摘要:

physicochemical topics, thereby encouraging scientists of different disciplines to contribute their varied viewpoints to a coiiinion theme. A recent Discussion is :- The Royal Soci- of Chemistry- No.75 lntraamolecwlar Kinetics No. 75 in the series, this publication is the result of a general discussion held at the University of Warwick in April 1983. Contents: The Spiers Meniorlal Lecture; Vibrational Redistribution within Excited Electronic States of Polyatomic Molecules Inrraniolecular R e h u t i o n o f 1. vcited States lsomerization of Intcrnal~ncrgy-selected Ions Kinetics of Ion-Molecule Collision Coinple\es in the Gas Phase, E\periinent and Theory lntrainolccular Decay 01' Soinc Open-shell Pulya t o niic Ca lions On tlic Theory u i Iiitrdniolccul~r I n e r g y Transfer Pulsed Laser Preparation and Ouaniuin Superposition Statc Evolution in ReguLtr and Irregular Systems A Ouantuiii-iiicclianical Internal-collision Model for State-sclcctcd Uniinolccular Decoiiiposilio n The Correspondence Principle and Intramolecular Dynamics lntrainoleculdr Dcphasiiig.t'icusecond Evolution of Wavepacket States in a Molecule with Int erinediate-casc level Struct urc Energy Conversion in van der Waals C'u~~iplc\c\ ol s-Tetrarine and Argon Tim-dependent Processes in Polyatuinic Molecules During and After Intense Intrarcd Irradiation Energy Distributions in tlic (.N(X'L+) bragnient froiii tlie Infrared Multiplepholun Dissociation ol' CI. ICN. A Coinparison between 1:xperiiiiental Results and the Predictions ot Statistical Theories of ChFO + Product Energy Partitioning in the Decoiii- position of State-selectively Excited HOON and IIOOD Low-power Inl-rarcd Laser I'hoiolysis o f Tetramethy ldioxetan Uniinolecular Reactions lnduccd by Vibrational Overtone Excitation Uniiiiolecular Decomposition of t-Butylhydro- peroxide by Direct Excitation of the 6-0 0-11 Stretching Overtone I'icosecond-jet Spectroscopy and Photoclieinistry.Energy Redistribution and its Iiiipact'on Coherence, Isoincrization, Ihssociatiun and Solvalioii knergy Redistribution in Large Molecules. Duect St ud y o f In1 rainolucular Rehxa lion in the Gas Phase with Picosecond Gating Rotation-dependent Intrainolecuhr I'r~)cessc.sofSO:(A'A.) in a Superwnic Jct Role of Rotation-Vibration Interaction in Vibrational Keh\ation. Energy Kcdistribution in k,xcitcd Singlet I~'ornialdc1iyde Sub-lhppler.Spectroscopy of Benrcnc in the "('liaiinel-lliree" Region Intraiiiulccular 1:lectronic Kclau~tion and I'liotois~)iiieruati[)n Processes in tlie lsuhted Azabenrene Molecules Pyridinc, Pyrazinc and I'yriiiiidinc Softcover 434pp 0 85186 658 1 Price f25.00 ($48.00) Rest of the World f26.00 RSC Members f 16.25 Faraday Discussions of the Chemical Society 7< lnrruniolei u h r Kincrit I Faraday Symposia are usually held annually and are confined to more specialiscd topics than Discussions, with particular reference to recent rapidly developing lines of rescuch. A recent Symposium is :- No.l?The Hydrophobic Interadion No. 17 in the series, this publication is the result of a symposium on The Hydrophobic Interaction held at the Uiiiversity of Reading in December 1982.Contents: Hydrophobic Interdctionr a llistaric.11 Per spect ivr llydrupliobic Ilydration Geometric Kelaution in Water. Its Role in Precise Vapour-pressure Measureiiients of the SolubilkdtiorI of Benzene by Aqueous Sodiuiii Octylsulphate Solutions Nuclear Magnetic Resonance R e b u t i o n Investigation of Tetrahydrofuran and Methyl Iodide Clathrdtes Infrared and Nuckar Magnetic Kcwnance Studies Pertaining to the (age Model t o r Solutions oS Acetone in Water Irothernial Transport Properties in Solutions o f S y mmet r ica I Tet ra-alk y hmnioniuiii Bromides Thermodynamics of Cavity I'oriiiaiion in Water. A Molecular Dynamics Study Molecular Librations and Solvent Oricnt- ational Correlations in Hydrophobic Phenomena Monte Carlo Computer Siniulation Study of the Hydrophobic Effect.Potential ot Mean Force for ECfir)gaq at 25 and SOv C Hydroplicibic Moments and Protein Structure Application 01' the Kirkwood-Buff Theory to the t'roblcin 01 Hydrophobic Interactions Ihentangleinent of Ilydrophubic and IFlcctrosi~tic Contributions t o the I.ilni Pressures O i Ionic Surfactants llydrophobir. Intcracliuns in Dilute Su lut io ns u t 1'0 1 y (vin y I a Ico lio I) ('onioriii;tiionaI and 1:unc.i ional I'ropertics of tiaeiiwglobin in Water+Alcohol Mixtures. Dependence o f Bull. Electrostatic and tlydrupliohic I n t c r x t i o n s upon ptl and KCI concentrations Softcover 24Opp 0 85186 668 9 Price f36.50 ($70.00) Rest of the World f38.50 RSC Members f 23.75 ORDERING RSC Members should send their orders to: The Royal Society of Chemistry.The Membership Officer. 30 Russell Square, Non-RSC Members The Royal Society of Chemistry, Distribution Centre, Blackhorse Road, L London WC1 B 5DT. Letchworth, Herts SO6 IHN, England. Faradaj Symposia of the Chemical Society hGi 17 I hc HI drophohr' Inrcrm rron 1 9 X ? (viii)physicochemical topics, thereby encouraging scientists of different disciplines to contribute their varied viewpoints to a coiiinion theme. A recent Discussion is :- The Royal Soci- of Chemistry- No.75 lntraamolecwlar Kinetics No. 75 in the series, this publication is the result of a general discussion held at the University of Warwick in April 1983. Contents: The Spiers Meniorlal Lecture; Vibrational Redistribution within Excited Electronic States of Polyatomic Molecules Inrraniolecular R e h u t i o n o f 1.vcited States lsomerization of Intcrnal~ncrgy-selected Ions Kinetics of Ion-Molecule Collision Coinple\es in the Gas Phase, E\periinent and Theory lntrainolccular Decay 01' Soinc Open-shell Pulya t o niic Ca lions On tlic Theory u i Iiitrdniolccul~r I n e r g y Transfer Pulsed Laser Preparation and Ouaniuin Superposition Statc Evolution in ReguLtr and Irregular Systems A Ouantuiii-iiicclianical Internal-collision Model for State-sclcctcd Uniinolccular Decoiiiposilio n The Correspondence Principle and Intramolecular Dynamics lntrainoleculdr Dcphasiiig. t'icusecond Evolution of Wavepacket States in a Molecule with Int erinediate-casc level Struct urc Energy Conversion in van der Waals C'u~~iplc\c\ ol s-Tetrarine and Argon Tim-dependent Processes in Polyatuinic Molecules During and After Intense Intrarcd Irradiation Energy Distributions in tlic (.N(X'L+) bragnient froiii tlie Infrared Multiplepholun Dissociation ol' CI.ICN. A Coinparison between 1:xperiiiiental Results and the Predictions ot Statistical Theories of ChFO + Product Energy Partitioning in the Decoiii- position of State-selectively Excited HOON and IIOOD Low-power Inl-rarcd Laser I'hoiolysis o f Tetramethy ldioxetan Uniinolecular Reactions lnduccd by Vibrational Overtone Excitation Uniiiiolecular Decomposition of t-Butylhydro- peroxide by Direct Excitation of the 6-0 0-11 Stretching Overtone I'icosecond-jet Spectroscopy and Photoclieinistry. Energy Redistribution and its Iiiipact'on Coherence, Isoincrization, Ihssociatiun and Solvalioii knergy Redistribution in Large Molecules.Duect St ud y o f In1 rainolucular Rehxa lion in the Gas Phase with Picosecond Gating Rotation-dependent Intrainolecuhr I'r~)cessc.sofSO:(A'A.) in a Superwnic Jct Role of Rotation-Vibration Interaction in Vibrational Keh\ation. Energy Kcdistribution in k,xcitcd Singlet I~'ornialdc1iyde Sub-lhppler. Spectroscopy of Benrcnc in the "('liaiinel-lliree" Region Intraiiiulccular 1:lectronic Kclau~tion and I'liotois~)iiieruati[)n Processes in tlie lsuhted Azabenrene Molecules Pyridinc, Pyrazinc and I'yriiiiidinc Softcover 434pp 0 85186 658 1 Price f25.00 ($48.00) Rest of the World f26.00 RSC Members f 16.25 Faraday Discussions of the Chemical Society 7< lnrruniolei u h r Kincrit I Faraday Symposia are usually held annually and are confined to more specialiscd topics than Discussions, with particular reference to recent rapidly developing lines of rescuch.A recent Symposium is :- No.l?The Hydrophobic Interadion No. 17 in the series, this publication is the result of a symposium on The Hydrophobic Interaction held at the Uiiiversity of Reading in December 1982. Contents: Hydrophobic Interdctionr a llistaric.11 Per spect ivr llydrupliobic Ilydration Geometric Kelaution in Water. Its Role in Precise Vapour-pressure Measureiiients of the SolubilkdtiorI of Benzene by Aqueous Sodiuiii Octylsulphate Solutions Nuclear Magnetic Resonance R e b u t i o n Investigation of Tetrahydrofuran and Methyl Iodide Clathrdtes Infrared and Nuckar Magnetic Kcwnance Studies Pertaining to the (age Model t o r Solutions oS Acetone in Water Irothernial Transport Properties in Solutions o f S y mmet r ica I Tet ra-alk y hmnioniuiii Bromides Thermodynamics of Cavity I'oriiiaiion in Water.A Molecular Dynamics Study Molecular Librations and Solvent Oricnt- ational Correlations in Hydrophobic Phenomena Monte Carlo Computer Siniulation Study of the Hydrophobic Effect. Potential ot Mean Force for ECfir)gaq at 25 and SOv C Hydroplicibic Moments and Protein Structure Application 01' the Kirkwood-Buff Theory to the t'roblcin 01 Hydrophobic Interactions Ihentangleinent of Ilydrophubic and IFlcctrosi~tic Contributions t o the I.ilni Pressures O i Ionic Surfactants llydrophobir. Intcracliuns in Dilute Su lut io ns u t 1'0 1 y (vin y I a Ico lio I) ('onioriii;tiionaI and 1:unc.i ional I'ropertics of tiaeiiwglobin in Water+Alcohol Mixtures. Dependence o f Bull. Electrostatic and tlydrupliohic I n t c r x t i o n s upon ptl and KCI concentrations Softcover 24Opp 0 85186 668 9 Price f36.50 ($70.00) Rest of the World f38.50 RSC Members f 23.75 ORDERING RSC Members should send their orders to: The Royal Society of Chemistry. The Membership Officer. 30 Russell Square, Non-RSC Members The Royal Society of Chemistry, Distribution Centre, Blackhorse Road, L London WC1 B 5DT. Letchworth, Herts SO6 IHN, England. Faradaj Symposia of the Chemical Society hGi 17 I hc HI drophohr' Inrcrm rron 1 9 X ? (viii)

ISSN:0300-9599    

DOI:10.1039/F198480FX021

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

AUTHOR INDEX xxxv Sabbatini, L., 1029 Sacco, A., 2669 Sanders, J. V., 571 Sangster, D. F., 291 1 Sarka, K., 521 Sasahira, A., 473 Sasse, W. H. F., 571 Satchell, P. W., 2395 Sato, K., 841 Sato, Y., 341 Savino, V., 759 Sayers, C. M., 1217 Schiller, R. L., 1257 Schmidt, J., 1 Schmidt, P. P., 2017 Schneider, H., 3275, 3285 Schulz, R. A., 489, 1323 Scott, J. M. W., 739, 1651, 2287, Scott, S. K., 3409 Segall, R. L., 2609 Sehested, K., 2929, 2969 Seidl, V., 1367 Sem, P., 297 Serratosa, J. M., 2225 Seyama, H., 237 Seyedmonir, S. R., 87, 2269 Shanahan, M. E. R., 37 Sheppard, A., 2999 Sherwood, P. M. A., 135, 2099, Shindo, Y., 879, 2199 Shiotani, H., 2145 Shizuka, H., 383, 341 Siekhaus, W. J., 61 Sircar, S., 1101, 2489 Smart, R. St C., 2957, 2609 Smith, I. M., 3021 Smith, R., 3233 Snow, R.L., 3463 Solar, S., 2929 Solar, W., 2929 Solymosi, F., 1841 Soma, M., 237 Soupart, J-B., 3209 Sourisseau, C., 3257 Spink, J. A., 3469 Spoto, G., 1875, 1891 Spotswood, T. M., 3147 Staricco, E. H., 2631 Stassinopoulou, K., 3095 Stedman, D. H., 285 Stout, D. R., 3481 Strohbusch, F., 1757 Strumolo, D., 1479 Struve, P., 813, 2167 Styring, M. G., 3051 Subramanian, R., 2405 2881, 3359 2549, 2867 Sundar, H. G. K., 3491 Sutcliffe, L. H., 669, 3021 Sutton, H. C., 2301 Sutton, L. E., 635 Suzuki, H., 803 Suzuki, T., 1925, 3157 Symons, M. C. R., 423, 1005, Szamosi, J., 1645 Szczepaniak, W., 2935 Takagi, Y., 1925 Takahashi, K., 803 Takahashi, N., 629 Takanaka, J., 941 Takao, S., 993 Takasaki, S., 803 Takegami, H., 1221 Tam, S-C., 2255 Tamamushi, R., 2751 Tamaru, K., 29, 1567, 1595 Tamilarasan, R., 2405 Tanabe, S., 803 Tanaka, K., 2563,2981 Tanaka, T., 119 Taniewska-Osinska, S., 1409 Tascon, J.M. D., 1089 Teo, H. H., 981, 1787 Tetenyi, P., 3037 Thomas, J. K., 1163 Thompson, L., 1673 Thomson, M., 1867 Thomson, S. J., 1689 Tiddy, G. J. T., 789, 3339 Tittarelli, P., 2209 Tominaga, T., 941 Tomkinson, J., 225 Tonelli, C., 1605 Toprakcioglu, C., 13,413 Tran, T., 1867 Trasatti, S., 913 Tripathi, A. D., 1517 Tronc, E., 2619 Troncoso, G., 2127 Truscott, T. G., 2293 Tsurusaki, T., 879 Tuck, J. J., 309 Turner, P. S., 2609 Tusk, M., 1757 Tvarbikova, Z., 2639 Tyrrell, H. J. V., 1279 Ueki, Y.. 341 Ueno, A., 803 Unno, H., 1059 Valencia, E., 2127 van de Ven, T. G. M., 2677 van Ommen, J. G., 2479 van Truong, N., 3275, 3285 Vargas, I., 1947 2767, 2803, 21 1, 1999 Vedrine, J.C., 1017 Veith, J., 2313 Velasco, J. R., 3429 Vesala, A., 2439 Vickerman, J. C., 1903 Vincent, B., 2599 Vinek, H., 1239 Vink, H., 507, 1297 Waghorne, W. E., 1267 Wagley, D. P., 47 Walker, R. W., 435, 3187, 3195, Wallington, T. J., 2737 Wang, G-W., 1039 Watkins, P. E., 2323 Watkiss, P. J., 1279 Watt, R. A. C., 489 Webb, G., 1689 Webster, B. C., 255, 267 Weiner, E. R., 1491 Wells, C. F., 2155. 2445 Wells, J. D., 1233 Whang, B. C. Y., 291 1 Whittle, E., 2323 Wichterlova, B., 2639 Wiesner, S., 3021 Wilhelmy, D. M., 563 Williams E. H., 3147 Williams, P. A., 403 Williams, R. J. P., 2255 Wokaun, A., 1305 Wolff, T., 2969 Wood, S. W., 3419 Woolf, L. A., 549, 1287 Wright, C. J., 1217 Wu, D. C., 1795 Wiirflinger, A., 3221 Wyn-Jones, E., 1915 Yamabe, M., 1059 Yamamoto, S., 941 Yamashita, H., 1435 Yamauchi, H., 2033 Yamazaki, A., 3245 Yariv, S., 1705 Yasumori, I., 841 Yeates, S.G., 1787 Yide, X., 969, 3103 Ylikoski, J., 2439 Yokokawa, T., 473 Yoneda, N., 879 Yonezawa, T., 1435 Yoshida, S., 119, 1435 Zambonin, P. G., 1029 Zanderighi, L., 1605 Zecchina. A., 2209, 2723, 1875, Zipelli, C., 1777 Zundel, G., 553 348 1, 2827 1891AUTHOR INDEX xxxv Sabbatini, L., 1029 Sacco, A., 2669 Sanders, J. V., 571 Sangster, D. F., 291 1 Sarka, K., 521 Sasahira, A., 473 Sasse, W. H. F., 571 Satchell, P. W., 2395 Sato, K., 841 Sato, Y., 341 Savino, V., 759 Sayers, C. M., 1217 Schiller, R. L., 1257 Schmidt, J., 1 Schmidt, P. P., 2017 Schneider, H., 3275, 3285 Schulz, R. A., 489, 1323 Scott, J. M. W., 739, 1651, 2287, Scott, S.K., 3409 Segall, R. L., 2609 Sehested, K., 2929, 2969 Seidl, V., 1367 Sem, P., 297 Serratosa, J. M., 2225 Seyama, H., 237 Seyedmonir, S. R., 87, 2269 Shanahan, M. E. R., 37 Sheppard, A., 2999 Sherwood, P. M. A., 135, 2099, Shindo, Y., 879, 2199 Shiotani, H., 2145 Shizuka, H., 383, 341 Siekhaus, W. J., 61 Sircar, S., 1101, 2489 Smart, R. St C., 2957, 2609 Smith, I. M., 3021 Smith, R., 3233 Snow, R. L., 3463 Solar, S., 2929 Solar, W., 2929 Solymosi, F., 1841 Soma, M., 237 Soupart, J-B., 3209 Sourisseau, C., 3257 Spink, J. A., 3469 Spoto, G., 1875, 1891 Spotswood, T. M., 3147 Staricco, E. H., 2631 Stassinopoulou, K., 3095 Stedman, D. H., 285 Stout, D. R., 3481 Strohbusch, F., 1757 Strumolo, D., 1479 Struve, P., 813, 2167 Styring, M. G., 3051 Subramanian, R., 2405 2881, 3359 2549, 2867 Sundar, H.G. K., 3491 Sutcliffe, L. H., 669, 3021 Sutton, H. C., 2301 Sutton, L. E., 635 Suzuki, H., 803 Suzuki, T., 1925, 3157 Symons, M. C. R., 423, 1005, Szamosi, J., 1645 Szczepaniak, W., 2935 Takagi, Y., 1925 Takahashi, K., 803 Takahashi, N., 629 Takanaka, J., 941 Takao, S., 993 Takasaki, S., 803 Takegami, H., 1221 Tam, S-C., 2255 Tamamushi, R., 2751 Tamaru, K., 29, 1567, 1595 Tamilarasan, R., 2405 Tanabe, S., 803 Tanaka, K., 2563,2981 Tanaka, T., 119 Taniewska-Osinska, S., 1409 Tascon, J. M. D., 1089 Teo, H. H., 981, 1787 Tetenyi, P., 3037 Thomas, J. K., 1163 Thompson, L., 1673 Thomson, M., 1867 Thomson, S. J., 1689 Tiddy, G. J. T., 789, 3339 Tittarelli, P., 2209 Tominaga, T., 941 Tomkinson, J., 225 Tonelli, C., 1605 Toprakcioglu, C., 13,413 Tran, T., 1867 Trasatti, S., 913 Tripathi, A.D., 1517 Tronc, E., 2619 Troncoso, G., 2127 Truscott, T. G., 2293 Tsurusaki, T., 879 Tuck, J. J., 309 Turner, P. S., 2609 Tusk, M., 1757 Tvarbikova, Z., 2639 Tyrrell, H. J. V., 1279 Ueki, Y.. 341 Ueno, A., 803 Unno, H., 1059 Valencia, E., 2127 van de Ven, T. G. M., 2677 van Ommen, J. G., 2479 van Truong, N., 3275, 3285 Vargas, I., 1947 2767, 2803, 21 1, 1999 Vedrine, J. C., 1017 Veith, J., 2313 Velasco, J. R., 3429 Vesala, A., 2439 Vickerman, J. C., 1903 Vincent, B., 2599 Vinek, H., 1239 Vink, H., 507, 1297 Waghorne, W. E., 1267 Wagley, D. P., 47 Walker, R. W., 435, 3187, 3195, Wallington, T. J., 2737 Wang, G-W., 1039 Watkins, P. E., 2323 Watkiss, P. J., 1279 Watt, R. A. C., 489 Webb, G., 1689 Webster, B. C., 255, 267 Weiner, E. R., 1491 Wells, C. F., 2155. 2445 Wells, J. D., 1233 Whang, B. C. Y., 291 1 Whittle, E., 2323 Wichterlova, B., 2639 Wiesner, S., 3021 Wilhelmy, D. M., 563 Williams E. H., 3147 Williams, P. A., 403 Williams, R. J. P., 2255 Wokaun, A., 1305 Wolff, T., 2969 Wood, S. W., 3419 Woolf, L. A., 549, 1287 Wright, C. J., 1217 Wu, D. C., 1795 Wiirflinger, A., 3221 Wyn-Jones, E., 1915 Yamabe, M., 1059 Yamamoto, S., 941 Yamashita, H., 1435 Yamauchi, H., 2033 Yamazaki, A., 3245 Yariv, S., 1705 Yasumori, I., 841 Yeates, S. G., 1787 Yide, X., 969, 3103 Ylikoski, J., 2439 Yokokawa, T., 473 Yoneda, N., 879 Yonezawa, T., 1435 Yoshida, S., 119, 1435 Zambonin, P. G., 1029 Zanderighi, L., 1605 Zecchina. A., 2209, 2723, 1875, Zipelli, C., 1777 Zundel, G., 553 348 1, 2827 1891

