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/F198480FX025

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/F198480BX027

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’ wishesas to whether an article 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 1500 words or word-equivalents. (9NOMENCLATURE 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 ‘ Systkme 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 I , 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.THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO.78 Radicals in Condensed Phases University of Leicester, 44 September 1984 , Organising Committee ~ Professor M. C. R. Symons (Chairman) Dr K. A. McLauchtan Professor Lord Tedder Dr R. L. Willson Dr G. 6. Buxton Dr T. A. Claxton 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 1 and reactivity with spectroscopic and mechanistic studies directly concerned with radicals. ' Fundamental aspects will be stressed, and particular attention will be given to new deuelopments 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 W1 V OBN THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY D EUTSCH E BUNS EN G ES E LLSC HAFT FU R PHY S I KALISCH E CH E M I E SOCICTC DE C H l M l E PHYSIQUE ASSOCIAZIONE ITALIANA D I C H l M l C A FlSlCA Joint Discussion Meeting on: Laser Studies in Reaction Kinetics Evangelische Akademie, Tutzing, West Germany, 2627 September 1984 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)THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM NO. 19 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 B. R. Jennings, Electro-optics Group, Department of Physics, Brunel University, Uxbridge UB8 3PH. Molecular Electronic Structure Calculations-Methods 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 CrystalsTHE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO. 80 Organising Committee : l Professor J. S. Rowlinson (Chairman) Mrs Y. A. Fish 1 Dr E. Dickinson Dr N. Parsonage ~ Dr R. Evans Dr D. A. Young Physical Interactions and Energy Exchange at the Gas-Solid Interface McMaster University, Hamilton, Ontario, Canada, 23-25 July 1985 Organising Committee 1 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 t o 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 The aim of the meeting is to discuss phase transitions at gas/liquid, liquid/liquid 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. Ryan, ICI New Science Group, PO Box 11, The Heath, Runcorn, Cheshire WA7 4QE Gas Kinetics Group 8th International Symposium on Gas Kinetics To be held at the University of Nottingham on 16-20 July 1984 Further information from Professor J. P. Simons, Department of Chemistry, University of Nottingham, Nottingham NG7 2RD Industrial Physical Chemistry Group Bacterial Adhesion To be held at Girton College, Cambridge on 11-1 3 September 1984 Further information from Dr I.D. 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Richards, Department of Pure and Applied Chemistry, University of Strathclyde, Glasgow G1 1 XL (vii)A recent Discussion is I- 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 Memorlal Lecture; Vibrational Redistribution within Excited Electronic States of Polyatomic Mo lccules Intramolecular Relaxation of Excited States of C b F h + 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 Transfer Pulsed Laser Preparation and Quantum Superposition State Evolution in Regular and Irregular Systems A Quantum-mechanical Internal-cohsion 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 Lrradiation Energy Distributions in the CN(X’I+) Fragment from the Infrared Multiplephoton Dissociation of CFjCN.A Comparison between Experimental Results and the Predictions of Statistical Theories Product Energy Partitioning in the Decom- position of State-selectively Excited HOON and HOOD Low-power lnl-rared Laser Photolysis of Tetramethyldioxetan Unimolecular Reactions Induced by Vibrational Overtone Excitation Unimolecular 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 Solva t ion Energy Redistribution in Large Molecules.Dlrect Study of Intramolecular Relaxation in the Gas Phase with Picosecond Gating Rotation-dependent Intramolecular Proccssesof S O > ( A ’ A ? ) in a Supersonic Jet Role of Rotation-Vibration Interaction in Vibrational Relaxation. Energy Redistribution in Excited Singlet Formaldehyde Sub-Doppler, Spectroscopy of Benzene in the “Channel-three” Region Intramolecular Electronic Relaxation and Photoisomerization Processes in the Isolated Azabenzene Molecules Pyridine, Pyrazine and Pyrunidine Softcover 434pp 0 85186 658 1 Price f25.00 ($48.00) Rest of the World f 26.00 RSC Members f 16.25 The Royal Society of Chemistry- Faraday Discussions of the Chemical Societ) No 7 7 Inrrurnolerulor hrncrri Faraday Symposia are usually held annually and are confined to more specialised topics than Discussions, with particular reference to recent rapidly developing lines of research.A recent Symposium is :- NO.~? The Hydrophobic Interactton 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 Precise Vapour-pressure Measurements of the Solubilization 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-alkylammonium 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 50° C Geometric Relaxation in Water. Its Role in Hydrophobic Hydration 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 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 WClS SDT. Letchworth. Herts SO6 IHN, England. Faraday Symposia of the Chemical Society No 1 7 The H ) drophobrc lnferutrron 19x2 (viii)

