摘要:

Gas Kinetics Group and Division de Chimie-Physique de la Societe Francaise de Chimie 9th International Symposium on Gas Kinetics To be held in Bordeaux, France on 20-25 July 1986 Further information from Dr R. Lasclaux, Lab. Photophys. Photochim. MolBculaire, Universite de Bordeaux I, 33405 Talence Cedex, France Poiymer Physics Group Biologically Engineered Polymers To be held at Churchill College, Cambridge on 21-23 July 1986 Further information from Dr M. J. Miles, AFRC,Food Research Institute, Colney Lane, Norwich NR4 7UA Polymer Physics Group with the British Rheological Society Deformation in Solid Polymers To be held at the University of Leeds on 9-1 1 September 1986 Further information from Dr J. V. Champion, Department of Physics, City of London Polytechnic, 31 Jewry Street, London EC3N 2EY ~~_____________ ~~~~ Carbon Group Carbon Fibres- P ro pe rt i es and A p p I i cat i o ns To be held at the University of Salford on 1 5 1 7 September 1986 Further information from The Meetings Officer, The Institute of Physics, 47 Belgrave Square, London SW1 X 8QX ~ ~~~~~~~~ ~ Division with the Surface Reactivity and Catalysis Group-Autumn Meeting Promotion in Heterogeneous Catalysis To be held at the University of Bath on 23-25 September 1986 Further information from: Professor F.S. Stone, School of Chemistry, University of Bath, Bath BA2 7AY (viii)Gas Kinetics Group and Division de Chimie-Physique de la Societe Francaise de Chimie 9th International Symposium on Gas Kinetics To be held in Bordeaux, France on 20-25 July 1986 Further information from Dr R.Lasclaux, Lab. Photophys. Photochim. MolBculaire, Universite de Bordeaux I, 33405 Talence Cedex, France Poiymer Physics Group Biologically Engineered Polymers To be held at Churchill College, Cambridge on 21-23 July 1986 Further information from Dr M. J. Miles, AFRC,Food Research Institute, Colney Lane, Norwich NR4 7UA Polymer Physics Group with the British Rheological Society Deformation in Solid Polymers To be held at the University of Leeds on 9-1 1 September 1986 Further information from Dr J. V. Champion, Department of Physics, City of London Polytechnic, 31 Jewry Street, London EC3N 2EY ~~_____________ ~~~~ Carbon Group Carbon Fibres- P ro pe rt i es and A p p I i cat i o ns To be held at the University of Salford on 1 5 1 7 September 1986 Further information from The Meetings Officer, The Institute of Physics, 47 Belgrave Square, London SW1 X 8QX ~ ~~~~~~~~ ~ Division with the Surface Reactivity and Catalysis Group-Autumn Meeting Promotion in Heterogeneous Catalysis To be held at the University of Bath on 23-25 September 1986 Further information from: Professor F. S. Stone, School of Chemistry, University of Bath, Bath BA2 7AY (viii)

ISSN:0300-9599    

DOI:10.1039/F198581FX001

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

xxxij AUTHOR INDEX Singh, Km. S., 751 Sircar, S., 1527, 1541 Slade, R. C. T., 847 Smith, I. G., 1095 Snelling, C. M., 1761 Sobczyk, L., 311 Siiderberg, D., 17 15 Solar, S., 1101 Solar, W., 1101 Soma, M., 485 Somorjai, G. A., 1263 Somsen, G., 1015 Sorek, Y., 233 Souto, F. A., 2647 Spencer, S., 2357 Spichiger-Ulmann, M., 7 13 Spoto, G., 1283 Spotswood, T. M., 1623 Srivastava, R. D., 913 Stachurski, J., 1447, 2813 Staricco, E. H., 1303 Stock, T., 2257 Stockhausen, M., 397 Stokes, R. H., 1459 Stone, F. S., 1255 Strachan, A. N, 1761 Strohbusch, F., 2021 Stuckless, J. T., 597 Su, Z., 2293 Subrahmanyam, V. S., 1655 Sugimoto, N., 1441, 2959 Suminaka, M., 2287 Suprynowicz, Z., 553 Sutcliffe, L. H., 679, 1467, 1215 Suzanne, J., 2339 Suzuki, H., 3117 Swallow, A. J., 1225 Symons, M.C. R., 433, 565, 727, 2131, 2775, 1095, 1963, 242 1 Takagi, Y., 1901 Takahashi, Y., 3 117 Takeshita, H., 2805 Tamilarasan, R., 2763 Tamura, K., 2287 Tanaka, T., 1513 Taniewska-Osinska, S., 695, Tascon, J. M. D., 939, 2399 Taylor, M. J., 1863 Taylor, N., 2357 Tejuca, L. G., 939, 2399, 1203 Teller, R. G., 1693 Tempere, J-F., 1357 Teramoto, M., 2941 Theocharis, C . R., 857 Thomas, J. K., 735 Tielen, M., 2889, 3049 Tindwa, R. M., 545 Tissier, C., 3081 Toi, K., 2835 Tokuda, T., 2835 Torrez-Mujica, T., 343 Townsend, R. P., 1071, 173 1, Trasatti, S., 2995 Treiner, C., 2513 Trenwith, A. B., 745 Trifiro, F., 1003 Troncoso, G., 1631, 1637 Tseung, A. C. C., 1883 Tuck, J. J., 833 Turner, J. E., 1263 Uemoto, M., 2333 Uma, K., 2733 Valencia, E., 1631. 1637 Valigi, M., 813 Vallmark, T., 1389 Van Oort, M.J. M., 3059 Varma, M. K., 751 Vattis, D., 2043 Vecli, A., 433 Veseli, V., 2095 Vink, H., 1677, 1725 Vliers. D. P., 2009 Vukovid, Z., 1275 3081, 1913 3127 Waghorne, W. E., 2703 Ward, A. J., 2975 Watanabe, H., 1569 Waugh, K. C., 3073 Weckstrorn, K., 2947 Weinberg, N. N., 875 Weingartner, H., 1031 Wells, C. F.. 801, 1057, 1401, White, M. A., 3059 Williams, J. O., 271 1 Williams, P. A., 2635 Williams, P. B., 3067 Williams, R. T., 847 Wojcik, D., 1037 Wood, G. L., 265 Wood, R. M., 273 Woolf, L. A., 769, 2821 Wright, C. J., 2067 Wright, J. P., 1471 Wright, T. H., 1819 Wurie, A. T., 2605 Yadav, G. D., 161 Yadava, R. D., 751 Yamaguchi, M., 1513 Yamaguti, K., 1237 Yamasaki, S., 267 Yamashita, H., 2485 Yamatera, H., 127 Yelon, W., 1693 Yoshida, S., 1513, 2485 Yoshikawa, M., 2485 Zambonin, P.G.. 621 zdanov, S. P., 2541 Zecchina, A., 1283 Zelano, V., 2365 Zhan, R. Y., 2083 Zhao, Z., 185 Zhulin, V. M., 875 Zilnyk, A., 679, 1215 Zulauf, M., 2947 Zundel, G., 1425, 2375 1985. 2145, 2475, 3091xxxij AUTHOR INDEX Singh, Km. S., 751 Sircar, S., 1527, 1541 Slade, R. C. T., 847 Smith, I. G., 1095 Snelling, C. M., 1761 Sobczyk, L., 311 Siiderberg, D., 17 15 Solar, S., 1101 Solar, W., 1101 Soma, M., 485 Somorjai, G. A., 1263 Somsen, G., 1015 Sorek, Y., 233 Souto, F. A., 2647 Spencer, S., 2357 Spichiger-Ulmann, M., 7 13 Spoto, G., 1283 Spotswood, T. M., 1623 Srivastava, R. D., 913 Stachurski, J., 1447, 2813 Staricco, E. H., 1303 Stock, T., 2257 Stockhausen, M., 397 Stokes, R. H., 1459 Stone, F. S., 1255 Strachan, A.N, 1761 Strohbusch, F., 2021 Stuckless, J. T., 597 Su, Z., 2293 Subrahmanyam, V. S., 1655 Sugimoto, N., 1441, 2959 Suminaka, M., 2287 Suprynowicz, Z., 553 Sutcliffe, L. H., 679, 1467, 1215 Suzanne, J., 2339 Suzuki, H., 3117 Swallow, A. J., 1225 Symons, M. C. R., 433, 565, 727, 2131, 2775, 1095, 1963, 242 1 Takagi, Y., 1901 Takahashi, Y., 3 117 Takeshita, H., 2805 Tamilarasan, R., 2763 Tamura, K., 2287 Tanaka, T., 1513 Taniewska-Osinska, S., 695, Tascon, J. M. D., 939, 2399 Taylor, M. J., 1863 Taylor, N., 2357 Tejuca, L. G., 939, 2399, 1203 Teller, R. G., 1693 Tempere, J-F., 1357 Teramoto, M., 2941 Theocharis, C . R., 857 Thomas, J. K., 735 Tielen, M., 2889, 3049 Tindwa, R. M., 545 Tissier, C., 3081 Toi, K., 2835 Tokuda, T., 2835 Torrez-Mujica, T., 343 Townsend, R.P., 1071, 173 1, Trasatti, S., 2995 Treiner, C., 2513 Trenwith, A. B., 745 Trifiro, F., 1003 Troncoso, G., 1631, 1637 Tseung, A. C. C., 1883 Tuck, J. J., 833 Turner, J. E., 1263 Uemoto, M., 2333 Uma, K., 2733 Valencia, E., 1631. 1637 Valigi, M., 813 Vallmark, T., 1389 Van Oort, M. J. M., 3059 Varma, M. K., 751 Vattis, D., 2043 Vecli, A., 433 Veseli, V., 2095 Vink, H., 1677, 1725 Vliers. D. P., 2009 Vukovid, Z., 1275 3081, 1913 3127 Waghorne, W. E., 2703 Ward, A. J., 2975 Watanabe, H., 1569 Waugh, K. C., 3073 Weckstrorn, K., 2947 Weinberg, N. N., 875 Weingartner, H., 1031 Wells, C. F.. 801, 1057, 1401, White, M. A., 3059 Williams, J. O., 271 1 Williams, P. A., 2635 Williams, P. B., 3067 Williams, R. T., 847 Wojcik, D., 1037 Wood, G. L., 265 Wood, R. M., 273 Woolf, L. A., 769, 2821 Wright, C. J., 2067 Wright, J. P., 1471 Wright, T. H., 1819 Wurie, A. T., 2605 Yadav, G. D., 161 Yadava, R. D., 751 Yamaguchi, M., 1513 Yamaguti, K., 1237 Yamasaki, S., 267 Yamashita, H., 2485 Yamatera, H., 127 Yelon, W., 1693 Yoshida, S., 1513, 2485 Yoshikawa, M., 2485 Zambonin, P. G.. 621 zdanov, S. P., 2541 Zecchina, A., 1283 Zelano, V., 2365 Zhan, R. Y., 2083 Zhao, Z., 185 Zhulin, V. M., 875 Zilnyk, A., 679, 1215 Zulauf, M., 2947 Zundel, G., 1425, 2375 1985. 2145, 2475, 3091

ISSN:0300-9599    

DOI:10.1039/F198581BX003

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985, 81, 11-18 Kinetics of the Reaction between Cyanide Ions and Tris(4-methyl- 1 , 10-phenanthroline)iron(n) Cations in Aqueous Solutions Analysis of Kinetic Data for this Reaction and for Solvolysis of Benzyl Chloride in Water in Terms of Isothermal, Isobaric and Related Isochoric Activation Parameters BY MICHAEL J. BLANDAMER,* JOHN BURGESS AND BARBARA CLARK Department of Chemistry, The University, Leicester LEI 7RH AND ROSS E. ROBERTSON Department of Chemistry, University of Calgary, Calgary, Alberta, Canada AND JOHN M. W. SCOTT Department of Chemistry, Memorial University of Newfoundland, St John’s, Newfoundland, Canada Received 13th January, 1984 Kinetic data for the title reaction and for the solvolysis of benzyl chloride in water are analysed to obtain isobaric and isothermal activation parameters.Isochoric parameters are defined and the rate constants described as functions of temperature, pressure and molar volume of the pure solvent, V;”. A quantity Afv/( V:) characterises the dependence of rate constant on temperature at constant V:. We counter claims that isochoric activation parameters necessarily provide an insight into the process of activation in chemical reactions. Isobaric-isothermal functions, within the limitations of transition-state theory, have a sounder basis in terms of thermodynamic analysis of chemical reactions. The dependence on temperature is reported for the rate constant of reaction between cyanide ions and tris(4-methyl- 1 ,lo-phenanthroline)iron(rr)cations in aqueous solution.The data supplement the set reporting the dependence of rate constant on pressure.l The two sets of data are analysed to obtain isobaric and isothermal activation parameters. The latter are used in the calculation of isochoric activation parameters. The significance of activation parameters calculated for isochoric conditions is examined in some detail. Numerous accounts of kinetic data refer to isochoric activation parameters. [For further references see ref. (2).] Nevertheless, their meaning has seemed, at least to us, to be obscure. Previously2 we posed two questions which in effect asked ‘what volume is held constant?’ and ‘what reference states are used in the definition of isochoric activation parameters?’. The latter question was prompted by reports of isochoric heat capacities of activation, AICF [cf.ref. (3), where the symbol ACg is used]. These questions were answered2 in part by reference to the treatment of chemical equilibria in solution. For example, the dependence of an equilibrium quotient Q(s1n : T : p ) on temperature can be characterised under the constraint that the pressure changes to hold the molar volume of the solvent, V,*, 1112 ISOBARIC, ISOTHERMAL AND ISOCHORIC PARAMETERS constant. The condition isochoric ( V,*) is therefore extrinsic to the system characterised by the equilibrium quotient Q(s1n: T : p ) . By way of contrast, the dependences of In Q on temperature at constant pressure and on pressure at constant temperature yield the intrinsic isobaric and isothermal functions (cf.the van’t Hoff equations). Here we extend the analysis to activation parameters characterising chemical reaction in solution. In addition to the well established isothermal and isobaric activation parameters, the dependence of rate constant on temperature can be characterised under the constraint that the pressure changes whereby the molar volume of the solvent remains constant. This dependence yields an isochoric ( V ; ) activation parameter where the isochoric condition is extrinsic to the reacting system. We also comment on a number of related isochoric parameters. Attention is confined to kinetic data for chemical reactions using water as the solvent. The necessary density data for water over extended ranges of temperature and pressure are a~ailable.~? However, the procedures described below can be readily extended in analyses of kinetic data for reactions when the solvent is a binary liquid Data for the dependence of the density of these mixtures on mole-fraction composition, temperature and pressure are becoming available,8* providing the basis of the calculations discussed below. At this stage we conclude that isochoric activation parameters are certainly as complicated as isothermal and isobaric activation parameters, a conclusion which is at odds with that advanced by other authors [cf.ref. (10) and (1 l)]. EXPERIMENTAL MATERIALS These were prepared and purified as described previous1y.l KINETICS The progress of the title reaction was followed spectrophotometrically.lq l2 A first-order rate constant was calculated using a Hewlett-Packard minicomputer which controlled the spectrophotometer and logged the absorbance as a function of time.12 RESULTS The reaction between cyanide ions and Fe(4Me-phen)tS, to give Fe(4Me-phen), (CN), plus free ligand, follows first-order kinetics in aqueous solution when [CN-] % [Fe(4Me-phen)f+] : - d[Fe(4Me-phen)f+]/dt = ko,,[Fe(4Me-phen)t+].Except at very high cyanide concentrations The k, term corresponds to rate-determining dissociation of the c0mp1ex.l~. l4 The k, term corresponds to bimolecular reaction between the complex and cyanide. Whether this occurs by direct attack at the iron or initially by attack at the coordinated ligand is a matter still not ~ett1ed.l~~ l6 Although in one or two special cases there is evidence for two-stage attack involving initial attack at the ligand and subsequent transfer of cyanide (or hydroxide) to the meta1,l7?l8 there is no evidence for two-stage kinetics at iron(IIEphenanthro1ine complexes except when strongly electron-withdrawing nitro or sulphonato groups are present.For the Fe(4Me-phen):+ +cyanide reaction the kinetics conform to a single-step process.M. J. BLANDAMER et al. 13 Table 1. Rate constants for the reaction between Fe(4Me-phen)i+ and cyanide in aqueous solution ([CN-] = 0.5 mol dm-3) and for Fe(4Me-phen)i' dissociation ([H,SO,] = 0.5 mol dmP3) T/K k1/10-5 s-l k,,,/10-3 s-' k2/10P3 dm3 mo1-l s-l 298 3.59 1.84 3.61 30 1 6.1 3a 2.78 5.44 305 1 1.7a 4.29 8.35 308 1 9.0b 6.68 13.0 312 38.2a 9.82 18.9 315 55.1 10.65 24.2 318 89.4" 16.42 31.1 a Interpolated or extrapolated values (cf.text). In good agreement with earlier results [ref. (l)]. We have established k , by monitoring dissociation of the complex in acid solution (0.5 mol dm-3 sulphuric acid) at 298,308 and 315 K (table 1); k, at other temperatures was obtained by interpolation or (3 18 K) extrapolation; the Arrhenius dependence yields an activation energy of 126 kJ mol-l. Rate constants kobs are reported in table 1 ; k, was obtained from (kobs-kl)/[CN-]. All k, and kobs values reported in table 1 represent the means of three independent consistent determinations; they are quoted to a slightly higher precision than their accuracy warrants to avoid any information loss from rounding-off in the statistical analysis. ACTIVATION PARAMETERS According to transition-state theory,lg for reaction in solution reactants are in equilibrium with the transition state.The rate constant for reaction is given by where we have assumed that the transmission coefficient is unity and the solution is ideal. Further, AlG*(T) =-RTlnSKe(T) =pP(sln;T)-Z(j= l ; j = vlvj/pp(sln;T) (4) At Vm(sln; T ; p ) = VF(s1n; T ; p ) - C ( j = 1 ; j = v) I vi 1 Vj"O(s1n; T ; p ) . (5) and Thus the standard equilibrium constant $K*( T ) for the equilibrium between reactants and transition state is dependent on temperature and independent of pressure. Further, K e ( T ) is related through eqn (4) to the chemical potentials of reactants and transition states in their solution standard states at temperature T and standard pressure p e .ISOCHORIC ACTIVATION PARAMETERS According to eqn (3), the dependent variable In ( k / T ) can be expressed as a function of independent variables T and p : In ( k / T ) = In ( k / T ) ( T ; p ) . (6) Similarly the molar volume of the (pure) solvent Vi;" can be defined by the same independent variables : Vi;" = Vi;"(T; p ) . (7)14 ISOBARIC, ISOTHERMAL AND ISOCHORIC PARAMETERS The argument follows that described2 for equilibrium quotients. Consider a given temperature 8 and pressure n, where the molar volume of the pure solvent is V,*(8; n). We assert that there exists a temperature 8 + A 8 at pressure n+An, where the condition holds, Anl being characteristic of the solvent. According to eqn (6) there also exists a quantity In (k/T) at (O+A8, n+An,).In other words we may compare In (k/T) for reaction in solutions under conditions where the molar volume of the solvent is the same, eqn (8). The gradient of the tangent in the ln(k/T)-T-p domain, conforming to the isochoric condition in eqn (8), can be calculated from the isobaric-isothermal v,*(e; n) = v,*(e+Ae; ~ + A Z J (8) gradients The latter equation can be rewritten in terms of the enthalpies and volumes of Here a,*, K,* and #I,* are the thermal expansivity, isothermal compressibility and isochoric thermal pressure coefficient of the pure solvent, respectively. An analogous line of argument can be used in conjunction with eqn (8) such that, following an increase in pressure from n to n+An, there exists a temperature 8+A8, where the following condition holds: (1 1) V,*(e; n) = V,*(6+A8,; n+Alt).Here A8, is characteristic of the solvent. Hence AW~(T;P) 1 A%H~(T;P) +- RT #I,* RT2 . Again the calculated quantity is isochoric (V,*), meaning isochoric with respect to the molar volume of the solvent. NUMERICAL ANALYSIS The dependence of (k/T) on temperature and pressure was fitted to the equation ln(k/T) = ln(k[6; n]/8)+a,(T-8)+a,(p-n)+a,(T-8)2 + a,(p - n)2 + a,(p - n) (T- 8) + a6(T- 8)2 ( p - n) (1 3) about a reference temperature 6 and reference pressure n. In practice the rate constant k was corrected using the density of the pure solvent as indicated in eqn (3). BENZYL CHLORIDE First-order rate constants for the solvolysis of benzyl chloride were taken from published data,209 21 including the data reported by Robertson and Scott.22 Estimates of parameters in eqn (13) are summarized in table 2.Derived activation parameters are set out in table 3. In table 2 VR = V,*(e;n). The negative volume of activation is consistent with that reported by Hyne.lo We also calculate that AjVw(sln:T:p) decreases with pressure, but the dependence is smalllo and on the borderline of statistical significance.M. J. BLANDAMER et al. 15 Table 2. Benzyl chloride in water: numerical analysisa (n/bar = 1380, 8/K = 323.4, VR/m3 mol-' = 17.31 x parameter estimate standard error ~~ ~ ~~ In (k[8; .]/8) - 13.7108 6.9 x 10-3 a,/K-' 0.10059 6.2 x 10-4 a3/K-2 -4.233 x lo-* 8.07 x 10-5 a,/ bar-' 7.242 x 4.7 x 10-9 a,/bar-' K-' 2.339 x 6.1 x 10-7 a, / bar-' 3.519 x lop4 6.41 x lop6 a,/bar-' KP2 2.8 x 6 x a Standard error = 2.37 x (degrees of freedom = 43).Table 3. Benzyl chloride in water: derived parameters property value reference temperature, B/K reference pressure, n/bar V*(H,O: 1: 8: n)/m3 mol-' In (k[8: n]/@ (obs) A: Vm(sln: 8: n)/m3 mol-1 [aAI Vm(sln: 8: n)/ap],/m3 mol-l bar [aAI V"(s1n: 8: n)/aT],/m3 mol-' K-' AjHa(sln:8:n)/kJ mol-l [aAiHa(sln: 8: n)/@]/J mol-l bar-' AlC,"(sln: 8: n)/J K-' mot1 {(a In k$[B: n]/aT) at VR]/K-' A:v/( Vf)/kJ mol-' {[aAJly( Vf)/dT] at VR}/J K-' rno1-l {[aA: Vm(sln: O:n)/dT] at VR>/m3 mol-l K-l 323.4 1380 17.31 x - 13.721 -(9.46+0.12) x lo-' -(3.89f2.5) x lo-, -(2.95+0.05) x 87.47k0.54 0.04 0.08 0.1055+0.0004 - 195 k 143 91.74f0.51 - 143f 122 9.76 Table 4. Kinetics of reaction between cyanide ions and tris(4-methyl- 1,lO-phenanthroline) iron@) in aqueous solutiona (B/K = 298.2, n/bar = 68, Vg/m3 mol-l = 18.02 x lo-,) parameter estimate standard error In [k(e: n)/e] - 7.3274 5.9 x a,/K-' 0.10704 5.8 x 10-3 a,/bar-l -2.8065 x 10-4 9.2 x 10-5 a 11 data points: standard deviation = 1.27 x 10-l.CYANIDE IONS AND METAL COMPLEX The data for this reaction cover smaller pressure and temperature ranges than for the previous example. However, the consistency obtained by the analysis for benzyl chloride gave support to the idea of using the same procedures for the second-order reaction involving cyanide ions (tables 4 and 5). Nevertheless fewer terms were statis- tically significant in the fitting of the kinetic data to eqn (13).16 ISOBARIC, ISOTHERMAL AND ISOCHORIC PARAMETERS Table 5. Derivation parameters for reaction between cyanide ions and tris(4-methyl-1,lO- phenanthroline)iron(Ir) in aqueous solution parameter value AIH"(s1n: 8:z)lkJ mol-l 79.4k4.3 At V"(s1n: 8: z)/m3 mo1-I (6.96 f 2.2) x lop6 [a In (k(8: ~ ) / 8 ) / a T ] at 0.1044 & 0.0052 A$w( V:)/kJ mo1-I 77.22 f 3.84 [a In (k$(O: z)/8)/a V:]/mol m-3 (4.09 & 0.11) x lo6 Both sets of kinetic data are used to calculate a quantity ASv/(VF) defined by eqn (14) which is expressed in J mol-l : An analogous quantity AS@( V,*) calculated using eqn (1 5 ) is expressed in m3 mol-1: As@( V,*> RT * (15) DISCUSSION Relative to the extensive literature concerned with reactions of organic solutes, the literature dealing with the effect of pressure on rates of reactions concerning metal complexes remains small.23-25 Within this general area, isochoric activation parameters have also attracted little attention, whereas these quantities have prompted considerable interest for organic reactions. The claims made for isochoric activation parameters are strong. Hills and Vianall suggest that isobaric are more complicated than isochoric activation parameters. Further, these authors1' suggest that the isobaric heat capacity of activationzs* 27 is particularly complicated. Whalleyz3 has argued that isothermal-isochoric parameters are more fundamentallo, 28 and 29 particularly for understanding the depen- dence of activation parameters on composition of solvent for reactions in binary aqueous mixtures. Holterman and Engberts7 find, however, no compelling evidence, at least for one reaction, for preferring either isobaric or isochroic activation par- ameters.Nevertheless, our concern with the enthusiasm for isochoric activation parameters is, we suggest, more fundamental. From the outset, the term isochoric is not used in the sense normally used in thermodynamics. As far as we can discern from published papers, the rate constant for a given chemical reaction always characterises the approach of a system to a minimum in G under isothermal-isobaric conditions rather than to a minimum in F under isothermal-isochoric conditions. Reported isochoric functions are calculated using, for example, eqn (lo), in which the properties of the pure solvent are used. In other words, isochoric means 'at constant molar volume of the pure solvent'.Hence we have used the term isochoric (V;"). Other isochoric functions have been commented on. Whalley30 has argued that the volume kept constant is 'the volume of an equilibrium mixture of initial and transition state '.M. J . BLANDAMER et al. 17 The volume of the reacting system is given by I/(sln: T : p ) = (l/Ml) Vl(sln: T : p ) + C ( j = 2; j = i)mj y(s1n: T:p)+m$ V&sln: T : p ) . (16) For this system ml < mj and mi for all reactants and products is dependent on time as is F(sln : T : p ) . We might envisage a system containing a pseudochemical equilibrium between reactants and transition state: V(hypothetica1 equilibrium; sln: T : p ) = (l/Ml) Vl(sln: T:p)+mfq Vfq(s1n: T : p ) +C ( j = 2; j = r ) I vj I mFq Vfq(sln: T : p ) . (17) An alternative method considers a solution in 1 kg of solvent comprising 1 mol of transition state and I vj 1 mol of reactants where the partial molar volumes of reactants and transition states equal the volumes at infinite dilution.Hence, by definition, Vm(sln: T : p ) = (l/Ml) VT(1: T : p ) + Vr(sln: T : p ) + C ( j = 2; j = r ) I vj I VY(s1n: T : p ) . (18) For example, in the case of a second-order reaction, with reacting solutes 2 and 3, Vm(sln: T : p ) = (l/Ml) V,*(l: T : p ) + Vr(s1n: T : p ) + VP(s1n: T : p ) + VP(s1n: T : p ) . (19) Unfortunately it is not immediately obvious how one might calculate the corresponding thermal and pressure coefficients for these solutions. Alternatively a set of isochoric (AIVm) parameters may be calculated; a pro- posal along these lines was made by Caldin.31 In effect, the volume of activation At V"(s1n: T : p ) is treated as a dependent variable defined, for a given system, by the independent variables T and p .The analysis described above is repeated where At Vm replaces V,* in eqn (7), (8) and (1 1). Then, for example, we may define the partial differential Irrespective of the isochoric condition, our contention is that these constant volumes are extrinsic to the system undergoing reaction. It may also be useful to calculate AZv/(VT) as defined in eqn (14), but this term is not the thermodynamic energy of activation. Nor is its temperature derivative the molar isochoric heat capacity of activation. Finally we accept that isochoric functions may provide an insight into mechanisms of reaction.However, the precise meaning of these quantities must first be established. We hope that we have achieved at least part of that task. For the moment, however, we conclude that isobaric activation parameters26* 27 are free from the ambiguity associated with isochoric parameters. We thank the S.E.R.C. for a maintenance award to B.C. We thank Dr E. Whalley for providing reprints of his papers. 2 FAR 118 1 2 3 4 6 6 7 8 9 10 11 12 13 14 15 18 17 18 18 20 21 2P 23 24 25 26 27 28 28 30 31 ISOBARIC, ISOTHERMAL AND ISOCHORIC PARAMETERS F. M. Mikhail, P. Askalani, J. Burgess and R. Sherry, Transition Met. Chem., 1981, 6, 51. M. J. Blandamer, J. Burgess, B. Clark and J. M. W. Scott, J. Chem. SOC., Faraday Trans. 1, 1984,80, 3359. B. T. Baliga, R.J. Withey, D. Poulton and E. Whalley, Trans. Faraday SOC., 1965, 61, 517. G. S. Kell and E. Whalley, Philos. Trans. R. SOC. London, Sect. A, 1965, 258, 565. G. S. Kell, J. Chem. Eng. Data, 1970, 15, 1 19. B. T. Baliga and E. Whalley, J. Phys. Chem., 1967, 71, 1 166. H. A. J. Holterman and J. B. F. N. Engberts, J. Am. Chem. SOC., 1982,104,6382. J. F. Nary, M. A. Simard, J. Dumont and C. Jolicoeur, J. Solution Chem., 1982, 11, 755. M. Nakagawa, Y. Miyamoto and T. Moriyoshi, J. Chem. Thermodyn., 1983,15, 15. D. L. Gay and E. Whalley, J. Phys. Chem., 1968,72,4145. G. J. Hills and C. A. Viana, in Hydrogen-bonded Solvent Systems, ed. A. K . Covington and P. Jones (Taylor & Francis, London, 1968), p. 261. M. J. Blandamer, J. Burgess, B. Clark, P. P. Duce and J. M. W. Scott, J. Chem. SOC., Faraday Trans. I , 1984,80, 739. D. W. Margerson and L. P. Morgenthaler, J. Am. Chem. SOC., 1962,84, 706. J. Burgess, Inorg. Chim. Acta, 1971, 5, 133. R. D. Gillard, Inorg. Chim. Acta, 1974, 11, L21; Coord. Chem. Rev., 1975, 16, 67; 1983, 50, 303. N. Serpone, G. Ponterini, M. A. Jamieson, F. Bolletta and M. Maestri, Coord. Chem. Rev., 1983,50, 209. M. J. Blandamer, J. Burgess and P. Wellings, Transition Met. Chem., 1979, 4, 95; 1981, 6, 364. J. A. Arce Sagiies, R. D. Gillard and P. A. Williams, Transition Met. Chem., 1979, 36, L411 and references therein. S. Glasstone, K. J. Laidler and H. Eyring, Theory of Rate Processes (McGraw-Hill, New York, 1941). M. J. Mackinnon, A. B. Lateef and J. R. Hyne, Can. J. Chem., 1970,48,2025. D. L. Gay and E. Whalley, Can. J. Chem., 1970,48,2021. R. E. Robertson and J. M. W. Scott, J. Chem. SOC., 1961, 1596. E. Whalley, Adv. Phys. Org. Chem., 1964, 2, 93. R. van Eldik and H. Kelm, Rev. Phys. Chem. Jpn, 1980, 50, 185. S. D. Hamann, Rev. Phys. Chem. Jpn, 1980,50, 147. R. E. Robertson, Progr. Phys. Org. Chem., 1967,4, 213. M. J. Blandamer, R. E. Robertson and J. M. W. Scott, Prog. Phys. Org. Chem., in press. B. T. Baliga and E. Whalley, Can. J. Chem., 1970, 48, 528. B. T. Baliga and E. Whalley, J. Phys. Chem., 1969,73, 654. E. Whalley, Ber. Bunsenges Phys. Chem., 1966, 70, 958. E. Caldin; quoted in ref. (7). (PAPER 4/069)