ISSN:0300-9599    

DOI:10.1039/F198480BX023

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

JOURNAL OF THE CHEMICAL SOCIETY FARADAY TRANSACTIONS, PARTS I AND I 1 The Journal of the Chemical Society is published in six sections, of which five are termed Transactions; these are distinguished by their subject matter, as follows: Dalton Transactions (Inorganic Chemistry). All aspects of the chemistry of inorganic and organometallic compounds ; including bioinorganic chemistry and solid-state inorganic chemistry ; of their structures, properties, and reactions, including kinetics and mechanisms; new or improved experimental techniques and syntheses. Faraday Transactions I (Physical Chemistry). Radiation chemistry, gas-phase kinetics, electrochemistry (other than preparative), surface and interfacial chemistry, heterogeneous catalysis, physical properties of polymers and their solutions, and kinetics of polymerization, etc.Faraday Transactions II (Chemical Physics). Theoretical chemistry, especially valence and quantum theory, statistical mechanics, intermolecular forces, relaxation phenomena, spectroscopic studies (including i.r., e.s.r., n.m.r., and kinetic spec- troscopy, etc.) leading to assignments of quantum states, and fundamental theory. Studies of impurities in solid systems. Perkin Transactions I (Organic Chemistry). All aspects of synthetic and natural product organic, organometallic and bio-organic chemistry, including aliphatic, alicyclic, and aromatic systems (carbocyclic and heterocyclic). Perkin Transactions II (Physical Organic Chemistry). Kinetic and mechanistic studies of organic, organometallic and bio-organic reactions.The description and application of physicochemical, spectroscopic, and theoretical procedures to organic chemistry, including structure-activity relationships. Physical aspects of bio-organic chemistry and of organic compounds, including polymers and biopolymers. Authors are requested to indicate, at the time they submit a typescript, the journal for which it is intended. Should this seem unsuitable, the Editor will inform the author. The sixth section of the Journal of the Chemical Society is Chemical Communications, which is intended as a forum for preliminary accounts of original and significant work, in any area of chemistry that is likely to prove of wide general appeal or exceptional specialist interest. Such preliminary reports should be followed up eventually by full papers in other journals (e.g.the five Transactions) providing detailed accounts of the work. NOTES I t has always been the policy of the Faraday Transactions that brevity should not be a factor influencing acceptability for publication. In addition however to full papers both sections carry at the end of each issue a section headed ‘Notes’, which are short self-contained accounts of experimental observations, results, or theory that will not require enlargement into ‘full’ papers. The Notes section is not used for preliminary communications. The layout of a Note is the same as that of a paper. Short summaries are required. The procedure for submission, administration, refereeing, editing and publication of Notes is the same as for full papers.However, Notes are published more quickly than papers since their brevity facilitates processing at all stages. The Editors endeavour to meet authors’ wishes as to whether anarticle is a full paper or a Note, but since there is no sharp dividing line between the one and the other, either in terms of length or character of content, the right is retained to transfer overlong Notes to the full papers section. As a guide a Note should not exceed I500 words or word-equivalents.NOMENCLATURE AND SYMBOLISM Units and Symbols. The Symbols Committee of The Royal Society, of which The Royal Society of Chemistry is a participating member, has produced a set of recommendations in a pamphlet ‘Quantities, Units, and Symbols’ (1975) (copies of this pamphlet and further details can be obtained from the Manager, Journals, The Royal Society of Chemistry, Burlington House, London W 1 V OBN).These recommendations are applied by The Royal Society of Chemistry in all its publications. Their basis is the ‘ Systeme International d’Unites’ (SI). A more detailed treatment of units and symbols with specific application to chemistry is given in the IUPAC Manual of Symbols and Terminology for Physicochemical Quantities and Units (Pergamon, Oxford, 1979). Nomenclature. For many years the Society has actively encouraged the use of standard IUPAC nomenclature and symbolism in its publications as an aid to the accurate and unambiguous communication of chemical information between authors and readers. In order to encourage authors to use IUPAC nomenclature rules when drafting papers, attention is drawn to the following publications in which both the rules themselves and guidance on their use are given: Nomenclature of Organic Chemistry, Sections A , B, C, D, E, F, and H (Pergamon, Oxford, 1979 edn). Nomenclature of Inorganic Chemistry (Butterworths, London, 197 1, now published by Pergamon).Biochemical Nomenclature and Related Documents (The Biochemical Society, London, 1978). A complete listing of all IUPAC nomenclature publications appears in the January issues of J. Chem. SOC., Faraday Transactions. It is recommended that where there are no IUPAC rules for the naming of particular compounds or authors find difficulty in applying the existing rules, they should seek the advice of the Society’s editorial staff.(ii)THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO 78 Radicals in Condensed Phases University of Leicester, 4-6 September 1984 Organising Committee Professor M C R Symons (Chairman) Dr K A McLauchlan Dr G B Buxton Professor Lord Tedder Dr T A Claxton Dr R L Willson I THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY DEUTSCH E BUNSENG ESELLSCHAFT FUR PHYSl KALISCH E CH EM1 E ASSOCIAZIONE ITALIANA D I CHIMICA FlSlCA Joint Discussion Meeting on: ~ SOCIETE DE CHlMlE PHYSIQUE Laser Studies in Reaction Kinetics Evangelische Akademie, Tutzing, West Germany, 24-27 September 1984 The discussion will be primarily concerned with the structure and reactions of radicals in liquids and solids It is designed to bring together theoretical work on structure, environmental effects and reactivity with spectroscopic and mechanistic studies directly concerned with radicals Fundamental aspects will be stressed, and particular attention will be given to new developments including measurement at short time intervals, special solvent effects, and the effects of external fields A special area for inclusion will be electron gain and loss processes including trapped and solvated electrons, electrochemical reactions, and specific electron capture and electron loss in low-temperature systems Photochemical charge-transfer processes will also be included The final programme and application form may be obtained from Mrs Y.A. Fish, The Royal Society of Chemistry, Burlington House, London W1V OBN I Organising Committee R.Ben Aim (Gif sur Yvette) G. Giacometti (Padova) P. Rigny (Gif sur Yvette) E. W. Schlag (Munchen) I. W. M. Smith (Cambridge) J. Troe (Gottingen) K. Welge (Bielefeld) The aim of this meeting is the discussion of the latest experiments and related theories in the field of laser studies of elementary chemical reactions in molecular beams, in the gas phase, and in the condensed phase. The discussion will include oral contributions and poster presentations. Further information may be obtained from: Professor Dr J. Troe, lnstitut fur Physikalische Chemie, Universitat Gottingen, Tammannstrasse 6, D3400 Gottingen, West Germany. Authors of accepted contributions will be required to provide a manuscript for publication in a special issue of the Berichte der Bunsengesellschaft fur Physikalische Chemie.The Faraday Division has a small fund to assist members with the expenses of attending this conference. Applications for a grant should be submitted to Mrs Y. A. Fish, The Royal Society of Chemistry, Burlington House, London WIV OBN, by 31 July 1984. (iii)1 THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY 1 SYMPOSIUM NO. 19 Molecular Electronic Structure Ca Icu lat ions- Met hods and ~ Applications University of Cambridge, 12-1 3 December 1984 N.B. Please note change of date Molecular electronic structure calculations have now developed into a powerful predictive tool and are necessary in several different fields to aid the understanding and interpretation of experimental observations. The meeting will review the current state of this rapidly developing discipline and will bring together experts on some of the most advanced methods and their applications.The meeting will provide an opportunity for discussion and comparison of the various techniques currently in use. It will therefore not only be a valuable forum for discussion among research workers in the field, but should also show the non-specialist what theoretical calculations can be expected to achieve now and in the near future. The preliminary programme may be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry, Burlington House, London W1V OBN THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO. 79 (in conjunction with the Polymer Physics Group) Polymer Liquid Crystals University of Cambridge, 1-3 April 1985 The object of the meeting will be to discuss all aspects of the developing subject of polymeric liquid crystals.The hope is to bring together scientists from the fields of conventional polymer science and monomeric liquid crystals who are active in this field. The discussion is aimed at understanding the following facets: (a) The chemical characteristics that give rise to polymer liquid crystalline behaviour. (b) The nature of the high local anisotropy of these systems and their structural organisation at the molecular, micron and macroscopic levels. (c) The physical properties and their industrial exploitation, with particular reference to the influence of external force fields such as flow, electric and magnetic fields.(d) The inter-relations of polymer liquid crystals with small-molecule mesophases, conventional flexible polymers and biopolymers which exhibit liquid-crystalline behaviour. Further information may be obtained from: Professor 6. R. Jennings, Electro-optics Group, Department of Physics, Brunel University, Uxbridge UB8 3PH.THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO. 80 Physical Interactions and Energy Exchange at the Gas-Solid Interface McMaster University, Hamilton, Ontario, Canada, 23-25 July 1985 Organising Committee : Professor J. A. Morrison (Chairman) Dr M. L. Klein Professor G. Scoles Professor W. A. Steele Professor F. S. Stone Dr R. K. Thomas The discussion will be concerned with certain aspects of current research on the gas-solid interface: elastic, inelastic and dissipative scattering of atoms and molecules from crystal surfaces, and the structure and dynamics of physisorbed species, including overlayers.Emphasis will be placed on the themes of physical interactions and energy exchange rather than on molecular-beam technology or the phenomenology of phase transitions on overlayers. The interplay between theory and experiment will be stressed as they relate to the nature of atom and molecule surface interaction potentials, including many- body effects. Further information may be obtained from: Professor J. A. Morrison, Institute for Materials Research, McMaster University, Hamilton, Ontario, Canada L8S 4M1 THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM NO.20 Phase Transitions in Adsorbed Layers University of Oxford, 17-1 8 December 1985 Organising Committee: Professor J. S. Rowlinson (Chairman) Dr E. Dickinson Dr R. Evans Mrs Y. A. Fish Dr N. Parsonage Dr D. A. Young The aim of the meeting is to discuss phase transitions at gas/liquid, liquidliquid and solid/fluid interfaces, and in other systems of constrained geometry or dimensionality less than three. Emphasis will be placed on molecularly simple systems, whereby liquid crystal interfaces and chemisorption phenomena are excluded. Contributions for consideration by the Organising Committee are invited and abstracts of about 300 words should be sent by 1 2 October 1984 to: Professor J. S. Rowlinson, Physical Chemistry Laboratory, University of Oxford, South Parks Road, Oxford OX1 3QZ.Full papers for publication in the symposium volume will be required by August 1985.FARADAY DIVISION INFORMAL AND GROUP MEETINGS Industrial Physical Chemistry Group The Metal-Polymer Interface To be held at Girton College, Cambridge on 10-1 2 July 1984 Further information from Dr T. G. 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Richards, Department of Pure and Applied Chemistry, University of Strathclyde, Glasgow G1 1 XL (vii)The Royal Society of Chemistry Faraday Discussions take place twice a year and are designed to cover the broad z t s of physicochemical topics, thereby encouraging scientists of different disciplines to contribute their varied viewpoints to a common theme. A recent Discussion is :- No.75 lntramolecular Kinetics No.75 in the series, this publication is the result of a general discussion held at the University of Warwick in April 1983. Contents: The Spiers Memorial Lecture; Vibrational Redistribution within Excited Electronic States of Polyatomic Molccules Intramolecular Relaxation of Excited States lsomerization of Internal-energy-selected Ions Kinetics of Ion-Molecule Collision Complexes in the Gas Phase, Experiment and Theory Intramolecular Decay of Some Open-shell Polyatomic Cations On the Theory of Intramolecular Energy Transf-er Pulsed Laser Preparation and Quantum Superposition State Evolution in Regular and Irregular Systems A Quantum-mechanical Internal-cabsion Model for State-selected Unimolecular Decomposition The Correspondence Principle and Intramolecular Dynamics Intramolecular Dephasing.Picosecond Evolution of Wavepacket States in a Molecule with Intermediate-case level Structure Energy Conversion in van der Waals Complexes of s-Tetrazine and Argon Time-dependent Processes in Polyatomic Molecules During and After Intense Infrared Irradiation Energy Distributions in the CN(X?E+) Fragment from the Infrared Multiplephoton Dissociation of CFJ CN. A Comparison between Experimental Results and the Predictions of Statistical Theories Of CbFb + Product Energy Partitioning in the Decom- position of State-selectively Excited HOON and HOOD Low-power Infrared Laser Photolysis of Tetramethyldioxetan Unimolecular Reactions Induced by Vibrational Overtone Excitation Unmolecular Decomposition of t-Butylhydro- peroxide by Direct Excitation of the 6-0 0-H Stretching Overtone Picosecond-jet Spectroscopy and Photochemistry. Energy Redistribution and its Impact’on Coherence, Isomerization, Dissociation and Solvation Energy Redistribution in Large Molecules.Dlrecl Study of Intrainolecular Relaxation in the Gas Phase with Picosecond Gating Rotation-dependent Intramolecular Processes of SO? (A’A?) in a Supersonic Jet Role of Rotation-Vibration Interaction in Vibrational Rela xa tion. Energy Redistribution in Excited Singlet Formaldehyde Sub-Doppler, Spectroscopy of Benzene in the “Channel-three” Region Intramolecular Electronic Relaxation and Photoisornerization Processes in the Isolated Azabenzene Molecules Pyridine, Pyrazine and Pyrmidine Softcover 434pp 0 85186 658 1 Price f25.00 ($48.00) Rest of the World f26.00 RSC Members f16.25 Faraday Discussions of the Chemical Society Il’o 75 lnrromoleculor Kinctru faraday Symposia are usually held annually and are confined to more specialised topics than Discussions, with particular reference to recent rapidly developing Lnes of research.A recent Symposium is :- No.17 The Hydrophobic 1nteracKon No. 17 in the series, this publication is the result of a symposium on The Hydrophobic Interaction held at the University of Reading in December 1982. Contents: Hydrophobic Interactions-a Historical Perspective Hydrophobic Hydration Geometric Relaxation in Water. Its Role in Precise Vapour-pressure Measurements of the Solubikation of Benzene by Aqueous Sodium Octylsulphate Solutions Nuclear Magnetic Resonance Relaxation Investigation of Tetrahydrofuran and Methyl Iodide Clathrates Infrared and Nuclear Magnetic Resonance Studies Pertaining to the Cage Model for Solutions of Acetone in Water Isothermal Transport Properties in Solutions of Symmetrical Tetra-akylammonium Bromides Thermodynamics of Cavity Formation in Water. A Molecular Dynamics Study Molecular Librations and Solvent Orient- ational Correlations in Hydrophobic Phenomena Monte Carlo Computer Simulation Study of the Hydrophobic Effect. Potential of Mean Force for C(CH4)jaq at 25 and SOo C Hydrophobic Moments and Protein Structure Application of the Kirkwood-Buff Theory to the Problem of Hydrophobic Interactions Disentanglement of Hydrophobic and Electrostatic Contributions to the Film Pressures of Ionic Surfactants Hydrophobic Interactions in Dilute Solutions of Poly(viny1 alcohol) Conformational and Functional Properties of Haemoglobin in Water+Alcohol Mixtures. Dependence of Bulk Electrostatic and Hydrophobic Interactions upon pH and KCI concentrations Softcover 24 Price f 36.50%0.00) Rest of the World f38.50 RSC Members €23.75 p 0 85186 668 9 ORDEAtNG RSC Members should send their orders to: The Royal Society of Chemistry. The Membership Officer, 30 Russell Square, NorrRSC Members The Royal Society of Chemistry, Distribution Centre, Blackhorse Road, L London WClB 5DT. Letchworth. Her- SG6 IHM, England. Faraday Symposia of the Chemical Society No 17 The Hi drophobrc Interoction 1982 (viii)

ISSN:0300-9599    

DOI:10.1039/F198480FP045

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984, 80, 1329-1340 Tin Oxide Surfaces Part 11 .-Infrared Study of the Chemisorption of Ketones on Tin(1v) Oxide BY PHILLIP G. HARRISON* AND BARRY M. MAUNDERS Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD Received 17th March, 1983 Infrared spectroscopy has been employed to study the chemisorption of a number of unsymmetrical ketones, RCOMe [R = C3H7, C,H,, (CH,),CH, (CH,),C, Ph], onto tin@) oxide. In every case, the final product was the surface carboxylate, RCO;(ads). The data suggested a mechanism involving initial coordination of the ketone to a surface tin site, followed by nucleophilic attack of a neighbouring hydroxyl group at the carbonyl carbon atom. Studies of the adsorption of ketones on metal oxides have largely been confined to acetone on alumina, rutile, magnesia, beryllia and nickel and calcium oxides.Three modes of reaction have been distinguished. Surface acetate is the most common chemisorption product,l-10 although coordinated acetone,ll* l2 mesityl oxide13-15 and surface enolatesll? l2 have also been observed. Both coordinated acetone and surface enolate have been proposed as intermediates in the formation of both mesityl oxide13 and surface a ~ e t a t e . ~ ~ lo No evidence has been found for the dissociative chemisorption of acetone on silica12 or germania,16 although hydrogen bonding to surface hydroxyl groups occurs. Other ketones which have been studied are halogenated acetones8> 1 7 9 or symmetrical ketones such as di-isopropy1ketone.ls Surface-bound carbonyl com- pounds have often been proposed as intermediates in the oxidation of hydrocarbons over tin(rv) oxide-containing catalysts. We have previouslylg briefly examined the chemisorption of acetone on tin@) oxide, where a surface acetate was also observed.In this paper, we report the chemisorption behaviour of a number of other unsymmetrical ketones towards tin(rv) oxide in order to elucidate the mechanism of oxidation of such species when bound to a tin(rv) oxide surface. EXPERIMENTAL The vacuum line and general procedure have been described previously.20 All spectra were recorded using Perkin-Elmer model 157G, 521 or 577 spectrophotometers. Tin(1v) oxide gel was prepared by precipitation from redistilled tin(rv) chloride.20 All ketone adsorbates were redistilled under nitrogen and then subjected to freeze-thaw cycles prior to use in order to remove dissolved gases.RESULTS PENTAN-2-ONE ADSORPTION Fig. 1 shows the spectra obtained from the action of pentan-2-one on an evacuated N m-2, 320 K). Exposure to pentan-2-one vapour N m-2) resulted in tin(1v) oxide disc (1.33 x (1.9 k N m-2, 320 K, 10 min) followed by evacuation (320 K, 13291330 I.R. STUDY OF KETONES ON TIN OXIDE 46- 95. 80 1800 1600 1400 1200 wavenumberlcm -' Fig. 1. Infrared spectra of pentan-2-one chemisorbed on tin(1v) oxide: (1) background of evacuated disc; (2) pentan-2-one vapour 320 K, 1.9 kN m-2; (3) exposed to pentan-2-one 320 K, 10 min, 1.9 kN m-2, then evacuated 320 K, 0.5 h, < 1.33 x N md2; subsequent evacuation (4) 320 K, 20 h, < 1.33 x N m-2; (6) N m-2; ( 5 ) 378 K, 2 h, < 1.33 x N m-2.513 K, 20 h, < 1.33 x absorption bands at 1510, 1408 and 1310 cm-l, a broad shoulder between 1770 and 1700 cm-l and a second shoulder at 1585 cm-l. Further evacuation (320 K, 20 h) removed the 1770-1700 cm-l shoulder and greatly reduced the intensity of the 1585 cm-l shoulder. Evacuation at higher temperatures (5 13 K, 20 h) greatly increased the intensity of the 1510 and 1408 cm-I bands and reduced the 1200-1000 cm-l band, because of the OH deformation, enabling two weaker bands at 1265 and 1105 cm-l to be discerned. In addition a weak shoulder at ca. 1600 cm-l was present on the high- wavenumber side of the 1510 cm-' band. The two shoulders at 1700-1770 and 1585 cm-l can be assigned to pentan-2-one hydrogen-bonded to surface hydroxyl groups and coordinated to Lewis-acidic tin sites, respectively.Shifts in carbonyl stretching frequencies of the order of 6 to 33 cm-l forP. G. HARRISON AND B. M. MAUNDERS 1331 68. 83 n 83 0 s 4-3 c ._ 89 E 2 4-3 75 1800 1600 1400 1200 wavenum ber/cm-' Fig. 2. Infrared spectra of butyric acid chemisorbed on tin(1v) oxide: (1) background of evacuated disc, (2) exposed to butyric acid vapour, 320 K, 0.5 h, < 0.1 kN m-2; subsequent evacuation (3) 320 K, 1.5 h, < 1.33 x N m-2; ( 5 ) N m-2, (4) 378 K, 20 h, < 1.33 x N mP2. 453 K, 3.5 h, < 1.33 x ketones interacting with hydroxyl groups of molecules have been ascribed by Denisov21 to the formation of a hydrogen bond of the type OH. - .O=C.Con- siderably greater shifts are observed when ketones coordinate to strong aprotic electron acceptors; for example, Filimonov22 reports that the v(C=O) stretching band of acetone shifts to 1625-1635 cm-l upon coordination with AlBr,. The 1510 and 1408 cm-l bands can be assigned to the v,,(COO) and v,(COO) stretching modes, respectively, of a surface butyrate formed by oxidative chemisorption of pentan-2-one, while the 1310 cm-l band can be attributed to a C-H deformation mode of the butyrate and/or coordinated hydrogen-bonded pentan-2-one. The former is most likely since an absorption band occurs in this position on a tin(rv) oxide disc exposed to butyric acid vapour and its intensity is unaltered after prolonged evacuation.1332 .R. STUDY OF KETONES ON TIN OXIDE 1 1 1 1 1 1 1 1 1 1 1 1800 1600 1400 1200 wavenumber/cm -I Fig.3. Infrared spectra of butan-2-one chemisorbed on tin(1v) oxide : (1) background of evacuated disc; (2) exposed to butan-2-one vapour, 320 K, 0.5 h, 4.3 kN m-2, then evacuated, 320 K, 1 h, c 1.33 x N m-2; subsequent evacuation (3) 393 K, 20 h, < 1.33 x N m-2. The weak band at 1265 cm-I and shoulder at ca. 1600 cm-l can tentatively be assigned to the (COH) deformation and (C=O) stretching modes, respectively, of free acid, formed by protonation of the carboxylate, coordinated through the carbonyl oxygen to surface Lewis acidic tin sites. The 1105 cm-l band can possibly be assigned to a (C-C) stretching mode. The assignment of the chemisorption product to a surface butyrate was confirmed by the adsorption of butyric acid vapour (< 0.1 kN m-2, 320 K, 0.5 h) on a similarly treated tin(1v) oxide disc (fig.2), where, after evacuation, absorption bands were observed at 1610, 1512, 1410 and 1310 cm-l. The 1610 cm-' band was reduced to a shoulder after evacuation at 453 K, at the same time two bands at 1265 and 1105 cm-l were observed. The 1512, 1410, 13 10 and 1105 cm-I bands are due to the butyrate, while the 1610 and 1265 cm-l bands are due to the v(C=O) stretching and (COH) deformation modes, respectively, of coordinated butyric acid. The increases in band intensity observed on evacuation at 5 13 K after pentan-2-one adsorption can be attributed either to oxidation of any remaining coordinated pentan-2-one or to pentan-2-one, adsorbed on the silica walls at lower temperatures, being desorbed then oxidised on the tin@) oxide surface.BUTAN-2-ONE ADSORPTION Exposure of an evacuated tin(1v) oxide disc (1.33 x N m+, 320 K) to butan-2-one vapour (320 K, 0.5 h, 4.3 k N m-2), then subsequently re-evacuated, exhibited strong infrared absorption bands at 1512 and 1410 cm-l, together withP. G. HARRISON AND B. M. MAUNDERS 1333 76 - 95 - 87- * 94- E a) E: +-I Y .- A 2 * 90. 8 3 . - l o o / o - 1 1 1 1 1 1 1 1 1 1 1 1800 1600 1400 1200 w avenum berlcm-' Fig. 4. Infrared spectra of 3-methylbutanone chemisorbed on tin(rv) oxide : (1) background of evacuated disc; (2) exposed to 3-methylbutanone, 320 K, 10 min, 1.5 kN m-2, (3) 3- methylbutanone vapour, 320 K, 1.5 kN mP2; subsequent evacuation (4) 320 K, 1.5 h, < 1.33 x Nm-2; (5) 373 K, 2 h, < 1.33 x lop4 Nm-2; (6) 453 K, 20 h, < 1.33 x N mp2.weaker bands at 1378 and 1300 cm-l and a broad shoulder centred at ca. 1610 cm-1 (fig. 3). Evacuation (393 K, 20 h) increased the intensity of the 1512 and 1410 cm-1 bands, the 1512 cm-l band being unsymmetrical and broader on the high-wavenumber side. The 1512 and 1410 cm-l bands can be assigned to the (COO) stretching modes, of a surface propionate, with the 1378 and 1300 cm-l bands being attributed to the symmetric methyl deformation and a C-H deformation mode, respectively. The shoulder at 1610 cm-l is due to the v(C-0) stretching mode of coordinated butan-2-one. The broad, unsymmetrical nature of the 15 12 cm-l band after evacuation at 393 K suggests that some coordinated ketone is still present.1334 I.R.STUDY OF KETONES ON TIN OXIDE 3-METHYLBUTANONE ADSORPTION Fig. 4 shows the spectra obtained when an evacuated tin(1v) oxide disc (1.33 x lo-* N m+, 320 K) was exposed to 3-methylbutanone vapour (320 K, 10 min, 1.5 kN m-2). Absorption bands were observed at 1680, 1610, 1530, 1460, 1412 and 1362 cm-l, along with a shoulder at ca. 17 10 cm-l. Subsequent evacuation removed the 1710 cm-l shoulder, reduced the 1680 cm-l band to a shoulder on the broad, strong 161 5 cm-l band, shifted from 16 10 cm-l. The 1530 cm-l band was shifted to 1515 cm-l and, like the 1412 cm-l band, became more distinct. The 1460 cm-l band, shifted to 1470 cm-l, and the 1362 cm-l band, shifted to 1370 cm-l, were both reduced in intensity. In addition weak bands were now apparent at 1290 and 1252 cm-l.Further evacuation at 373 and then 453 K reduced further the intensities of the 1680 cm-l shoulder and the 1610, 1470 and 1370 cm-l bands, while the 151 5 , 1412 and 1290 cm-l bands all increased in intensity, with the 1290 and 1252 cm-l bands shifting to 1300 and 1265 cm-l, respectively. Absorption bands also became apparent at 11 80 and 1100 cm-l. The 1710 cm-l shoulder can be assigned to the v(C=O) mode of the vapour-phase ketone, or possibly the same species weakly hydrogen-bonded to the surface. The 1680 cm-l band is in the position expected for the v(C=O) stretching mode of the ketone coordinated, through the carbonyl oxygen, to Lewis-acidic tin sites. The 1610 cm-l band could be ascribed to ketone coordinated to a second, stronger, Lewis-acid site, or possibly to a carbon-carbon double-bond stretch of an enolate structure, a third, and perhaps the most probable, assignment could be due to coordinated isobutyric acid, formed by oxidation of the ketone to the isobutyrate followed by protonation to give the free acid, since a band occurs in this position upon isobutyric acid absorption.The bands at 1515 (shifted from 1530 cm-l upon evacuation) and 1412 cm-1 are due to the v(C00) stretching modes, respectively, of a surface isobutyrate, while the 1460 and 1362 cm-l bands can be assigned to the S,,(CH,) and S,(CH,) deformation modes, respectively, of the ketone and possibly isobutyrate. The decrease in intensity, upon evacuation, of the 1680, 1460 and 1362 cm-l bands is evidence for the removal of the majority of the coordinated ketone.The shift of the latter two bands to slightly higher wavenumbers, 1470 and 1370 cm-l, reflects the smaller surface coverage by the ketone and hence less steric restraints upon the deformation modes. The increase in intensity, with evacuation, of the 1515, 1412 and 1290 cm-l bands and the appearance of the 1180 and 1100 cm-l bands confirm that these are all due to the same species. The assignment of the species to the iso- butyrate was confirmed by adsorption of isobutyric acid vapour. The observation of a band at 1375 cm-l after adsorption indicates that the 1370 cm-l band observed on 3-methylbutanone adsorption must be partially due to the isobutyrate. The weak band at 1625 cm-l could be due to a (COH) deformation mode associated with the free acid, for which the 1610 cm-l band is assigned to the v(C=O) stretching mode.The spectra from both 3-methylbutanone and isobutyric acid both show the presence of a shoulder at ca. 1610 cm-l on the 15 15 cm-l band together with the weak 1265 cm-l band. The 1180 and 1100 cm-l bands can tentatively be assigned to carbon-carbon skeletal vibrations of the isopropyl unit [(CH3)2C],23 while the 1290-1 300 cm-l band is probably due to a (C-H) deformation mode. 2,2-DIMETHYLBUTAN-3-ONE ADSORPTION Fig. 5 shows the spectra obtained when an evacuated tin(1v) oxide disc (1.33 x N mP2, 320 K) was exposed to 2,2-dimethylbutan-3-one vapour (320 K, 10 min, 2.0 kN m-2). A broad, strong absorption band was observed with itsP.G. HARRISON AND B. M. MAUNDERS 1335 1800 1600 1400 wavenumberlcm-' Fig. 5. Infrared spectra of 2,2-dimethylbutan-3-one chemisorbed on tin(1v) oxide; (1) background of evacuated disc, (2) 2,2-dimethylbutan-3-one vapour, 320 K, 2.0 kN m-2; (3) exposed to 2,2-dimethylbutan-3-one, 320 K, 10 min, 2.0 kN m-2; subsequent evacuation (4) 320 K, 1 h, < 1.33 x N m-2; ( 5 ) 388 K, 48 h, < 1.33 x lop4 N m-2; (6) 493 K, 2 h, N mp2; (7) 623 K, 2 h, < 1.33 x < 1.33 x N m-2. maximum at 1680 crn-l, together with absorption bands at 1470,1360 and 1 140 cm-l. Subsequent evacuation (320K) left bands at 1630 (broad and strong), 1480 and 1370 cm-l, along with a slight shoulder at 1530 cm-l and a weak band between 1430 and 1385 cm-l. Prolonged evacuation (388 K, 60 h) produced a change in the spectrum with two broad, overlapping bands appearing at 1555 and 1520 cm-l, along with other bands at 1480, 1410, 1370, 1220 and 1170 cm-l.A broad band, with its maximum at 1515 cm-l, remained after evacuation at 493 K for 10 min, along with a shoulder on the low-wavenumber side at 1480 cm-l. The absorption bands at 1410 (now shifted to 1415), 1370, 1220 and 1170 cm-l still remained. From a comparison with the infrared spectrum of the vapour-phase ketone, which exhibited strong absorption bands at 1728, 1485, 1370, 1360 and 1140 cm-l, the absorption bands at 1630, 1480 and 1370 cm-l, remaining after evacuation at 320 K, can be assigned to the v(C=O) stretching, S,,(CH,) and S,(CH,) deformation modes,1336 I.R. STUDY OF KETONES ON TIN OXIDE 60 78 I I I I I I I I I I I 1800 1600 1400 1200 wavenumber/cm -' Fig.6. Infrared spectra of acetophenone chemisorbed on tin@) oxide: (1) background of evacuated disc; (2) exposed to acetophenone vapour, 320 K, 10 min, c 0.1 kN m-2; subsequent evacuation (3) 320 K, 20 h, < 1.33 x N m-2; (4) 393 K, 20 h, c 1.33 x N m-2. respectively, of coordinated 2,2-dimethylbutan-3-one, where the ketone is coordinated via the carbonyl oxygen to Lewis sites on the oxide surface. The very broad band at 1680 cm-l observed in the presence of the vapour is due to overlapping of the free v(C=O) stretching, hydrogen-bonded and coordinated v(C=O) stretching vibrations of the ketone. At 320 K there is very little evidence for any other surface species, only the very weak bands at 1530 and 1430-1385 cm-l.However, at higher evacuation temperatures absorption bands of a new surface species appear as those due to the coordinated ketone disappear. The new bands at 1515-1520 and 1415 cm-1 can be assigned to the v,,(COO) and v,(COO) stretching modes, respectively, of a surface 2,2-dimethylpropionate, while the 1480 (shoulder) and 1370 cm-l bands can be assigned to 6,,(CH,) and 6,(CH,) deformation modes, respectively. Although adsorp- tion of the corresponding acid was not carried out, the two bands at 1220 and 1170 cm-l lend weight to the proposed structure, since they occur in the expected position for the skeletal vibrations of the 2,2-dimethyl group.23P. G. HARRISON AND B. M. MAUNDERS 1337 ACETOPHENONE ADSORPTION Fig. 6 shows the spectra obtained when an evacuated tin(1v) oxide disc (1.33 x lop4 N m-2, 320 K) was exposed to acetophenone vapour (320 K, 10 min, < 0.1 kN mP2).Strong absorption bands were observed at 1655, 1640, 1590, 1565 (shoulder), 1265 and 12 15 cm-l, with weaker bands at 1490,1445, 1390 (broad), 1360, 1180 and 1160 cm-l. The 1640 cm-l bands can be attributed to the v(C=O) stretching mode of the ketone coordinated or hydrogen bonded to the surface [cf. v(C=O) of the liquid of 1690 cm-’1. The 1590 and 1445 cm-l bands are due to the aromatic-ring vibrations, while the 1360 cm-l band can be ascribed to the d(CH,) deformation. After evacuation at 320 K strong absorption bands remained at 1590, 1490 and 1390 cm-l with weaker bands at 1560, 1445, 1220, 1180 and 1160 cm-l, and a shoulder on the high-wavenumber side of the 1590 cm-l band. The large decrease in intensity of the 1640 cm-l band indicates the removal of the majority of the coordinated ketone.Other aromatic-ring stretches are still quite strong, indicating that the aromatic portion of the molecule is still present, but the methyl deformation band at 1360 cm-l has been removed. Two previously weak bands, at 1490 and 1390 cm-l, have increased in intensity, thus it can be concluded that the acetophenone has been chemisorbed on the surface with the loss of the methyl group. The two bands at 1490 and 1390 cm-l can be ascribed to the va,(COO) and v,(COO) stretching modes, respectively, of a surface benzoate group. Further evacuation for 20 h at 393 K completely removed the shoulder due to coordinated ketone, leaving strong, sharp, absorption bands at 1590 and 1445 cm-l due to aromatic-ring stretching vibrations, at 1495 and 1395 cm-l due to the two (COO) stretching modes and at 1180 and 1160 cm-l due to the (C-H) out-of-plane deformations of the aromatic ring.Absorption bands for the carboxylates observed from adsorption of all ketones are summarised in table 1. DISCUSSION The tin(rv) oxide surface can be considered as consisting of Lewis-acidic and -basic sites, the Lewis-acidic sites being the exposed Sn ions and the basic sites the surface hydroxyl groups or oxide ions. Weak Brsnsted-acid sites can also be detected from ammonia adsorption and from the observation of absorption bands attributable to butyric and isobutyric acid after penton-2-one and 3-methylbutanone adsorption, respectively.The mechanism of surface carboxylate formation’ from adsorption of the ketones is uncertain, though it is likely that chemisorption proceeds via the nucleophilic attack of a reactive surface hydroxyl group at the carbonyl carbon of the ketone, following coordination to an unsaturated surface tin ion (scheme 1). R R’ R f?’ R R‘ ‘-0 L II JI 0 .. 0 /sn. ,Sn, 0 0 0 Scheme 9, 2”. 0 0 0 1.CI w w 00 Table 1. Vibrations of surface carboxylates derived from adsorbed ketones 2,2-dimethyl- ketone: pentan-2-one butan-2-one 3-methylbutanone butan-3-one acetophenone corresponding adsorbed 2,2-dime thyl- carboxylate : butyrate propionate isobutyrate propionate benzoate assignments ; 1600(sh) 1590 ring vibration 3 1512 1515 1515 1490 - va,(COO) % 1510 1445 ring vibration 1408 1410 1412 1415 1390 - v,(COO) 0 1378 1375 1370-1 - 375 - 3 v(c=o)a 5 - ! - - 1610 - das(CH& 1480-1485 - - - - - - 6s(CH,) 6(COH)a 2 d(C-H) skeletal vibration of 2,2-dimethyl group - 1290-1300 1250 - 1310 1300 - - - - 1170 - \ 0 1220-1 225 - 1265 - - - - 1 180 6(C-H) out-of-plane of 1160 ) aromatic ring - - ?$ - U m - - - skeletal vibration - - - - - ) of isopropyl group - 1180 1100 - - v(C-C) - - - - 1105 - - a These bands arise from acid formed by protonation of the carboxylates, see text.P.G . HARRISON AND B. M. MAUNDERS 1339 In our previous studieP of the adsorption of [2H,]acetone on tin(1v) oxide we observed the formation of a surface acetate and suggested that an enolate intermediate may be involved due to the presence of absorption bands at ca.1600 and 1240- 1250 cm-l which were assigned to the v(C=C) stretching and 6(COH) deformation modes, respectively, of the enolate species. In this investigation absorption bands were observed in similar positions with pentan-2-one and 3-methylbutanone adsorptions ; however, as similar absorption bands were observed with adsorption of the corre- sponding acids to the observed carboxylates, butyric and isobutyric acids, respect- ively, the assignment of these bands to coordinated protonated acid seems more reasonable in this case, formed by the abstraction of hydrogen from an adjacent sur- face hydroxyl group (scheme 2). R I R 0 - H ‘ r ’ Scheme 2. All the ketones employed in the present study were unsymmetrical ketones of the type RCOMe.However, in every case, the surface carboxylate formed was that derived from the group R and in no case was a surface acetate observed. Thus, pen tan- 2-one, butan- 2-one, 3 -met h ylbu tanone, 3,3 -dime t h ylbu tan-2-one and ace to- phenone gave a surface butyrate, propionate, 2-methylpropionate7 2,2-dimethyl- propionate and benzoate, respectively. The underlying reason for this observation may be thermodynamic or kinetic in origin. A comparison of the relative heats of formation of the two possible sets of products [i.e. RCO;(ads) + CH, and MeCO;(ads) + RH (assuming that the tin-oxygen bond remains the same in each case, and using data for the free acid, which should parallel those of the carboxylates)] shows that, where data are a~ailable,~, the marginally favoured thermodynamic products would be expected to be MeCO;(ads) + RH by ca.3-33 kJ mol-l. Similarly, simple bond dissociation energy estimates2, of the two bonds, R-COMe and Me-COR, shows that the bond required to be broken to give formation of surface acetate, i.e. R-COMe, is consistently the weaker of the two by ca. 2-2 1 kJ mol-l. Thus, contrary to observation, thermodynamic arguments favour the formation of surface acetate and methane. Hence, it may be concluded that the reactions are kinetically controlled, which is consistent with the nucleophilic process shown in scheme 1. Although in this study the constitution of the gas phase was not investigated, met- hane was confirmed to be a constituent of the gas phase in the adsorption of acetone.19 M.L. Hair and I. D. Chapman, J . Phys. Chem., 1965,69, 3949. H. Knozinger, Adv. Catal., 1976, 25, 184. H. Knozinger, H. Krietenbrink, H. D. Muller and W. Schulz, Proc. 6th. In?. Congr. Catalysis (The Chemical Society, London, 1976), p. 183. P. Fink, Rev. Roum. Chim., 1969, 14, 81 1 . W. Schulz and H. Knozinger, J . Phys. Chern., 1976, 80, 1502. A. V. Kiselev and A. V. Uvarov, Surf. Sci., 1967, 6, 399. N. E. Tretyakov and V. N. Filimonov, Kinet. Catal., 1970, 11, 815. H. Miyata, Y. Toda and Y. Kubokawa, J. Catal., 1974, 32, 155. ’ A. V. Deo, T. T. Chuang and J. G. Dalla Lana, J. Phys. Chem., 1971,75, 234.1340 I.R. STUDY OF KETONES ON TIN OXIDE lo H. Miyata, M. Wakamiya and Y. Kubokawa, J. Catal., 1974, 34, 117. l1 A. A. Kadushin, Yu. N. Rufov and S. Z. Roginskii, Kinet. Catal., 1967, 8, 1356. J. C. McManus, Y. Harano and M. J. D. Low, Can. J. Chem., 1969,47, 2545. l 3 D. M. Griffiths and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1978,74,403. l4 I. R. Shannon, I. J. S. Lake and C. Kemball, Trans. Faraday SOC., 1971,67, 2760. l5 R. Fujii, J. Chem. SOC. Jpn, Pure Chem. Sect., 1948, 69, 151. l6 M. J. D. Low, N. Madison and P. Ramamurthy, Surf. Sci., 1969, 13, 238. D. M. Griffiths and C. H. Rochester, J. Chem. SOC., Faraday Trans. I , 1977, 73, 1913. H. Knozinger, Forschungsber. Whrtech (Pundesminist, Verteidigung, 1976) (brivg-FBWT 7620, Luft-Raumfahrt, Teilz), p. 53. l9 E. W. Thornton and P. G. Harrison, J. Chem. SOC., Faraday Trans. 1, 1975, 71, 2468. 2o E. W. Thornton and P. G. Harrison, J. Chem. SOC., Faraday Trans. I , 1975,71, 461. 21 G. S. Denisov, Dokl. Akad. Nauk SSSR, 1960, 134, 1131. 22 V. N. Filimonov, D. S. Bystrov and A. N. Terenin, Opt. Spektrosk., 1957, 3, 480. 23 L. J. Bellamy, Infrared Spectra of Complex Molecules (Methuen, London, 1958). 24 The Handbook of Chemistry and Physics (CRC Press, Boca Raton, Florida, 1980). (PAPER 3/419)

