induced surface-catalysed reaction^^^-^^ the production of acetone and acetonitrile can be accounted for by the direct heating of the cell window and by convection currents bringing hot molecules onto the cell wall. Under conditions which do not permit surface chemistry the progress of the reaction and mass balance performed by gas chromatography coupled with sampling via a valveJ. POLA, M. FARKACOVA, P. KUBAT AND A. TRKA A 1503 t 3000- 2 2000- 1000 - (b 1 (c ) Fig. 3. Temperature distribution for (a) SF, (1.3 kPa) (T,,, = 4407 K), (b) SF, (1.3 kPa) + argon, (T,,, = 2578 K) and (c) SF, (1.3 kPa)+ helium (T,,, = 547 K) computed for the cell with Y = 18 mm and x = 100 mm, wall temperature 300 K, irradiance 0.3 W mm-2, beam radius 1 mm.The thermal conductivity was taken to be independent of temperature and the excess of diluent was sufficient to consider the thermal conductivity of the mixtures to be equal to that of the pure diluent. Table 1. Mean effective temperature, ( T ) , estimated for different modes of 2NP decomposition 2NP decomposition <T)a/K E reactant log /kcal reactant + SF, mode ( A s-l) mol-l T / K ref. +SF, + Ar partly heterogeneous 11.3 40 523-1100 2,22 910 1000 reactionb elimination of HONO" five-centre molecular 11.5 42 ca. < 700 10 760 840 C-N bond fissionC 17.5 62 ca.> 700 10 710 680 a Values of ( T ) = (EEA-E~NP)/(~.~R [(logk,,p/log kEA) -log (A~NP/A EA)]} where log (Als-') = 12.6 and E/kcal mol-l = 48 for EA decomposition via the six-centre molecular elimination of CH,CO,H obtained in the temperature range 725-883 K were taken from ref. (22). Experimental. Calculated using thermochemistry and estimated activation energy for the reverse reaction. and based on the determination of absolute amounts of 2NP, propene and methane revealed that 2NP, yielding nitrogen oxides and water, completely decomposes into propene, which in turn partly (3-6 % ) undergoes decomposition to yield methane. Water, not detectable by g.c., was found to be absorbed on NaCl windows by i.r. spectroscopy as a weak absorption band at 1640 and 3400 cm-1.26 The laser-induced decomposition apparently occurs in a cylindrical volume with an axis identical to that of the reactor and positioned between the entrance window and x = 12 mm, its diameter being less than 8 mm.This assumption is consistent with the observation of almost 70 % absorption of the laser-beam energy after passage through1504 Fig. 2NP (A). PYROLYSIS OF 2-NITROPROPANE f 4. Dependence of the 2NP and EA decomposition progress upon time for the systems +SF, (O), 2NP+SF,+Ar (O), 2NP+SF,+He (A), EAfSF, ( x ) and EA+SF,+Ar The curves relate to a laser output of 6 W and reactant (1.3 kPa) + SF, (2.1 kPa) or reactant (1.3 kPa) + SF, (2.1 kPa) + diluent (36.6 kPa) mixtures. Table 2. Heterogeneous pulsed pyrolysisa of 2NP productsb T/K CH,C(O)CH, CH,CHO CH,CN CH,C(NO),=CH, - - 695 0.80 820 0.59 850 0.14 0.05 0.23 870 0.05 0.03 0.09 910 0.03 0.03 0.04 930 0.03 0.05 0.02 1010 0.0 1 0.03 0.02 1030 0.03 0.06 0.02 - - - 2.3 0.96 0.58 0.10 0.02 - a Residence time 0.47 s.* Molar amount relative to propene. the shorter (r = 18 mm, x = 14 mm) reactor filled with 0.5 kPa of SF, and 39.4 kPa of helium. The dependence of the decomposition of 2NP on time is shown in fig. 4, together with that of the reference reaction, decomposition of EA. The rate of both decompositions decreases upon addition of diluent gas (argon or helium). The greater decrease in rate observed with helium (in its presence EA does not decompose at all) is consistent with its higher thermal conductivity. Relative (mean) rate constants k,,,: kEA calculated for reactant + SF, (4.2) and reactant + SF, + Ar (2.8) systems allow one to estimate the mean effective temperature, ( T ) ; its value calculated for three different mechanistic paths to 2NP decomposition, namely, partly heterogeneous, molecular and homolytic bond-fission, are gathered in table 