ISSN:0300-9599    

DOI:10.1039/F19858100011

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1985, 81, 19-36 Study of Methanol and Water Chemisorbed on Molybdenum Oxide BY JONG s. CHUNG, RAUL MIRANDAT AND CARROLL 0. BENNETT* Department of Chemical Engineering, University of Connecticut, Storrs, Connecticut 06268, U.S.A. Received 16th January, 1984 Methanol and water chemisorbed on finely divided molybdenum oxide have been studied by infrared spectroscopy and by dynamic adsorption and desorption techniques. Isolated hydroxyl and two kinds of water on the surface desorb at ca. 150 "C, while hydrogen-bonded hydroxyls desorb at 350 "C. Most of the methanol chemisorbs dissociatively, even at room temperaure, and at least three forms of chemisorbed methanol exist : methanol dissociated into methoxy and hydrogen across Mo=O (form C), methoxy on the vacancy of terminal-bonded oxygen Mo=O (form A) and methoxy on the vacancy of bridge-bonded oxygen Mo-0-Mo (form B).Most of the form C desorbs reversibly as methanol below 110 "C and a part of the form A is desorbed as methanol with a peak at 110 "C during temperature-programmed desorption. The rest of the methoxy groups are decomposed into formaldehyde and CO at ca. 260 "C. The mechanism of the selective oxidation of methanol on Mo-Fe-0 mixed oxides has been extensively studied since the early There is a growing concern about the importance of the methoxy group as a precursor of various products.276-8 However, the presence of the species has not been confirmed by infrared spectroscopy, since the surface area of conventionally made MOO, or molybdate is very low.Also it is known that water in the gas phase is a strong inhibitore and that the selectivity for the products depends on the hydroxyl content on the catalyst surface.l0 This paper presents a basic study of the chemisorption of water and methanol on MOO, in order to contribute to our understanding of the nature of the reaction intermediates on active sites. Infrared spectroscopy and dynamic adsorption/ desorption techniques have been used to study water and methanol adsorbed on unsupported MOO,. Finely divided MoO,ll prepared in a flame reactor is suitable for the infrared study because of its high surface area and good transmission qualities. EXPERIMENTAL CATALYST PREPARATION The catalyst consisted of finely divided unsupported MOO, particles made in a flame reactor and obtained from Teichner's laboratory in Lyon.ll The B.E.T.surface area after extensive treatment in oxygen at 380 "C was 27 m2 g-l and the average particle size measured by X-ray line broadening was 450 A. The reduced MOO, for the adsorption experiments was obtained by passing 3.6% methanol in helium over the catalyst at 150 "C, then passing helium for < 5 min to clean the surface. The reduced catalyst was cooled to room temperature quickly in -f Present address : Department of Chemical Engineering, University of Louisville, Louisville, Kentucky 40292, U.S.A. 19 2-220 CHEMISORPTION OF H20 AND CH,OH ON MOO, order to minimize oxygen diffusion from the bulk. The desired degree of reduction could be controlled by the length of the reduction with methanol. A higher degree of reduction was possible by repeating the reduction and cleaning cycle.The degree of reduction on the surface of the catalyst was measured by the normalized absorbance at 1320 cm-l, as described previously.12a Oxodized MOO, refers to the catalyst obtained by heating at 380 "C overnight in oxygen and cooling to room temperature in oxygen. DYNAMIC ADSORPTION AND DESORPTION EXPERIMENTS EXPERIMENTAL SYSTEM The experimental system consisted of the feed system, which was the same as that for the infrared system, the reactor and the mass spectrometer with a minicomputer. The differential reactor and tubing were made of SS316. The grain size of the pressed catalyst particles was between 60 and 100 mesh. The analysis of gases was made with an on-line magnetic-sector-type (12 in* radius analyser) nuclide mass spectrometer which was connected to a MINC[(R)DEC] minicomputer to follow up to 3 peaks per second.Since various components had contributions to several of the monitored peaks, their individual cracking patterns were determined using pure components in order to convert the collected data into concentrations against time. EXPERIMENTAL PROCEDURE Just before an adsorption experiment MOO, was subjected to another temperature increase up to 400 "C in 3 % 0, +He to remove hydrated water. The catalyst was then cooled, purged with helium for 10 min and an adsorbate admitted. The pulse duration was 4 min, a time sufficient to attain a steady state. Subsequently, helium was introduced into the reactor for 5 min.Still in flowing helium, the temperature of the reactor was increased at an initial rate of 60 "C min-l and an average rate of 23 "C min-l until desorption activity ceased. ANALYSIS The analysis of the reactor inlet and outlet was accomplished by monitoring peaks at mass/charge ratios of 1 8(H,O), 28(CO), 30(CH,O), 3 1 (CH,OH) and 40(Ar), representing all the species present below 100 "C. Sensitivities stayed in the 1 ppm range with white noise of k2 ppm. In order to eliminate some residual noise (60 Hz), 100 samples were taken in 35 ms and averaged to get each peak. In order to calculate the amount of methanol adsorbed (or desorbed) during the step input of gaseous methanol (or helium), a tracer component, Ar, in the gas mixture was used as a standard curve of step change in the reactor system.After the concentration of Ar was set equal to that of 3.6% methanol, the difference in the response curves between adsorbate and Ar was used to calculate the amount of methanol adsorbed (or desorbed). INFRARED METHODS The whole system was the same as that described When the chemisorbed intermediates of methanol and water were observed, methanol and water were passed through both the sample and reference cell and the infrared beam was balanced carefully between the sample and reference cell by adjusting the distance between the surface of the catalyst (or the stainless-steel mirror in the reference cell) and the infrared window. In order to reduce the severe transmission loss in the high-frequency region, above 3 100 cm-', CaF, of 1 mm thickness was used for the window instead of ZnS.A low concentration of oxygen in the methanol mixture also results in the loss of transmission in the high-frequency region, especially at high temperatures, because of the reduction of MOO,. To avoid this, methanol in oxygen was used during the chemisorption of methanol on MOO,. The optical system was purged with dry air to minimize the contamination of infrared bands by moisture. Real transmission values of the samples are < 5% in the 4000 cm-l region and ca. 10-2074 in the region from 2000 to 1000 cm-l. * 1 in = 2.54 x m.J. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 21 RESULTS AND DISCUSSION INFRARED SPECTRA OF HYDROXYL GROUPS Fig. 1 shows the OH bands of water adsorbed and irreversibly held in helium at various temperatures.Bianchi et al.13 have observed OH bands of adsorbed water on the same catalyst and concluded that water did not chemisorb dissociatively. This conclusion was based on the bands observed at 3500 and 1610 cm-l. Besides these bands we detected more bands at 3685-3705,3375,3255 and 1625 cm-l. The impurity bands above 3705 cm-l result from moisture in the optical chamber. Even though dry air was purged into the optical chamber, the contamination could not be removed completely because the balanced beam was greatly magnified in order to observe the weak OH bands of water. This contamination became more severe when OH bands of chemisorbed methanol were observed. The band at 3685-3705cm-l is also confirmed by observing OD bands with deuterated methanol (vide infra) and is assigned to isolated hydroxyl.The band at 3500cm-l is assigned to physically adsorbed water. The bands at 1690 and 1640 cm-l have already been assigned to an impurity and to a structural band of MOO,, re~pective1y.l~ A close look at fig. 1 reveals that the bands at 3480-3500 and 1610 cm-l disappear at ca. 100 "C and the bands at 3585 and 1625 cm-1 remain up to 185 "C. Therefore these two sets of bands correspond to water chemisorbed undissociatively, since they have water deformation bands at 16 10 and 1625 cm-l. There remain no corresponding water deformation bands for the bands at 3375 and 3255 cm-l. Dihydrated MOO, has coordinated and interlayer water bands at 3520, 3410,3200, 3140 and 1615 cm-l, and monohydrated molybdenum oxide has coordinated water bands at 3430 and 1605 crn-l.l4 However, it is known that upon heating, the dihydrate and mono- hydrate bands disappear at 80 and 130 OC, respectively.Therefore the two bands at 3375 and 3255 cm-l cannot be water bands of hydrated MOO, because they remain up to 350 "C. The assignments of these two bands will be discussed later. In order to find proper chemisorption sites for the two chemisorbed waters which have bands at ca. 3480 and 1610 cm-l and ca. 3580 and 1625 cm-l, we have studied the effect of traces of water in the carrier gas on the chemisorption of water on MOO,; the results are shown in fig. 2. Carrier-grade helium and zero-grade oxygen (Aero All Gas Co., Hartford, CT) were used without further purification, so that they contained traces of water.The band at 3400 cm-l is developed by the water in the oxygen and the band at 3500 cm-l by the water in the helium. If we assume that there is less moisture impurity in helium than in oxygen the results indicate that oxygen has a blocking effect for the chemisorption of water on reduced sites. This leads us to the conclusions that the water band at ca. 3500-3580 cm-l in fig. 1 and 2 corresponds to water chemisorbed on a reduced site of MOO, and that the band at ca. 3460- 3480 cm-l corresponds to water on an oxygen of MOO,. The results also imply that there are vacancies on the oxidized MOO,. Fig. 3(A) shows the OH bands which are developed after passing methanol+ oxygen over MOO, at 27°C. All of the OH bands which are developed by the adsorption of water on MOO, are also found here at the same frequencies.The absence of any new hydroxyl bands indicates that methanol dissociates to form methoxy and hydrogen on the surface of MOO,. The main difference between the OH bands produced by water and by methanol is that water adsorption shows a maximum peak at 3500 cm-l, whereas methanol adsorption shows a maximum at 3375 cm-l. Almost no increase in the water deformation band at 1625 cm-l is observed at room temperature, suggesting that the bands at 3375 and 3255 cm-1 do not arise from water22 CHEMISORPTION OF H20 AND CH,OH ON MOO, Fig. 1. Infrared bands of hydroxyl group and water in OH stretching (left) and water in deformation region (right) on MOO,, produced by water adsorption and desorption at (a) 320, (b) 250, (c) 185, (d) 145, (e) 90, cf) 50 and (g) 27 "C.(-) 4.6 Torr water in helium and ( * * - -) after a purge with helium for 20 min. The ordinates of the spectra are displaced to avoid overlapping of traces. molecules. Hydroxyl bands near 3375 cm-l are also observed with Ti0215 and a-Fe,0,.16 Hydroxyl bands developed by the chemisorption of methanol at temperatures > 27 "C are shown in fig. 3(B). When methanol in oxygen is present in the gas phase at 60 "C, the intensities of the water peaks are greater than those of the bands at 3375 and 3325 cm-l, indicating that water desorption is the rate-determining step at this temperature. As the temperature increases, however, the amount of chemisorbed water decreases sharply and almost disappears at 200 "C in the presence of 3.6% methanol in oxygen.Unfortunately it is impossible to observe the OH bands in the presence of lower oxygen concentrations than 3.6% CH,OH + 96.4% oxygen, especially at high temperatures, because of the severe loss of the background transmission at high wavenumber. The shift in the hydroxyl band at 3375 to 3395 cm-l above 100 "C is probably related to the reduction of MOO,. In order to see the effect of the reduction of MOO,, OD bands were studied because there was still enough transmission in the OD vibration region with reduced MOO,.J. S. CHUNG, R. MIRANDA AND C. 0. BENIWTT 23 wavenumber/cm-' Fig. 2. Inhibition effect of oxygen on the chemisorption of water on MOO, observed in the OH stretching region at 27 "C. (- - - -) Background of MOO,. (a) He, 2 h ; (b) He, 1 day; (c) 0,, 1 day; ( d ) 0,, 2 days.Fig. 4(A) shows OD bands on the oxidized (upper part) and 10% reduced MOO, (lower part) after D,O, CH,OD and CD,OD have been passed over the catalysts for 20 min at 27 "C, and fig. 4(B) shows the OD bands remaining on both oxides after a purge with helium. On the oxidized MOO,, all the OD bands which correspond to the five OH bands in fig. 1 and 3 are observed as shown on the lower part in fig. 4(B). On the reduced MOO,, four of them are shifted to higher wavenumbers compared with those on the oxidized MOO,, but the band 2435 cm-l is not shifted. Table 1 shows the band positions of all OH and OD bands on both oxidized and reduced MOO,. We have already assigned the two water bands based on the results in fig. 2.With the results obtained so far and based on the structure of MOO, we will now assign the three hydroxyl bands. MOO, is known to have a layered structure17 with a protruding terminal double- bonded oxygen (Mo=O), two single-bonded oxygens shared by two Mo and another three shared by three Mo (Mo-0-Mo). From now on we will use OT to denote the double-bonded oxygen in Mo=O, 0, for the bridge-bonded oxygen in Mo-0-Mo, and V, and V, for their corresponding oxygen vacancies. For water and methanol to produce hydroxyl on the surface they must be dissociated as follows: H,O+V+OL + HOV+HOL (1) (2) where the subscript L represents lattice oxygen to distinguish it from the oxygen of adsorbates, V represents an oxygen vacancy of either VT or V, and * is used to represent a site on the surface on which methoxy chemisorbs.According to reactions (1) and (2), hydroxyl can be produced by an oxygen vacancy upon decomposition of chemisorbed water, while this is not true for methanol unless chemisorbed methanol H,COH + * + 0, -+H,CO* + HOL24 CHEMISORPTION OF H,O AND CH,OH ON Moo, 3900 3700 3500 3300 3100 2900 wavenum ber/cm-' Fig. 3. For legend see facing page. can break the CO bond as easily as water does the OH bond. Breaking the OH bond in water is much easier than breaking the CO bond in methanol on the reduced sites of MOO,. Evidence for this is provided by the result that the irreversibly chemisorbed hydroxyls produced by D,O chemisorption have higher intensities than those formed by methanol chemisorption, especially when MOO, is reduced, as shown on the upper part of fig.4(b). On the oxidized catalyst, the hydroxyls produced by CD,OD show the highest intensity among the three adsorbates shown in fig. 4(b), indicating that breaking the CD bond, in other words the'production of surface formaldehyde, occurs even at room temperature. Once the surface is covered with hydroxyls formed by either reaction (1) or (2), they may be linked with each other by multiple hydrogen bondings: H H I I I I I I I I 0. .. . . . H . . . . . O . . . .. . H . . . . . .O Mo-0-Mo-0- Mo VT VB The hydrogen-bonded hydroxyl on V, (or HOB) is affected more strongly than the free hydroxyl on V, (or HOT), since the former interacts directly with neighbouring terminal oxygens. As the catalyst is reduced, oxygen vacancies are formed on the OT and OB sites so that the chance of hydrogen bonding between hydroxyls is reduced.J.S. CHUNG, R. MIRANDA AND C. 0. BENNETT /-I_ 3395 I 3700 3500 3300 3100 2900 wavenumber / cm- l 25 Fig. 3. (A) Hydroxyl bands produced by adsorption and desorption of methanol on MOO, at 27 "C. (-) 3.6% methanol in oxygen and ( * * .) after purging with helium for (a) 35, (b) 25 and ( c ) 6 min. The ordinates of spectra are displaced to avoid overlapping of traces. (B) As (A) but taken at (a) 200, (b) 160, (c) 100 and (d) 60 "C. (. - - .) After a purge with helium for 20 min. Thus the hydroxyl band observed at 3685-3740 cm-l is assigned to free hydroxyl on V, and the band at 3375-3430 cm-l to multiple hydrogen-bonded HOB.The band at 3235-3255 cm-l is at an unusually low wavenumber for a hydroxyl band. Chromia18 showed the same kind of broad band at ca. 3280cm-l, but no assignment was given. On MOO, it probably results from hydroxyl perturbed by hydrocarbon impurities. Strong interactions between hydroxyl and chemisorbed hydrocarbons develop wide bands at 3300cm-l on ~i1ica.l~ Oxidized MOO, has impurity bands at 1950, 1930 and 1880cm-l, which are similar to the bands of a molybdenum carbonyl complex.20 However, the mechanism of interactions between hydroxyl and impurity bands is not clear. The band could be assigned to a hydroxyl in the interlayer of MOO,, because only this band is not shifted upon the reduction of MOO, and no corresponding water deformation band is observed.Molybdenum oxide is known to form hydrogen bronzes H,MoO, (0 < x < 2). Recent studies,' show that H is held between the MOO, layers forming OH, and for higher concentration of H forming OH, groups.26 CHEMISORPTION OF H20 AND CH,OH ON Moo, 2592 2800 2700 2600 2500 2400 2300 wavenumber/cm- 1 Fig. 4. For legend see facing page. INFRARED SPECTRA OF METHOXY GROUPS In the case of methanol chemisorbed on other oXides22-2s the chemisorbed form of methanol is usually regarded as methoxy after being compared with the bands of metal methoxides. It has been assumed that the methoxy on metal oxides is formed either by a condensation reaction or by the opening of the metal-oxygen bond:22-26 H CH3 I 1 CH, H I I 0 0 I I M+CH,OH+M + M. /"\ M (3') Generally three strong bands are observed at ca.3000,292&2960 and 28 15-2860 cm-l and assigned to overtones or combinations of CH, deformation (6) and asymmetric (v,) and symmetric (v,) stretching, respectively. In the case of silica,24 4 bands areJ. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 27 (B) 2690 I I ox; d ; sed M O 03 2735 26kO . 2435 2580 I 2500 wavenumber/cm-' Fig. 4. (A) Infrared bands of OD on reduced and oxidized MOO, at 27 "C after adsorptions of (a) 3.6% CH,OD+He, (b) 3.6% CD,OD and (c) 5 Torr D,O. (B) As (A) for irreversibly bound OD remaining after a purge with helium for 20 min. Table 1. Wavenumbers and assignments for the irreversible hydroxyl bands at 27 "C wavenumber/cm-l oxidized MOO, reduced MOO, OH OD OH OD (3685) 2735 (3740)" 2770 3580 2650 (363 1) 2690 3480 2580 (3534) 2618 3375 2500 (3529) 2540 3255 2435 (3287) 2435 assignment free hydroxyl of HOT water in oxygen vacancies water on oxygen of MOO, hydrogen-bonded HOB hydroxyl perturbed by hydrogen impurities, or in interlayer of MOO, " The number in parenthesis is a value calculated with the ratio of v(OH)/v(OD) = 1.35.28 CHEMISORPTION OF H20 AND CH,OH ON Moo, t 2955 3100 3000 2900 2800 2700 wavenumber/crn-l Fig.5. Infrared spectra of methanol in the gas phase (a) and chemisorbed on MOO, (b) at 27 "C. (-) 3.6% methanol + oxygen, ( * * a ) after purging with helium for 20 min and (- - -) MOO,. observed above 2900 cm-l, and the band at 2928 cm-l is assigned to an overtone or combination of CH, deformations. Fig. 5 shows the CH stretching bands of gas-phase methanol and methanol adsorbed on MOO, at 27 "C.According to the assignment of others the bands at 3006, 2955-2965 and 2845-2854 cm-l in fig. 5 are assigned to 6, v, and v,, respectively. For the two weak bands at 2930 and 2830 cm-l, we prefer the assignment as v, and v, of a different type of methoxy, since these bands do not seem to correspond to any reasonable combination of those of the strong methoxy. Also, the intensity ratio and the distance between these two weak bands are similar to those between the bands at 2955 and 2845 cm-l. The second weak v, at 2830 cm-l has not been observed on other oxides. The infrared spectra of [2H,]methanol chemisorbed on MOO, also confirm the existence of the two weak bands, as shown in fig. 6. All the corresponding bands developed by chemisorbed methanol are detected by the chemisorption of [2H,]methanol, although the band intensity at 2177 cm-l is weaker than the intensity of the corresponding band 2955 cm-l in fig.5. Wavenumbers and assignments of the observed bands are given in table 2. Fig. 7 shows the bands in the CH, deformation region developed by the chemi- sorption of methanol on molybdenum oxides. Other bands which are not specified in fig. 7 result from moisture in the optical chamber. The band at 1430 cm-l is alsoJ. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 29 , I I' 2 2 4 0 I I I I 2000 2300 2200 2100 wavenum ber/cm-l Fig. 6. Infrared spectra of chemisorbed CD,OD in the CD stretching region at 27 "C. (-) 3.6% methanol+ helium, (. * * .) after purging with helium for 20 min and (- - -) MOO,.Table 2. Wavenumbers and assignments of observed bands of chemisorbed methanol and [2H,]methanol at 27 "C CH,OH CD,OD a b a b assignment 3006 - 2955 2965 2925 2928 2845 2854 - 2825 1467 1467 1445 1445 13781 1345 2248 2284 overtone or combination of CH, deformation, form A 2225 overtone or combination of CH, deformation, form B 2177 2185 v,,(CH,), form A 2136 2139 vas(CH3), form B 2075 2085 v,(CH,), form A 2068 v,(CH,), form B - - - - CH, deformation, form A - - CH, deformation, form B CH deformation, formaldehyde - - a In the presence of 3.6% methanol in the gas phase. After purging with helium.30 CHEMISORPTION OF H,O AND CH,OH ON Moo, 1445 I I I I 1550 1450 1350 wavenumber/cm- ' Fig. 7. Infrared spectra of chemisorbed CH,OH in the CH, deformation region at 27 "C with 3.6% CH,OH+He.(a) After catalyst is oxidized at 350 "C for 2 h, (b) reduced with 3.6% methanol+ helium at 145 "C for 12 min and (c) reduced under the same conditions as (b) for 45 min. observed on haematite22 and is assigned to a vibrational or electronic transition produced by depletion of structural oxygen atoms from the surface. Both chemisorbed formaldehyde and formic acid can develop the bands at 1378 and 1345 cm-l. In the case of MOO,, these are assigned to formaldehyde since the band at 2785 cm-l, which belongs to chemisorbed formaldehyde, is observed in fig. 6 (2033 cm-l as CD). The two bands at 1467 and 1445 em-l are CH, deformation bands, but it is not clear whether they represent two different species of methoxy or not.Although the methanol molecule shows three CH, modes, usually only one band is observed at 1465-1475 cm-l in the case of other oxides.26-28 There is an exception: haematite22 shows two 6(CH,) bands at 1460 and 1440 cm-l. However, they are regarded as arising from one type of methoxy, because only one set of v(CH) at 2920 and 2815 cm-1 is observed. Fig. 8 is further evidence of the existence of two different types of methoxy. As the degree of reduction of MOO, increases, the intensity ratio of the band at 1445 cm-l to that at 1467 cm-l increases. This ratio increases even more in the case of methanol held irreversibly after purging with helium, as shown on the right-hand side of fig. 8. Of course the total amount of methoxy on the surface increases as the reduction of MOO, increases.Note that the intensity of the formaldehyde band at 1378 em-l decreases as the reduction increases. This results from a deficiency of surface oxygen capable of taking hydrogen from chemisorbed methoxy. that oxygen of Mo=O is more labile than that in Mo-0-Mo on the surface and that the intensities of the weak CH stretching bands at 2930 and We haveJ. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 31 I I I 1550 1450 1350 wavenumber/cm-' i ; 1378 . . I . ( C ) : I , , .' I , , I 1467 j W'.. . : , . L, 1448 wavenumber/cm-' I I I 1550 1450 1350 Fig. 8. Effect of reduction on relative strength of two different species of methoxy observed at 27 "C in the presence of 3.6% CH,OH + He (-) and after a purge with helium for 20 min ( - . . .).(a) Oxidized at 350 "C for 2 h, (b) 3% reduced on the surface and (c) 10% reduced on the surface. 2830 cm-l increase with increased reduction of MOO,. Therefore the 6(CH3) band at 1445 cm-l is connected to the v(CH) bands at 2930 and 2830 cm-l, and the 6(CH3) band at 1467 cm-l to the v(CH) bands at 2965 and 2854 cm-l. Since the Mo=O is more labile than Mo-0-Mo, it is natural to assume that the abundant methoxy with bands at 2965,2854 and 1467 cm-l chemisorbs on the oxygen vacancy of Mo=O (V,) and the methoxy with weak bands at 2930, 2830 and 1445 cm-l chemisorbs on the oxygen vacancy of Mo-0-Mo (V,). When MOO, is reduced at 150 "C, the oxygen vacancy of Mo-0-Mo (V,) is also generated and this is related to the increase of the bands at 2930, 2830 and 1445 cm-l.The results in fig. 8 show that methoxy chemisorbs on V, more strongly than on V,. Takezawa and Kobayashi30 have correlated the CH stretching bands of methoxy chemisorbed on metal oxides with the acidity (electronegativity) of the oxide; on more acidic oxides the CH stretching bands of methoxy are located at higher frequencies. Therefore the V, site is more electronegative (acidic) than the V, site because methoxy on V, has CH stretching bands at higher frequencies than that on V,. This can also be explained by the geometric effect. The methoxy on V, has more tendency to hydrogen bonding with neighbouring 0, than that on V,, which may cause CH stretching bands of methoxy on V, to shift to lower frequencies. The lower frequencies and higher probability of hydrogen bonding in the methoxy chemisorbed on V, indicate that this species is probably responsible for the pro-32 CHJZMISORPTION OF H,O AND CH30H ON MOO, duction of the observed byproducts such as dimethyl ether, dimethoxymethane and methyl formate when combined with methoxy on the V,, as follows: CH3 I I 0 ----- CH3 I 0 MO- - CHSOCHB (DME) CH3 I I CHOOCH3(MF) .CH - 0 --___ 0 MO - I The CH stretching bands of methoxy observed above 27 "C are shown in fig. 9. The results show that, especially in the presence of methanol+oxygen in the gas phase, the frequencies of the bands shift to higher values as the temperature increases. In the case of methoxy groups chemisorbed on other oxides, no band shift is observed as a function of temperat~re.~*9~~ Therefore the band shifts observed at low temperatures and in the presence of methanol in the gas phase probably indicate the existence of another type of methoxy or methanol. The existence of methanol adsorbed undissociatively is negligible if we recall that there is no new hydroxyl band observed on the chemisorption of methanol in contrast to those formed on the chemisorption of water, Methanol could decompose into methoxy and hydrogen across Mo=O as follows: HqC 0 0 Mo-00-MO + CH,OHT-- Mo-O-Mo II I1 \ / O II (4) form c Some loss (ca.10%) of Mo=O upon adsorption of methanol has been observed even at temperatures < 100 "C.lab After adsorption, however, a switch to helium returns the loss to < 10% of its original value, indicating that reaction (4) is reversible. If reaction (4) is reversible, form C will disappear after flushing the surface with helium at high temperatures.Therefore a negligible shift of the bands at 110 "C after a switch from methanol to helium indicates that the bands observed at 110 "C in the presence of helium in fig. 9 represent the pure methoxy groups chemisorbed in the V, and V,: CH3 I 0 0 I II Mo-0-MO V,, 2964 and 2858 cm-l form A H3 I 0 c o II I II Mo-0-MO VB, 2938 and 2834 cm-l. form BJ. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 33 2958 ~~ 3100 3000 2900 2800 2700 wavenumber/crn-' Fig. 9. Infrared spectra of methanol chemisorbed on MOO, in the CH stretching region at (a) I 10, (b) 90 and (c) 50 "C. (-) 3.6% methanol + oxygen and ( . . - .) after purging with helium for 20 min. The ordinates of spectra are displaced to avoid overlapping of traces.DYNAMIC ADSORPTION AND DESORPTION MEASURED BY MASS SPECTROMETRY Because of appreciable reduction of molybdenum oxide in the presence of methanol above 100 "C,lZa this experiment was limited to temperatures < 100 "C. Fig. 10(a) and (b) show the results obtained during the adsorption of CH,OH at 30 "C and temperature-programmed desorption of the chemisorbed species. Even at 30 "C, water [I11 in fig. lO(b)] is produced when methanol+ 5% oxygen is passed over MOO,. This indicates that methanol is dissociated into methoxy and hydrogen and that some of the lattice oxygens are so labile they produce water: HOL+HOL + H,O+OL. ( 5 ) Upon a switch to helium, methanol chemisorbed reversibly [I in fig. 10 (a)] and water [IV in fig.lO(b)] are desorbed. The amount of formaldehyde and CO desorbed are negligible. The catalyst temperature is then raised with helium (region A'). First methanol [I1 in fig. lO(a)] is observed with a peak maximum at 110 "C, followed by a water peak at 155 "C. Formaldehyde and CO are detected at higher temperatures with peak maxima at 235 and 245 "C, respectively. Another water peak appears as a shoulder at 265 "C. The results, together with the results obtained by infrared spectroscopy, show that there are at least three kinds of chemisorbed methoxy: reversible, irreversible and desorbed into methanol, and irreversible and decomposed34 5 4 n s c 3 .- Y (d 0 & - E o 2 1 ' 0 CHEMISORPTION OF H,O AND CH,OH ON MOO, I A 1 A' ....- ...... I-* 0 3.5 7:0 10.5 14 time/min 0 .8 ' . O 0 . 2 I 0 3: 5 710 10.5 14.0 time/min Fig. 10. (a) Dynamic adsorption and desorption of methanol on 118 mg MOO, with flow rate of 0.0013 mol min-l. Region A, pure helium; region B, 3.6% methanol+5% oxygen+3% Ar + balance helium. Region A' represents temperature-programmed desorption with helium. (-.--) Water, (. - * -) methanol and (-) Ar tracer. (b) Same as (a): (--.--.) water, (-) formaldehyde and ( - - - a ) CO. into formaldehyde and CO upon desorption. The desorption of water with a peak at 155 "C during the temperature-programmed desorption is related to the disappear- ances of the two water bands chemisorbed on oxidized and reduced sites (3480 and 3580 cm-l) and the isolated hydroxyl band (3685-3705 cm-l) in fig. 1. The hydroxyl bands at 3375 and 3255 cm-l in fig.1 correspond to the desorption of water with a peak at 265 "C in fig. lO(b). Fig. 11 shows the amounts of the three different kinds of chemisorbed methanol as a function of adsorption temperature. The total amount of methanol measured by the infrared optical densities is made to agree with that found by mass spectrometry at 30 "C. The agreement at higher temperatures is good. The total irreversible methanol (t.i.m.) obtained by mass spectrometry is the sum of methanol desorbed at ca. 110 "C and the methanol desorbed as formaldehyde and CO (DM) in fig. 1O(a) and (b). The t.i.m. obtained by i.r. spectroscopy is based on the optical density of the CH stretching bands after a switch to helium at the adsorption temperature. However, using the same calibration for the optical densities, the methanol remaining adsorbed after a helium purge (t.i.m.) measured by the infrared method appears to be much less than that measured by mass spectrometry.The smaller t.i.m. measured by i.r. spectroscopy than that measured by the volumetric method probably indicates that, after purging with helium, a part of the methoxy held irreversibly transforms into other kinds of hydro- carbons which were not detected by infrared spectroscopy. The distance between the curve of the total amount of methanol and that of the t.i.m. is the amount of methanol absorbed reversibly at the adsorption temperature. This decreases rapidly with temperature and becomes zero at temperatures > 130 "C. This agrees with the results in fig.9, where the shifts in the band positions of CH stretching bands after a switch from methanol to helium become smaller as theJ. S. CHUNG, R. MIRANDA AND C. 0. BENNETT 35 0 30 50 70 90 T / T Fig. 11. Amounts of three different types of methanol on MOO, as a function of temperaure. Circles, volumetric; squares, i.r.; 0 and m, total amount; 0 and 0, tirn., total amount of irreversible chemisorption left on the surface after purging with helium; @, DM, the amount desorbed as formaldehyde and CO during temperature-programmed desorption. temperature increases and are negligible at 110 "C. Therefore the reversible methanol desorbed below 110 "C in fig. 10(a) corresponds to the methanol dissociated into methoxy and hydrogen across Mo=O (form C) and reaction (4) is reversible.This type of chemisorption must have a weak bond strength between methanol and MOO, and compiicated hydrogen bonding is expected. The extent of hydrogen bonding decreases with increasing temperature, as shown by the increase in the frequency of the CH stretching bands in fig. 9 up to 110 "C. The desorption of methanol with a peak at 110 "C is probably related to the methoxy chemisorbed on acidic V,, and the methoxy which is decomposed into formaldehyde or CO is related to the methoxy on the less acidic V, and/or V, sites. The chemisorption behaviour of methanol above 150 "C is quite different from that below 110 "C because of the reduction of MOO,. This will be discussed in another paper. CONCLUSIONS This work has led to the following conclusions.(1) Chemisorption of water or methanol develops three hydroxyl bands and two water bands. All except the band at 3255 cm-l are shifted to higher frequencies upon reduction of MOO,. (2) The isolated hydroxyl band and the two water bands disappear at ca. 150 "C, while the hydrogen-bonded hydroxyl band remains up to 350 "C. ( 3 ) Most of the methanol chemisorbs dissociatively and at least three fonns of chemisorbed methanol exist : methanol dissociated into methoxy and hydrogen cross Mo=O (form C), methoxy on the oxygen vacancy of Mo=O (form A) and methoxy on the oxygen vacancy of Mo-0-Mo (form B). (4) The bond strength between methoxy and MOO, decreases in the order: form B > form A > form C. ( 5 ) Form C reversibly desorbs as methanol36 CHEMISORPTION OF H,O AND CH,OH ON MOO, below 110 "C and a part of form A is desorbed as methanol above 110 "C.The rest of form A and form B desorb as formaldehyde and CO at ca. 240 "C. (6) There exist some oxygen vacancies capable of forming methoxide even after MOO, is oxidized at 380 "C overnight. We thank the National Science Foundation for support under grant no. CPE-7091018. G. K. Boreskov, Proc. 3rd Znt. Congr. Catal. (Elsevier, Amsterdam, 1964), p. 213. N. Pernicone, F. Lazerrin, G. Liberti and G. Lanzavecchia, J. Catal., 1969, 14, 293. J. Novakova, P. Jim and V. Zavadil, J. Catal., 1970, 17, 93. N. Evmenko ind Ya. Gorokhovatskii, Kinet. Katal., 1969, 10, 1299. R. Mann and K. Hahn, Znd. Eng. Chem., Process Des. Dev., 1970, 9, 43. J. Edwards, J. Nikcolaidis, M. B. Cutlip and C. 0. Bennett, J. Catal., 1977, 50, 24. M. Niwa, M. Mimtani, M. Takahashi and Y. Murakami, J. Catal., 1981, 70, 14. C. J. Machiels and A. W. Sleight, J. Catal., 1982, 76, 238. N. Pernicone. F. Lazerrin and G. Lanzavecchia, J. Catal., 1968, 10, 83. P. Jiru, M. Krivanek, J. Novakova and B. Wichterlova, Proc. 4th Znt. Congr. Catal. (Nauka, Moscow, 1969), p. 1919. l1 M. Formenti, F. Juillet, P. Meriaudeau, S. J. Teichner and P. Vergon, J. Colloid Interface Sci., 1972, 39, 79. l2 (a) C. 0. Bennett, ACS Symp. Ser., 1981, 178, 1; (6) J. S. Chung, Ph.D. Dissertation (University of Connecticut, 1984). l3 D. Bianchi, J. L. Bernard, M. Camelot, R. Bernali-Chaoui and S. J. Teichner, Bull. Soc. Chim. Fr., l4 J . R. Gunter, J. Solid State Chem., 1972, 5, 354. l5 M. Primet, P. Pichat and M. Mathieu, J. Phys. Chem., 1971, 75, 1216. 1980, 7-8, 1-275. C. H. Rochester and S. A. Topham, J. Chem. Soc., Faraday Trans. I , 1979,75, 1073. G. Anderson and A. Magnelli, Acta Chem. Scand., 1950, 4, 793. A. Zecchia, S. Coluccia, E. Guglieminotti and G. J. Ghiotti, J. Phys. Chem., 1971, 75, 2774. le L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, New York, 1966), p. 285. 2o R. T. Jernigan and G. R. Dobson, Inorg. Chem., 1972, 11, 81. 21 P. G. Dickens, J. J. Birtill and C. J. Wright, J. Solid State Chem., 1979, 28, 185; 1979, 29, 367. 22 G. Busca and V. Lorenzelli, J. Catal., 1980, 66, 155. 23 R. G. Greenler, J. Chem. Phys., 1962, 37, 2094. 24 B. A. Morrow, J. Chem. Soc., Faraday Trans. I , 1974, 70, 1527. 25 J. C. McManus, K. Matsushita and M. J. D. Low, Can. J. Chem., 1969,47, 1077. 26 R. 0. Kagel and R. G. Greenler, J. Chem. Phys., 1968,49, 1638. 27 E. W. Thornton and P. G. Harrison, J. Chem. Soc., Faraday Trans. 1, 1975,71,2468. 28 A. A. Davydov, V. M. Shchekochikhin, P. M. Zaitsev, Ya. M. Shchekochikhin and N. P. Keier, 29 K. Miranda, J. S. Chung and C. 0. Bennett, Proc. 8th Znt. Congr. Catal. (Springer, Berlin, 1984) 30 N. Takezawa and H. Kobayashi, J. Catal., 1973, 28, 335. Kinet. Katal., 1971, 12, 694. p. 111-347. (PAPER 4/088)