ISSN:0300-9599    

DOI:10.1039/F19848001329

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1984,80, 1341-1356 Tin Oxide Surfaces Part 12.-A Comparison of the Nature of Tin(rv) Oxide, Tin(rv) Oxide-Silica and Tin(rv) Oxide-Palladium Oxide : Surface Hydroxyl Groups and Ammonia Adsorption BY PHILIP G. HARRISON* AND BARRY M. MAUNDERS Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD Received 28th April, 1983 The chemical nature of the surfaces of tin(1v) oxide, tin(1v) oxide-silica and tin(1v) oxide-palladium oxide have been compared by surface dehydroxylation, deuterium exchange and ammonia adsorption. Physisorbed and hydrogen-bonded molecular water is lost from all three oxides at temperatures up to ca. 350 K. Further significant loss of water does not occur until ca. 480 K, when condensation of adjacent surface hydroxyl groups occurs.At higher temperatures, a sharp band at 371 5 cm-l is observed for tin(1v) oxide-silica, which is assigned to the v(0H) stretching vibration of isolated surface Si-OH groups. This band is present even at 723 K, when essentially all other surface hydroxyl groups had desorbed. The thermal dehydroxylation behaviour of tin@) oxide and tin(1v) oxide-palladium oxide is similar, although serious transmission problems occur above 523 K for tin(1v) oxide. All three oxides undergo surface deuteration with deuterium oxide at 523 K, although exchange with tin(1v) oxide-silica is not complete until 633 K. Ammonia adsorption demonstrates that tin(1v) oxide and tin(rv) oxide-palladium oxide are predominantly Lewis acidic, although weak Bronsted acidity can be observed in the presence of water vapour.Strong Bronsted-acid sites can be produced by protonation with hydrogen chloride. Surface amide groups are formed on tin@) oxide-palladium oxide after evacuation at 500 K, probably on palladium sites. Tin(1v) oxide-silica exhibits both Lewis and Bronsted acidity, the amount of Bronsted acidity with respect to Lewis acidity decreasing with increasing temperature. Our earlier studies of the chemisorption properties of tin(rv) oxide have included a wide range of adsorbent molecules.1 It is well known that the chemical properties of oxide materials can be modified, sometimes substantially, by the inclusion of other metals or metalloids. We have previously compared the chemisorption behaviour of tin(rv) oxide modified by the inclusion of some first-row transition-metal ions;2 here we report a detailed comparitive study of the nature of the surfaces of tin(rv) oxide, tin(rv) oxide-silica and tin(rv) oxide-palladium oxide.EXPERIMENTAL Tin(1v) oxide, tin(1v) oxide-silica and tin(1v) oxide-palladium oxide were prepared by (C0)-precipitation of the appropriate (mixture of) the metal chloride solution in water by AnalaR aqueous ammonia solution. Assuming that complete conversion to metal oxide occurred, the coprecipitated gels had the stoichiometry SnO, - 0.25Si0, and SnO, * 0.02PdO. Self-supporting discs of the tin(1v) oxide or coprecipitated gels were prepared by finely grinding ca. 60 mg of the oxide with an agate mortar and pestle, drying the powder in air at ca.373 K for 0.5-1 h, then pressing it in a 2.5 cm diameter stainless-steel die with a compacting pressure of 34 MN 13411342 10 - - 8 - h $ 5 - * 6 - E - 3 4 - : - 2 - TIN OXIDE SURFACES TIK Fig, 1. Thermogravimetric results for tin(rv) oxide-silica. Plot of percentage weight loss against temperature of evacuation. Prior to use the discs were pretreated by evacuation in the infrared cell for at least 2 h at room temperature then for at least 3 h in the temperature range from room temperature to ca. 760 K at a vacuum pressure of 1.33 x N m-2. For pretreatment above 523 K, evacuation was usually followed by treatment with oxygen (1.33-3.32 kN m-2) at the evacuation temperature for at least 1 h, then re-evacuation accompanied by cooling of the disc. Adsorbate adsorption was carried out in the range of the ambient temperature of the spectrometer beam (Tab z 320 K) to 760 K.Specific surface-area measurements using the B.E.T. method by adsorption of nitrogen at 77 K gave values of 270 m2 g-' after room-temperature evacuation reducing to 215 m2 g-l after 17 h at 483 K for the tin(rv) oxide-silica sample, and 190 m2 g-' reducing to 99 m2 g-l after 16 h at 423 K for the tin(rv) oxide-palladium oxide sample. Infrared spectra were recorded using a Perkin-Elmer 577 spectrometer. RESULTS AND DISCUSSION SURFACE DEHYDRATION AND DEHYDROXYLATION Results of a thermogravimetric analysis (t.g.a.) of the tin(1v) oxide-silica sample are shown in fig. 1. The slight loss up to 323 K can be explained by removal of loosely held molecular water.Removal of hydrogen-bonded water can explain the loss at 348 K, while the large loss in weight beginning with evacuation at ca. 480 K can be attributed to the process of condensation of hydroxyl groups. Even after evacuation at 723 K loss of weight was observed, which suggests that the gel was still partially hydroxylated. The infrared spectra (fig. 2) parallel the t.g.a. observations. Prior to evacuation an intense absorption band was present at 1625 cm-l due to the v, bending mode of molecular water. In water vapour this band has been reported at 1595 crn-l,, while for monomeric, dimeric and polymeric water molecules in a nitrogen matrix it has been reported at 1600, 1620 and 1633 cm-l, re~pectively.~ On evacuation at the intensity of this band was greatly reduced and shifted to 1610 cm-l.This is in agreement with the idea of the molecular water being loosely held to the surface. The remaining band at 16 10 cm-l suggests that some water is more firmly held, presumably by hydrogen bonding to surface hydroxyl groups or oxide ions, or held at strongly Lewis-acidic centres. On evacuation, the broad intense v(0H) absorption between 3760 and 2000 cm-l due to molecular water and surface hydroxyl groups decreasedP. G. HARRISON AND B. M. MAUNDERS 1343 80 - 60 E Q) c * +-I .g 40 E c c 20 0 4000 3500 3000 2500 2 000 Fig. 2. Infrared spectra of tin(1v) oxide-silica at (1) Tab and after evacuation at (2) Tab, (3) 373 K for 4 h, (4) 483 K for 3 h, (5) 544 K for 15.5 h, (6) 625 K for 50 h and (7) 723 K for 5 h. Spectra (6) and (7) are displayed to lower wavenumber by 100 cm-l.wavenumber/cm-' 80 60 % Q) u c * .z 4 0 E 2 U 2 0 0 1 4000 3500 3000 2500 2000 wavenum ber/cm-' Fig. 3. Infrared spectra of tin(rv) oxide-palladium oxide at (1) Tab and after evacuation at (2) 373 K for 15.5 h, (3) 488 K for 23 h, (4) 530 K for 24 h, (5) 573 K for 24 h and (6) 723 K for 23 h. in intensity, and a shoulder appeared at ca. 3700 cm-l which may be assigned to the v(0H) of isolated surface Si-OH groups. Raising the evacuation temperature had little effect on the intensity of the broad band until the disc was calcined at 483 K, when a significant reduction occurred. This indicates the removal of surface hydroxyl groups by elimination of water, and agrees well with the increase in weight loss from1344 TIN OXIDE SURFACES * I 4000 3500 3000 2500 2 000 Fig.4. Infrared spectra of tin(1v) oxide-silica exposed to deuterium oxide: (1) sample evacuated at 708 K for 10 h followed by oxygen treatment; (2) after exposure to D,O vapour (1.33 kN m-,) for 14 h at 568 K and evacuation; (3) after further exposure to D,O vapour (1.33 kN m-2) for 48 h at 633 K and evacuation. wavenum ber/cm-' the thermogravimetric results observed around this temperature. Continued evacuation at increasing temperatures further reduced the intensity of the broad band. At the same time the shoulder at ca. 3700 cm-l developed into a fairly sharp band at 3715 cm-l, but with some broadening on the low-wavenumber side. This band, although reduced in intensity, was still present after evacuation at 723 K.The behaviour of tin(1v) oxide and tin(rv) oxide-palladium oxide is very similar, and both exhibit a broad intense band between 3700 and 2000cm-l [v(OH)], a medium-intensity band at ca. 1630 cm-l (v, bending mode of molecular water) and a broad medium-intensity band between ca. 1250 and 1100 cm-l [a(Sn-OH)]. Room-temperature evacuation removed the v, modes at ca. 1630 cm-l and decreased the intensity of the v(0H) band. Little further change in the intensity of this band takes place until heating in vacuo at ca. 483 K, when a significant decrease in intensity occurs because of the condensation of adjacent surface hydroxyl groups. The tin(1v) oxide-palladium oxide surface was still partially hydroxylated after evacuation at 723 K (fig. 3). Evacuation of tin(1v) oxide at 523 K or above produced discs of very poor optical transmittance which varied in colour from dark brown to black with increasing temperatures.Although not restoring the original white colour of the disc, oxygen treatment did restore it partially, the discs becoming pale yellow in colour. Subsequent prolonged evacuation at Tab again resulted in loss of transmittance accompanied by darkening of the disc. Similar discolouration upon thermal treatment was also observed for the tin(1v) oxide-silica and tin(rv) oxide-palladium oxide gels, although for these materials the problem was not serious. SURFACE DEUTERATION Fig. 4 shows the infrared spectra of a tin(rv) oxide-silica disc outgassed at 708 K and then treated with D,O vapour. Hydrogen-deuterium exchange was rapid but notP.G . HARRISON AND B. M. MAUNDERS 1345 - I 1700 1500 1300 11 00 wavenumberlcm-' Fig. 5. Infrared spectra of ammonia adsorbed on tin(rv) oxide: (1) starting surface evacuated (320 K, 24 h); (2) duringexposure to ammonia (1.33 kN m-2, 320 K); (3) subsequent desorption (320 K, 15 min). complete until treatment with the vapour at 633 K. A sharp absorption band at 2740 cm-l can be assigned to the 0-D stretching vibration of isolated SOD groups, while the broad band centred at 2525 cm-l is due to the deuteroxyl groups perturbed by hydrogen bonding. The observed shift for the isolated 0-D stretching vibration (v,/v, = 1.36) is in very good agreement with the calculated shift ( v , / v , = 1.37). Essentially, complete H-D exchange occurs with tin(rv) oxide-palladium oxide at 523 K, to give a broad band centred at ca.2400 cm-l attributable to the hydrogen- bonded deuteroxyl stretching vibrations. A slight shoulder was present at 2640 cm-l which can be ascribed to isolated SnOD groups. Again the observed shift in the isolated 0-D stretching vibration (v,/v, = 1.38) is in very good agreement with the expected shift. The hydroxyl deformation mode would be expected at ca. 870 cm-l; however, this was not observed owing to bulk oxide adsorptions. SURFACE ACIDITY: ADSORPTION OF AMMONIA Fig. 5 shows the infrared spectra of a tin@) oxide disc evacuated (< 1.33 x N m-2) at Tab for 24 h and subsequently exposed to ammonia vapour (1.33 kN m-2) at Tab. Three new absorption bands appeared in the presence of the ammonia vapour at 1620, 1470 and 1240-1250 cm-l.The 1470 cm-l band could be removed by evacuation for a few minutes at qb but was restored on exposure to water vapour ; re-evacuation again removed this band. Evacuation also decreased the intensity of the 1620 and 1240-1250 cm-l bands, with small shifts to 1615 and1346 TIN OXIDE SURFACES 71 - 74 E s 77 5 2 65 e, .- * - 10 - 1700 1500 1300 1100 wavenumberlcm-' Fig. 6. Infrared spectra of ammonia adsorbed on tin@) oxide during exposure to ammonia (1.33 kN mW2, 320 K) on discs heated under vacuum at (1) 320, (2) 369, (3) 413 and (4) 495 K. 1235 cm-l, respectively. The effect of the presence of water vapour was difficult to judge, the water-bending mode being in the same region as the 1615 cm-l band.However, subsequent evacuation left both the 1615 and 1235 cm-l bands slightly reduced in intensity. The broad, intense hydroxyl stretching band prevented the observation of any N-H stretching vibrations. Increasing the thermal pretreatment temperature of the tin(1v) oxide disc had two effects on the spectra. First, the intensity of the 1470 cm-l band decreased with increasing evacuation temperature and was completely removed after evacuation at 495 K (fig. 6). Secondly, the hydroxyl stretching band, which exhibits a maximum at 3200 cm-l after evacuation at 495 K, shifts to lower frequency by ca. 150 cm-l. Subsequent evacuation restored the band to its original position with a slightly increased intensity. Attempts to study this at higher pretreatment temperatures were unsuccessful due to a severe loss in transmittance of the disc upon admitting ammonia vapour.On evacuation of ammonia vapour from the cell the intensities of both the 161 5 and 1235 cm-l bands decreased with increasing evacuation temperature. A tin(1v) oxide disc was heated under vacuum at 367 K, treated with ammonia vapour (4.0 kN m-2, Tab), evacuated and subsequently exposed to D20 vapour (1.33 kN mP2, Tab) (fig. 7). The 1615 and 1245 cm-l bands were greatly reduced in intensity and new bands were observed at 1435, 1205, 1140, 1075,980 and 900 cm-l. Subsequent evacuation had little effect on the spectrum apart from the followingP. G. HARRISON AND B. M. MAUNDERS 1347 - I 1700 1500 1300 1100 900 wavenumber/cm -' Fig. 7. Infrared spectra of a tin(1v) oxide disc heated under vacuum at 369 K and then given the following treatments: (1) exposed to ammonia, followed by evacuation at 320 K; (2) during exposure to D,O vapour (1.33 kN m-,, 320 K); (3) subsequent evacuation (320 K, 10 min).1800 1600 1400 1200 1000 wavenumber/cm -' Fig. 8. Infrared spectra of tin(1v) oxide: (1) starting surface of disc evacuated (485 K, 17 h); (2) during exposure to hydrogen chloride (1.33 kN m-,, 320 K); (3) subsequent evacuation (320 K, 2 h); (4) during exposure to ammonia (1.33 kN m-,, 320 K); (5) subsequent evacuation.1348 TIN OXIDE SURFACES changes: the 1245 cm-l band became more distinct and increased in intensity, the 1205 cm-l band was reduced in intensity and shifted to 1 180 cm-l, the 1075 cm-l band shifted to 1035 cm-l and the 980 cm-l band was removed. In order to characterise bands due to the adsorbed ammonium cation, a tin(1v) oxide disc was protonated by treatment with hydrogen chloride vapour and then exposed to ammonia vapour (fig.8). In the presence of hydrogen chloride three new bands appeared in the spectrum at 1620, 1410 and 1100 cm-l. Evacuation removed the 1410 cm-l band and reduced the intensity of the 1620 cm-l band, which shifted to 1605 cm-l. Subsequent exposure to ammonia vapour produced strong absorption bands at 1620, 1410 and 1290 cm-l, a broad shallow band at ca. 1750 cm-l and a shoulder on the high-wavenumber side of the 1410 cm-l band. Pumping off the vapour phase reduced the intensity of all the bands, removing the 1750 cm-l band and the shoulder on the 1410 cm-l band completely; the 1620 cm-l band shifted to 1605 cm-l while the 1290 cm-l band shifted to 1270 cm-l.The 1615 and 1235 cm-l bands of adsorbed ammonia can be assigned to the v4(E) or 6,, and the v2(Al) or 6, modes of coordinatively bonded ammonia molecules, re~pectively.~*~ The decrease in intensity and shifts of the two ba1:ds from 1620 and 1240-1250 cm-l in the presence of the vapour to 1615 and 1235 cm-l in its absence suggest that there are two adsorbed ammonia species: one held on the surface by weak hydrogen bonding, and readily removed upon evacuation, the other being more strongly held, either coordinated or hydrogen-bonded to the surface. The position of the v, mode is consistent with the ammonia being coordinated to Lewis-acid tin sites,5 but the decrease in intensity of the bands with increasing evacuation temperature suggests that the ammonia is hydrogen-bonded to surface hydroxyl groups.The 1470 cm-l band is in the correct position for the bending mode of the ammonium ion (NHZ);6 its disappearance on evacuation and reappearance in the presence of water vapour suggests that it results from an interaction between water, or hydroxyl ions, and ammonia. However, on the protonated tin(1v) oxide surface the v4(&) or 6,, mode of the ammonium cation occurs at 1410 cm-l. The difference of 60 cm-l can be attributed to the strength of the Bronsted-acid site and the electrostatic interaction or ammonium ions with the surface. The energy required for the bending vibration of the 1470cm-l band is greater than that for the 1410cm-l band; therefore the interaction of the ammonium ions with the surface is less.This is as expected, since H,O should act as a weak acid, giving NH,+ and -OH ions adsorbed on the surface, while HCl should act as a strong acid, giving the NH: and C1- ions. A stronger electrostatic interaction would therefore be expected between the cation and the chloride ion adsorbed on the surface. Hydrogen bonding of ammonia to surface hydroxyl groups can be seen from the shift in position of the hydroxyl stretching band. The band position is restored on evacuation, which suggests that the remaining ammonia is coordinatively bonded to unsaturated tin ions. No N-H stretches due to coordinated ammonia or amide groups could be discerned against the intense hydroxyl stretching band.Rapid H-D exchange occurs on admitting deuterium oxide to an ammonia-treated tin(rv) oxide disc. The observed infrared bands can all be explained in terms of coordinated or hydrogen-bonded NH,, NH,D, NHD, and ND, species (table 1). In the presence of the D,O vapour the main species was adsorbed ND,. However, on evacuation the intensities of bands related to ammonia species with one or more hydrogens present were increased. The band shifts for deuterated ammonia agreed well with observed shifts on cobalt monoxide and calcium oxide;5 for NH,D v(NH,)/v(NH,D) M 1.08, for NHD, v(NH,)/v(NHD,) M 1.19, for ND, v(NH,)/v(ND,) z 1.34. A band was observed at ca. 2400 cm-l that could be due toP. G. HARRISON AND B. M. MAUNDERS Table 1. Observed vibrational bending modes of NH, D3-,(x = 0,1,2,3) species adsorbed on tin(rv) oxide band position/cm-' in the presence of D20 vapour after evacuation assignment 1600 (sh) 1435 (br) 1250 (vw) 1205 (s) 1 140 (vw) 1075 (w) 980 (sh) 900 (vs) 1349 the N-D symmetric stretching mode of ND, hydrogen-bonded to surface oxide ions.However, no other bands due to N-D stretches were discernable against the strong OD absorption band. The 1605 and 1 100 cm-l bands observed on a tin(rv) oxide disc exposed to hydrogen chloride vapour and subsequently evacuated can be assigned to the v, and v2 bending modes, respectively, of the hydroxonium cation. They occur in similar positions to the bending modes of the isoelectronic and isostructural ammonia molecule. Bands for the hydroxonium ion in strong mineral acids have been reported in similar p0sitions.~9 The nature of the 1410 cm-l band, observed prior to evacuation, is more difficult to assign.It occurs in the same position as the v, bending mode of the ammonium cation. However, as it disappears on evacuation it is unlikely to be due to traces of contaminating ammonia. One possibility for its cause is a doubly protonated water molecule involving both lone pairs of the oxygen. However, this is very unlikely, the H402+ ion having only been considered theoretically, whilst in practice it is unlikely to be stable. A more likely explanation would be the formation of an H,OZ, H7O; or H90: species, all of which are known to exist.9 The formation of the 1410 cm-l band on subsequent exposure to ammonia vapour is due to the v4(F2) bending mode of the ammonium cation, NHZ.The Bronsted acid responsible for its formation is the hydroxonium ion, since the 1 100 cm-l band disappears. The shoulder on the 1410cm-l band is probably due to the same species postulated for the 1470 cm-1 band on the untreated disc. Evacuation removed the shoulder and shifted the 1620 and 1290 cm-l bands to 1610 and 1260 cm-l, respectively. These two bands are due to the v4(E) and v2(Al) bending modes, respectively, of the coordinated or hydrogen-bonded ammonia. The 1750 cm-l is also removed by evacuation. A band has been observed in this region for ammonium halideslO-l= (other than phase I NHJ), where it has been attributed to the ~4(F2) + v, (rotary lattice vibration) combination band of the ammonium cation, when the ammonium ion is unable to freely rotate.No one simple adsorbed NH, structure can explain the results. On the untreated disc the ammonia may be hydrogen-bonded to surface oxide or hydroxyl groups or, more probably, coordinated to unsaturated tin sites. Similar bands are observed on the protonated surface, the main difference for adsorbed NH, groups being the position of the S, band, 1235 cm-l on the untreated and 1270 cm-l on the protonated1350 TIN OXIDE SURFACES -1800 1600 1400 1200 wavenum berlcm-’ Fig. 9. Infrared spectra of ammonia adsorbed on tin@) oxide-palladium oxide: (1) starting surface evacuated (320 K, 16 h); (2) during exposure to ammonia (2.6 kN m-2, 320 K); (3) subsequent evacuation (320 K, 1 h); (4) evacuation (365 K, 40 min); (5) evacuation (500 K, 2.5 h).surface. This can be explained as being due to coordinated ammonia molecules also involved in hydrogen bonding, uiz. H 1 H t Sn Sn Sn Sn \ /sn\ /s”\ untreated surface protonated surface \o/ Lo/ \() 0’ 0 0 0 Weak Bronsted-acid sites can be postulated on the untreated oxide surface from the appearance of the 1470 cm-l band in the presence of water vapour. Protonation of the surface causes strong Bronsted-acid sites. The infrared spectra of a tin(rv) oxide-palladium oxide disc evacuated at room temperature then treated with ammonia vapour are shown in fig. 9. Three new bands appeared in the presence of the vapour at 1612, 1470 (shoulder) and 1230 cm-l, the shoulder at 1470 cm-l being removed on evacuating the cell.Exposure to water vapourP. G. HARRISON AND B. M. MAUNDERS 1351 I I I I I I I 1800 1600 1400 1200 wavenumber/cm -' Fig. 10. Infrared spectra of tin(1v) oxide-palladium oxide exposed to hydrogen chloride and then ammonia: (1) starting surface of disc evacuated (320 K, 40 h); (2) during exposure to hydrogen chloride (9.31 kN m-2, 320 K); (3) subsequent evacuation (320 K, 1 h); (4) during exposure to ammonia (9.31 kN m-2, 320 K); (5) subsequent evacuation (320 K, 1 h). caused the 1612 and 1230 cm-l bands to shift to 1620 and 1250 cm-l, respectively, and restored a band at 1450-1460 cm-l. Subsequent evacuation at 500 K reduced in intensity the 1612 and 1230 cm-1 bands, now shifted slightly to 1610 and 1235 cm-l, and produced a new band at 1510 cm-l. The strong hydroxyl stretching vibrations prevented the observation of any N-H stretching modes.When a tin(1v) oxide-palladium oxide disc was heated in vacuu and treated with oxygen at 563 K, the intensity of the hydroxyl stretching band increased and shifted slightly to lower frequency in the presence of the vapour, and new bands were observed at 1615 and 1225 cm-l. Evacuaticn of the cell did not restore the OH band to its original intensity but reduced in intensity the 1615 and 1225 cm-l bands, shifted to 1605 and 1230 cm-l, and a very weak band was discernible at ca. 535 cm-l. Evacuation at 393 K removed the 1605, 1535 and 1230 cm-l bands. A tin(rv) oxide-palladium oxide disc exposed to hydrogen chloride vapour at the ambient beam temperature exhibited three new bands at 1615, 1405 and 1100 cm-l (fig.10). Pumping off the vapour phase removed the 1405 cm-l band and decreased the intensity of the 1615 and 1100 cm-l bands, moving them to 1580 and 1090 cm-l, respectively. Subsequent exposure to ammonia vapour produced bands at 1730 (broad1352 TIN OXIDE SURFACES I I 1 I I 1800 1600 1400 1200 wavenum ber/cm -' Fig. 11. Infrared spectra of ammonia adsorbed on tin(1v) oxide-silica: (1) starting surface of tin@) oxide-silica evacuated at 320 K; (2) during adsorption of ammonia (1.33 kN mU2, 320 K); (3) subsequent evacuation (320 K, 24 h); (4) disc heated under vacuum at 483 K, during exposure to ammonia (1.33 kN mP2, 320 K); (5) disc heated under vacuum at 693 K, during exposure to ammonia (2.6 kN m-2, 320 K); (6) after evacuation of (5) and during exposure to water vapour (1.33 kN m-2, 320 K).and weak), 1615 (strong), 1410 (very strong) and 1290cm-1 (strong) but removed the 1090 cm-I band. A shoulder on the high-frequency side of the 1410 cm-l band was also observed. Evacuating the cell removed the 1730 cm-l band and the shoulder on the 1410cm-l band, and the 1615, 1410 and 1290cm-l bands were reduced in intensity and shifted to 1605, 1405 and 1270 cm-l, respectively. The 1090 cm-I band was restored but not to its original intensity. The behaviour of tin(1v) oxide-palladium oxide toward ammonia and hydrogen chloride vapour is very similar to the behaviour of tin(1v) oxide. As such, the 1612 and 1230 cm-l bands, produced on adsorption of ammonia, can be assigned to the v4(E) and v2(A,) bending modes, respectively, of coordinatively bonded, or hydrogen- bonded, ammonia molecules, while the shoulder at 1470 cm-' can be ascribed to the bending mode of an ammonium species, formed from weak Bronsted-acid sites.The formation of the band at 1510 cm-l, and decrease in the 1612 and 1230 cm-l bands, on evacuation at 500 K can be ascribed to the formation of surface amido groups from the adsorbed ammonia. Bending vibrations of surface amido groups have been variously reported between 1470 and 1570 cm-l on different aluminas, silicas,P. G. HARRISON AND B. M. MAUNDERS 1353 4 000 3500 3000 2500 wavenumber/cm-' Fig. 12. Effect of ammonia adsorption of the hydroxyl stretching band of tin(rv) oxide-silica heated under vacuum at 693 K: (1) starting surface; (2) during exposure to ammonia vapour (2.6 kN m-2, 320 K); (3) subsequent evacuation (320 K, 4.5 h).aluminium phosphate and germania gels.5 In conjuction with the decrease in the 16 12 and 1230 cm-l band intensities, slight shifts in position to 1610 and 1235 cm-l occurred, thus indicating that there were two forms of adsorbed ammonia present, one of which reacted to form the amide. A possible precursor could be a coordinatively held ammonia molecule involved in hydrogen bonding to surface oxide ions. The intense nature of the hydroxyl stretching band made it impossible to judge whether or not there was a concomitant increase in its intensity with the appearance of the 1510 cm-l band. However, a definite increase in intensity was observed when a disc heated and oxygen treated at 563 K was exposed to ammonia vapour.The 1615 and 1100 cm-l bands, shifting to 1580 and 1090 cm-l on evacuation, observed on exposure of a tin(rv) oxide-palladium oxide disc to hydrogen chloride vapour, can be assigned to the vq and v2 bending modes, respectively, of the hydroxonium cation. The 1 730,16 15,14 10 + shoulder and 1290 cm-l bands observed on subsequent exposures to ammonia vapour and their behaviour on evacuation are analogous to the bands observed on a similarly treated tin@) oxide disc; accordingly the band assignments are the same. The infrared spectra of a tin@) oxide-silica disc evacuated and exposed to ammonia vapour at Tab exhibited three new absorption bands at 1620, 1455 and 1245 cm-l (fig. 11). On evacuation the 1455 cm-l band was reduced in intensity but not removed.Evacuation had little affect on the other bands. Three bands at 1620, 1465 and 1240 cm-l were observed on exposure to ammonia of a disc heated under vacuum at 483 K, although the intensity of the 1465 cm-l band was considerably less 45 FAR 11354 TIN OXIDE SURFACES 1800 1600 1400 1200 wavenum ber/cm Fig. 13. Infrared spectra of hydrogen chloride then ammonia vapour absorption on a tin(1v) oxide-silica disc heated under vacuum at 693 K: (1) starting surface; (2) during adsorption of hydrogen chloride (2.0 kN m-2, 320 K); (3) subsequent evacuation (320 K, 1 h); (4) during adsorption of ammonia (2.66 kN m-2, 320 K); (5) subsequent evacuation (320 K, 15 h). than that of the unheated disc under similar conditions of ammonia vapour.After evacuation at 693 K, the 1465 cm-l band was very weak and shifted to ca. 1490 cm-l. The bands at 1620 and 1245 cm-l were also weaker and shifted to 1615 and 1240 cm-l, respectively. Exposure to water vapour produced a fairly intense band at 1445 cm-l which was removed on subsequent evacuation. The formation of a broad intense band was also observed, with its maximum at 3200 cm-l, on the tin(xv) oxide-silica disc heated under vacuum at 693 K (fig. 12). A tin@) oxide-silica disc heated under vacuum at 693 K and exposed to hydrogen chloride vapour exhibited two weak bands at 1620 and 1410 cm-l (fig. 13). Subsequent exposure to ammonia vapour produced absorption bands at 1750 (broad and shallow), 1610 (medium), 1410 (very strong) and 1280 cm-l (strong). The intensity of the 16 10,141 0 and 1280 cm-l bands were reduced by evacuation and the 1750 cm-l band removed.The 1620 and 1245 cm-l bands can be attributed to the v4(E) and v2(A1) bending modes, respectively, of coordinatively bonded or hydrogen-bonded ammonia mol- ecules, while the 1455 cm-l band can be ascribed to the ~4(F2) bending mode of the ammonium cation. Such a band has been observed in this region on silica-alumina mixed-oxide systems.69 l3 Unlike the other two oxides the 1455 cm-l band is notP. G. HARRISON AND B. M. MAUNDERS 1355 removed by evacuation. The effect of increased pretreatment temperature is to reduce considerably the number of Bronsted-acid sites, as observed from the weaker band intensity. The shift of the v4 band intensity to higher frequency suggests a weaker interaction between the NH,+ ion and the surface at the higher evacuation temperatures.Bronsted-acid sites could be reformed by exposure to water vapour. The 1410 cm-l band observed after ammonia adsorption on the disc treated with hydrogen chloride and heated to 693 K can be assigned to the v,(F,) bending mode of the ammonium cation, involved in a stronger interaction with the surface than the ammonium cation produced on the untreated oxide. The 3200cm-I band can be assigned to the N-H stretching vibrations of the adsorbed ammonia. The broad nature of the band made identification of the nature of the N-H stretches difficult, although the band is at too low a frequency for it to be due to amido N-H stretches. No evidence was seen for amido groups in the N-H bending region.DISCUSSION Infrared spectroscopy of ammonia adsorption on metal oxides has been employed as a method of studying the nature of acidic sites on oxide^.^^ 6~ 14-31 Pure oxides usually exhibit only Lewis l4? 2 o v 2 3 y 25 although this is not always the 2 4 9 29 while mixed oxides exhibit Bronsted acidity.l59 2 4 9 26* 27 The infrared absorption bands of ammonia adsorbed on tin(1v) oxide at 1615 and 1235 cm-l clearly indicate that ammonia is coordinated to Lewis-acidic sites on the tin(1v) oxide surface, i.e. to surface tin sites. Similar bands have been observed for ammonia adsorption on GeO2,ls Ti0,59 2 o q 21* 31 (both rutile and anatase), A1203,5914922-24731 Zn0,29931 Ni0,5731 Mg0,5931 Zr0,,5931 Be0,5931 Ga,O,,,l Ta,05,31 Cr,O, and Fe20,.6* 25 The shift to higher frequency of the 6,(NH,) deformation mode (at 950 cm-I in free ammonia) on coordination has been employed as a guide to the Lewis acidity of metal centres in these oxides.The observed order [6,(NH,) mode in parentheses], Al,O, (1 3 15-1283 cm-l), Al,O, * SiO, (1 3 10-1280 cm-l) > Ga,O, (1275 cm-l) > TiO, (1225 cm-l) > ZnO (1 220 cm-l), Be0 (1220 cm-l) > Cr,O, (1 220 cm-l) > ZrO, (1 200- 1 160 cm-l) > NiO (1 1 80-1 1 30 cm-l) > MgO (1 170 cm-l) > Ni,O, (1 145 cm-l) > COO (1 130 cm-l), therefore ranks the Lewis acidity of the bare tin sites of tin(rv) oxide between those of Ga,O, and TiO,. Analogous bands observed for ammonia on tin(1v) oxide-palladium oxide (1 605-1 6 10 and 1230-1235 cm-l) and tin(1v) oxide-silica (1620 and 1240-1245 cm-l) indicate that the strengths of the Lewis-acid sites are slightly less and slightly stronger, respectively, than those on the pure oxide.Weak Bronsted-acid sites are observed for all three oxides. An important difference, however, is the appearance of the band at 1510-1 535 cm-l, together with an increase in intensity of the hydroxyl stretching band, after evacuation at 500 K, which is attributed to dissociative chemisorption giving rise to surface amido groups on palladium sites similar to those found on palladium metal.30 Several other metals also chemisorb ammonia to give surface amido 22* 23* 32 including silica, which does not exhibit any Lewis acidity towards amrnonia.l7-l9 CONCLUSIONS Tin(rv) oxide is predominately Lewis acidic, with weak Bronsted acidity evolving in the presence of water vapour.Strong Bronsted-acid sites can be produced by protonation with hydrogen chloride. 45-21356 TIN OXIDE SURFACES The nature of tin@) oxide-palladium oxide is quite similar to tin(1v) oxide, exhibiting predominately Lewis acidity, although weak Bronsted acidity can be observed in the presence of water-vapour. Strong Bronsted-acid sites can be produced by treating the surface with hydrogen chloride vapour. The main difference is that surface amido groups are formed on tin(1v) oxide-palladium oxide after evacuation at 500 K, whereas no bands ascribable to surface amido groups were observed on tin(rv) oxide under any conditions. It seems reasonable, therefore, to assume the amide is forming on palladium sites. Unlike tin@) oxide or tin(1v) oxide-palladium oxide, tin@) oxide-silica exhibits both Lewis and Bronsted acidity, the amount of Bronsted acidity with respect to Lewis acidity decreasing with increasing evacuation temperature.P. G. Harrison and E. W. Thornton, J. Chem. Soc., Faraday Trans. 1, 1976, 72, 1310; 1317; 2484; and references cited therein. P. G. Harrison and E. W. Thornton, J. Chem. SOC., Faraday Trans. I , 1978, 74, 2703. M. Van Thiel, E. D. Becker and G. C. Pimentel, J. Chem. Phys., 1957, 27, 486. K. Nakamoto, Infrared Spectra of Indrganic and Coordination Compounds (Wiley, London, 1970). A. A. Tsyganenko, D. V. Pozdnyakov and V. N. Filimonov, J. Mol. Struct., 1975, 29, 299. M. L. Hair, Infrared Spectroscopy in Surface Chemistry (Arnold, London, 1967). ' D. E. Bethel1 and N. Sheppard, J. Chem. Phys., 1953, 21, 1421. C. C. Fenso and D. F. Hornig, J. Chem. Phys., 1955, 23, 1464. I. Olovsson, J. Chem. Phys., 1968, 49, 1063. lo E. L. Wagner and D. F. Hornig, J. Chem. Phys., 1950, 18, 296; 305. l1 R. C. Plumb and D. F. Hornig, J. Chem. Phys., 1953, 21, 366. l2 R. C. Plumb and D. F. Hornig, J. Chem. Phys., 1955, 23, 947. l3 L. H. Little, Infrared Spectra of Adsorbed Molecules (Butterworths, London, 1966). l4 R. P. Eischens and W. A. Pliskin, Adv. Catal., 1958, 10, 1. l5 J. B. Pen, Discuss. Faraday Soc., 1971, 52, 55. l6 M. J. D. Low and K. Matsushita, J. Phys. Chem., 1969, 73, 908. l8 N. W. Cant and L. H. Little, Can. J. Chem., 1965, 73, 1252. IQ B. A. Morrow and I. A. Cody, J. Phys. Chem., 1976,80, 1098. 2o G. D. Parfitt, J. Ramsbottom and C. H. Rochester, Trans. Faraday Soc., 1971,67, 841. 21 M. Herrmann and H. P. Boehm, 2. Anorg. Allg. Chem., 1969, 73, 1368. 22 A. A. Tsyganenko, D. V. Pozdnyakov and V. N. Filimonov, Usp. Fotoniki, 1975,5, 150. 23 J. B. Pen, J. Phys. Chem., 1965, 69, 231. 24 A. A. Davydov and Yu. M. Shchekochikhin, Kinet. Catal., 1969, 10, 523. 25 G. Blyholder and E. A. Richardson, J. Phys. Chem., 1962, 66, 2599. 26 J. E. Mapes and R. P. Eischens, J. Phys. Chem., 1954,58, 1059. 27 N. W. Cant and L. H. Little, Nature (London), 1966, 69, 21 1. 28 M. M. Mortland, J. J. Fripiat, J. Chaessidon and J. Uytterhoeven, J. Phys. Chem., 1963, 67, 248. T. Morimoto, H. Yani and M. Nagao, J. Phys. Chem., 1976,80,471. 30 D. V. Pozdnyakov and V. N. Filimonov, Kinet. Catal., 1972,13, 522. 31 N. E. Tret'yakov and V. N. Filimonov, Kinet. Katal., 1973, 14, 803. 32 G. M. Zhabrova and E. V. Egorov, Russ. Chem. Rev., 1961,30, 338. J. B. Peri, J. Phys. Chem., 1966, 70, 2937. (PAPER 3/667)