ISSN:0300-9599    

DOI:10.1039/F198480FP053

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1984’80, 1673-1688 Interaction Forces in Dispersions Containing Non-ionic Surfactants BY LAURENCE THOMPSON Unilever Research, Port Sunlight Laboratory, Quarry Road East, Bebington, Wirral, Merseyside L63 3JW Received 25th April, 1983 Rates and equilibrium positions of aggregation have been measured for dispersions of monodisperse polystyrene latex particles containing the non-ionic surfactant octaoxyethylene glycol dodecyl ether (C12E8). Measurements cover a range of particle sizes and temperatures. Rates are observed which, even in the presence of sufficient electrolyte largely to suppress electrostatic interactions, can be significantly below the diffusion-controlled ‘ rapid ’ rate. Data are interpreted in terms of both primary and secondary minima.The temperature sensitivity of the aggregation appears to arise from variations in the attractive rather than the repulsive contribution, these variations resulting from the effect of temperature on the level of adsorption. The attractive effects are of an essentially short-range nature, and because of this primary- minimum aggregation is more sensitive to temperature than secondary-minimum aggregation. The data have been fitted to potential-energy functions consisting of an attractive van der Waals contribution and a repulsion arising from the surfactant. Making reasonable allowance for the attractive contribution leads to the conclusion that a weak repulsive force is operative at separations considerably in excess of the physical reach of the adsorbed surfactant.Here an arbitrary model for the repulsion has been used which comprises the sum of two exponential decays, a relatively strong short-range component (decay length 0.2 nm) and a weak long-range component (decay length 3.0 nm). Non-ionic surfactants of the alcohol polyethoxylate type are well known as stabilisers of colloidal dispersions. Their effectiveness can decrease with increasing temperature, and flocculation temperatures are often observed below the cloud point of the s~rfactant.l-~ Flocculation is associated with a region of attraction in the plot of potential energy against separation caused by domination of van der Waals forces. It is distinct from the flocculation occurring at the &temperature of a stabilising p ~ l y m e r , ~ - ~ which involves attraction arising from polymer-polymer interaction.The effect is attributable to temperature-induced changes in the adsorbed layer. Reduction of the area per adsorbed molecule, for example, may reflect a reduction of the repulsive forces associated with the surfactant while simultaneously increasing the attractive contribution through its effect on the adsorbed layer’s density.’ The purpose of the work described in this paper is to develop a model for the repulsive force originating within layers of adsorbed non-ionic surfactant so that the relative importance of attraction and repulsion changes in bringing about temperature-induced aggregation can be defined. The temperature dependence of the equilibrium aggregation state of polystyrene latex dispersions containing the surfactant octaoxyethylene glycol n-dodecyl monoether (C,,E,) was described in earlier w0rk.l Conditions were such that the range of the double-layer interactions was very short, being suppressed by electrolyte addition.The aggregation state was related to the effective thickness, 6, of the adsorbed layer 16731674 INTERACTION FORCES IN NON-IONIC SURFACTANTS through an approach based on the second virial coefficient of the particles in the dispersion. A plot of 6 against temperature was constructed which indicated that 6 decreased with increasing temperature from a value of 6.2 nm at 15 "C to 2.2 nm at 43 "C. The lower figure is consistent with the likely dimensions of the C1&8 molecule (random coil length 2.5 nm),8 whereas the larger figure is significantly in excess of even its fully extended length (3.9 nm).The long-range forces implied at the lower temperatures were thought to be of similar origin to the 'solvent-ordering' forces between surfactant or lipid bilayer~,~-ll i.e. they are separate from the osmotic repu1sion12 generated in overlapping adsorbed polymer layers. The main problem with the original interpretation is its foundation on a hard-sphere model for the surfactant's repulsion [fig. l(a)]. This gives a (minimum) estimate of the range of the repulsion but no estimate at all of its dependence on separation. The possibility of obtaining such an estimate from a combination of the equilibrium data and aggregation rate measurements emerged with the realisation that flocculation of the systems can occur at rates below the rapid (Smoluchowski) rate, even in the presence of excess electrolyte. This implies the presence of a maximum in the plot of potential energy against separation [fig.1 (b)], which in itself indicates the limited scope of the hard-sphere approwh. Note that aggregation rates which are slower than the diffusion-controlled (rapid) rate do not necessarily imply the existence of a barrier to aggregation. The rate constant (k,) for aggregation into minima of restricted depth ( Vmin) can be expressedl3 kl=kR 1-exp- Vmin ( kT where kR is the 'rapid' rate constant. An alternative treatment14 suggests an equation of the form k1= kR [1 - ~ X P (Vmin - V*,in)/kTI where Pmin is the value of Vmin below which (using a phase-change model for the aggregation process) aggregation does not occur.Here, however, the observed levels of aggregation imply sufficiently high values of Vmin or (Vmin- Pmi,) that in the absence of a potential-energy barrier k, -+ kR in all cases. Fig. 1 (b) is of similar form to curves typically encountered with charge-stabilised colloids, showing both primary and secondary minima separated by a potential-energy maximum. In the electrostatic case the primary minimum is usually deep and is associated with strong, irreversible and complete aggregation. Here the adsorbed layer restricts the depth of the minimum and weak, reversible and partial aggregation may occur at a rate determined by the magnitude of any potential-energy barrier which may be present. Where a barrier exists, its magnitude is related to the aggregation rate by the expressi on1 9 where the stability constant W is equal to kR/kexptl, the ratio of the rapid rate con- stant (4kT/3q) to the experimental rate constant, q is the viscosity of water, a is the particle radius, and DZ and D,, are the bulk diffusion coefficient of the particles and its value at centre-centre separation, r, respectively:16 V(r) is the potential energy of interaction obtained by summation of attractive and repulsive contributions.Where equilibrium aggregation is observed the depth of the energy well can beL. THOMPSON 1675 v* /” Fig. 1. Schematic plots of potential energy against separation. (a) Hard-sphere repulsion model for particle interactions showing van der Waals attraction V,.(b) Realistic model showing primary and secondary minima separated by a po tential-energy maximum. obtained through the virial-coefficient approach used in the work1? 3* l7 in which the potential-energy function is related to the aggregation state by the expression -= n0 1+2zn, /mr2[exp(F)-l]dr - V(r) - J R where no and n, are the initial and equilibrium particle concentrations, respectively, and R is the shortest distance of centre-centre approach of particles, effectively the inner edge of the energy well. Combination of the ‘energy-barrier ’ and ‘energy-well’ data from a range of particle sizes should in principle yield information on the form of V(r). Since V(r) is the sum of an attractive van der Waals contribution and the repulsive energy that is to be investigated, then provided that a reasonable estimate of the attraction can be made, a model for the repulsion will emerge.To this end, the original equilibrium studies have been extended and the appropriate aggregation rate data have been obtained. EXPERIMENTAL PROCEDURES Aggregation processes were monitored as follows. Test tubes containing the unaggregated dispersions were placed in constant-temperature cabinets. Settling problems were generally avoided by locating the tubes in a rack which was set to rotate end-over-end at ca. 5-10 r.p.m. for 1 min in 15 min. At higher rotation speeds the aggregation state of the dispersion may be affected.” Samples were taken at suitable intervals and the particle concentration was determined using an automatic particle counter which was essentially a laser-illuminated flow ultramicroscope.18~ The samples were withdrawn from the tubes by a 2 cm3 plastic syringe.The shear forces involved in this procedure and in the injection of the samples into the particle counter do not affect the aggregation state of the dispersion (order-of-magnitude variations in injection speed have no effect on the particle count obtained). The errors associated with the counting procedure are close to the statistical prediction nk where n is the number of particles sampled per unit time. Typically the average of ten counts each of ten seconds duration is used in which ca. 1000-2000 particles per 10 s are sampled. This leads to standard deviations of order 2.5%.1676 INTERACTION FORCES IN NON-IONIC SURFACTANTS Rates and equilibria were obtained from plots of reciprocal particle concentration against time.In both types of experiment similar initial particle concentrations were employed throughout [(2-3) x lo8 ~ m - ~ ] . MATERIALS Monodisperse polystyrene latex dispersions were prepared in this laboratory by emulsion polymerisation and were dialysed extensively before use. Particle sizes were determined by electron microscopy or photon correlation spectroscopy. The latex, which in previous work was erroneously assigned a diameter of 1.6 pm, has been reassessed at 1.76 pm (& 10.5%). The other dispersions were more monodisperse, having standard deviations within 5 %. Inorganic reagents were of analytical reagent grade. Octaoxyethylene glycol n-dodecyl monoether (C,,E,) was prepared in this laboratory and was shown by gas-liquid chromatography to be 99% pure.The absence of a minimum in the surface-tension curve indicated the absence of more surface-active impurities. The cloud point of C&, is 78 "C in water and 77 "C in 10-1 mol dm-3 MgSO,, the experimental medium. Water for use with the particle counter was prepared by normal distillation of deionised water from caustic soda/potassium permanganate, followed by a second distillation involving a ' hot spot' to break the water film and minimise particle carryover. This produced water with a very low background particle correction at a counter sensitivity setting appropriate to the smallest particles used (0.25 pm diameter). RESULTS AND DISCUSSION Fig. 2 shows the equilibrium aggregation data of Thompson and Prydel for 0.25, 0.5, 1.0 and 1.76 pm diameter polystyrene latex dispersions stabilised by C,,E,.Excess electrolyte was included to suppress double-layer repulsion. In addition it shows some extra data in the same range of size and temperature together with similar data for 3.8 pm diameter polystyrene particles. The original work indicates that larger particles were less stable to temperature increase than small ones. This was thought to be so because the range of van der Waals attraction increases with size whereas the repulsion due to an adsorbed layer, when viewed as anything like a hard sphere, does not. The residue of attraction at large separations therefore increases with particle size and aggregation is favoured.The new data for dispersions of 3.8 pm particles do not fit this framework in that an increased resistance to complete aggregation was found. Partial aggregation was observed at all temperatures to an extent which bears an unexpected oscillatory temperature dependence (discussed later). The switch to partial aggregation leads to a suspicion that, in terms of fig. 1 (b), the secondary- rather than a primary-minimum mechanism may be involved for these large particles. Further evidence for this emerges from the aggregation-rate data, which are shown in table 1 in the form of the stability ratio W. W is the ratio kR/kexptl, k, is the diffusion-controlled rate constant at the appropriate temperature and kexptl is the experimental rate constant generally determined from the initial slope of the l/n against time curves.Three types of behaviour are apparent. First, the results for the 0.25, 0.5 and 1.0 pm particles indicate the presence of a barrier to flocculation. This implies that aggregation involves a primary minimum. Apart from the 1 .O pm latex, the values of W exhibit little systematic temperature dependence. This appears to be linked to the oscillatory equilibrium behaviour of the 3.8 pm latex and will be dis- cussed later. For the 1.76 and 3.8 pm particles two further modes of behaviour are seen, both of which are typified by an initial rate which is close to the rapid rate. Values for W of up to ca. 2.5 are normally expected for ' rapid ' aggregation because of well known hydrodynamic effects.'* This is in line with the values observed for the 1.76 pm latex.The 3.8 pm latex at low temperature aggregates more quickly, and this may be because it aggregates at relatively greater separations and is therefore free of theseL. THOMPSON 1677 15 35 55 75 TI" C Fig. 2. Plot of equilibrium aggregation state against temperature for C,,E,-stabilised polystyrene latex dispersions in 10-l mol dm-3 MgSO,. 0, 0.25; 0, 0.5; 8, 1.0; V, 1.76 and A, 3.8 pm diameter. no and nco represent initial and equilibrium particle concentrations, respectively. Table 1. Aggregation-rate showing the stability ratio, W, for various temperatures and particle sizes TIT diameter h m 15 20 25 35 40 50 ~~ a a a 46 62 35 0.25 - 40 60 58 55 0.5 1 .o - - 141 118 45 24 1.76 2.1 1 .8b 1.6' 1 .gb 3.2' 1.9 3 . P 0.7 - 0.92 1.8 2.9 2.6 - 31" 64c 406c 72c a Insufficient aggregation to produce reliable data; rates, respectively, where 'type 3' behaviour is observed; 5.5 and 5.2, respectively.and refer to initial and second stage at 60 and 70 "C values of W were effects. In addition there are indications that aggregation of the 3.8 pm latex is slower at high temperature than is allowed for by the simple viscosity correction used here. The observation of diffusion-controlled aggregation indicates that no barrier to flocculation exists, implying that either the secondary minimum becomes important or that the barrier to primary-minimum aggregation has disappeared. The larger particles (3.8 ,urn) exhibit no further complications, and this behaviour will be referred to as 'type 2'.Examination of the details of the aggregation rate curves involving 1.76 pm particles reveals a more complex situation (type 3). After the initial rapid phase (a) in table 1, a period of slower aggregation (b) is observed so that the overall shape of the type 3 curves is quite different to that of type 2. The rate constants quoted under (b) in table 1 were derived from the slope of the appropriate part of the l / n1678 INTERACTION FORCES IN NON-IONIC SURFACTANTS 0 3 1 n E --- 3 v 0. 0.51 I I I I I 0 10 20 30 40 50 tlh Fig. 3. Aggregation-rate data for 1.0 pm diameter polystyrene latex dispersion at 35 "C in 1.86 x lo-' mol dm-3 C,,E, and mol dm-3 MgSO,. (-) Theoretical curve based on initial rate constant, k = 7 x cm3 0, and equilibrium constant [eqn (A 3)], K = 2.31 x cm3.against time curve. For type 3 curves equilibrium may or may not eventually be reached depending on temperature. This suggests concurrent primary (slow) and secondary (rapid) aggregation mechanisms. By extension of the trend from simple primary minimum for the smaller particles through mixed primary/secondary for the 1.76 pm particles, it follows that the largest particles aggregate through a purely secondary-minimum mechanism rather than through a mechanism involving dis- appearance of the barrier to the primary minimum. The likely change of mechanism is supported by the completely different temperature dependence that the 'type 2' data exhibit. Analysis of all of the rate curves was required to confirm this pattern because it was not known what shape the rate curves for reversible aggregation should adopt.(1.e. how quickly does deviation from the initially second-order process occur, and is the apparent difference between types 2 and 3 just a quirk of the kinetics?) For this reason the kinetics of reversible aggregation processes have been analysed (see Appendix) and an expression derived which describes the particle concentration as a function of time. This expression contains the initial rate constant and the equilibrium constant, both of which are determined experimentally. The rate curves predicted by this treatment have been compared with the data. Fig. 3-6 give examples of the three different kinds of behaviour. Fig. 3 shows a 'second-order' plot of reciprocal particle concentration (1 / n ) against time for a dispersion of 1 .O pm diameter particles at 35 "C (type 1).The coincidence of theory and data supports the idea of slow flocculation into a primary minimum of restricted depth, uncomplicated by the presence of a significant secondary minimum. The rate treatment in the Appendix predicts the form of the rate curve remarkably well despite the assumption of non-multibonded, essentially linear aggregation that it contains. This is consistent with the findings of the earlier work' in which the equilibrium treatment [eqn (2)], which contains the same assumption, was found valid to levels of equilibrium aggregation where n,/n, = 2-3, after which it rapidly becomes ineffective. Here n,/n, reaches only ca. 1.5 at equilibrium.L. THOMPSON 1679 Fig.4. Aggregation-rate data for 3.8 pm diameter polystyrene latex at 25 "C in 1.86 x lo-, mol dm-3 C12E8 and 10-1 mol dm-3 MgSO,. (-) Theoretical fit based on initial rate constant, k = 6.7 x 10-l2 cm3 s-l, and equilibrium constant [eqn (A 3)], K = 1.29 x cm3. 0 0 0 0 0 0 100 200 300 400 tlmin Fig. 5. Aggregation of 1.76 pm diameter polystyrene latex in 1.86 x lo-, mol dm-3 C12E, and 10-l mol dm-3 Mg SO, at 40 "C. The line is a theoretical fit for the rapid portion of the curve based on an initial rate constant, k = 2.76 x 10-l2 cm3 s-l, and an estimated equilibrium position, K = 1.31 x cm3. Fig. 4 shows a similar plot for a dispersion of 3.8 pm polystyrene particles at 25 "C (type 2). A good fit is obtained using a rapid initial rate and the observed equilibrium position, supporting the secondary-minimum model.Fig. 5 and 6 show type 3 behaviour, for which it was necessary to propose concurrent primary- and secondary-minimum flocculation. Fig. 5 shows an example in which equilibrium was not achieved and the abrupt transition from the initially rapid aggregation to a clearly defined process about twenty times slower can be seen.1680 INTERACTION FORCES IN NON-IONIC SURFACTANTS 0 5 10 15 ' I 2 5 tlh Fig. 6. Aggregation of 1.76 pm diameter polystyrene latex in 1.86 x lo-, mol dmP3 C,,E, and 10-l mol dm-3 Mg SO, at 20 "C. (---) Theoretical fit based on initial rapid rate, k = 2.0 x 10-l2 cm3 s-l, and observed equilibrium position, K = 1.35 x lo-* cm3. (. . . . .) Theor- etical fit for 'rapid' portion of the curve based on initial rate constant, k = 2.0 x cm3 s-l, and estimated equilibrium position, K = 5.09 x low9 cm3.(-) Theoretical fit for 'slow' portion of the curve based on initial rate constant, k = 1.98 x cm3 s-l, and K = 6.5 x cm3. Initial particle concentration, no, in this case was taken as the intercept of the calculated curve rather than the true value. A theoretical fit for the whole curve cannot be attempted because of the absence of a final equilibrium. If, however, it is assumed that the discontinuity in the curve represents the limit of secondary-minimum aggregation, a theoretical fit for the early time evolution of the process becomes possible and reasonable agreement is obtained. Fig. 6 shows an example in which equilibrium was achieved. The calculated curve based on the initial rapid rate, and the final equilibrium position does not fit the data in that more rapid achievement of equilibrium was predicted than was observed.This confirms that two processes are at work, and (as explained in the Appendix) an approximate fit should be possible if these processes are treated separately. By examination of fig. 6 an estimate can be made of the position of the equilibrium that would have been reached in the absence of the slow primary-minimum process. A combination of the equilibrium constant associated with this position and the initial rapid rate again enables a good fit for the early time evolution of the process to be obtained. An approximate rate constant for the primary-minimum (slow) part of the process was obtained from the slope of the appropriate part of fig.6 and this was used, together with the final position of equilibrium, to model the time evolution of the remainder of the aggregation. Here the fit is poor, and while this fact may contain some message of a fundamental nature it is more likely to be a result of the level of guesswork involved in separating the two processes and will not be pursued. REPULSION MODEL FOR ADSORBED NON-IONIC SURFACTANT In the previous section a qualitative explanation for the data has been established. The aim of this section is to place this interpretation on a more quantitative footing by developing a model for the repulsive interactions between non-ionic stabilised particles. In deriving such a model reasonable allowance must be made for attractiveL.THOMPSON 1681 van der Waals interactions. Here the expressions given by Vincent2O9 21 for particles surrounded by a double sheath of adsorbed material have been used. The Hamaker function of the polystyrene itself was obtained from the Lifshitz computations of Richmond.22 The average value for A,, in the important separation range was 8.2 x J. The double sheath constituting the adsorbed layer contained an inner sheath of pure hydrocarbon (All = 5.8 x J),21 and an outer sheath which was a mixture of polyoxyethylene and water (All = 6.9 x and 3.7 x J, respectively).21 The thickness used for the outer sheath was in general the ‘random- coil’ length of 1.4 nm. The thickness of the inner sheath and the constitution of the outer sheath are determined by the area per molecule.The adsorption data were those obtained in earlier work. An appropriate mathematical form for the repulsion between adsorbed non-ionic surfactant layers is not easily defined because there is no clear understanding of the nature of the repulsion other than the rather nebulous term ‘hydration force’. In a sense this is unimportant because the function of the model is to define the magnitude and separation dependence of the repulsion. This can undoubtedly be achieved by a variety of mathematical forms, but it is useful to apply one which allows ready comparison with the results of other workers. If it is assumed that the forces involved here are those responsible for maintaining a well defined spacing between non-ionic surfactant bilayers in lamellar phase and between other lipid bilayers, then the work of Parsegian et aL9-11 provides a starting point.These workers used osmotic-pressure measurements to obtain the force between lecithin bilayers as a function of their separation. Their results indicated a simple exponential relationship between the pressure, P , and the bilayer separation, d,: P = Po exp (- dw/A). (3) For egg lecithin the value of p0 was 7.05 x lo8 N m-2 and the decay length L was 0.25 nm.ll Note that this type of relationship has been predicted23 through order- parameter considerations. Using Derjaguin’s approximation, eqn (3) can be expressed as an energy between spheres, when the relationship becomes V = zaA2P. (4) Attempts to fit this equation to the aggregation data were unsuccessful because an appropriate combination of the three extrema in fig.1 (b) could not be generated. The model needed to achieve this requires a strong short-range Component to restrict the primary minimum, together with a weaker long-range component to generate a potential-energy barrier. The range of the second component must not, however, be so great that it precludes the presence of a significant secondary minimum with the larger particles. The most obvious way to achieve this is to use the sum of two exponentials so that where h is the separation between the outer edges of the adsorbed layers on adjacent particles. All of the aggregation data, regardless of particle size and temperature, can be fitted to this model using a combination of constants differing only slightly from the following P,, = 5 x lo6 N m-2, PO2 = 2 x lo4 N m-2, A, = 0.2 nm and A2 = 3.0 nm.Despite the longer decay length of 3 nm, these forces are very much weaker than those observed between lecithin bilayer~.