1.The ArrheniusJ. POLA, M. FARKACOVA, P. KUBAT AND A. TRKA 1505 parameters of the first two reactions yield values of ( T ) for the system reactant + SF, which are lower than for the system reactant + SF, +argon. In contrast, and in accord with the calculated temperature distributions (fig. 3) and reaction rates (fig. 4) for the systems in question, the higher ( T ) value for the system without argon is obtained by considering Arrhenius parameters for the C-N bond-fission mechanism.It thus appears that from the two possible (and competinglO) mechanisms of 2NP decomposition operating in the photosensitized experiments, i.e. the molecular and C-N bond-fission mechanisms, the latter is more important. In fact, the true temperatures in irradiated systems have to be higher than the effective temperature, and the occurrence of the C-N bond-fission mechanism in the photosensitized reaction is in agreement with the findinglo that C-N bond fission is favoured over molecular decomposition at temperatures > 700 K. Such an inference, taken together with the product analysis, supports a reaction scheme different from that valid for surface decomposition : CH, * CH(N0,) - CH, + [CH,-CH-CH,] CH, - CH=CH, + HNO, ---+--- NO, 2HN0, -+ H,O + NO +NO,.The surface reaction at temperatures < 700 K proceeds, via the molecular elimin- ation of HONO, followed by a series of reactions between 2NP and the decomposition products of nitrous acid, yielding acetone, acetonitrile and formaldehyde. At higher temperatures the decomposition of HONO was implieds to take place either by the bimolecular route or by a reaction with the wall. The product distribution of surface pyrolysis of 2NP at temperatures > 700 K is given in table 2. Noticeable yields of acetonitrile and acetone confirm the assumption of their surface formation during the photosensitized reaction. The most striking feature emerging from the comparison of the laser-powered homogeneous pyrolysis and the heterogeneous pulsed pyrolysis is the surface-facilitated formation of 2-nitropropene : surface CH,-CH-CH, + CH,-C=CH, I -Hz I NO2 NO2 This reaction obviously proceeds via dehydrogenation and should merit further attention since this non-polar type of reaction is favoured on a metallic surface over the competing modeslo (molecular and radical) of the elimination of HNO,.The yield of 2-nitropropene obtained in the pyrolysis at 820 K is 2.3 times higher than that of propene. The surface pyrolysis of 2NP may thus prove to be more efficient for preparative purposes than thermal methods in the production of propene. C. Frejacques, C.R. Acad. Sci., 1950, 231, 1061. V. V. Dubinin, G. M. Nazin and G. B. Manelis, Izv. Akad. Nauk SSSR, Ser. Khim., 1971, 1412. K. A.Wilde, Ind. Eng. Chem., 1956, 48, 769. P. Gray, A. D. Yoffe and L. Roselaar, Trans. Furaday Soc., 1955, 51, 1489. G. N. Spokes and S. W. Benson, J. Am. Chem. Soc., 1967,89,6030. ' K. Glanzer and J. Troe, Ber. Bunsenges. Phys. Chem., 1974, 78, 121. Kirk-Othmer Encyclopedia of Chemical Technology (Interscience, New York, 2nd edn, 1967), vol. 13, p. 881. * T. E. Smith and J. G. Calvert, J. Phys. Chem., 1959,63, 1305.1506 PYROLYSIS OF 2-NITROPROPANE G. N. Nazin, G. B. Manelis and F. 1. Dubovickij, Usp. Khim., 1968, 37, 1443. lo R. Shaw, Znt. J. Chem. Kinet., 1973, 5, 261. l1 W. M. Shaub and S. H. Bauer, Znt. J. Chem. Kinet., 1975, 7, 509. l2 J. Tardieu de Maleissye, F. Lempereur and C. Marsal, C.R. Acad. Sci., 1972, 275, 11 53. l3 J. T. Knudson and G. W. Flynn, J. Chem. Phys., 1973,58, 1467. l4 K. Klusakk, P. Schneider and V. Masak, J. Chromatogr., 1982, 244, 125. l5 S. Daren, M. Levy and D. Vofsi, Br. Polym. J., 1975, 7, 247. l6 N. Kornblum and J. W. Powers, J. Org. Chem., 1957, 22, 455. l7 R. N. Zitter, D. F. Coster, A. Cantoni and J. Pleil, Chem. Phys., 1980, 46, 107. I s Y. S. Toulokian, P. E. Liley and S. C. Saxena, Thermophysical Properties of Matter (Plenum Press, l9 R. N. Zitter, D. F. Coster, A. Cantoni and R. Ringwelski, Chem. Phys., 1981, 57, 11. 