ISSN:0300-9599    

DOI:10.1039/F19858100019

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985, 81, 3 7 4 8 Electron Spin Resonance Spectroscopy of Surface Species formed upon Adsorption of Nitrogen Oxides and Oxygen on High-surface-area NiO-MgO and Coo-MgO Solid Solutions BY VALERIO INDOVINA, * DANTE CORDISCHI, STEFANO FEBBRARO AND MANLIO OCCHIUZZI Centro di Studio del CNR su ‘Struttura ed Attivita Catalitica di Sistemi di Ossidi’, Department of Chemistry, University of Rome, Rome, Italy Received 27th February, 1984 Solid solutions of high-surface-area (h.s.a.) NiO-MgO (Ni content 0.1-5 atoms per 100 Mg atoms) were activated by heating under vacuum at 1173 K for 5 h before exposure to 0,, NO,, NO or N,O at temperatures in the range 298-815 K. With all gases, depending on the temperature of adsorption, two species were observed by electron spin resonance spectroscopy.Both are axial signals of nickel species on the surface of MgO, as shown by broadening experiments with 0,. The e.s.r. analysis, the chemical conditions under which the species are formed and their thermal stability allow the assignment of the two species to (i) a coordinatively unsaturated Ni3+ (Ni:+) and (ii) Ni3+ ions in a distorted octahedral surface complex (Ni:+ * * L). Adsorption of NO at 178-298 K produces a (Ni . . . NO),+ adduct, NO:- and Mg2+. * * NO. Solid solutions of h.s.a. Coo-MgO (Co content 0.1-5 atoms per 100 Mg atoms) activated as for NiO-MgO were exposed to NO, or NO. The labile cobalt nitrosyl adduct formed is tentatively assigned to Co2+. .(NO),; NO;- species and Mg2+- * .NO are also formed. Some total and reversible adsorptions (removed by evacuation at the temperature of adsorption) were determined volumetrically.Reflectance spectra were taken under the same conditions for the formation of the paramagnetic species. Surface processes by which the various species were formed are discussed. Solid solutions of transition-metal ions in oxide matrices are useful model catalysts for the investigation of fundamental aspects of heterogeneous catalysis. Simple reactions such as the decomposition of N,O, the CO+O, reaction and H,-D, equilibration have been studied on high- and low-surface-area Coo-MgO ~atalysts.l-~ NiO-MgO samples have also been For an understanding of the catalytic behaviour, we need to know (i) the electronicconfiguration of the active transition-metal ion and (ii) the nature and the reactivity of the species formed on the surface of the catalysts upon adsorption of reagents and products.With this aim in mind, we have previously studied the adsorption of N,0,4 0,8 and on h.s.a. Coo-MgO solid solutions. In this paper we present the data for the adsorption of NO,, NO, N,O and 0, on h.s.a. NiO-MgO and for the adsorption of NO, and NO on h.s.a. Coo-MgO. The results for the two systems are complementary, since the surface species which cannot be detected by electron spin resonance spectroscopy in the case of cobalt can be detected in the case of nickel and vice versa, because of the electronic structure of the two ions. Evidence is presented for the 3 + oxidation state, which has often been assumed to be formed in the elementary steps of oxidation reactions on catalysts containing Co or Ni.3738 E.S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO Investigation of the adsorption by volumetric techniques allows evaluation of the significance of the e.s.r. observations with respect to the total amount of chemisorbed gas. EXPERIMENTAL MATERIALS The h.s.a. NiO-MgO solid solutions were prepared by impregnation of magnesium hydroxide with a solution of nickel nitrate (Erba, RP). The magnesium was the same as used before to prepare high-purity MgO.* The samples were decomposed slowly by raising the temperature to 800 K before carrying out the final activation at 1173 K for 5 h under a pressure of < lo-, N m-,. This procedure gives a good dispersion of nickel ions in solid solution of MgO.” The h.s.a.COO-MgO solid solutions were prepared as described for NiO-MgO. Details of the preparation and characterization are given in ref. (8). All specimens are designated as MN or MC followed by a number giving the transition-metal content as atoms per 100 Mg atoms. B.E.T. surface areas (rn, g-l), determined by N, adsorption at 77 K, are indicated in parentheses after each catalyst: MN 0.1 (253), MN 1 (215) and MN 5 (193); MC 0.1 (213), MC 1 (200) and MC 5 (170). NO, N,O, NH, and CO, were purified by double distillation under vacuum. High-purity 0, (Air Liquide, 99.95%) was used. NO, was prepared before each experiment by reacting NO with an excess of 0, and by subsequent distillation under vacuum. All gases were dried before admission to the adsorption chamber.APPARATUS AND PROCEDURE A weighed amount of the solid solution (ca. 0.1 g) was placed in a silica bulb equipped with a side e.s.r. tube and activated under vacuum at 1173 K, as specified above, before exposure to the various gases. For the adsorption measurements, the sample was contacted with the gas at a pressure chosen so as to have in all cases a final pressure of ca. 600 N m-,. Pressure readings were made with a pressure transducer (MKS, Baratron) capable of detecting variations of 0.01 N m-,. The adsorption was considered complete when two successive readings at 5 min intervals did not differ by > 0.5 N rn-, (hereafter called ‘total adsorption’). The samples were subsequently evacuated for 10 min at the same temperature and the amount adsorbed (always at the same temperature) gave a measure of the amount of gas which could be easily desorbed (‘reversible adsorption’).Occasionally, after NO or N,O adsorption, the gas phase was analysed by mass spectrometry (Micromass 601, VG). The e.s.r. spectra at 77 or 298 K were recorded at X-band frequencies on a Varian E-9 spectrometer. The absolute number of spins was determined from electronically integrated spectra using Varian ‘strong pitch’ as standard [(3 & 1) x lo1’ spin m-l]. The u.v.-visible reflectance spectra were recorded at room temperature with a Beckman DK I-A spectrometer using a silica activation chamber with a side cell equipped with an optical window. Activations and adsorptions were carried out in situ. RESULTS NiO-MgO SOLID SOLUTIONS E.S.R.SPECTRA The e.s.r. spectrum observed for activated MN samples consists of an isotropic signal at g = 2.21 due to Ni2+ ions in octahedral sites.12 Because a large fraction of Ni2+ ions depart from exact cubic symmetry, in these high-surface-area materials the linewidth of the e.s.r. signal is larger (AH,, = 1500 G) than that measured for low-surface-area MN solid solutions (80-1 60 G). l2 Upon adsorption of NO,, NO, N,O or O,, the formation of various paramagnetic species is observed (see table 1). For each species table 1 also reports the g values. UponV. INDOVINA, D. CORDISCHI, S. FEBBRARO AND M. OCCHIUZZI 39 Table 1. Surface species on high-surface-area NiO-MgO samples, as detected by e.s.r. spectroscopy gas adsorbed species g values NO, NO N,O 0, adsorption of a given gas, the formation or not of a species is specified in table 1 by the symbol + or -, respectively. For each gas the temperature of adsorption and pressure under which the various e.s.r.signals were observed will be specified in the next section, where the concentration of the radicals will also be reported. E.s.r. signals of the species NO;- and Mg2+. -NO (table 1) were observed by Lunsford13 upon adsorption of NO on thermally activated MgO. Since the intensities of the e.s.r. signals of these two species are the same on MN samples as on pure MgO, no specific role can be envisaged for the nickel ions in their formation. The e.s.r. signal of the 0; species on the surface of pre-irradiated MgO (Mg2+ - - * O;, table 1) has also been reported previ0us1y.l~ However, in this case a specific role exists for the transition-metal ion since the 0; species are not formed on thermally activated Mg0.15916 The e.s.r.spectra of the species designated in table 1 as A, B and C are reported in fig. 1. Spectra (a) and (c) correspond to species A and C, respectively. Spectrum (b) consists of two signals, namely species A and B. Species A and B, when occurring together, can be distinguished since the signal of species A undergoes reversible and complete broadening upon exposure to oxygen whereas that of species B is only slightly broadened. The g values of the three axial signals are reported in the stick diagram of fig. 1 . FORMATION, CONCENTRATION AND STABILITY OF SPECIES A, B AND C Species C is considered first.This is formed by adsorption of NO on MN samples at 298 K or lower temperatures (down to 178 K). The intensity of the signal, which depends on the pressure of NO, reaches a maximum of 40 x 10l5 spin rn-, on MN 1 and 6.5 x 1015 on MN 0.1 at a pressure of 2 kN m-,. Species C is very labile and its e.s.r. signal disappears on evacuation for 10 min at 298 K. It is useful to compare this amount with the amount adsorbed as determined by volumetry on the MN 1 sample. The total adsorption at 298 K is 790 x 1015 molecule rn-, and includes both irreversible (NO;-) and reversible forms, the latter amounting to 180 x 1 OI5. Thus, theconcentration of the radical is less than the amount of NO reversibly chemisorbed at 298 K. Species A and B are formed with NO,, NO, N,O and 0, at the temperatures specified in table 2.For each gas, table 2 also reports for MN 1 and MN 0.1 the concentration of radicals, species A and €3, as a function of the temperature. The results show that species A and B are formed with NO, even at 298 K. Progressively higher temperatures are required for NO, N,O and O,, in that order. Both species are stable at 298 K, i.e. the intensities of their e.s.r. signals are unaffected by evacuation40 E,S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO .-i 2.36 2 . 0 2 species A species B 2.25 2.11 2.28 2.15 species C - - Fig. 1. E.s.r. spectra at 77 K on an MN 1 sample activated by heating under vacuum at 1173 K and exposed to: (a) 0, at 723 K, (b) NO, at 298 K and (c) NO at 298 K. The group of lines marked with an asterisk belongs to the Mg2+- * .NO and NO:- species.Table 2. Concentration of species A and B obtained upon adsorption of NO,, NO, N,O and 0, on NiO-MgO at various temperatures concentration/ spin m-2 NO2 NO N2O 0 2 T / K A B A B A B A B sample MN 1 298 70 50 466 572 717 - - 815 MN 0.1 298 6.6 8.4 572 - - - - - - - - 2.2 0.0 0.0 9.0 2.5 3.0 2.0 2.6 25.0 14.0 0.0 5.0 - - - - - - - - - 0.0 0.0 0.0 0.8 0.0 0.0 0.0 0.2 0.0 4.9 3.2 0.9 4.0 0.9 0.0 - - 4.0 - - - at this temperature. The concentration of species A + B is 4-8 times larger on MN 1 than on MN 0.1. The thermal stability of species A and B was further investigated as follows. The two species were formed on an MN 1 sample in two different ways: (i) by exposure to NO, at 298 K and (ii) by exposure to N,O at 572 K.The samples were thenV. INDOVINA, D . CORDISCHI, S. FEBBRARO AND M. OCCHIUZZI 41 0-4P 100 o--o--- t \ 8ol 2o t O\ o+o 1 , , , , , Q \ 0 0= T/K 273 473 673 873 Fig. 2. Thermal stability of species A and B, obtained either by exposure to N,O at 573 K (0 and 0, respectively) or NO, at 298 K (A and V, respectively). The samples were subsequently heated under vacuum at progressively higher temperature. Intensities are reported in arbitrary units as a percentage of the initial intensity (= 100). Table 3. Concentration of species A and B after adsorption of CO, or NH, concentration/ spin mP2 treatmenta A B A+B N,O at 573 K 25 10 35 N,O at 573 K 22 1 1 33 NH, at 298 Kb 18 16 34 CO, at 298 Kb 12 28 40 a Before exposure to N,O the MN 1 sample was activated by heating under vacuum at 1173 K.Before exposure to CO, or NH, the sample was outgassed at 298 K. evacuated at progressively higher temperatures, recording the e.s.r. spectra after each step. In fig. 2, the intensities of the signals for species A and B, reported as a percentage of the initial intensity (= loo), are plotted as a function of the evacuation temperature. Both species possess the same thermal stability and are stable up to ca. 573 K. At 673 K the original intensities are reduced by one-half and at 780 K the signals are destroyed. In other experiments, MN 1 samples containing a known amount of species A and B were exposed to CO, or NH, at a pressure of 13.5 kN m-2. The treatment provokes a decrease of the concentration of species A and an increase of that of species B, while the total concentration of radicals remains constant (table 3).OXYGEN ADSORPTION The total adsorption of 0, on MN I , activated by heating under vacuum at 1 173 K, was as indicated in parentheses at the various temperatures: 298 K (7 x 1015 molecule m-2), 428 K (19 x 1015), 563 K (22 x and 717 K (30 x 1015). At 298 K, the42 E.S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO amount reversibly adsorbed was 5 x 1015 molecule m-2 and the amount irreversibly adsorbed was 2 x 1015. Species A and B were not formed at 298 K (table 2). The 0; species was observed at low concentration (0.9 x 1015 spin rn-,). Thus, oxygen is mostly chemisorbed in diamagnetic forms : probably 0;- for the reversible adsorption and 02- for the irreversible adsorption.The MgO matrix does not chemisorb oxygen at 298 K (< 0.5 x 1015 molecule m-2). REFLECTANCE SPECTRA The spectrum obtained for an MN 5 sample after activation under vacuum at 1 173 K consists of five bands. These have been observed by other authors and assigned to bulk Ni2+ ions (8300, 14000 and 23000 cm-l) and surface Ni2+ probably in C4v symmetry (5000 and 23000 cm-l).ll* l7 After adsorption of N,O, NO, and 0, a broad and intense band is observed. This disappears upon evacuation at ca. 870 K. A similar band was observed by Hagan et al.ll on h.s.a. MN samples following adsorption of 0,. Note that the broad adsorption is observed under the same conditions for the formation of species A and B and, moreover, the temperature at which the absorption band disappears is not far from that at which species A and B are destroyed.COO-MgO SOLID SOLUTIONS E.S.R. SPECTRA Upon adsorption of NO (6 kN m-,) on MC 1 or MC 5 samples activated by heating under vacuum at 1173 K for 5 h, intense e.s.r. signals are observed (fig. 3). These consist of a broad band at 7 = 2.12 showing a poorly resolved hyperfine structure arising from interaction v ;ch the cobalt nucleus [MC 1, spectrum (a)]. In addition, as in the case of NO adsorption on MN, the e.s.r. signals of species adsorbed on sites of the MgO matrix can be recognized: NO:- and Mg2+. * .NO. The species is very labile and is destroyed by a few minutes evacuation at 298 K. Thus, the stability is the same as for species C of the MN system. The intensity of the signal for the MC 1 sample corresponds to 23 x 1015 spin rn-, at 298 K and 58 x 1015 at 178 K.The total adsorption of NO is 930 molecule m-, at 298 K and the reversible adsorption 370 molecule md2 at the same temperature (MC 1). The reasons for the NO radical concentration being smaller than the reversible adsorption of NO are as reported above for the MN system. The maximum radical concentration on the MC 5 sample is 1.0 x 1017 spin m-,. ADSORPTION OF 0, AND N,O The adsorption of 0,8 and N,O has been previously investigated in our laboratory on h.s.a. MC samples. Upon adsorption of O,, the most significant results, recalled here in order to underline some important similarities with the MN system, concern the formation of Co3+ - * 0; (for this species, up to 1 .O x 10" spin m-, were detected) and that of 0; and 0; species chemisorbed on sites of the matrix.The total adsorption of oxygen, as determined volumetrically, was larger than the concentration of radicals. Therefore, 0;- and 0,- are also thought to be present on the surface.8 Adsorption of N,O leads to the formation of 0-, 0; and 0; chemisorbed on sites of the matrix and to the formation of Co3+ - - -0;. The formation of 0- species suggests an intermediate species Co3+. - -O-, which was not detected as it is very labile .4V. INDOVINA, D . CORDISCHI, S. FEBBRARO AND M. OCCHIUZZI 500 G 1 43 -I- 2.007 - 1.996 1.89 c o 1 2 12 A = 94G Fig. 3. E.s.r. spectrum at 77 K of (a) MC 1 and (b) MC 5 heated under vacuum at 1173 K and exposed to NO. The lines at g, = 1.996 and g,, = 1.89 are due to the Mg2+.* . NO species and that at g, = 2.007 to NO:-. REFLECTANCE SPECTRA The reflectance spectrum of an MC 5 sample activated under vacuum at 1173 K was nearly identical to that previously reported for similar samples by Hagan et all1 The broad background absorption, obtained by exposure to 0, at 298 K,ll is also detected upon adsorption of N,O and NO, both at 298 K. Evacuation at 298 K does not affect the spectra. Evacuation at progressively higher temperatures shows that the broad absorption disappears at a temperature in the range 723-850 K. DISCUSSION THE NATURE OF SPECIES A AND B The e.s.r. axial signals of species A and B are due to different forms of nickel with spin 1/2. The two species are on the surface of the MgO matrix, as shown by the fact that their e.s.r.signals are reversibly broadened upon addition of 0,. In principle, two oxidation states are possible: Ni+ or Ni3+. In the following it will be shown that species A and B are surface Ni3+ ions. The assignment relies on both spectroscopic and chemical evidence. Spectroscopic evidence will be considered first. In most cases, relying on g values alone it is rather difficult to distinguish between the low-spin d7 configuration (Ni3+) and the d9 configuration (Ni+). In fact, the orbitally degenerate ground state (2E2s)44 E.S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO is not split in a cubic or trigonal crystal field, but the degeneracy is removed by the Jahn-Teller effect (either static or dynamic). Thus, for both configurations an axial spectrum with g values in the range 2-2.4 is expected.However, the effect of the tetragonal distortion on the principal values of the g tensors is opposite for the two configurations. In particular, for an elongated octahedron gll < gl is expected for a low-spin d7 ion and gI1 > g , for a d9 ion; the reverse is predicted for a compressed octahedron. Thus, given that in the present case gll < g,, the only alternative left is between Ni3+ ions in an elongated octahedron and Ni+ in a compressed octahedron. The latter configuration can be ruled out since we have substantiated that species A and B are on the surface where a compressed octahedral form appears to be very unlikely. Chemical arguments in favour of Ni3+ rather than Ni+ arise from an inspection of (a) the conditions in which species A and B are formed and (b) their thermal stability.In particular, oxidation of the surface Ni2+ ions of the MN solid solutions is expected upon exposure to NO,, NO, N20 and 0,. More specifically, oxygen atoms appear to be involved in the oxidation of Ni2+ since mass-spectrometric analysis shows the formation of N, when NO or N,O are contacted with the MN samples. In other words, NO and N,O decompose during the formation of species A and B. Analogously, with O,, oxygen atoms arise from the cleavage of the 0-0 bond during the activation of molecular oxygen. As far as the thermal stability of species A and B is concerned, this is found to be nearly identical to that previously observed for Ni3+ ions present in the bulk of MgO-NiO-Li,O solid solutions, in which the Ni3+ oxidation state was induced by the presence of lithium?, The significance of the comparison relies on the fact that for both systems (h.s.a.MN and MgO-NiO-Li,O) the decay of the Ni3+ signal (either bulk or surface) is related to desorption of oxygen from the surface of the samples. Having established that species A and B are two different Ni3+-containing species, the next step toward the identification of species A and B is as follows. The anisotropy of the g values of species A and B cannot be fully explained by using the equations given by Lacroix et al. for low-spin d7 ions undergoing a Jahn-Teller effect.’* The equations have been successfully used to assign isotropic or axial signals observed in several matrices for bulk Ni3+ ions undergoing either static or dynamic Jahn-Teller In the case of species A and B, Ni3+ ions are in an octahedral field with strong axial distortion, the cubic field being near to the cross-over region between the ,EB and the 4qg states (fig.4). A similar situation has been discussed previously for other octahedral Ni3+ complexes by Reinen et aL21 The analysis, which applies to any symmetry with a quaternary axis, givesgil and g,values as functions of S,/r and S,, ?/<, in which r is the spin-orbit coupling constant (fig. 4). By following this analysis it is found that the g values of species A are fitted by S2/t = 6 and = 2.45 and those of species B by S,/t = 0.7 and S2.4/t = 3.30. Thus, the S, parameter, which increases with increasing axial field, is higher for species A than for species B, suggesting that species A consists of coordinatively unsaturated Ni3+ (Ni;+) and species B of surface Ni3+ in an octahedral distorted complex (NiE+ - * - L).Again, the assignment is consistent with the chemical behaviour observed upon exposure to NH, or CO, ; i.e. NiE+ is transformed into Ni;+ - * L by adsorption of either NH, or CO, schematically, Nig+ + L --+ NiE+ - - L). In the case of N,O, NO or 0, adsorption, the ligand L is thought to be 02- (Ni3+- * -02-). The fact that the e.s.r. parameters of species B are largely independent of the particular L implies that the axial component of the crystal field is of the same order of magnitude for the various ligands. This is possible since ligands such as those used here have been shown to be linked to the surface by bonds whose strength is comparable to that of the 02-ions of the lattice.,,V.INDOVINA, D. CORDISCHI, S. FEBBRARO AND M. OCCHIUZZI 45 2 / Eg / Fig. 4. Splitting of the lowest terms, zEg(r&,ei) and 4Tg (rig,ei), in a strong tetragonal field near to the cross-over region. Finally, evidence has been given for the previous assignment of the broad band in the reflectance spectra of NiO-MgO solid so1utions.l1Vl7 In fact, species A and B, whose behaviour parallels that of this reflectance band, have been shown to consist of Ni3+ ions. NITROSYL ADDUCTS OF Ni AND Co Surface species formed on MN (species C) and MC (broad e.s.r. signal at g = 2.12) upon adsorption of NO have the typical behaviour of surface nitrosyl complexes, such as high lability. For species C, gll > gl.The situation is the opposite to that observed for species A and B and, therefore, by using the same arguments given above for species A and B, species C can be assigned to a surface complex of a nickel ion in the dg configuration, i.e. to a Ni+ complex. Accordingly, the simplest formulation of the complex would be Nil * -NO+. This formulation has been adopted previously for mononitrosyl Ni adducts in zeolites, with g values similar to those found here for species C.23- 24 The assignment also relies on i.r. spectra which showed the typical bands of nitrosyl ligands (ca. 1800 cm-l). On the other hand, the e.s.r. spectra of Ni adducts in zeolites have been analysed using the equations: These apply to a dg ion in a tetragonal elongated octahedron. On the basis of simple crystal-field theory a ratio Ail/Al < 1 is expected, as found in the case of Ni+ in various However, as shown in table 4, the All/A1 ratios calculated from the g values of various nitrosyl adducts are greater than unity.The discrepancy can be attributed to the fact that the Ni-NO bond is, to some extent, covalent, thus accounting for the apparent failure of crystal-field theory. Thus, in the absence of hyperfine structure in the e.s.r. signal of species C and of a more detailed description of the structure of the nickel adduct, it is better to formulate species C as (Ni. * Coming now to the broad band showing hyperfine interaction with the Co nucleus, 26 and also for several complexes of rather than as Ni+.* .NO+.46 E.S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO Table 4. Spectroscopic parameters of Ni nitrosyl adductsa system 811 g, 4 / A l b ref. NP-Y zeolite 2.33 2.169 2.04 24 NiII-Y zeolite 2.34 2.171 2.00 23 NiO-MgO 2.28 2.15 2.14 this work Ni+ in NaF 2.766 2.1 14 0.59 25 Ni+ in SrF, 2.597 2.092 0.60 26 a The data are compared with those obtained for Ni+ in the matrix of NaF and SrF,. A/\/AL = 4kl-ge)/kl/ -ge)* Table 5. Surface species observed by e.s.r. on h.s.a. NiO-MgO and COO-MgO solid solutions on adsorption of 0,, N,O and NO surface species gas adsorbed NiO-MgO CoO-MgOa a The data with 0, and N,O for the COO-MgO system are taken from ref. (4) and (8). this cannot be due to the analogous mononitrosyl adduct since this, in the case of Co, is diamagnetic or of even spin.Possible paramagnetic complexes are : Co+ - * NO, Co3+ - - - NO and Co2+ - - - (NO),. A tentative assignment to Co2+ - - * (NO), is suggested by (i) the known tendency of Co to coordinate more than one molecule1* and (ii) the fact that the adsorption of NO is higher on MC than on MN (by a factor of ca. 2). Moreover, no evidence for a Co3+ - - - NO species was found in a special experiment performed by adsorbing NO on an h.s.a. MC doped with Li in order to induce the formation of Co3+: the intensity of the e.s.r. signal was the same as obtained after NO adsorption on a normal MC sample. SURFACE REACTIONS The mechanisms of 0, adsorption and N20 decomposition have been discussed previously in the case of COO-MgO solid ~ o l u t i o n s .~ ~ ~ The data presented here for the related NiO-MgO system, along with the additional information for COO-MgO, provide further evidence for the mechanisms reported previ0usly.~9 In particular, the T3+ and T3+ * - 0,- which could not be observed when T was cobalt have now been detected with the nickel system. The reverse is true for the T3+. -0; species which was observed in the case of cobalt but not in that of nickel. Table 5 lists the surface species formed on the MN and MC systems upon adsorption of 0,, N,O and NO. From a qualitative view point, the nature of the surface species points to a strict analogy between the two systems. However, differences exist as regards the temperature at which the surface species are formedV. INDOVINA, D.CORDISCHI, S. FEBBRARO AND M. OCCHIUZZI 47 and their stability. These differences can be used to explain the higher catalytic activity of COO-MgO as compared with NiO-MgO for the decomposition of N2028 and the CO+O, reaction.' This aspect will be discussed elsewhere. Note, however, that the absence of 0- and 0, species on the MN surface (0, only being observed, table 5) can be accounted for by the higher temperature at which the N,O decomposition is carried out on the MN catalysts as compared with the MC catalysts. In previous work from our group the thermal stability of 0- and 0; species was found to be lower than that of O;(O- < 0; < O T ) . ~ ~ ~ Coming now to the adsorption of NO, the surface species observed can be accounted for by the following reactions. In particular, whereas the elementary steps for the formation of Nii+ and Ni:+ * .* 02- cannot be given, the overall process can be shown by: 2 NO(g) + 2 NiE+ + 0:- + Ni:+ + Ni;+ - - * 0,- + O& + N,(g). The formation of NO;-, (Ni - - - NO),+ and Mg2+ - - NO is given by : NO(g)+Oi-+NO;- (1) NO(g) +NiE+ + (Ni * - - NO),+ (2) NO(g)+Mgi+ + Mg2+- - *NO. (3) The analogous surface reactions (1)-(3) are also observed for Coo-MgO. Note that oxygen species (0-, 0; and 0;) are not seen upon adsorption of NO on both systems. A possible cause is that once formed these species react with gas-phase NO, yielding NO; and NO;. In agreement with this hypothesis, the concentration of radicals is found to be smaller than the total adsorption of NO. CONCLUSIONS Two main conclusions may be drawn from these results: (i) Ni3+ and Co3+ ions are formed on the surface of NiO-MgO and COO-MgO solid solutions upon adsorption of the appropriate gases and (ii) nearly all the Ni2+ and Co2+ ions on the surface can be oxidised to Ni3+ and Co3+, respectively.The first conclusion follows from the e.s.r. evidence and the second can be substantiated as follows. A surface analysis by X.P.S. of the MN 29 and MC 30 systems has shown that the surface composition is only slightly different from that of the bulk. Therefore, the surface concentration of Ni2+ or Co2+ is 1.2 x 1017 atom m-, in MN 1 and MC 1 samples and 0.12 x 1017 in the MN 0.1 sample. Comparison of these surface compositions with the concentrations of paramagnetic species leads to conclusion (ii).In fact, the maximum concentration of Ni:+ + Nii+ - - L was 1.2 x 1017 spin m-, on MN 1 and 0.15 x 1017 on MN 0.1 ; the concentration of Co3+ - - - 0; was 1 .O x 10'' spin rnP2 on MC 1. A. Cimino and F. Pepe, J. Catal., 1972, 25, 362. V. Indovina, A, Cimino and F. Pepe, Gazz. Chim. Ztal., 1980, 110, 13. V. Indovina, A. Cimino, M. Inversi and F. Pepe, J . Catal., 1979,56, 396. 75, 21 77. A. Cimino, R. Bosco, V. Indovina and M. Schiavello, J . Catal., 1966, 5, 271. A. Cimino, V. Indovina, F. Pepe and M. Schiavello, J . Catal., 1969, 14, 49. V. Indovina, A. Cimino and F. Pepe, 9th. Iberoamerican Symposium on Catalysis, 1984, Lisbon, Portugal. * V. Indovina, D. Cordischi, M. Occhiuzzi and A. Arieti, J. Chem. SOC., Faraday Trans. 1 , 1979,48 E.S.R. STUDY OF NO, AND 0, ON NiO-MgO AND COO-MgO D.Cordischi, V. Indovina, M. Occhiuzzi and A. Arieti, J. Chem. SOC., Faraday Trans. I , 1979, 75, 533. D. Cordischi, V. Indovina and M. Occhiuzzi, J. Chem. Soc., Faraday Trans. I , 1980, 76, 1147. lo V. Indovina, D. Cordischi and M. Occhiuzzi, J. Chem. Soc., Faraday Trans. I , 1981, 77, 81 1. l1 A. P. Hagan, M. G. Lofthouse, F. S. Stone and M. A. Trevethan, Preparation of Catalysts II, ed. B. Delmon, P. Grange, P. Jacobs and G. Poncelet (Elsevier, Amsterdam, 1975), p. 417. l2 A. Cimino, D. Cordischi, G. Guarino and A. Micheli, Trans. Faraday Soc., 1971,67, 1776. l3 J. H. Lunsford, J. Chem. Phys., 1967, 46, 4347. l4 A. J. Tench and P. Holroyd, Chem. Commun., 1968,471. l5 V . Indovina and D. Cordischi, Chem. Phys. Lett., 1976, 43, 485. l6 D. Cordischi, V. Indovina and M. Occhiuzzi, J. Chem. Soc., Faraday Trans. I , 1978, 74, 456. l7 M. Lo Jacono, A. Sgamelotti and A. Cimino, Z. Phys. Chem. N.F., 1970,70, 179. Is J. Tanaka, I. Shindo and M. Tsukioka, J. Phys. Soc. Jpn, 1980, 49, 120. 2o D. M. Hannon, Phys. Rev., 1967, 144, 366. 21 D. Reinen, C. Friebel and V. Propach, Z. Anorg. Allg. Chem., 1974,408, 187. 22 D. Shopov, A. Andreev, P. Pumpalov, J. Catal., 1970, 9, 398. 23 P. H. Kasai, R. T. Bishop Jr and D. McLeod Jr, J. Phys. Chem., 1978,82-83, 279. 24 C. Naccache and Y. Ben Taarit, J. Chem. Soc., Faraday Trans. I , 1973, 69, 1475. 25 J. W. Orton, Rep. Prog. Phys., 1959, 22, 204. 26 P. J. Alonso, C. J. Gonzales, H. W. Den Hartog and R. Alcala, Phys. Rev. B, 1983, 27, 2722. 27 P. A. Ayscough, Electron Spin Resonance in Chemistry (Methuen, London, 1967) p. 191, table 6.8. 28 A. Cimino, Chim. Ind. (Milan), 1974, 56, 27. 2s A. Cimino, B. A. De Angelis and G. Minelli, J. Electron Spectrosc. Relat. Phenom., 1978, 13, 291. 30 A. Cimino, B. A. De Angelis and G. Minelli, Surf. Interface Anal., 1983, 5, 150. R. Lacroix, V. Hochli and K. A. Miiller, Ber. Tag. Schweiz. Phys. Ges., 1964, 37, 327. (PAPER 4/334)