~-ll Eqn ( 5 ) is purely empirical, so that it has no theoretical implications concerning the origin of the repulsion. Some observations concerning the origin of the short-range component can, however, be made. They1682 INTERACTION FORCES IN NON-IONIC SURFACTANTS 12 6 Fz, bh * o - 6 -12 0 4 a 12 16 20 separation/nm Fig. 7. Plots of potential energy against separation for different-sized polystyrene particles. The repulsion model [eqn (5)] used the following parameters: Po, = 5 x lo6 N m-2, Al = 0.2 nm and P,, = 2 x lo4 N m-,. A, varied as follows with particle diameter.(a) 3.8 pm, 3.5 nm; (b) 1.76 pm, 3.0 nm; (c) 1.0 pm, 3.0 nm; (d) 0.5 pm, 3.1 nm and (e) 0.25 pm, 3.3 nm. derive from the exclusion from the calculations of an electrostatic repulsion term arising from the native charge of the particle. The original justification for this was that at the high electrolyte concentration used, and in the presence of an adsorbed layer, the range of the electrostatic term was likely to be insignificant. Fitting the aggregation data, however, has produced a model which contains just such a short-range component. It follows that this component contains at least a contribution from the electrostatic term. Indeed, the repulsion arising from a potential of ca. 10 mV (at the outer edge of the adsorbed layer) is approximately equivalent to the above values of Pol and 2,. 10 mV is not an unlikely figure, and the resulting implication that the short-range component may even arise wholly from this source demands an investigation involving electrokinetic measurements of non-ionic stabilised particles.Such an investigation is not trivial in either an experimental or a theoretical sense and it will be considered separately. Fig. 7 shows plots of potential energy against separation calculated for each of the particle sizes used in the experiments. The headgroup area for the adsorbed surfactant (0.65 nm2) refers to 35 "C. The curves have been fitted to within kT [through eqn (1) and (2)] to the kinetic and equilibrium data at this temperature and they are consistent with the mechanisms discussed in the last section.The models used differed only in the values of &, which ranged from 3.0 to 3.5 mm, and in the precise detail of the primary minimum. Fig. 7 does not reflect the restriction placed on the depth of the primary minimum. This is because the balance of forces at the short primary-minimum separations is too model-sensitive for serious curve fitting and we can only establish the feasibility of an appropriately restricted minimum. In fact in all cases shown it is possible to eliminate the primary minimum altogether by a relatively small change in the short-range component through either Pol or L. Both the short-range model sensitivity and the restriction of the primary minimum are better demonstrated when eqn ( 5 ) is used to investigate the temperature sensitivity of the particles (1.76 pm) that exhibit the most complex behaviour.L.THOMPSON 1683 Fig. 8. Plots of potential energy against separation for 1.76 pm diameter polystyrene particles showing the effect of temperature (through area per molecule). (a) 1 nm2 per molecule (15 "C), (b) 0.8 nm2 per molecule (25 "C), (c) 0.55 nm2 per molecule (40 "C). This parameter only affects the attractive component of the potential energy. The repulsion model was that used for fig. 7. Fig. 8 shows the effect on the calculated plot of potential energy against separation for 1.76 pm particles produced when the area/adsorbed molecule is given its experimentally determined temperature dependence, C,,E, adsorption increases significantly between 15 and 35 "C, the area per adsorbed molecule changing from 1 to 0.58 nm2.Above this temperature adsorption remains relatively constant up to the limit of the available data at 50 O C . l This changes the attractive van der Waals part of the interaction potential, because increasing the adsorbed layer's density increases its contribution to the attraction. The repulsion was allowed to remain independent of temperature for the purposes of fig. 8, i.e. the changes in the plot of potential energy against separation shown in fig. 8 are brought about entirely by the effect of temperature on the interparticle attraction. These changes may be summarised as follows. At low temperatures there is no primary minimum, and aggregation is associated with the secondary minimum.As the temperature increases the primary minimum develops and the potential-energy barrier becomes smaller. This is broadly in line with experiment where rapid, weak aggregation at low temperature gives way to slow, weak then slow, strong and finally rapid, strong aggregation at high temperature. This final phase indicates that, in practice, the potential-energy barrier is eliminated completely. Note that although the main features of the transition are predicted, the detailed behaviour is more complex. This is thought to involve changes in the structure of the adsorbed layer which are not dealt with by the simple model used here. This point is discussed later. Fig. 8 contains two implications which are central to an understanding of colloid stabilisation by non-ionic surfactants. First, since it is close to explaining the observed temperature effects solely on the basis of adsorption-induced changes in attraction, it follows that the repulsive contribution is to a first approximation independent of temperature. At first sight it seems contradictory to suggest that the adsorbed layer's repulsion remains constant while the level of adsorption increases, because this apparently involves a decrease in headgroup area which in turn reflects a decrease in1684 INTERACTION FORCES IN NON-IONIC SURFACTANTS headgroup repulsion.In this case, however, it is likely that the change in area per adsorbed molecule is brought about by configurational factors (e.g. flat or vertical orientation of the hydrocarbon chain) rather than by a fundamental change in headgroup repulsion.The phase diagram for C,2E,24 supports this view in that it demonstrates the existence of hexagonal liquid-crystal phase over the whole of temperature range studied here This would not be permitted by the packing requirements of the phase if the headgroup area changed drastically, and indeed X-ray diffraction measurements26 indicate an almost temperature-insensitive headgroup area (0.52 nm2) for this system. Additional evidence for a change in adsorbed-layer configuration is contained in the aggregation data and will be discussed later in this paper. Temperature-independent hydration forces have also been observed between mica ~ h e e t s . ~ ' ~ 28 These forces were attributed to the presence of highly hydrated metal ions, and they followed an exponential decay pattern with a typical decay length of 1 nm compared with the 3 nm found here and the 0.25 nm found in Parsegian's work.When the electrolyte was hydrochloric acid, hydration forces were not observed. It seemed unlikely that the present results could be attributed to ion hydration because there was no apparent mechanism for attaining the required high levels of specific counter-ion adsorption. Nevertheless, aggregation-rate experi- ments have been conducted in 0.4 mol dm-3 HCl, and essentially similar results were obtained to those quoted here for 10-1 mol dm-3 MgSO,. This shows that the long-range component of the repulsion is unaffected. As mentioned earlier the short-range component may be sensitive to zeta potential and hence to pH, but this would significantly affect only the primary minimum equilibrium.Note also that little or no difference was produced by substitution of 0.4 mol dme3 NaCl. The second important implication of fig. 8 is that the temperature effect is basically short-range in nature. This is because adsorbed layers of the kind used here cause only a short-range perturbation of the attraction of the underlying particle. Primary- minimum aggregation equilibria are greatly susceptible, whereas rates are less so because they are controlled by an energy barrier at greater separations. Secondary- minimum aggregation is hardly affected at all. This is why aggregation of the largest (3.8 pm) latex, which according to fig. 7 operates through a secondary-minimum mechanism, is relatively insensitive to temperature in the range in which smaller particles aggregate.Fig. 2 shows that this latex finally aggregates at ca. 70 "C, and this leads to a suspicion that the repulsion forces are not completely independent of temperature. Extension of the calculations to this temperature could not be made in the absence of relevant adsorption data. The temperature sensitivity of the repulsive force may provide an interesting area for future research which may best be tackled using a surfactant which has a less temperature-sensitive area per adsorbed molecule. Finally, attention is drawn to some details of both rate and equilibrium data for which the model resulting in fig. 7 and 8 fails to account despite its success in explaining the main features of the aggregation behaviour.It is not surprising that such details exist, because the attraction expression used employs an idealised view of the adsorbed layer's structure. By examining the difference between the model predictions and the data it is possible to draw qualitative conclusions about adsorbed-layer structure. Fig. 8 predicts that for 1.76 pm particles increasing temperature will cause a progressively decreasing barrier to flocculation together with a slightly increased secondary minimum depth. The predicted effect is similar in direction but smaller in magnitude for the smaller particles. This is because of the different separations relative to particle size at which the potential-energy barriers are situated, Fig. 7 shows that the barriers are at about the same absolute separations, so that those associated with the smaller particles are at greater relative separations and are therefore less sensitiveL.THOMPSON 1685 15 2s 35 4s 55 TI" C Fig. 9. Temperature dependence of the secondary minimum: A, 3.8 pm diameter polystyrene latex (secondary minimum); 0, 1.76 pm polystyrene latex (combined primary and secondary minimum); 0, 1.76 pm polystyrene latex (secondary-minimum contribution). to the short-range temperature effect. The model for 0.25 pm particles predicts a variation in barrier height of only ca. 1.5 W i n the 35-50 "C range. The experimental data in table 1 indicate that the temperature effect on barrier heights is generally less than predicted being negligible for the smaller particles.This and the presence of a maximum in W for the 1.76 pm particles indicates an effective increase in repulsion with temperature relative to the model used. The equilibrium data give rise to a similar conclusion, It is recalled that the equilibrium aggregation state of the 3.8pm latex exhibited an oscillatory temperature dependence with least aggregation at ca. 40 "C. It has been concluded that the mechanism of aggregation involved a secondary minimum whereas the aggregation of 1.76 pm particles involved both primary and secondary minima. When the secondary-minimum part of the latter process is crudely separated out (by inspection of the rate curve) its temperature dependence is seen in fig. 9 to follow the same trend as the larger particles with least aggregation at ca.35 "C. This effect was not immediately apparent for the 1.76 pm particles because it was overshadowed by the highly temperature-sensitive primary-minimum deepening caused by increased area per adsorbed molecule. To explain this observation it is necessary to propose some effect which increases the net repulsion slightly (by 1 or 2 k T ) in the 25-35 "C range. A change in the configuration of the adsorbed layer is one possibility which has the merit of consistency with the discussion earlier in this paper where a difference was noted between the temperature sensitivity of the area per adsorbed molecule and the headgroup area in liquid crystals. The difference implied that a configurational change actually caused the area per molecule variation that gives rise to the temperature sensitivity of the colloid behaviour.The present approach is unable to define changes in the adsorbed layer because a given aggregation effect may arise from a large number of conceivable changes. It can, however, assess the feasibility of likely reorganisations by comparing their calculated consequences with the aggregation data. Of the possible changes a reorientation of the adsorbed hydrocarbon chains from flat to vertical seems the most likely. The1686 INTERACTION FORCES IN NON-IONIC SURFACTANTS consequences of this to the interaction energy are of the right order to explain the experimental results. To illustrate this, the equilibrium level of secondary-minimum aggregation alone has been calculated for 1.76 and 3.8 pm particles, first using the ‘condensed-sheath’ model exactly as defined for fig.7, and secondly substituting a vertically orientated hydrocarbon-chain/water layer of depth I .3 nm.8 In both cases the extended-chain model predicts less aggregation (a larger nln,), so that n/n, = 0.90 and 0.60 for the 1.76 and 3.8 pm dispersions, respectively, compared with values of 0.78 and 0.42 for the condensed-sheath model. APPENDIX KINETICS OF REVERSIBLE AGGREGATION Reversible aggregation can be represented as Pmer +jmw + (P +j)rner* It consists of a series of associations and decompositions involving aggregates of different sizes and morph~logies.~ This situation has recently been discussed by Nir and bent^^^ in terms of the distribution of aggregate types formed and its consequences.A reasonably simple, and for the present systems effective, kinetic model can, however, be extracted if it is assumed that no selective aggregation or dissociation takes place, i.e. that the values of p and j are unimportant. (This assumption is also implicit to the classical Smoluchowski treatment of flocculation kinetics and to the equilibrium treatment used here.) In this circumstance the aggregation rate is given by dn dt -- = k,n2-k-, b where n is the concentration of aggregates of all kinds including singlets, k, and k-, are the forward and reverse rate constants, respectively, and b is the concentration of aggregation ‘bonds’. Where aggregates are not multibonded (as would be the case if they were linear in the ‘string of beads’ sense), b = no-n where no is the original singlet concentration.Eqn (A 1) now becomes dn = Jdt. - I k , n2+ k-, n- k-, no Solution of this integral gives the total aggregate concentration as a function of time, so that where q = kZ,(l +4Kn0) 2k, no + k-,( 1 + 2/ 1 + 4Kn0) 2k, no + k-,( 1 - 2/ 1 + 4Kn0) d = tqt+ln and the aggregation equilibrium constant K = k,/k-,. From eqn (A 2) the time evolution of the aggregation process can be derived if the initial rate constant k, can be determined before aggregate dissociation causes deviation from second-order kinetics and if the equilibrium constant K can be obtained. K can be determined experimentally by measuring initial and equilibrium particle concentration and using the equilibrium form of eqn (A 1): Comparison of eqn (A 3) with eqn (2) shows that the equilibrium constant can be expressed in terms of the second virial coefficient.L. THOMPSON 1687 Where two potential-energy minima separated by an energy barrier are concerned, the situation is more complex in principle but is amenable to some simplification.The aggregation is now represented as (1) Pmer +jmer * (P + A s where s and p refer to secondary and primary bonds, respectively. The rates, R, of processes (1) and (2), respectively, are given by R, = k , n2( 1 - W-l) - k-, b, (A 4) and R, = k , rt2 - k-, 6,. (A 5 ) Step (1) only concerns those collisions not energetic enough to go over the barrier, so that the (1 - W-l) term is necessary to prevent the total forward rate from exceeding the diffusion-controlled rate. The total rate expression analogous to eqn (A 1) now becomes dn -- - k , n2( 1 - W-l) - k-, b, + k , n2 - k-, bp.dt Without some means of expressing b, and b, in terms of the overall aggregate concentration n, a rate equation analogous to eqn (A 2) cannot be derived. Where k, % k,, however, the two processes can be treated independently without too much inaccuracy. 1.e. in the early stages of the aggregation the essentially diffusion-controlled secondary-minimum process dominates since, when t = 0, b = 0, (A 7) dn dt -- - n2 [k,( 1 - W-l) + k,] z k , n2. The primary-minimum aggregation cannot strictly be discussed in terms of a diffusion argument until the secondary-minimum equilibrium has been achieved, because a steady state of diffusion does not exist. After the secondary-minimum process has achieved equilibrium (A 8) and only the slower primary-minimum process need be considered.Even then the application of a simple second-order rate model may appear suspect because of possible effects that the accumulation of particles in the secondary minimum may have on the rate of primary-minimum aggregation. Prieve and Rucken~tein~~* 32 have, however, concluded that accumulation at the secondary minimum does not affect the rate of primary-minimum aggregation. k,( 1 - W-l) n2 = k-, bs L. Thompson and D. N. Pryde, J . Chem. Soc., Faraday Trans. I , 1981,77, 2405. * V. A. Volkov and L. F. Komova, Kolloidn. Zh., 1978, 40, 337. Z. Haq and L. Thompson, Colloid Polym. Sci., 1982, 260, 212. D. H. Napper, J. Colloid Interface Sci., 1970, 32, 106. R. Evans and D. H. Napper, Kolloid Z. Z. Polym., 1973, 251, 329. R. Evans and D. H. Napper, Kolloid Z. Z. Polym., 1973, 251,409. C. Tanford, Y. Nozaki and M. F. Rohde, J. Phys. Chem., 1977,81, 1555. D. M. Le Neveu, R. P. Rand, V. A. Parsegian and D. Gingell, Biophys. J . 1977, 18, 209. lo A. C. Cowley, N. L. Fuller, R. P. Rand and V. A. Parsegian, Biochemistry, 1978, 17, L3163. l 1 V. A. Parsegian, N. L. Fuller and R. P. Rand, Proc. Natl Acad. Sci. USA, 1979, 76, 2750. l 2 B. Vincent, Adv. Colloid Interface Sci., 1974, 4, 193. l 3 P. Richmond and A. L. Smith, J . Chem. Soc., Faraday Trans. 2, 1975,71, 468. l4 C. Cowell and B. Vincent, J . Colloid Interface Sci., 1983, 95, 573. l5 J. Th. G. Overbeek, in Colloid Science, ed. H. R. Kruyt (Elsevier, Amsterdam, 1952), vol. 1, p. 285. l6 L. A. Spielman, J. Colloid Interface Sci., 1970, 33, 562. I’ L. Thompson and A. L. Smith, J. Chem. Soc., Faraday Trans. I , 1981, 77, 557. l8 P. McFadyen and A. L. Smith, J. Colloid Interface Sci., 1973, 45, 573. l9 P. G. Cummins, E. J. Staples, L. Thompson, L. Pope and A. L. Smith, J. Colloid Interface Sci., 1983, ’ D. H. Everett and J. F. Stageman, Faraday Discuss. Chem. Soc., 1978, 65, 230. 92, 189.1688 INTERACTION FORCES IN NON-IONIC SURFACTANTS 2o D. W. J. Osmond, B. Vincent and F. A. Waite, J. Colloid Interface Sci., 1973, 42, 262. 21 B. Vincent, J. Colloid Interface Sci., 1973, 42, 270. 22 P. Richmond, Chem. Ind., 1977, 1, 792. 23 S. Marcelja and N. Radic, Chem. Phys. Lett., 1976, 42, 129. 24 D. J. Mitchell, G. J. T. Tiddy, L. Waring, T. Bostock and M. P. McDonald, J. Chem. SOC., Faraday 25 D. J. Mitchell and B. W. Ninham, J. Chem. SOC., Faraday Trans. 2, 1981, 77, 601. 26 I. G. Lyle, personal communication. 27 R. M. Pashley, J. Colloid Interface Sci., 1981, 80, 153. 28 R. M. Pashley, J. Colloid Interface Sci., 1981, 83, 531. 29 J. Bentz and S. Nir, J. Chem. SOC., Faraday Trans. I , 1981 77, 1249. 30 A. Marmur, J. Colloid Interface Sci., 1979, 72, 41. 31 E. Ruckenstein, J. Colloid Interface Sci., 1978, 66, 531. 32 D. Prieve and E. Ruckenstein, J. Colloid Interface Sci., 1980, 73, 539. Trans. I, 1983, 79, 975. (PAPER 3/653)