2o V. N. Sidelnikov, A. K. Petrov, N. N. Rubcova, Yu. N. Samsonov and Yu. N. Molin, Zzu. Sib. Otd. 21 Yu. N. Samsonov, A. K. Petrov, A. V. Baklanov and V. V. Vizhin, React. Kinet. Catal. Lett., 1976, 22 S . W. Benson and H. E. O’Neal, Kinetic Data on Gas Phase Unimolecular Reactions (National 23 Z. Karny and R. N. Zare, Chem. Phys., 1977,23, 321. 24 W. D. Farneth, P. G. Zimmerman, D. J. Hogenkamp and S. D. Kennedy, J. Am. Chem. SOC., 1983, 25 G. Salvetat and M. Bourene, Chem. Phys. Lett., 1980, 12, 348. 26 V. C. Farmer and J. D. Russell, Spectrochim. Acta, 1962, 18, 461. 27 W. D. Emmons, W. N. Cannon, J. W. Dawson and R. M. Rose, J. Am. Chem. Soc., 1953,75, 1993. 28 R. L. Abbott, U.S. Patent, 1966, 3, 255, 263. New York, 1970), vol. 111. Akad. Nauk SSSR, Ser. Khim. Nauk, 1976, 33. 5, 197. Standard Reference Data Series, National Bureau of Standards, Washington D.C., 1970). 105, 1126. (PAPER 3/ 1474)

ISSN:0300-9599    

DOI:10.1039/F19848001499

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1984, 80, 1507-1515 Photocatalytic Hydrogen Evolution from Aqueous Hydrazine Solution over Precious-metal/Anatase Catalysts BY YOSHINAO OOSAWA National Chemical Laboratory for Industry, Higashi, Yatabe, Tsukuba, Ibaraki 305, Japan Received 23rd August, 1983 Aqueous hydrazine solution has been photocatalytically decomposed over precious- metal/anatase catalysts yielding only H,, N, and NH,. H,/N, (molar ratio of H, evolved to N, evolved) is usually ca. 1 and does not depend on the reaction conditions (pH, metal, reactant concentration). However, the value is > 1 when another hole scavenger (CH,OH) is present. N,/NH, is CQ. 0.5 and does not change in the presence of CH,OH. On the basis of these results, the reaction scheme is considered to be as follows: the electrons photoexcited into the conduction band of anatase are consumed through H, formation and the positive holes generated in the valence band are consumed through the simultaneous formation of N, and NH,. The dependence of the reaction rate on the concentration of N,H, and on the metal suggests that the rate-determining step is the formation of H, over the metal. In recent years, photocatalytic H, evolution from ~ a t e r l - ~ or organic- compound + water systems6* using precious-metal/semiconductor powders has been widely investigated from the standpoints of photo-to-chemical energy conversion assisted by light energy, reaction mechanism and so on.On the other hand, few reportssy have been published on photocatalytic H, evolution from systems with an inorganic electron donor other than water, although in some systems high photo- catalytic reaction rates and some simple products are expected, making it easier to determine the products and follow the reaction.Hydrazine, which is an inorganic electron donor, can reduce protons and produce H, thermodynamically. However, it has been reported that the thermal-catalytic decomposition of an acidic aqueous solution over platinum black proceeds according to the following reaction with no evolution of H,:l0 3N,H, + N, + 4NH,. In a previous communicationS photocatalytic evolution of H, from an aqueous solution of N2H,*2HC1 in the presence of Pt/TiO, has been reported: 2N,H4 + N, + H, + 2NH,. On the basis of preliminary experiments it has been proved that this reaction has the following features: (i) the rate of reaction is fast, (ii) the reaction produces only a few products (H,, N, and NH,) and so it is easy to obtain good mass balance and charge balance, which, despite their importance, have been reported for only a few cases of photocatalytic reactions in aqueous solution, and (iii) the successive reaction of primary products is minimal, which is rather different from organic electron donors.Therefore, this reaction system is considered to be suitable for an investigation of the features of photocatalytic reactions. This paper presents the results of a study of the photocatalytic decomposition of N2H4 in aqueous solution over precious-metal/anatase catalysts. 