ISSN:0300-9599    

DOI:10.1039/F19858100037

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1985, 81, 49-60 Photophysics of the Excited Uranyl Ion in Aqueous Solutions Part 4.-Quenching by Metal Ions BY HUGH D. BURROWS,* AUGUSTO C. CARDOSO, SEBASTI~O J. FORMOSINHO* AND MARIA DA GRACA M. MIGUEL Chemistry Department, University of Coimbra, 3000 Coimbra, Portugal Received 5th March, 1984 The quenching of the excited state of the uranyl ion by metal ions in aqueous solutions has been studied under conditions where (UOi+)* decay is biexponential. The effect of metal ions always follows Stern-Volmer behaviour both for real quenching and for the second process, which is suggested to involve reversible crossing via solvent exchange between two energetically close excited states. With Tl+, Ag+, Fez+, Pb2+, Mn2+, Ce3+ and Ni2+ quenching is suggested to occur by electron transfer.Theoretical calculations using a quantum-mechanical tunnelling model support a mechanism involving an inner-sphere exciplex. With Eu3+ preferential quenching of the emitting U* state is suggested to involve enhanced hydrogen-atom abstraction by (UOi+)* following complexing and overlap of europium and uranyl orbitals. Initial fluorescence enhancement by metal ions in the uranyl system is observed in both dynamic and static studies. In cases where there is no complexing in the ground state with UOi+, this initial enhancement is interpreted in terms of an effect of the metal ion on the initial relaxation of the emitting (UOi+)* state. The time-resolved fluorescence spectra of UOi+, together with other data on the photophysics of (UOg+)*, suggest that the UOi' ground state and the lowest excited state, X*, have a slightly bent geometry (ca.176"), whereas the second excited state, U*, is linear. The quenching of excited uranyl ion by metal ions in aqueous solutions has been studied exten~ively.l-~ The mechanism of quenching with several ions involves an electron-transfer process, but energy transfer has also been invoked with E u ~ + , ~ where electron transfer is not energetically feasible. Marcus theory does not seem to be applicable to these electron-transfer proce~ses.~ It is suggested that this rules out an outer-sphere mechanism, and Marcantonatos has revealed that the quenching processes depend on the chemical ionization energy and on the spatial extension of the orbital containing the electron to be tran~ferred.~ In all these experiments, either in the steady-state or by dynamic studies, uranyl decay was considered to be a single exponential. However, (UOi+)* is known to undergo biexponential decay under certain experimental conditions and this has been attributed to a reversible crossing between two energetically close excited The states have a different electronic nature, n3,& and n3,4E, and the reversible crossing between the two states has been attributed to a solvent-exchange process.Since solvent can play a significant role in the mechanisms of quenching of (UOi+)* by metal ions, we have investigated such quenching processes under conditions where the decay of the excited uranyl ion is biexponential. When analysed in terms of a reversible-crossing mechanism, the decays permit the study of the effect of the metal ion on quenching and on solvent-exchange processes.Such studies have also helped elucidate the mechanism of enhancement of uranyl fluorescence (particularly in aged solutions) caused by ions such as Ce3+, Zn2+ and 3 49 FAR 150 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION EXPERIMENTAL Excited uranyl decays were studied using a nanosecond flash-photolysis apparatus with a pulsed N, laser (A,,, = 337 nm) at constant laser intensity. Fluorescence spectra were run on a Spex Fluorolog model 111 instrument and absorption spectra were obtained using a Shimadzu UV-240 spectrophotometer. Solutions were prepared with triply distilled water from uranyl nitrate (analytical grade).All metal ions were taken as their nitrate salts and were of analytical grade. RESULTS FLUORESCENCE QUENCHING Fluorescence decays were obtained for aqueous solutions of uranyl nitrate (0.02 mol dm-3, pH 2.8, T = 20 "C) in the absence and in the presence of varying concentrations of other metal nitrates under conditions where the decays were biexponential. Both components were observed to decay more rapidly with an increase in metal-ion concentration. The decays are analysed according to the following ki scheme : U * r X * The effect of metal ions always follows Stern-Volmer plots, k,+k2[Q], for the rate constants of irreversible decays (true quenching) and for the rates of reversible crossing. Nitrate ion has also some effect on the rates' but its effect is much smaller than that of the metal ions; NO; has the following quenching rate constants (from addition of NaNO,): kro; = 8.0 x lo4 dm3 mol-1 s-l; k,Noi = 4.5 x lo4 dm3 mol-l s-l.The rate k, decreases (1.6 x lo5 dm3 mol-1 s-l) up to [NO;] = 0.3 mol drn-,. Table 1 presents the rate constants for 'quenching' by T1+, Ag+, Fe2+, Pb2+, Mn2+, Ce3+, Ni2+ and Eu3+ together with those for autoquenching by UO;+.' The rate constants were estimated in terms of metal-ion concentrations and not in terms of activities because we have used low ionic strengths. Marcantonatos did not find a great difference between the two kinds of rate^,^ although for the less efficient quenchers the increase in ionic strength (p) may lead to a small increase ( x 2) in the estimated quenching rate constants.To compare the present results with other published data it is useful to estimate the average quenching rate constants. The rates are estimated from Stern-Volmer plots of the biexponential decays with rate constants k, and k, and amplitudes C, and Cz.5 The average quenching rate constant is obtained from the Stern-Volmer plots of the average decay constants kav = (C, k, + C, k,)/(C, + C2). Table 1 also presents these values, which are in good agreement with other published data,1-4 with the exception of Ce3+, which has a rate constant ca. 20 times higher than that obtained by steady-state measurements.2 This discrepancy can be attributed to an enhancement in fluorescence which will be discussed later. Acidity has been found to have virtually no effect on the autoquenching rates of uranyl within a pH range 2.0-4.0.The same seems also to be true for Ag+, which does not undergo hydrolysis within this pH r e g i ~ n . ~ Table 2 confirms also the small increase in the quenching rates with an increase in ionic strength at different pH values. FLUORESCENCE ENHANCEMENT Enhancements in the fluorescence intensities have previously been reported in steady-state measurements with aged solutions of uranyl ion and Ce3+, Zn2+ andH. D. BURROWS, A. c. CARDOSO, s. J . FORMOSINHO AND M. DA G. M. MIGUEL 51 Table 1. Stern-Volmer rate constants" for the quenching of the excited uranyl ion by metal ions at 20 "C and pH 3 ion k"4" T1+ Fez+ Pb2+ Mn2+ Ce3+ Ni2+ Eu3+ Ag+ uo;+ 1600 1700 390 24 2.5 1.6 1.1 1.3 10 1900 1400 3 70 22 2.0 1.8 1.5 4.0 ca.0.19 850 900 130 10 8.5 5.0 0.8 2.6 19 I100 700 150 15 6.0 4.5 0.6 7.3 ca. 0.4 3500 1600 590 14 17 6.0 0.76 35 " Determined as described in ref. (5) with an estimated error of &20% ; units lop6 dm3 mol-1 s-1. Table 2. Rate constants for quenching of (UO;+)* by Ag+ as a function of pH and ionic strength 1 O9 dm3 mol-l s-' P /mol dmP3 PH 2 PH 3 PH 4 - - 0.016 1.1" 1 .oc 0.06 1.6" 1.5b 1 .6b 1 .@ 0.1 2.1" 1.76 - 1 .7c 0.2 - 2.4c 0.355 2.2b - - - - - " M. D. Marcantonatos and M. Deschaux, Chern. Phys. Lett., 1980, 76, 359. Our results by dynamic studies. Our results by steady-state measurements. CO~+.~-* With Ag+ we have been able to observe an enhancement under steady-state conditions, but only at pH 4 and p 0.1 mol dmP3. At a lower pH only the normal quenching effect was observed.Attempts were made to detect complex formation at pH 3 by conductivity measurements. However, within experimental error (3 % ) there is no difference in conductivity between the sum of conductivities of the separated ions and that of the mixed solution, suggesting that the degree of complexing, if any, is not very big. Marcantonatos has also found negative deviations in the Stern-Volmer plots for the quenching of (UOi+)* by relatively high concen- trations of Ag+ lo and Ce3+4 under conditions where there are some changes in the absorption spectra of UOE+. We have also observed under dynamic conditions an increase (up to ca. 30%) in the initial fluorescence intensity of (UOi+)* with the addition of all the metal ions (T1+, Ag+, Fe2+, Mn2+, Ce3+ and Eu3+ were investigated).Since we are monitoring the fluorescence intensity at a particular wavelength (A = 485 nm), the apparent enhance- ment could be caused by a change in the fluorescence spectra accompanying the addition of metal ions. In order to elucidate this problem time-resolved emission spectra of (UOE+)* were recorded. As fig. 1 shows, the fluorescence spectrum of (UOi+)* does not change significantly in overall area, but suffers a red shift from 3-252 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION 460 4 50 500 A/nm Fig. 1. Time-resolved spectra of UOi+ fluorescence: 0, 0 and 0, 300 ns. 470 nm to 480 nm with time. This shift is much faster than the overall excited-state decay and occurs with a lifetime of 200 f 30 ns for [UOi+] = 0.02 mol dm-3.In the presence of metal ions the emission spectra seem to be relaxed within 100 ns (the time resolution of our detection system). DISCUSSION FLUORESCENCE ENHANCEMENT Both static and dynamic studies of uranyl fluorescence reveal an enhancement in initial luminescence intensity in the presence of metal ions. When there are changes in the absorption spectrum of UO;+ on addition of the metal ion, this increase can be attributed to an increase in the extinction coefficients. Such an interaction between UO!+ and the metal ion is favoured by hydrolysis of the uranyl ion, which decreases the latter’s charge, and by an increase in ionic strength, which decreases the repulsion between the ions. This is particularly evident with Ag+, which causes an enhancement in the uranyl fluorescence under steady-state conditions only at pH 4 and high ionic strength ( p 3 0.1 mol dm-3).Under these conditions a new absorption band is observed at 435 nm. Nitrates of Ce3+ and Co2+ are known also to form complexes with uranyl ion under certain conditions,ll and with Fe2+ Marcantonatos4 has shown that the absorption spectra of Fe2+ and UO;+ in aqueous solutions are non-additive. The enhancement of fluorescence under certain experimental conditions can be caused by an increase in light absorption and by an increase in the radiative rate constants of the hydroxo-oxo complexes between UOi+ and the metal ion which supersedes the current quenching effect. The initial enhancement of the fluorescence intensity observed in dynamic studies is interpreted in terms of effect of the metal ions on the shift in the emission spectrum of (UOg+)*.With (UO,2+)* the shift in the spectrum is relatively slow (ca. 200 ns) and the fluorescence intensity at 485 nm increases slightly with time up to ca. 150 ns. However, in the presence of metal ions such an increase is not observed, revealing that in the presence of metal ions the relaxation of the fluorescence spectrum has a rate constant > lOlo dm3 mol-1 s-l.H. D. BURROWS, A. C. CARDOSO, S. J. FORMOSINHO AND M. DA G. M. MIGUEL 53 GEOMETRY OF THE URANYL ION AND POSSIBLE STRUCTURES OF GROUND- AND EXCITED- STATE AQUO-URANYL SPECIES Comparison of the UOi+ absorption spectral data of Bell and Biggersl, and the emission spectra from the U* and X* states5 reveals that whereas the X* state shows a shift or stabilization of 420 cm-l after being populated through light absorption, the U* state exhibits a perceptibly larger shift of 870 cm-l.The difference in these stabilizations is equal to the red shift (450 cm-l) observed in the time-resolved emission spectra in the present study (fig. 1). An attractive interpretation of these observations is that the UOi+ ground state and X* have the same equilibrium geometry, whilst the geometry of the U* state is slightly different. Whilst the geometry of O=U=02+ in the ground state has generally been considered to be linear,13 there has been some controversy as to why this should be so when the isoelectronic species Tho, possesses a bent geometry (8 = 122&20).14 Recent theoretical calculation^^^ support the view that the bare, gas-phase uranyl cation is linear whereas Tho, is strongly bent, since the Sforbitals, which are dominant in UOi+, prefer linear geometries while the 6d orbitals, dominant in Tho,, prefer bent geometries.However, calculations on naked UO;+ species may not be completely appropriate for this species in solution surrounded by ligands in the equatorial position. The uranyl ion seems to be able to accommodate between 3 and 6 ligands in the equatorial plane. Recent magnetic circular dichroism studies on aqueous uranyl perchlorate so1utions15 indicate that the dominant species is [U0,(OH,),]2+. This is supported by X-ray diffraction and n.m.r. studies.lG Having five equatorial ligands might be expected to lead to a slight distortion of the uranyl species from linearity.This has been observed experimentally in both neutron17 and X-ray18 diffraction studies of single crystals of uranyl compounds. We wish to suggest that ground-state UOi+ and X* are both slightly bent, whereas U* is linear under our experimental conditions. The degree of distortion may be expected to be relatively small. For a harmonic U=O oscillator the 450 cm-l energy shift is calculated to correspond to a deviation from linearity of ca. 4"; the displacement of the oscillator, x, is taken as x = (1/2) no'/ 180 where I is the length of the U-0 bond and 8' the angular change from linearity. The suggested 0-U-0 bond angle (176") is intermediate between those observed in crystal structures of uranyl perchlorate heptahydrate (161 ')la and [U0,(OH),(OC(NH,),)4](N0,), (177.7 +0.4").li Support for such changes in geometry come from considerations of the oscillator strengths of the transitions. The lowest absorption bands of UOi+ are considered to be Laporte-forbidden.lS Since there is no change in the geometry of a molecular species during an electronic transition, the degree to which the transition is forbidden will be smaller for a transition where the initial state is slightly bent and larger for a linear species.The calculated oscillator strengths are f(U*)/f(X*) = 3.7 (absorption) andf(U*)/f(X*) = 2.3 (emi~sion),~ in agreement with the suggestion that the ground state and X* are slightly bent, whereas U* is linear. Since the relaxation of the U* state illustrated in fig.1 occurs on a relatively long timescale it can be suggested to result from some chemical change. Since the solvent-exchange process in (UO?j+),*, has more associative character in X* and more dissociative character in U*,i it is possible that this relaxation involves an increase in the primary hydration sphere from 5 to 6, a process which may be54 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION Table 3. Metal-ion effect on the non-radiative processes of the electronic states U*(ni &) and X* ( x i dt) of the uranyl ion ion k,&/kF Ag+ 1.1 1.28 Fe2+ 1.05 0.90 Mn2+ 1.26 1.04 Ce3+ 0.9 1.1 Eu3+ 3.4 6.5 UOi+ ca. 7.0 ca. 6.0 facilitated by the presence of metal ions. Such changes in hydration number have previously been postulated in excited states of Eu~+.~O ELECTRON-TRANSFER MECHANISM To discuss the effect of metal ions on reversible crossing and irreversible rates of decay in (UO,Z+>* it is convenient to compare the relative rate constants for each of the non-radiative processes in different states.Table 3 allows a distinction to be made between two kinds of mechanism. The first is relative to the metal ions where an electron transfer between Mn+ and (UO,Z+)* is energetically possible (TI+, Ag+, Fe2+, Pb2+, Mn2+, Ce3+ and Ni2+) and where the 'quenching' effect is the same in both U* and X*: for these ions ku&/kz and kF/kF are close to 1.0. In the second mechanism, found in Eu3+ and UOg+, the metal ion has a much stronger effect on the non-radiative rates of U* than of X*. For the first group of ions the quenching process is currently interpreted in terms of an electron-transfer pro~ess.~-~ Direct evidence for electron transfer has been obtained in the cases of Mn2+ and Ce3+.2v21 As table 3 shows, the effect of the metal ions on the irreversible rates of decay of (UOi+)* is of the same order of magnitude as the effect on the rates of solvent exchange. This fact suggests that both effects should have a common nature.A possible mechanism that will interpret such a situation is an inner-sphere electron- transfer mechanism where (UOg+)* and the metal ion Mn+ are linked by a water molecule. Electron transfer to the (UOg+)* ion leads to quenching. The transfer of a water molecule to the metal ion leaves (UOg+)* with one less water molecule in the first coordination shell, and the subsequent entrance of another H 2 0 molecule into the hydration shell of (UO;+)* can change the electronic nature of the excited uranyl ion.A kinetic scheme which represefits such a mechanism for the process from the state U* is Uzq + Mtg' kd k- d 7 (U**H,O*M"+),q (U* - H 2 0 * Mn+)aq represents an inner-sphere exciplex that leads to quenching via an electron-transfer process; the metal-ion-induced U* -+ X* transition is a non- quenching process. We have not considered the electronic-state transition within the inner-sphere exciplex : (U* - H,O - Mn+),, f (X* - H,O Mn+)aq. In fact such a transition occurs via a solvent-exchange mechanism and consequently requires the dissociation of the exciplex in order to be induced by the metal ion.H.D . BURROWS, A. C. CARDOSO, S. J. FORMOSINHO AND M. DA G. M. MIGUEL 55 1 I I I 4.0 4-5 5.0 5.5 6.0 (IleV)” Fig. 2. Correlation between the logarithm of the quenching rate constants k,& (0) and k,& (0) and the square root of the ionization potentials of various metal ions. (1) Tl’, (2) Ag+, ( 3 ) Fez+, (4) Pb2+, (5) Mn2+, (6) Ce3+ and (7) Ni2+. According to the proposed mechanism the rate constant for the electron-transfer For the reversible crossing induced by the metal ion Equivalent expressions can be established for the state X*. Marcantmato9 has considered the formation of an intimate pair of solvated ions, P*, prior to the formation of an exciplex. A poor correlation was found between the quenching rates and the equilibrium constant of P*, expressed in terms of the molar refractivity of the metal ions.However, other factors, such as coulombic repulsion, can be relevant, and this is invoked to explain the poor character of the correlations. Owing to the impossibility of assessing with some accuracy the effect of the nature of the ions on intimate pair formation, we will consider its equilibrium constant invariant to the metal ions and equal to 1 dm3 mol-l. The quenching rates k,& and k,& decrease with an increase in the ionization potential, I, of the metal ions in the vapour phase, as reported in ref. (2). Fig. 2 presents correlations between logk: or logk? and Pie. A linear plot is found for metal ions with high ionization energies ( I > 30 eV). For metal ions with I < 30 eV the rate of quenching is diffusion controlled.A similar correlation (fig. 3) is also found for the56 4 a cv -Y en - 7 6 5 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION 1 A 2 h A 4.0 4.5 5.0 5.5 6.0 Fig. 3. Correlation between the logarithm of the rate constants k? (A) and k? (A) and P. (1) T1+, (2) Ag+, (3) Fez+, (4) Pb2+, (5) Mn2+, (6) Ce3+ and (7) Ni2+. (I/eV)x rates of induced U* + X* transitions, kp and k?, but with a different slope. Such correlations can be interpreted if both processes are due to an electron-transfer process which occurs by a tunnelling mechanism. For the tunnelling of an electron of mass m through a rectangular barrier of height I and width Y the transmission factor, T, is22, 23 T = exp [ -47~(2rnI)l/~ r/h] log kQ = log Y - 5 . 2 ~ ( 2 m I ) l / ~ r / h (3) and logkQ is a linear function of PI2.For a three-dimensional cubic barrier (4) where v is a frequency factor. Estimation of values of Y from the experimental slopes adds support to the proposed mechanism for the effect of metal ions on the excited uranyl ion. The value found for the induced solvent-exchange process is 4 A (fig. 3), which is close to the diameter of a water molecule (3 A).24 In contrast, a value of r = 5 A can be estimated for the electron-transfer process from the slope of fig. 2. This is close to the diameter of the water molecule plus the radius of the uranium 5forbital in UOi+ (0.57 A).25 These values suggest that when charge is transferred to the uranium atom, reduction of UOi+ occurs and there is quenching, whilst charge transfer to the bridging water molecule between Mn+ at (UOi+)* leads to a breaking of the (UOg+)* .-.OH, bond such that there is no net oxidation and reduction. The (UOi+)* species, thus deprived of a water molecule in its hydration shell, can capture another water molecule and be transformed into the other state, X* or U*. The similar behaviour of the U* and X* states follows from them having empty nu orbitals of similar energy. 9 l9H. D . BURROWS, A. C. CARDOSO, S. J. FORMOSINHO AND M. DA G. M. MIGUEL 57 The use of gas-phase ionization potentials in correlations of (UOi+)* quenching has been questioned on the grounds that these values are higher than the nu energy in the uranyl ion.26 Covalent interactions or polarization effects may have a significant effect on bringing the energy of the electron to be transferred above the nu energy of (UOi+)*, as can be seen by the facts that hard? cations with relatively low gas- phase ionization energies24 such as Ba2+ ( I = 35.5 eV) and Cs+ ( I = 25.1 eV, k , < 3 x lo5 dm3 mol-1 s-l) have very much lower quenching rates than soft cations such as Ni2+ ( I = 35.16 V) or Cu2+ (36.83 V).However, attempts to correlate quenching rates with either the ionization energies of aquo-metal ions26 or with redox potentials of the metal ions3 have been unsuccessful. Similarly, although Marcantonato~~ has found a correlation between quenching rates of some metal ions and E;, the energy necessary to overcome electrostatic repulsion in a (U02,+)*-Mn+ exciplex, we did not find such a correlation when other ions, such as Fe2+ and Ce3+, are included.Where overall electron-transfer occurs, the use of vapour-phase I values in estimating barrier widths for electron tunnelling from metal ions to (UOg+)* can be justified as the process is very fast, and consequently does not allow reorientation of the water molecules in the hydration shell.8 A correlation between quenching rates and either redox potentials or ionization energies of aquo-ions would be expected if the electron transfer was a non-radiative process involving solvent-ion vibrational or multiphonon 28 However, in the reduction of (UO;+)* to UO; there is a significant increase in the U=O bond length, estimated to be ca. 0.3 A.29330 Such a bond-length change requires participation of the U=O vibration in any non-radiative transitions, and the Franck-Condon factors, and consequently rate constants,31 for such a process are low owing to the large reduced mass of the oscillator and to the large energy barrier.Further insight into the quenching process comes from consideration of the cases of Ag+ and T1+, where the rates are diffusion controlled. Combining eqn (1) and (2) gives and since in these cases kg,+ k& % k-, then Further, for diffusion-controlled processes between ions 8RT o , / R T 300011 exp (w,/RT) - 1 k , = ~ (7) where 11 is the medium viscosity and w, the work for the formation of the collision complex.32 This is given as Z1Z2 Ne2 Er cur = where E is the dielectric constant for water and Y is the charge separation. From the data for Ag+ and T1+ for the states U* at X*, at pH 2.8, we can estimate k, = 2.6 x lo9 dm3 mol-1 s-l, and using this and taking r = 3 x lo-* cm for the electron-transfer process gives 2, 2, = 0.7.Since there is negligible hydrolysis of Ag+ or T1+ under these condition^,^^^ 34 this value indicates that (UOi+) must be significantly hydrolysed at this pH, in agreement with earlier suggestions.6 Activation energies were determined for quenching of (UOi+)* by Ag+ and Mn2+ (table 4). Although the solvent modes do not seem to play an explicit role in the ?- The terms hard and soft mean non-polarizable and polarizable, respectively.3758 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION Table 4. Activation energies for the quenching of (UOi+)* by metal ions (EJkJ mol-l) AgZ 25 27.5 a a Mn+ 57 42 44 59 Eu3+ 29.5 11.5 35.5 11.5 uo;+ 24 24.5 66.5 50 a Non-linear Arrhenius plot.electron-transfer process, a decrease in the H,O-metal and H,O-uranyl ion distances with an increase in energy would decrease the barrier width r for electron transfer and consequently increase the rate of electron tunnelling. An activation energy would be observed and would be highest for the highest barrier height. Consequently it is no surprise that quenching by Mn2+ will have a higher activation energy than that by Ag+. Some support for the tunnelling mechanism comes from the fact that with Ag+ strongly non-linear Arrhenius plots are observed (fig. 4) for ki& and k9. QUENCHING BY Eu3+ Eu3+ quenches the U* (n;&) state of excited uranyl ions more strongly than the X* (n",;) state.Moriyasu et al., have stated that quenching of (UO;+)* by Eu3+ occurs via energy transfer, because luminescence was observed for Eu3+. Although we have observed such luminescence, more than 95% under the present conditions can be accounted for by direct excitation of the Eu3+ ion. Consequently an energy-transfer mechanism does not account for the quenching process in water, although it may be important in more viscous solvents or in glasses. Furthermore, electron transfer is not energetically possible and a heavy-atom quenching effect is expected to be similar for both U* and X* states. Europium ion has a quenching effect similar to the autoquenching of the uranyl This would suggest that quenching involves the overlap of atomic orbitals of Eu3+ and molecular orbitals of UOi+ and consequently that H,O hydration molecules are also labile in Eu3+, since Eu3+ has a very high coordination number and there is a fast equilibrium between Eu(H,O);+ and EU(H,O);+.~O The labile character of the H,O molecules is also evident in changes in the inner coordination sphere that are observed for some lanthanide ions, including Eu3+, upon direct electronic e~citation.,~ In the U* state with an n3,& configuration the half-filledf,, orbital has all the lobes pointing in the equatorial direction and can overlap one of the half-filled or empty 5f,, orbitals of Eu3+, which have energies comparable to those of UO;+.However, in the ~ $ 6 ; configuration of the state X*, the half-filled &(fk ,) orbital makes an angle of 45" with the equatorial plane and the overlap is much weaker.So, as with UOi+, Eu3+ affects the state U* more strongly. In the exciplex [(UOi+)* * Eu3+Jaq the rates of reversible crossing, for the state U*, have enthalpies and entropies of activation of AH+ = 35.5 kJ mol-1 and AS? x 4 J K-l mol-l. When such values are compared to those in (UOi+)* (AH? = 67 kJ mol-l, AS7 = 75 J K-l mol-l) it seems that Eu3+ increases ki by increasing the association character of solvent e~change.~ This effect can be attributed to an increase in the radius and charge of the exciplex. The irreversible decays in (UO;+)* have been attributed to hydrogen abstraction from coordinated H,O molec~les.~~ 7 9 26 On increasing the association character of theH. D. BURROWS, A.C. CARDOSO, S. J. FORMOSINHO AND M. DA G. M. MIGUEL 59 24 23 4 c - 22 21 201 - - - - I 1 1 I I 1 24 23 22 a c e 21 2c 15 3.0 3.2 3.4 3.6 1 0 3 ~ 1 ~ A, ki and A, k,. Fig. 4. Arrhenius plots of the quenching rate constants for Ag+: 0 , k,; 0, k,; coordinated water molecules on the solvation shell of the exciplex, Eu3+ increases the number of H,O molecules available for hydrogen abstraction by (UOt+)* and consequently the frequency factor. The activation energy is similar to that in the excited uranyl ion at pH 3. Since for ((UOi+)* * Eu3+) the frequency factor is very high ( A z 1 x 1013 s-l) it seems that hydrogen abstraction from H,O is controlled by the frequencies of metal-water bonds and OH modes, rather than diffusion of water molecules to the U=O bonds. Other rare-earth ions such as Gd3+ and Yb3+ which have empty orbitals of higher energy than thef orbitals of UOi+, because they have no low-lying ( < 22000 cm-l) excited electronic are much less efficient (2-20 times) quenchers than E u ~ + . ~60 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTION The quenching ability of such ions is probably associated with the orbital overlap between the lowest unoccupied orbitals of (UOt+)* and the highest occupied orbitals of the metal ions.We are grateful to INIC and GTZ for financial support. R. Matsushima, H. Fujimori and S. Sakuraba, J. Chem. Soc., Faraday Trans. I , 1974, 70, 1702. H. D. Burrows, S. J. Formosinho, M. G. Miguel and F. Pinto Coelho, J. Chem. Soc., Faraday Trans. I , 1976, 72, 163. M. Moriyasu, Y. Yokoyama and S.Ikeda, J. Znorg. Nucl. Chem., 1977, 39, 2205. M. D. Marcantonatos, J. Chem. Soc., Faraday Trans. 1, 1979,75, 2252. S. J. Formosinho, M. G. Miguel and H. D. Burrows, J. Chem. SOC., Faraday Trans. I, 1984,80, 1717. M. G. Miguel, S. J. Formosinho, A. C. Cardoso and H. D. Burrows, J. Chem. Soc., Faraday Trans. I , 1984,80, 1735. S. J. Formosinho and M. G. Miguel, J. Chem. Soc., Faraday Trans. 1, 1984,80, 1745. H. D. Burrows, S. J. Formosinho, M. G. Miguel and F. Pinto-Coelho, Mem. Acad. Cigncias Lisboa, 1976, 19, 185. L. G. Sillen and A. E. Martell, Stability Constants of Metal-ion Complexes [Chem. SOC. Spec. Publ. 17 and 25, Suppl. 1 (The Chemical Society, London, 1964 and 1971)l. lo M. D. Marcantonatos, Znorg. Chim. Acta Lett., 1977, 25, 87. l1 S. S. Gupta and S.D. Marwah, J. Prakt. Chem., 1964, 24, 83; 1965,29, 1. l 2 J. T. Bell and R. E. Biggers, J. Mol. Spectrosc., 1965,'18, 247. l3 H. D. Burrows and T. J. Kemp, Chem. Soc. Rev., 1974,3, 139. l4 S. D. Gabelnick, G. T. Reedy and M. G. Chasanov, J. Chem. Phys., 1974, 60, 1167; W. R. Wadt, l5 C. Gorller-Walrand and W. Cohen, Chem. Phys. Lett., 1982, 93, 82. l6 M. Aberg, D. Ferri, J. Glaser and I. Grenthe, Znorg. Chem., 1983, 22, 3986. l7 N. K. Dalley, M. H. Mueller and S. H. Simonson, Znorg. Chem., 1972, 11, 1840. l9 C. K. Jrargensen and R. Reisfeld, Chem. Phys. Lett., 1975, 35, 441. 2o H. Kanno and J. Hiraishi, Chem. Phys. Lett., 1980, 75, 553. 21 T. Rosenfeld-Grunwald, M. Brandels and J. Rabani, J. Phys. Chem., 1982, 86, 4745. 22 K. I. Zamaraev and R. F. Khairutdinov, Russ. Chem. Rev., 1978, 47, 518. 23 See for example R. P. Bell, The Tunnel Eflect in Chemistry (Chapman and Hall, London, 1980), 24 S. W. Mayer, J. Phys. Chem., 1963, 67, 2160. 25 R. L. Belford and G. Belford, J. Chem. Phys., 1961, 34, 1330. 26 M. D. Marcantonatos, Znorg. Chim. Acta Lett., 1977, 24, 53. 27 R. A. Marcus, J. Chem. Phys., 1965, 43, 679 and references therein. 28 J. Jortner, J. Chem. Phys., 1976, 64, 4860. 29 N. K. Dalley, M. H. Mueller and S. H. Simonsen, Znorg. Chem., 1971, 10, 323; I. I. Chernyaev, Complex Compounds of Uranium, transl. from Russian by L. Mantel (Davey, New York, 1966). 30 M. G. Miguel, Ph.D. Thesis (Universidade de Coimbra, 1984). 31 H. D. Burrows and S. J. Formosinho, J. Chem. Soc., Faraday Trans. 2, 1977, 73, 201. 32 P. Debye, Trans. Electrochem. Soc., 1942, 82, 265; M. Z. Eigen, Z. Phys. Chem. ( Wiesbaden), 1954, 33 U. Von Meyenburg, 0. Siroky and G. Schwarzenbach, Helv. Chim. Acta, 1973, 56, 1099. 34 J. Burgess, Metal Ions in Solution (Ellis Horwood, Chichester, 1978). 35 M. D. Marcantonatos, M. Deschaux and J. J. Vuilleumier, Chem. Phys. Lett., 1981,82, 36; 1982,91, 36 G. H. Dieke and H. M. Crosswhite, Appl. Opt., 1963, 2, 675. 37 R. G. Pearson, J. Am. Chem. SOC., 1963,85, 3533. J. Am. Chem. Soc., 1981, 103, 6053. N. W. Alcock and S. Esperbs, J. Chem. Soc., Dalton Trans., 1977, 893. p. 21. 1, 176. 149. (PAPER 4/371)