ISSN:0300-9599    

DOI:10.1039/F19848001341

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1984,80, 1357-1365 Tin Oxide Surfaces Part 13.-A Comparison of Tin(1v) Oxide, Tin@) Oxide-Palladium Oxide and Tin(rv) Oxide-Silica: an Infrared Study of the Adsorption of Carbon Dioxide BY PHILIP G. HARRISON* AND BARRY M. MAUNDERS Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD Received 28th April, 1983 The adsorption of carbon dioxide onto tin(1v) oxide, tin(1v) oxide-palladium oxide and tin(1v) oxide-silica heat-treated in the temperature range 320-703 K has been studied by infrared spectroscopy. The surface species formed depend on the pretreatment temperature of the oxide. Bicarbonate and unidentate carbonate are formed at low pretreatment temperatures on all three oxides, but the higher Bronsted acidity of tin(rv) oxide-silica allows more bicarbonate formation than the other two oxides.At pretreatment temperatures 2 473 K the major adsorption products are unidentate carbonate and bidentate carbonate, together with small amounts of ‘organic’ or bridging carbonate. Complementary to the ammonia-adsorption study described in the previous paper1 is the investigation of the chemisorption of carbon dioxide onto the same three oxides, tin@) oxide, tin@) oxide-palladium oxide and tin(1v) oxide-silica. Ammonia adsorption indicated significant differences in the behaviour of these three oxides, in particular the higher Bronsted acidity of tin(1v) oxide-silica and the ability of surface palladium sites to induce dissociative chemisorption of ammonia, producing surface amido groups.Adsorption studies employing carbon dioxide should provide further information regarding the comparative nature of the surfaces of these oxides, especially ‘ acid-base ’ sites, where a reactive surface (hydr)oxide closely adjoins an ‘ exposed ’ or incompletely coordinated metal ion, which are also important in several cat a1 y tic processes. RESULTS The adsorption of carbon dioxide on a tin(1v) oxide disc pretreated at 320 K gave rise to only very weak bands centred at ca. 1575-1 580 and 1430-1440 cm-l, regardless of the pressure of carbon dioxide. Heat pretreatment at 390 K (fig. l), however, produced adsorption bands of increased intensity at 1575-1580 and 1430-1440 cm-l, along with a new, weak band at 1370-1375 cm-l and a weak shoulder at 1460 cm-l.Pumping off the carbon dioxide totally removed the 1575-1 580 cm-l band and greatly reduced the 1430-1440cm-1 band, leaving only a weak broad band at 1450cm-l, while the 1370-1 375 cm-l band was largely unaffected. The infrared spectra became more complex when the pretreatment temperature was raised to 473 K (fig. 2), with absorption maxima occurring at 1700-1720,159O-1595, 1430,1365-1 370, 1295-1 300 and 1220 cm-l, along with the shoulder at 1460 cm-l. The 1590-1595 cm-l band was sharp on the high-wavenumber side but broad on the low-wavenumber side, with a slight shoulder at ca. I550 cm-l. Again, on evacuating the cell only the 1365-1 370 cm-l 13571358 58 83 83 81 n g 5 Q) c.’ .g c 2 w 67 - 10 O/O - TIN OXIDE SURFACES I 1 1 I I I I I 1800 1600 1400 1200 wavenum ber/cm -1 Fig.1. Infrared spectra of carbon dioxide adsorbed at 320 K on tin(rv) oxide evacuated at 390 K. (1) Starting surface; carbon dioxide pressure: (2) 1.6, (3) 4.39 and (4) 11.04 kN m-*; (5) subsequent desorption (390 K, 1 h). band was unaffected, and a weak band at 1450 cm-l remained. Although the intensity of the absorption bands appears greatly reduced on a disc pretreated at 623 K compared with a 473 K pretreated disc, the actual difference is not that great. When the absorption bands at ca. 1590 cm-l were converted from percentage transmittance to absorbance units, the peak heights were measured as 0.159 and 0.131 units for the 473 and 623 K pretreated discs, respectively, at a carbon dioxide pressure of 1.33 kN m-2. The spectra obtained with carbon dioxide adsorption on tin(1v) oxide-palladium oxide discs were similar to those formed on the pure tin(1v) oxide discs, with some slight differences in the band positions. The other main differences were that on the 320 K pretreated disc, the resulting spectra were more intense than on the similarly treated tin(1v) oxide disc, and resembled more the spectra of the tin(1v) oxide disc pretreated at 390 K, with a band being present at 1380 cm-l and the 1595 cm-l band being broader on the low-wavenumber side with a shoulder at ca.1555 cm-l. After evacuation of the cell only the 1380cm-l band and a weak band at 1450cm-l remained. Pretreatment of a disc at 393 K (fig. 3) had little effect on the number of absorption bands but increased the intensity of the 1440 cm-l band, with respect to the 1590 cm-l band, and the 1378 cm-l band.More intense absorption bands at 1440 and 1378 cm-l remained upon subsequent evacuation in comparison with the bandsP. G . HARRISON AND B. M. MAUNDERS 1359 V I I I I I I 1800 1600 1400 1200 w avenumberlcm -' Fig. 2. Infrared spectra of carbon dioxide adsorbed at 320 K on tin(1v) oxide evacuated at 473 K. (1) Starting surface; carbon dioxide pressure: (2) 1.33, (3) 4.39 and (4) 10.11 kNm-2; (5) subsequent desorption (473 K, 1.5 h). remaining on the 320 K pretreated disc. In a similar manner to the tin(rv) oxide disc, a tin@) oxide-palladium oxide disc pretreated at 483 K gave rise to new absorption bands at 1700-1725 cm-l (broad and weak), 1300 cm-l (weak) and 1222 cm-l. The 1595-1 600 cm-l band, shifted from 1590 cm-l, appeared sharper on the high- wavenumber side than before.The intensity of the 1222 cm-l band was significantly greater when the pretreatment temperature was raised to 593 K (fig. 4) and a slight shift in the 1700-1725 cm-l band to ca. 1740 cm-l was observed. Carbon dioxide adsorption on tin(rv) oxide-silica differed from the previous two oxides in so far as that on the discs pretreated at 320 and 383 K (fig. 5) only absorption bands at 1600-1590 and 1440-1448 cm-l were observed, with the slight shoulder on the 1440-1448 cm-l band being present as on the previous two oxides. No absorption band was observed in the 1370-1 380 cm-l region until the pretreatment temperature was raised to 473 K, when a weak band appeared at 1380 cm-l, along with very weak bands at 1700-1720 and 1350 cm-l and a band at 1222 cm-l, in a similar fashion to the bands formed on tin(1v) oxide and tin(1v) oxide-palladium oxide.Also, similarly to the previous two oxides, the 1590-1595 cm-l band now appeared sharper on the high-wavenumber side. Pretreatment at 593 K (fig. 6) increased the intensities of the 1700-1720 cm-l band, shifting it to 1760-1780 cm-l, and the 1370-1380 cm-l band. The 1350 cm-l band was effectively removed. Finally, unlike the other two oxides,1360 TIN OXIDE SURFACES 52 75 1800 1600 1400 1200 wavenum berlcm-' Fig. 3. Infrared spectra of carbon dioxide adsorbed at 320 K on tin(1v) oxide-palladium oxide evacuated at 393 K. (1) Starting surface; carbon dioxide pressure: (2) 1.73, (3) 5.19 and (4) 851 kN m-2; subsequent desorption at (5) 320 K, 10 min and (6) 393 K, 2.5 h.when the pretreatment temperature was raised to 703 K the 1370 cm-l band became quite strong and sharp in comparison with the 1600 cm-l band. The 1440 cm-l band was also replaced by a broad band centred at 1460 cm-l, the 1222 cm-l band was removed and the 1760-1780 cm-l band was further shifted to ca. 1800 cm-l. DISCUSSION Carbon dioxide may be absorbed on oxide surfaces in many ways: physically adsorbed CO,, a carboxylate-type group, uncoordinated carbonate, a unidentate or bidentate carbonate group, an ' organic-type ' carbonate group and a bicarbonate group have all been reported by various authors.2-10 The bands which occur at ca. 1580-1590 and 1430-1450 cm-l on all three oxides at low pretreatment temperatures can be assigned to the v4 and v, modes, respectively, of a surface bicarbonate species.Miller and Wilkinsll have reported bicarbonate bands at 1660-1632 and 1410-1300 cm-l for some inorganic bicarbonate species, whilst Parkynsgp lo has reported bicarbonate bands at ca. 1640 and 1480 cm-l for carbon dioxide adsorbed on alumina. This assignment is in conflict with our previousf2 assignments for carbon dioxide adsorption on tin@) oxide, in which we failed to assign the 1580-1585 cm-l band and incorrectly assigned the band at 1222 cm-l to the S(C0H) vibration of a surface bicarbonate, which we suggested did not form on oxideP. G. HARRISON AND B. M. MAUNDERS 1361 n 40 9 42 5 42 E 42 2 2 Y Y .- * - 10 % - 33 1800 1600 1400 1200 wavenumber/crn -' Fig.4. Infrared spectra of carbon dioxide adsorbed at 320 K on tin(1v) oxide-palladium oxide evacuated at 593 K. (1) Starting surface; carbon dioxide pressure: (2) 1.60, (3) 5.45 and (4) 8.91 kN m-2; (5) subsequent desorption (320 K, 0.5 h). surfaces pretreated below 508 K. Since bicarbonate formation might be expected to occur with greater facility on more highly hydroxylated surfaces, i.e. those pretreated at low temperatures, we consider that our previous suggestion is in error, and that bicarbonate is only formed in significant quantities on discs pretreated at low temperatures, and is only present in small amounts on high-temperature (2 473 K) pretreated tin(rv) oxide. The 1370-1380 cm-l band, observed on tin(1v) oxide and tin(1v) oxide-palladium oxide, can be assigned to the v,(A,) mode of a unidentate carbonate group [v(C-O~~)+V(C-O~)] for which the v5(B2) mode [v(C-OII)] is involved in the 1430-1450cm-l band.This assignment is in good agreement with the unidentate carbonate bands assigned by Nakamoto13 for some cobalt complexes, and agrees with previous similar assignments.12 The vz(A1) mode, reported by Nakamoto at 1070-1050cm-1, is not observed, and is presumably obscured by the bulk oxide absorptions. The size of the 1370-1 380 cm-l band, compared with the 1580-1 590 cm-l band, suggests that the unidentate carbonate is a minor product. The shoulder at 1555 cm-l is probably due to a surface carboxylate species, the corresponding symmetric stretching mode for which may be the weak band at 1300 cm-l.The shoulder at 1460 cm-l could be due to the v5 mode of a second unidentate carbonate species. A further complicating factor is that an uncoordinated carbonate ion is expected to exhibit an antisymmetric stretching vibration in the1362 59 63 73 72 h E Q, E Y * -g C E Y 63 - 10 Ole- - TIN OXIDE SURFACES 1 5J-J I I I I 1 I I 1800 1600 1400 1200 wavenumber/cm -1 Fig. 5. Infrared spectra of carbon dioxide adsorbed at 320 K on ti@) oxide-silica evacuated at 383 K. (1) Starting surface; carbon dioxide pressure: (2) 1.33, (3) 4.66 and (4) 7.71 kN m-*; (5) subsequent desorption (383 K, 3 h). 1440-1460 cm-l region. The symmetric stretching frequency for this species occurs at too low a wavenumber to be observed, so no further evidence exists for an uncoordinated carbonate structure.With higher pretreatment temperatures the 1580-1 595 cm-I band shows a marked change, shifting slightly in position and becoming sharper on the high-wavenumber side; also a sharp band appears at 1220-1222 an-’ in each case. These two bands can be assigned to the v,(A,) [v(C-011) + v(C-O,)] and v5(B2) [v(C-0,) + S(0, * CO,,)] modes of a bidentate carbonate species, respectively. The band positions are in good agreement with assignments of bidentate carbonate bands by Nakamoto13 for cobalt complexes. The formation of surface bidentate carbonate at these pretreatment temperatures, but not below, can be justified on the grounds that these are the temperatures at which surface hydroxyl groups are condensing to eliminate water, and so leave surface oxide groups with adjacent bare metal cations.High pretreatment temperatures will also give rise to ‘ strained’ oxide linkages, which can also react with carbon dioxide to produce the bridging or ‘organic’ type of surface carbonate responsible for the weak absorption bands at 1700-1725 cm-l. On tin(1v) oxide-palladium oxide the 1700-1725 cm-l band shifts to 1740 cm-l on a disc pretreated at 593 K, while on the tin(xv) oxide-silica it shifts to 1800 cm-l for a 703 K pretreated disc. At the same time on the 473-483 K pretreated discs weak bands were observed at 1295-1 300 cm-l [tin@) oxide and tin@) oxide-palladiumFig. 6. at 593 P. G. HARRISON AND B. M. MAUNDERS 1363 1800 1600 1400 1200 wavenurn ber/crn Infrared spectra of carbon dioxide adsorbed at 320 K on tin(rv) oxidesilica evacuated K.(1) Starting surface; carbon dioxide pressure: (2) 1.46, (3) 5.73 and (4) 9.18 kN m-2; (5) subsequent desorption (320 K, 0.75 h). Table 1. Carbon dioxide adsorbed on tin@) oxide pretreatment band position/cm-l temperature/K 320 - - 1575-1580 (w, br) - 473 1700-1720 (w, br) 1590-1595 (vs) - 1460 (sh) 623 1700-1720 (w, br) 1585 (m) - 1460 (sh) 390 - - 1575-1580 (s) 1460 (sh) VIII I, VI I 11, I11 ~~ 320 1430-1440 (w, br) - - - 473 1430 (vs) 1365-1 370 (w) 1295-1300(~) 1220 623 1430 (m) 1365-1370(~) 1295-1300(~) 1220 - - 390 1430-1440 (m) 1370-1 375 (vw) IV V IX VII Assignments: I, vq bicarbonate; 11, v, bicarbonate; 111, uncoordinated carbonate ion?; IV, vg unidentate carbonate; V, v, unidentate carbonate; VI, v, bidentate carbonate; VII, v, bidentate carbonate; VIII, C=O stretch of ' organic-type' carbonate; IX, CO, asymmetric stretch of ' organic-type' carbonate.w o\ P Table 2.Carbon dioxide adsorbed on tin(1v) oxide-palladium oxide ~ ~~ pretreatment band position/cm-l temperature/K 320 - - 1595 (s) 1555 (sh) 1460 (sh) 1440 (m) 1380 (w) - - 393 - 1590 (s) 1560 (sh) 1460 (sh) 1440 (s) 1375-1380 (w) - - VIII I, VI I X 11, I11 IV V IX VII - 483 1700-1 720 (vw, br) 1600 (s) - 1560 (sh) 1460 (sh) 1440 (s) 1380 (w) 1300 (w) 1222 (m) - 1740 (w, br) 1590 (s) - 1560 (sh) 1460 (sh) 1430 (s) 1380-1385 (w) 1300 (w) 1222 (vs) - 4 2 A Assignments: I-IX, as table 1; X, CO, asymmetric stretch of carboxylate anion. 2 e Table 3. Carbon dioxide adsorbed on tin@) oxide-silica 9 $2 pretreatment band position/cm-l I/] temperature/K 320 - - 1595-1 600 (s) 1460 (sh) 1445-1450 (m) - - - - - 1460 (sh) 1440 (s) - - - 1590 (s) 383 473 1700-1720 (vw, br) 1592 (vs) - 1460 (sh) 1435-1440 (VS) 1380 (vw) 1350 (vw) 1222 (m) - 1222 (m) 593 1760-1780 (v, br) 1595 (vs) - 703 1800 (w, br) 1600 (s) - 1460 (m) 1430 (sh) 1370 (m) - - - VIII I, VI I 11, I11 IV V IX VII 1460 (sh) 1435-1440 (VS) 1380 (w) Assignments: as table 1.P. G.HARRISON AND B. M. MAUNDERS 1365 oxide] and 1350 cm-1 [tin(rv) oxide-silica] which can be assigned to the corresponding antisymmetric CO, stretch, Such shifts in the position of the v(C=O) stretching band can be rationalised in terms of increased covalency of the metal-oxygen bond of the carbonate group. The separation between the v(C-0,) and v(C-OII) stretching modes of carbonate species is sensitive to the mode of bonding and is larger for bidentate carbonate than for unidentate ~arb0nate.l~ With increasing covalency of the metal-oxygen bond the separation between the two bands increases further, so that with completely covalent bonding (as in dimethylcarbonate) the separation is ca.600 cm-l. The main effect of covalency is in the C-0,, bond, whilst the C-0, bond, despite being attached to the metal, is hardly perturbed. In the present case the Si-0 bond is significantly more covalent than the Sn-0 bond. Thus on tin@) oxide-silica pretreated below 703 K, the ‘organic-type’ carbonate can be presumed to bridge two tin ions. However, above 703 K the observed shift can be rationalised in terms of a bridged carbonate involving at least one C-0 bond attached to silicon, which has become available owing to dehydroxylation of SiOH groups at this temperature.The smaller shift to higher wavenumbers for this band, in the case of carbon dioxide adsorption on tin(rv) oxide-palladium oxide pretreated at 593 K, can be taken to indicate a smaller increase in the covalent character of the M-0 bond than occurs on tin(rv) oxide-silica. Our previous assignments of the sharp band at I222 cm-l to the d(C0H) deformation mode of a surface bicarbonate, made by analogy to the 1233 cm-l band observed after carbon dioxide adsorption on al~mina,~ is obviously erroneous. The revised assignment to the v, mode of a bidentate carbonate is preferred since the band does not appear until higher pretreatment temperatures.* As the pretreatment temperature is raised, the surface bidentate carbonate structure becomes more favoured. This is particularly clear in the case of adsorption on tin(rv) oxide-palladium oxide. The reason for this is uncertain, although it may be due to a steric effect, whereby increased dehydroxylation at higher temperatures permits more adsorption sites with less steric hindrance, allowing two oxygens to come in close approach of the surface. A second possibility is that a certain amount of weak hydrogen bonding may occur with a unidentate carbonate, which prevents the formation of bidentate carbonate. The observed bands for all three oxides are summarised in tables 1-3, with appropriate assignments. P. G. Harrison and B. M. Maunders, J. Chem. SOC., Faraday Trans. I , 1983, 80, 1341. J. H. Taylor and C. H. Amberg, Can. J. Chem., 1961, 39, 535. J. B. Peri, J. Phys. Chem., 1966, 70, 3168. 4 R. P. Eischens and W. A. Pliskin, Adu. Catai., 1957, 2, 662. 5 C. E. O’Neill and D. J. C. Yates, Spectrochim. Acta, 1961, 17, 953. ? Y . Fukuda and K. Tanabe, Bull. Chem. SOC. Jpn, 1973,46, 1616. * J. V. Evans and J. L. Whateley, Trans. Faraday SOC., 1967, 63, 2769. N. D. Parkyns, J. Chem. SOC. A, 1969,410. lo N. D. Parkyns, J. Phys. Chem., 1971, 75, 526. l1 F. A. Miller and C. H. Wilkins, Anal. Chem., 1952, 24, 1253. 12 E. W, Thornton and P. G. Harrison, J. Chem. SOC., Faraday Trans. I , 1975,71,461. l 3 K. Nakamoto, Injrared Spectra of Inorganic and Coordination Compounds (Wiley, London, 1970). l4 M. L. Hair, Infrared Spectroscopy in Surface Chemistry (Arnold, London, 1967). l5 A. Guest and P. G. Harrison, unpublished data. M. Courtois and S. J. Teichner, J. Catai., 1962, 1, 121. (PAPER 3/668) * Recent work on the adsorption onto tin(1v) oxide from synthetic air containing 30-50 ppm levels of CO has demonstrated unequivocally that the band at 1222 cm-’ is not due to a surface bicarbonate species since identical spectra are obtained from both hydroxylated and deuteroxylated surface^.'^