ISSN:0300-9599    

DOI:10.1039/F19848001673

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. Sac., Faraday Trans. I, 1984, 80, 1689-1704 Sulphur-35 Radiotracer Studies of the Effect of Hydrogen Sulphide on a Molybdenum Disulphide Catalyst in the Hydrogenation of Buta- I ,3-diene BY KENNETH C. CAMPBELL,* MOHAMMAD L. MIRZA, SAMUEL J. THOMSON AND GEOFFREY WEBB Department of Chemistry, University of Glasgow, Glasgow G12 SQQ, Scotland Received 20th June, 1983 Hydrogenation of buta-l,3-diene on a molybdenum disulphide catalyst has been studied in order to determine the change in the pattern of butene product distribution with catalyst treatment (with hydrogen, a buta-l,3-diene + hydrogen mixture, air, hydrogen sulphide or a thiophene + hydrogen mixture). Catalyst treatment with sulphur compounds (H,S or thiophene + H,) resulted in a butene product distribution closer to thermodynamic equilibrium proportions than for a catalyst treated with hydrogen or air.35S has been used as a tracer in H,S to measure the quantity of H,S necessary to change the catalytic function on fresh and hydrogen-treated catalysts and to observe the extent of sulphur isotopic exchange of gas-phase H,S with MoS,-lattice sulphur and with adsorbed H,S. Evidence is presented for the participation in the hydrogenation reaction of hydrogen present as adsorbed HS groups on the MoS, surface, and the change of catalytic function with H,S treatment has been interpreted in terms of the change in distribution of these HS groups. Extensive use has been made of thiophene in catalytic hydrodesulphurisation studies as a simple but typical heterocyclic sulphur compound. Much information has been obtained from cobalt-molybdenum sulphide systems,l> and results with other mixed sulphide systems show that it is possible to generalise to a large extent from results obtained on cobalt-molybdenum cataly~ts.~ An interesting feature to emerge from these studies comes from a consideration of the generally accepted mechanism for thiophene hydrodesulphurisation put forward by Kolboe and Amberg.* They suggested that the first step in the primary reaction pathway is C-S bond cleavage to form b~ta-1,3-diene.~ This idea opens a vast field of interest in using buta-l,3-diene as a probe, by studying its catalytic hydrogenation on an unsupported molybdenum disulphide catalyst, with particular attention being paid to the isomeric butene products. Investigation of the hydrogenation of butadiene (which term will be used throughout this paper to refer exclusively to buta-l,3-diene) over a molybdenum disulphide catalyst has shown that two distinct types of catalyst surface can be characterised in terms of the pattern of distribution of the various n-butenes in the initial products, typically as shown in table 1, where we have adopted the nomenclature of Wells and coworkers.6.The type-A surface is characteristic of a freshly prepared catalyst which has been subjected to several butadiene hydrogenation reactions. Type-B behaviour resulted when a type-A catalyst was sulphided, either by being treated with hydrogen sulphide or by being allowed to serve as a catalyst for the hydrodesulphurisation of thiophene.The rate of butadiene hydrogenation was very similar on both types of surface. 16891690 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE Table 1. Butene distributions type A type B but- 1 -ene 58% 30% trans-but-2-ene 28% 50% cis-but-2-ene 14% 20% This paper is concerned with a quantitative study in which radioactive sulphur has been used as a tracer in hydrogen sulphide to observe its uptake by the catalyst, its exchange with sulphur in the molybdenum sulphur lattice and exchange between adsorbed hydrogen sulphide and inactive gas-phase hydrogen sulphide. EXPERIMENTAL CATALYST PREPARATION Ammonium thiomolybdate was prepared by passing hydrogen sulphide into a solution of 110 g AnalaR ammonium molybdate in 400 cm3 of 36 wt% ammonia solution at 65 "C.The ammonium thiomolybdate was filtered off, dried and decomposed by being heated to 1000 "C in flowing oxygen-free nitrogen for 48 h. A yellow sublimate, presumed to be sulphur, was formed in the cooler regions of the tubing outside the furnace. The nitrogen flow was maintained while the product cooled to room temperature. Electron diffraction confirmed the product to have the molybdenite (MoS,) structure. The surface area (B.E.T.) was 4.3 m2 g-l. MATERIALS Hydrogen (British Oxygen Co.) was purified by passing it over reduced 5% Pd/W03 catalyst at ambient temperature to remove oxygen, then through anhydrous magnesium perchlorate to remove water. Hydrogen sulphide (B.D.H. Laboratory Chemicals Ltd), 99.7% pure, was degassed and purified by trap to trap distillation.Radioactive [35S]hydrogen sulphide was obtained in 5 mCi samples from Radiochemical Centre, Amersham, with a specific activity of 10 mCi mmol-'(0.37 GBq mmol-l) and was suitably diluted before use with inactive hydrogen sulphide. Buta-1,3-diene (B.D.H. Laboratory Chemicals Ltd), 99% pure, was condensed at 77 K and degassed before use. No impurities were then detectable by gas chromatography. APPARATUS The apparatus was a conventional vacuum system employing a mercury diffusion pump and consisting of a ca. 100 cm3 flat-bottomed reaction vessel in an electrically heated furnace, a mercury manometer, storage vessels for reactants, a spiral trap which could be cooled in liquid nitrogen to separate hydrocarbon products from unreacted hydrogen and sample vessels for transferring these hydrocarbons to a separate gas-chromatography apparatus.Two chromato- graphy columns in series were used, a 5 m column packed with 33% dimethylsulpholane on 30-60 mesh firebrick (to separate n-butane, but-1-ene, trans-but-2-ene, cis-but-2-ene and buta-1,3-diene) and a 9 m column containing 5% S.E. 30 silicone oil on 30-60 mesh firebrick (to resolve trans-but-2-ene and hydrogen sulphide). Helium (35 cm3 min-l) was used as carrier in conjunction with a katharometer detector, and the columns operated at 20 "C. Calibrations were made with standard samples. The effluent from the detector was passed through a flow cell to measure [35S]H,S activity by means of a Mullard MX 168/01 thin end-window Geiger-Muller counter coupled to an I.D.L. Ltd autoscaler-timer frequency ratemeter and potentiometric chart recorder.Counter reproducibility was verified and decay corrections made (35S, ti = 87 d) by frequent calibration with standard [35S]H,S samples.K. C. CAMPBELL, M. L. MIRZA, S. J. THOMSON AND G. WEBB 1691 60 h 50 20 0 10 20 30 LO 50 60 -AP Fig. 1. Variation of butene distribution with extent of reaction on a fresh molybdenum disulphide catalyst at 350 "C: 0, but-1-ene; a, trans-but-2-ene; 0, cis-but-2-ene. PROCEDURE A weighed sample of molybdenum disulphide catalyst (ca. 500 mg) was distributed over the flat bottom of the reaction vessel and evacuated for 30 min during which the temperature was raised to 350_+2 "C. A mixture of butadiene (50+ 1 Torr*) and hydrogen (l50+2 Torr) was admitted (either immediately or after catalyst treatment with H,S etc.) and the progress of hydrogenation observed with the manometer.Products were separated from unreacted hydrogen when the pressure had fallen 10 Torr and examined by gas chromatography. This standardised procedure was adopted in all experiments except where otherwise indicated in the results section. RESULTS HYDROGENATION OF BUTADIENE ON A FRESH CATALYST Five successive butadiene hydrogenation reactions were performed on a fresh sample (0.576 g) of molybdenum disulphide. The product distribution for the first reaction was but-1-ene, 30%; trans-but-2-ene, 44%; and cis-but-2-ene, 26%. During the course of the first three reactions there was then a transition to conventional type-A behaviour (but-1-ene, 58%; trans-but-2-ene, 24%; cis-but-2-ene, 18 %).After the catalyst had achieved this type-A state, further butadiene hydrogenation reactions were performed in which the products were extracted at varying pressure falls. The reactions were carried out in a random order. Within the range of pressure fall up to 30 Torr, the product distribution was found to be constant, independent of the extent of reaction (fig. 1). CATALYST PRETREATMENT WITH HYDROGEN A fresh catalyst (0.576 g) was heated to 350 "C under vacuum and was then treated with two changes of hydrogen (300 Torr) at 350 "C for a total of 4 h. A series of four standard butadiene hydrogenation reactions was carried out. The catalyst showed * 1 Torr = 101 325/760 Pa.1692 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE Table 2.Change of butadiene hydrogenation product distribution with catalyst pretreatment no. of but-1-ene but-2-ene but-2-ene type catalyst treatment results (%) (%> translcis assignment H2 H2S H,, then air H,S, then H, H,S, H,, air H,S, H,, air, H, H,S, H,, air, H,, H,S thiophene + H, thiophene + H2, then N, at 750 "C H,, air, H,S 4 5 8 7 5 4 4 10 3 4 62.6 54.6 27.9 28.4 63.2 61.9 32.9 30.0 63.4 32.7 37.4 45.4 72.1 71.6 36.8 38.1 67.1 70.0 36.6 67.3 1.54 1.18 2.42 2.44 2.34 2.75 2.15 2.47 2.00 1.62 A B reproducible type-A behaviour throughout the series, with the average distribution as shown in table 2, row (a). The rapid transition of the fresh catalyst to type-A behaviour can evidently be brought about either by hydrogen or a hydrogen+butadiene mixture, which suggested that the initial different behaviour might have been due to surface oxygen acquired by exposure to the atmosphere.To investigate this, another sample of catalyst (0.570 g) was treated with hydrogen (400 Torr, 6 h, 350 "C), evacuated (30 min, 350 "C), then exposed to air (45 Torr, 2 h, 350 "C). Five standard butadiene hydrogenation reactions then gave the result shown in table 2, row (b). The but- 1 -ene and trans-but-2-ene showed a typical type-A distribution, except that the trans-but- 2-enelcis-but-2-ene ratio was considerably lower than normal. However, the results give no basis for the belief that surface oxygen caused different selectivity. CATALYST PRETREATMENT WITH HYDROGEN SULPHIDE A fresh sample of catalyst (1 .OO g) was heated under vacuum at 350 "C for 30 min.A large excess of hydrogen sulphide (20 Torr) was admitted, allowed to remain for 1 h, then pumped away and the catalyst evacuated for 30 min. Eight standard butadiene hydrogenation reactions, with analysis of products after 20 Torr pressure fall, showed a consistent type-B behaviour with product distributions as in table 2 (c). The initial rate, however, diminished during the series from 18.2 Torr min-l in the first reaction to 9.0 Torr min-l in reactions 7 and 8. The catalyst was treated with two changes of hydrogen (ca. 200 Torr) at 420 "C for a total period of 10 h, in an attempt to remove the sulphur which had been introduced as hydrogen sulphide. The catalyst was evacuated for 15 min at 420 "C and while the temperature was lowered to 350 "C.A further series of standard reactions revealed [table 2(d)] no change in the type-B behaviour. A marked change was observed in the catalytic function, however, when the catalyst was exposed to air (100 Torr) at 350 "C for 1 h followed by evacuation at 350 "C for 30 min. The catalyst had reverted to type A, and further treatment with hydrogen (200 Torr, 350 "C, 4 h) did not alter this [table 2(e) and cf)]. Re-exposure to hydrogen sulphide (20 Torr, 350 "C, 30 min) followed by evacuation (15 min) restored the type-B behaviour [table 2 ( g ) ] .K. C. CAMPBELL, M. L. MIRZA, S. J. THOMSON AND G. WEBB 1693 0 10 20 30 LO H2S uptake/pmol g-I Fig. 2. Butene product distribution as a function of hydrogen sulphide u>-ake: 0, but-1-ene; a, trans-but-2-ene; 0, cis-but-2-ene. INTRODUCTION OF SULPHUR via THIOPHENE HYDRODESULPHURISATION A fresh sample of catalyst (1.002 g) was treated with two changes of a 1 : 4 mixture of thiophene vapour and hydrogen (250 Torr, 350 O C , 1.5 h).The reaction products were not analysed at this stage. The catalyst was evacuated at 350 "C for 30 min and a series of five standard butadiene reactions performed, with analysis of products after 20 Torr pressure fall. The results, table 2(h), show type-B behaviour the same as if the catalyst had been treated with hydrogen sulphide. An attempt was then made to remove the sulphur (introduced from the thiophene) under more drastic conditions. The catalyst was placed in a vitreous silica tube in a furnace and oxygen-free nitrogen passed over it (1 h at ambient temperature, then 5 h at 750 "C).A pale yellow sublimate, presumed to be sulphur, was observed on the cooler parts of the tube. The nitrogen flow was maintained while the catalyst was cooled completely to ambient temperature; the catalyst was then transferred back to the reaction vessel. A further series of butadiene hydrogenation reactions [table 2 ( j ) ] showed that the catalyst had reverted to type-A behaviour. EFFECT OF HYDROGEN PRESSURE ON PRODUCT DISTRIBUTION Using a hydrogen-treated type-A catalyst a series of butadiene hydrogenation reactions (butadiene pressure 50 Torr) with initial hydrogen pressures varying in the range 50-280 Torr showed no variation in product distribution. Similarly, with a hydrogen sulphide-treated type-B catalyst the product distribution was independent of initial hydrogen pressure over the range 100-250 Torr.TRACER EXPERIMENTS WITH [35S]H2S QUANTITY OF HYDROGEN SULPHIDE TO CHANGE CATALYTIC FUNCTION (i) Fresh molybdenum disulphide catalyst. In these experiments a fresh sample of molybdenum disulphide (0.5704 g) was treated with successive small doses of [35S]H2S and several standard butadiene hydrogenation reactions performed after each dose. It was thus possible to determine the amount of hydrogen sulphide necessary to change1694 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE 0 5 10 15 20 H2 S u&-ike/pmol gzl Fig. 3. Butene product distribution as a function of hydrogen sulphide uptake on catalysts which had been used for butadiene hydrogenation (round symbols) or had been treated with hydrogen at 350 "C (square symbols).(Key as fig. 2.) the catalytic behaviour completely from type A to type B. The results are presented in fig. 2 in terms of the butene distribution relative to the hydrogen sulphide uptake in pmol per g of catalyst. Although 27 pmol g-l was sufficient to change the catalytic behaviour completely from type A to type B, hydrogen sulphide uptake continued beyond this amount. However, doses of hydrogen sulphide up to 27 pmol g-l were completely taken up, but beyond this limit each sample was only partially adsorbed and adsorption may have been pressure dependent. The catalyst was treated with hydrogen (300 Torr, 500 "C, 5 h), evacuated for 15 min, then the temperature was lowered to 350 "C and three successive standard butadiene hydrogenations performed. This was followed by exposure to air (25 Torr, 350 "C, 1 h), 30 min evacuation, then two further standard butadiene hydrogenation reactions which revealed a return to type-A behaviour.No 35S activity was released either in the products of the butadiene hydrogenation reactions or as a result of the treatment of the catalyst with hydrogen or air. It was also observed in the hydrogen sulphide adsorption that, after the adsorption limit had been reached, the specific activity of the hydrogen sulphide which was not taken up by the catalyst remained unchanged. This showed that there was no exchange of sulphur between the catalyst and the gas-phase hydrogen sulphide; i.e. the sulphur in the hydrogen sulphide had not been isotopically diluted by inactive sulphur from the molybdenum disulphide lattice.(ii) Used or hydrogen-treated molybdenum sulphide catalysts. Fig. 3 shows the results of other series of reactions with samples of catalysts which had been used for butadiene hydrogenation (round symbols, 0.577 g catalyst) or had been treated with hydrogen at 350 "C (square symbols, 0.575 g catalyst) before being exposed to hydrogen sulphide. The results are comparable and reveal that the transition to type-B behaviour required less hydrogen sulphide than for the fresh catalyst. Finally, a sample of fresh catalyst (0.578 g) was treated with hydrogen (400 Torr, 350 "C, 6 h), evacuated, then exposed to air (20 Torr, 350 "C, 2 h). Five standard butadiene hydrogenation reactions were performed.9.74 pmol of hydrogen sulphideK. C. CAMPBELL, M. L. MIRZA, S. J. THOMSON AND G . WEBB 1695 Table 3. Interaction of hydrogen sulphide with molybdenum disulphide catalyst expt. 1 expt. 2 expt. 3 weight of MoS,/g [35S]H,S admitted/pmol [35S]H,S taken up/pmol H, returned to gas phase/pmol H, formed: H,S taken up H,S admitted for exchange/pmol H,S remaining in gas phase/pmol relative specific activity exchangeable 35S on catalyst/pmol g-' total H,S uptake at saturation/pmol g-l 35S uptake as percentage of total H,S to saturate /pmol g-l 0.308 47.7 44.5 144.2 11.5 13.3 13.3 30.1 144.2 100 0.26 0.41 0.598 0.471 39.7 65.8 39.7 65.8 66.3 139.7 14.1 18.7 45.3 11.5 12.1 122.8 54.0 0.36 0.28 0.27 were then admitted, of which 8.85 pmol were taken up (15.3 pmol g-l) and a further four standard butadiene hydrogenation reactions carried out ; the results are shown in table 2(k).The feature of only partial uptake of hydrogen sulphide, after the transition to type-B behaviour was complete, was observed with these catalysts also. HYDROGEN SULPHIDE ADSORPTION AND ISOTOPIC EXCHANGE [35S]hydrogen sulphide was used as a tracer to study (i) the limit to the extent of hydrogen sulphide uptake by the catalyst, (ii) the amount of hydrogen produced as a result of this hydrogen sulphide uptake and (iii) the extent to which sulphur introduced to the surface by the adsorption of hydrogen sulphide could be exchanged with gas-phase hydrogen sulphide. In experiment 1 a fresh sample of catalyst was heated to 350 "C and kept at this temperature under vacuum for 1 h.It was exposed to an excess of [35S]hydrogen sulphide for 2 h and the gas-phase material was transferred quantitatively to the gas chromatograph for analysis. Non-radioactive hydrogen sulphide was then introduced into the catalyst vessel to investigate whether exchange would occur with the 35S now present on the surface from the adsorbed [35S]H,S. The hydrogen sulphide was allowed to remain in the vessel for 30 min before it was removed for analysis, and its specific activity was then determined relative to that of the original [35S]hydrogen sulphide. The results are summarised in table 3. Experiment 2 was conducted under similar conditions to experiment 1 except that the amount of [35S]hydrogen sulphide was deliberately chosen to be only about half that necessary to saturate the surface.After allowing interaction to take place for 2 h, non-radioactive hydrogen sulphide was admitted in sufficient quantity to saturate the surface and leave a small excess in the gas phase which could be examined for exchanged radioactive sulphur. These results also are summarised in table 3. Table 3 also shows the results of experiment 3 which was conducted to verify the amount of hydrogen produced by the interaction of hydrogen sulphide with the catalyst under conditions of near-saturation with hydrogen sulphide. The investigation as in experiment 1 was then conducted on a sample of molybdenum disulphide which had been pretreated with hydrogen. A fresh sample of catalyst was kept at 350 "C under vacuum for 1 h, then a measured amount of hydrogen was1696 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE Table 4.Hydrogen sulphide uptake on molybdenum disulphide pretreated with hydrogen weight of catalyst/g 0.758 1 hydrogen uptake/pmol 40.70 /pmol g-l 53.70 hydrogen sulphide admitted/pmol 99.29 hydrogen sulphide uptake pmol 24.96 /pmol g-' 32.9 admitted and allowed to remain in contact with the catalyst at 350 "C for 1.5 h. After the hydrogen uptake had been measured, the catalyst vessel was evacuated for 30 min and a measured amount of hydrogen sulphide was introduced and left for 2 h at 350 "C. Table 4 shows the results which were obtained. BUTADIENE HYDROGENATION REACTION RATES Butadiene hydrogenation on a fresh catalyst proceeded typically with an initial rate in the range 3-10 Torr min-l per g of catalyst.Treatments with hydrogen, air or hydrogen sulphide all increased this rate (but only by a factor of 2 or less). In the experiments in which the catalyst was progressively treated with several small doses of hydrogen sulphide, the gradual change in the butene distribution was accompanied by a gradual change in the hydrogenation activity. When at a certain stage of sulphur uptake no further modification in the butene distribution was observed, then at the same stage the hydrogenation activity also ceased to increase. Thereafter, on all catalysts the rate diminished gradually with the number of butadiene hydrogenation reactions performed, and this may possibly be attributable to the accumulation of carbonaceous material on the surface.Some reactivation could be brought about by treatment with hydrogen at 500 O C , but in general the activity could not be fully restored to its original value. To obtain some insight into the reliability of the initial rate measurements, a series of butadiene hydrogenation reactions was performed at various temperatures in the range 3 14-424 "C on a run-in sample of molybdenum disulphide which showed type-A behaviour. An Arrhenius plot, using the initial rates measured from pressure against time plots, gave 48.5 3.0 kJ mol-l for the energy of activation. Using this value it can readily be shown that between the limits of the temperature control in standard butadiene hydrogenation reactions (350 f 2.0 "C) a 6% variation in reaction rate would be expected.In view of the uncertainties imposed by temperature control, surface deposits and the difficulty of measuring initial rates with good accuracy from pressure against time plots, we do not attempt to attach any significance to the rather small changes in rate produced by the various catalyst treatments. However, there is certainly no marked change in hydrogenation activity accompanying the change from type-A to type-B product distributions. DISCUSSION PRODUCT DISTRIBUTIONS IN RELATION TO CATALYST TREATMENT Apart from the anomalous behaviour during the first three hydrogenation reactions performed on a fresh molybdenum disulphide catalyst, the product distributions conformed to a clear pattern. Catalysts which had been treated with sulphur showed a type-B product distribution, while those which had not gave a type-A distribution.K.C. CAMPBELL, M. L. MIRZA, S. J . THOMSON AND G. WEBB 1697 Table 5. Distribution of butene isomers but-1 -ene trans-but-2-ene cis-but-2-ene (%> (%> (%I observed (fig. 1) 21 47 calculateda 16 52 32 32 a Data from ref. (9). It was immaterial whether the sulphur was introduced by the adsorption of hydrogen sulphide, or by allowing the hydrodesulphurisation of thiophene to take place on the catalyst. Once type-B behaviour had been established by such sulphur treatment, the catalyst could be made to revert to type A by treatment with air at 350 "C or by heating it to 750 "C in a stream of oxygen-free nitrogen. Removal of elemental sulphur was evident at 750 "C. Air at 350 "C possibly removed sulphur as sulphur dioxide, although none was detected in the gas phase in the 35S-tracer experiments.A similar effect of air, but in the opposite sense (favouring but-2-ene formation), has been observed in the hydrogenation of butadiene on a nickel phosphide catalyst.