ISSN:0300-9599    

DOI:10.1039/F19858100049

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. 1, 1985, 81, 61 -68 Correlation of the Crystal Structure of Titanium Dioxide Prepared from Titanium Tetra-2-propoxide with the Photocatalytic Activity for Redox Reactions in Aqueous Propan-2-01 and Silver Salt Solutions BY SEI-ICHI NISHIMOTO, BUNSHO OHTANI, HIROSHI KAJIWARA AND TSUTOMU KAGIYA* Department of Hydrocarbon Chemistry, Faculty of Engineering, Kyoto University, Sakyo-ku, Kyoto 606, Japan Received 8th March, 1984 Titanium dioxide (TiO,) has been prepared by the hydrolysis of titanium tetra-2-propoxide, followed by calcination at various temperatures (T,) up to 1000 "C. The content and crystallite size of anatase in the TiO, powders increased upon increasing T, up to 550 "C. In the T, range 550-600 "C a mixture of anatase and rutile was obtained.A further increase in T, resulted in TiO, of rutile structure only. The photocatalytic activities of these TiO, powders for redox reactions were evaluated in the following systems: (1) aqueous propan-2-01 solution, (2) aqueous Ag,SO, solution and (3) aqueous Ag,SO, solution containing propan-2-01. The anatase TiO, showed photocatalytic activity in all these systems, the activity increasing with crystal growth. In aqueous propan-2-01 solution the activity is dramatically enhanced by partial coverage of the TiO, with platinum black. The photocatalytic activity of the rutile TiO, powder was comparable to or even greater than that of anatase when the reaction system included the silver salt, but was negligibly small for aqueous propan-2-01 solution regardless of the partial Pt coverage.Photocatalysed reactions such as the dehydrogenation of alcohol~l-~ or the photo-Kolbe reaction of carboxylic using dispersed TiO, particles have been widely investigated. In many cases various kinds of commercially available TiO, powders have been used as photostable catalysts with sufficient oxidizing and reducing abilities.1°-13 Although the correlation between the physical properties, e.g. bulk crystal and surface structures, and the photocatalytic activity for oxidizing water has been reported for the polycrystalline TiO, electrode systems,14 TiO, suspension systems are still a subject of investigation. We have recently characterized the photocatalytic activity of TiO, powders prepared from Ti(SO,), by hydrolysis and calcination a t various temperatures up to 1000 "C.l5 The activity of such TiO, powders, when mixed with platinum black, for the dehydrogenation of propan-2-01 in aqueous solution depends on the crystal structure; i.e. the activity of anatase TiO, was adequate whereas that of rutile was negligible. The superior photocatalytic activity of anatase compared with rutile has previously been observed for the oxidation of liquid propan-2-01 in the presence of 0,.l6 Unfortunately, a small amount of residual sulphate ions caused considerable lowering of the photocatalytic activity of the TiO, powders obtained from Ti(SO,), because they decreased the pH of the aqueous s u ~ p e n s i o n . ~ ~ In order to characterize the intrinsic photocatalytic activities it is therefore desirable that TiO, powders free from such contamination are prepared by an alternative method and are subjected to further investigation.This paper describes the effects of the physical properties of anion-free TiO, powders prepared from titanium tetra-2-propoxide [Ti(OPr),]l on the photocatalytic activity for redox reactions in aqueous solution. 6162 PHOTOCATALYTIC ACTIVITY OF TITANIUM DIOXIDE EXPERIMENTAL PREPARATION OF TiO, POWDER Titanium tetra-2-propoxide [Ti(OPr),] was supplied by Wako Pure Chemicals and distilled before use (b.p. 92.0-94.0 "C, 0.47 kPa). A mixture of Ti(OPr), (100 cm3, 0.34 mol) and propan-2-01 (1 80 cm3) was added dropwise to an ice-cooled mixture of propan-2-01 (450 cm3) and distilled water (150 cm3) with vigorous stirring. The resulting white precipitate of titanic acid was filtered off, washed repeatedly with distilled water and precalcined for 24 h at 120 "C in air.The TiO, powder thus obtained was subjected to further calcination in air at various temperatures with an electric furnace equipped with a programmed controller (Ohkura EC 53/2 PB). The heating was first performed for 4 h at a constant rate of ca. 1-4 "C min-l until a specified temperature (T,) was attained and continued for 5 h at T,, after which cooling was performed at the same constant rate as in the case of increasing temperature. X-RAY DIFFRACTION ANALYSIS The crystal structures of the TiO, powders were determined by an X-ray diffraction method, using a Rigaku Geigerflex 2013 diffractometer (target, Cu; filter, Ni; 35 kV; 20 mA; scanning speed, 1" min-l). The contents of anatase and rutile in the TiO, were evaluated by integration of the most intense peaks 28 = 25.4" [d = 0.352 nm, the (011) plane of anatase] and 27.3" [d = 0.325 nm, the (1 10) plane of rutile], respectively, by reference to CaCO, as an internal standard.18 The calibration curves for the anatase and rutile were obtained using commercially available anatase (Merck) and rutile (prepared by heating the Merck TiO, powder at 1200 "C for 10 h in aiP), respectively.The mean crystallite size (L) was determined from the broadening (8) of the most intense line in the X-ray diffraction pattern, after corrections for the Ka doublet and instrumental broadening based on the Scherrer equation20 (L = kl/@cos 8, where A is the radiation wavelength, 8 is the Bragg angle and k = 0.90).PHOTOREACTION A finely ground TiO, powder (50 mg), with or without platinum black (typically 5 wt % , Nakarai Chemicals), was suspended in distilled water (5.0 cm3) or aqueous Ag,SO, solution (0.025 mol drn-,, 5.0 cm3) in a glass tube (18 mm dia. x 180 mm, transparent for exciting-light wavelengths > 300 nm). The suspension was purged with Ar for at least 30 min and sealed off with a rubber cap. Propan-2-01 (38 mm3, 0.50 mmol) was injected through the cap by a syringe. The Ar-purged TiO, suspension was irradiated under magnetic stirring at room temperature with a merry-go-round apparatus equipped with a 400 W high-pressure mercury arc (Eiko-sha 400). PRODUCT ANALYSIS A portion (0.2 cm3) was withdrawn from the gas phase (30.0 cm3) of the sealed sample and subjected to analysis for volatile products such as H, and 0,, using a Shimadzu GC 4A gas chromatograph equipped with t.c.d.and a 5A molecular-sieve column (3 mm diameter x 3 m) with Ar carrier at 100 "C. Propan-2-01 and acetone were analysed with a Shimadzu GC 6A gas chromatograph equipped with f.i.d. and polyethylene glycol 20M on a Celite 545 column (3 mm diameter x 2 m) with N, carrier at 90 "C. The procedure and apparatus for the determination of the amount of deposited Ag have been described e1~ewhere.l~ RESULTS AND DISCUSSION PHYSICAL PROPERTIES OF THE TiO, POWDERS PREPARED FROM Ti(OPr), The weight of TiO, powder obtained on calcination, relative to the weight before calcination at 120 "C, decreased with increasing c, attaining a constant value of ca.50%. The phase transition from anatase to rutile was observed in the range 600-650 "C. This transition temperature is considerably lower than that (750-800 "C) for TiO, prepared from Ti(S0,),.l5 The mean crystallite size of the anatase (LA) is plotted against T, in fig. 1. The value of LA did not appreciably change at T, d 550 "C,S-I. NISHIMOTO, B. OHTANI, H. KAJIWARA AND T. KAGIYA 200 150 { 100 4 50 01 I 0 200 400 600 800 1000 T, 1°C Fig. 1. Variation in the mean crystallite size of anatase (LA) as a function of calcination temperature (T,). 63 0 200 400 600 800 1000 T , 1°C Fig. 2. T,-dependent yields of H, (YH2, 0) and acetone (I;CH3)2C0, 0 ) on the irradiation (10 h) of TiO, (50 mg) suspended in aqueous propan-2-01 (38 mm3, 500 pmol) solution (5.0 cm3) under Ar. but rapidly increased in the narrow T, range 55&610 "C.The mean crystallite size of the rutile obtained at T, 2 600 "C was estimated to be > 200 nm, although its exact value could not be determined using the Scherrer equation. PHOTOCATALYTIC ACTIVITY OF TiO, AND Ti0,-Pt IN AQUEOUS PROPAN-2-OL SOLUTION The TiO, powders, prepared as above, when suspended in aqueous propan-2-01 solution and irradiated at Aex > 300 nm under Ar produced H, and acetone. Fig. 2 shows that yields of both H, ( YH2) and acetone ( Y(CH3)2CO) over an irradiation period of 10 h were strongly dependent on the calcination temperature T,. Although the reproducibility of the data in fig. 2 was relatively poor (ca. 30%), a trend is evident64 PHOTOCATALYTIC ACTIVITY OF TITANIUM DIOXIDE 240 200 160 -.1 g 120 5;- 80 40 0 t I ) 0 200 400 600 800 1000 T, 1°C Fig.3. T,-dependent yields of acetone (qCH3)2CH, 0 ) and H2(YH2, 0) on the irradia- tion (1 h) of 5 wt % Pt loaded Ti02 (50 mg) suspended in aqueous propan-2-01 (38 mm3, 500 pmol) solution (5.0 cm3) under Ar. that YH, and Y(CH3),C0 increased with increasing T, until the maximum values ( YH2 w 7 pmol and Y(cH3)2C0 w 10 pmol) were attained at T, = 600 "C. For the treatment at T, > 610 "C, YH2 and qCH,),CO decreased to a greater extent relative to the maximum values. TiO, for T, = 1000 "C, which consists only of rutile, was virtually ineffective for the formation of H, and acetone. As has been well doc~mented,l-~? 21 partial coverage of TiO, ( T , = 610 "C) with a small amount of Pt (up to 5 wt%) enhances the rate of H, formation (RH2).A saturation limit of RH2 was observed in the range of Pt coverage from 2.5 to 5.0 wt% , which was 100-fold greater than without Pt. Fig. 3 illustrates the T, dependences of YH2 and qCH3)2C0 for 1 h photoirradiation of the TiO, powders covered with 5 wt% Pt (Ti0,-Pt). The profiles of these T, dependences are seen to be essentially identical to those in fig. 2, suggesting that the photocatalytic activity of a given TiO, particle is intrinsically determined by the solid properties of the particle but enhanced by the surface Pt. It is evident from fig. 3 that the yields of H, and acetone from Ti0,-Pt are equal within experimental error (reproducibility & 5% ), regardless of q.The equivalence of H, and acetone produced by the photodecomposition of propan-2-01 catalysed by anatase-Pt has been also demonstrated by Teratani et aL3 The net photoreaction in the present system is therefore represented as follows : hv > 800 nm (CH,),CHOH - (CH,),CO + H,. Ti0,-Pt In absence of Pt, YHz was less than I;CH3)2C0. The lower yield of H, is accounted for by the reduction of TirV on the illuminated TiO, catalyst to form Ti11J22 as a non-catalytic side reaction, because the white suspension of TiO, was observed to turn grey during the photoirradiation : hv > 300 nm i(CH,),CHOH +TiIV i(CH,),CO + Ti"' + H+. (3)S-I. NISHIMOTO, B. OHTANI, H. KAJIWARA AND T. KAGIYA 65 Thus, in the absence of partial Pt coverage, the two photoreactions (2) and (3) proceed in competition.It is also plausible that H,, as a photoproduct, reduces TiIV to TiIII. Fig. 2 provides an estimate that at most 60% of the photogenerated reducing species, i.e. electrons photoexcited to the conduction band of TiO,, would be consumed for the self-reduction of TiO, and formation of TiIII. In contrast, the presence of Pt on the TiO, surface could prevent electron trapping of TiIV into Ti111. Moreover, the Pt is responsible for the efficient charge separation of the photogenerated electron-hole pairs in the Ti02,10 thereby promoting the reduction of protons to H, and the oxidation of propan-2-01 to acetone: (TiO,) + hv -+ e- + h+ (4) Pt e-+ H+ -+ Pt-H iH2 (5) h++$(CH,),CHOH -+ ~(CH,),CO+H+. ( 6 ) Reactions (4)-(6) are in accord with the stoichiometry of reaction (2).By reference to the structural evidence, it is clear from fig. 2 and 3 that the TiO, powders containing anatase (T, = 120-620 "C) show sufficient photocatalytic activity to produce H, and acetone, while those containing rutile only (T, b 650 "C) show negligible activity even when covered with Pt. The activity of anatase-containing TiO, powders increased with increasing q. For the TiO, powders (T, = 600-650 "C) containing both anatase and rutile, the activity decreases with decreasing content of anatase. The negligible activity of rutile in aqueous propan-2-01 solution is attributed to the disadvantageous energetics for the reduction of protons to H, compared with anatase : because a rutile electrode exhibits a more positive flat-band potential than an anatase electrode, the energy of photoexcited electrons in the conduction band of rutile is expected to be lower than that of those in anata~e.~~,, PHOTOCATALYTIC ACTIVITY OF TiO, IN AQUEOUS SOLUTIONS OF SILVER SALT Photoirradiation (Aex > 300 nm) of aqueous suspensions of these TiO, powders containing Ag,SO, led to the formation of 0, and the deposition of Ag metal on the TiO, parti~1es.l~.l9 The T, dependences of the yields of O,( YO,) and the Ag deposit ( YAg) over an irradiation period of 1 h are shown in fig. 4. A linear relationship with a slope of 0.23 was obtained between the YO, and YAg, in accord with the following net photoreaction: hv > 300 nm 4Ag++2H20 - 4Ag+02+4H+. TiO, (7) Support for the release of H+ in this scheme was obtained by the observation of a rapid decrease in the pH (from ca.4 to 2) of the suspension during photoirradiation. A slightly larger amount of Ag deposit compared with the stoichiometry in reaction (7) would be attributed to the partial photoadsorption of another product, O,, on the TiO, s ~ r f a c e . ~ ~ - ~ ~ Fig. 4 shows that both Y,, and YO, are relatively small and nearly independent of T, when TiO, powders prepared in the lower T, region (120-550 "C) are used. Upon raising T, from 550 to 600 "C the photocatalytic reaction became more rapid to give eventually maximum values of YAg z 110 pmol and YO, z 25 pmol. Y,, and Yo, then decreased and approached constant values of 80 and 20 pmol, respectively, in the T, range 800-1000 "C. The surface area of rutile probably decreases but activity per unit area is constant in this T, range.Note that in the presence of Ag+ ions the66 PHOTOCATALYTIC ACTIVITY OF TITANIUM DIOXIDE 100 80 3 60 x 40 20 0 0 200 400 600 800 1000 T, 1°C Fig. 4. T,-dependent yields of Ag metal (YAP, 0) and 0, (Yo,, a) on irradiation (1 h) of an aqueous TiO, suspension in Ag,SO, solution (250 pmol Ag+, 5.0 cm3) under Ar. TiO, powders containing the rutile structure only, which were practically inactive in aqueous propan-2-01 solution without Ag+, showed a larger activity than the anatase TiO, (T, < 600 "C). In particular, the mixed anatase-rutile powder (600 6 &/"C < 620) showed the highest activity for the formation of Ag and 0,, in contrast to the activity for propan-2-01 dehydrogenation (fig.2 and 3). These facts clearly demonstrate that the photocatalytic activity of rutile is essentially comparable to that of anatase in certain photoreaction systems, as in this case where the more reducible Ag+ ions but not protons can react with the photogenerated electrons of rutile. Furthermore, the observed highest activity of the anatase-rutile mixture demonstrates that the crystal structure alone cannot explain the activity. It is likely that the surface area25 and the porosity29 also have significant effects on the photocatalytic activity of TiO, powders. PHOTOCATALYTIC ACTIVITY OF TiO, IN AQUEOUS SOLUTIONS OF SILVER SALT AND Addition of propan-2-01 to the aqueous suspension of TiO, containing Ag,SO, resulted in the oxidation of propan-2-01 into acetone together with the formation of 0, and the deposition of Ag metal.Fig. 5 shows variations of the product yields YAg, Yo, and qCHs)2C0 over the 1 h photoirradiation as a function of T,. Formation of 0, by the TiO, powders at T, < 550 "C was negligible, while a small amount of 0, (ca. 5 pmol) was obtained for T, 600 "C. YAg and ~ C H , ) , C O increased with increasing T, to attain their maxima of ca. 125 and 50 pmol at T, = 650 "C, and then decreased toward constant values on further increase in The total yield of oxidation reduction products satisfied a stoichiometry given by 2 Y;CH3)2C0 + 4YO2 = YAg. Clearly, the oxidation of both water [reaction (7)] and PROP AN-2-OL (> 650 "C). propan-2-01 hv > 300 nm (CH3),CHOH + 2Ag+ (CH,),CO + 2Ag+ 2H+ (8) TiOl is involved in the present system, although the proportion of water oxidation is much smaller.It is seen from fig. 5 that the TiO, powders at T, >, 550 "C, which consist onlyS-I. NISHIMOTO, B. OHTANI, H. KAJIWARA AND T. KAGIYA 67 140 120 100 80 5 1 x 60 40 20 0 0 200 400 600 800 1000 T, 1°C Fig. 5. T,-dependent yields of Ag metal (YAP, O), acetone ( Z;CH3)2C0, 0 ) and 0, (Yo,, 0 ) on irradiation (1 h) of an aqueous TiO, (50 mg) suspension in Ag,SO, solution (250 pmol Ag+, 5.0 cm3) containing propan-2-01 (38 mm3, 500 pmol) under Ar. of anatase, give rise to propan-2-01 oxidation almost exclusively according to reaction For T, > 600 "C the profile of the T,-dependent activity in this system (fig. 5 ) is similar to that without propan-2-01 (fig.4), although propan-2-01 predominantly undergoes oxidation instead of water. Since the TiO, powders in this T, range contained an increasing proportion of rutile as T, increased (fig. I), this similarity seems to originate largely from the action of the rutile. In contrast, the apparent photocatalytic activity of the TiO, powders at & < 550 "C, which contain only anatase, is at least three-fold enhanced by the addition of propan-2-01. This is clearly a result of the oxidation of the added propan-2-01 that would occur at the illuminated anatase TiO, more readily than that of water. Compared with the results for TiO, prepared from Ti(SO,),, the activity of TiO, powders from Ti(OPr), seems to be greater in this photoreaction system: YAg values for the TiO, powders from Ti(OPr), ( T , = 600 "C, LA = 53 nm) and Ti(SO,), ( T , =700 "C, LA = 37 nm)15 were 73 pmol for 1 h irradiation and 87 pmol for 2 h irradiation, respectively.ENHANCEMENT OF THE PHOTOCATALYTIC ACTIVITY OF TiO, POWDERS BY REDUCIBLE (3). OR OXIDIZABLE SPECIES As described above, the photocatalytic activity of TiO, powders suspended in aqueous solution depends on both the crystal structure and the solution species to be oxidized or reduced by the photogenerated hole (h+) or electron (e-), respectively. The latter effect was clearly demonstrated by the different photocatalytic activities in the three photoreaction systems. Almost no reaction occurred on the photoirradiation of an aqueous TiO, suspension in the absence of propan-2-01 or Ag+. However, the addition of propan-2-01 to this system led to the formation of a small amount of H,68 PHOTOCATALYTIC ACTIVITY OF TITANIUM DIOXIDE and acetone, some of which was produced by a non-catalytic process.When Pt black was loaded on the TiO, powders, the activity increased to yield stoichiometric amounts of H, and acetone. These facts show that an easily oxidizable species such as propan-2-01 is more effective for the trapping of the photogenerated hole, which has little ability to oxidize water. In addition, the Pt loading enhanced the electron trapping by H+ and depressed the self-reduction of TiIV to T F , especially in the case of anatase. The effect of Ag+ was also evident: TiO, could oxidize water to 0, with the aid of the easy reduction of the Ag+. A further enhancement of the catalytic activity of TiO, was observed on the addition of both propan-2-01 and Ag,SO,, particularly in presence of anatase (see fig.4 and 5 ) . We thank Prof. Satohiro Yoshida (Kyoto University) for his valuable advice on X-ray diffraction measurements. We also thank the Instrumental Analyses Research Centre of Kyoto University for permission to use atomic absorption spectrometers. Both referees are thanked for their careful and constructive reports. T. Kawai and T. Sakata, J. Chem. SOC., Chem. Commun., 1980,694; Nature (London), 1980,286,474; Chem. Lett., 1981, 81. P. Pichat, J-M. Herrmann, J. Disdier, H. Courbon and M-N. Mossanega, Nouv. J. Chim., 1981, 5, 627. S. Teratani, J. Nakamichi, K. Taya and K. Tanaka, Bull. Chem.SOC. Jpn, 1982, 55, 1688. K. Domen, S. Naito, T. Onishi and K. Tamaru, Chem. Lett., 1982, 555. B. Kraeutler and A. J. Bard, J. Am. Chem. SOC., 1978, 100, 2239; 5985. H. Reiche and A. J. Bard, J. Am. Chem. Soc., 1979, 101, 3127. I. Izumi, F-R. F. Fan and A. J. Bard, J. Phys. Chem., 1981,85, 218. H. Yoneyama, Y. Takao, H. Tamura and A. J. Bard, J. Phys. Chem., 1983,87, 1417. 1982, 86, 172. 'I H. Reiche, W. W. Dunn, K. Wilbourn, F-R. F. Fan and A. J. Bard, J. Phys. Chem., 1980,84, 3207. lo A. J. Bard, J. Photochem., 1979, 10, 59; Science (Washington D.C.), 1980, 207, 139; J. Phys. Chem., l1 A. J. Nozik, Annu. Rev. Phys. Chem., 1978, 29, 189. l 2 M. S. Wrighton, P. T. Wolczanski and A. B. Ellis, J. Solid State Chem., 1977, 22, 17. l 3 M. A. Fox, Acc. Chem. Res., 1983, 16, 314. l4 C. Stalder and J. Augustynski, J. Electrochem. Soc., 1979, 126, 2007. I5 S. Nishimoto, B. Ohtani, A. Sakamoto and T. Kagiya, Nippon Kagaku Kaishi, 1984, 246. l6 R. B. Cundall, R. Rudham and M. Salim, J. Chem. SOC., Faraday Trans. I , 1976,72, 1642. l8 R. A. Spurr and H. Myers, Anal. Chem., 1957,29, 760. l 9 S. Nishimoto, B. Ohtani, H. Kajiwara and T. Kagiya, J. Chem. SOC., Faraday Trans. I , 1983,79,2685. *O H. Klug and L. E. Alexander, X-Ray Diyraction Procedures (Wiley, New York, 2nd edn, 1974), 21 J. Disdier, J-M. Herrmann and P. Pichat, J. Chem. SOC., Faraday Trans. I , 1983, 79, 651. z2 A. D. Buss, M. A. Malati and R. Atkinson, J. Oil Colour Chem. Assoc., 1976, 59, 369. 23 M. V. Rao, K. Rajeshwar, V. R. Pai Verneker and J. DuBow, J. Phys. Chem., 1980, 84, 1987. 24 A. Mills and G. Porter, J. Chem. SOC., Faraday Trans. I , 1982, 78, 3659. 28 E. Borgarello, J. Kiwi, M. Gratzel, E. Pelizzetti and M. Visca, J. Am. Chem. Soc., 1982, 104, 2996. 26 A. H. Boonstra and C. A. H. A. Mutsaers, J. Phys. Chem., 1975, 79, 1694. 27 G. Munuera, V. Rives-Arnau and A. Saucedo, J . Chem. SOC., Faraday Trans. I , 1979,75, 736. 28 A. R. Gonzalez-Elipe, G. Munuera and J. Soria, J. Chem. SOC., Faraday Trans. I , 1979, 75, 748. 29 L. Kruczynski, H. D. Gesser, C. W. Turner and E. A. Speers, Nature (London), 1981,291, 399. M. R. Harris and G. Whitaker, J. Appl. Chem., 1962, 12, 490. p. 618. (PAPER 4/384)