15071508 PHOTOCATALYTIC HYDROGEN EVOLUTION EXPERIMENTAL MATERIALS N2H4*H20, N2H4.2HC1, NH,C1 and CH,OH were all reagent grade.Deionized water was used without further purification. Purest-grade anatase (TP-2) was purchased from Fuji Titan. Each precious-metal/anatase was prepared by an impregnation H, reduction method (H,, 101 kPa, 200 "C, 4 h) from anatase and an aqueous solution of the precious-metal chloride (except for OsO,). X-ray powder diffraction measurements (Rigaku, Geigerflux, Cu Ka) confirmed that the crystal structure was not changed by loading of the precious metals. Average particle diameters of the anatase and the Pt/anatase (lop2 wt/wt) were measured using a centrifugal automatic particle analyser (Horiba, CAPA-500) and were both 0.26 pm. Their specific surface areas were measured by the B.E.T.method (Shimadzu surface-area analyser 2205) with Ar as an adsorbate and were 18.3 m2 g-l and 17.8 m2 g-l, respectively. APPARATUS AND PROCEDURE The reaction was performed in a photocell (108 cm3) with a rectangular-parallelepiped lower part (35 x 35 x 60 mm) and a septum. Usually, the reaction mixture, consisting of precious- metal/anatase (10 mg) and an aqueous solution of reactant (30 cm3 of 0.1 mmol drn-,) in an argon atmosphere, was stirred and irradiated using a 500 W ultra-high-pressure mercury lamp (Ushio) at ca. 40 "C in the stationary state. The dark reaction was performed at 50 "C. ANALYSIS Gaseous products (H, and N,) were analysed quantitatively by gas chromatography (Yanaco G 180; MS-13X column). The liquid reaction mixture was filtered to remove solid photocatalyst, which was washed with HCl(l0 cm3 of 1 mol dm-3) in order to desorb reactant and products adsorbed on it.The filtrates and the washings were combined together and analysed." N,H, was analysed quantitatively by titration with KIO, so1ution.12 Concentrated NaOH was then added to the mixture and the resulting solution was distilled. NH, was analysed qualitatively by the indophenol method', and quantitatively by acid-base titration', of the distillate. U.v.-visible spectrophotometric analysis (Shimadzu U.V. 240) ascertained that no reaction mixture except for anatase had absorption in the wavelength region transmitted through the Pyrex glass ( A > ca. 300 nm). RESULTS REACTION OVER Pt/ANATASE Fig. 1 shows the time dependence of the volume of gases evolved when the aqueous solution of N2H4 or N,H4*2HC1 was irradiated in the presence of Pt/anatase wt/wt). The gas evolution rate was higher when N2H4 was used as a reactant than when N,H4*2HC1 was used.The volume of the gases is expressed in cm3 at s.t.p. (0 "C, 101 kPa) unless otherwise stated. No reaction occurred when a 420 nm cut-off filter was used or in the dark. When water (30 cm3) was irradiated in the presence of the Pt/anatase (lo-, wt/wt), a small amount of H, was evolved: 0.17 cm3 after a 19 h reaction. So H, evolution in the absence of N,H, (or N,H4*2HCl) may be due to photocatalytic decomposition of or photoassisted decomposition of water. l5 Products were analysed in a few reactions. Table 1 shows the results of the analysis. No compounds other than H,, N,, NH, and N2H4 were found by gas-chromatographic analysis with MS-13X (1 m, 50 "C) and chromosorb 103 (2.25 m, 170 "C).It is clear that the reactions presented in table 1 proceed photocatalytically, because they fulfil all of the following requisites: (i) dark reaction does not occur, (ii) the reaction stops when irradiated through a 420 nm cut-off filter (the band gap of anatase is 3.2 eV, therefore light of energy greater than that of the band gap is cut of€), (iii) mass balanceY. OOSAWA 1509 18 16 14 6 2 12 c: v) m 5 --. U 0 w v) w 10 - % I 8 Lr 0 w 5 6 9 4 2 / / i Y I I I I I I I I 1 2 3 4 5 6 7 8 time/h Fig. 1. Photocatalytic H, and N, evolution over Pt/anatase wt/wt). Reaction conditions: Pt/anatase, 10 mg; reactant, 30 cm3 of 0.1 mol dm-3 N2H4, A, H, and A, N,; reactant, 30 cm3 of 0.1 mol dm-3 N2H4.2HC1, 0, H, and 0, N,; (---) when the 420 nm cut-off filter was used.