ISSN:0300-9599    

DOI:10.1039/F19848001357

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraahy Trans. I , 1984,80, 1367-1376 Properties of Y-type Zeolites with Various SiliconlAluminium Ratios Obtained by Dealumination with Silicon Tetrachloride Distribution of Aluminium and Hydroxyl Groups and Interaction with Ethanol BY LUDMILA KUBELKOV~~,* VLASTIMIL SEIDL,~ JANA NOVLKOVA, S O ~ A BEDNMOVL AND PAVEL JfRe J. Heyrovskg Institute of Physical Chemistry and Electrochemistry, Czechoslovak Academy of Sciences, Machova 7, 121 38 Prague 2, Czechoslovakia Received 28th March, 1983 Y-type zeolites with %/A1 ratios of 3.420, prepared by the dealumination of NaY (%/A1 = 2.5) with SiCl,, have been characterized using X-ray diffraction analysis, sorption- capacity data and by obtaining infrared spectra of the framework, OH groups and adsorbed pyridine. Dealuminated zeolites exhibit good crystallinity and some aluminium released from the lattice remains in extralattice positions in the zeolite cavities or in an amorphous phase to provide strong electron-acceptor sites.The zeolites were obtained directly in the H form; further exchange with NHZ affected their proton-donor properties only slightly. In addition to weakly acidic or non-acidic hydroxyls (bands at 3740 and 3620 cm-l), strongly acidic framework hydroxyls were present (bands at 3630 and 3560 cm-l) in amounts dependent on the amount and distribution of aluminium. An analysis of substances released during temperature- programmed desorption from zeolites with preadsorbed ethanol revealed the influence of dealumination on the catalytic activity: the maximum of evolved ethylene as a main product was reached on moderately dealuminated zeolites (Si/Al = 3.4-6) with an appreciable number of both strong electron-acceptor and strong proton-donor sites.In 1980 Beyer published a new method for the thermochemical modification of zeolites1 consisting of the substitution of aluminium in the lattice by silicon using the reaction of the zeolite with SiCl,. Studies have so far concerned the struct~rel-~ of the highly dealuminated Y zeolites obtained by this method (Si/Al = 20-50), their selectivity for adsorption of n-hexane, n-butane, benzene, ammonia and water,l a demonstration of the presence of strong proton-donor sites4 and a comparison with HZSM-5 and HZSM- 1 1 with regard to isomerization and hydrocracking of n-decane on Pt-loaded zeolite^.^ The possibility of the modification of mordenites has also been rep~rted.~ In this work we deal with Y zeolites with Si/Al ratios in the range 3-20, taking into account the distribution of aluminium between the lattice and extralattice positions, the electron-acceptor properties and the types and acidity of hydroxyl groups.[This subject was considered very briefly in ref. (6).] The catalytic activity is demonstrated by the data obtained from the temperature-programmed desorption (t.p.d.) of ethanol. t Present address: Institute of Chemical Technology, Department of Mineralogy, Suchbatarova 5, 166 28 Prague 6, Czechoslovakia. 13671368 PROPERTIES OF Y-TYPE ZEOLITES EXPERIMENTAL Dehydrated NaY (provided by the Research Institute for Oil and Hydrocarbon Gases, Bratislava) was dealuminated by SiCl, according to the procedure described in ref.(1) at reaction temperatures of 620-770 K for a period of 1-2 h. After purging with dry nitrogen the product was washed until chloride ions were no longer present in the wash water (AgNO, test) and was dried at 320 K. Letters E-B (in order of decreasing A1 content) were used to designate the dealuminated zeolites (table 1). Some samples were subjected to cationic exchange with a Table 1. Chemical composition and sorption capacity of dealuminated Y zeolites Si/A1 100 Na/A1 capacity zeolite (molar ratio) (molar ratio) /mmol (Ar) g-la NaY NH,Y Y-E Y-E/H Y-D Y-D/H Y-c Y-C/H Y-B Y-B/H Y-Ab 2.55 2.50 3.45 3.40 4.9 5.1 5.7 5.75 19.9 19.1 43.4b 99 27 12 7 6 3 5 2 13 7 - 10.9 - 9.2 9.3 9.2 9.4 9.4 - - - a Per g of dry sample; obtained from H.K. Beyer. 0.5 mol dm-3 NH,NO, solution at 350 K. The usual designation NH,Y has been used for the initial Y zeolite (HY after deammoniation and dehydration) and the terms E/H to B/H for the dealuminated zeolites. Term A is reserved for the sample kindly provided by H. K. Beyer (table 1). Prior to measurements of sorption capacity, infrared spectra of OH groups and of adsorbed pyridine and temperature-programmed desorption of ethanol, the samples were dehydrated and deammoniated at 670 K and lo-, Pa. The sample composition was obtained by chemical analysis, and the sorption capacity was determined by adsorption of Ar at 78 K and 13.3 kPa. X-ray powder patterns were measured on a Geigerflex difractometer, and the unit-cell edge dimension was obtained by averaging the values calculated from the individual diffractions in the range t9 = 245".The mid-infrared (i.r.) spectra of zeolites contained in KBr pellets (0.1 wt %) were recorded on a Nicolet MX-1E Fourier-transform infrared spectrometer. The i.r. spectra of the hydroxyl groups and adsorbed pyridine were measured on a Beckman IR-7 spectrometer using zeolite plates of 5-8 mg cm-2 thickness. Interaction with pyridine was carried out for 30 min at 430 K and pyridine pressure of 0.56 kPa. Then the weakly adsorbed species were removed by evacuation at 520 K for 45 min and the spectrum was recorded after cooling to ambient temperature. The numbers of strong proton-donor and electron-acceptor sites were evaluated from the heights of the 1545 and 1455 cm-l bands of the pyridinium ions and pyridine-aluminium complexes, respectively, with the assumption that the extinction coefficients remain constant. The values used,' cpyH+ = 9.7 x cm2 molecule-l [in good agreement with value reported in ref.(S)] and cpyAl = 30.5 x cm2 molecule-l, were determined by the adsorption of successive doses of pyridine on H,,Na,Y zeolite in its hydroxylated and dehydroxylated forms. The conversion of ethanol was followed in the t.p.d. experiments. 10 pmol of alcohol were adsorbed at ambient temperature onto 0.1 g of zeolite. After adsorption for 1 h the sample wasL. KUBELKOV~, v. SEIDL, J. NOVAKOVA, s. BEDNA~OVA AND P. J ~ R ~ J 1369 heated at a rate of 6 K min-'. The maximum pressure of the desorbed vapours did not exceed 10-1 Pa.The desorbates were removed by the vacuum system of an M 1305 mass spectrometer which was used for their analysis and by auxiliary diffusion and ion vacuum pumps. RESULTS AND DISCUSSION CHEMICAL ANALYSIS AND SORPTION CAPACITY The results of chemical analysis (table 1) demonstrate that the reaction of NaY zeolite with SiCl, and subsequent washing with water are accompanied by a decrease in the mole fraction of aluminium [Al/(Si+Al)] from 28% in NaY to 4% in the dealuminated sample Y-B. This method can thus be successfully used for the preparation of zeolites of both high' and intermediate silica contents. It also follows from table 1 that the Na/Al ratio in the initial zeolite approaches unity, in agreement with the fact that aluminium in synthetic Y zeolites is localized in the lattice and sodium ions compensate the lattice charge.In dealuminated zeolites this ratio is about one order of magnitude lower, and thus sodium does not play an important role in charge compensation. The low ratio shows that the zeolites were obtained directly in the acid form, which was confirmed by the i.r. spectra of hydroxyl groups in the dehydrated samples. Consequently, NH, exchange did not increase the acidity substantially (table 1). The sorption-capacity values correspond to good crystallinity of the dealuminated samples (table 1). A decrease in this value relative to the initial zeolite corresponds to data published for stabilized zeolite^.^^ lo An even lower value was cited in ref.(1 1) for a highly dealuminated Y zeolite prepared by combined hydrothermal treatment and acid leaching E8.6 mmol (Ar) g-l]; sample A, provided by H. K. Beyer, yielded a value comparable with that obtained for our samples (table 1). Although the decrease in sorption capacity can be explained by lattice contraction and filling of the cavities by Al-containing species, the presence of a small amount of amorphous phase cannot be excluded for any of the dealuminated zeolites. X-RAY DIFFRACTION ANALYSIS AND I.R. SPECTRA OF SKELETAL VIBRATIONS X-ray powder patterns with very sharp and well defined peaks confirmed the relatively perfect structural arrangement of the dealuminated Y zeolites. All the diffraction peaks of NaY zeolite were found in the powder patterns of the dealuminated zeolites, with shifts in position corresponding to the respective changes in unit-cell dimensions and with changes in relative intensity in character with the literature data.l Similarly, the i.r.spectra of the skeletal vibrations contained the same type of bands as Nay, as can be seen from fig. 1. It is known that substitution of aluminium in the zeolite lattice by silicon leads to a contraction of the lattice and to a change in the bond orders, appearing as a shift of the bands of most of the skeletal vibrations to higher wavenumbers. A linear dependence has been foundl2?l3 between the position of a given band and the mole fraction of aluminium in the lattice for the faujasite-type synthetic zeolites and Y zeolites dealuminated by chelates. Recently this has been confirmed for stabilized zeolites and for hydrothermally dealuminated Y zeolites through a comparison of the values of the lattice Si/Al obtained from i.r.data and from the results of magic- angle-spinning nuclear magnetic resonance using 29Si.147 l5 In fig. 2 the position of the band arising from the internal asymmetric stretching vibration of the lattice tetrahedral2 is plotted against the mole fraction of aluminium determined by chemical1370 PROPERTIES OF Y -TYPE ZEOLITES 10.3 I I 1 I I L 1100 9 00 700 500 wavenum ber/cm Fig. 1. Mid-i.r. spectra of the zeolites NaY (-) and dealuminated Y-C (---) and Y-B ( . . .). analysis, for X and Y zeolites obtained by direct synthesis and for dealuminated Y zeolites. A reasonable linear relationship is found (fig.2) for data from synthesized zeolites (with Si and A1 in the lattice) and from highly dealuminated Y zeolites, while the data of samples Y-C to Y-E exhibit a marked deviation. This fact can be explained by the suggestion that some A1 in the moderately dealuminated zeolites is located in the extralattice positions. The ratios of the number of Si atoms to the number of A1 atoms in the skeleton (Si/Als), which correspond to the respective band positions according to the above relationship, are listed in table 2. These values indicate that only 70-60% of the total amount of A1 is located in the skeleton of moderately dealuminated zeolites. The same Si/Al, ratios are retained even after NH,+ exchange. This conclusion is further supported by an analysis of the dependence of the unit-cell dimension (a) on A1 content.Breck16 published the following relationship for synthetic zeolites of the faujasite type with an Si/Al ratio of 1-3: a = 0.000868NA,+2.419~; NA1 = 192/(1 +Si/Al,). (1) Application of this relationship to the system studied here is thus an approximation; nonetheless, the experimentally determined unit-cell dimensions of moderately dealu- minated zeolites are systematically appreciably lower than the values calculated from eqn (1) using Si/Al ratios found by chemical analysis (table 2). Consequently, the value of Si/Al, calculated from the experimental values of the unit-cell dimensions is in reasonable agreement with the ratios obtained from i.r. data (table 2). Note that no substantial differences were found either in the positions of the i.r.bands or in the values of the unit-cell dimension of the highly dealuminated zeolites Y-B (Si/Al = 20) and Y-A (Si/Al = 43.5) (see fig. 2 and table 2). The aluminium content is apparently so low that the applicability of the above relationships is dubious.L. KUBELKOVA, V. SEIDL, J. NOVAKOVA, S. BEDNhfiOVA AND P. J i R e 1371 h ? \ ' 0 '\\ \ \ \ \ 0 1 2 3 lOAI/(Si + Al) Fig. 2. Dependence of the position (a), height (b) and product of height and half-bandwidth (c) of the asymmetric stretching vibration band of lattice tetrahedra on the mole fraction of A1 in the following zeolites: 0, synthesized X and Y; 0, NaY and dealuminated Y-E, Y-D, Y-C and Y-B; A, dealuminated Y-A. Table 2. Unit-cell dimensions (a), Si/Al ratios obtained from chemical analysis and Si/Al, ratios obtained from X-ray diffraction and i.r. data %/A1 a/nm Si/Al, (chemical sample analysis) exptl theor.X-ray i.r. NaY 2.5 2.469 1 2.4667 2.35 2.45 Y -E 3.4 2.4456 2.4570 5.3 5.2 Y-D 4.9 2.441 1 2.4473 6.6 7.2 Y-c 5.7 2.4323 2.4440 11.6 9.9 Y-B 19.9 2.4264 2.427 1 21.8 - Y-A" 43. 5" 2.427 1 2.4248 - - a Obtained from H. K. Beyer.1372 PROPERTIES OF Y-TYPE ZEOLITES Fig. 2 also demonstrates that the intensity of the band produced by internal asymmetric vibrations of the lattice tetrahedra increases with decreasing aluminium content. Simultaneously, the half-bandwidth, Av;, decreases ; nonetheless, the product of these two values also increases. These values depend on a change in the dipole moment and thus reflect the ionic content of the skeletal bonds.The trend found is in complete agreement with the results of theoretical calculation^,^^ according to which exchange of aluminium for silicon in the skeleton increases the ionic character of the bonds. HYDROXYL GROUPS The individual types of hydroxyl groups and their proton-donor properties were characterized using i.r. spectra and their interaction with pyridine. Fig. 3 depicts the spectra of hydroxyl groups in the HY zeolite and in the dealuminated Y-E/H to Y-B/H zeolites after dehydration and deammoniation at 670 K (the spectrum of sample Y-A is given for comparison). The participation of OH groups in the formation of stable pyH+ ions is apparent from fig. 4, where changes in the spectra of the HY zeolite and the dealuminated Y-C/H zeolite are shown.The numbers of strong proton-donor and electron-acceptor sites, calculated from the heights of the pyridine-species bands, are compared in table 3 with the numbers of lattice and extralattice aluminium atoms determined from the Si/Al ratio obtained by chemical analysis and from the lattice Si/Al, ratio derived from mid-i.r. data. These results indicate that dealumination has a marked effect on the properties of the Y zeolite and that several types of OH groups appear in dealuminated zeolites. (a) Strongly acidic hydroxyls with bands at 3630 and 3560 cm-l, analogous to the structural OH groups in the HY zeolite (bands at 3645 and 3550 cm-l). These were present in considerably smaller amounts in moderately dealuminated zeolites than in HY zeolite; highly dealuminated samples contained very few of these groups.In the HY zeolite only hydroxyls in large cavities formed stable pyH+ species, while all the structural OH groups of dealuminated zeolites took part in the formation of these species (fig. 4). Nevertheless, the number of stable pyridinum ions was always lower than the number of lattice aluminium atoms not compensated by Na+ cations (table 3). (b) Dealuminated zeolites contain a new type of OH group characterized by a band at ca. 3620 cm-l. The number of these groups decreased with dealumination (fig. 3). They were non- or only weakly acidic or were inaccessible to the pyridine molecule (fig. 4). It can thus be assumed that the appearance of these hydroxyls is closely related to the presence of extralattice aluminium.(c) Non-acidic SiOH hydroxyls with a band at 374Ocm-l were found in all the zeolites in amounts increasing with dealumination (fig. 3). Highly dealuminated zeolites contained a large number of these groups apparently as a result of the formation of structural defects and the presence of an amorphous SiO, or AlSiO phase. Some of them may exhibit acidic properties (fig. 4), probably evoked by aluminium. EXTRALATTICE ALUMINIUM It has been shown recently8~1a-20 that firmly bound pyridine complexes can be formed with both extralattice aluminium species (stabilized Y and AlHY zeolites) and lattice aluminium atoms after the release of the structural hydroxyls by dehydroxyl- ation. The structural hydroxyls of the HY zeolite are, however, stable at a temperature of 670 K: dehydroxylation hardly occurs, so that the zeolite contains only a negligible number of strong electron-acceptor sites (table 3).It has been found that dealuminated zeolites exhibit even greater thermal stability; thus the large number of strong electron-acceptor sites found in moderately dealuminated samples apparentlyL. KUBELKOVA, v . SEIDL, J. NOVAKOVA, s. BEDNAROVA AND P. ~ i ~ f i 1373 I 1 I I I _....... . . . . . . ..- .... ..*...... .....* ........ .' 6 . - . . . . . 1 I I I I 3500 3700 wavenurnberlcm -' Fig. 3. 1.r. spectra of the OH groups of (1) HY, (2) Y-E/H, (3) Y-D/H, (4) Y-C/H, (5) Y-B/H and (6) Y-A. provides evidence of extralattice aluminium (table 3). As the adsorption of pyridine measures the coordinatively unsaturated and accessible aluminium, the number of such sites may be lower than the total number of extralattice aluminium atoms (table 3).The formation of pyAl complexes in highly dealuminated Y-B/H (Si/Al = 20) and Y-A (Si/Al = 43.5) samples confirms the presence of extralattice aluminium species in these substances, in agreement with suggestion of cationic A1 implied from 29Si and 27Al n.m.r. studies., The appearance of extralattice aluminium is most probably caused by hydrolysis of chloride complexes that were not removed from the zeolite during its reaction with SiCl, and subsequent purging with dry nitrogen. The reaction with SiCl, can be formally described by the equation1 Na,(AlO,), (Si02)y + SiCl, +NaCl + AlC1, + Na,-l (A102)z-1 (Si02)u+1.PROPERTIES OF Y-TYPE ZEOLITES I I I I 1 I 1 3500 3600 3700 I I I 3500 3600 3700 wavenumber/cm -' Fig.4. 1.r. spectra of the OH groups of (a) HY and (b) Y-C/H after activation at 670 K in VQCUO (---), after interaction with pyridine at 430 K and evacuation at 520 K (---) and after evacuation at 680 K (. . . .). Table 3. Number of aluminium species (At, total, A,, skeletal and Aex, extralattice) and number of strong proton-donor sites (H), electron-acceptor sites (L) and Na ions ( N ) in the HY zeolite and dealuminated Y zeolites HY 18.8 19.3 0 7.9 0.15 5.1 Y-E/H 17.0 11.1 5.9 4.2 2.2 1.3 Y-D/H 12.4 8.8 3.6 5.4 1.9 0.4 Y-C/H 11.3 6.5 4.8 3.8 2.0 0.2 Y-B/H 3.9 - 0.9 0.75 0.3 Y -A 2.0 - - - - 0.2d 0.25 a Chemical analysis; from mid-i.r. data; from the heights of the 1455 and 1545 cm-l bands; from the height of the 1488 cm-l band.During washing with water, hydroxide complexes of aluminium are formed and transferred into solution, depending on the acidity of the suspension. From this point of view, different pH values of the suspension seemingly account for the absence of a decrease in the extralattice aluminum content with dealumination in our moderately dealuminated zeolites (table 3); for Si/Al = 3.4-5.7 the pH of the first suspension had values of 4.8, 2.5 and 2.8 for zeolites Y-E, Y-D and Y-C, respectively. The X-ray photoelectron spectroscopy data showed that the surface of dealuminated zeolites was enriched in aluminium; however, the surface Si/Al ratio did not exceed twice the bulk value. Extralattice aluminium is thus a bulk phenomenon.Because ofL. KUBELKOVP;, v. SEIDL, J. NOVAKOVA, s. BEDNA~~OVP; AND P. J ~ R ~ J 1375 A B/H C/HD/H E/H HY dealuminatedi J- b . 1 .1 .1 Y W c .C g -E 0.1 0.2 0.3 Al/(Si + Al) Fig. 5. Amounts of ethanol (a) and ethylene (b) released during the temperature-programmed desorption of preadsorbed ethanol on HY and dealuminated Y zeolites plotted as a function of Al/(Si + Al). the relatively low number of strongly acidic hydroxyls, this aluminium apparently helps to compensate the lattice charge by being present as cations or oxide clusters in the zeolite cavities; it may also be present in the amorphous phase. It is important from a catalytic point of view that this aluminium is a source of strong electron-acceptor sites. REACTIVITY The above results indicate that the overall amount of aluminium and its distribution affect both the electron-acceptor and proton-donor properties of the dealuminated zeolites, which are very different from those of the initial HY zeolite.Their influence on the catalytic activity is demonstrated by the transformation of preadsorbed ethanol. Fig. 5 shows the amount of unreacted alcohol and of ethylene, the main reaction product, as a function of the ratio Al/(Si+Al) for the zeolites studied. The amounts of desorption products were obtained from t.p.d. measurements by integrating the areas below the respective curves. Unreacted ethanol was released in the temperature range 250-470 K, ethylene at 470-570 K. From these data it follows that moderately dealuminated zeolites with an appreciable number of strong proton-donor and electron-acceptor sites have the greatest efficiency in ethanol transformation.The HY zeolite, with a large number of proton-donor sites and a negligible number of strong electron-acceptor sites, exhibits much lower activity, in a similar manner to the highly dealuminated zeolites with a low number of both types of site. The amount of extralattice A1 may thus be considered an important factor influencing the activity of the zeolites studied.1376 PROPERTIES OF Y-TYPE ZEOLITES CONCLUSIONS The reaction of NaY zeolite with SiCl, at 620-770 K followed by washing with water yielded Y zeolites exhibiting Si/A1 ratios of 3.4-20 and good crystallinity. Some of the aluminium removed from the lattice remains in extralattice positions in the zeolite cavities where it helps to compensate the skeletal charge.Aluminium may also be contained in the amorphous portion of the system. This phenomenon is apparently a result of hydrolysis of the chloride complexes of aluminium left in the zeolite after reaction with SiCl, and purging with nitrogen. The extralattice aluminium may be in either the cation or oxide form and acts as a strong electron acceptor. The number of A1 atoms and their distribution also affect the number and type of OH groups. In addition to non-acidic SiOH hydroxyls (the band at 3740 cm-l) and hydroxyls probably related to the presence of the extralattice A1 (the band at ca. 3620 cm-l), dealuminated zeolites contain strong proton-donor sites : structural hydroxyls corresponding to bands at 3630 and 3560 cm-l are similar to those present in the HY zeolite.Compared with stabilized Y zeolites, dealuminated zeolites exhibit other acid properties and different hydroxyl compositions, and thus constitute a related system with new properties. Moderately dealuminated zeolites were most active in the interaction with ethanol, resulting in the formation mainly of ethylene; these zeolites contain large numbers of both extralattice aluminium atoms and strong proton-donor sites. We thank Dr H. K. Beyer for providing the highly dealuminated Y-A zeolite and for stimulating discussions. H. K. Beyer and I. Belenykaya, in Catalysis by Zeolites, ed. B. Imelik, C. Naccache, Y. Ben Taarit, J. C. Vedrine, G. Coudurier and H. Praliaud (Elsevier, Amsterdam, 1980), p.203. J. M. Thomas, G. R. Millward, R. Ramdas, L. A. Bursill and M. Audier, Faraday Discuss. Chem. SOC., 1982, 72, 345. J. Klinowski, J. M. Thomas, C. A. Fyfe, G. C. Gobbi and J. S. Hartman, Znorg. Chem., 1983,22,63. P. A. Jacobs, J. A. Martens, J. Weitkamp and H. K. Beyer, Faraday Discuss. Chem. Soc., 1982, 72, 353. J. Klinowski, J. M. Thomas, M. W. Anderson, C. A. Fyfe and C. G. Gobbi, Zeolites, 1983, 3, 5. P. Jhfi, in Metal Microstructures in Zeolites, ed. P. A. Jacobs, N. I. Jaeger, P. J i f i and G. Schulz-Ekloff (Elsevier, Amsterdam, 1982), p. 137. B. Wichterlova, J. Novakova, L. Kubelkova and P. Jifi, Proc. 4th Znt. Con5 on Zeolites, ed. L. V. C. Rees (Heyden, London, 1980), p. 373. J. Datka, J. Chem. Soc., Faraday Trans. 1, 1980.76, 2437. Z. TvarEkova, V. Patzelova and V. BosaEek, React. Kinet. Catal. Lett., 1977, 6, 433. lo J. Novakova, L. Kubelkova and B. Wichterlova, Coll. Czech. Chem. Commun., 1980, 45, 2143. l1 V. Bosatek, V. Patzelova, Z. Tvarfiikova, D. Freude, U. Lohse, W. Schirmer, H. Stach and l2 E. M. Flanigen, ACS Monogr., 1976, 171, 80. l3 P. Pichat, R. Beaumont and D. J. Barthomeuf, J . Chem. SOC., Faraday Trans. 1, 1974, 70, 1402. l4 G. Engelhardt, U. Lohse, A. Samoson, M. Magi, M. Tarmak and E. Lippmaa, Zeolites, 1982, 2, 59. l5 G. Engelhardt, U. Lohse, V. Patzelova, M. Magi and E. Lippmaa, Zeolites, 1983, 3, 233. l6 D. W. Breck, Zeolite Molecular Sieves (Russ. Trans., Mir, Moscow, 1976), p. 100. l7 S. Beran, and J. Dubsky, J. Phys. Chem., 1979, 83, 321. l9 K. M. Wang and J. H. Lunsford, J . Catal., 1972, 24, 262. 2o V. BosaEek, J. Brechlerova and M. Kfvanek, in Adsorption of Hydrocarbons on Microporous Adsorb- ents ZZ, Preprints of the Workshop, ed. W. Schirmer and H. Stach (Academy of Sciences of the GDR, Berlin, 1982) vol. 2, p. 26. H. Thamm, J. Catal., 1980, 61, 435. J. W. Ward, ACS Munugr., 1976, 171, 118. (PAPER 3/502)