ISSN:0300-9599    

DOI:10.1039/F19858100061

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1985,81, 69-82 Adsorption and Conductivity Studies in Oxychlorination Catalysis Part 4.-Effect of Adsorption on the Conductivity of Copper(1) Chloride Films BY PETER G. HALL, RICHARD A. H A N N , ~ PHILIP HEATON f AND DAVID R. ROSSEINSKY* Department of Chemistry, University of Exeter, Stocker Road, Exeter EX4 4QD Received 12th March, 1984 Thin copper(1) chloride films on glass have been prepared, and the effects of adsorbed nitrogen, oxygen and ethylene on conductivity have been studied using special electrode configurations to maximise the effects. Compositions have been examined by secondary-ion mass spectrometry (SIMS) and X-ray photoelectron spectroscopy (XPS), porosity by scanning electron micrography (SEM) and net reaction by microreactor studies.The conductor is electronic, the 0.53 eV activation energy agreeing with literature values. Adsorbed ethylene gives an irreversible decrease in conductivity, oxygen a reversible one, while nitrogen, physisorbed, is ineffective. The constancy of activation energy despite the adsorption of gases implies that adsorption results mainly in fewer charge carriers rather than a change in mechanism. From the conductivity results oxygen is reversibly chemisorbed, at higher temperatures possibly becoming included within the lattice. In ~tudiesl-~ of oxychlorination catalysis, interactions of absorbed ethylene and other gases with copper(1, 11) chloride and similar systems have been examined. This work is now extended to the effects of adsorption on the conductivity of copper(1) chloride in film form.The active-site formulation of adsorption proposed by Taylor4 and elaboratored by Kevan5 and Boudart6 has led on to the view that catalytic and semiconductor properties are intimately related.:-" Concomitantly, ideas of charge transfer in catalysisl2* l3 have been examined1"-': with but tenuous experimental support until re~ent1y.l~- l9 Adsorption can change the work function of crystalline semiconductors20 and the consequent alteration of the contiguous space-charge region is expected to affect the magnitude of the surface conductivity.21- 22 The only marginal effect of ethylene adsorption on the conductivity of bulk powder compactions of copper(1, 11) ~hlorides,~ has led us now to examine adsorption effects on thin films of CuCl, where the surface/bulk ratio is greatly enhanced.Adsorbates were nitrogen, ethylene, water and oxygen, and their effects on the d.c. conductivity and its time and temperature dependence were observed. Characterisation of the solid was effected by scanning electron microscopy while chemical composition was established by SIMS and XPS. The interaction of CuCl with C1, to give CuIl-doped material was studied by Harrison and ~ o w o r k e r s . ~ ~ - ~ ~ The temperature dependence of CuCl conducti- 2 7 , 29--36 has been variously depicted as either curved or bilinear Arrhenius behaviour. Some theoretical studies have been attem~ted.~' 42 The dependence on potential and pressure has been studied ;43 high pressure introduces metallic c o n d u ~ t i o n ~ ~ ~ 44-49 and the possibility of s~perconduction.~~~ 51 Single crystals t Present address : Electronics Group, ICI, Runcorn.1 Present address: Johnson Matthey Research Centre, Blounts Court, Sonning Common, Reading. 6970 OXYCHLORINATION CATALYSIS have29v44v 4 5 9 5 2 9 5 3 conductivities o = 2 x to lo-* C2-l cm-l at ca. 300 K, which have been compared with compaction values.3o CuCl as solid ele~trolyte~~ (or doped semiconductor) has been examined.2g+ 55-59 There is scope for debateso* as to whether band conduction or hopping predominates. For CuCl and doped CuCl no simple Cole-Cole dielectric dispersion characteristic of pure hoppings2 is observable, and the band form~lation~~ seems the more tenable. In mixed valence, e.g. CuII-doped CuCl, systems, the Fermi level, conductivity and adsorption properties will be determined by the valence compo~ition.~~ EXPERIMENTAL SAMPLE PREPARATION AND MOUNTING Copper(1) chloride, prepared as before,23 was purified using a vacuum depositer (Creative Vacuum Services, Worthing, Sussex).The white powder was placed in a molybdenum sample boat and brought up to deposition temperature at CQ. mmHg,* so that it just began to deposit on glass slides placed above it. The evaporating/heating current was removed and a solid plug of CuCl was left, free of volatile impurities. A frame supporting four clean microscope slides was placed ca. 3 in. above the plug of CuCl in a sample boat. The ends of the slides were masked and the chamber was evacuated to 8 x mmHg. Current was passed through the sample boat via a Variac until visible melting began.The evaporating current was then kept constant to prevent vaporisation of any impurities with higher melting points. A white film was allowed to form slowly on the slides, until opaqueness indicated a continuous layer. For a completely opaque layer depositions from several plugs of CuCl were required. Gold electrodes were then deposited on the surface, using a mask which left a strip (1 mm or 2.5 cm wide) of copper(r) chloride free of gold [fig. 1 a)]. An alternative electrode configuration was also used where platinum was deposited on a microscope slide first. Then a discontinuity was scored across the middle of the slide, before copper(1) chloride was deposited over the platinum [fig. l(b)]. Coated slides were kept under vacuum, rapidly transferred to a nitrogen-filled dry box, dried over P20,, and mounted in a cell with 4 painted leads of gold wire.The evacuated mmHg) cell was immersed in a water thermostat. Alternatively, for high temperatures, the sample side was transferred to a stainless-steel cell where it was mounted with silver paint onto gold wire. The stainless-steel shell was closed by means of a rubber or metal O-ring and evacuated to lod5 mmHg. A close-fitting furnace allowed the sample to be thermostatted over the range 25-400 "C, as monitored by a chrome-alumel thermocouple positioned inside the cell. CONDUCTIVITY MEASUREMENTS AND GAS DOSING Direct-current measurements were taken using either a polarograph (Bruker, model EMS) with an ( X , Y) chart recorder, or a voltage supply (DEB Electronics Ltd) with an ammeter (Keithley Instruments 610C electrometer), giving currents generally of the order of A.Between readings the sample was shorted to avoid any build up of charge. A potential range of k 3 V was commonly used. Known pressures of gas (nitrogen, oxygen or ethylene) were connected to the high-temperature conductivity cell, to a pressure of 0.57 atmt (chosen so as not to exceed the seal specification). For dosing with H20 a small glass tube was filled with glass wool and ca. 2 cm3 of added distilled water was degassed by a freeze-thaw cycle under vacuum, repeated twice. The sample was then opened to the manifold and subject to the saturated vapour pressure of water at ca. 17 "C. SECONDARY-ION MASS SPECTROMETRY AND OTHER TECHNIQUES SIMS was performed by Dr J.Myatt at the ICI Research Laboratories using a Riber SIMS instrument. The mass spectrometer gave both a negative-ion and a positive-ion spectrum. CuCl * 1 mmHg x 133.322387 Pa. t 1 atm = 101 325 Pa.P. G. HALL, R. A. HANN, P. HEATON copper (I) chloride microscope slide AND D . R. ROSSEINSKY copper (1) chloride < d i scon t inu i t y microscope platinum sIide (6) 71 Fig. 1. (a) Gold/copper(I) chloride substrate/gold configuration. (b) Substrate Pt/superposed CuCl/substrate Pt configuration. was deposited on 12 mm diameter glass discs to provide samples for this technique. It was possible to vary the sample temperature and dose with a gas (oxygen) within the machine. All spectra were necessarily recorded under ultra-high vacuum (ca.1 0-lo Torr).* X-ray photoelectron spectroscopy (XPS) was performed at the Analytical Laboratory, ICI on a Kratos ES200 instrument. Samples of CuCl on glass rectangles of size 2 cm x 0.5 cm were analysed and exposed to oxygen at various temperatures within the machine. Scanning electron microscopy was performed on a Cambridge S600 microscope. The microreactor used in Part 33 was used to study the bulk reaction of CuCl with 0,. Dry helium was the carrier gas and temperature thermostatting was effected by a chromatography oven. Conventional mass spectra were obtained on a VG Micromass MM16F spectrometer. RESULTS AND DISCUSSION Using the low-temperature glass cell the conductance, R-l, of a film of CuCl with platinum electrodes [fig.l(b)] was measured from the time of mounting and evacuation. The result is shown in fig. 2, with an initial strong decrease in R-l with time, but stabilisation to within an order of magnitude after several days. A similar decrease was observed for the alternative [fig. 1 (a)] configuration. This decrease was attributed to annealing rather than an outgassing process. After 1 month the CuCl layer was shown by electron micrography of a step where part of a layer had broken away from on top of another, exposing a layer cross-section, to be a very ordered microcrystalline array. The deposited layer was 5-6 pm thick, and commonly total thicknesses were estimated to lie between 1 and 20 pm. Such film, being on top of the electrodes, was not very robust and became partially detached from the metal on first exposure to vacuum.Therefore in subsequent experiments use of the configuration shown in fig. 1 (b) was favoured and care was taken to leave the sample for several (2-3) days before conducting experiments. The linearity or otherwise of the voltage against current plot produced by the chart recorder indicates the deviation from Ohm's law of a sample, depending on the vacuum. Initial vacua of only ca. lop4 mmHg resulted in non-Ohmic behaviour. At high voltages (> 10 V) the white sample blackened and the resistance suddenly decreased by several orders of magnitude, similar observation~~~ being attributed to a phase change. However, a perfectly Ohmic response was obtained for vacua better than mmHg at potentials + 3 v. At near-to-laboratory temperatures with a 15 min equilibration time at each temperature, values of E, = 0.46 eV for ascending temperature (1 3-39 "C) and 0.64 eV for descending temperature were obtained.The discrepancy is not unexpected as many solids are ' structure-sensitive' in this temperature range owing to imperfect structure * 1 Torr = 101 325/760 Pa.72 - 3 r OXYCHLORINATION CATALYSIS - 9 I 1 I I I 13 26 39 52 tlh Fig. 2. Change of log(conductance) with time for configuration of fig. 1 (b). or lack of Extending the temperature range with an icebath gave a linear graph of In R-l against 1/T [see fig. 3(a)], with E, = 0.53 eV, between the extremes. The high-temperature cell allowed measurements over a wider range of temperature. 45 min for temperature stabilisation and 15 min for equilibration were allowed.Some deterioration of vacuum from the usual mmHg occurred at > 200 "C, but no effect on E, was discovered. This was later traced to a nitrogen-producing organic impurity which fortunately played no further discernible role. From fig. 3(b) an Ea value of ca. 0.5 eV is obtained and compared (table 1) with literature values, falling well within the upper group. When two activation energies for conduction are reported in the literature, that with E, z 1 eV above 180 "C is attributed to ionic conduction (cationic conduction by a Frenkel defect mechanism), while that with Ea z 0.5 eV below 140 "C is attributed to hole conduction, allowing27, 30 an estimate of the percentage ionic conduction at any temperature. No upturn in E, above 180 "C is seen.Harrison and PrasadZ7 report that the electrode material caused the d.c. values above 190 "C to fall by 40 and 80% below the a.c. value for Ag and Pt electrodes, respectively, and doubtless our Au electrodes also suppress ionic carriers and the consequent upturn. Thus the electronic conductivity of a thin film of CuCl can be studied by restricting the temperature range to below 140 "C to minimise the ionic contribution, the former being potentially more sensitive to gas effects than ionic conductivity. Using the configuration of fig. l(b), no apparent effect on the resistance of the sample was observed on exposure to oxygen, nitrogen, helium or ethylene. Prolonged exposure was complicated by the obscuring of small changes by the background decrease due to aging.The effect of adsorption of water vapour on R-l is shown in table 2. During exposure the film physically disintegrated. There was slight reversibility in that when the sample was re-evacuated after 2 h, readings steadied at 29% above the original R-l value. Further studies with H,O were not attempted because of the destruction of the film. Since the micrographs had indicated very little porosity in the copper(1) chloride, gaseous penetration to the main conducting pathways between the electrodes isP. G. HALL, R. A. HANN, P. HEATON AND D . R. ROSSEINSKY 73 -21.3 -22.0 - - I C 7 -22.7 1 s M 3 - 2 3 . 1 3.3 3.5 103KIT - 13 -16 - x .. 5 -19 M - - 22 1 1 1 I 1.6 2.2 2.8 3 . 4 103K/T Fig. 3. (a) Activation energy at low (ambient) temperature.(b) Activation energy up to 280 "C: 0, ascending temperature ; 0, ascending temperature (same sample). Table 1. Activation energy for electronic conduction in CuCl activation energy for electronic temperature conduction/eV range/"C ref. 0.53 < 280 this work 0.51-0.59 < 140 27 0.54 90- 130 55 0.51 < 225 at 20 kbar 29 0.39 < 220 44 0.37 < 280 30 doubtful and therefore the electrode configuration of fig. 1 (a) was used subsequently with the low-temperature glass cell. In a preliminary experiment, using an electrode separation of 2.5 cm, the results in table 3 were obtained. These results all come from the same slide and are therefore subject to possible error due to a previously unreversed effect. Accuracy in measuring R-l was difficult ( f 10%) because the relatively wide electrode separation gave low current, and the experiment was repeated using new slides for each gas, with electrode separations reduced to 1 mm.Fig. 4 and 5 show the effects of 0, and C,H, on the conductance. The addition and evacuation of N, had no discernible effect. The aging behaviour of the film characterised before74 OXYCHLORINATION CATALYSIS Table 2. Effect of water vapour on conductance of CuCl duration of exposure to change in water vapour/ h R-' (%) 0 0.5 1 2 0 15 23 35 Table 3. Preliminary observations of effects of gases on conductance of CuCl change in R-l on gas exposure to gas change in R-l on re-evacuation helium none none ethylene decrease of 30-50% after 4 h none oxygen nitrogen none none decrease of 20% after 4 h ca.0.4 of decrease reversed 0 t l h Fig. 4. Effect of oxygen on the conductance of CuCl film : (A) oxygen added ; (B) re-evacuation. admission of the gas serves to indicate the value to which the gas should return on re-evacuation. Some scatter in results is attributed to experimental error in the current measurement, due to the need for high sensitivity. In the case of oxygen, a decrease of ca. 40% over 7.25 h is apparently totally reversible, rather than only partially as earlier indicated in table 3 when no account was taken of aging. The slow rate of change of R-l with time confirms the possibility that diffusion through the film is slow. Ethylene (fig. 5 ) causes an irreversible decrease of ca. 25%. The adsorption results for ethylene (Part 2)2 indicate it will chemisorb onto anyP.G. HALL, R. A. HANN, P. HEATON AND D . R. ROSSEINSKY 75 0 \ A 0 20 40 60 ti h Fig. 5. Effect of ethylene on CuCl conductance: (A) ethylene added; (B) re-evacuation. impurity Cu2+ sites, in localised adsorption. The 25% decrease suggests that the Cu2+ impurity is responsible for a large proportion of the conduction mechanism. CuCl has been cla~sified~~y 39 as a p-type semiconductor with acceptor levels lying just above the valence band. If these acceptor impurity levels are filled by electrons from an adsorbate molecule, no electrons can be excited from the valence band and no positive holes are left as carriers, hence conductivity decreases. Wolkenstein13 has defined ‘ weak ’ and ‘strong’ chemisorption; only in the latter case is an effect on conductivity to be expected.Nitrogen is also included in the category of localised adsorption but bonding is so weak (van der Waals only) that conduction is not altered. Ionosorption is a possibility for adsorbate bonding when an electron from the conductor band or hole from the valence band becomes captured or injected by the surface species, but calculation^^^ show that pure inosorption is likely to occur only for 0, adsorbate. Considerable work has been carried out and reviewedl31 66-68 on the adsorption of oxygen on semiconducting oxides but not chlorides. Ionosorption is generally irreversible, being associated with chemisorption, and causes an increase in conductivity for p-type semiconductors. Because of the adsorption data3 of Part 3 (qo = - 7 kJ mol-l on CuC1, qo = - 16 kJ mol-1 on CuCl,, A; H = 6.82 kJ mol-l) and the observed decrease in conductivity, ionosorption is not suspected.The mechanism by which 0, decreases R-l for CuCl is unclear: while the reversibility does not preclude localised bonding, adsorption possibly being both weak yet specific to particular sites, incorporation of the adsorbate into the chemical lattice seems likely. The slow change of R-l with time is consistent with such absorption of oxygen. As a further investigation, the effect of the gases on the activation energy of conduction E, was established. Using the high-temperature cell, a fresh sample was76 OXYCHLORINATION CATALYSIS Table 4. Activation energies on exposure to gases (eV). condition (gas) vacuum exposed re-evacuated oxygen 0.65 & 0.06 0.57 0.02 0.62 ethylene 0.61 0.65 0.65 nitrogen 0.60 f 0.02 0.61 0.66 a & indicates mean of T-increasing and T-decreasing runs, the error limits representing observed differences. mounted and outgassed for ca.3 days to allow aging effects to stabilise. Then Ea was determined by ascending and descending temperature runs. A gas was then admitted at room temperature and left on the sample for ca. 12 h, and the activation energy again established by ascending and descending temperature runs. Finally the sample was re-evacuated and outgassed at room temperature for > 12 h to a vacuum of better than mmHg before a single ascending temperature run. 1 h was allowed for equilibrium at each reading. The results are shown in fig.6 and table 4. For oxygen and ethylene R-l again decreased on admission of the gas, while for nitrogen it remained constant. For oxygen there appears to be no increase in R-l on re-evacuation, although we omitted to take readings immediately after re-evacuation or to characterise the aging effect. No substantial change was observed in E, for any gas, and so decreases in conduction must be attributed to a removal of carriers rather than a change in mechnism. The reaction of oxygen with copper(1) chloride was further investigated by a range of techniques, starting with SIMS. Spectra were recorded for a copper(1) chloride sample before any oxygen was added in order to gauge sample purity. The negative-ion spectrum [see fig. 7(a)] shows large peaks for [Cll-, [ClJ, [CuClI- and [CuCl,]- with very few other peaks.This technique does not give information on the valence state of the copper ions. The positive-ion spectrum was considerably more complex [see fig. 7(c)], because of identifiable impurities including Na+, Al+, K+, [CuOH]+ and traces of a nitrogen-containing contaminant. A sample was exposed to ca. 1 0-1 Torr of oxygen at room temperature and at 150 "C for 1 h. The negative-ion spectrum was unchanged except for a [CuClCNI- peak. A positive-ion spectrum was not taken, as the entities which indicate a reaction has taken place, e.g. [CuOCl], would probably produce negative rather than positive ions. Samples were subjected to high vacuum (ca. 10-lo Torr) when mass spectra were taken, thus removing any weakly absorbed material. The sample was then heated for 30 min at 300 "C under ca.10-1 Torr of oxygen and outgassed overnight before spectra were taken. The positive-ion spectrum [see fig. 7 (c)] indicated that the organic impurity had largely been removed. There were strong sodium impurity levels due to diffusion from the soda-glass microscope slide and the high ion yield of alkali metals generally. The negative-ion spectrum [see fig. 7(b)] showed peaks for [CuO,]-, [CuClCNI-, [CuOCl]-, [CuOI-, [CNI- and [CNOI-; it is concluded that at 300 "C a reaction has occurred incorporating oxygen and CN- into the surface. At lower temperatures the inclusion of oxygen within the surface cannot be discounted, as a vacuum removes reversibly held species. The XPS studies indicated that copper@) chloride samples were virtually free of copper(r1) chloride (fig.8). The prepared samples remained stable to atmospheric oxidation over 2 h. Quantitative comparison of the copper and chloride ion peak areas0 pu P. G. HALL, R. A. HANN, P. HEATON AND D. R. ROSSEINSKY ( I - W I - x) 801 - - ? m - 9 0 Fr M . m 0 + ,? N I I I I I h Y - m 7778 I OXY CHLORINATION CATALYSIS CUCli CuCN' CUCI- I i I I 3x104 -f K* c 1- I 3 Ao3 I ! lo3 I 106 I 3 x 1 0 ~ I Fig. 7. Secondary-ion mass spectra of CuCl: (a) negative-ion spectrum of untreated showing absence of impurity, (b) negative-ion spectrum after preheating in 0, at 300 "C and (c) positive- ion spectrum after preheating in 0, at 300 "C.P. G. HALL, R. A. HANN, P. HEATON AND D. R. ROSSEINSKY 2 5 c M 0 0 2 20 rc.n E 5 1 5 - % - 3 a en 1- g l o - 5 L. 0 79 -- - 234 2 50 266 2 82 energyIeV Fig. 8. CuCl and CuCl, (reference) X-ray photoelectron spectra. 0 80 160 240 T/"C Fig. 9. Extent of reaction of 0, with CuCl from microreactor studies. Flow rate : lJ,0.32 cm3 s-l ; 0, 0.25 cm3 s-' (corrected to 0.32 cm3 s-l). shows a deficiency of chlorine assuming the 1 : I stoichiometry of CuC1. A simple explanation is offered by the presence of some zero-valent copper metal on the surface, as has been reported69 for samples outgassed at 350 "C. Another plausible possibility is that some of the copper(1) is present as a species other than chloride, e.g. the oxide. However, intrinsic differences between sample and standards could introduce errors of 10-2004 in inferred stoichiometries.80 OXYCHLORINATION CATALYSIS Using a microreactor to study the bulk reaction of copper(1) chloride with oxygen, injections of oxygen of constant size (25 x lops dm3) were made over ca.6 g of solid. Assuming negligible adsorption the eluted peak area should be constant (see dashed line in fig. 9). However, as the temperature was increased, the amount of oxygen emerging from the solid decreased. It appears that even at room temperature some small extent of reaction is occurring. Above 200 "C eluted peaks showed tailing, suggesting an additional process above this temperature. The colour of the white solid appeared to darken, and therefore the formation of an oxychloride was suspected. Two samples, copper(1) chloride and copper(1) chloride exposed to a continuous stream of oxygen at 250 "C for ca.40 h, which caused a substantial grey discolouration, were subjected to mass-spectrometry analysis. The spectra were taken with a sample-probe temperature of 250 "C. At 450 "C the vapour of CuCl has been shown to contain substantial amounts of the cyclic trimer CU,C~,~~, 71 and also a t e t ~ a m e r . ~ ~ A later mass-spectrometry study of the vapour in equilibrium with solid CuCl(280-430 "C) has indicated comparable concentrations of Cu3Cl, and Cu,Cl, molecules and a smaller concentration of CU,C~,.~~ Peaks of associated species up to Cu,Cli only were observed by us, with characteristic isotope-splitting patterns. The spectrum of sample (ii) was identical to that of (i) and therefore it was concluded that mass spectrometry is too insensitive to examine the possible oxychloride, probably because the temperature is too low to volatilise the grey-black impurity.Previous attempts to characterise this oxygen/CuCl product have used X-ray diffracti~n,~~ electron spin resonance spectroscopy and mass spe~trometry.~~ The original proposition of an oxychloride being formed at high temperatures was made by The situation was initially complicated by reports of Cu,OCl, formation from copper(I1) 76 The work of Allen and Clarks9 established that CuCl, does not react with oxygen while CuCl forms CuO.CuC1, (or CuO/CuCl,). There are also reported studies carried out in solvents at lower temperatures (5-1 10 "C) of the reversible uptake of oxygen by Cu177 and a change in valency from CuI to CUI*.~* CONCLUSIONS The value of the activation energy of conduction (0.53 eV) obtained for the prepared thin films of copper(1) chloride indicated that the electronic conduction of 'pure' copper@) chloride was being measured, despite the nitrogen-containing impurity of probably organic origin and an apparent non-stoichiometry in the Cu to C1 ratio as shown by XPS.Owing to the non-porous nature of the film, the electrode configuration was important if the effects of gas on conductivity were to be observed. The absence of an effect at room temperature and above with helium or nitrogen was as expected for gases which are only physically adsorbed. The behaviour of R-l with ethylene is consistent with irreversible chemisorption. The site to which C,H, attaches is responsible for a large proportion of the charge carriers. Adsorption results indicate that ethylene chemisorbs to Cu2+.Therefore Cur* appears to be present as an impurity site and these acceptor impurity levels participate in the electronic conduction mechanism. The reversible behaviour of oxygen in decreasing conductivity was not expected for a p-type semiconductor. Absorption is tentatively suggested as being responsible since microreactor studies indicated that incorporation of oxygen into the solid was occurring. Its characterisation by SIMS failed, possibly because of the high vacuum used in detection. The colour change and SIMS results indicate that at higher temperatures (> 200 "C) an irreversible uptake of oxygen occurs. This could explain why the oxygen effect was not reversible after temperature cycling.As no significant change in E, was observed for any gas, decreases in conduction were attributed toP. G. HALL, R. A. HA", P. HEATON AND D. R. ROSSEINSKY 81 removal of charge carriers by donation of electrons from the adsorbate into the acceptor impurity levels, rather than a change in mechanism. We thank the S.E.R.C. and Imperial Chemical Industries p.1.c. for the CASE award under which the work was performed, and the latter for access to analytical and spectral services arranged by Dr J. Wolstenholme. We also thank both referees for their helpful comments. P. G. Hall, P. Heaton and D. R. Rosseinsky, J . Chem. Soc., Faraday Trans. I , 1984, 80, 2777. P. G. Hall, P. Heaton and D. R. Rosseinsky, J .Chem. Soc., Faraday Trans. I , 1984, 80, 2785. P. G. Hall, P. Heaton and D. R. Rosseinsky, J . Chem. Soc., Faraday Trans. I , 1984, 80, 3059. H. S. Taylor, J . Chem. Phys., 1927, 15, 624. T. Kwan, Adv. Catal., 1954, 6, 67. M. Boudart, J . Am. Chem. Soc., 1952,74, 1531. P. Aigrain and C. Dugas, Z . Elektrochem., 1952, 56, 363. K. Hauffe and H. J. Engel, Z . Elektrochem., 1952, 56, 366. P. B. Weisz, J . Chem. Phys., 1952, 20, 1483. lo P. B. Weisz, J . Chem. Phys., 1953, 21, 1531. l 1 S. R. Morrison, Adv. Catul., 1955, 7, 259. l2 T. E. Madey, J. T. Yates, D. R. Sandstrom and R. J. H. Voorhoeve, in Treatise on Solid State Chemistry, ed. N. B. Hannay (Plenum Press, New York, 1976) vol. 6B, p. 59. l 3 Th. Wolkenstein, The Electronic Theory of Catalysis on Semiconductors (Pergamon Press, Oxford, 1963); Adu.Catal., 1960, 12, 189. l4 P. J. Fensham, Q. Rev. Chem. Soc., 1957, 11, 227. l5 K. J. Miller, J . Chem. Ed., 1971, 48, 582. l6 A. Clark, The Theory of Adsorption and Catalysis (Academic Press, London, 1970). l7 J. J. Herrmann, J. Portefaix, M. Forissier, F. Figueras and P. Pichat, J . Chem. Soc., Faraday Trans. l8 M. Takata and H. Yanagida, Yogyo Kyokai Shi, 1979,87, 13. l9 N. V. Suntsov, V. I. Arkharov, V. N. Konev and A. G. Miloslavskii, Dokl. Akad. Nauk S.S.S.R., *O A. R. Plummer in The Electrochemistry of Semiconductors, ed. P. J. Holmes (Academic Press, London, 21 A. Many, Y. Goldstein and N. B. Grover, Semiconductor Surfaces (North-Holland, Amsterdam, 22 D. R. Frankl, Electrical Properties of Semiconductor Surfaces (Pergamon Press, Oxford, 1967).23 P. G. Hall, M. Parsley, D. R. Rosseinsky, R. A. Hann and K. C. Waugh, J . Chem. SOC., Faraday 24 L. G. Harrison and C. F. Ng, Trans. Faraday Soc., 1971, 78, 1787. 25 L. G. Harrison and C. F. Ng, Trans. Furaday Soc., 1971, 67, 1801. 26 L. G. Harrison and C. F. Ng, Trans. Furaday Soc., 1971, 67, 1810. 27 L. G. Harrison and H. Prasad, Trans. Faraday Soc., 1974, 70, 471. 28 C. F. Ng and K. S. Leung, J . Catal., 1981, 67, 410. 29 R. S. Bradley, D. C. Munro and P. N. Spencer, Trans. Faraday Soc., 1969, 65, 1912. 30 L. G. Maidanovskaya, I. A. Kirovskaya and G. L. Lobanova, Izv. Akad. Nauk SSSR, Neorg. 31 Y. W. Hsueh and R. W. Christy, J . Chem. Phys., 1963, 39, 3519. 32 J. B. Wagner and C. Wagner, J. Chem. Phys., 1957, 26, 1597. 33 C. Tubandt, Handb.Experimentalphys., 1932, 12, 383. 34 A. K. Shukala, S. Chattopadhyay and C. N. Rao, Cryst. Lattice Defects, 1976, 6, 191. 35 K. D. Becker, G. W. Herzog, D. Kanne, H. Richtering and E. Stadler, Ber. Bunges. Phys. Chem., 36 R. E. Malpas, Ph. D. Thesis (Exeter University, 1977). 37 K. S. Song, J . Phys. (Paris), 1967, 28, 195. 38 K. S. Song, J . Phys. Chem. Solids, 1967, 28, 2003. 39 A. B. Kunz, R. S. Weidman, J. Boettger and G. Cochran, Int. J . Quantum Chem., Quantum Chem. 40 L. Kleinman and K. Mednick, Phys. Reti. B, 1979, 20, 2487. 41 P. C. de Mello, M. Hehenberger, S. Larsson and M. Zerner, J . Am. Chem. Sac., 1980, 102, 1278. 4 2 H. A. Cerdeira, Solid State Commun., 1980, 35, 109. I, 1979, 75, 1346. 1979, 245, 1372. 1962). 1965). Trans. I , 1983, 79, 343.Muter., 1967, 3, 936. 1970, 74, 527. Symp., 1980, 14, 585. 4 FAR 182 OXYCHLORINATION CATALYSIS 43 C. Divakar, M. Mohan and A. K. Singh, Solid State Commun., 1980, 34, 385. 44 C. W. Chu, S. Early, T. H. Geballe, A. Rusakov and R. E. Schwall, J. Phys. C, 1975,8, 1241. 45 C. W. Chu, A. P. Rusakov, S. Huang, S. Early, T. H. Geballe and C. Y. Huang, Phys. Rev. By 1976, 18, 2116. 46 G. J. Piermarhi, F. A. Mauer, S. Block, A. Jayaraman, T. H. Geballe and G. W. Hull Jr, Solid State Commun., 1979,32,275. 47 G. J. Piermarini, F. A. Mauer, S. Block, A. Jayaraman, T. H. Geballe and G. W. Hull Jr, High Pressure Sci. Technol., Proc. 7th Znt. AIRAPT Con$ (National Bureau of Standards, Washington D.C., 1980), vol. 1, p. 395. 48 C. W. Chu, S. Early, T. H. Geballe and C.Y. Huang, J. Less-Common Metals, 1978, 62, 463. 49 B. Batlogg, J. P. Remeika and R. G. Maines, Solid State Commun., 1981, 38, 83. 50 G. C. Vezzdi and J. Bera, Phys. Rev. B, 1981, 23, 3022. 52 A. V. Joshi and L. B. Wagner Jr, J. Electrochem. Soc., 1975, 122, 1071. 53 J. Rivera, L. A. Murray and P. A. Hoss, J. Cryst. Growth, 1967, 1, 17. 54 A. V. Joshi, in Fast Zon Transport in Solid, ed. W. van Goo1 (North-Holland, Amsterdam, 1973). 55 T. Matsui and J. B. Wagner Jr, J. Electrochem. Soc., 1977, 124, 610; 941 ; 937. 58 S. Geller, J. R. Akridge and S. A. Wilber, Phys. Rev. B, 1979, 19, 5396. 57 T. Takahashi, 0. Yamamoto, S. Yamada and S. Hayashi, J. Electrochem. Soc., 1979, 126, 1654. 58 T. Takahashi, 0. Yamamoto and A. Sawai, J. Appl. Electrochem., 1978,8, 161. 5g T. Jow and J. B. Wagner Jr., J. Electrochem. Soc., 1979, 126, 1963. 6o N. F. Mott, Nuovo Chim., 1958, 10, 7, suppl. 312. 61 I. G. Austin and N. F. Mott, Science, 1970, 168, 71. 62 D. R. Rosseinsky and J. S. Tonge, J. Chem. Soc., Faraday Trans. 1, 1982,78, 3595. 63 P. Day, Endeavour, 1970, 29, 45. 64 N. F. Mott and R. W. Gurney, Electronic Processes in Ionic Crystals (Dover, New York, 1964). 65 M. Green and M. J. Lee, in Solid State Surface Science, ed. M. Green (Marcel1 Dekker, New York, 68 S. J. Thomson and G. Webb, Heterogeneous Catalysis (Oliver and Boyd, Edinburgh, 1968), p. 59. 67 F. S. Stone, A h . Catal., 1962, 13, 1. 68 D. 0. Hayward and B. M. W. Trapnell, Chemisorption (Buttenvorths, London, 2nd edn, 1964), p. 261. 6e J. A. Allen and A. J. Clark, J. Appl. Chem., 1966, 16, 327. 70 L. Brewer and N. L. Lofgren, J. Am. Chem. Soc., 1950,72, 3038. 71 C. H. Wong and V. Shomaker, J. Phys. Chem., 1957,61, 358. 72 H. M. Rosenstock, J. R. Sites, R. J. Walton and R. Baldock, J. Chem. Phys., 1955, 23, 2442. 73 Guido, G. Balucci, G. Gigli and M. Spoliti, J. Chem. Phys., 1971, 55, 4566. 74 €3. Deacon, Brit. Pat. 1948 (1866). 75 A. E. Korvezee, Reel. Trav. Chim., 1931,50, 1085. Von K. Jellinek and A. Rudat, 2. Anorg. Allg. Chem., 1926, 155, 73. 77 J. E. Bulkowski, P. L. Burk, M-F. Ludmann and J. A. Osborn, J. Chem. Soc., Chem. Commun., 1977, 498. G. Davies, M. F. El-Shazly, D. R. Kozlowski, C. E. Kramer, M. W. Rupich and R. W. Slaven, in Inorganic Compounds with Unusual Properties, ed. R. B. King, Adv. Chem. Ser. no. 173 (American Chemical Society, Washington D.C., 1979), vol. 2, p. 178 and references therein. 1969), vol. 1. (PAPER 4/404)