* The results shown in fig. 1 reveal that the hydrogenation of butadiene on type-A molybdenum disulphide proceeds, during the course of a single reaction, with marked preference for but-1-ene formation in the early stages, but that the distribution later changes abruptly to acomposition close to the thermodynamic equilibrium proportions calculated for 350 0C,9 as shown in table 5. For the standard hydrogenation mixture used (50 Torr butadiene + 150 Torr hydrogen), a pressure fall of 50 Torr represents complete conversion of butadiene to butenes, so that the completion of the change in the product distribution corresponds to the stage in the reaction when nearly all of the butadiene has been used up.Similarly, significant amounts of n-butane were formed only when the total pressure had fallen by 45 Torr or more. Evidently the sites of activity for butadiene hydrogenation are also active for isomerisation and hydrogenation of butenes, but when both are present the butadiene competes favourably for the available sites and effectively prevents re-equilibration or further hydrogenation of the butene isomers. However, but-2-enes do occur in the initial products and the but- l-ene/but-Zene ratio is independent of the extent of reaction in the pressure-falLrange 0-25 Torr.The but-2-enes are therefore not formed by the sequence butadiene -+ but-1-ene --+ but- 2-enes, but are formed simultaneously with but-I-ene. This view is supported by the results of Wells and co-workers,1° who found identical deuterium distributions in the three butene products when butadiene reacted with deuterium over MoS, at 573-623 K. It is also consistent with our finding that the product distributions are independent of initial hydrogen pressure, which would not be the case if butadiene hydrogenation to butenes and butene isomerisation occurred simultaneously. We conclude that discrimination is in favour of but- 1-ene, and that thermodynamic equilibrium of the butenes can only occur when the amount of butadiene remaining has decreased sufficiently for significant butene readsorption to become possible. Hydrogen treatment was without effect in changing the function of either the type-A or type-B catalyst: in particular it did not interact with a type-B catalyst to form hydrogen sulphide and cause it to revert to type-A behaviour. The fresh molybdenum sulphide catalyst therefore cannot be a true sulphur-rich type-B catalyst because the change to type-A behaviour could be brought about at 350 "C by treatment with 56 FAR 11698 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE Fig.4. Disposition of sulphur atoms about a molybdenum atom. 5-S --- -- I 1 s (a) (b) Fig, 5. Molybdenum disulphide adsorption sites: (a) B or 2M site; (b) C or 3M site. hydrogen alone. Although on the fresh catalyst the but-1-ene: but-2-ene ratio was similar to type-B behaviour, the trans:cis ratio for the but-2-enes (1.69) was much nearer to the thermodynamic equilibrium value (1.63) than for typical type-B behaviour.Thus it seems likely that the fresh catalyst is really showing type-A behaviour but with subsequent isomerisation of the but-1-ene, whilst on the run-in catalyst this isomerisation has become so inhibited as to be negligible. CATALYST STRUCTURE The hydrogenation of olefins on molybdenum disulphide has been extensively studied by Tanaka and Okuharall and has recently been reviewed.12 To explain the different function of molybdenum disulphide catalysts after treatment with hydrogen sulphide we begin with their model, which was developed from the concept of coordinative unsaturation introduced by Siegel.13 We introduce the refinement of making a distinction between sulphur atoms coordinated to three molybdenum atoms (as in the bulk of the material) and those coordinated to only two or one (as on exposed edges in the layer structure).Using the symbol X to indicate a fully coordinated sulphur atom and S to indicate an unsaturated sulphur atom, the disposition of sulphur atoms about any molybdenum atom in the bulk can be represented as in fig. 4(a), but a six-coordinate molybdenum on an edge, with exposed face represented by the shaded area, will be as in fig. 4(b). Some of these S-sulphur positions are likely to be unoccupied. Two vacancies will be common and will constitute a B or 2M site, while three vacancies, as at a corner, will result in a C or 3M site (fig.5). Sulphur atoms in X situations need not be considered so it is convenient to indicate the reactive sites as a plane view of the exposed face, as in the lower part of fig. 5.K. C. CAMPBELL, M. L. MIRZA, S. J. THOMSON AND G . WEBB 1699 J--i H S Scheme 1. INTERACTIONS OF H, AND H,S WITH THE CATALYST We now make the postulate that the S-sulphur atoms are capable of bonding to hydrogen atoms and that hydrogen in such SH groups is reactive, as we have shown in an earlier p~b1ication.l~ It may be that the circumstances under which this can occur demand the presence of vacant hydrogen adsorption sites at molybdenum atoms, e.g. S S HS S and this is implicit in the interpretation of earlier work where the presence of ,MH or 3MH sites is invoked13 if these sites are formed from dissociated H, molecules.Treatment of the catalyst with hydrogen sulphide provides an alternative means of producing HS groups on the surface, and the interaction of H, with ,M and 3M sites on H,S-treated and untreated catalysts is compared in scheme 1 (the sulphur atoms originating from MoS, are ringed for clarity, but are chemically identical with those introduced from hydrogen sulphide). Points of contrast between H, and H,S adsorption (in equivalent amounts) are that H,S increases the population of HS groups on the surface more so than H,, and that H,S reduces the coordinative unsaturation with respect to sulphur of the molybdenum by introducing another sulphur atom. Since olefin hydrogenation has been inferred to occur on 3M sites,15-18 hydrogen sulphide would be expected to act as a catalyst poison by conversion of these sites to ,M or 'M, and this has indeed been found16 to be the case for ethylene hydrogenation at ambient temperatures : ethylene hydrogenation and ethylene-D, exchange were negligible after the catalyst had been exposed to hydrogen sulphide.In contrast, for the hydrogenation of butadiene at the higher temperature used in our 56-21700 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE P S ' butadiene , + but -1 -ene HS H Scheme 2. work, we found no poisoning (rather a slight increase in rate) when the catalyst was exposed to hydrogen sulphide. Two other points of contrast were apparent between our high-temperature work and the published ambient-temperature results for butadiene hydrogenation.First, at room temperature the product was found to be but-1-ene with nearly 100% selectivity until all the butadiene had reacted: only then did isomerisation to but-2-enes and hydrogenation to n-butane commence.18 However, at 350°C a yield of 40-50% but-2-ene (fig. l), the typical type-A behaviour, was observed from the beginning of the reaction. Secondly, at ambient temperature, maintenance of the molecular identity of the hydrogen was observed, i.e. [,H,]but-l-ene from D, and [,Hl]but-l-ene from HD as the sole products.ls9 l9 Deuterium, introduced to the catalyst as D,S, became more widely distributed than this when a butadiene + hydrogen mixture was allowed to interact on the catalyst at 350 OC.14 TYPE-A AND TYPE-B BEHAVIOUR : THE REACTION MECHANISM All of the foregoing observations are evidence for the greater lability of reactive hydrogen at 350 "C, and in particular we believe14 that hydrogen in HS groups takes part in the reaction.As a consequence, on our evacuated MoS, catalyst, as well as 1, 2 addition of hydrogen to give but-1-ene, some 1,4 addition to form but-2-enes can occur by the interaction of an adsorbed isobutenyl intermediate on a 3M site with hydrogen on the more distant HS groups of adjacent ,M sites, as in schemes 2 and 3. Formation of both cis and trans isomers of but-2-ene is permitted by rotation about the C2 to C3 bond in the isobutenyl adsorbed intermediate. Under the standard conditions used in our work (10-20% conversion), once the butene products have been desorbed, they cannot be readsorbed owing to competition from further butadiene adsorption ; hence a steady product distribution independent of conversion over the limits observed is possible even though this is markedly different from the thermo- dynamic equilibrium ratios for butenes.On the hydrogen sulphide-treated catalyst the possibility of 1,4 addition is considerably greater owing to the denser population of HS groups on the surface, hence the typical type-B behaviour showing amounts of but-2-enes nearer to thermodynamic equilibrium. This reactivity of hydrogen in a different adsorption situation is analogous to the participation2O9 21 of support hydrogen in hydrogenation reactions on alumina- or silica-supported metal catalysts. In the present case the two types of site are much more intimately mixed, which should facilitate the transfer processes.If under workingK. C. CAMPBELL, M. L. MIRZA, S. J . 'I'HOMSON AND G. WEBB 1701 S m H, H I fi ,SH s s S S H H S *M %l > + but -2-ene Scheme 3. conditions the S-H bond is weaker than the Mo-H adsorption bond, this raises the interesting possibility that reactions which involve transfer of hydrogen atoms will show a preference for one of these adsorption modes over the other. For example, deuterium exchange or isomerisation of but- 1 -ene proceeding via an isobutenyl intermediate might take up a hydrogen atom from an SH group and subsequently lose it, or another hydrogen atom, to form a Mo-H bond. The continuing activity of the catalyst would then be sustained by transfer of hydrogen atoms from Mo--H sites to S-H groups, and the activation energy necessary might be such that this would be the rate-determining step of the overall process.This is a possible explanation of why we observe large changes in product distribution without any significant change in reaction rate. To explain the lack of a poisoning effect by hydrogen sulphide in spite of the reduction in coordinative unsaturation of the molybdenum, we think it likely that sulphur which is not coordinated to three molybdenum atoms is sufficiently loosely held at 350 "C for the butadiene to compete adequately for adsorption sites; i.e. butadiene is able to create 3M sites by promoting the diffusion of sulphur. We envisage this diffusion as taking place between neighbouring sites rather than as an extensive, long-range mobility, because this is ruled out by the exchange results to be discussed later.There is considerable evidence in support of this. (1) On the basis of experimental observations a similar process has been 23 to explain the hydrogenation of acetylene on a sulphided nickel catalyst at 119 "C. (2) At 500-800 "C molybdenum disulphide is an effective catalyst for the decomposition of hydrogen s ~ l p h i d e ~ ~ to hydrogen and elemental sulphur. (3) Even at 350 "C we have observed hydrogen to be produced slowly as a consequence of hydrogen sulphide adsorption on molybdenum disulphide (table 3). (4) At 350 "C gas-phase hydrogen sulphide will exchange with radioactive sulphur previously taken up on the surface as H,S to the extent of ca.30 pmol g-l (table 3), and this amount is in excess of the amount required to change the catalyst function from type A to type B. This is discussed in the next section in relation to the hydrogen sulphide adsorption and exchange results. HYDROGEN SULPHIDE ADSORPTION AND EXCHANGE A sample of molybdenum disulphide which had attained the steady type-A state (either by being used as a butadiene hydrogenation catalyst or by treatment with hydrogen) showed a limit to its hydrogen sulphide uptake of 32.9 pmol g-l (table 4). In contrast, much more hydrogen sulphide was taken up by a fresh catalyst1702 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE (> 120 pmol g-l, table 3). Hydrogen sulphide uptake is thus markedly reduced by the preadsorption of hydrogen, which lends support to the idea of heterolytic dissociative adsorption H2S+H + SH for which there is ample other e ~ i d e n c e .~ ~ - ~ ~ On the used catalyst ca. 15 pmol g-l was sufficient hydrogen sulphide to bring about the change of function from type A to type B (fig. 3), whereas ca. 27 pmol g-l was required for a fresh catalyst: the first 8 pmol g-l of this was taken up by the catalyst before it had been exposed to the reaction mixture containing hydrogen. We think it likely that the extra hydrogen sulphide necessary to bring about the change in function on the fresh catalyst is dispersed over sites which are not available on the run-in catalyst. The fresh and the used catalysts have the common feature that further hydrogen sulphide samples were not taken up completely once adsorption had proceeded beyond the limit necessary to produce a complete change in the catalytic function.This, together with the fact that < 25% of the sulphur introduced to the catalyst as [35S]H2S is exchangeable with gas-phase hydrogen sulphide, suggests that there is more than one type of hydrogen sulphide adsorption. Comparison of the results of isotopic tracer experiments 1 and 2 gives further insight into this. If the non-exchangeable sites were occupied preferentially as the available surface was progressively filled with H2S molecules, no exchange would have been expected in experiment 2, because only 54% of the adsorptive capacity had been satisfied when the last of the [35S]H2S was taken up. In fact 3.10 pmol of [35S]H2S appeared in the gas phase in the exchange, so a strict sequential occupation of sites in order of adsorption strength does not occur.It is therefore interesting to compare the behaviour of the radioactive and non-radioactive portions of the hydrogen sulphide adsorption in experiment 2 as in the following calculations. Experiment 1 .-Surface saturated with [35,S'lH2S. Let x pmol g-l be exchangeable surface sulphur : * * H2S admitted for exchange = 13.32 pmol = 43.25 pmol g-l (catalyst). Observed relative specific activity = 0.41 x = 0.41 4 3 . 2 5 - k ~ whence inactive H2S and an excess of H2S added for exchange. x = 30.05 pmol g-l. Experiment 2.-Surface partly covered with [35sJH2S, coverage compIeted with Excess of H2S for exchange = 11.5 pmol = 19.23 pmol g-l catalyst.Let y pmol g-l of H2S introduced as [35S]H2S, and z pmol g-l of that introduced as non-radioactive H2S be exchangeable. Observed relative specific activity = 0.27 = 0.27. Y 19.23 + y + zK. C. CAMPBELL, M. L. MIRZA, S. J. THOMSON AND G. WEBB 1703 By comparison with experiment 1, assuming the exchangeable portion is proportional to the total uptake of H2S, y+z x 122.8 144.2 -=- whence whence and y + z = 25.59 = 0.27 Y 19.23+25.59 y = 12.10 pmol g-l z = 13.49 pmol g-l = 0.54 39.7 [35S]H2S adsorbed - total H,S adsorbed - 39.7+45.3 - 11.5 -- [ 35 S]H 2S exchanged total H2S exchanged - y + z = 0.47. These ratios are so similar that there is no clear evidence that there was any significant difference in the ease of exchange of hydrogen sulphide adsorbed at different coverages.It is inferred that hydrogen sulphide must become adsorbed at the first vacant site which the molecule encounters and is not extensively mobile on the surface. Accordingly, we conclude that in the surface SH groups only the hydrogen is capable of migration on the surface. We are grateful to the Ministry of Education of Pakistan for a Central Overseas Training Scheme Award (to M. L. M.). 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 F. E. Massoth, Adv. Catal., 1978, 27, 265. C. H. Amberg, in Proc. Climax Third Znt. Con$ on the Chemistry and Uses of Molybdenum (Climax Molybdenum Co., Ann Arbour, 1980), p. 180. P. Grange, Catal. Rev., 1980, 21, 135. S. Kolboe and C. H. Amberg, Can. J . Chem., 1966,44, 2623. P. J. Owens and C. H. Amberg, Ado. Chem. Ser., 1961, 33, 182. M. George, R. B. Moyes, D. Ramanarao and P. B. Wells, J. Catal., 1978, 52, 486. J. J. Phillipson, P. B. Wells and G. R. Wilson, J. Chem. SOC. A, 1969, 1351. F. Nozaki and R. Adachi, J. Catal., 1975, 40, 166. F. D. Rossini et al., Selected Values of Physical and Thermodynamic Properties of Hydrocarbons and Related Compounds (American Petroleum Inst. Research Project 44, Carnegie Press, Pittsburgh, Pa, 1953), p. 737. M. R. Blake, M. Eyre, R. B. Moyes and P. B. Wells, Proc. 7th Znt. Congr. Catal., Tokyo, 1980 (Elsevier, Amsterdam, 1981), p. 591. K. Tanaka and T. Okuhara, Catal. Rev. Sci. Eng., 1977, 15, 249. P. C. H. Mitchell, in Catalysis (Specialist Periodical Report, The Royal Society of Chemistry, London, 1981), vol. 4, p. 175. S. Siegel, J. Catal., 1973, 30, 139. J. Barbour and K. C. Campbell, J. Chem. SOC., Chem. Commun., 1982, 1371. K. Tanaka, T. Okuhara, S. Sat0 and K. Miyahara, J. Catal., 1976, 43, 360. T. Okuhara, K. Tanaka and K. Miyahara, J. Catal., 1977, 48, 229. T. Okuhara, T. Kondo, K. Tanaka and K. Miyahara, J. Phys. Chem., 1977,81,90. T. Okuhara and K. Tanaka, J. Chem. SOC., Faraday Trans. I , 1979, 75, 1403. T. Okuhara, T. Kondo and K. Tanaka, Chem. Lett., 1976, 717. J. A. Altham and G. Webb, J . Catal., 1970, 18, 133. K. C. Campbell and J. Mooney, J . Chem. Soc., Faraday Trans. 1 , 1980, 76, 2332. A. Takeuchi, K. Tanaka and K. Miyahara, Chem. Lett., 1974, 171.1704 RADIOTRACER STUDIES ON MOLYBDENUM DISULPHIDE 23 A. Takeuchi, K. Tanaka, 1. Toyoshima and K. Miyahara, J . Catal., 1975, 40, 94. 24 K. Fukuda, M. Dokiya, T. Kameyama and Y. Kotera, Ind. Eng. Chem. Fundam., 1978, 17, 243. 25 F. E. Massoth, J. Catul., 1975, 36, 164. 26 C. J. Wright, S. Sampson, D. Fraser, R. B. Moyes, P. B. Wells and C. hekel, J. Chem. SOC., Faraday 27 P. Ratnasamy and J. J. Fripiat, Trans. Faraduy SOC., 1970, 66, 2897. Trans. 1, 1980, 76, 1585. (PAPER 3/ 1042)