(N and H) and charge balance hold well and (iv) the turnover number (moles of N2H4 consumed/moles of photocatalyst)* is > 1. Table 2 shows the average evolution rate of H, and N, for the initial few hours over various precious-metal/anatase catalysts ( lo-, wt/wt). Neither the reaction with the cut-off filter nor the dark reaction at 50 "C proceeded. Therefore, the gas evolutions presented in table 2 are also assumed to proceed photocatalytically. In almost all cases, H,/N, was ca. 1. The gas-evolution rate decreased in the following order: Pt > Pd > Ir > Rh > 0 s > Ru. This order seems to reflect the magnitude of hydrogen overvoltage required to obtain a current density of 2 mA cm-, in 0.05 mol dm-3 H2S04 (Pt < Ir < 0 s and Pd < Rh < Ru).16 REACTION IN THE PRESENCE OF CH,OH In order to obtain information on the photocatalytic decomposition of N,H, in aque- ous solution, the reaction was performed in the presence of another electron donor, i.e.CH,OH. Neither photocatalytic H, evolution from CH,OH + H,O(v/v = 1 / 1) in the presence of anatase nor the dark reaction in the presence of Pt/anatase (lo-, wt/wt) proceeded. The results of the analysis are shown in table 1. The mass balance of nitrogen held fairly well in both cases. Note that N,/NH3 is close to 0.5 * Turnover number (moles of reactant consumed/moles of photocatalyst used) is used as a guide as to whether the reaction is stoichiometric or catalytic.1510 PHOTOCATALYTIC HYDROGEN EVOLUTION Table 1.Analysis data reactant N,H, 2HC1 N2H4. 2HC1 N2H4 * 2HCP N2Hdd + CH,OHey f + CH30He,Q reaction time/h N2H, consumed/mmol product/mmol H2 N, NH3 N reactant N product H reactant H product e- consumeda pf consumedb turnover numberc mass balance/mequiv. charge balance/mequiv. 8 0.84 0.42 0.46 0.79 1.68 1.71 3.36 3.21 1.63 1.84 6.7 6.5 0.95 0.56 0.50 0.83 1.90 1.83 3.80 3.64 1.95 2.00 7.6 15 0.66 0.47 0.31 0.58 1.32 1.20 - 5.3 ~ 15 0.23 0.68 0.14 0.24 0.46 0.52 - 1.8 Reaction conditions: 0.1 mol dm-3 N2H, or NZH,*2HC1; amount of reactant, 30 cm3; a Calculated as 2H,+ NH,. Calculated as [moles of N2H, consumed 500 W ultra-high-pressure Hg lamp (Ushio) was photocatalyst, Pt/anatase ( lo-, wt/wt), 10 mg.per rnol of Pt/anatase ( lo-, wt/wt) used]. used. Calculated as 4N2. 450 W Xe lamp (Oriel) was used. 1 0.1 rnol dmV3 CH,OH. 1 rnol dm-, CH,OH. Table 2. H, and N, evolution rates (cm3 h-l) over precious-metal/anatase catalysts pho tocatalyst H2 N2 P t / ana tase 2.06 2.07 Ir/anatase 0.80 0.8 1 Rh/ana tase 0.70 0.74 Os/anatase 0.061 0.068 Ru/anatase 0 0 anatase 0 0 Pd / an at ase 1.58 I .55 Reaction conditions: reactant, 0.1 mol dm-, N,H4*2HC1; photocatalyst, 10 mg. even in the presence of a large excess of CH,OH. On the other hand, H,/N, > 1 in the absence of CH,OH. DEPENDENCE OF THE REACTION ON N,H, CONCENTRATION The dependence of the H, evolution rate on the concentration of N,H, is shown in fig. 2. In these experiments, a small amount of the Pt/anatase (lo-, wt/wt, 2 mg) was used in order to keep the conversion and the decrease in the concentration of reactant small even in the most dilute solution.In this concentration region, the rate was higher at lower concentrations and became constant at concentrations > 0.1 mol drn-,. HJN, was 1 even at a concentration of 10 rnol dmP3 : reaction time, 14 h; H, evolved, 21.9 cm3; N, evolved, 20.8 cm3; H,/N, = 1.05.Y. OOSAWA 2.0; - I s - 1.6- 5 1 - e, Y 2 1 . 2 - 3 0 . 8 - c - .