ISSN:0300-9599    

DOI:10.1039/F19858100069

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1985, 81, 83-89 Adsorption and Conductivity Studies in Oxychlorination Catalysis Part 5.-Temperature-programmed Desorption BY PETER G. HALL,* PHILIP HEATON? AND DAVID R. ROSSEINSKY Department of Chemistry, University of Exeter, Stocker Road, Exeter EX4 4QD Received 15th March, 1984 The desorption of ethylene, 1,2-dichIoroethane (EDC), carbon dioxide and other adsorbates from CuCl, has been studied using temperature-programmed desorption (t.p.d.). A chemi- sorptive centre corresponding to an ethylene concentration of 5 x lo+' mol mP2 has an energy of activation for desorption (Ed) of 73 kJ mol-l. Another chemisorptive site has a value of 57 kJ mol-l for E d when the ethylene coverage is 10-6-10-7 mol mP2. These magnitudes are closely related to previously reported heats of chemisorption, suggesting adsorption to be only weakly activated.EDC is shown to desorb readily whereas CO, has three distinct values of E d (85, 94 and 97 kJ mol-l). An assumed value of 1013 s-l for the pre-exponential factor Ad is appropriate in Redhead's method of analysis ; the kinetics are consistent with unimolecular decomposition. In this paper we report results relating to the desorption of species from copper(I1) chloride; adsorption behaviour has been discussed in Parts 2 and 3. Temperature- programmed desorption is used to evaluate activation energies of desorption and surface coverages. These can be correlated with results obtained by gas-adsorption chromatography (g.a.c.) to complete the energetics of adsorption-desorption for ethylene oxychlorination.EXPERIMENTAL APPARATUS The temperature-programmed desorption technique consists of the following steps : ( 1) catalyst pretreatment, (2) preadsorption of the adsorbate, (3) evacuation after preadsorption to remove the physically adsorbed gas, (4) programmed desorption of the residual chemisorbed gas into the stream of a carrier gas, ( 5 ) detection of the desorbed gas in the carrier and (6) trapping and analysis of the desorbed gas to establish its identity. The apparatus was designed to allow these procedures except for step (3), i.e. there was no vacuum facility, and the experiment relied on the adsorbent being held at the boiling point of the adsorbate for sufficient time to remove physically adsorbed gas. Another requirement of the apparatus was that it should contain a cryogenic section where th.e adsorbent could be cooled down to ca.-200 "C and then the temperature increased in a controlled manner up to ca. 300 "C. This was achieved by the novel use of small aluminium arid brass blocks closely fitted around the sample tube which were cooled by immersing them in a Dewar of liquid nitrogen and then allowed to warm up inside a temperature-programmable oven. The apparatus is summarised schematically in fig. 1. In contrast to the g.a.c. experiments, in order to minimise the possibility of readsorption of t Present address : Johnson Matthey Research Centre, Blounts Court, Sonning Common, Reading. 8384 DESORPTION STUDIES IN OXYCHLORINATION CATALYSIS Fig. 1. Schematic diagram of t.p.d.apparatus. Key: A, adsorbate cylinder/regulator; C, catalyst; CR, chart recorder; He, helium cylinder/regulator ; I, injection head; K, katharometer ; MS, molecular sieve; R, restrictor; SV, switching valve; T, thermocouple/digital voltmeter/chart recorder; Z, oven. the desorbing materials' a short plug of copper(~i) chloride (ca. 1 cm long, 0.5 g) was used in the t.p.d. experiments. The catalyst was preheated at 250 "C for 1 h under helium to allow comparison to g.a.c. results. After cooling to room temperature, the exit side of the sample tube was connected directly to a katharometer detector and the apparatus was leak-tested. With the catalyst held at a low temperature T, by liquid nitrogen and by-passed by the carrier gas, a stream of adsorbate was introduced into the flow.All adsorbates were of standard laboratory grade as in Part 2. The pneumatic switching valve was then used to divert the helium+adsorbate stream over the catalyst. When the adsorbate was indicated by the detector to be in excess of that amount which the catalyst could adsorb, the adsorbate supply was closed off. For liquid adsorbates the switching valve was not necessary, a syringe being used to inject into the carrier-gas stream over the catalyst. The sample temperature was maintained at until the detector response had returned to close to the baseline. (No vacuum facility was available to remove totally all physically adsorbed material.) The temperature of the catalyst was then raised at a linear rate p using the metal block and programmable oven.The natural heating rates for each metal block had to be determined in separate experiments so that the ramping of the oven could be matched to ensure a continuation of the initial nearly linear rate. To check the identity of desorbing species the katharometer was replaced by a glass column of 15% MS 550 silicone on Chromosorb W. AW-DMCS at 80 "C with a flame ionisation detector. This column separated ethylene from all chlorination products, and for the heating rates generally used allowed sampling every 10 "C. RESULTS AND DISCUSSION It was found that the different metal blocks gave a range in/3 of 4-12 K min-l, which could be extended to 80 K min-l by removing the block completely, but at higher rates there was considerable error in reading T, and B.The apparatus has therefore provided a more suitable heating rate than the very rapid exponential rate previously used2 for eliminating thermal gradients in the catalyst.P. G. HALL, P. HEATON AND D. R. ROSSEINSKY 85 time c- c temperature Fig. 2. Desorption spectrum for ethylene from copper(r1) chloride. Table 1. Activation energies for the desorption of C,H, from CuC1, 1 2b 3 4 5 6 7 8 9 10 187.85 218.45 215.95 206.7 216.2 21 1.2 212.15 187.65 213.15 194 0.1307 0.3436 0.5464 0.4902 0.472 1 1.2165 0.1910 0.0538 0.1286 0.0538 53 59.8 58.2 55.9 58.7 55.7 59 54 60 55.8 246.7 0.0555 71.5 263.15 0.2469 73 C - - - - - - 257.15 0.4386 75 254.15 0.1626 71.4 257.1 0.1032 73.6 - - - a Runs 1-6 on same sample of CuC1,; runs 7-10 on another sample. Run 2 shown in fig.2. - Denotes peak not clearly distinguishable. RESULTS FOR THE DESORPTION OF ETHYLENE FROM COPPER(II) CHLORIDE For a constant heating rate, the position of peak maxima were independent of the preadsorption temperature T,. A value of -90 "C for T, was chosen to avoid the inherent problems of lower temperatures with physisorbed and liquid ethylene (b.p. = - 104 "C), which appears generally as a peak at ca. -92 "C. Maintaining the sample at -90 "C for 10 min caused the loss of the majority of weakly held species. Thus the desorption chromatogram shown in fig. 2 does not start with a flat baseline. The main features of the spectrum are a major peak at ca. - 60 "C and a smaller peak shoulder at ca. - 10 "C. Sampling by g.c. at ca. 10 "C intervals showed that all desorbing species in this temperature region were pure ethylene, i.e.there was no86 DESORPTION STUDIES IN OXYCHLORINATION CATALYSIS L l I l I 1 I I I I 1 I 1 I I I 1 I l 1 I I I I I l I I I I I I TPC Fig. 3. Desorption spectrum from copper(rr) chloride. -80 -60 -40 -20 0 20 40 60 80 dissociative adsorption. However, an important observation revealed by sampling at high sensitivity is that the concentration of ethylene never reaches zero but reduces to approximately one-seventieth of the value at the main peak maximum. The actual heating curve also appears above the peaks, a twin channel recorder being used. Each desorption was continued up to 120 "C and then the temperature was reduced to T, again and the process repeated for different heating rates; therefore a sample was only preheated to 250 "C once, before any ethylene adsorption.The value of 120 "C as an upper limit was used because of the complicating effects of chemical reaction and disproportionation of the cupric chloride above 150 "C, which was detected by the katharometer as a continuous increase in baseline. The results for a series of heating rates are given in table 1, with activation energies for desorption Ed calculated by Redhead's method3 with an assumed value of 1013 s-l for the pre-exponential factor Ad. Since E d values show a random variation with p, the assumption for Ad is valid. The values of the two activation energies for desorption are 57k4 and 73+2 kJ mol-l. This compares with 73+ 1 and 83 k 1 kJ mol 1-1 reported2 for higher exponential heating rates.After calibration the area of the major peak corresponded to a surface concentration I' of ca. 10-6-10-7 mol m-2. The uncertainty in r arises from the resolution of the peak from shoulders (see dotted line in fig. 2) and the unknown extent of tailing. The amount of C2H, desorbing with E d = 73 kJ mo1-l was equal to a surface concentration of ca. 5 x mol m-2. RESULTS FOR THE DESORPTION OF 172-DICHLOROETHANE FROM COPPER(I1) CHLORIDE For preadsorption temperatures below the melting point (-35 "C) of 1,2- dichloroethane (EDC) it was found that the extensive desorption on melting, even for a 1 mm3 dose size, decreased the sensitivity of the experiment to an unacceptable level. Choosing T, to be -20 or 20 "C overcome this problem, but a longer length of time than for ethylene was required for physically adsorbed material to be desorbed; 30 min was allowed.The effect of heating was to cause a slow steady increase in the amount of EDC desorbing up to its boiling point (83 "C) and then a slow decrease with no other maximum before heating ceased at 150 "C. Sampling at ca. 20 "C intervals revealed that only EDC was desorbing. In one experiment, the solid was held at T, = 20 "C for 4 h to allow all weakly adsorbed species to desorb; on heating, no EDC was seen. The conclusion is that EDC is only physisorbed on copper(r1) chloride and readily desorbs. RESULTS FOR THE DESORPTION OF CARBON DIOXIDE FROM COPPER@) CHLORIDE = -90 "C, a desorption spectrum was obtained from the katharometer detector as shown in fig. 3. In preliminary experiments using ethylene as adsorbate forTable 2.Activation energies for the desorption of CO, from CuCl, 1 2 3 4 6 7 8 9 10 11 12 13 294 290 294 295 316 319 302 310.5 309 304 302 28 1 280 0.0575 0.0555 0.0529 0.0779 0.1035 0.0983 0.2131 0.23 15 0.23 15 0.1267 0.2584 0.1004 0.0992 85 84 85.5 85 90.5 91.5 84.5 87 86 86.5 84 80 80 33 1 33 1 332.5 331.5 333 333.5 326.5 0.0667 0.06 17 0.0667 0.08 17 0.1 138 0.1075 0.1945 96 96 96.5 95.5 95 95.5 92 - 338 321.5 3 20 318 - 0.258 0.2564 0.0952 0.09 16 - 94.5 90 92 91.5 a - - 339.5 344.5 345 342 347 345 347 342 332 332 - - 0.0641 0.093 1 0.1263 0.2203 0.254 0.265 0.273 0.2564 0.0980 0.0926 - 98.5 99 98.5 96 97 96.5 97 95.5 95.5 95.5 - a Denotes peak not clearly distinguishable. Run 5 spectrum shown in fig.3. 9 z b P w88 DESORPTION STUDIES IN OXYCHLORINATION CATALYSIS Comparison with the spectrum from an f.i.d. indicated that all peaks above 0 "C were inorganic. The peaks at ca. 40, 60 and 75 "C were identified by control experiments as being due to the desorption of carbon dixoide. The spectra obtained for the desorption of CO, from CuCl, were independent of the presence of C2H4 (and vice versa) and the value of T,, provided that T, was below the sublimation temperature (- 78 "C) of carbon dioxide. The results for a series of heating rates are given in table 2, with activation energies for desorption calculated by Redhead's method3 with an assumed value of 1013 s-l for Ad. The table shows it is not necessarily possible to distinguish all 3 peaks on each run.The average values of the activation energies for desorption were 85, 94 and 97 kJ mol-l. There was no apparent dependence of Ed on /3 so the assumption of 1013 s-l for Ad was justified. No calibration for the amount of CO, adsorbed was carried out, but it was a factor of 10, greater than the amount of C,H, (fig. 3). The magnitudes of Ed indicate non-dissociative adsorption of CO,. RESULTS FOR THE DESORPTION OF OTHER ADSORBATES In order to identify the desorption peak at ca. 10 "C, oxygen and water vapour adsorbates were studied. Owing to the proximity of the melting point, water was unsuitable for t.p.d. in the temperature range of interest. Oxygen preadsorbed on CuCl, gave no peaks, although CO, impurities in the oxygen gave a weak background of peaks between 40 and 80 "C.Dosing of copper(]) chloride with oxygen at - 196 "C gave small peaks at - 135 and 3 "C in one run, and for another sample peaks at - 145 "C (B = 0.2688 K s-l) and 6.5 "C (p = 0.2268 K s-l) were obtained. When copper(r) chloride was dosed with CO, at -90 "C a single peak at 3.5 "C (B = 0.2315 K s-l) was obtained. Thus the unidentified peak could be due to the desorption of CO, or 0, from CuCl present by disproportionation during pretreatment. On heterogeneous surfaces Ad will vary, and consequently kd also varies. Distortion of curves attributed to varying Ed is often due to varying kd. An interesting analysis of factors that may influence Ad for first-order kinetics is given by Peter~nann.~ An evaluation of readsorption may be gained by the effect on Tm of the flow rate F.However, T, is approximately proportional to -log Fso a large variation is required. Despite these assumptions Ed values in the present work were obtained to a reproducibility of S%, indicating the useful nature of the simple Redhead analysis method. The important parameter is T,; an error of + 5 % gives an error of + 5 % in Ed. A variation of f 5% in gives an error of only 0.1 % in Ed. CONCLUSIONS The desorption spectrum for ethylene from copper@) chloride, shown in fig. 2, clearly indicates that besides physically adsorbed C2H4 there are two other types of chemisorbed ethylene, neither being dissociatively adsorbed. The amount of material adsorbed remained roughly constant for different preadsorption temperatures, and thus the two peaks cannot be attributed to coverage-dependent repulsions between adsorbed molecules.One of these adsorption sites has a surface concentration r of ca. 5 x mol rn-, and an Ed value of 73 kJ mol-l. The results in Part 2 show that qst at this region of r is 65 kJ mol-l. The surface concentration of the other site is ca. 10-6-10-7 mol m-,. Table 10 in Part 2 clearly shows that qst is nearly constant at 44 kJ mol-l for this region of r. Thus t.p.d. and g.a.c. lead to the conclusion that other than physically adsorbing sites there are two types of sites, as also observed by Par~ley,~ which have energetics summarised by fig. 4. The values of the activationP. G. HALL, P. HEATON AND D. R. ROSSEINSKY 89 65 i‘ k J mol-’ 4 1 4 3 k J mol-’ I I I 57 k J mol-’ Fig. 4. Adsorption-desorption energy diagram for ethylene on copper(I1) chloride.energy for adsorption Ei are 8 and 13 kJ mol-l, the low value of which indicates that surface diffusion is not necessary for desorption, but the non-zero values confirm that the C2H4 adsorption is chemisorption. These conclusions ignore the broad width of the major desorption peak and the observation that the concentration of desorbing C2H, did not reach zero below 300 “C. The broadness indicates surface heterogeneity, with a range of higher energy sites, the surface concentration of which could not be determined. These higher-energy sites could be due to edges, dislocations and other surface imperfections, the extremely low surface concentration of which generally render them of little catalytic interest.using T, and the peak width at half-maximum only, gave Ed = 71.5 kJ mol-1 for the higher-energy site of C,H, on CuCl,, which is in very good agreement to 73 kJ mol-1 obtained by Redhead’s analy~is.~ However, for the other sites it gave Ed = 15 kJ mol-l, which is too low because of the peak shape being broadened by surface heterogeneity. Results for EDC indicate that it is only physically adsorbed and desorbs readily. Carbon dioxide has a high activation energy (ca. 100 kJ mol-l) for desorption when preadsorbed on CuC1, below its sublimation temperature. When dosed at higher temperatures such as those used in g.a.c. it readily desorbs. An alternative Financial assistance for this work was provided by an S.E.R.C. - Imperial Chemical Industries PLC (Runcorn) CASE award. We also thank Dr R. A. Hann (I.C.I.) and Dr T. Tribbeck (I.C.I.) for helpful discussions and advice. R. J. Cretanovic and Y. Amenomiya, in Catalysis Reviews, ed. H. Heinemann (Marcel Dekker, New York, 1972), vol. 6, p. 21. P. G. Hall, M. Parsley, D. R. Rosseinsky, R. A. Hann and K. C. Waugh, J. Chem. SOC., Faraday Trans. 1, 1983, 79, 343. P. A. Redhead, Vacuum, 1962, 12, 203; Trans. Faraday SOC., 1961, 57, 641. L. A. Petermann, Nuovo Cimento, Suppl., 1967, 5, 364. M. Parsley, Ph. D. Thesis (Exeter University, 1979). L. D. Schmidt, Catal. Rec. Sci. Eng., 1974, 9, 1 1 5. ’ D. Edwards Jr, Surf. Sci., 1975, 54, 1 . (PAPER 4/42 1 )