ISSN:0300-9599    

DOI:10.1039/F19848001689

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984,80, 1705-1 7 15 Metachromasy in Clay Minerals Sorption of Acridine Orange by Montmorillonite BY R. COHEN AND S. YARIV* Department of Inorganic and Analytical Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel Received 8th July, 1983 The adsorption of acridine orange by H-, Na-, Mg-, Al- and Cu-montmorillonite has been studied by visible and infrared spectroscopy and by X-ray diffraction methods. Adsorption of the dye takes place by the mechanism of cation exchange. The maximum amount of adsorbed dye depends on the exchangeable metallic cation, decreasing in the order Na > H > Cu > Mg > Al. The adsorbed cationic dye is located in the interlayer space. Adsorption leads to metachromasy of the dye molecule and a shift of the absorption bands to lower wavelengths in the visible range.Two types of associations are formed between montmorillonite and acridine orange: (i) a monolayer of the dye situated in the interlayer space, with the aromatic rings parallel to the aluminosilicate layer and (ii) a bilayer in the interlayer space or tilting of the cationic dye relative to the aluminosilicate layer. In the association of type (i) metachromasy cannot be attributed to polymerization of the organic dye but must be caused by interactions between the oxygen plane of the aluminosilicate and the aromatic dye. The spectra of aqueous solutions of several cationic dyes are dependent on the concentration of the dye.l In the visible region these spectra are characterized by several maxima, the intensities of which depend on the dye concentration.The absorption band of the dye at the longer wavelength is characteristic for dilute solutions and is usually assigned as the a band. Increasing dye concentrations results in the gradual replacement of the a band by a band with a shorter wavelength which is assigned as the Further increase in the dye concentration may cause substitution of the /? band by another diffuse band closer to the blue range of the spectrum. This band is assigned as the y band. The a band, is attributed to the monomeric form of the dye while the p and y bands are attributed to dimeric and higher aggregates of the dye, respectively.6 The dye aggregation is caused mainly by the interaction between the n electrons of the aromatic rings of the dye cations.This spectroscopic phenomenon is called metachromasy and cationic dyes which show this effect are known as metachromic dyes.’ Similar changes in the spectra of the aqueous dye solutions may also be obtained by the addition of salts. This has been interpreted as being due to the ‘salting out’ of the Addition of polyelectrolytes to a dilute dye solution also causes changes in the spectrum which are similar to those obtained by increasing the concentration of the dye. The latter has recently been widely studied’ and it is now accepted that the cationic dyes are adsorbed by polyelectrolytes. Since the concentration of the dye is higher on the polyelectrolyte surface than in the bulk solution the absorbed dye species form dimers and higher aggregates at the adsorption sites, thus giving rise to the metachromic effects in the spectrum.In a previous publication the adsorption of methylene blue by montmorillonite was 17051706 ADSORPTION OF DYE ON MONTMORILLONITE described.12 This adsorption takes place by a cation-exchange mechanism. In the above system it was shown that metachromasy occurs even when the interlayer space of montmorillonite is not large enough to account for the aggregation of dye molecules. For that specific system it was concluded that metachromasy was caused by n interactions between the dye cations and the oxygen plane of the aluminosilicate layer of the montmorillonite. In that investigation12 only the behaviour of a few metallic mono-ionic montmorillonites were studied and of these only copper mont- morillonite was studied in detail.However, a more basic investigation is necessary to confirm the existence of interactions between the oxygen plane of the mineral and the n orbitals of the dye cations in order to explain metachromasy in clay minerals. In the present study the adsorption of another dye cation, namely acridine orange (AO), by montmorillonites pretreated with various inorganic cations was investigated. The cations Na+, H+, Mg2+ and A13+ were chosen to represent mono-, di- and tri- valent cations common in nature. In addition, one transition element (Cu) was also chosen. Very little work has been done so far on the effect of montmorillonite on the spectrum of AO. Recently Yamagishi and Soma13 studied the adsorption of N-alkylated A 0 by Na-montmorillonite. In their study they showed that adsorption of cationic dyes which belong to the acridine orange family also results in meta- chromasy.However, they considered the clay mineral as a normal polyelectrolyte and believed that the interaction between the negatively charged mineral and the dye cation is a pure electrostatic attraction and that metachromasy resulted from the electronic interaction between neighbouring adsorbed cations. In the light of the findings obtained from the study of the adsorption of methylene blue by mont- morillonite it seems necessary to clarify in detail the mechanism of metachromasy of A 0 after its adsorption by montmorillonite. acridine orange (AO) EXPERIMENTAL MATERIALS Wyoming bentonite (montmorillonite) was supplied by Wards and ground to 80 mesh.Cation-exchanged montmorillonites were prepared by equilibrating 1 g natural clay with 100 cm3 0.02 normal chloride solution of the corresponding cation. The excess of salt was washed out after 24 h.14 In addition to the principal cation, the montmorillonites contain other exchangeable cations. For example, the following cations were released into an aqueous solution from Cu-montmorillonite treated with 1 mol dm-3 LiC1: Na = 8.7, K = 22.0, Ca = 10.0 and Cu = 77.0 mequiv. per 100 g clay. By this method stable suspensions were obtained which are suitable for spectrophotometric determinations. Acridine orange was supplied by B.D.H. Inorganic chemicals were Baker analytical reagents. METHODS ABSORBANCE MEASUREMENTS Spectra of aqueous suspensions of the various mono-ionic montmorillonites in the concen- tration range 0.01-0.0008 wt % in the presence of various concentrations of acridine orange were studied using a Cary 14 double-beam spectrometer in the 700-300 nm range.The optical path in most determinations was 10 mm. Aliquots of the various suspensions were separated by centrifuging and the spectra of the supernatants were also recorded.R. COHEN AND S. YARIV 1707 CATION-EXCHANGE DETERMINATIONS Analysis of the metallic cations in the supernatants of the montmorillonites treated with A 0 were carried out by atomic absorption (Perkin-Eln-~ model 403). INFRARED AND X-RAY MEASUREMENTS Oriented samples of AO-treated montmorillonites were prepared by drying suspensions of cation-exchanged montmorillonites to which various amounts of A 0 had been added.The supporters for these samples were polyethylene films and glass slides for infrared and X-ray measurements, respectively. Infrared spectra were recorded on a Perkin-Elmer 457 grating spectrophotometer. X-ray diffraction patterns were obtained on a Philips diffractometer using monochromatic Cu Ka radiation. Diffractions were determined after leaving the samples for 7 days in a vacuum desiccator containing Mg(N0,),.6H20 (40% humidity) at room temperature and after 3 h in a vacuum oven at 150 "C. RESULTS AND DISCUSSION The adsorption of A 0 by montmorillonite takes place largely by a cation-exchange mechanism, Exchangeable metallic cations are released into the solution when the clay is equilibrated with AO.Some representative data are presented in tables 1 and 2. From table 2 it may be seen that up to a certain concentration of A 0 no exchangeable aluminium is released into the equilibrium solution from Al- montmorillonite and exchangeable copper is released into the solution from Cu- montmorillonite only to a small extent. Cation exchange takes place between A 0 and other metallic cations which are still present in the exchange positions of the clay (such as Na+, K+, Mg2+ and Ca2+) even after the preparation of the monoionic samples. ABSORPTION SPECTRUM OF ACRIDINE ORANGE IN AQUEOUS SOLUTION Metachromasy of A 0 in aqueous solutions has been previously described.l? 3 9 4 f l5 The effect of concentration on the absorption spectrum of A 0 in aqueous solutions is shown in fig 1.The most dilute solution gives an absorption with a maximum at 492 nm and a shoulder at 470 nm, assigned as bands a and p, respectively. More concentrated dye solutions demonstrate metachromasy, i.e. show a change in the relative absorbances of the two bands, the intensity of the a band becoming higher than that of the /l band. A further increase in the dye concentration results in the displacement of the /? band to wavelengths below 470 nm while the a band remains at 492 nm. EFFECT OF MONTMORILLONITE ON THE ABSORPTION SPECTRUM OF ACRIDINE ORANGE The intensities of the absorption bands decrease with the ageing period of the suspensions. This is associated with the flocculation of the clay particles and is observed mainly in those batches where flocculation is high.Absorbance values which are included in the following figures were obtained 15 min after the preparation of the clay-dye suspensions. However, ageing had no effect on the location of the a and bands. By comparison between fig. 1 and 2 it is obvious that the adsorption of A 0 by montmorillonite results in the combination of two spectroscopic effects. These are (i) metachromasy and (ii) partial extinction of the absorption bands. Metachromasy occurs as soon as A 0 is adsorbed by montmorillonite, even if the adsorption takes place from very dilute solutions of the dye, which in the absence of the clay do not show this effect. A series of montmorillonite-A0 suspensions was1708 ADSORPTION OF DYE ON MONTMORILLONITE Table 1.Metallic cations (in mequiv. per 100 g montmorillonite) released into the equilibrium suspensions of cation-exchanged montmorillonites treated with acridine orange (1 50 mequiv. per 100 g montmorillonite) major exchange- able cation in montmorillonite Na+ K+ Mg2+ Ca2+ Cu2+ Al3+ Na 44 12 13 23 -- - < 33 12 36 9 - - Mg c u ca. 0 0.2 10 5 63 - A1 < 33 5 7 7 - 44 Table 2. into the Major exchangeable metallic cations (in mequiv. per 100 g montmorillonite) released equilibrium suspension of cation-exchanged montmorillonites treated with various amounts of acridine orange major exchangeable cations in montmorillonite acridine orange/mmol dm-3 per 100 g montmorillonite Mg2+ Cu2+ A13+ 15 30 60 90 120 150 a 0 15 27 8 0 32 7 0 37 12 0 39 25 27 36 63 44 - a Not determined.prepared (in 25 cm3) which contained a constant amount of A 0 (lob3 mmol dmM3) and increasing amounts of Na-montmorillonite. Visible spectra of the suspensians and of the supernatant obtained after separation of the clay by centrifuging were recorded. It was found that the addition of small amounts of Na-montmorillonite to the dye solution resulted in a deep coloration of the clay phase. The supernatants of these suspensions have a spectrum characteristic of free (unadsorbed) AO. The presence of free A 0 is also shown in the clay-A0 suspensions by the location of the a and /? bands at ca. 490 and 470 nm, respectively. At large clay concentrations the dye is completely adsorbed by the clay phase and no A 0 remains in the supernatant after centrifuging. The whole series can be regarded as a titration of an aqueous solution of A 0 by Na-montmorillonite. An ‘end-point’ may be defined as the stage at which all the A 0 has been adsorbed by the clay.The minimum concentration of clay necessary to reach this end point was found to be 0.0016%. This corresponds to the adsorption of 135 mmol dmP3 A 0 per 100 g clay. Aqueous suspensions of samples containing ratios of concentrations of A 0 and clay at the end-point or the stage after the end-point have a and /? bands located at 2 500 and < 470 nm, respectively. At the end-point the/? band makes the major contributionR. COHEN AND S. YARIV 1709 L I I I I 400 450 500 550 600 wavelength/nm Fig. 1. Effect of concentration on the absorption spectrum of aqueous solutions of acridine orange: (a) 0.9 x mol dmP3 (in 1.0 mm cell) and mol dm-3 (in 10.0 mm cell), (b) 6 x (c) 30 x mol dm-3 (in 1.0 mm cell).to the absorption and has its maximum at 467 nm. With increasing clay concentration the p band shifts to lower values, reaching a minimum at 460 nm when the clay concentration is 0.0024%, corresponding to an adsorption of 100 mmol dmP3 A 0 per 100 g clay. At still higher clay concentrations the band shifts again to longer wavelengths reaching 467 nm at a clay concentration of 0.0048%, corresponding to 50 mmol dm-3 A 0 per 100 g clay, and is no longer changed even at a clay con- centration of 0.008%. The adsorption of the dye cations by the clay affects the absorbance of both a and p bands in the spectrum of AO, resulting in a considerable decrease in the intensities of both bands, whose absorption intensities are dependent on the clay concentration. During the titration they decrease up to a certain concentration of clay and then increase again.The minimum absorbance of both bands takes place at a clay concentration of 0.028%, which corresponds to an adsorption of 85.7 mmol dm-3 A 0 per 100 g clay (fig. 3).1710 ADSORPTION OF DYE ON MONTMORILLONITE wavelength/nm Fig. 2. Effect of Na-montmorillonite on the absorption spectrum of 2.4 x mol dm-3 acridine orange: (a) 0.0008% clay, (6) 0.0016% clay, (c) 0.0024% clay and (6) 0.0080% clay. ADSORPTION OF ACRIDINE ORANGE BY MONTMORILLONITES TREATED WITH DIFFERENT CATIONS A series of montmorillonite-A0 suspensions was prepared (in 25 cm3) which contained increasing amounts of A 0 and a constant amount of montmorillonite saturated with one of the cations, Na+, H+, Mg2+, A13+ or Cu2+.Spectra of the suspensions and the supernatants were recorded. A whole series can be regarded as the titration of cation-exchanged montmorillonite by an aqueous solution of AO. The conclusions drawn from these titration systems are summarized as follows : (1) in the first stage of the titration the dye is totally adsorbed and the supernatant is colourless, (2) at a certain concentration the clay becomes saturated by A 0 (this stage in the titration may be defined as the end-point) and (3) further addition of A 0 results in coloration of the supernatant. The location of the maximum of the B band is dependent on the amount of A 0 present in the suspension.As the titration proceeds the /3 band shifts first to shorter wavelengths, reaching a minimum wavelength which is dependent on the exchangeable cation, and then shifts again to longer wavelengths. These results are summarized in table 3.R. COHEN AND S. YARIV 171 1 I .o Q) c 0.5 2 % e.p. 9 1 . . ‘ I band band (Y ,0-- -v- - - - - - / I ‘ \ I I I I I 1 I I 0 5 10 15 20 Na-montmorillonite (0.0 l%)/cm3 Fig. 3. Effect of the addition of an aqueous suspension of Na-montmorillonite (0.01 %) on the absorbance of bands cc and /3 in the spectrum of 2.4 x mol dm-3 acridine orange. The total volume of the suspension is 25 cm3. Table 3. Adsorption of acridine orange by cation-exchanged montmorillonites and the location of the /? band” major exchangeable cation in montmorillonite H+ Na+ Mg2+ cu2+ ~ 1 3 + lowest value of p band 459 460 465 463 468 before e.p.(in nm) mmol dm-3 A 0 per 100 g clay 90 100 75 90 55 mmol dm-3 A 0 per 100 g clay 120 135 90 110 55 /? band at e.p. (in nm) 465 462 466 467 468 /? band after the addition of 465 465 469 470 472 150 mmol dmP3 A 0 per 100 g clay ” e.p. = end-point. Changes in the location of the a band also provide information about AO-clay adsorption. As long as the added A 0 is completely adsorbed by the clay, the band maximum appears at a wavelength above 500 nm. As soon as free A 0 is present in the suspension the maximum of the a band is shifted towards 490nm. The first suspension to have a spectrum with an a band below 500 nm was considered to be that of a post end-point.Therefore the end-point for the following titrations is regarded as the amount of A 0 in the last suspension to give an a at > 500 nm. Regarding the accuracy of the estimation of the end-point, it should be mentioned here that the addition of A 0 solution to the clay suspension was carried out with increments of 0.25 cm3. The calculated amounts of A 0 adsorbed by 100 g of the various cation-exchanged montmorillonites at the end-points are summarized in table 3, where it is shown that the amount of A 0 adsorbed by the clay depends on the1712 ADSORPTION OF DYE ON MONTMORILLONITE acridine orange/cm3 Fig. 4. Absorbance of the /? band as a function of the amount (in cm3) of acridine orange (3 x mol dm-3) in 25 cm3 of an aqueous suspension of cation-exchanged montmorillonites (0.004 %).Table 4. Basal spacings (in nm) of cation-exchanged montmorillonites treated with acridine orange major exchangeable cation in montmorillonite ~~~ ~ mmol dm-3 acridine H+ Na+ Mg2+ CU2+ ~ 1 3 + orange added per 100 g clay a b a b a b a b a b 0 1.28 1.21 1.32 1.24 1.47 1.30 1.58 1.58 1.47 1.36 15 1.30 1.24 1.39 1.32 1.47 1.33 1.26 1.23 1.42 1.34 30 1.30 1.30 1.35 1.30 1.42 1.32 1.26 1.26 1.45 1.36 60 1.60 1.37 1.52 1.34 1.52 1.32 1.56 1.32 1.47 1.38 90 1.60 1.52 1.58 1.39 1.58 1.50 1.60 1.30 1.67 1.47 120 1.63 1.61 1.66 1.57 1.58 1.53 1.58 1.3OC 1.67 1.57 150 1.685 1.63 1.77 1.3lC 1.59 1.52 1.58 1.32c 1.70 1.60 a Equilibrated for one week at room temperature under an atmosphere of 40% humidity. Heated 3 h in uucuo at 150 "C.A very broad and asymmetric peak. exchangeable cation. It is highest with Na- and lowest with Al-montmorillonite. Bodenheimer and HelleP and Brindley et l8 observed a similar dependence during the adsorption of methylene blue by mono-ionic montmorillonites. The dependence was attributed to a primary adsorption which takes place at the edges of the particles. This kind of coverage may limit the extent of cation exchange. Thus,R. COHEN AND S. YARIV 1713 Table 5. Ring vibration (in cm-l) in the i.r. spectra of acridine orange sorbed on various cation exchanged montmorillonites major exchangeable cation in montmorillonite mmol dm-3 A 0 per 100 g clay Cu2+ Mg2+ AI3+ 15 1602 1602 1602 30 1600 1601 60 1600 1598 1598 90 1598 1596 1597 120 1598 1595 1595 150 1598 1595 1594 the tactoid size of the mono-ionic montmorillonite, which depends on the exchangeable metallic cation, governs the degree of cation exchange.The smaller the tactoid, the higher the extent of cation exchange. The effect of the concentration of A 0 on the intensities of the adsorption bands is shown in fig. 4, where the absorbance, A, is plotted against the amount (in cm3) of AO. From fig. 4 the ratio dA/dV, where dA is the increment in absorbance and dV is the increment in amount of AO, can be calculated. This ratio is small at the beginning of the titration. It becomes smaller with the continuation of the titration, but before the end-point it increases. After the end-point, when the added amounts of A 0 are not adsorbed by the clay but remain in the aqueous phase, the ratio dA/d V becomes high and, except for Al-montmorillonite, it remains constant even after the addition of 1.5 cm3 above the end-point.The graph which describes absorbance against cm3 A 0 after the end-point is a straight line. The lines obtained after the end-points, during the titrations of the various cation-exchanged samples, are almost parallel. Al-montmorillonite behaves in a different manner. Before the end-point the ratio dA/dV is very small. After the end-point it increases slightly, but only after the addition of 3 cm3 of A 0 solution to the clay suspension does dA/dV become as high as that found for the other clay samples. The small increment of dA/d Vin the titration region between 1.75 and 3.00 cm3 A 0 may be due to further adsorption of A 0 by Al-montmorillonite after the end-point.Thus, the A 0 which is added after the end-point partly remains in the aqueous phase and partly is adsorbed by the clay phase. By comparing this observation (table 3) with the cation-exchange study (table 2) it may be concluded that A1 is desorbed from Al-montmorillonite only after the end-point, when an excess of A 0 is present in the equilibrium system. X-RAY STUDY Oriented specimens of the clay samples taken at various stages of the titration were examined by X-ray diffraction under ambient conditions after being equilibrated at 40% humidity and dried at 150 "C in a vacuum oven. The results are summarized in table 4. By comparing the c spacings obtained before the thermal treatment with those obtained after this treatment, it is obvious that the interlayer space of samples equilibrated at 40% humidity contain water, the amount of which depends on the exchangeable metallic cation.With increasing amounts of A 0 the c spacings either1714 ADSORPTION OF DYE ON MONTMORILLONITE remain unchanged (with H, Na and Cu) or decrease (with Mg and Al) up to a certain amount and then increase again. From table 4 a number of conclusions may be drawn : (1) The dependence of the c spacing on the amount of A 0 present may serve as a proof that the sorbed cationic dye is located in the interlayer space. (2) Up to a certain degree of saturation where the c spacing is ca. 1.3 nm a monolayer of A 0 is formed in the interlayer space with the aromatic rings parallel to the aluminosilicate 1 a ~ e r .l ~ At this stage of the titration there is no possibility for any kind of aggregation of the dye cations in the interlayer space. (3) With large amounts of A 0 the c spacings of the thermally dehydrated samples are > 1.5 nm in all cases except for the Cu-montmorillonite, indicating the presence of a bilayer or tilting of the cationic dye relative to the aluminosilicate 1 a ~ e r . l ~ Under such conditions it may be possible that the interlayer space is populated by aggregates of the cationic dye, but so far there is no conclusive evidence for this. The transition between monolayer and bilayer (or tilting) as seen from the X-ray data corresponds to the minimum wavelength of the /3 band during the titration (compare tables 3 and 4). INFRARED STUDY Infrared spectra of clay samples taken at various stages of the titration were recorded under ambient conditions. A ring vibration at 1600 cm-l seems to be reliable for the purpose of following the changes in the character of the interactions in which the cationic dye takes part.This band appears at 1590 cm-l in the spectrum of A 0 microcrystals obtained by the drying of a few drops of an aqueous dye solution on a piece of polyethylene. It is assumed that under these conditions a spectrum characteristic at an aggregate of A 0 is recorded. For a clay sample containing a small amount of A 0 this band appears at 1602 cm-l and is shifted to lower wavenumbers with the addition of increasing amount of A 0 to the clay. Some representative results are given in table 5.When these results are compared with those of the X-ray and visible spectra it may be concluded that the band which appears at a high wavenumber is characteristic of a monomeric species, forming 7t interactions with the oxygen plane of the alumino- silicate layer. The shift of this band to lower wavenumbers may be attributed to the aggregation of the cationic dye in the interlayer space. CONCLUSIONS Our results indicate the occurrence of at least two types of associations formed between montmorillonite and AO. According to the X-ray data in the first type of association only a single monolayer of A 0 is located in the interlayer space of the montmorillonite, the aromatic rings being parallel to the aluminosilicate layer. The c spacing permits that in the second type the interlayer space will be populated by dimers or other forms of aggregated AO, but this has not been proved.The existence of two types of associations is supported by visible and infrared spectroscopy. Metachromasy indicates the occurrence of 7t interactions. When associations of the first type are formed, a 7t interaction must occur between the oxygen plane and the aromatic dye. However, the nature of the x interaction in the second type, which may or may not include aggregated AO, is not yet clear. We thank the Research Fund of the Faculty of Science of The Hebrew University of Jerusalem for financial support. The chemical analyses were performed by Mrs Luba Kristol of the Analytical Services Laboratory of the Hebrew University.R. COHEN AND S . YARIV 1715 V. Vitagliano, Interaction between Cationic Dyes and Polyelectrolytes, in Chemical and Biological Applications of Relaxation Spectrometry, ed. E. Wyn-Jones (D. Riedel, Dordrecht, 1975), pp. 437-466. E. Rabinowitch and L. F. Epstein, J. Am. Chem. SOC., 1941, 63, 69. V. Zanker, Z. Phys. Chem., 1952, 199, 225. M. E. Lamm and D. M. Neville Jr, J. Phys. Chem., 1965, 69, 3872. L. Michaelis, J. Phys. Colloid Chem., 1950, 54, 1 . J. F. Padday, J. Phys. Chem., 1967,71, 3488. M . Schubert and A. Levine, J. Am. Chem. SOC., 1955,77, 4197. G. Schwarz, S. Kmose and W. Balthasar, Eur. J. Biochem., 1970, 12, 454. G. Schwarz, and W. Balthasar, Eur. J. Biochem., 1970, 12,461. lo G. R. Haugen and E. R. Hardwich, J. Phys. Chem., 1963, 67, 725. l1 G. Barone, L. Costantino and V. Vitagliano, Ric. Sci., Parte 2, Sez. A. 1964, 34, 87. l 2 S. Yariv and D. Lurie, Zsr. J. Chem., 1971, 9, 537. l3 A. Yamagishi and M. Soma, J. Phys. Chem., 1981,85, 3090. l4 S. Yariv, L. Heller, Z. Sofer and W. Bodenheimer, Zsr. J. Chem., 1968, 6, 741. l5 V. Vitagliano, 0. Ortona and M. Parrilli, J. Chem. SOC., 1978, 82, 2819. l6 W. Bodenheimer and C. Heller, Zsr. J. Chem., 1968, 6, 307. l7 G. W. Bnndley and T. D. Thompson, Zsr. J. Chem., 1970,8, 409. l8 Ph. Th. Hang and G. W. Brindley, Clays Clay Miner., 1970, 18, 203. R. Green-Kelly, Trans. Faraday SOC., 1955, 51, 412. (PAPER 31 1 175)