& I ? - '0 151 1 I I 0.01 0.03 0.1 0.3 1 3 10 20 concentration/mol dm-3 Fig. 2. Dependence of H2 evolution rate on NzH4 concentration. Reaction conditions: Pt/anatase wt/wt), 2 mg. Table 3. Photocatalytic decomposition of NH, 0.4 Pt Pd Rh Ir, Ru, 0 s p hotoca tal y s t /anatase /anatase /anatase /anatase reaction time/h 5 6 23 4-1 5 volume of gas/cm3 0.52 0.29 0.46 < 0.01 0.28 < 0.01 0.21 < 0.01 H2 N2 Reaction conditions: reactant 0.1 mol dm-, NH4C1; photocatalyst, 10 mg.DECOMPOSITION OF NH, NH, was formed in the photocatalytic reaction of aqueous N,H, solution over Pt/anatase as shown in table 2. NH, formed photocatalytically is also able to decompose phot~catalytically.~~ A few experiments were performed to confirm this point. NH,Cl and precious-metal/anatase ( lop2 wt/wt) were used as reactant and photocatalyst. The results are shown in the table 3. Neither the reaction with the cut-off filter nor the dark reaction proceeded. Therefore the gas evolution in table 3 is considered to proceed photocatalytically. The gas evolution rates with NH4C1 were much lower than those with N,H,.Therefore the contribution from the photocatalytic decomposition of NH, is negligibly small compared with that of N,H4 under the present reaction conditions. DEPENDENCE OF H, EVOLUTION RATE ON THE Pt LOADING RATIO Eight kinds of Pt/anatase with different Pt loading ratios were prepared and H, evolution over them was investigated. The results are presented in fig. 3. H, evolution was detected even when Pt/anatase (lo+ wt/wt) was used. The H, evolution rate increased with increasing Pt loading ratio and reached saturation above a loading ratio of lo-, wt/wt. It was observed that the H, evolution rate is almost constant over a wide range of Pt loading ratios. The Pt coverage in Pt/anatase ( 5 x wt/wt) was estimated to be ca.4%: the Pt surface area estimated from the half-width of the X-ray diffraction peak was 0.71 m2 g-l and the B.E.T. surface area was 17.5 m2 g-l. Therefore the cause of the saturation in the H, evolution rate at low Pt loading ratios1512 PHOTOCATALYTIC HYDROGEN EVOLUTION Pt/anatase (wt/wt) Fig. 3. Dependence of H, evolution rate on Pt loading ratio. Reaction conditions: Pt/anatase, 10 mg; 30 cm3 of 0.1 mol dm-3 N,H,. is not caused by a decrease in the number of photons striking the anatase surface. Note that even Pt/anatase with a very low Pt loading ratio showed high activity: e.g. Pt/anatase ( wt/wt) showed an activity of approximately one-fourth that of the Pt/anatase (lo-, wt/wt). Further study is required to understand these points.DISCUSSION REACTION SCHEME A few standard redox potentialsls for the present reaction system are presented in fig. 4, together with the bottom of the conduction band (Gb)l9 and the top of the valance band (Kb)” of anatase at pH 7. Their relative location changes little with a change of pH. Reactions (1)-(3) correspond to the three redox pairs in fig. 4. Aque- ous hydrazine solution is thermodynamically unstable and can be decomposed thermodynamically according to reaction (4) or (5) by combination of reactions (1) and (2) or reactions (1) and (3), respectively. As has been stated already, however, these reactions did not take place thermocatalytically under the reaction conditions used in the present study. N, + 5H+ + 4e- N,H; (1) 2H+ + 2e- + H, (2) N,H: + 3H+ + 2e- + 2NHi (3) N,H4 + N2+2H, AGO = -22 kcal mol-1 (4) N,H4 --+ 4N2 + QNH, AGO = - 46 kcal mol-l.( 5 ) When the anatase is irradiated by light of energy greater than that of the band gap, the electrons photoexcited into the conduction band can reduce protons producing H, [reaction (2)] or reduce N2H4 producing NH, [reaction (3)] thermodynamically. On the other hand, the positive holes generated in the valence band can oxidize N,H, producing N, [reaction (l)] thermodynamically. Following three features of the reaction can be used to elucidate the photocatalytic reaction scheme : (i) H,/N, is ca. 1 over a wide range of reaction conditions (pH, loaded metal, N2H4Y. OOSAWA 1513 -1 0 w x z ;> b. 2 1 --- 2 3 anatase Fig. 4. Energetic correlation between a