ISSN:0300-9599    

DOI:10.1039/F19848001367

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. 1, 1984,80, 1377-1389 Pulse Radiolysis of p-Hydroxycinnamic Acid in Aqueous Solution BY KRZYSZTOF BOBROWSKI? Radiation Laboratory, University of Notre Dame, Notre Dame, Indiana 46556, U.S.A. Received 20th April, 1983 Using pulse radiolysis with optical detection it is shown that the hydroxyl radical OH reacts with p-hydroxycinnamic acid in acidic and near-neutral solutions predominantly by addition to the unsaturated substituent, forming benzyl-type radicals. Above pH 9 the spectrum obtained is identical to those obtained from the reaction of p-hydroxycinnamic acid with 0-, Cl;, Br; and CH,CHO radicals. It is suggested that OH radicals attack the ionized hydroxyl group under alkaline conditions to give phenoxyl derivatives. In addition, the spectral data presented show a significant influence of the unsaturated substituent on the position of the absorption maxima in the case of phenoxyl-type radicals, i.e.A = 545 and 595 nm with extinction coefficients 1500 and 1800 m2 mol-l, respectively. The hydroxyl radical is known to react with substituted benzenes predominantly by addition to the yielding hydroxycyclohexadienyl radicals, and not by interaction with the substituent. However, little information is available in the literature concerning the reactivity of OH radicals with substituted benzenes containing unsaturated substituents. In this case two possible reaction pathways can occur, i.e. addition to the aromatic ring and/or addition to the double bond. Earlier studies10-14 concerned with the pulse radiolysis of styrene and methylstyrene confirmed both possibilities for the sites of addition but were in disagreement as to the pattern of the distribution. Our previous paper15 has shown a strong preference of the OH radical for attack at the double bond in the substituent in the case of cinnamic acid.Addition of OH radicals to aromatic rings carrying an additional hydroxy group is of particular interest, because of the possibility of elimination of water from the OH adduct to form phenoxyl-type radicals. Such uncatalysed, acid-catalysed and base-catalysed water- elimination reactions have been reported for phenols,ls methylated benzenes,17 and p-dihydroxybenzenes,21 3,4-dihydro~ytoluene,~~ adrenaline,lg9 2o 4-t-butyl- 1,2- dihydr~xybenzene,~~ 4-t-butyl- 1,2-q~inone,~~ adrenal~ne~~ and p-hydroxyphenyl- propionic a ~ i d .~ ~ p ~ ~ The yield of forming phenoxyl radicals can serve as a measure of the OH radical attack on the ring position. The purpose of this investigation was to obtain additional information about the relative extent of ring/double-bond OH addition, and p-hydroxycinnamic acid was chosen as an appropriate compound. EXPERIMENTAL The solutions were irradiated by 5 ns electron pulses from an ARC0 LP-7 linear accelerator. The dose per pulse was such as to produce only 1-3 pmol dm-3 of radicals in order to minimize their second-order decay. Dose effects were measured with pulses of up to 50 ns duration. Optical detection and signal averaging were carried using the computer-controlled pulse- ? On leave of absence (1979-1981) from the Institute of Nuclear Research, 03-195 Warsaw, Poland.13771378 PULSE RADIOLYSIS OF p-HYDROXYCINNAMIC ACID A/nm Fig. 1. Absorption spectrum (corrected for bleaching of the parent compound, dashed line) of the intermediates obtained after pulse radiolysis of aqueous solutions of p-hydroxycinnamic acid mol dm-3), saturated with N,O at pH 6.0 (immediately after the pulse). radiolysis apparatus described previo~sly.~~ The change in absorbance (AOD) following electron irradiation was followed as a function of time by kinetic spectrophotometry and expressed in terms of an extinction-coefficient parameter (e’) given by the relationship E’ = (AOD x K)/(G x dose), where K is a multiplying factor so chosen that E’ for (SCN), in the same cell is 7600 dm3 mol-l cm-l at 475 nm in N,O-saturated aqueous solutions and G is the yield per 100eV.28 Digitized absorbance data were recorded on magnetic tape and kinetic analyses were made by computer fitting of the curves.2v p-Hydroxycinnamic acid (predominantly in the transform) was obtained from Aldrich and was of the highest purity commercially available.Sodium phosphates (monobasic and dibasic), sodium hydroxide and t-butyl alcohol were analytical grade from Mallinckrodt . Sodium chloride and sodium bromide were from Alfa Division. Ethylene glycol and perchloric acid were obtained from Fischer. In order to minimize thermal oxidation (phenols are extremely sensitive to oxygen in alkaline solutions) all solutions were prepared in the following manner and were not allowed to come into contact with air.First the water was saturated with N,O for 30min and then the appropriate amount ofp-hydroxycinnamic acid was added. During further saturation the KOH was introduced rapidly. Fresh solutions were prepared before each irradiation. The N,O was used in order to remove oxygen and to convert eiq into OH by the reaction N,O+e,, -+N,+OH+OH-. t-Butyl alcohol (0.5 mol dmV3) was used as OH scavenger. The pH of the solution was adjusted with NaOH, HClO,, Na,HPO, and NaH,PO,. Optical absorption spectra of the solutions before irradiation were recorded on a Cary 219 spectrophotometer. The absorption spectra were corrected for the depletion of solute (S) in the wavelength range where the solute absorbed lightK. BOBROWSKI I""'"'''''''''' 4 - ai 5 3- -2 51.D ai 0 OH HC=CH-C< @ +OH OH \ 1379 400 500 600 X/nm Fig. 2. Absorption spectrum (corrected for bleaching of the parent compound, dashed line) of the intermediates obtained after pulse radiolysis of aqueous solutions of p-hydroxycinnamic acid mol dm-3) saturated with N,O at pH 3.0 (immediately after the pulse). (usually < 350 nm). G(p-OHCA),,,,,, was assumed to be 6.0. Owing to the existence of two functional groups (OH and COOH) which can lose a proton, p-hydroxycinnamic acid exists in three forms depending on pH of the solution. A more detailed examination of the optical absorption spectra as a function of pH showed three distinctive regions with pK values pK, = 4.0 and pK, = 9.1 which are in agreement with literature values (Suom.Kemistil. B, 1965, 38, 291). RESULTS REACTION OF OH RADICALS WITH p-HYDROXYCINNAMIC ACID A pulse-radiolysis study of nitrous oxide-saturated aqueous solutions of p- hydroxycinnamic acid (p-OCHA) at pH 3.0-11.5 was carried out using optical absorption for the detection of transients. Under these conditions ca. 90% of the primary radicals are available as OH radicals for reaction with p-hydroxycinnamic acid. The transient absorption spectra obtained at various pH values are given in fig. 1-3. In almost neutral solutions (pH 6.0) a strong sharp absorption band was observed at 0.7 ps after the pulse with an absorption maximum at 3, = 335 nm together with broad shoulder with maximum at ca. 405 nm (fig. 1). The rate of the formation of the transient at 335 nm followed first-order kinetics.From the fitting of the exponential increase of the 335nm absorption after the pulse and the concentration of p- hydroxycinnamic acid (in the range 2 x mol dm-3) a rate constant k = 8.2 x lo9 dm3 mol-l s-l was calculated for the reaction. This value is almost equal -2 xI380 PULSE RADIOLYSIS OF p-HYDROXYCINNAMIC ACID 16 14 12 i 10 e 0- L l r l r l l l l l r l r l r r l l 400 50 0 600 h/nm Fig. 3. Absorption spectrum (corrected for bleaching of the parent compound) of the intermediates obtained after pulse radiolysis of aqueous solutions of p-hydroxycinnamic acid mol dm-3) saturated with N20 at pH 10.0 (immediately after the pulse). l 4 t I I 2 [H+l/10-5 mol dmV3 Fig. 4. Plot of the observed rate constant kexptl of the fast process against H+ concentration.Inset: pH dependence of the absorption of the intermediate taken 10 ps after the pdse at A = 345 nm in aqueous solution of p-hydroxycinnamic acid saturated with N20.K. BOBROWSKI 1381 to the value obtained for cinnamic acid.15 Similar results were observed at pH 3.0 with a possible absorption maximum located at (or below) 3, = 340 nm (fig. 2). The spectral region below 340 nm becomes inaccessible because the monitoring light is absorbed by the parent compound itself. A broad band was observed in the spectrum at 425 nm together with weak absorption bands at A = 545 and 595 nm. On the other hand transient absorption spectra taken 0.7 ,us after the pulse in alkaline solution (pH 10.0) show broad bands at 470 nm and below 380 nm; again the region below 380 nm was inaccessible because of absorption by the parent compound itself (fig.3). Instead of the very weak absorption bands observed in acid and neutral solution, two intense sharp maxima were observed at 545 and 595 nm, respectively. As one can see, the reaction of OH radicals with p-hydroxycinnamic acid produces absorption spectra which differ in the positions of their absorption maxima as well as in their intensities. Thus a more detailed examination of the effect of varying the pH on the spectra was carried out. The transient absorption at A = 345 nm is shown as a function of pH (inset) in fig. 4. The pK value of the transient observed at this wavelength is ca. 4.9 + O . 1, which is higher than the value of pKl ofp-hydroxycinnamic acid.As in the case of cinnamic acid,15 in a very narrow region of pH (4.4-6.0) the absorption spectrum ofp-hydroxycinnamic acid undergoes a very similar small change during the first 5 ,us after the pulse. In the region 330-345 nm the transient absorption decays fast in a first-order reaction which is independent of dose and solute concentration but slightly dependent on H+ concentration. The experimental rate constants were plotted as a function of [H+] and a straight line was obtained (see fig. 4). Such H+ dependence strongly suggests an equilibrium of the form OH 1 8 HC-CH-c OH I 40 H C-CH-C I OH OH The experimental rate constant kexptl for a reaction of this type is, of course, equal to k,+k,[H+]. From the plot presented in fig. 4 rate constants of 3.8 x 1O1O dm3 mol-1 s-l and 2.9 x lo5 s-l were determined for kf and k,, from which K = k,/k, = 1.33 x mol dm-3 was calculated corresponding to a pK, of 4.9, in excellent agreement with value estimated from the dependence of absorption on pH at A = 345 nm.REACTION OF 0- RADICAL IONS WITH p-HYDROXYCINNAMIC ACID The radical ion 0- was produced by irradiating 1 mol dm-3 NaOH solutions saturated with N,O. Under these conditions essentially all primary radicals (99 %) were converted to 0-. The transient absorption spectrum obtained with irradiated solutions of p-hydroxycinnamic acid (1 x mol dm-3) is presented in fig. 5. The spectrum recorded at pH 14 indicates some contribution of the OH adduct at ca. 470 nm. Sharp maxima at 545 and 595 nm are also present and are similar to those previously observed in solutions of p-hydroxycinnamic acid at pH 10.With phenols the 0- radicals have been ~uggested~O7~l to react by direct electron transfer, and the predominant product is phenoxyl radical. The reaction of 0- with a phenoxide-type ion (-0-C,H,-CH=CH-COO-) was followed directly by examining the buildup of1382 PULSE RADIOLYSIS OF p-HYDROXYCINNAMIC ACID I6 - I4 - 12 - W 10- e 2 D ; 8 - .- * id W - 6 - 4 - 2- L I 1 1 1 1 1 1 " 1 1 " 1 1 I I 400 500 600 h/nm Fig. 5. Absorption spectrum (corrected for bleaching of the parent compound, dashed line) of the intermediates obtained after pulse radiolysis of aqueous solutions of p-hydroxycinnamic acid ( mol drnp3) saturated with N,O at pH 14.0 (40 p s after the pulse). transient absorption at A = 595 nm.The rate constant determined from the reaction periods observed over the concentration range 5 x 10-4-5 x mol dm-3 was 3.1 x lo8 dm3 mol-l s-l (after correction for the reaction of OH with p-hydroxycin- namic acid). This rate constant was found to be one order of magnitude greater than that expected for addition of 0- to the aromatic ring.31 These facts suggest that this reaction also occurs through transfer of an electron from the phenoxide anion to 0-. However, the spectrum obtained is quite different from the spectra of phenoxyl-type radicals; usually they absorb in the region 380-420 nml6* 3 0 ~ 32-37 with well defined maxima at ca. 390 and ca. 410 nm. The shape of the observed spectrum resembles that of the phenoxyl radical but is red-shifted by ca.170 nm. REACTIONS OF OTHER OXIDIZING RADICALS (Br;, Cl; AND CH,CHO) WITH p-HYDROXYCINNAMIC ACID It is necessary to determine if the spectrum obtained through the reaction of 0- withp-hydroxycinnamic acid is due to a phenoxyl-type radical or to a different species. It appeared to us that observation of the spectra formed by the reaction of other oxidizing radicals with p-hydroxycinnamic acid could throw some light on the origin of the spectrum. Of the oxidants mentioned above, Br; is the most reactive after C1, and seemed the most convenient to study because it could be produced over a wide pH range.32 However, even in this case the reaction rate with undissociated phenols is too slow, so studies could be made only above the pK, of the phenol. To provide other oxidizing radicals38 studies were also made with C1; (at pH 3) and with theK.BOBROWSKI 1383 I I I I I 400 500 600 X/nm Fig. 6. Absorption spectrum of the intermediates obtained after pulse radiolysis of aqueous solutions containing I x mol dmW3 ofp-hydroxycinnamic acid and 1 mol dm-3 KBr at pH 11.5 and saturated with N,O. Inset: plots of the change in relative absorbance with time monitored at the maximum of the absorption band at 595 and 385 nm (near the maximum of the absorption band of Br;). formylmethyl radical, CH,CH0,39v 40 which has received some recent attention. Use of the SO, radical*l as a direct oxidant of carboxylated phenoxide ions and for the generation of C1; above pH 3 was unsuccessful because of the thermal reaction between sodium persulphate and p-hydroxycinnamic acid.Reaction of Br; with carboxylated phenoxide, -O--C,H,-CH=CH' 0 '0- (pK, 9. lo), was carried out at pH 11.5 in N,O-saturated solutions containing 1 mol dmF3 KBr. At wavelengths > 500 nm the optical spectrum was essentially identical to that found with 0- as a reactant (fig. 6). Below 500 nm there was much less absorption than in the previous cases, since any OH radical formed reacts with Br- (because of the high concentration of the latter and the high rate constant for its reaction with OH). Thus there was no possibility of a contribution by the OH adduct to the benzene ring in such conditions. Clearly direct oxidation of carboxylated phenoxide ions by Br; radicals was observed at A = 595 nm (see inset in fig. 6). As can be seen, the dependences of the two normalized plots of the change in relative absorbance with time monitored at both wavelengths gave symmetrical curves with the same rate constant, k = 9.9 x lo8 dm3 mol-l s-l.Similar experiments were performed at pH 2.9 in N,O-saturated solutions contain- ing 1 mol dm-3 Cl- in order to minimize the competition by p-hydroxycinnamic acid1384 PULSE RADIOLYSIS OF p-HYDROXYCINNAMIC ACID [p-OHCA]/10-3 mol dm-3 Fig. 7. Plot of the observed rate constant kexpt, of the formation of 595 nm absorption against concentration of p-hydroxycinnamic acid in aqueous solutions containing 1 mol dm-3 of ethylene glycol at pH 1 1.5, saturated with N,O. Inset: experimental trace followed at A = 595 nm in aqueous solutions of p-hydroxycinnamic acid (2 x rnol dm-3) at pH 1 1.5 containing 1 rnol dm-3 ethylene glycol and saturated with N,O.for hydroxyl radicals. In this case the only reactive species which can oxidize p-hydroxycinnamic acid is the C1; radical. The intense absorption of C1; below 380 nm was observed immediately after the pulse. Simultaneously with the decay of the C1, radical an absorption spectrum was formed which was essentially identical to those found with Br; and 0- as reactants (at wavelengths > 500 nm). From the plot of the pseudo-first-order rate constants for the decay of Cl, as a function of p-hydroxycinnamic acid concentration, the second-order rate constant for the reaction of C1; with p-hydroxycinnamic acid was derived as 2.9 x loa dm3 mol-1 s-l. the formylmethyl radical can be produced from ethylene glycol via H abstraction As previously HOCH2CH20H + OH --+ HOCH,cHOH + H20 followed by a rapid elimination of water in alkaline solutions: OH- -OH- HOCH,cHOH - HOCH,CHO- - cH2CH0.The resultant radical was found to oxidize various phenols by direct electron t r a n ~ f e r . ~ ~ ~ * ~ In our case the formylmethyl radical produced by the reaction of OH with 1 mol dm-3 ethylene glycol at pH 11.5 was allowed to react with p-hydroxycin- namic acid [(0.5-2) x mol dm-3]. The inset in fig. 7 shows the experimental trace obtained for the p-hydroxycinnamic acid + ethylene glycol system in water monitored at 1 = 595 nm. In this case a build-up of absorbance is observed due to the formation of the oxidation product of p-hydroxycinnamic acid. By using a large excess of ethylene glycol (1 mol dm-3) the direct reaction of OH radicals withp-hydroxycinnamicK.BOBROWSKI 1385 1 I 1 I 400 450 500 h/nm Fig. 8. Absorption spectrum at pH 1 of the hydrogen adduct to p-hydroxycinnamic acid mol dm-3) saturated with N, and containing 0.5 mol dm-3 t-butyl alcohol to remove OH radicals. acid was prevented. A plot of kexpt against p-hydroxycinnamic acid concentration yielded a straight line (fig. 7). From the slope of this line kp.0CHA+CH2CH0 was estimated as 7.7 x lo7 dm3 mol-l s-l. This value is typical of such an electron-transfer reaction (values are usually in the range 107-109 dm3 mol-1 s - ' ) . ~ O The specttal absorption maxima (i.e. 545 and 595 nm) of the transient were found to be the same as those of the transients formed in previous systems.REACTION OF H ATOMS WITH p-HYDROXYCINNAMIC ACID The reaction of H atoms withp-hydroxycinnamic acid was studied after the removal of most of the OH radicals through reaction with t-butyl alcohol at pH 1. The spectrum observed immediately after the pulse has an absorption maximum around 2 = 425 nm (fig. 8). The spectral region below 360 nm becomes inaccessible because of the absorption of the monitoring light by the parent compound itself. DISCUSSION SPECTRUM OF THE PI-ENOXYL-TYPE RADICAL The absorption spectra (fig. 5 and 6) represent transients which are formed by direct electron transfer. The similarity between the optical absorption spectra at il > 500 nm observed with 0-, Br,,Cl, and CH,CHO radicals as oxidants strongly suggests that direct oxidation of the phenoxide ion from p-hydroxycinnamic acid is involved.In each case the spectrum is similar in shape to that of phenoxyl radical but with a strong 46 FAR 11386 PULSE RADIOLYSIS OF P-HYDROXYCINNAMIC ACID 0 4 HC=CH-C 'OH t OH /p '0- HC=CH-C @ + OH It -H* I I+H* 0 4 9 + OH \ 0- +Hi0 0- U - 2 OH- Scheme 1. r e d - ~ h i f t ~ ~ in the position of the two maxima. The extinction coefficients of the phenoxyl-type radical can be determined directly from absorption spectra produced in a solution co,ntaining 1 mol dm-3 Br- and 1 mmol dm-3 ofp-hydroxycinnamic acid. Under these conditions any OH radical should react with Br- because of the high concentration of the latter and the high rate constant for its reaction with OH. As was mentioned earlier, Br; is quite reactive and direct oxidation of phenoxide-type ions by Br; should proceed with the yield G(Br;) = G(0H) = 6.0.The extinction coefficients calculated on the basis of complete reaction between Br; and p- hydroxycinnamic acid are E,,, = 1500 and E,,, = 1800 m2 mo1-l. PROPOSED REACTION MECHANISM The experimental results are explained in terms of three processes: (a) addition of OH to the unsaturated substituent leading to formation of benzyl-type radicals, (b)K. BOBROWSKI 1387 addition of OH(H) to the benzene ring leading to the formation of OH(H) adducts (in the first case these may decay by elimination of water, yielding phenoxyl-type radicals) and (c) direct attack of OH at the ionized phenolic group, giving also phenoxyl derivatives (see scheme 1).The wavelength maxima and approximate relative intensities of the transient bands obtained under different conditions are summarized in table 1. The observed bands Table 1. Transient spectra obtained by pulse radiolysis in aqueous solutions of p-hydroxycinnamic acid (1 0-3 mol dm-3) additive PH transient absorption maxima"/nm none none none with with with with 1 mol dm-3 NaOH 1 mol dm-3 KBr 1 mol dm-3 KCl 1 mol dm-3 ethylene glycol 3.0 < 360b (23) 6.0 330 (102) 10.0 < 380* (14) 14.0 < 380b (10) 1 1.5 absorption due to Br; 2.9 absorption due to C1; - 11.5 42Y (13.5) 41OC (13.5) 47OC (8.0) 47OC ( 10.5) no absorption no absorption no absorption 545 (2.5) 595 (3.0) 545 (2.5) 595 (3.0) 545 (10.0) 595 (12.5) 545 (1 3.0) 595 (14.5) 545 (1 5.0) 595 (18.0) 545 (13.0) 595 (14.5) 595 (13.0) 545 (1 1 .O) a The number in parenthesis indicates the relative band height.Not determinable due to the strong absorption of the analysing light by the parent compound. It is possible that the wavelength region corresponds to both the cyclohexadienyl and hydroxycyclohexadienyl derivatives, but the strong overlapping between OH addition and H addition products makes the estimation of the yield very uncertain. fall into three groups. The strong absorption near 340 nm for p-hydroxycinnamic acid corresponds to the characteristic benzyl-type radicals. In the pH region where both functional groups are protonated the initial attack of hydroxyl radicals is by addition to the unsaturated substituent [reaction (1 b)], and a relatively sharp and strong absorption band in the region 360 nm is assigned to species 16.Very weak absorption in the regions 410-470 and 540-600 nm testifies to a low yield of OH addition to the aromatic ring, if any, followed by acid-catalysed elimination of the water molecule and generation of phenoxyl-type radicals [reaction (1 a)]. In the pH region where only the carboxyl group is deprotonated the spectrum obtained is practically identical to the previous one. This shows that in acidic and near-neutral solutions OH addition to the unsaturated substituent, leading to benzyl-type radicals, is the predominant process between the substrate and the initial radiolysis products of water. The fast decay observed during the first 5ps after the pulse represents (as in the case of cinnamic acid)15 an equilibration process between species I b and I1 6, i.e.by reactions with constants kf and k,. This conclusion is supported by the fact that the decay is dependent on H+ concentration, the agreement in the values of K, obtained from the kinetic treatment of the equilibration process and the dependence of the optical density observed at il = 345 nm as a function of pH. As in the case of fully protonated acid 46-21388 PULSE RADIOLYSIS OF p-HYDROXYCINNAMIC ACID molecule a very low yield of phenoxyl-type radicals is observed at this pH. Thus using values of the extinction coefficient for the phenoxy radical calculated in this work one can estimate that reactions ( l a ) and (2a) account for ca. 15% of the total hydroxyl radical yield.On increasing the pH to ca. 9, i.e. above pK2 (fig. 3), intense absorption was observed at A 2 500 nm which was very weak in acid and neutral solutions. From the similarity of the optical absorption spectra to those observed in systems containing Br;, Cl;, CH,CHO and 0- it is concluded that phenoxyl-type radicals from p- hydroxycinnamic acid are formed. In alkaline solution (at pH lo), in which the hydroxy group in p-hydroxycinnamic acid is ionized, OH radicals may both add to the benzene ring and react with the OH group simultaneously. In this context note that the water-elimination reaction16 via the hydroxycyclohexadienyl radical occurs with uncatalysed elimination rate constants of lo3 s-l. However, such reactions are acid/base catalysed with subsequent elimination rate constant k = lo5 s-l.If reactions (1 a) and ( 3 a) were significant then the yield of phenoxyl-type radicals in the acid and alkaline regions should be very close. The large difference in yields of the phenoxyl-type radical in both acid/neutral and alkaline solutions is taken to mean that the transient phenomena observed in both regions make the occurrence of reactions ( 1 a), (2a) and ( 3 a) followed by the elimination of water very unlikely in this case. Our observations regarding the large increase in the yield of phenoxyl-type radical on going from low (fig. 2) to high pH (fig. 3 ) clearly show that OH radicals attack the ionized hydroxy group under alkaline conditions to give phenoxyl derivatives. Such a reaction was proposed by Chrysoch0os~~9 26 for p-hydroxyphenylpropionic acid, tyrosine and polytyrosine in aqueous media.The wavelength region corresponding to benzyl-type radicals indicates that these are formed to some extent, but the strong absorption of the analysing light below 380 nm by the parent compound makes an estimation of the absorption maximum impossible. However, based on the known extinction for the phenoxyl-type radical and the absorption value of phenoxyl radicals formed at pH 10, the proposed direct electron-transfer reaction between the ionized hydroxy group and hydroxyl radical accounts for ca. 70% of the total hydroxyl radical yield. The for- mation of phenoxyl radicals through step (3a) can also be excluded for the following reason. If step (3a) were to be an important path, one should observe predominantly the formation of the hydroxycyclohexadienyl radical followed by the delayed formation of the phenoxyl radical with simultaneous decay of the hydroxycyclohexadienyl derivative (as was observed in some systems).On the contrary, in our case absorption due to phenoxyl radicals occurs at the same time as very weak absorption in the region 410-470 nm. We therefore conclude that formation of phenoxyl radicals occurs mainly through direct electron transfer between the OH radical and the ionized hydroxy group in the ring and not via a hydroxycyclohexadienyl derivative. CONCLUSIONS Three conclusions may be derived from these results. First, the addition of an unsaturated side chain can cause a large red shift in the absorption spectrum of phenoxyl-type radicals.Secondly, as in the case of cinnnamic acid, addition of hydroxyl radicals to the unsaturated substituent can effectively compete for addition to the benzene ring at both acid and neutral pH. Thirdly, phenoxyl-radical formation from p-hydroxycinnamic acid is negligible unless the phenol is ionized, thus allowing electron transfer to take place.K. BOBROWSKI 1389 Helpful discussions with Dr P. Neta are gratefully acknowledged. The research described herein was supported by the Office of Basic Energy Sciences of the Department of Energy. This is document no. NDRL-2352 from the Notre Dame Radiation Laboratory. The paper was presented in part at the 28th IUPAC Congress held in August 1981 at Vancouver, B.C., Canada. P. Neta and L. M.Dorfman, Adv. Chem. Ser., 1968, 81, 222. C. R. E. Jefcoate and R. 0. C. Norman, J. Chem. SOC. B, 1968, 48. P. Neta, M. Z. Hoffman and M. J. SimiC, J. Phys. Chem., 1972,76, 847. N. V. Raghavan and S. Steenken, J. Am. Chem. SOC., 1980,102, 3495. N. Selvarajan and N. V. Raghavan, J. Phys. Chem., 1980,84, 2548. 0. Volkert, W. Bors and D. Schulte-Frohlinde, Z. Naturforsch., Tie1 B, 1961, 22, 480. P. Neta, Radiat. Res., 1973, 56, 201. S. Steenken and P. O’Neill, J. Phys. Chem., 1978,82, 372. A. Samuni and P. Neta, J. Phys. Chem., 1973,77, 1629. lo A. J. Swallow, Adv. Chem. Ser., 1968, 81, 499. 0. Brede, W. Helmstreit and R. Mehnert, J. Prakt. Chem., 1974, 316, 402. l 2 0. Brede, J. Bos and R. Mehnert, Radiochem. Radioanal. Lett., 1979, 39, 259. l3 0. Brede, J. Bos, W.Helmstreit and R. Mehnert, 2. Chem. (Leipzig), 1977, 17, 447. l4 C. Schneider, Adv. Chem. Ser., 1967, 91, 219. l5 K. Bobrowski and N. V. Raghavan, J. Phys. Chem., 1982,86,4432. l6 E. J. Land and M. Ebert, Trans. Faraday SOC., 1967, 63, 1181. K. Sehested, H. Cofitzen, H. C. Christensen and E. J. Hart, J. Phys. Chem., 1975, 79, 310. W. Bors, M. Saran, C. Michel, E. Lengfelder, C. Fuchs and R. Spottl, Int. J. Radiat. Biol., 1975, 28, 353. l9 M. Gohn, N. Getoff and E. Bjergbakke, Int. J. Radiat. Phys. Chem., 1976, 8, 533. 2o M. Gohn, N. Getoff and E. Bjergbakke, J. Chem. SOC., Faraday Trans. 2, 1977, 73, 406. 21 G. E. Adams and B. D. Michael, Trans. Faraday SOC., 1967, 63, 1 171. 22 M. Gohn and N. Getoff, J . Chem. Soc., Faraday Trans. I , 1977, 73, 1207. 23 H. W. Richter, J. Phys. Chem., 1979, 83, 1123. 24 W. Bors, M. Saran and C. Michel, J. Phys. Chem., 1979,83, 2447. 25 J. Chrysochoos, Radiat. Res., 1968, 33, 465. 26 J. Chrysochoos, J. Phys. Chem., 1969, 73,4188. 27 L. K. Patterson and J. Lillie, Int. J. Radiat. Phys. Chem., 1974, 6, 129. 28 R. H. Schuler, L. K. Patterson and E. Janata, J. Phys. Chem., 1980,84, 2088. 29 R. H. Schuler and G. Buzzard, Int. J. Radiat. Phys. Chem., 1976,8, 563. 30 P. Neta and R. H. Schuler, J. Am. Chem. SOC., 1975, 97, 912. 31 P. Neta and R. H. Schuler, Radiat. Res., 1975, 64, 233. 32 K. M. Bansal and R. W. Fessenden, Radiat. Res., 1976, 67, 1. 33 K. Kemsley, J. S. Moore, G. 0. Phillips and A. Sosnowski, Acta Vitaminol. Enzymol., 1974, 28, 263. 34 J. L. Redpath and R. L. Willson, Int. J. Radiat. Biol., 1975, 27, 389. 35 E. J. Land, G. Porter and E. Strachan, Trans. Faraday SOC., 1961, 57, 1885. 36 H. Joschek and L. I. Grossweiner, J. Am. Chem. SOC., 1966,88, 3261. 37 P. K. Das and S. N. Bhattacharyya, J. Phys. Chem., 1981,85, 1391. 38 K. Hasegawa and P. Neta, J. Phys. Chem., 1978,82, 854. 39 S. Steenken, J. Phys. Chem., 1979,83, 595. 40 S. Steenken and P. Neta, J. Phys. Chem., 1979, 83, 1134. 41 H. Zemel and R. W. Fessenden, J. Phys. Chem., 1978,82, 2670. 4 2 K. M. Bansal, M. Gratzel, A. Henglein and E. Janata, J. Phys. Chem., 1973, 77, 16. 43 S. Steenken and P. Neta, J. Phys. Chem., 1982, 86, 3661. (PAPER 3/634)