ISSN:0300-9599    

DOI:10.1039/F19848001705

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1984, 80, 1717-1733 Photophysics of the Excited Uranyl Ion in Aqueous Solutions Part 1 .-Reversible Crossing BY SEBASTIAO J. FORMOSINHO~ AND MARIA DA GRACA M. MIGUEL Chemistry Department, University of Coimbra, 3000 Coimbra, Portugal AND HUGH D. BURROWSt Chemistry Department, University of Ife, Ile-Ife, Nigeria Received 25th July, 1983 Upon excitation at 337 nm by a nitrogen laser the uranyl ion undergoes biexponential decay from pH 1 to pH 4, this being more evident at the higher pH values. At pH 3 the biexponential is evident for [UOE+] from 2 x to 0.1 mol dm-3. The decay rates depend on [UO;+], the laser intensity and the temperature. Temperature also affects the characteristics of the decay, giving single-exponential decay temperatures < 6 "C.Under steady-state conditions comparison of the fluorescence spectra at low and high temperatures reveals another emitting state of uranyl that has a spectrum which is similar to but weaker than the normal uranyl luminescence and shows a red shift of ca. 300 cm-l. Several possible kinetic models for analysing the experimental data are discussed and it is concluded that the data can be best interpreted in terms of a reversible crossing between two states of uranyl, U* ca. 300 cm-l higher than X*, both of which decay by unimolecular kinetic processes. The oscillator strengths of the two states determined from the fluorescence decays and spectra are in good agreement with absorption spectral data. The model allows the estimation of the rate constants of the different processes at different temperatures, uranyl concentrations and excitation intensities.At pH 3, [UOg+] affects only the rate of decay of the state U*, whilst an increase in the excitation intensity decreases the rate of the reversible crossing. The activation energies of all the processes suggest that they are of a chemical nature. It is suggested that the reversible decays are due to a solvent exchange process and the irreversible decays are due to hydrogen abstraction from coordinated water molecules. The excited uranyl ion continues to attract the interest of photochemists and spectroscopists,l but in spite of the vast literature on the subject there remain a number of puzzling aspects about the decay of (UOi+)* in aqueous solutions. (UOi+)* has been found to decay exponentially in water over many lifetimes,2v but, under certain experimental conditions, dual luminescence and non-exponential decays have also been reported.*^ Such features were interpreted by Marcantonatos in terms of exciplex formation between (U02H2+)* and UOi+ via a complex mechanism where hydrogen abstraction from coordinated water plays an important role? However, the exciplex mechanism does not agree with experimental data that we have recently obtained.We show that a mechanism involving reversible crossing between two almost degenerate states of (UOi+)* is in good agreement with both our experimental data and the results of Marcantonatos. -f Visiting Professor at the University of Coimbra. 17171718 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS EXPERIMENTAL Excited uranyl decays were studied using nanosecond flash-photolysis apparatus (Applied Photophysics) with a pulsed (3.0 ns) N, laser (Lambda Physics).The exciting wavelength was 337 nm and the laser pulse had an average intensity of 2 mJ. Emission was monitored at right angles to the laser beam using a photomultiplier. Signals were recorded on Nicolet model 204 digital transient time recorder. Since the laser intensity affects the decay rates, the laser intensity was monitored simultaneously with the fluorescence intensity; except where otherwise stated the laser intensity was constant within 10% for each set of experiments. Normally decays were run at room temperature, but for temperature studies the cell was thermostatted in a water bath giving temperatures k0.5 "C.Solutions were not normally degassed as oxygen has no effect on the luminescence decays. Luminescence spectra were run on a Spex Fluorolog model 11 1 fluorimeter. Absorption spectra were run on a Shimadzu UV-240 absorption spectrophotometer. The fluorescence quantum yield of uranyl solution at pH 3 was determined with respect to fluorescein (0.1 mol dm-3 NaOH) for excitations at 366 nm; the fluorescence quantum yield of the standard is 0.92.6 Since the uranyl ion and fluorescein emit in the same wavelength region no correction for the photomultiplier response was introduced. (UOi+)* fluorescence yields at other pH values and at different exciting wavelengths were determined with respect to the yield of the uranyl solutions ([UO2,+] = 0.02 mol dm-3) at pH 3 and Aex = 366 nm.Solutions were prepared from uranyl nitrate of the purest grade commercially available in triply distilled water ; the pH was adjusted with nitric acid and was found to maintain its value after excitation of the uranyl solutions. Fresh solutions were used for all the experiments. RESULTS . EXCITED-STATE DECAY Upon excitation at 337 nm the luminescence decay of (UO:+)* followed for 4-5 half-lives is non-exponential (fig. 1) for pH 1-4. Except where otherwise stated, we will here present the results at pH 3 where both exponential components have similar intensities at room temperature. At this pH the dimer (UO,),(OH)i+ does not exceed 5% of 0.1 mol dm-3 UOi+ since the equilibrium constant is [(UO,), (OH):+] [H+I2/[UOi+l2 = mol dm-3.7 The effect of pH on the photophysics of uranyl will be discussed in Part 2 of this series and the effects of temperature and D,O in Part 3.These non-exponential decays can be fitted fairly closely (20%) by the sum of two first-order decays, with rate constants k , and k,, over a uranyl concentration range from 2 x to 0.1 mol dm-3. Within experimental error no effect was found on the overall decay of UOi+ in the presence of 0, in equilibrium with the atmosphere. However, both the decay rates and the relative intensities of the two emitting components are slightly dependent on the wavelength of detection, Aem (see table l), with rates showing a small decrease with increasing &In* INTENSITY OF EXCITATION The fluorescence decay was studied as a function of the laser intensity, Iex, which was varied by a factor of 4 by employing neutral density filters.Fig. 2 shows that the decay of the fast component, k,, decreases with increasing laser intensity whereas k , is constant. The intensity of both components increases linearly with the intensity of excitation, but the intensity of the first component, I,, has a stronger dependence on Iex. Consequently the ratio of the intensities of both components, IJI,, increases with increasing laser intensity . TEMPERATURE The effect of temperature on the uranyl decay was studied between 3 and 50 "C for [UOi+] = 0.02 mol dm-3. Changing the temperature causes drastic changes in theS. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1719 I I I I I I 1 I I I I 0 2 4 6 8 10 12 14 16 18 20 tlW I 0 L 8 12 16 20 24 tlw Fig.1. Logarithmic plots and reconvolution of the luminescence decay of UOi+ in aqueous solutions at pH 3.0; [UO;'] = 0.02 mol dm-3, T = 20 O C , (---) fast-decay component. decay characteristics and the decay rates. At temperatures < 6 "C the decay can be represented by a single exponential, which seems to correspond to the long-lived component of the biexponential decay observed at higher temperatures. Table 2 shows the decay rates and 11/12 increasing with increasing temperature. The ratio 11/12 approaches zero at 6 "C. AUTOQUENCHING Uranyl is known to suffer autoquenching39 although such evidence was obtained under conditions where uranyl decay was considered to be represented by a single exponential.We have studied the effect of UOi+ concentration on (UOi+)* lumin- escence decay over the concentration range 0.005-0.1 mol dmP3. As table 3 shows, the decay rates increase with increasing [UOz+]. Both decays follow good Stern- Volmer plots but whereas k , depends strongly on [UOi+] with a quenching rate k, = 3.7 x lo6 dm3 mol-l s-l (20 "C), k, has a much smaller quenching rate,1720 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS Table 1. Excited uranyl decay as a function of the wavelength of detection; [UOg+] = 0.02 mol dm-3, pH 3.0 _ _ ~ ~ ~ ~ ~ I/nm 460 470 485 500 kJ106 s-l 6.0 5.7 5.5 5.4 k,/105 S-1 1.10 1 .oo 0.95 0.90 11 /I2 1.75 1.40 1.10 0.90 30 I 2 0 10 9 8 " I v) vl 2 --. - * 7 6 0 I I I 100 200 300 1lasC-r Fig. 2. Intensities and rate constants for the biexponential decay of (UOi+)* as a function of the intensity of excitation, lex; pH 2.8, [UOi+] = 0.04 mol dm-3.S.J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1721 Table 2. Temperature dependence of excited uranyl decay; [UOi+] = 0.02 mol dmP3, pH 3.0 T/"C 9 12 15 20 26 36 50 k,/ lo5 s-l 4.00 4.30 5.0 6.3 8.90 12.00 16.10 k2/ lo5 s-' 1.10 1.10 1.15 1.10 1.40 1.85 3.20 11/12 0.42 0.60 0.82 0.94 1.11 1.23 1.25 Table 3. Autoquenching decay rates and intensities of uranyl excited state; pH 3.0, T = 20 "C [UOi+]/mol dm-3 5 x 1 x 2 x lop2 4 x lop2 6 x 8 x 1 x 10-I k/105 s-l 7.30 7.40 7.80 8.70 9.30 10.20 11.30 I1 /I2 0.50 0.60 0.70 0.90 1 .oo 1.15 1.25 k2/105 s-l 1.10 1.20 1.25 1.40 1.70 1.75 2.00 I 1 1 I 0.02 0.OL 0.06 0.08 0.1 [ UO$'l /mol dm-3 Fig.3. Stern-Volmer plot for the autoquenching of excited uranyl ion under stationary conditions (excitation at 406 nm); [UOg+] = 0.02 mol dm-3, pH 3, E = 8.6 dm3 mol-' cm-l. k, = 1 x lo6 dm3 mol-1 s-l. The ratio 11/12 increases by a factor of 2.5 within this concentration range. Both the autoquenching rates and the intensity ratio Il/Z2 increase with increasing temperature. The autoquenching of uranyl was also studied under stationary conditions. However, since the intensity of light absorbed, labs, increases with increasing [UOi+], the Stern-Volmer plot was obtained for ( I ~ ~ / z ~ b s ) / ( z ~ ~ / z a ~ s ) as a function of [uoi+], where I,, is the fluorescence intensity and Zabs the intensity of light absorbed; the Z" data were obtained at [UOg+] = 0.005 mol dm-3.The amount of absorbed light was1722 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS 4 50 5 00 550 6 00 X fnm Fig. 4. Fluorescence spectra of uranyl ([UOt+] = 0.05 mol dm-3) at (---) 5 and (-) 23 "C, pH 3.1. (. - - .) Spectrum obtained from subtraction after normalization at 485 nm. estimated from Beer's law with E = 8.6 dm3 mol-1 cm-l. Fig. 3 shows that the data obtained at 20 "C follow a reasonable Stern-Volmer plot with K,, = 26 dm3 mol-l. In order to obtain the quenching rate an average (UOi+)* decay rate needs to be estimated. The average decay rate can be calculated as k,, = k , + I , k,)/(Zl + I , ) and with a value of k,, = 3 . 3 ~ lo5 s-l the autoquenching rate constant is k , = 8.0 x lo6 dm3 mol-l s-l. This quenching rate is 2.2 times higher than the quenching rate determined under dynamic conditions.FLUORESCENCE SPECTRA Marcantonato~~ has obtained an emission spectrum by subtracting the uranyl emission at [UO$+] = 10+ mol dm-3 from the spectrum at 6 x mol dm-3 after normalizing both spectra at 485 nm. This residual emission at pH 2, which is ca. 10 times weaker than the normal fluorescence spectrum, was attributed to an exciplex emission. Since temperature effects were found in the fluorescence intensity of (UOi+)*, we decided to study the fluorescence spectra as a function of temperature. Fig. 4 presents the fluorescence emissions at 23 and 3 "C. It is clear that the spectrum at room temperature has a series of shoulders on the low-energy side of the bands of the fluorescence spectrum of uranyl at low temperature.Upon normalizing both spectra to the band at 485 nm, subtraction of the low-temperature spectrum from the room-temperature spectrum produces a new spectrum which is also shown in fig. 4.S. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D . BURROWS 1723 This new spectrum, X*, is virtually identical to that for (UOi+)* emission (3 "C) but is red shifted by ca. 300 cm-l. Its area is 4.3 times smaller than that for uranyl fluorescence and it is similar to that reported in the literat~re.~ Although Marcantonatos reports a ca. 120 cm-l shortening of the vibrational progression in the spectrum of X*, we were not able to find such an effect. DISCUSSION MODELS OF BIEXPONENTIAL DECAY Three likely models for biexponential decay were considered.These are analysed below. TWO STATES CONNECTED BY REVERSIBLE CROSSING A general kinetic model for biexponential decay can be established on the basis of a reversible crossing between two excited states U* and X* ki where k, and k, are the total decay rates of U* and X* and ki and kr are the constants for the reversible crossing. Assuming that only U* can be populated directly on excitation, the evolution of a system excited in this state is described by the differential equations --- d[U*l - (k, + ki) [U*] - kr[X*] - labs(t) dt -~ d[X*l = (k, + k,) [X*] - ki[U*]; dt In the case of short-pulse excitation, integration of these equations leads to the solutions [U*(t)] = [U*(O)] [C, exp (- k, t) + C, exp (- k, t)] (1) [X*(t>l = [ U * ( O ) l y ki [exp ( - k2 t ) - exp ( - k, t)] kl k2 and Thus the excited U* state decays according to a biexponential law with two lifetimes k;l and kil.The state X builds up with the higher rate k, and then decreases with the rate constant k,. The ratio of the fluorescence intensity of the two components is Stationary-state luminescence studies are used very frequently to investigate the influence of quenchers on the luminescence yields. Under stationary conditions d[U*]/dt = d[X*]/dt = 0 and we can write1724 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS and In the presence of a quencher, Q, that has quenching rates k, and kb, respectively, for U* and X*, Stern-Volmer linear plots can be observed when kb[Q] -g kx+kr. Under these conditions the emission intensity of U* in the presence, I, and absence, kJQ1 I,, of quencher is 1 0 -= 1 + I k, + ki - ki k,/(kx + k,) ‘ When this Stern-Volmer equation is compared with the normal one in the absence of reversible crossing we realize that, under stationary conditions, reversible crossing increases the quenching rate with respect to the dynamic methods, as if it was a case of static quenching. TWO STATES CONNECTED BY IRREVERSIBLE CROSSING (EXCIPLEX MODEL) First exciplex model: Instead of being formed in a radiative transition X* can be an excimer or e~ciplex.~ The most simple mechanism for the formation of such a species is U*+U +E* + U + U kD k E k” u * + u.In absence of reversibility for the first reaction, biexponential decays can be observed if both U* and E* species emit at the wavelength of detection.In the case of short-pulse excitation, eqn (1) allows the establishment of the following equations: [u*(t)] = [U*(o)] exp { -(ku + kDIUl) t> cu*(0)l kDIU1 (exp ( - k , t ) - exp { - (ku + k,[U]) 1 ) ) . EE*(t)l = ku + k D [ q - k, The overall decay Z*(t) is therefore (7) where a is the ratio of emission intensities of E* and U* at the detection wavelength. Rearranging this expression - akDIU1 exp (- k, I)]. (8) k,+k,[Ul-k, The ratio of the intensity of the two components is I, 1-a k,-kE +- I2 a akD[U]’ _-- - (9)S. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1725 Under stationary-state conditions and if only state U* suffers quenching then the (10) Stern-Volmer equation is kq[Ql - I + I ku + k,[UI* Since ( k , + k,[U])-l is the lifetime of U*, static and dynamic quenching are expected to have the same quenching rate.Second exciplex model: Another possible kinetic model involving the formation of E* is the following: kD kE k-D U* + U e E* + U +U k-D u*+ u. When only the emission of U* is observed its decay is again biexponential where [U*(t)] = [U*(O)] [C; exp (- k; t)+ Ci exp (- ki t)] 2ki, 2 = (k, + k&] + k-, + kE) d { ( k u + k,[U]- k-D- kE)2 + 4k-D k,[U]} (1 1) and TWO-INDEPENDENT-STATES MODEL Finally we can consider the possibility that the states U* and X* are two independent excited states that can be formed immediately upon excitation. The two states emit in the same wavelength region and decay upon short-pulse excitation: [U*(t)l = [U*(O)l exp (- ku t ) (1 3) [X*(t)] = [X*(O)] exp (- k , t).Under stationary conditions the intensities of luminescence are, in the absence of quencher, k; I L b s k , I’ = ~ where k f and k; are the radiative rate constants of U* and X*. The ratio of the total emission intensities in the absence and in the presence ofa quencher that only quenches em If then Again the Stern-Volmer equation is obeyed, but now since B > 0 the apparent quenching is lower than the quenching measured in dynamic studies for U*. As table 1 shows, the emission intensities of the two components of the biexponential decay1726 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS are of similar magnitude and so p x 1. Under these conditions the quenching rate is half the value found in dynamic experiments. Although there are various criteria that can distinguish between these mechanisms, the comparison between static and dynamic quenching in a situation where just one of the states suffers quenching is one of the most powerful criteria.COMPARISON BETWEEN EXPERIMENTAL DATA AND THE EXCIPLEX MODEL Marcantonato~~! has interpreted the biexponential decay of (UOi+)* in terms of dual luminescence from (UOi+)* and from the exciplex (U,04H4+)*. This proposal was based on the fact the luminescense intensity of uranyl solutions is not a linear function of the rate of light absorption. The luminescence spectrum of the exciplex was also obtained from the fluorescence spectra at different uranyl concentrations. Such a mechanism was also apparently in agreement with the autoquenching process, because, according to eqn (7), only the first component of the biexponential decay suffers the quenching process; the second component, k,, is independent of uranyl concentration. Furthermore, at room temperature and low uranyl concentration, 3 x mol dm-3 (pH 1.9), the decay became a single exponential.In spite of such facts we have found a considerable number of experimental results that arouse serious doubts as to the role of the exciplex mechanism in uranyl photophysics. The first difficulty with the exciplex mechanism is the fact that biexponential decay is observed at all emission wavelengths between 460 and 520 nm. However, according to the spectrum reported by Marcant~natos,~ biexponential behaviour should only have been observed for wavelengths 2 490 nm. Although Marcantonatos4 was able to observe single-exponential decay at low uranyl concentrations (3 x loa3 mol dm-3) we were able to observe biexponential decay at 20 and 40 "C at lower concentrations, [UOi+] = 2 x Marcantonatos' results could be due to the fact that a lower pH and excitation intensity was used.According to eqn (9) in the first exciplex mechanism 11/12 should be a linear function of [UOi+]-l. However, the opposite effect is observed since I J I , increases with increasing [UO;+]. With IJI, = 0.7 ([UOi+] = 0.02 mol dm-3) and the rate constants k, = 7.8 x lo5 s-l, k, = 1.25 x lo5 s-l and k , = 4.0 x lo6 s-l, as found in the autoquenching of uranyl, the ratio of E* and U* emission intensities at the detection wavelength should be a = 5.4. Consequently strong enhancement of the overall fluorescence with increasing UOi+ concentration should have been observed.This is not the case under either dynamic or stationary irradiation conditions. Furthermore, the model predicts that the static and dynamic autoquenching rates should have the same value, contrary to the experimental observations. The first exciplex model is thus not able to explain the fluorescence decay of the excited uranyl ion. With the values of k,, k , and 11/12 given above, eqn (1 1) and (12) reveal that for the second exciplex model the maximum value of kD is 9 x lo5 s-l. Such a value implies a minimum stabilization energy for the exciplex E* of AH M -28 kJ mol-1 and consequently a red shift of at least 2300cm-l was expected for the second fluorescence emission.However, a red shift of only 300 cm-l was observed. Conse- quently the second exciplex mechanism cannot interpret the present experimental data. mol dm-3. REVERSIBLE CROSSING versus THE TWO-INDEPENDENT-STATES MODEL The distinction between ' reversible crossing' and the ' two-independent-states model' can be made by a comparison of dynamic and stationary uranyl quenching with a quencher that only quenches one of the states. As far as the ' two-independent- states model' is concerned the autoquenching data of uranyl seem to fulfil thisS. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1727 0 0.05 0.1 [ UOq'] /mol dmd3 Fig. 5. Analysis of the autoquenching data of table 3 within a reversible-crossing model: @,k,+ki; A, k,+k,; M, kik,. purpose. This also seems to be valid for the reversible-crossing model, because the analysis of the experimental data according to eqn (l), (3) and (4) allows estimation of k,+ki, k,+k, and k i k , as functions of uranyl concentration.As fig. 5 shows, only the state U* suffers quenching by UO;', with a quenching rate of 3.8 x lo6 dm3 mo1-I s-l at room temperature; the state X* suffers no quenching during its lifetime. According to the models presented a static quenching rate higher than the dynamic quenching rate is only accounted for by the reversible-crossing model. Eqn ( 5 ) predicts an increase in the quenching rates from dynamic to stationary experiments of 2.0 in close agreement with the experimental value of 2.2. For a two-state model the stationary quenching rate would have been ca.2 times slower than the dynamic quenching rate. Furthermore, the difference in the two quenching rates is not amenable to an explanation through a static-quenching mechanism since the dimerization of UO;+ at pH 3 does not exceed 5% at the highest uranyl concentration. Further support for the reversible-crossing mechanism comes from the temperature studies. A decrease in temperature causes a large decrease in the intensity of the first component of the decay curve that seems to disappear at temperatures close to 6 "C. Such a decrease can be easily explained in terms of a reversible-crossing mechanism, since, according to eqn (4), this decrease can be caused by a decrease of k , and/or k,+ki which are/is faster than the decrease of k , caused by the decrease in temperature.In contrast, within the ' two-independent-states model ' such a decrease cannot be attributed to the temperature dependence of k , and k,. It could be explained1728 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS I I I I 0 100 2 00 300 Fig. 6. Effect of laser intensity on the decay rates of (UO?j+)* decay in terms of a reversible- crossing mechanism; [UO;'] = 0.04 mol dm-3, pH 2.8: ., ku; 0, k,; A, ki; A, k,. 4aser Table 4. Activation energies for the decay rates in excited uranyl ion, EJkJ mol-I; [UO2,+] = 0.02 mol dm-3 kU 24 24 67 50 in terms of a ground-state equilibrium between the two states, but equilibrium between UOi+ and one of its polymeric species is not significant at pH 3. The same kind of interpretation was also required to explain the increase of IJ12 with increasing [UO;+].ANALYSIS OF DATA IN TERMS OF THE REVERSIBLE-CROSSING MODEL Within the present reasoning experimental data should be analysed in terms of a fluorescence decay of state U* initially excited and in quasi-equilibrium with another state X* that does not emit at the wavelength of detection. Biexponential decays of (UO;+)* allow the experimental determination of the rates of exponential decays k , and k , and the ratio of the intensity of the two components, Il/12. With these three parameters eqn (3) and (4) allow us to calculate the rates of overall decay of the states U* and X*, respectively k , + ki and k , + k,, and the product of the rates of reversible crossing, k,k,. However, for a convenient study of uranyl photophysical processes it is necessary to differentiate between the reversible rates andS.J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1729 the other rates of decay. Since the build up and decay of state X* cannot be followed directly through emission, in spite of a search for this possibility, we have to make this separation in terms of an average of the extreme values of k, and k,. If we designate the experimental values k,+ki = y , k,+k, = 6 and kik, = E (1 5 ) the maximum possible value for ki is k,(max) = y ; the minimum value is k,(min) = ~ / d since k,(max) = 6. The geometric average of these two ki values is k?" = 2 / ( ~ ~ / 6 ) kFv = 2 / ( ~ 6 / y ) . and by the same reasoning This averaging procedure satisfies the experimental condition that k?" k:" = E .As we see in the following discussion, this procedure seers to provide a reasonable estimate for the reversible rate constants and allows an e: imation of the decay rates k , and k,. Fig. 1 shows the reconvolution of a fluorescence decay with the calculated rates k,, k,, ki and k,, and the agreement with the experimental curve is good. Fig. 6 shows an analysis of the effect of the laser intensity on the different decay rates of (UOi+)*. The results reveal that at pH 2.8 the laser intensity has no effect on the decay rates of (UOi+)*, k, and k,. Since the quantum yield of UOi+emission at pH 3.0, determined with respect to fluorescein, is 4F = 1.1 x ([UO:+] = 0.02 mol dmF3), the rates k, and k, are virtually non-radiative rates of decay. However, ki and k, suffer a large decrease with increasing I,,,,, .Temperature has a different effect on the rates of reversible crossing and on the rates of decay. The rates ki and k, have normal Arrhenius-type behaviour for temperatures c 30 "C (table 4). At temperatures > 30 "C the rates tend asymptotically to their maximum values of ki (max) x 6.5 x lo5 s-l and k,(max) = 5.5 x lo5 s-l. In contrast the rates k, and k , are independent of temperature for temperatures Z 21-26 "C; k, even has a small negative apparent activation energy in this temperature region, E, = - 12.5 kJ mo1-l. At higher temperatures k, and k, have normal activation energies with E,, = 24 kJ mol-' and a pre-exponential factor of 2 x lo9 s-' (fig. 7). The autoquenching process can also be analysed within a reversible-crossing scheme.The results presented in fig. 8, calculated from data at 21 "C, are typical for the other temperatures. Uranyl has no effect on the rates k, and k, of the state X* but has a normal quenching effect on the decay of the state U*, which follows good Stern-Volmer kinetics. The quenching rate is always higher for ki than for k,. The temperature also increases the autoquenching rates for ki and k,, E,(ki) = 46 kJ mol-1 and E,(k,) = 25 kJ mol-l. Although [UO;+] has virtually no effect on k , and k, a more detailed study as a function of temperature reveals that k, has a small but positive quenching rate at 12 "C, k,(k,) = 7.5 x lo5 dm3 mol-l s-l, identical to the one for k,, but the effect decreases with increasing temperature. The rate k , suffers a small decrease up to [UOi+] = 0.02 mol dmF3 and then stays constant.These results cannot be interpreted in terms of the exciplex mechanism proposed by Marcant~natos,~. but in contrast our model can interpret his experimental observations. Since UOi+ only quenches the state U* and one of the components of the biexponential decay, it becomes clear that the intensity of fluorescence is not a linear function of the light absorbed when [UOi+] is varied. Such deviation increases with increasing [UOi+] and becomes noticeable when U* is quenched by ca. 20%. At room temperature this corresponds to [UOi+] = 0.018 mol dm-3, in agreement with 51 F A R 11730 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS 13 I 11 3.0 3.2 3 .L 3.6 1 0 3 KIT 13 n I v1 \ W - c - 12 11 I I I I I I 3.0 3.2 3.L 3.6 lo3 KIT Fig.7. Effect of temperature on the decay rates of (UOg+)* decay for a reversible-crossing mechanism; [UOi+] = 0.02 mol dm-3, pH 3.0: 0, k,; 0, k,; A, ki; A, k,. the observations of Marcantonatos. Since the autoauenching decreases only the fluorescence of U*, any emission of X* can be estimacd from the difference in the fluorescence spectra at high and low uranyl concentrations, as found by Marcant~natos.~ Temperature has a similar effect and consequently it is no surprise that the spectrum of fig. 4 is identical to the one reported by Marcantonatos. This spectrum can be attributed to the state X*. Within this hypothesis the ratio of theS. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1731 5 10 [ U 0 l 2 ] / 1 Od2 mol dm -3 Fig.8. Effect of uranyl concentration on the decay rates of (UOif)* at 20 "C and pH 3 ; quenching rates k,(ki) = 2.6 x lo6 dm3 mol-l s-' and k, (k,) = 1.3 x lo6 dm3 mol-l s-l: 0, k,; A, ki; 0, k,; A, kr. spectral areas of fig. 4 together with the measured yield of emission, I . 1 x lop3, allows an estimation of the fluorescence yields of U* and X*. The ratio of the integrated intensities is Iu*/Ix* = 4.23 and consequently = 0.89 x and &* = 0.21 x At room temperature and with [UOf+] = 0.02 mol dm-3 the rates of decay are: ku = I .2 x lo5 s-l, ki = 2.9 x lo5 s-l, kx = 1.6 x lo5 s-l and kr = 3.9 x lo5 s-l and because kU, h* = k, + ki - ki &k, +- k,) the radiative rate constant for U* is k$ = 1.8 x lo2 s-l. The yield of X* emission is kg ki 1 Qx* = ~ kx i- kr ku + ki - ki kr/(k, i- k,) and k$ = 0.80 x lo2 s-l.The Einstein formula e22.3 x el fTW2 krad = - where ei are the degeneracies of the electronic states andfis the oscillator strength of the electronic transitions, allows evaluation of the oscillator strengths for these emissions. With e,/e, = 29 and hv = 2.55 eV the oscillator strengths are in good agreement with the absorption results of Bell and BiggerslO (see table 5 ) in spite of the fact that their results were obtained in perchloric acid solutions and ours in uranyl nitrate solutions at pH 3. Their agreement gives further support to the reversible- 57-21732 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS Table 5. Electronic origins and oscillator strengths of the excited uranyl ion absorption a emission E/cm-l f E/cm-l f 22 050 3.8 x 22 looc - 21 330 1.67 x 20 460 1.3 x 20 580 0.45 x 20 160 0.56 x a Electronic origins in the first absorption band.1° This work.See Part 2 of this series. crossing mechanism between the two lowest excited states of (UOi+)* and seems to rule out again an exciplex nature for the second emission. The difference in U* and X* emissions is 300 cm-l, whereas for absorption it is higher (750 cm-l). The effect of changing the counter-ion from NO; to Cloy and of changing the pH may be responsible for such a difference. Another plausible cause of such an effect could be the fact that the coordinated water molecules are rearranged strongly in a non-Franck-Condon fashion in the absorption of U*. The presence of a weakly luminescent state X* with a first emission band at 496 nm accounts for the slight decrease in k, with increasing wavelength of detection (table l), because X* builds up initially with a rate constant k, [eqn(2)].If the detection wavelength was centred in one of the X* bands the decay rate k, could at most decrease by 25%. This does not occur under our experimental conditions where for detection at 485 nm the contribution of X* is not higher than 5%. An increase in the intensity of the excitation decreases the rates, ki and k,, of reversible crossing as fig. 6 shows. This effect can explain the single-exponential decays2v3 observed for (UOi+)* when excited with laser pulses ca. 20 times more intense than ours. However, under low-intensity excitations, similar to those achieved using single-photon counting a p p a r a t ~ s , ~ ~ uranyl decays are biexponential.ON THE NATURE OF THE RADIATIONLESS PROCESSES The rates of reversible crossing ki and k, have quite large activation energies (5&67 kJ mol-l) for a purely physical non-radiative process where the activation energies are '< 15 kJ mol-l. Therefore these processes are probably associated with some kind of chemical processes. It may be important for the elucidation of this question that Ikeda et d . l l have reported a rate for water exchange of 9.8 x lo5 s-l in the ground state of uranyl at 25 "C. These n.m.r. studies also provide an activation enthalpy of 42 kJ mol-l for the exchange. Although ki and k, refer to an excited state both the activation energies and the rates of exchange are of a similar magnitude to these values, which suggests that the reversible crossing could be due to solvent exchange in both states U* and X*.The rates of the irreversible decays can be interpreted in terms of a tunnelling process12913 since the decay rates k, and k, have an activationless region at low temperatures and an activation region with apparent activation energies of 24 kJ mo1-l up to 50 "C. Burrows and Formosinhol* have studied theoretically within a tunnelling model hydrogen abstraction from H,O molecules by (UOi+)*, and such a process may correspond to the rates k, and k,. Nevertheless, elucidation of the nature of the reversible and irreversible decays requires further studies and will be more fully discussed in Part 3 of this series.S. J. FORMOSINHO, M. DA G. M. MIGUEL AND H. D. BURROWS 1733 This work was supported by INIC through the Research Centre QC-1. We thank G.T.Z. for the gift of the Spex-Fluorolog fluorimeter used in this work and the referees for useful comments. C. K. Jurgensen and R. Reisfeld, Struct. Bonding (Berlin), 1982, 50, 121. R. J. Hill, T. J. Kemp, D. M. Allen and A. Cox, J. Chem. Soc., Faraday Trans. I, 1974, 70, 847. P. Benson, A. Cox, T. J. Kemp and R. Sultana, Chem. Phys. Lett., 1975, 35, 195. M. D. Marcantonatos, Znorg. Chim. Acta, 1978, 26, 41. M. Deschaux and M. D. Marcantonatos, Chem. Phys. Lett., 1979, 63, 283; M. D. Marcantonatos, J. Chem. Soc., Faraday Trans. I , 1980, 76, 1093. G. Weber and F. W. J. Teale, Trans. Faraday SOC., 1957, 53, 646. H. S. Dunsmore, S. Hietanen and L. G. Silkin, Acta Chim. Scand., 1963, 17, 2644; R. N. Sylva and M. R. Davidson, J . Chem. SOC., Dalton Trans., 1979, 465. H. D. Burrows, S. J. Formosinho, M. G. Migueland F. Pinto-Coelho, J . Chem. Soc., Faraday Trans. I, 1976, 72, 163. C. K. Jerrgensen, J . Lumin., 1979, 18/19, 63. lo J. T. Bell and R. E. Biggers, J . Mol. Spectrosc., 1965, 18, 247. l 1 Y. Ikeda, S. Soya, H. Fukutomi and H. Tomiyasu, J . Znorg. Nucl. Chem., 1979,41, 1333. l2 J. Jortner, J . Chem. Phys., 1976, 64, 4860. l 3 E. Buhks and J. Jortner, J . Phys. Chem., 1980,84, 3370. l4 H. D. Burrows and S. J. Formosinho, J . Chem. Soc., Faraday Trans. 2, 1977, 73, 201. (PAPER 3/ 1282)