ISSN:0300-9599    

DOI:10.1039/F19858100083

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1985,81, 91-104 Reaction of Catalase with Ethylhydrogen Peroxide BY MORDECHAI L. KREMER Department of Physical Chemistry, The Hebrew University of Jerusalem, Jerusalem 9 1904, Israel Received 14th March, 1984 C,H,OOH reacts with catalase in a basically irreversible reaction in the course of which the species called compound (I) is formed and decomposed. The formation of compound (I) is preceded by the formation of a precursor complex which is able to react with a further molecule of C,H,OOH to yield an inactive biperoxy complex. The biperoxy complex causes a diminution of the extent of formation of compound (I) at hgh [C,H,OOH]. As a consequence, compound (I) can never be formed quantitatively. Some of its physical constants can, nevertheless, be evaluated.Compound (I) with C,H,OOH appears to retain C,H,OH in its structure. The study of the reactions of enzymes with pseudosubstrates has played an important role in efforts to elucidate the mechanism of enzyme action. It has often been hoped that the deviation of the enzyme reaction from its natural path will provide clues regarding certain steps in the otherwise unresolvable (or only partially resolvable j complexity of steps which comprise the natural course of enzyme catalysis. With this purpose in mind, the reactions of catalase with H202 and substituted peroxides have been studied, the former being the natural substrate for ~atalase.l-~ By using derivatives of H202 as substrates, the course of the reaction changes, resulting, when alkyl groups are used as substituents, in oxidation of the alkyl C,H,OOH -+ CH,CHO + H20 (A) and no evolution of oxygen.A characteristic feature of the reaction of catalase with both H202 and substituted peroxides is the formation of optically distinct intermedi- ates. In the commonly accepted interpretation, the (green) primary intermediates [called compound (I) (C,) because they are the first to appear in the spectrum], which are formed with mono- or un-substituted hydrogen peroxide, are regarded as essen- tially identical. Moreover, the decrement of the molar absorptivity of catalase upon formation of C, with CH,OOH or with C2H,00H has been used to determine the concentration of CI formed with H202.8* In the present investigation the basis of this cross-determination will be examined in some detail.One of the assumptions on which the above procedure is based concerns the possibility of a complete saturation of catalase by excess alkylhydroperoxide to form C,. This assumption was questioned by Brill and Williams, who found that the degree of saturation of catalase by C2H,00H depended in a complex fashion on the excess of C2H,00H applied.1° Details of their observations will be given in the experimental section. The study of the model reaction with the pseudo-substrate created some additional problems which have to be solved before any conclusions can be drawn regarding various aspects of the enzymic reaction of catalase with H202. First and foremost among these problems is the question of the detailed mechanism of the catalase + C2H,00H reaction. It is the objective of the present work to construct a 91 4-292 REACTION OF CATALASE WITH C,H,OOH scheme for the reaction.On the basis of this scheme the degree of saturation of catalase and the question of the determination of the physical constants of CI will be discussed. EXPERIMENTAL The basic experimental observations on the reaction can be summarized as follows. By adding a dilute solution of C,H,OOH, with an initial concentration (x,) < ca. 50 pmol dm-3, to a dilute solution of catalase, its total concentration on a haematin basis (e,) being ca. 5 pmol dm-3, there is a transient decrement of the optical absorption (AA) at the Soret peak of catalase (A. = 405 nm). After AA has reached a maximum (A&,,) there is a gradual regeneration of the original catalase spectrum.The time scale of events is ca. 1-5 min. There is a variation of the rate of regeneration of the spectrum with the type of catalase used, it being higher with horse-liver catalase and lower with bacterial catalase.ll9 l2 With x, > 50 pmol dm-3, the original catalase spectrum is not restored. Instead, a new species, denoted ethylhydroperoxicatalase, compound (11), is formed to a varying extent, depending on the conditions of the experiment. These changes occur on a time scale of several tens of minutes. Therefore, the best conditions for a saturation of catalase haematins (E) by C,H,OOH to form C, quantitatively were considered to be (a) short times (< 2 min) and (b) x, 50 pmol dm-3.10 Using this strategy, Brill and Williams undertook a systematic investigation of the catalase C,H,OOH reaction.’, They found that AA,,, first increased with increasing x,, but above x, x 300 pmol dm-3, a further increase of x, caused a decrease of AA,,,.Furthermore, they could show that even under optimal conditions for the formation of C, (largest AA,,,), at x, x 300 pmol dm-3 (a large excess of C,H,OOH over E), there was still free cutaluse haematin in the system. No explanation for this observation was offered. Also, no specific mechanism of the reaction was forwarded. It was suggested that the reactions which occurred in the system were basically irreversible. This suggestion contradicted an earlier assumption of Chance and Schonbaum, who considered the analogous reaction of catalase with CH300H to be a reversible, equilibrium reaction.12 Later work of Chance and Schonbaum supported the hypothesis of irreversibility, as they were able to show the formation of acetaldehyde from ethylhydrogen peroxide during its reaction with catalase.,9 1 3 9 l4 Irreversibility will be assumed also in the present analysis, the immediate aim of which is the construction of a mechanism which can account for the observations of Brill and Williams.Details of the method of calculation applied are given in the following section. METHOD OF CALCULATION A certain mechanism was assumed and the relevant set of rate equations was integrated numerically, using the DGEAR routine of the IMSL computer program library. By inserting a given set of initial concentrations, rate constants and molar absorptivity parameters, the routine calculated the instantaneous concentrations of the various species and of the optical absorption decrements (AA) at a preset array of points of time. The calculations were continued until AA reached a maximum (AAmax).At this point the calculations were stopped and a check was made to make sure that the program reached AA,,, in no less than 50 steps. In case this condition was not met, the calculations were repeated with a more densely set array of time points. In this manner it was ensured that the program did not miss the largest value of AA during a given run. Calculations were performed for the set of values of e, and x, used by Brill and Williams [ref. (lo), table 13. The quantity m,,, = AAmax/eo was evaluated and compared with the experimental results (column 3 of table 1, vide infra).The parameters were then varied in order to reach agreement between the calculations and the experiment. The standard optimization procedure, based on the method of ‘least residual squares’, was applied. In order to avoid the simultaneous variation of all parameters, which makes such calculations extremely difficult, the followingM. L. KREMER 93 procedure was adopted. First, it was assumed that at most two absorbing species (C,,C,) were formed in the system and that the deviations of their respective molar absorption coefficients from that of catalase were AE, and A E ~ . Next, the expression log m,ax(calc) = log + log ([PI + (AE~/AEJP~I/~OI (1) was evaluated. p1 and p 2 are the respective concentrations of C, and C,.By choice, C, was taken to be the species responsible for the major part of the absorbance change. [Eqn (1) could be simplified in cases in which the absorbance change was attributed to a single species only.] In a series of calculations AE,/AE, was held constant, whereas the factor log AE, appeared as an additive constant, which did not influence the shape of the log~,,x(calc) against x, curves. (Strictly speaking, the data of Brill and Williams do not lie on a single curve, since e, was not held strictly constant in their experiments, but they can be said to do so in an approximate sense.) Regarding the group of rate-constant parameters, their number could always be reduced by one, since the value of the computed maxima depended only on the ratios of rate constants and not on their absolute values.Consequently, the value of one of them (chosen as the rate constant of formation of the principal intermediate) could be arbitrarily fixed ( I dm3 mol-1 s-l). The other rate constants were then evaluated relative to this reference value. As the mechanisms considered involved, at most, three independent rate constants, the usual fitting procedure consisted of the adjustment of two rate parameters and of a coordinate shift parameter, representing log A&,. In this manner a partial minimum, corresponding to the chosen value of AE~/AE,, was found. The calculations were repeated for different values of AE~/AE, until a true minimum was found. The converged values of the parameters were then further refined and their standard deviations determined using the BMDPAR program of the BMDP computer program library.DERIVATION OF THE MECHANISM The criterion for a successful mechanism was seen in an explanation of the fact that as x, increased at constant e,, the values of AA,,, (reached in each experiment) passed through a maximum (AAMAX). First, the assumption was made that the phenomenon was caused by ethanol contamination of the ethylhydrogen peroxide reagent. Assuming a minimal set of reactions involving the formation and decomposition of an intermediate C, and its additional reaction with C2H50H, it was found that such a scheme could not account for the existence of AA,,,. An increase of the fraction of ethanol in the reagent caused a decrease of AA,,, at all values of x,, without causing the appearance of AAMAx.Next, catalatic-type reactions were tried, i.e. schemes which involve a reaction of CI with C2H500H to free catalase haematin. These reactions also failed to predict the existence of AAMAX. (An increase in the rate of freeing catalase increased the rate of its combination with C2H,00H, so that there was no net effect on [C,].) This result is in accordance with the lack of 0, evolution in the system.? These negative results indicated that the maximum was probably caused by an irreversible inactivation of catalase during the reaction with C,H,OOH. It could occur either via a direct reaction with C, or through a reaction involving a precursor of C,. First, the direct reaction of C, (I) with C,H,OOH was investigated. The details of this mechanism [mechanism (A)] are given in Appendix A.As this mechanism failed ZC,H,OH + 0,. t A catalatic reaction is equivalent to the disproportionation of C,H,OOH : 2C,H,OOH +94 REACTION OF CATALASE WITH C,H,OOH to account for the existence of AA,,,, the other possibility, involving a precursor complex, was investigated. The mechanism considered consisted of the following steps : where C,,, denotes a precursor complex to C,, C,, is an (inactive) biperoxy derivative of catalase haematin and lower-case letters denote concentrations. By assuming a rapid equilibrium between E, C,H,OOH and CPRC and further that at any instant only a small fraction of E is tied up in CPRC, we can write the rate equations as follows : = k;(e - p - q) (x, - p - 2q-S) - k , p dt dq - = ki(e - p - q) (x, - p - 24 - s), dt ds dt = k2P - (3) (4) where ki = kJK, kj = k3/K and The results of computer simulations based on mechanism (B) are summarized in table 1 and shown in fig.1. In fig. 1 , AAmax/eO is plotted as a function of the ratio of the concentrations of C,H,OOH and catalase haematin. Both the experimental and the calculated data show that AAmax/e, first increases then decreases with increasing x,/e,, having a maximum at x,/e, z 55. The rate constants were determined only up to a common multiplication factor (see Appendix B). The following values were calculated relative to the chosen reference of k; = 1 dm3 mol-l s-l: k, = (4.16k0.72) x s-l, ki = (8+ 12) x lo2 dms rnol-, s-l and eE-cI = (8.02f0.65) x lo4 dm3 mol-1 cm-l. The data show that while k, and cI, referring to the main catalytic path and intermediate, could be determined with reasonable accuracy, the accuracy achieved in the determination of ki and cBp, being representative of a minor branch of the reaction path under the experimental conditions, was considerably worse. The value of cE-cBP could not be determined with any degree of accuracy. The calculated standard deviation was much larger than the value itself [cE-cBP = - (8 & 70) x lo3 dm3 mol-l cm-l].It may be assumed that cBP has a value very near cE itself at 1 = 405 nm. Thus, according to the present explanation, the decrease of AAmax/eO at high x,/e, is due to the increased formation of CBP. As EE-EBP is much smaller than cE -c1 (and possibly has a negative sign), the formation of CBp causes an increase in the absorbance.In order to determine the absolute values of the rate constants, at least one AA against time curve must be known. Brill and Williams made a recording of AA as K = (e, - p - 4) (x, - p - 2q - s)/z.M. L. KREMER " 5- E d I 0 4 E E P 4 - E! m 0 --- n 0 u -. X il 9 'f 95 0 6 6 @ 3 - 3 I I I I I I I I 10 20 30 40 50 60 70 80 90 Fig. 1. Plot of AAmaX/eo as a function of xo/eo: 0, experimental results; +, calculated results. Experimental data of Brill and Williams [ref. (lo), table 13; T = 25 "C, pH 7.25 and A = 405 nm; calculations based on mechanism (B); k', = 1 dm-3 mol-l s-l, k, = 4.16 x s--l , and E ~ - E ~ ~ = ki = 7.87 x 10, dms rnol-, s-l, E ~ - E ~ = 8.02 x lo4 dm3 mol-l cm-l -7.9 x lo3 dm3 mo1-I cm-l. Table 1.Maximum optical adsorption decrements per unit concentration of catalase haematin. Data of Brill and Williams [ref. (lo), table 11. pH 7.25, T = 25 "C, A = 405 nm. Parameters used in calculations: k', = 1 dm3 mol-1 s-l, k, = 4.16 x lop5 s-l, ki = 7.87 x 10, dms rnol-, s-l, E ~ - E ~ = 8.02 x lo4 dm3 mol-l cm-l and E ~ - E ~ ~ = -7.9 x lo3 dm3 mol-1 cm-l 5.04 32 5.08 98 6.36 29 5 5.25 286 6.32 560 3.05 4.93 5.30 5.43 4.80 3.05 4.93 5.36 5.37 4.80 a function of time using e, = 6.32 pmol dmp3 and x, = 295 pmol dm-3 [fig. 1. of ref. (lo)]. No entirely satisfactory simulation of this curve could be achieved by taking the above set of relative rate constants and absorption coefficient parameters and adjusting the scaling factor of the rate constants.It is thought that the initiation of the reaction by Brill and Williams was not fast enough to ensure instantaneous mixing of the reactants. In fact, 65% of the total absorbance change had already occurred before the first observation was made (10 s). This factor may have influenced some96 REACTION OF CATALASE WITH C,H,OOH of their absorbance readings. Therefore, their measurement of the total absorbance change is considered to be more accurate than their recording of its temporal behaviour . Approximate values of the absolute rate constants which optimized, within limitations, the fit between the experimental and calculated curves, were as follows : k; = 1.7 x lo2 dm3 mol-1 s-l, k, = 7.1 x s-l and ki = 1 x lo5 dm6 rnol-, s-l.DISCUSSION The overall decomposition of C,H,OOH is, in effect, a dehydration process C,H,OOH -+ CH,CHO + H20 and it is remarkable that it is catalysed by an oxidation-reduction catalyst like catalase. With the participation of a catalyst, however, the dehydration reaction can be resolved into steps which do involve oxidation and reduction a C,H,OOH + E -+ EO + C,H,OH b EO + C,H50H -+ E + CH3CH0 + H,O or C C,H,OOH + E -+ {EO * C,H,OH} d (EO - C,H,OH)-+ E + CH3CH0 + H20. Mechanism (B) involves steps (c) and (d) as the path of catalysis, but an alternative mechanism, based on reactions (a) and (b) may also be envisaged: K E + C,H,OOH CpRC kz C; + C,H,OH -+ E + CH,CHO + H20 A series of calculations, based on mechanism (C), has been carried out. They have shown that the residual square sum, defined as {S = Z[logbAmax(calc.) - log aA,,,(e~ptl)]~}, was considerably higher for mechanism (C) (1.42 x lo-*) than for mechanism (B) (4.49 x became 1.37 x lo5 dm3 mol-1 cm-l, which implies a negative molar absorption coefficient for compound (I).[The molar absorption coefficient of free catalase haematin at 405 nm is 1.01 x lo5 dm3 mo1-l cm-l, from fig. 2 of ref. (lo)]. From this failure of mechanism (C), in contrast to mechanism (B), to provide an adequate basis for an interpretation of the experimental data we conclude that C,, rather than C;, is the correct description of compound (I), i.e. that compound (I) from C,H500H retains C,H,OH as part of its structure. This conclusion agrees with the statement of Nicholls and Schonbaum Furthermore, the converged value of E~M.L. KREMER 97 that no evidence has been found for the liberation of alcohol upon the formation of compound (I).3 Mechanisms (B) and (C) contain the essential elements of scheme (11) of Schonbaum and Chance with path 2 or 3 ~ p e r a t i n g : ~ ~ E+CH,CHO+H,O 1 E+CH,CH,OOH e ( E . C H , C H , O O H ) EO + CH,CH,OH scheme (I) E+CH,CH,OOH + (E.CH,CH,OOH) + (EO.CH,CH,OH) ‘ k / k3 E + CH,CHO + H,O E O+ CH,CH,OH scheme (11) The present analysis shows that, at least in the case of C,H,OOH, only path 2 of scheme (11) is operative. (Also the reverse transformation EO * CH,CH,OH -, E.CH,CH,OOH, questioned by Schonbaum and Chance, is now ruled out.) The precursor complex C,,, = E * CH3CH200H becomes significant kinetically only at high x,, where it presents a point of branching off from the catalytic path to form the inactive intermediate CBp.Another question concerns the extent of conversion of catalase into compound (I) during the reaction. To investigate this point a series of computer runs was performed at a constant e, and at increasing x,. In each of these runs pmax and the simultaneous values of q and of free catalase haematin remaining in the reaction mixture were determined. (Note that p and AA reach their maxima at different times during the reaction.) Calculations were based on mechanism (B) and on the set of (relative) rate constants obtained in the previous section. The results obtained at e, = 5 pmol dmW3 are shown in fig. 2. Fig. 2 shows that as x, increases, the values of pmax pass through a maximum (pMAX).The values OfpMAX obtained at different e, are shown in table 2, where qMAX, and [&AX, are the concentrations of the respective species when p = PMAX. The data in table 2 show that there is an inherent limitation to the extent of formation of C,. When e, is in the range 2.5 - 10 pmol dm-3 only cu. 69% of catalase haematins can be converted under optimal conditions into compound (I). Cu. 18% of the haematins are in the inactive biperoxy complex, while nearly 13 % of the haematins are free. An important point to be considered concerns the fact that the value of p at the maximum of AA is always less than pmax, since at the maximum of AA, p is already decreasing. Thus, the values of p, corresponding to AAmax, are even lower than those given by the curve in fig.2. By adding 0.085 cm3 0.013 mol dm-3 KCN to a ‘steady-state mixture’ (i.e. when AA has reached its maximum) of 3.00 cm3 4.2 pmol dmP3 catalase haematin and 0.052 cm3 0.017 mol dm-3 C2H,00H, Brill and Williams found that cu. 5-10% of catalase haematins were free. A separate computer simulation was run with the appropriate initial concentrations of the above experiment (e, = 4.17 pmol dm-3, x, = 290 pmol dm-3). The calculations were stopped when AA reached its maximum98 REACTION OF CATALASE WITH C,H,OOH 1 I00 xo /pmol dm-3 200 300 400 Fig. 2. Maximum concentration of compound (I) as a function of x,; e, = 5 pmol dm-3; mechanism (B), rate parameters as in fig. 1. rable 2. Maximum formation of compound (I) [calculated on the basis of mechanism (B: ki = 1 dm3 mol-l 0, k, = 4.16 x lo5 s-l and k; = 7.87 x 10, dms rnol-, s-l e0 X O PMAX QMAX, p p PMAX QMAX.p IE1MAX, /pmol dmW3 /pmol dmP3 /pmol dm-3 /pmol dm-3 /pmol dm-3 / e , l e , /e0 2.5 230 1.734 0.447 0.323 0.694 0.179 0.129 5 23 5 3.464 0.905 0.63 1 0.693 0.181 0.126 10 240 6.914 1.821 1.265 0.691 0.182 0.127 value. At this point the following results were obtained: p = 2.873 and q = 0.868 pmol dm-3, giving [Elfree = 0.429 pmol dm-3, i.e. 10.3% of e,. This result is in excellent agreement with the above estimate of Brill and Williams.? The decrease of the molar absorption coefficient accompanying the change E + C, was found to be 8 . 0 2 ~ 104dm3mol-1cm-1 at d =405nm. Thus, E, becomes 2.1 x lo4 dm3 mol-1 cm-l, which is considerably lower than the value of 4.5 x lo4 dm3 mol-l cm-l given by Brill and Williams (their fig.2). The difference stems from the different values of the degree of conversion of E into C, used in the two calculations. Brill and Williams, on the basis of their CN- binding experiments, used a value of 92.5 % as the degree of conversion of catalase, by ignoring any species other than E or CI in the system. The present value also differs from that given by Chance [4.9 x lo4 dm3 mol-1 cm-l, ref. (8), fig. 2, calculated on haematin basis], who assumed 100% conversion into C,. t The results of Brill and Williams should be corrected slightly upward, if the finite dissociation constant of the E-CN complex [KECN = 22 pmol dm-3, ref. (9)] is considered. Thus, for example, a calculated value of 7.5% free catalase haematin should become 8%.M.L. KREMER 99 c a t a l y t i c range non-catalytic range I < n 3 50 100 I 5 0 xo /pmol dm-' Fig. 3. Concentrations of the components of the reaction mixture at the end of the reaction; eo = 5 pmol dm-3; mechanism (B), k; = 1 dm3 mol-1 s-l, k, = 4.33 x s-' and kj = 1.17 x lo3 dms rnol-, s-l. 50 100 150 xo/pmol dm-3 c a t a l y t i c range n o n - c a t a l y t i c range , Fig. 4. Relative concentrations of the components of the reaction mixture at the end of the reaction; e = 5 pmol dm-3; mechanism (B), all parameters as in fig. 3.100 REACTION OF CATALASE WITH C2H,00H The decrease of the molar absorption coefficient of bacterial catalase upon forming the analogous ‘Compound (I)’ species with H202 is 2.5 x lo4 dm3 mol-l cm-l at 1 = 405 nm.15 It differs substantially from AcI = 8.02 x 104 dm3 mol-1 cm-1 calculated above for the catalase-thylhydroperoxi-compound (I).This result does not support the view that there is an identical depression of the Soret band absorption of catalase upon forming different compound (I) species with alkyl hydrogen peroxides and with H202.8 According to the present discussion [retainment of C2H50H by ethylhydroperoxi-compound (I) and differences in the spectra], it appears that ‘compound (I) ’ species obtained with different substrates are not identical. The various steps and rate constants describe only the first part of the reaction, up to the maximum of the decrease in the absorbance. In spite of the inherent uncertainty, it is of interest to discuss the predicted course of the reaction in its later stages.In the following calculations it will be assumed that E~ - E ~ , has a small positive value. This assumption is not excluded by the above results and seems to agree with observations made on the system during the later phases of the reaction. (The parameters used in the calculations are those obtained before the BMDPAR refinement procedure. There has been only a slight change of the rate parameters as a result of this refinement so that the essential features of the following results remain unaffected by it.) In fig. 3 and 4 the final concentrations of the various constituents of the reaction mixture are plotted as a function of x,, and s,, x, and qe denote the final concentrations of CH3CH0, C,H,OOH and CBP, respectively. The ‘end of the reaction’ was defined here as a state in which either x became < 0.1 % of x, (low x, range), or the concentration of free catalase haematin falls to < 0.01% of e, (high x, range).(Under these conditions no substantial changes in the system could further take place.) Fig. 3 shows the absolute concentrations of the various components of the system at varying x,. In fig. 4 the concentrations relative to x, and e,, respectively, are given. Fig. 3 and 4 show that around x, = 100 pmol dm-3, where the formation of C,, becomes, to a good approximation, quantitative, there is a striking change in the distribution of the products of the reaction. There is an abrupt decrease in the final concentration of aldehyde (s,).Its formation is not quantitative even at x, < 100 pmol dm-2 (its curve of formation lags behind the straight line r which corresponds to the quantitative formation of CH3CH0 from C2H500H), but at x, = 100 pmol dm-3 there starts a sharp decline of s, as x, is further increased. At the same time, peroxide appears in the final state of the system in increasing concentrations. If we regard the reaction as essentially catalytic when > 90% of C2H,00H is converted to CH3CH0, then there is a sharply defined region when the reaction is essentially catalytic (x, < 100 pmol dm-3) and another region (x, > 100 pmol dmP3) where it is essentially non-catalytic. At low x,, the formation of CBp in the system can be neglected. Under these circumstances, the absorbance of the system will return at the end of the reaction to that of the free enzyme [fig.5, curves (a) and (b)]. Note the relatively flat maxima of curves (a) and (b), which in some interpretations of the corresponding experimental curves were thought to represent a ‘steady state, existing for a limited amount of time’. In reality, the system is far removed from any such situation. As noted above, the system is also not in an equilibrium. Consequently, the initial part of the curves cannot be treated as being a pre-steady state or an approach to an equilibrium or to represent quantitative bimolecular formation of compound (I). Any evaluation of the rate constant for the formation of CI on these assumptions must lead to erroneous r e s ~ l t s .~ ~ l1 Considering the run with x, = 5 pmol dm3, the curve x(a) shows that there is a considerable drop in the concentration of ethylhydrogen peroxide until the maximumM. L. KREMER 101 0.20 0 15 - I 5 2 0.10 Q 0 0 5 \ I 3 4 m 3 2 - 0 2 3 Y I I I I I t l s Fig. 5. Time dependence of AA at various x,; e, = 5 pmol dmP3; x, = (u) 5, (b) 10, (c) 20 and ( d ) 50 pmol dm-3; mechanism (B), k; = 6.0 x 10, dm3 mol-l s-l, k , = 2.6 x lop2 s-l, kh = 7.0 x lo5 dm6 mo1F2 s-l, C ~ ; - - E ~ = 8.17 x lo4 dm3 mol-l cm-' and E~~ = 8.2 x lo3 dm3 mol-l cm-'. 100 200 300 400 500 decrement of the absorption is reached (ca. 90 s). According to the calculations, ca. 77% of the original ethylhydroperoxide are still present. This feature of the reaction differs from that observed with the catalase-H,O, system, where at the maximum of AA, [H,O,] is practically nil.,? 18* l9 As x, is increased, the concentration of C,, in the system increases. Beyond a certain value of x,, the absorbance at the end of the reaction will not return to that of free catalase, but will remain at some residual value corresponding to the permanent typing up of part of the catalase haematins in the inactive biperoxy complex [curves In general, the calculated curves of fig.5 agree qualitatively with the behaviour of various alkylhydroperoxikcatalase systems, at low as well as at high initial concen- trations of the alkylhydroperoxides.ll9 l6 There is an interesting parallelism between the conditions of formation of C,, and of compound (11) (high x, and long reaction times).16 There may thus exist a relationship between these two species.? It should also be added that the present analysis cannot distinguish between the existence of an actual biperoxy complex or (4 and (41.t In the case of H,O,, the rate constant of reaction of the precursor complex with H,O, is probably very low. In the course of many catalytic cycles C,, may, however, accumulate. These are exactly the conditions under which compound (11) is observed in the catalase-H,O, system. Repeated cycles of reaction are provided here by continuously generating H,O, with the notatin-glucose-0, system.''102 REACTION OF CATALASE WITH C,H,OOH of some products of a reaction between the constituents of this complex, as long as the enzyme remains in an inactive form. Summarizing, the present investigation shows that the catalase-ethylhydrogen peroxide system is complex and that it is not possible to convert catalase quantitatively, under any circumstances, into compound (I).(At low x, the concentration of the alkylhydroperoxide is insufficient to saturate catalase, and at high x, part of the catalase is converted into the inactive biperoxy complex.) The formation of compound (I) can, however, be optimized and (at least some of) its physical constants can be determined. APPENDIX A The reactions comprising mechanism (A) are as follows: ki E +C,H,OOH 4 C, (e0-P-q) X ke C, --+ E + CH,CHO + H20 P s k3 C, + C2H5OOH + CBp. 4 At low x , step (2) predominates over step (3): the scheme is then a catalytic mechanism for the decomposition of C2H500H with a gradual inactivation of the catalyst through step (3).By increasing x , the catalytic step (2) becomes suppressed and the scheme becomes a two-step mechanism for the formation of C,, via the intermediate C,. We are interested in an analytical solution of the rate equations at high x (compared with eo). This condition is fulfilled in most of the experiments. By assuming that x remains constant during the reaction at its initial value (xo) and by introducing the variable u = eo - q, we can write the following matrix representation of the rate equations: This is a set of simultaneous linear first-order differential equations whose solutions are and The reduced absorbance decrement 6 = A A / ( E ~ - E , ) can be written as 6 = p + t r q (A 4) where E, = ( E ~ - E , ~ ) / ( E ~ - E ~ ) .By introducing the expressions for p and q from eqn (A 2) and (A 3) into (A 4), we obtain 6 = Er eo + (eo/ v'B) [(kl xo + Er 12) ~ X P (4 1) - (k1 xo + Er A) ~ X P (22 t)l- (A 5 ) By differentiating 6 with respect to time, and equating the time derivative to zero, we obtain the following expression for t,,,, the time when 6 has reached its maximum: In(=)M. L. KREMER 103 where u = k,/k,, a = t [ v + ( v 2 - 4a):] b = k,/k,, u = 1 +a+ (b/xo), and p = t [ u - (u2 - 4a):I. By introducing t,,, into eqn (A 5 ) we obtain an expression for 6,,, 6max = Er eo + S,,, depends on the ratios of rate constants a and b and on xo. By differentiating 6,,, with respect to u (and hence with respect to x,) we obtain where 2+2 2(1 +z) F(z) = z-'ln(l+z)-- and z = (a-p)/(P-a&,). It follows from eqn (A 9) that - 1 is the lower limit of z.Since tmax must be real and positive, (a-m,)/(P-ae,) must be positive (and greater than one). F(z), on the other hand, is always negative, in the allowed range of z. Thus, dS,,,/du is always negative, and because of the inverse relationship between xo and u, d6,,,/dx0 is positive. Mechanism (A) predicts, therefore, a monotonous increase of a,,, with x,, in contradiction to the experiment. To complete the proof, direct numerical integrations of the rate equations were carried out under conditions in which constancy of x during a run could not be assumed. The computations have shown that for all values of rate constants and extinction coefficients tried, 6,,, increased monotonously with increasing xo, in the same way as found in the analytical solution.APPENDIX B The reduced absorbance decrement is given by 6 = p+crq. The condition for the maximum of 6 is given by dP dq --f&,- = 0 dt dt which can be written in the alternative form Since at finite times dq/dt > 0, the condition for the maximum becomes dP -+&, = 0. d4 dp/dq can be obtained from eqn (1) and (2) and s can also be expressed as a function of q Eqn (B 5 ) and (B 6) are first-order linear differential equations for the calculation of p and s104 REACTION OF CATALASE WITH C,H,OOH as functions of q. From the form of these equations it can be deduced that the solution will depend on the ratios of the rate constants but not on their absolute values. Prof. Shalom Baer and Prof. Maurice Cohen are thanked for helpful discussions. B. Chance, in The Enzymes, ed. J. B. Sumner and K. Myrback (Academic Press, New York, 1951), vol. 2, part I, p. 482. B. Chance, in Investigation of Rates and Mechanisms of Reactions, ed. S . L. Friess, E. S. Lewis and A. Weissberger (Interscience, New York, 1963), part 11, p. 1314. P. Nicholls and G. R. Schonbaum, in The Enzymes, ed. P. D. Boyer, H. Lardy and K. Myrback (Academic Press, New York, 1963), vol. 8, p. 147. A. S. Brill, in Comprehensive Biochemistry, ed. M. Florkin and E. Stotz (Elsevier, Amsterdam, New York, 1966), vol. 14, chap. X, p. 447. G. R. Schonbaum and B. Chance, in The Enzymes, ed. P. D. Boyer (Academic Press, New York, 3rd edn, 1976), vol. XIII, p. 363. K. G. Stem, J. Biol. Chem., 1936, 114, 473. D. Keilin and E. F. Hartree, Proc. R. SOC. London, Ser. B, 1936, 119, 141. B. Chance, J. Biol. Chem., 1949, 179, 1331. B. Chance and D. Herbert, Biochem. J., 1950,46,402. lo A. S. Brill and R. J. P. Williams, Biochem. J., 1961, 78, 253. l1 B. Chance, J. Biol. Chem., 1949, 179, 1341. l2 B. Chance and G. R. Schonbaum, J. Biol. Chem., 1962,237,2391. l3 G. R. Schonbaum, Wenner-Gren Symp. Struct. Funct. Oxidation-Reduction Enzymes, 1970, p. 48. l4 G. R. Schonbaum, 9th Int. Congr. Biochem., 1973, p. 49. l5 M. L. Kremer, Isr. J. Chem., 1975, 13, 91. l6 B. Chance, J. Biol. Chem., 1949,180, 865. B. Chance, J. Biochem., 1950,46, 387. B. Chance, Acta Chem. Scand., 1947, 1, 236. lS E. Zidoni and M. L. Kremer, Arch. Biochem. Biophys., 1974, 161, 658. (PAPER 4/41 3)

ISSN:0300-9599    

DOI:10.1039/F19858100091

出版商:RSC    

年代:1985

数据来源: RSC