ISSN:0300-9599    

DOI:10.1039/F19848001377

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1984,80, 1391-1407 Interaction between Zeolites and Cluster Compounds Part 2.-Thermal Decomposition of Iron Pentacarbonyl on Zeolites BY THOMAS BEIN? Institute fur Physikalische Chemie der Universitat Hamburg, Laufgraben 24, D-2000 Hamburg 13, Federal Republic of Germany AND PETER A. JACOBS* Centrum voor Oppervlaktescheikunde en Colloidale Scheikunde, Katholieke Universiteit Leuven, Kardinaal Mercieplaan 92, B-3030 Leuven (Heverlee), Belgium Received 20th April, 1983 Thermal decomposition in a thermobalance of Fe(CO), adsorbed on alkali-metal, hydrogen-Y, dealuminated Y, L and omega zeolites proceeds stepwise via slow decarbonylation at low and high temperatures, separated by a fast endothermic reaction. Average CO/Fe ratios have been determined after each step.From i.r. results the former intermediates are assigned to species bearing bridging CO, whereas reaction products with CO/Fe < 1 are associated with highly unsaturated carbonyl clusters in strong interaction with the zeolite. The thermal stability of zeolite/Fe(CO), adducts as well as of the intermediates increases with the electron-donor properties of the matrix and can be rationalized using the Sanderson electronegativity concept. Iron loadings ranging from 2.4 wt % in zeolite L up to 10 wt % with NaY and HY are obtained by decomposition in inert atmosphere. Under vacuum conditions loss of metal up to 50% is observed. Metallic iron clusters are the final decomposition products in alkali-metal zeolites, as probed by NO adsorption. In HY part of the metallic iron is oxidized to FeI1 ions, which are located at cation positions.Attempts to prepare a mononodal metal particle-size distribution on zeolites need no .further justification, since these materials may have unusual properties in metal catalysis. This is certainly true for Fe zeolites, which by such classical methods of preparation as hydrogen or sodium vapour reduction of iron@)-exchanged samples show multinodal distributions and/or chemically contaminated metal phases. 1-3 In the present work the thermal decomposition of sorbed iron pentacarbonyl, Fe(CO),, was chosen as an alternative route. Zeolites are suitable supports as a result of their variable defined cage dimensions and cation-exchange capacity, which can be used to adjust the electronic properties.*$ In the first part of this series equilibrium sorption data of Fe(CO), on several zeolites were reported.6 It seems that for room-temperature sorption a critical pore diameter of ca.0.67 nm is needed in order to be able to reach complete pore filling. The sorbate-zeolite matrix interaction, and consequently the packing of iron pentacarbonyl, was found to be determined mainly by the nature of the zeolite charge-compensating cations rather than by the structure of the pores. The gas-phase decomposition of Fe(CO), occurs stepwise. The process is initiated by a slow release of the first CO ligand followed by a fast decomposition of the species t On leave from : Centrum voor Oppervlaktescheikunde en Colloidale Scheikunde, Katholieke Univer- siteit Leuven, De Croylaan 42, B-3030 Leuven (Heverlee), Belgium.13911392 DECOMPOSITION OF Fe(CO), ON ZEOLITES Table 1. Unitcell composition of the zeolites zeolites unit-cell composition FAU Na55.5 "55.5 si136.5 '384 CS-FAU H-FAU "37 Na18.5 A155.5 si136.5 O384 H50 Na5.5 A155..5 si13fJ.5 '384 FAU* Si192 0 3 8 4 MOR Na, Al, Si,, O,, LTL &.8 "8.8 si27.2 '72 MAZ Na6.B "6.8 ' h . 4 '72 formed this way. This is the result of differences in bonding energies of the CO ligands to the iron atom. The thermal decomposition of Fe(CO), impregnated on supports is known to give a dispersed FeO phase,* although the fate of the sorbed carbonyl during thermal decomposition is not known in detail. On thoroughly dehydroxylated MgO, adsorbed Fe(CO), is decomposed uia mononuclear or polynuclear species, the latter being the more thermo~table.~ Iron dodecacarbonyl, Fe,(CO),,, sorbed on the same support partially covered with surface hydroxyl groups, produced a hydridoundeca- carbonyl complex, HFe,(CO),,, which upon thermal treatment is transformed into FeII ions and superparamagnetic iron particles.1° Partial oxidation of iron also seems to occur on other acidic supports.The evolution of molecular hydrogen on y-aluminall and hydrogen-Y zeolitel21 l3 during thermal treatment of the pentacarbonyl-support adduct is considered to be evidence for this reaction. Recently it was established using volumetric methods that the thermal decomposition of Fe(CO), on y-aluminall and of Fe,(CO),, on NaY zeolite14 occurs in a stepwise fashion. The same was found to be true for Fe(CO), adsorbed on NaY or HY zeo1ites,l5 although on the latter support Fe,(CO),, may be one of the early intermediates.N.m.r. methods indicate that the interaction of these intermediates with the hydrogen-Y zeolite depends upon the method used for their formation.ls In order to be able to prepare a mononodal iron metal particle-size distribution it was the aim of the present work to elucidate the thermal-decomposition processes of Fe(CO), on several large-pore zeolites, either with acidic or non-acidic properties. Attention was focussed more particularly on the thermal stability of sorbed Fe(CO),, the nature and reactivity of intermediates and the influence of the zeolite itself on this behaviour. Thermoanalytical techniques were used to follow the decomposition in a quantitative way and in situ i.r.spectroscopy to probe the carbonyl-zeolite interactiom8 EXPERIMENTAL MATERIALS All zeolites used were of synthetic origin. The unit-cell composition of the dry zeolites was determined by atomic absorption spectrometry and is given in table 1. A mnemonic code was used to identify the structure type,17 and the pretreatment procedures are given elsewhere in more detail.6 NaY was from Strem Chemicals (FAU) and TMA-omega from Linde (MAZ). CsY (Cs-FAU) was obtained from NaY by ion exchange. HY (H-FAU) was obtained by thermal decomposition of NH4C1-exchanged NaY at 720 K. Dealuminated Y (FAU*) was prepared from steamed NH4Y (at 820 K), subsequently treated with SiCl, vapour at 820 K and then calcined at 1220 K.Linde L (LTL) and Na-mordenite (MOR) from Norton were used without pretreatment. Before use all zeolites were stored over saturated NH,Cl solution in order to ensure constant humidity.T. BEIN AND P. A. JACOBS 1393 Iron pentacarbonyl from Ventron (99.5%) was cold-distilled in the dark and stored over molecular sieve 5A. METHODS Thermoanalytical d.t.a. and d.t.g. measurements were made in a Mettler Thermoanalyser 2 under helium purge, mostly in the 10-1 mg range, the error of the measurements being < 0.1 mg. Samples of 10 & 2 mg zeolite and the same amount of dried Al,O, as reference were placed in Pt crucibles of 7 mm height and 4 mm diameter. The samples were degassed by heating at a rate of 2 K min-1 up to 720 K, keeping the temperature at this value for 10 h.Loading of the samples with Fe(CO), proceeded at 293 K in a stream of dry helium (2.8 dm3 h-l). The partial pressure of the carbonyl was 0.4 kN m-,. Thermal decomposition was also carried out in a stream of helium at the same flow rate. Before increasing the sample temperature, the samples were purged at room temperature for 1 h. The arrangement of the sample holder in the balance is such that the carrier stream is directed to the top of the crucible containing the sample but does not pass through the bed of the sample, so that sorption or desorption from the sample is governed by diffusion processes. Weight changes of samples loaded with pentacarbonyl were also determined in McBain balances under a dynamic vacuum of lo-, N m-2. After contacting the sample with 0.88 kN m-2 of sorbate for 24 h, the samples were decarbonylated stepwise at different temperatures.Degassing was performed at 333 K for 5 h, at 373 K for 15 h, at 413 K for 22 h and at 493 K for 20 h. The heating rate applied was 0.5 K mind'. The infrared spectra were recorded with a Perkin-Elmer 580 B spectrometer in the 4000-1200 cm-l range. The scan mode used provided a resolution better than 1.5 cm-l. The zeolite powder was pressed to self-supporting wafers of cu. 5 mg cm-, at 0.1 GN m-2. A quartz cell with 80 mm path length and equipped with CaF, windows of 3 mm thickness was placed in the spectrometer and was connected to a greaseless gas dosing and handling system. The zeolite wafers were degassed in this cell at 720 K and lo-, N m-2 for 2 h.They were loaded at 293 K with Fe(CO), vapour as follows. The carbonyl was frozen in liquid air, outgassed and allowed to warm up until the desired pressure was reached. Cu. 10% saturation was reached after contacting the zeolite wafers with 3 N m-2 of Fe(CO), for 10 s, whereas saturation was obtained by contacting the wafers with 0.88 kN rn-, of Fe(CO), for 1 h. The sample loaded with pentacarbonyl was then degassed for 10 min and heated at 10 K min-l after addition of 60 kN m-2 of helium to the cell. RESULTS AND DISCUSSION VACUUM DECOMPOSITION OF IRON PENTACARBONYL-ZEOLITE ADDUCTS Representative weight-loss curves of zeolites, loaded with Fe(CO), at room tem- perature, during thermal treatment in vacuu are given in fig. 1. These data obtained in a McBain balance show that for LTL, MAZ and FAU zeolites the greatest weight loss occurs between 330 and 370 K and the decomposition is complete around 420 K.H-FAU and MOR samples behave in a similar way. The colour of the samples becomes grey or black depending on the degree of iron loading. Upon further heating to 600 K, in no case was a decrease in weight found. This proves that the zeolite heated until 420 K is always CO-free. Fig. 1 also shows that the steady weight obtained for LTL and MAZ corresponds within the experimental error to the weight expected when all the sorbed Fe(CO), is decarbonylated. It can be shown that for FAU zeolites 25% of the iron is sublimated from the sample, while for H-FAU and MOR iron losses of 50% occur. If decomposition is carried out in a continuous-flow reactor using a purge of an inert gas, a similar loss of iron is found.Iron loadings varying between 1 and 8 wt % have been reported using thermal vacuum decomposition of Fe(CO),-HY adducts,l3? l6 which confirms that iron loss is extremely dependent on the decomposition procedure used. Obviously the one-dimensional channel structure of MAZ and LTL zeolites prevents extensive loss of iron. Since the free diameter of the pores in MAZ1394 DECOMPOSITION OF Fe(CO), ON ZEOLITES 2 0 1 0 0 degassing temperature/K Fig. 1. Weight-loss curves upon thermal decomposition under vacuum in a McBain balance of zeolites loaded with Fe(CO),. The horizontal lines represent the theoretical levels in the case that no iron is lost during desorption. and LTL zeolites only slightly exceeds the kinetic diameter of the sorbate, reduced mobility and effective pore-mouth blocking seems to hinder the loss of Fe(CO), during the decomposition procedure.DECARBONYLATION OF F~(CO),-ZEOLITE ADDUCTS IN A THERMOBALANCE The thermoanalytical changes occurring during heating of Fe(CO),-zeolite adducts are shown in fig. 2 (A) for the FAU and Cs-FAU samples and in fig. 2 (B) for the H-FAU and FAU* zeolites. In every case an endotherm d.t.a. effect occurs which is accompanied by a relatively fast decrease in sample weight. For FAU and Cs-FAU samples [fig. 2(A)] the sample weight starts to decrease slowly but then suddenly drops when the endothermic process occurs. After this endothermic event the decarbonylation again proceeds slowly, as can be seen from the slow decrease in weight.As shown in table 2, MAZ and LTL behave similarly. The H-FAU and FAU* matrices [fig. 2(B)] show different behaviour. On these samples the decarbonylation is complete when the endothermic reaction has stopped. On FAU*, the rate of weight loss increases gradually and a distinction between the slow and fast decarbonylation processes can only be made formally. With H-FAU a very broad endotherm is found in which two steps can be distinguished. When the characteristic d.t.a. temperatures ( K , T, and are the temperatures at which the d.t.a. peak starts, is at its maximum and ends, respectively) and the temperatures at which decarbonylation is complete are compared for the faujasite-type zeolites (table 2), the following sequence for the stability of sorbed Fe(CO), as well as for the intermediates is obtained: FAU* < H-FAU < FAU < CS-FAU.T.BEIN AND P. A. JACOBS 1395 (A) 1 40 - 30 20 - - 10 - 40 30 20 10 ...................... 1 ; ................ CS-FAU . . .. # '..\ ... j d.t.a.0.l K I ! 20 10 ....................... I 1- .................... FAU * ; ' y .I' . . . ,,.**' /" !,/ 1 d.t.a.O.l K I . ] d.t.g.0.1 mg min-' l l l l l l l , l , ' ! , l , , I , / , , 300 4 00 500 TIK Fig. 2. Thermogram of the decarbonylation of Fe(CO),-zeolite adducts performed at a heating rate of 2 K min-l: (-) weight-loss curve, (. . .) d.t.g. signal; (---) d.t.a. endotherm. Table 2. Characteristic temperatures during thermoanalysisa of Fe(CO),-zeolite adducts d.t.a. endotherm T/K at which reached (T,) Tm/K G/K constant weight is TIK zeolite (start) (peak) (end) FAU 39 1 414 424 553 CS-FAU 400 414 419 573 H-FAU 367 380b 420 420 400b FAU* 369 39 1 395 395 LTL 383 (420)e 448 563 MAZ 363 378 383 443 a With 10k2 mg of sample, using a heating rate of 2 K min-l.Broad endotherm consisting of two ill-defined maxima. Very weak endotherm. Since the cage dimensions are almost similar in these samples, it seems that the nature of the cation determines the stability of the unperturbed as well as the partially decomposed Fe(CO),-zeolite adducts. A comparison of the LTL and MAZ structures indicates higher stability for the LTL-carbonyl adduct (table 2). This may be the effect of the charge-compensating1396 7 1800- E 1 rd P E 1900 - > B 2000 - DECOMPOSITION OF Fe(CO), ON ZEOLITES - 400 - - - -380 - - - - 360 r I 1700 \cs - FAU c d , 3.50 3.7 5 4.0 4.25 Sanderson electrogegativity Fig.3. Correlation between the Sanderson electronegativity of the zeolite and (a) the temperature Ti at which the fast decarbonylation of the Fe(CO),-zeolite adduct starts (a) and (b) the stretching frequency of residual CO ligands after almost complete decarbonylation of the adduct (A). cation (K in LTL as against Na in MAZ) or of the smaller pore diameter in the LTL structure, which may result in a more effective pore mouth blocking. In order to quantify the influence of changes in the chemical composition of the zeolites, the characteristic temperatures of the d.t.a. endotherm have been correlated with their average electronegativity calculated according to Sanderson.laV l9 There is now ample evidence in the literature that variation in zeolite properties can be explained reasonably well and even predicted using this formalism.*^ 5 9 2o As an example, in fig.3(a) a correlation between and the Sanderson electronegativity is shown to hold for the homologous series of FAU structures, while values for LTL and MAZ deviate. Two effects emerge which seem to determine the stability of the adduct: the chemical composition of the zeolite and the size and geometry of the zeolite channels. For the FAU series the stability of the adducts decreases with decreasing basicity of the samples. A more extensive discussion on this will be found in the discussion of the i.r. results. Iron loadings obtained after complete decomposition of the adducts (table 3) range from 1 to 10 wt % .For the FAU samples, the loading is ca. 10% , except for FAU* on which a significant loss of carbonyls is found during the heat treatment. The iron loadings on the other structures (2.4% on MAZ, 2.8% on LTL and 1% on MOR) are sufficiently high to use the iron for catalytic applications. INTERMEDIATES FORMED UPON THERMAL ACTIVATION OF Fe(CO),-zEoLIm ADDUCTS From table 3 it also follows that, for the decomposition procedures used, the samples lose the same relative amounts of carbonyl at the specific temperatures determined by the endothermic effect. During the pre-endothermic period, i.e. before the characteristic endotherm is observed, most samples lose 20-30% of total sorbate weight.At the end of the endotherm between 60 and 70% has disappeared. When complete decarbonylation has occurred the residual iron on the samples is betweenT. BEIN AND P. A. JACOBS 1397 Table 3. Decarbonylation of Fe(CO),-zeolite adducts in a thennobalance:" (a) the fraction (I;) of the original loading remaining adsorbed at the characteristic d.t.a. temperatures (q, T f and c); (b) the average stoichiometries ( X ) of [Fe(CO)& formed during decarbonylation at the characteristic d.t.a. temperatures T T f T, iron loadingb sample F X F X F X (wt %) FAU 68 2.8 33 0.3 29 0.0 9.7 & 0.7 Ca- FAU 74 3.2 37 0.6 26 - 0.2 5.8 2.0 FAU+ LTL 72 3.2 35 0.4 27 -0.1 2.8 MAZ 81 3.7 42 0.9 33 0.3 2.4 H-FAU 86 4.0 38 0.7 31 0.2 10.5 - 13 - - - - a With 10 f 2 mg of sample, using a heating rate of 2 K min-l.Iron metal per weight of Taken at the end of the first endotherm at 413 K dry zeolite after complete decarbonylation. [see fig. 2(b)]. 26 and 33 % of the original weight of sorbate. Only sample FAU* behaves exceptionally, since not only CO but also an important amount of iron is lost. The data presented allow us to state the following. (1) During a thermoanalytic experiment, Fe(CO), sorbed on zeolites can be decarbonylated without any detectable loss of iron. Only on a dealuminated Y zeolite is iron loss significant during this treatment. Note that during vacuum decomposition of the carbonyl, iron loss is always important. (2) Decarbonylation occurs stepwise on the alkali-metal-cation-exchanged zeolites. On average, [Fe(CO),], species are formed during a slow decomposition step.A subsequent fast reaction transforms these species to [Fe(CO),], intermediates. x is between 0.3 and 0.8 and rn and n represent respective nuclearity. Detailed average stoichiometries for the intermediates [Fe(CO),ly can be derived from the data of table 3 (a) using the relation indicated. The results are given in table 3 (b). (3) The H-FAU-carbonyl adduct behaves differently compared to the zeolites containing alkali-metal cations. The first slow release of one CO liganc! per atom of iron is followed by two fast endothermic processes, involving the removal of ca. three and one CO ligands per Fe, respectively. From the homogeneous gas-phase decomposition of iron pentacarbonyl, it was derived that the first Fe-C bond scission should be rate determining.' Its bond energy amounts to 201 kJ mol-l, which exceeds by far the average Fe-C bond energy of 117 kJ mol-l.' These results, as well as those from a recent study of the laser-induced photodissociation of Fe(CO), vapour,21 provide evidence for a non-statistical decom- position process.The Fe-C bond strength of the Fe(CO), intermediates, determined using laser photoelectron spectrometry,22 are 2.4, 0.2, 1.4 and 1 eV for Fe(CO),, Fe(CO),, Fe(CO), and Fe(CO),, respectively. This implies that decomposition of Fe(CO), vapour occurs via a slow removal of the first CO ligands. Essentially the same is observed in the present work for Fe(CO), adsorbed on zeolites. On the contrary, thermogravimetric and i.r. results indicate that no such interme- diates exist when Fe,CO,, in a KBr matrix is thermally dec~mposed.~~ On NaY zeolite,1398 DECOMPOSITION OF Fe(CO), ON ZEOLITES however, iron dodecacarbonyl decomposition also seems to occur stepwise, as shown by i.r. spectroscopy and temperature-programmed desorption.l4 The following values of x are observed at the temperatures indicated: 343 K 388 K 423-473 K 4 4 2 . 8 d 6 1.2 ---+ 0. The release of 1.2 CO per Fe was associated with the occ~rrence~~ of a dismutation reaction of Fe,(CO),, into Fe2+ and Fe3(CO)T; ions instead of a simple decarbonylation reaction. Although the sequence for the values of x closely resembles the sequence obtained in this work for decomposition of Fe(CO), on alkali-metal-cation zeolites, including faujasite, the dismutation reaction is discarded for reasons explained below.The thermal decomposition of Fe(CO), on y-alumina under a flow of inert gas was also followed using temperature-programmed desorption.'l In this way iron losses up to 50%, and as a result low FeO loadings (up to 0 . 3 wt %), are found. The number of ligands ( x ) per Fe were found to vary as follows: 5+2&0. Evolution of CO started at 333 K, reached a maximum at 413 K and slowly continued up to 670 K. It may be concluded from all these results that the decomposition of iron carbonyls, supported or not, starts with slow evolution of a few CO ligands, followed by fast release of the major part of the ligands. With most support materials, complete removal of residual CO occurs only slowly. Sensitive monitoring of the accompanying thermal processes is needed to detect these stepwise variations in the number of CO ligands.This is most probably the reason why different authors sometimes obtain conflicting data. fast slow APPARENT ACTIVATION ENERGY FOR DECOMPOSITION OF Fe(CO), ON FAU From the variation of the temperature at which a d.t.a. maximum occurs with the heating rate, the apparent activation energy of the process can be determined,24 provided the volume of the sample cell is small and the reaction rate is m~derate.~, This was true for the present experiment. When the heating rate was varied between 0.2 and 2 K min-l an apparent activation energy of 100 _+ 20 kJ mol-1 was calculated for the decomposition of Fe(CO), on FAU in the temperature region of the d.t.a.endotherm. From deposition experiments of Fe(CO), on a value of 84 kJ mol-1 was found for the apparent activation energy of pentacarbonyl decomposition. For Fe(CO), on y-alumina a value of 126 kJ mol-l could be determined.ll It seems that the apparent activation energy of Fe(CO), decomposition is independent of the substrate and represents for a major part the slowest step, namely the initial release of two CO ligands. The value obtained differs significantly from the dissociation energy reported for the removal of the first ligand from gaseous Fe(CO),, i.e. 201 kJ mol-l.' Since decomposition of Fe(CO), starts at 333 K in contact with metallic iron,27 an autocatalytic step seems to be involved in this reaction. This is confirmed by t.p.d. results. l1 I.R.INVESTIGATION OF THE THERMAL DECOMPOSITION OF IRON PENTACARBONYL-ZEOLITE ADDUCTS In situ thermal decomposition of the adducts was carried out in a static helium atmosphere. The spectral changes in the CO-stretching region were recorded system- atically during decarbonylation. Representative spectra for the adducts of Fe(CO), with FAU, Cs-FAU, LTL and H-FAU zeolites are shown in fig. 4-7, respectively.T. BEIN AND P. A. JACOBS 1399 1945 J 1 I 1 1 I I 1 I I !2 21 20 19 10 17 wavenumber/lO-* em-' Fig. 4. 1.r. spectra of the CO-stretching region during decarbonylation in inert atmosphere of Fe(CO),-FAU adducts: (A) FAU zeolite degassed at 720 K, (B) the adduct at 10% saturation, (C) the saturated adduct heated at 373 K for 15 min in 60 kN mb2 of helium, (D) heated at 420 K for 15 min under helium, (E) vacuum degassed at 420 K, (F) heated at 470 K for 70 min after introduction of the same amount of helium.BEHAVIOUR OF ALKALI-METAL-CATION ZEOLITES During the decomposition of Fe(CO), on the zeolites FAU, Cs-FAU and LTL (fig. 4-6) the original CO-stretching bands of the adduct in the region of linearly bound CO gradually decrease in intensity, while in the region of bridged CO (below 1900 cm-l) the i.r. absorption increases. The latter bands are found at 1860 cm-l for FAU, at 1875 cm-l for LTL and at 1818 cm-l for Cs-FAU. As a result of their intermediate appearance in the decomposition process, these bands can be associated with the existence of the [Fe(CO),], fragments evidenced by the d.t.a. experiments. The question arises as to whether cluster formation occurs already with these fragments prior to complete decarbonylation.The CO-stretching bands in the frequency range below 19OOcm-l can alternatively be assigned to bridged CO ligands or anionic carbonyl species, respectively. The interpretation of the data presented by assuming bridged CO is preferred for the following reasons. (i) Formation of anionic species in dehydrated alkali-metal zeolite would require a dismutation reaction providing cations such as Fe2+. These have never been detected after complete decomposition of the adducts. (ii) Assuming1400 DECOMPOSITION OF Fe(CO), ON ZEOLITES wavenumber/10-2 cm-I Fig. 5. Spectra of CO stretching for the Fe(CO),-Cs-FAU adduct: (A) the adduct at 10% saturation, (B) the saturated adduct heated up to 400 K in helium, (C) heated at 470 K for 40 min, (D) at 470 K for 2 h, (E) degassed at 720 K.the intermediate formation of anionic species, it would be difficult to explain the ultimate generation of FeO particles after complete decomposition as observed by magnetic mea~urements.~~ (iii) With a localized negative charge on a certain inter- mediate, the striking influence of the zeolite cations on the CO-stretching frequencies would also be difficult to explain (see below). However, only complementary experiments such as in situ magnetic measurements would rule out completely the assignment of the low-frequency bands to anionic mononuclear complexes.4o In analogy to the d.t.a. results the bands which remain on these zeolites at relatively high temperatures must be assigned to highly unsaturated [Fe(CO),], species, with an average stoichiometry of x = 0.5.These bands are seen at 1895, 1875 and 1713 cm-l for FAU, LTL and Cs-FAU, respectively. Although the bands for Cs-FAU are in a region where surface carbonate or carboxylate species 29 they cannot be assigned to these species, since the vibrations expected in the 1500-1300 cm-l range are not observed. An explanation for the influence of the nature of the matrix on the frequency of these bands will be advanced below. MAZ and MOR zeolites adsorb only a limited amount of Fe(CO),. Upon decomposition, these adducts are expected to behave in the same way as FAU, although this could not be verified owing to the low intensity of the residual CO bands.The existence of a dismutation reaction of Fe,(CO),, adsorbed on FAU has been invoked14 duringT. BEIN AND P. A. JACOBS 1401 I **’ 2116 Fig. 6.1.r. spectra of the decarbonylation of Fe(CO),-LTL: (A) the adduct at 10% saturation, (B) saturated adduct in helium, (C) heated at 420 K for 1 h, (D) for 2 h, (E) heated at 470 K for 30 min, (F) heated up to 550 K, (G) degassed at 720 K. moderate heating at 333 K. Evidence for this reaction was based only on the similarity of the CO stretching frequencies of the expected dismutation product [Fe,(C0),,]2- and the heated Fe,(CO),,-FAU adduct, respectively. For the following two reasons this assumption is questionable: (i) the assignment of some of the carbonyl bands ob- served to particular species such as [Fe,(C0),J2- is arbitrary, since the CO-stretching patterns of carbonyl complexes are sensitive to solvent effects14 and often are found in the frequency region considered here; (ii) Fe2+ ions, which are expected as the other dismutation product, have been observed neither after decomposition of a Fe,(CO),,- FAU adduct14 nor after decomposition of a Fe(CO),-FAU adduct (this work).BEHAVIOUR OF THE CATION-FREE FAUJASITE Upon decomposition of the FAU*-Fe(CO), adduct, the i.r. bands in the CO- stretching region gradually decrease in intensity and only minor frequency shifts and changes in peak intensities occur. It has already been concluded6 that in this adduct only a small interaction exists between matrix and sorbate. This explains why the major amount of carbonyl is desorbed upon heating.The [Fe(CO),], intermediates are not formed, since stabilisation by the matrix is lacking.1402 DECOMPOSITION OF Fe(CO), ON ZEOLITES I"" 2120 \ --.*, wavenumber/l O-' cm-' Fig. 7.1.r. spectra of the decarbonylation of Fe(CO),-H-FAU : (A) the adduct at 10% saturation, (B) the saturated adduct, (C) heated in helium at 400 K for 10 min, (D) for 25 min, (E) for 32 min, (F) for 16 h at 400 K, (G) degassed at 720 K. BEHAVIOUR OF ACIDIC FAUJASITE Upon decomposition of the H-FAU-carbonyl adduct, most of the broad low- frequency CO bands which are found for the cationic forms of this zeolite are not generated (fig. 7). Only one new band at 1875 cm-l, of minor intensity, appears in this frequency range during the decomposition process. This low-frequency band disappears even faster than the bands around 2000 cm-l.The low number and low thermal stability of the CO bands provide evidence for another but simpler decomposition pattern. The intermediate formation of Fe,(CO),, has been postulated in this case.l37 l5 Inspection of the OH region during adduct decomposition provides supplementary evidence for this mechanism (fig. 8). Upon adsorption a hydrogen bond of moderate strength is formed with all protons available in the supercage: the 3645 cm-l band shifts to 3550 cm-l. After complete decarbonylation both the 3645 and 3550 cm-l bands are only partially restored [fig. 8 (C)], which indicates that deprotonation of the lattice has already occurred at these low temperatures. The extent to which this occurs depends very much on the decomposition conditions.After heating at 420K for 45 min almost 75 % of the original intensity is restored [fig. 8 (C)], while after prolonged heating at 400 K for 960 min only 25% of the OH groups resist [fig. S(D)]. The extentT. BEIN AND P. A. JACOBS 1403 I I I I I I I I I 38 36 34 32 30 Fig. 8. Interaction of the OH groups of H-FAU with Fe(CO),: (A) OH groups after degassing at 720 K, (B) saturation with Fe(CO),, (C), the adduct heated at 420 K for 45 min, (D) the adduct heated at 400 K for 16 h. wavenumber/10-2 cm-' to which the OH groups disappear is a measure of the amount of FeJ1 ions formed during the acid reaction. A possible reaction is the oxidation of sorbed carbonyl :139 30 Fe(CO), + 2H+-+H2 + Fez+ + 5CO. The data of fig.8 show that this cannot be the major pathway. Indeed, after prolonged heating at a low temperature more OH groups are removed than after shorter heating at a higher temperature [fig. 8(C) and (D)]. This is only possible if the oxidation of FeO clusters is a major reaction pathway: FeO + 2H++H2 + Fe2+. AN ATTEMPT TO RATIONALIZE THE INFLUENCE OF THE MATRIX UPON THE DECOMPOSITION OF THE ADDUCT There is general agreement in the l i t e r a t ~ r e ~ l - ~ ~ as to the influence of the number of electron-donor ligands (L) on the frequency of the CO bands in L, Fe(CO),-,. When the value of x increases, a decrease of the CO-stretching frequency is found. The increased stability of Fe(CO), in FAU compared with H-FAU was explained by1404 DECOMPOSITION OF Fe(CO), ON ZEOLITES -**.,,,:..'.'..**... *... ...... ..+."I ........ r... ,...C.. ....... -...., 1 I 1 I I 1 I 1 22 21 20 19 18 17 wavenumber/10-2 cm-1 Fig. 9. 1.r. spectra of the stretching region of NO, sorbed at 293 K on decarbonylated Fe(CO),-zeolite adducts: (A) the carbonyl-LTL adduct heated at 550 K, (B) addition of NO at 293 K (lo4 N md2), (C) room-temperature degassing at 10+ N m-2 for 10 min, (D) the carbonyl-FAU adduct heated at 470 K, (E) the same as for (B), (F) the same as for (C), (G) the carbonyl-Cs-FAU adduct heated at 550 K, (H) the same as for (B), (I) the same as for (C). increased backbonding in the Fe-CO bond, resulting from increased electron-donor properties of the matrix.15 This hypothesis is confirmed in the present work. Indeed a monotonous correlation exists between the Sanderson electronegativity of the zeolite and the thermal stability of the adducts [fig.3(a)]. When the zeolite-Fe(CO), adduct is more stable, the CO frequencies are found at lower wavenumbers. Another monotonic correlation is established between the CO frequencies and the Sanderson electronegativity values of the zeolite [fig. 3(b)]. It follows that when the overall electron-donor properties of the zeolite matrix (as determined by their overall chemical composition) increase, the corresponding CO-stretching vibrations of the [Fe(CO),], intermediates are found at lower wavenumbers. PROBING OF THE DECARBONYLATED ADDUCTS WITH NO Upon room-temperature addition of NO (10 kN m-2) to the decarbonylated adducts of iron pentacarbonyl and the alkali-metal-cation-exchanged zeolites (LTL, FAU and Cs-FAU), spectra of remarkable similarity are observed (fig.9). Apart from the stretching frequency of gas-phase NO at 1875 cm-l, two typical bands areT. BEIN AND P. A. JACOBS 1405 A 1878 \ r '.. 1813 .... ..... &...'' ..... .................. .......................... ................. I I I I 1 I I 1 1 22 21 20 19 18 17 wavenumber/lO-' cm-' Fig. 1O.I.r. spectra of NO adsorbed at 293 K on a decarbonylated Fe(CO),-H-FAU adduct: (A) the adduct heated at 430 K, (B) addition of NO at 293 K ( lo4 N m-2), (C) room-temperature degassing at N m-2 for 10 min. observed, between 1810 and 1800 cm-l and between 1735 and 1710 cm-l. Bands at 1810 and 1720cm-l with the same relative intensities have been reported for NO chemisorbed on Fe suspended in The assignment of the two bands has been firmly established :36 NO chemisorbed on metallic iron absorbs at 1720 cm-l, while the band around 1800 cm-l is assigned to NO sorbed on surface-oxidized iron particles.It is evident that since in zeolites NO disproportionation was rep~rted,~' the surface of the iron phase may be oxidized by N,O formed this way. The spectra of NO sorbed on the Fe/H-FAU system are given in fig. 10. The most significant peaks are at 1878, 1813 and 1765 cm-l. Comparable spectra were obtained with H-FAU-iron dodecacarbonyl adducts.13 Based on earlier work3* on Fe-Y zeolites, the 1878 and 1765 cm-l bands have been assigned to [Fe(N0)I2+ complexes of high and low spin, respectively. The band at 1813 cm-l was assigned13 to NO sorbed on FeO, which is in agreement with the present interpretation for the 1800 cm-l band on alkali-metal-cation zeolites. This assignment allows us to conclude the following with regard to the nature of the Fe after decomposition of the adducts.(i) On H-FAU a considerable amount of FeII is located in cationic positions (high- and low-spin complexes with NO) as a result of the oxidation of the metal phase and irreversible removal of the acidic protons. This phenomenon is absent on alkali-metal-cation zeolites since no Brsnsted-acid sites are1406 DECOMPOSITION OF Fe(CO), ON ZEOLITES present. (ii) The surface of the remaining iron cluster on H-FAU is easily oxidizable by NO at room temperature, while this is only partially true for Fe clusters on alkali-metal-cation zeolites.The present data do not permit an explanation of this difference. (iii) The interaction of NO with iron clusters is stronger than with surface-oxidized clusters; this becomes evident when spectra B and C, E and F, H and I (fig. 9) are compared. CONCLUSIONS The combined use of thermogravimetric and i.r. methods allows one to depict the decomposition of Fe(CO),-zeolite adducts as a stepwise decarbonylation reaction, consisting of slow (s) and fast ( f ) phenomena. For alkali-metal zeolites of different structure it occurs as follows: For HY zeolites the following sequence emerges : The fast phenomena are endothermic, characterized by an apparent activation energy of lOO+20 kJ mol-l. 1.r. results show that the intermediates cannot be exclusively of a mononuclear nature.The stability of the Fe(CO), intermediates as well as their CO-stretching frequency increases with the electron-donor properties of the substrate, in this way enhancing the Fe-CO bond strength through increased backbonding. Use of NO as probe indicates that in H-faujasites FeII ions can take cationic positions and remove the acidic protons. T. B. acknowledges grants from the Deutscher Akademischer Austauschdienst, the Belgian Ministry for Education (Ministerie van Nationale Opvoeding en Neder- landse Cultuur) and the Alfried Krupp von Bohlen und Halbach-Stiftung. P. A. J. acknowledges a permanent research position as Senior Research Associate from the N.F.W.0-F.N.R.S. (Belgium). Discussions with Priv.-Doz. Dr F. Schmidt and Prof.W. Gunsser (University of Hamburg) are much appreciated. Financial support from the Belgian Government (Concerted Actions on Catalysis, Diensten Wetenschapsbe- leid) is also acknowledged. Y-Y. Huang and J. R. Anderson, J . Catal., 1975,40, 143. F. Schmidt, W, Gunsser and J. Adolph. ACS Symp. Ser., 1977,40, 291. W. Gunsser, J. Adolph and F. Schmidt, J. Magn. Magn. Muter., 1980, 15-18(II), 1 1 15. W. J. Mortier, J. Catal., 1978, 55, 138. P. A. Jacobs, W. J. Mortier and J. B. Uytterhoeven, J. Inorg. Nucl. Chem., 1978, 40, 1919. Th. Bein and P. A. Jacobs, J. Chem. Soc., Faraday Trans. 1, 1983,79, 1819. G. P. Smith and R. M. Laine, J . Phys. Chem., 1981,85, 1620. A. Terenin and L. Roev, Spectrochim. Acta, 1959, 11, 946. E. Guglielminotti, A. Zecchina, F. Boccuzzi and E.Borello, in Growth and Properties of Metal Clusters, ed. J. Bourdon (Elsevier, Amsterdam, 1980), p. 165. lo F. Hugues, B. Besson, P. Bussiere, J. A. Dalmon, J. M. Basset and D. Olivier, Nouo. J. Chim., 1981, 5, 207. l 1 A. Brenner and D. A. Hucul, Inorg. Chem., 1979, 18, 2836. l2 D. Ballivet-Tkatchenko, G. Coudurier, H. Mozzanega and I. Tkatchenko, in Fundamental Research in Homogeneous Catalysis, ed. M . Tsutsui (Plenum Press, New York, 1979), vol. 3, p. 257.T. BEIN AND P. A. JACOBS 1407 l3 D. Ballivet-Tkatchenko and G. Coudurier, Inorg. Chem., 1979, 18, 558. l4 D. Ballivet-Tkatchenko, G. Coudurier and Nguyen Duc Chau, in Metal Microstructures in Zeolites, l5 Th. Bein, P. A. Jacobs and F. Schmidt, in Metal Microstructures in Zeolites, ed. P. A. Jacobs, P. Kirii, ed. P. A. Jacobs, P. Jiru, N. Jaeger and G. Schulz-Ekloff (Elsevier, Amsterdam, 1982), p. 123. N. Jaeger and G. Schulz-Ekloff (Elsevier, Amsterdam, 1982), p. 1 11. J. B. Nagy, M. Van Eenoo and E. G. Derouane, J. Catal., 1979, 58, 230. Pittsburgh, 1978). l7 D. H. Olson and W. M. Meier, Altas of Zeolite Structure Types (IZA, Polycrystal Book Service, lR R. T. Sanderson, J. Coll. Sci. Teach., 1972, 1, 16; 47. l9 R. T. Sanderson, Chemical Bonds and Bond Energy (Academic Press, New York, 2nd edn, 1976). 2o P. A. Jacobs, Catal. Rev. Sci. Eng., 1982, 24, 415. 21 J. T. Yardley, B. Gitlin, G. Nathanson and A. M. Rosan, J. Chem. Phys., 1981, 74, 370. 22 P. C. Engelking and W. C. Lineberger, J. Am. Chem. Soc., 1979, 101, 5569. 23 R. Psaro, A. Fusi, R. Ugo, J. M. Basset, A. K. Smith and F. Hugues, J. Mol. Catal., 1980, 7, 51 1. 24 H. E. Kissinger, Anal. Chem., 1957, 29, 1702. 25 K. Akita and M. Kase, J. Phys. Chem., 1968, 72, 906. 26 H. E. Carlton and J. H. Oxley, AIChE J., 1965, 11, 79. 27 A. Mittasch, Z. Angew. Chem., 1928, 41, 831. 28 P. A. Jacobs, F. H. Van Cauwelaert, E. F. Vansant and J. B. Uytterhoeven, J. Chem. SOC., Faraday 29 P. A. Jacobs, F. H. Van Cauwelaert and E. F. Vansant, J. Chem. SOC., Faraday Trans. I , 1973, 69, 30 J. Dewar and H. 0. Jones, Proc. R. SOC. London, 1905, 76, 569. 31 M. Bigorgne, J. Organometal. Chem., 1970, 24, 21 1. 32 B. F. G. Johnson, J. Lewis and M. V. Twigg, J. Chem. Soc., Dalton Trans., 1974, 241. 33 M. Poliakoff, J. Chem. SOC., Dalton Trans., 1974, 210. 34 P. Poliakoff and J. J. Turner, J. Chem. SOC., Dalton Trans., 1974, 2276. 35 C. Blyholder and M. C. Allen, J. Phys. Chem., 1965, 69, 3998. 36 H. Bandow, T. Onishi and K. Tamaru, Chem. Lett. (Jpn), 1978, 83. 37 C. C. Chao and J. H. Lunsford, J. Am. Chem. SOC., 1971,93, 71. 38 J. W. Jermyn, T. J. Johnson, E. F. Vansant and J. H. Lunsford, J. Phys. Chem., 1973, 77, 2964. 39 F. Schmidt, Th. Bein, U. Ohlerich and P. A. Jacobs, Proc. Sixth Int. Zeolite Conf., Reno 1983 Trans. I , 1973, 69, 1056. 2 130. (Butterworth Scientific Press, Sevenoaks, 1983, in press). (PAPER 3/635)