few standard redox potentials,18 the bottom of the conduction band (&b)l@ and the top of the valence band (&b)'@ of anatase at pH 7.concentration), (ii) N,/NH, is ca. 0.5 and the value does not change even in the presence of another electron donor (CH,OH) and (iii) the photocatalytic decom- position does not proceed over anatase without loaded metal, which is different from the case of NH,OH.,O At first sight, feature (i) can be explained by the combination of reactions (l), (2) and (3); that is (2)+(3)-(1). However, it is difficult to explain features (ii) and (iii) using this reaction scheme. At first, it is reasonable to assume that N, is formed through the hole-consumption pathway, because N, is the only oxidation product in the present reaction system.Seemingly, there are two kinds of reduction product: H, and NH,. It may also be assumed that H, is formed through the electron-consumption pathway, because H, evolution was only possible by loading precious metals which have a low H, overpotential, as shown in table 2. On the other hand, problems arise if it is assumed that NH, is formed in the electron-consumption pathway. Feature (ii) is very significant. When CH,OH is added to the N2H4 system, some of the positive holes generated by irradiation will react with CH,OH. Then some electrons generated by irradiation will not recombine with positive holes. If NH, were formed in the electron-consumption pathway, some of the electrons not recombining with positive holes would be used in the formation of both H, and NH,.As a result, N,/NH, would be < 0.5, the value obtained in the absence of CH,OH. On the other hand, H,/NH, would be 2, the value obtained in the absence of CH,OH. However, the results shown in table 1 contradict these expectations. Therefore feature (ii) shows that NH, is not formed in the electron- consump tion pathway . Concerning feature (iii), it is useful to compare it to the case where NH,OH is a reactant.,O When NH,OH was used as the reactant, the photocatalytic reaction proceeded even over anatase without any loaded metal, yielding N,, N,O and NH,. In that case, NH, was formed in the electron-consumption pathway, because NH, is the only reduction product. If NH, could have been formed in the electron- consumption pathway for N2H, system also, the photocatalytic reaction [reactions (1) and (3)] over the anatase should have occurred.In practice, however, the photocatalytic reaction over anatase does not proceed. Therefore, feature (iii) also supports the idea that NH, formation over anatase cannot be an electron-consumption pathway in the N2H4 system. 50 FAR 11514 PHOTOCATALYTIC HYDROGEN EVOLUTION As NH, formation cannot be an electron-consumption pathway, H, formation should be the only electron-consumption pathway. Therefore, on the basis of the above discussion, it is reasonable to assume that N, and NH, are formed together in the hole-consumption pathway. It is tentatively postulated that tetrazane (N,H,) is formed as a reaction intermediate. In several reaction systems it has been assumed that positive holes are consumed through the formation of hydroxy radical and through direct reaction with reactant.,l This may also be the case in the present reaction system.Therefore the main reaction will be represented by the reaction scheme shown below. anatase ec(e1ectron) + p+(hole) hv (electron-consumption pathway) H++e-+H' 2H' + H, (over precious metal) (hole-consumption pathway) H,O+p+ -+ H++OH' N,H, + p+ + N,Hi + H+ N,H, +OH' + N2Hi + H20 2N2Hi + N,H,[ = NH,(NH),NH,] N,H, -+ N, + 2NH,. This reaction scheme can explain the three features of the reaction described above. In some cases: H2/N2 < 1. It can be explained by assuming that the following reaction scheme is a side reaction. over precious metal. 