ISSN:0300-9599    

DOI:10.1039/F19848001391

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1984, 80, 1409-1414 Enthalpy of Solution of Calcium Chloride in Aqueous Mixtures of Methanol, Ethanol and Propan- 1-01 at 298.15 K BY STEFANIA TANEWSKA-OSINSKA* AND JOLANTA BARCZYNSKA Department of Physical Chemistry, University of Lodz, 91-416 Lodz, Poland Received 27th May, 1983 Enthalpies of solution of CaC1, in aqueous mixtures of methanol, ethanol and propan-1-01 have been measured over the entire range of mixed-solvent compositions. Plots of the standard enthalpy of solution against composition exhibit a maximum in the water-rich region and a minimum in the alcohol-rich region. Enthalpic pair-interaction coefficients have also been calculated: these were found to be positive for electrolyte-alcohol pairs in water and negative for electrolyte-water pairs in alcohols.Solutions of electrolytes in water + organic-solvent mixtures have been investigated by many A number of studies conducted in our laboratory have been devoted to the physico-chemical properties of NaI in aliphatic alcohols and their mixtures with ~ a t e r . ~ - ~ We are currently concerned with 1 : 2-type electrolytes, and the subject of the study reported here was CaC1,. We present enthalpies of solution of CaCI, in aqueous mixtures of methanol, ethanol and propan-1-01 obtained calori- metrically at 298.15 K. EXPERIMENTAL Puriss-grade anhydrous CaCl, produced by POCh Gliwice (Poland) was dried in a glass-tube furnace under a continuously flowing atmosphere of gaseous HC1 at 530 K. The HC1 gas stream was then replaced by argon and heating was continued for an additional 0.5 h, after which the product was furnace-cooled under argon.The CaC1, was then transferred into a specially designed container from which glass ampoules were filled with the desired doses of the salt. The degree of reagent purity was established by analytic methods. The content of C1- ions in the anhydrous CaC1, was determined potentiometrically and found to be 99.9%. All the alcohols used in the study were puriss grade produced by POCh Gliwice (Poland) and were purified by standard methods.* Measurements of the enthalpies of solution of anhydrous CaCl, in the water + alcohol mixtures were performed with the aid of an isoperibol calorimeter. The calorimetric vessel had a capacity of ca. 180 cm3. A balanced Wheatstone bridge with an NTC-type thermistor temperature indicator with a resistance of 40 k n at T = 293.15 K was used as the measuring instrument.The voltage of the unbalanced bridge was determined using a MERA V-534 (Poland) digital voltmeter. The overall temperature sensitivity of the calorimeter was found to be ca. 1 x lop4 K, and heat effects could be measured with a precision of ca. 0.5%. The heat effects resulting from breaking ampoules in the reaction vessel were found to be negligible. 14091410 SOLUTION ENTHALPY OF CaCl, IN AQUEOUS ALCOHOLS RESULTS AND DISCUSSION The measurements of the heat of solution of CaCl, in aqueous mixtures of methanol, ethanol and propan-1-01 were performed at 298.15 K over the whole composition range. The isotherms AHm =f(rn), where rn is the molality of CaCl, obtained exhibit decreasing exothermicity of the thermal effect of solution with increasing CaCl, concentration.As an example table 1 shows the integral enthalpies of solution of CaC1, in water+methanol mixtures containing 5 and 15 mol% methanol. Table 1. Integral enthalpies of solution (AH,/kJ mol-l) or CaC1, in water + methanol mixtures at 298.15 K 0.0014 - 80.0 0.0019 0.0038 - 79.0 0.0020 0.0047 - 79.3 0.0023 0.0048 - 78.8 0.0024 0.0085 - 78.4 0.0027 - - 0.0029 - - 0.0040 AH0 = - 80.4 & 0.3 0.0040 - AH0 - - 79.6 - 79.1 - 79.3 - 79.2 - 79.0 - 78.8 - 78.3 - 78.9 z= - 80.6 & 0.5 a Mol% of alcohol. Concentration of CaC1, in mol kg-l of solvent. AH* is the standard enthalpy of solution of CaCI, & the standard deviation of the mean.The measured enthalpies of solution of CaCl, were extrapolated to infinite dilution by the method proposed by Criss and C ~ b b l e . ~ The dielectric constants and densities were taken from the literature.lOJ1 The standard solution enthalpies of CaCl, in mixtures of the three alcohols with water are presented in table 2 and fig. 1. In both systems analysed in the present study, i.e. CaCl, + H,O + EtOH and CaC1, + H,O + n-PrOH, there are maxima in AH0 (fig. 1) corresponding to ca. 12 and 8 mol% alcohol, respectively. Maxima in the standard solution enthalpies of electrolytes are characteristic of all water + alcohol systems investigated thus 57 6 v 12* l3 Their appearance is attributed to the stabilizing effect of a small addition of alcohol on the three-dimensional structure of water.3* 14* l5 Although within experimental error the positions of the maxima in A H 0 (CaCl,) (fig.1) for aqueous ethanol and propan-1-01 mixtures are the same as those for other electrolyte^,^ the shapes and heights of the peaks are different. These differences may be due to the higher valence of the Caz+ cation or to the greater number of ions arising from one ‘molecule’ of the electrolyte. An especially distinct difference between the behaviour of the electrolytes studied earlier3* 4 9 l6 and CaCl, is observed in methanol + water mixtures. As can be seen in fig. 1, the CaC1, + H,O + MeOH system exhibits no maximum in A H e . The flat shape of the function AH* =f(mol% methanol) in the range 0-17 mol% methanol (fig. 1) suggests the existence of two thermal effects which compensate each other.First, theres. TANIEWSKA-OSINSKA AND J. B A R C Z Y ~ K A 141 1 Table 2. Standard enthalpies of solution (AH*/kJ mo1-l) of CaC1, in water + methanol, water +ethanol and water + propan-1-01 mixtures at 298.15 K water + methanol water + ethanol water + propan-1 -01 &a A H 0 A H 0 &a A H 0 - 0 2 3 5 7 15 17 20 30 40 50 70 80 85 90 93 95 96 100 - 80.8 - 80.7 - 80.4 - 80.5 - 80.3 - 80.6 - 80.3 -81.3 -85.5 - 92.0 - 94.1 - 103.5 - 107.4 - 110.7 - 112.5 - 113.0 - 1 13.9 -113.6 - 107.9 0 5 8 10 12 15 20 30 40 50 60 70 75 80 85 90 93 1 00 - 80.8 - 78.0 - 76.4 - 74.6 - 74.0 - 75.6 - 77.2 - 82.1 - 87.6 -91.4 -97.5 - 102.7 - 104.8 - 105.1 - 106.3 - 104.2 - 101.0 -91.7 - 0 5 6 8 9 10 15 20 30 40 50 60 70 80 85 90 100 - 80.8 - 76.6 - 75.9 - 74.6 - 74.8 - 75.2 - 77.7 -81.4 -87.1 -93.5 - 98.9 - 103.7 - 107.7 - 112.4 - 105.6 -98.8 - 86.3 a Mol% of alcohol.is the structure-making effect that addition of methanol exerts on water, giving rise to an increased solution enthalpy observed in other water + alcohol +electrolyte systems. We can assume that the other effect, accompanied by a decreased enthalpy of solution, represents the partial incorporation of methanol molecules in the ion hydration envelopes. From 17 mol% methanol in water upwards, the solvent structure in CaCl, solutions becomes disturbed in a similar way to the case of other electrolytes.** l6 In each case this is due to an increased exothermic thermal effect. We have calculated the interaction coefficients17 of ion-alcohol-molecule pairs in water (table 3) on the basis of the enthalpies of electrolyte transfer from water to water + alcohol mixtures determined by us.For the sake of comparison, table 3 also contains pair-interaction coefficients for aqueous alcohol solutions of NaI and NaCl. The positive values of the coefficients h,, for all three electrolytes suggest that ions interact with alcohol molecules only weakly, which may mean that they interact more strongly with water molecules. A comparison of the pair-interaction coefficients for NaCl-alcohol and NaI-alcohol pairs in water shows them to be similar. Thus in the three alcohols studied changing the anion has no effect on the values of these coefficients. On the other hand, their positive values increase on increasing the length of the non-electrolyte chain. The pair-interaction coefficients for CaC1,-alcohol pairs in water increase more rapidly with increasing length of the alcohol molecule than do those for NaCl or NaI.The values of the coefficients h,, for all three electrolytes with propan-1-01 in water1412 SOLUTION ENTHALPY OF CaCl, IN AQUEOUS ALCOHOLS 0 mol % alcohol 5 0 I Fig. 1. Standard solution enthalpy of CaCl, in alcohol + water mixtures as a function of solvent composition: 0, water +methanol; 0, water +ethanol; A, water +propan-1-01. Table 3. Enthalpic electrolyte-non-electrolyte pair-interaction coefficients (hxy/J kg molP) in aqueous solution electrolyte non-electrolyte solvent hXY ref. NaCl NaI CaC1, NaCl NaI CaCl, NaCl NaI CaC1, MeOH MeOH MeOH EtOH EtOH EtOH n-PrOH n-PrOH n-PrOH I50 157 49 290 298 I62 370 395 340 18 16 this work 18 16 this work 18 16 this works.TANIEWSKA-OSINSKA AND J. B A R C Z ~ S K A 1413 are identical within experimental error. One can thus suppose that the structures of the solvation shells of the three ions studied here are identical or almost identical. It has also been found that the pair-interaction coefficients h,, for CaC1,-MeOH and CaC1,-EtOH pairs in water have considerably smaller positive values than those for NaI and NaCl solutions in the same solvents, suggesting that interactions between CaCl, and alcohol molecules (especially methanol) in water are stranger than is the case with NaI and NaCl. Quite possibly this effect may hinder the ordering of water structure by methanol molecules.Further analysis of the plots presented in fig. 1 shows that A H e exhibits minima in the range of high alcohol contents in the mixed solvent. These minima in the standard solution enthalpy of the electrolyte in mixtures of the first three alcohols with water have been observed for the first time in this work. The composition of the mixtures corresponding to the minima are shifted toward lower alcohol contents in the following order : methanol > ethanol > propan- 1-01. However, similar minima were observed for mixtures of butyl alcohols and water in the presence of NaI,5 although not for NaI solutions in mixtures of the first three alcohols with water. The appearance of these minima may reflect changes either in the structhre or in the interactions of water + alcohol mixtures studied in the range of high alcohol contents.The effect of NaI on the structures of aqueous methanol, ethanol and propan-1-01 mixtures is probably too small to reveal the properties of the mixed solvent. On the other hand addition of CaCl,, which contains a greater number of ions and a bivalent cation, makes the change distinct because of the dominant ion-solvent interaction. The appearance of minima in the standard enthalpies of solution of CaCl, in mixtures the first three alcohols with water may be attributed to the formation of a certain kind of alcohol-water associate, which could be similar to those proposed by Franks and Ives for the t-butyl alcohol+water mixtures.19 If this supposition is justified, then we can compare the position of the minima observed in the presence of CaCl, (fig. 1) with those of NaI in t-butyl alcohol + water mixture^.^ The minimum in the standard enthalpy shifts from ca.95 mol% alcohol in methanol + water mixtures to ca. 75 mol% alcohol in t-butyl alcohol + water mixtures. The pair-interaction coefficients calculated for CaC1,-water pairs in alcohols (table 4) have negative values; this has not been observed to date for other salts in Table 4. Enthalpic electrolyte-water pair-interaction coefficients (h,,/J kg molP) in non-aqueous solvents electrolyte non-electrolyte solvent hX, ref. NaI H2O MeOH 583 20 NaI H2O EtOH 532 20 CaCl, H2O -1083 this work EtOH - 1167 this work NaI H2O n-PrOH CaC1, H2O n-PrOH - 1225 this work CaC1, H2O MeOH 174 20 the first three alcohol + water mixtures.With increasing length of the alcohol chain, the values of the pair-interaction coefficients become increasingly more negative. The affinity of this ion for water molecules probably irlcreases too. Further work on 1:2 electrolytes in alcohol solutions is in progress.1414 SOLUTION ENTHALPY OF CaC1, IN AQUEOUS ALCOHOLS C. M. Slansky, J. Am. Chem. Soc., 1940, 62, 2340. K. P. Mischenko and G. M. Poltoratskii, in Problems of Thermodynamics and Structure of Aqueous and Non-aqueous Electrolyte Solutions (Plenum, New York, 1972). G. A. Krestov, in Thermodynamics of Ionic Processes in Solution (Khimiya Press, Leningrad, 1973). A. Dadgar and M. R. Taherian, J. Chem. Thermodyn., 1977,9, 71 1. S. Taniewska-Osinska and H. Piekarski, J. Solution Chem., 1978, 7, 891. S. Taniewska-Osinska, H. Piekarski and A. Kacperska, in Thermodynamics and Structure of Solutions (Ivanovo, USSR, 1976), vol. 4, p. 123. S. Taniewska-Osinska and H. Piekarski, in Thermodynamics and Structure of Solutions (Ivanovo, USSR, 1979), vol. 6, p. 15. * A. Weissberger, E. S. Proskauer, J. A. Riddick and E. E. Toops Jr, Organic Solvents (Interscience, New York, 1955). C. M. Criss and J. W. Cobble, J. Am. Chem. Soc., 1961,83, 3223. lo Y. Y. Akhadov, Dielectric Properties of Binary Solutions (Pergamon Press, Oxford, 1981). l1 Y. Tashima and Y. Arai, Mem. Fac. Eng., Kuushu Univ., 1981,41, 215. l2 N. Dollet and J. Juillard, J. Solution Chem., 1976, 5, 77. l3 Y. Pointud, J-P. Morel and J. Juillard, J. Phys. Chem., 1976, 80, 2381. l4 0. Ya. Samoilov, Zh. Strukt. Khim., 1966, 7, 15; 175. l5 G. Nemethy and H. A. Scheraga, J. Phys. Chem., 1962,66, 1773. l8 H. Piekarski, Can. J. Chem., 1983,61,2203. W. G. MacMillan and J. E. Meyer, J. Chem. Phys., 1945, 13, 276. l8 G. Perron, D. Jolly and J. E. Desnoyers, Can. J. Chem., 1978,56, 552. l9 F. Franks and D. J. G. Ives, Q. Rev., 1966, 20, 1 . 2o H. Piekarski, A. Piekarska and S. Taniewska-Osinska, Can. J. Chem., in press. (PAPER 3/863)

ISSN:0300-9599    

DOI:10.1039/F19848001409

出版商:RSC    

年代:1984

数据来源: RSC