1 H' + N2H4 + N2Hj N2Hi + H' + 2NH,, In some cases, on the other hand, H,/N, > 1.It can be explained by assuming that the following reaction scheme is a side reaction or by assuming the decomposition of the NH, formed photocatalytically : N,H, + 2p+ -+ N2H; + H+ N2H; + N,H, + H+ N2H2 + 2p+ + N, + 2H'. RATE-DETERMINING STEP The dependence of the reaction rate on reactant concentration often provides useful information about the rate-determining step of the reaction. Nevertheless, the dependence of the rate of photocatalytic H, evolution on the reactant concentration is reported only for a few cases. In the case of N2H4 solution, the rate decreased with increasing concentration above 0.1 mol dm-3. It is clear that N,H4 depresses H, evolution in this concentration region in the present reaction system. This means that the N,H,-consumption step is not rate-determining.Moreover, as described already, the order of the H, evolution rate over precious-metal/anatase corresponded well to the order of magnitude of the hydrogen overvoltage of these precious metals. ThisY. OOSAWA 1515 suggests that the rate-determining step is H, formation on the surface of the precious metal. Electrolytic H, formation over precious metals is considered to be a combination of hydrogen atoms adsorbed on the surface of the precious metal. On the surface of these metals loaded on anatase, N,H, may also be adsorbed. The dependence of the H, evolution rate on the N,H, concentration depicted in fig. 2 may be understood through the inhibition effect of N,H, on the combination of hydrogen atoms on the Pt surface.1 H. Yoneyama, M. Koizumi and H. Tamura, Bull. Chem. SOC. Jpn, 1979, 52, 3449. J-M. Lehn, J-P. Sauvage and R. Ziessel, Nouv. J. Chim., 1980, 4, 623. S. Sat0 and J. M. White, Chem. Phys. Lett., 1980, 72, 83. E. Borgarello, J. Kiwi, E. Pelizetti and M. Gratzel, J. Am. Chem. SOC., 1981, 103, 6324. K. Domen, S. Naito, T. Ohnishi, K. Tamaru and M. Soma, J. Phys. Chem., 1982, 86, 3657. T. Kawai and T. Sakata, Nature (London), 1980, 286,474. Y. Oosawa, J. Chem. SOC., Chem. Commun., 1982, 221. E. Borgarello, K. Kalyanasundaram and M. Gratzel, Helu. Chim. Acta, 1982, 65, 243. lo L. F. Audrieth and B. A. Ogg, The Chemistry of Hydrazine (Wiley, New York, 1951). l1 C. A. Streuli and P. R. Averell, The Analytical Chemistry of Nitrogen and Its Compounds (Wiley, New 'I P. Pichat, J-M. Herrman, J. Disdier, H. Courbon and M-N. Mozzanega, Nouu. J . Chim., 198 1,5627. York, 1970), part I. R. A. Penneman and L. F. Audrieth, Anal. Chem., 1948,20, 1058. l3 W. T. Bolleter, C. J. Bushman and P. W. Tidwell, Anal. Chem., 1961, 33, 592. l4 J. E. Devries and E. St Clair Ganz, Anal. Chem., 1953, 25, 973. l5 S. Sat0 and J. M. White, J. Phys. Chem., 1981, 85, 592. l6 S. Srinivasan and F. J. Salzano, Int. J. Hydrogen Energy, 1977,2,53. The order of H,-selectivity shown in ref. (20) (Pd > Rh > Ru and Pt > Ir > 0 s ) also corresponded well to the magnitude of their hydrogen overpotential. Encyclopedia of Electrochemistry of the Elements, ed. A. J. Bard (Marcel Dekker, New York, 1978), vol. viii. l7 Q-S. Li, K. Domen, S. Naito, T. Onishi and K. Tamaru, Chem. Lett., 1983, 321. l9 M. V. Rao, K. Rajeshwar, V. R. P. Vernekker and J. Dubow, J. Phys. Chem., 1980, 84, 1987. 2o Y. Oosawa, J. Phys. Chem., in press. When NH,OH-HCl(aq) (30 of 0.1 mol dm-3) was irradiated by a 500 W ultra-high-pressure Hg lamp over a photocatalyst [precious-metal/anatase (lop2 wt/wt), 10 mg], H,, N,, N,O and NH, were formed. Relative reaction rate (normalized to anatase) and H, selectivity [(amount of electron consumed for H, evolution)/(amount of electron consumed for the reduction)] were as follows: anatase, 1, 0; Ru/anatase, 0.55, 0.002; Os/anatase, 0.18,O; Rh/anatase, 0.44,0.17; Ir/anatase, 0.59,0.12; Pd/anatase, 1.02,0.69; Pt/anatase, 0.91,0.37. 21 K. Jones, Comprehensive Inorganic Chemistry, ed. A. F. Trotman-Dickenson et al. (Pergamon Press, Oxford, 1973), vol. 2. (PAPER 3/ 1490) 50-2

ISSN:0300-9599    

DOI:10.1039/F19848001507

出版商:RSC    

年代:1984

数据来源: RSC