ISSN:0300-9599    

DOI:10.1039/F19848001717

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984, 80, 1735-1744 Photophysics of the Excited Uranyl Ion in Aqueous Solutions Part 2.-Acidity Effects between pH 0.5 and 4.0 BY MARIA DA GRACA M. MIGUEL, SEBASTIAO J. FORMOSINHO* AND ALJGUSTO C. CARDOSO Chemistry Department, University of Coimbra, 3000 Coimbra, Portugal AND HUGH D. BURROWST Chemistry Department, University of Ife, Ile-Ife, Nigeria Receiued 25th July, 1983 Fluorescence quantum yields and decays of excited uranyl ion in aqueous solutions have been studied over the pH range from 0.5 to 4.0. The fluorescence yield decreases between pH 0.5 and 2, is constant up to pH 2.5 and then varies again with a maximum at pH 3.5. The pH dependence of the emission decay is complex but when analysed in terms of a reversible-crossing mechanism reveals that the rate of reversible crossing has a maximum at pH 3 whereas the rate of irreversible decay has a maximum at pH 2 and a minimum at pH 3. All these effects have been interpreted in terms of several aquo, hydroxo-aquo and polynuclear uranyl cations (UO?j+)*, [UO,(OH)+]*, [UO,(OH),]* and [(UO,), (OH);]*, with a pH-dependent distribution curve and with different rates of fluorescence decay.Good agreement between the decay and the stationary-fluorescence data is found within a reversible-crossing kinetic scheme between two uranyl states that are energetically very close, a higher U* state and a lower X* state. Increasing uranyl concentration increases only the rate of decay of state U*, except at pH 1 where some effect is detected in the X* decay.The acidity effect on the autoquenching rate constant is not very pronounced. The laser intensity does not affect the rates of irreversible decay of both states, but affects the rate for the reversible transition between U* and X*. This effect depends on the pH. The photophysics of the uranyl ion in aqueous solutions is dependent on a great number of fact0rs.l One of these factors is the acidity of the solution, which strongly affects the fluorescence quantum yields and lifetime^.^-^ In the first part of this series5 we interpreted the photophysics of (UOi+)* in terms of reversible crossing between two excited electronic states U* and X* : ki u* e x* Here we quantum report a study of the effect of acidity on the decay and the fluorescence yield of the uranyl ion over the pH range from 0.5 to 4.0.The present results extend considerably the acidity range of previous studies and reveal new features that are accounted for by the reversible-crossing kinetic scheme. t Visiting Professor at the University of Coimbra. 17351736 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS 1 2 3 4 5 PH Fig. 1. Relative fluorescence yield (0) of uranyl as a function of pH, under stationary-state conditions; (-) fluorescence yield calculated through eqn (1) from the pH dependence of uranyl decay rates; [UOi+] = 0.02 mol drn-,, [NO;] = 0.55 mol dm-3 and A,, = 337 nm. EXPERIMENTAL Excited uranyl decays were studied using nanosecond flash-photolysis apparatus with a pulsed N, laser. Fluorescence spectra were run on a Spex Fluorolog model 11 1 fluorimeter and absorption spectra were recorded on a Shimadzu UV-240 spectrophotometer.Solutions were prepared with triply distilled water and from uranyl nitrate of the purest grade commercially available; the pH was adjusted with HNO, and NaOH for pH > 3.0. The fluorescence quantum yields were determined at different pH with respect to the quantum yield at pH 3.0, [UOg+] = 0.02 mol drn-,, Aex = 366 nm. Further experimental details have been presented elsewhere. RESULTS AND DISCUSSION FLUORESCENCE QUANTUM YIELDS UNDER STATIONARY CONDITIONS The fluorescence quantum yield of (UOi+)* was studied as a function of pH under stationary conditions for excitations at 337 nm. Fig. 1 shows that & decreases from pH 0.5 to pH 2, increases up to pH 3.5 and decreases again at pH 4.& is independent of the wavelength of excitation up to pH 3, but at pH 3.5 and 4 different yields are obtained for excitations at 337 nm (the laser exciting wavelength for the fluorescence decay) and at 366 nm. At this latter excitation the relative yields increase with increasing pH (pH > 2.5) with relative yields 43,S6(pH 4)/&3pH 3) = 3.6. Since the absorption spectrum of UO;+ changes drastically at pH > 3.5 because of several polynuclear uranyl species obtained by condensation reactions between hydroxocomplexes,6 this wavelength dependence can be attributed to the presence of some strongly fluorescent species that are only excited around 366 nm. The emission- excitation matrix plot (fig. 2) for the uranyl emission at different pH values supports this view.The pH dependence reported in fig. 1 is in agreement with the previous observations of Marcantonato~~? * within the pH region studied (0.5-2.5). Marcantonatos was ableM. DA G. M. MIGUEL, S. J. FORMOSINHO, A. C. CARDOSO AND H. to interpret his observations in terms of a complex mechanism pH-dependent hydrogen abstraction from H20 *UO;+ + H 2 0 f *U02H2+ + HO' and, after several steps, the formation of an emitting exciplex proposed kinetic scheme gave an expression for & D. BURROWS 1737 involving initially (U204H4+)*. The A + B [H+] " = C+D[H+] where A , B, C and D are independent of pH. Although such an expression can be used to fit data up to pH 2.5 it is clear that the mechanism of Marcantonatos is unable to interpret the acidity dependence for pH > 2.5, Moriyasu et aL7 interpret the pH dependence of uranyl decay in terms of a hydrogen-abstraction reaction of (UOi+)* that is slower from H,O+ than from H20.Although such an interpretation could again account for the experimental data at pH -= 2.5 it is unable to explain the pH dependence at higher pH. Our observations can only be fully interpreted after a discussion of the pH dependence of the excited uranyl ion decay rates according to the scheme of reversible cro~sing.~ Fig. 2 reveals that at low pH (ca. 1) uranyl emission has a high-energy shoulder at 452 nm for excitation between 380 and 400 nm. Such an emission has an energy of 22 100 cm-l, which is virtually identical to the third electronic origin (22050 cm-l) found by Bell and Bigger9 in the first absorption band of UOg+. Since significant intensity can be hidden under the most intense emissions at higher wavelengths an accurate quantum yield cannot be provided for such an emission.A rough comparison of the areas reveals that the high-energy emission is ca. 100 times weaker than the normal fluorescence emission, i.e. $F x With the oscillator strength for this transition (3.8 x lo+) reported by Bell and Biggers the non-radiative transition rate for this state is estimated to be 4 x lo7 s-l. FLUORESCENCE DECAY The fluorescence decay of uranyl excited at 337 nm by a short laser pulse was studied between pH 0.5 and 4. Since the fluorescence decay is dependent on the laser intensity, except where otherwise stated the decays were studied at a constant (within 10%) laser intensity.The decays were biexponential and analysis in terms of a reversible- crossing mechanism allowed the determination of the decay rates k,, k,, ki and k,.5 Fig. 3 presents the pH dependence of the different decay rates for [UOi+] = 0.1 mol dm-, at a constant NO; concentration (0.55 mol drn-,). As fig. 3 shows, the rates of the irreversible decays, k, and k,, have an identical pH dependence increasing from pH 0.5 up to pH 2, decreasing to pH 3 and increasing again at pH 4. This dependence contrasts with the pH dependence of the rates of the reversible crossing, ki and k,, both of which have maxima at pH 3.0. Similar variations have been found for lower uranyl concentrations down to 5 x mol dm-3. The most significant differences are above pH 3.The minimum for the rates k , and k , occurs at a higher pH (ca. 3.5) at low concentrations and the increase is not so steep around pH 4. Excited uranyl is known to suffer a significant autoquenching process. This autoquenching process at pH 3 affects only the decays rates of state U*, k , and ki. The same effect is apparent at other pH values (table 1) with the exception of pH 1, whereas k , also suffers significant autoquenching and ki suffers a very small decrease with increasing [UOi+]. Nevertheless, the autoquenching rate constants are not strongly pH dependent.1738 4.49E 05 x U .- 2 U .- O.OOE 00 45 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS I I I I 00 563.00 676 .OO position/nm 1 1.18E .06 ; I - O.OOE 00 450.00 563.00 posit ion/nm Fig.2.(a) and (b). For legend see facing page. 676 .OO The effect of the intensity of the laser was also studied at pH 1 and pH 4. The results at pH 4 resemble those at pH 35 for ki and k,, which decrease with increasing laser intensity, but more strongly at pH 4 than at pH 3. However, k, increases slightly with increasing laser intensity (ca. 1.4 for 3 times the laser intensity), but k , also decreases in a fashion similar to that of ki. At pH 1, however, ki and k , pass through a minimum (fig. 4), but k , and k , are again independent of the laser intensity.M. DA G. M. MIGUEL, S. J. FORMOSINHO, A. C. CARDOSO AND H. D . BURROWS 1739 6.36E .06 O.OOE .OO 450.00 563.00 position/nm 6 76 .OO Fig. 2. Emission-excitation matrix plot of uranyl nitrate aqueous solutions as a function of pH; [UOt+] = 0.02 mol dm-3 at room temperature: (a) pH 1; (b) pH 3; (c) pH 4.HYDROLYSIS OF URANYL Aquo-cations of high charge tend to act as acids in solutiong and UOi+ is no exception. Hydrolysis constants, pKl = 4.1-4.3, have been reported for UO:+.lo In general, the acid dissociation constant for the loss of a second proton from an aquo-cation is 10-100 times smaller than that for the loss of the first p r ~ t o n . ~ Hydroxo-0x0-aquo-cations are formed by uranyl which polymerizes at pH > 3 .06 with equilibrium constants for the dimer of ca. 10+ mol dm-3.67 Consequently at different pH different uranyl ions are present in the aqueous solutions, and these ions can have different rates of decay. The rates of reversible decay of the excited uranyl ion have been attributed to reversible crossing between two electronic states of different R, caused by a solvent- exchange rne~hanism.~ Not much is known about solvent exchange in excited states of metal ions9 nor about the possible influence of coordinated hydroxyl groups.Nevertheless, water exchange at FeOH2+ takes place much more rapidly than at Fe&ll and in U(H20)4,+ the rate of water exchange increases with decreasing acidity.12 A similar situation may also occur for the excited uranyl ion and may explain the increase in ki and k, with increasing pH up to pH 3. To our knowledge nothing is known of solvent exchange in polymer cations, but because of the 0x0, hydroxo and water bridges established between the different metal ions, we expect that solvent exchange rates are lower in such species, since fewer coordinated molecules are avail- able for the exchange.This seems to be the case for uranyl at pH 4. The rates of the irreversible decays, k, and k,, can be attributed to a hydrogen- abstraction reaction from H 2 0 by an excited uranyl ion or to a purely physical radiationless transition where the 0-H vibrations of the coordinated water play a significant role as accepting modes. However, since the 0-H frequencies are not pH1740 5 - 4 - * lv) 3 - 2 In --- 2 - PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS 3 - - I v) In 2 2 - Y' I I I 1 I 0 ' 1 2 3 4 5 PH I I I I I I 0 1 2 3 4 5 PH Fig. 3. pH dependence of the decay rates (a) and [NO;] 0 1 2 3 4 5 PH 't 0 Y 1 2 3 4 5 PH k,, (b) kx, (c) ki and ( d ) k,; [UO;+] = 0.1 mol dm-3 = 0.55 mol dm-3.Table 1. Autoquenching decay rates (s-l) at room temperature as a function of pH PH 1 .o 2.0 3.0 4.0 kll 1.0 x 106 0.96 x lo6 1.3 x lo6 0.97 x lo6 ki Oa 0.95 x lo6 2.6 x lo6 2.2 x 106 kX 1.7 x lo6 0 0 0 kr 0 0 0 0 a Q 2 x 105 S-1.M. DA G. M. MIGUEL, S. J. FORMOSINHO, A. C. CARDOSO AND H. D. BURROWS 1741 l5 t I 1 100 20 0 30 0 Fig. 4. Effect of laser intensity on the rate constants of uranyl decay at pH 1 ; [UOif] = 0.02 mol dm-3: 0, k,; 0, k,; A, ki; A, k,. Ilaser dependent for the different uranyl cations, the pH dependence of k, and k, does not support the latter mechanism for our experimental conditions. Rates of hydrogen abstraction will be dependent on the possibility of the excited uranyl ion abstracting a hydrogen atom from the water molecules loosely coordinated in the equatorial plane, or from water that comes into the hydration shell in an axial direction.The frequency factor and possibly the hydrogen-abstraction rates are higher for the former situation. Since the presence of OH groups can alter the water structure around UOi+, the hydrogen-abstraction rates can also be pH dependent. In order to test the possibility that such qualitative ideas will interpret quantitatively the data of fig. 3, a quantitative assessment of all those factors was established, assuming that all the acid-base equilibria for excited uranyl ion were faster than the fluorescence decay. The following set of equilibria? was considered : *UOi+ + H,O + *UO,(OH)+ + H+ = 2 x lo-, *U02(OH)++HH,0+ *UO,(OH),+H+ = 4 x 2 *UO,(OH),+H,O+*(UO,),(OH)~+H+ = 1 x lop3.Fig. 5 presents the distribution diagram for all these different species as functions of pH. Since NO; is present in solution some NO; complexes are also present but were not considered in the mechanism since the data do not allow a distinction between free uranyl ion and nitrate complexes. Furthermore, with an equilibrium constant t These species have not been established to be present, but are chosen on the basis of simplicity. Products of the first two equilibria should be capable of fast solvent exchange whilst polymer species undergo only slow exchange.1142 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS *uo**+ \ * UOZ (OH), PH Fig. 5. Distribution diagram for excited uranyl ions as a function of pH; [UO~+],,,,, = 0.1 mol dm-3.K* = 0.36 dm3 mol-l for the complex (U0,NO;)*13 we can estimate that at very low pH (< 0.5) (UO:+)* is at most 20% complexed with NO; ([NO;] = 0.55 mol dm-3 and [UOi+] = 0.1 mol dm-3). With those equilibria and the following rates for k, and ki k,/i05 s-1 2.3 5.8 0.1 10 *uo;+ * UO,(OH)+ * UO,(OH), *(UO,),(OH), ki/ lo5 s-l 1.7 5.2 4.5 0.3 good agreement between the calculated and the experimental rates was found (fig. 6). At lower [UOi+],,,,, the fraction of the polynuclear ion decreases and consequently there is a shift to pH ca. 3.5 in the minimum of the k, curve and the increase at pH 4 is not so strong. These facts are also in agreement with the experimental observations. The pK values for the hydrolysis of the excited uranyl ion are ca.2.5 units lower than the pK reported for the ground state. The increase in acid strength upon electronic excitation is a situation common to many other molecules. However, with uranyl, because there is not a great shift in the maximum of the absorption spectra with pH, except for pH > 3.5, the enthalpy contribution for such an increase is notM. DA G. M. MIGUEL, s. J. FORMOSINHO, A. c . CARDOSO AND H. D. BURROWS 1743 4 - - I v1 P 3 - >- a 2 . I 4 . I v1 VI 2 3 . a Y 2 . 0 1 2 3 4 5 0 1 2 3 4 5 PH PH Fig. 6. Calculated (-) and experimental (0) rates (a) k, and (b) ki for excited uranyl decay as a function of pH; [UO~+],,,,, = 0.1 mol dmP3. very high. At most we can estimate it to be 0.5 pK units. The remaining difference can be attributed to an increase in the entropy of the acid-base equilibrium in the excited state of ASAH* -ASAH = 37.5 J K-l mol-l.This value is not unreasonable since the change in entropy for the dissociation of a proton is very high for the aquo-cations of uranium; A S = 138-150 J mol-1 K-I for U4+.lo For UOq+ the situation is very unsatisfactory because the choice of values for the enthalpy is very large (AH = 44-87 kJ mol-l) and consequently A S ranges from 25 to 159 J K-' mol-1 at 298 K. The increase in entropy because of the closer packing of the hydration shell in the excited states seems to be in agreement with the data for solvent exchange in UO;+. For (UO?j+)* we have rates of solvent exchange at room temperature of (2-5) x lo5 s-l. Nuclear magnetic resonance studiesL4 of the kinetics of H,O-exchange process in the equatorial positions of UO,(H,O)q+ allowed an estimation of a rate for all the coordinated water of 3.8 x lo6 s-' at 25 "C [9.5 x lo5 s-'; (U02)(H20)2,t with n = 41, a value that is approximately an order of magnitude higher than the rates of exchange in the excited states.Further support for the hydrolysis of the excited uranyl ion comes from the fluorescence spectra (fig. 2), which are more sensitive to pH than the absorption spectra. Furthermore, the quenching studies of (UOi+)* with Ag+ as a function of ionic strength at pH 2 show a dependence of the quenching rate on the ionic strength. This dependence reveals that in the quenching process there is an intervention of two ions with a + 1 charge.I5 Fig.5 shows that at pH 2 [UO,(OH)+]* is the dominant species, but even the other relevant species, (UOi+)* and UO,(OH),*, are present in equal amounts. At higher pH values (3 and 4) the fluorescence in the presence of Ag+ is perturbed by the formation of polynuclear species between uranium and silver.16 Under stationary conditions the fluorescence spectra of the uranyl ion contains the emission of the excited states U* and X*.5 For a reversible-crossing kinetic scheme the overall fluorescence yield is given by k;' + kj k",(k, + k,) '' = k, + ki - ki k,/(k, + k,)1744 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS where k , are the radiative rate constants of the two excited states of the uranyl ion (kg = 1.8 x lo2 s-l and kE = 0.80 x lo2 s - ~ ) .~ In order to test the model of reversible crossing the pH dependence of the fluorescence yields can be estimated through eqn (1) using the experimental decay rates. The agreement between the experimental and the calculated bF is good at all pH except at pH 0.5 (fig. 1). However, at [UOi+] = 0.02 mol dm-3 no decay values were experimentally determined at this pH. The values employed in the calculation where extrapolated from the experimental curves. The absolute yield at pH 3 is now lower, &. = 0.7 x than that previously reported (& = 1.1 x [NO;] = 0.04 mol dmP3) since the NO; concentration is higher (0.55 mol dm-3) and NO; increases some of the decay rates of the excited uranyl At a higher uranyl concentration (0.1 mol dm-3) the increase of bF around pH 3.5 is much less than at lower concentrations.In conclusion, we have shown that the effect of acidity on the photophysics of the uranyl ion can be attributed to the different decay rates of the several ionic species caused by the hydrolysis of (UOi+)*. All the data support the previously proposed reversible-crossing mechanism. This work was supported by INIC through the Research Centre QC-1. We thank G.T.Z. for the kind gift of the fluorimeter used in this work and one of the referees for useful suggestions. C. K. Jrargensen and R. Reisfeld, Struct. Bonding (Berlin), 1982, 50, 121. M. D. Marcantonatos, Inorg. Chim. Acta, 1977, 25, L101. M. D. Marcantonatos, J. Chem. Soc., Faraday Trans. 1, 1979,75, 2273. M. D. Marcantonatos, J. Chem. Soc., Faraday Trans. 1, 1980,76, 1093. S. J. Formosinho, M. G. Miguel and H. D. Burrows, J. Chem. Soc., Faraday Trans. 1, 1984, 80, (3/ 1282). M. Mavrodin-Tirabic, Rev. Roum. Chim., 1973, 18, 73; S. Pocev and G. Johansson, Acta Chem. Scand., 1973, 27, 2146. ’ M. Monyasu, Y. Yokoyama and S. Ikeda, J. Inorg. Nucl. Chem., 1977,39, 221 1. J. T. Bell and R. E. Biggers, J . Mol. Spectrosc., 1965, 18, 247. J. Burgess, Metal Ions in Solution (Ellis Horwood, Chichester, 1978), chap. 9-1 1. and 1971). lo L. G. Sillen and A. E. Martell, Stability Constants of Metal Ions (The Chemical Society, London, 1964 l1 M. R. Judkins, Ph.D. Thesis (University of California, 1967) [quoted in ref. (9), p. 336 and 3471. l2 C. Kiener, G. Folcher, P. Rigny and J. Virlet, Can. J . Chem., 1976, 54, 303. l3 M. D. Marcantonatos, M. Deschaux and F. Celardin, Chem. Phys. Lett., 1980, 64, 144. l4 Y. Ikeda, S. Soya, H. Fukutomi and H. Tomiyasu, J . Inorg. Nucl. Chem., 1979,41, 1333. l5 H. D. Burrows, S. J. Formosinho, M. G. Miguel and F. Pinto-Coelho, J. Chem. Soc., Faraday Trans. l6 H. D. Burrows, A. C. Cardoso, S. J. Formosinho and M. G. Miguel, to be published. I, 1976, 72, 163. (PAPER 3/ 1283)

ISSN:0300-9599    

DOI:10.1039/F19848001735

出版商:RSC    

年代:1984

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1984, 80, 1745-1756 Photophysics of the Excited Uranyl Ion in Aqueous Solutions Part 3.-Effects of Temperature and Deuterated Water: Mechanisms of Solvent Exchange and Hydrogen Abstraction from Water in Excited States BY SEBASTIAO J. FORMOSINHO* AND MARIA DA GRACA M. MIGUEL Chemistry Department, University of Coimbra, 3000 Coimbra, Portugal Receiued 25th July, 1983 The effects of temperature and D20 on excited uranyl decay in aqueous solution have been studied at several pH values and uranyl concentrations. The data are interpreted in terms of reversible crossing between two excited states of uranyl with different electronic distributions. The highest state, called U*, has a 71; 4; configuration whereas the lowest state, called X*, is nf, 8; (n = 1).The reversible decay is interpreted in terms of a solvent-exchange process. The values of AH+ and AS+ for this process reveal that solvent exchange proceeds by different mechanisms for the U* and X* states. The activation parameters are pH dependent, but the U* state always has more dissociative character for solvent exchange, whereas the X* state has more associative character. A weak effect on the solvent-exchange rates was found in D20. The rates of the irreversible decays are attributed mainly to a hydrogen-abstraction reaction from water molecules by excited uranyl ions. The activation energies and the frequency factors are pH dependent, revealing that hydrogen abstraction from the coordinated equatorial H 2 0 and from the more loosely bound water in axial positions can occur.The temperature and the deuterium isotope effects on the decays were analysed in terms of the current theories of radiationless transitions for the involvement of the 0-H and U=O stretching modes in the abstraction process. The autoquenching processes and the effect of the exciting intensity on the decays are in agreement with the proposed mechanisms. The photophysics of the excited uranyl ion in aqueous solution depends on many factors, including the acidity, uranyl concentration, nature and concentration of anions, intensity of excitation, temperature and deuterium isotope effect.l. * In spite of the complexity of such a system, the dynamic and stationary behaviour of uranyl fluorescence can reasonably be interpreted in terms of a reversible-crossing process between two almost isoenergetic electronic states, U* and X*, of [UOit,,,]*.Each of the excited states also decays by an unimolecular irreversible-decay pr0cess.l Decays of both kinds of excited states seem to be of a chemical nature. The reversible crossing is attributed to a solvent-exchange mechanism1 and the irreversible decays to a hydrogen-abstraction reaction from H,01-5 with rates that are pH dependent. A better understanding of the nature of such a process requires study of the effect of temperature and of D,O on the photophysics of excited uranyl at different pH values and at different uranyl concentrations. This paper presents the results of such studies and reveals that although in terms of the rates of decay both states U* and X* behave in a similar manner, under certain experimental conditions they can be involved in different types of mechanisms for solvent-exchange and hydrogen-abstraction reactions.17451746 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS EXPERIMENTAL Excited uranyl decays were studied on a nanosecond flash-photolysis apparatus with an N, laser (Aex = 337.1 nm) at constant laser intensity. For temperature studies the solutions were thermostatted in a water bath giving temperatures which were constant to f0.5 "C. Solutions in H,O were prepared with triply distilled water; solutions in D,O were prepared with deuterium oxide (Sigma, 99.8% D). Uranyl nitrate was the purest grade commercially available. Further experimental details have been given elsewhere.RESULTS The fluorescence lifetime of uranyl nitrate in water was measured as a function of temperature and uranyl concentration at various pH values between 1 and 4. Decay of the uranyl excited state in all cases was biexponential, but could be fitted to a model involving reversible crossing between the two lowest excited states of the uranyl ion1 ki u*ex*. kr Increasing temperature increased both decay rates, which followed good Arrhenius plots independent of the pH of the solution. However, at certain pH values a limiting rate was observed at temperatures > 40 "C (fig. 1). The decays were analysed in terms of absolute rate theory for a first-order rate, assuming that the decay arises from exchange of all coordinated solvent molecules, using the expression nkT h k,,, = -exp (ASTIR) exp (- A H t / R T ) where n is the number of coordinated water molecules in the excited uranyl ion.Entropies and enthalpies of activation calculated assuming a hydration number of 4, identical to that suggested for the ground state,6 are presented in table 1 for pH values between 1 and 4. Since all of the exchange rates are considered as first-order processes, the AS? values in table 1 are strictly applicable only to a dissociative-exchange reaction ( S , 1). However, if exchange occurs by a second-order process ASt will only decrease by 17 J K-l m01-l.~ The effect of added nitrate ion (0-1 mol dmP3) on the decay of the uranyl excited state was studied at pH 3 and 20 "C. The biexponential decay was analysed in terms of the reversible-crossing mechanism.The results presented in fig. 2 reveal that NO; increases ki and decreases k,. The decay of laser-excited uranyl ion luminescence was also studied at pH 3 in D,O solutions. The decay was biexponential at all temperatures over the range studied (3.5-56 "C) and was analysed in terms of a reversible-crossing mechanism. The rates of reversible crossing are smaller in D,O than in H,O, but show a temperature- dependent isotope effect. At 10 "C k,,,(H,O)/k,,,(D,O) = 1.0-1.1; the ratio then attains a maximum of 1.3 at 30 "C and decreases at higher temperature. Study of the effect of temperature on ki and k, (fig. 3) allowed an estimation of AH? and ASt as in H,O (table 1). The effect of temperature on the rates of the irreversible decays in the lowest excited states of uranyl u * e x * was also studied in H,O at different pH values and in D,O at pH 3.In H,O, the rates k, and k , increase with increasing temperature and follow good Arrhenius plots. AtS. J. FORMOSINHO AND M. DA G. M. MIGUEL 1747 13 h - I v) \ -Y, c I 12 1' I I I I I I .o 3.2 3.3 3 . 4 1 0 3 KIT Fig. 1. Arrhenius plots for the rates of reversible crossing ki (A) and k , (A) at (a) pH 1 and (b) pH 4; [UOi+] = 0.02 mol dm-3 and [NO;] = 0.15 mol drnp3. low [NO;] (0.04 mol dm-3) the rates are independent of temperature for temperatures Z 20 OC.' In D20, k, shows good Arrhenius behaviour but k , does not follow an Arrhenius plot. The solvent isotope effect for k, and k , at pH 3 is presented as a function of temperature in fig. 4. Nitrate ion concentration (fig.2) does not affect the rate k , but slightly increases the rate constant k,. The activation energies, E,, and the pre-exponential factors, A , for k , and k , in H,O are presented in table 2 for different pH values; k, in D20 has E, = 18.0 kJ mol-1 and A = 1.5 x lo8 s-l.1748 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS Table 1. Enthalpies and enthropies of activation for the rates of reversible crossing for the excited uranyl ion in H,O 1 .o 62 58 52 29 2.0 46 8.5 33.5 - 37.5 3 .O 67 75 50 17 3.0 (D20b solution) 46 2 27 - 58 4.0 42 - 17 29 - 58.5 a Estimated for the exchange of four solvent molecules; for six molecules A S is lower by Value obtained from glass electrode but not corrected for the change in the 3.5 J mo1-l K-l. autoprotolysis constant on going from H 2 0 to D,O.10 8 7 6 v) 0 - --- * 4 2 0 4 - I v) " 3 2 . * 2 I I I I I 0.2 0.4 0.6 0.8 1.0 [NO;l,/mol dm-3 ( b ) 0.2 0.4 0.6 0.8 1.0 [NOJ,/rnol dm-3 Fig. 2. Effect of NO; on the decay rates of excited uranyl at 20 "C and at pH 3.0; [UOi+] = 0.02 mol dm-3: A, ki; A, k,; a, k , ; 0, k,.S. J. FORMOSINHO AND M. DA G. M. MIGUEL 1749 14 13 h " I m --. v Y E - 12 11 lo3 KIT Fig. 3. Arrhenius plot for the rates of reversible crossing ki (A) and k, (a) in D,O at pH 3 ; [UOi+] = 0.02 mol drnM3. 2 0 20 40 60 TIo C Fig. 4. Temperature dependence of the isotope effect ratio in H,O and D,O for the rates of irreversible decay k, (0) and k, (0) at pH 3.1750 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS Table 2. Activation energies and pre-exponential factors for the rates k , and k , in H,O 1 .o 31 7 x 10'0 13.5 7.5 x 107 2.0 44 3 x 1013 42 4 x 1013 3 .O 24 2 x 109 24 2 x 109 4.0 30.5 4.5 x 1Olo 19 4 x lo8 DISCUSSION SOLVENT-EXCHANGE REACTIONS SOLVENT EXCHANGE versus TUNNELLING Analysis of the biexponential decay of the uranyl ion in water gave activation parameters for reversible excited-state crossing between U* and X*.The activation enthalpies determined for such processes are too large for a purely physical radiationless transition from U* to X* since U* is ca. 300 cm-l above X*,l and consequently such a physical transition to an almost isoenergetic state should have almost negligible AH+ (< 15 kJ mol-l). Furthermore, for a physical process k, would be expected to have AH! higher than AH] by ca.4 kJ mol-l, contrary to experimental observations (table 1). Ikeda et aZ.6 have studied by n.m.r. the kinetics of the water-exchange process for the equatorial positions of ground state of UO;t,,,. The exchange rate for all the coordinated molecules at 25 "C is estimated to be 4 x lo6 s-l with AH+ = 42 kJ mol-l. Both the rate constant and the activation enthalpy are comparable to the values found for ki and k,, supporting the idea that the reversible-crossing process may be due to a solvent-exchange mechanism. Further support comes from the effect of [NO;] on the exchange rates at pH 3 (fig. 2). Nitrate ion increases ki and decreases k,. Because solvent exchange is a special case of a nucleophilic substitution reaction in metal complexes, for a dissociative solvent-exchange mechanism the negative charge of the NO; ligand and its larger size, when compared with H,O or OH-, are predicted to increase the rate of solvent exchange.8 However, for an associative ( S , 2) mechanism NO, can decrease the rate of exchange.As table 1 shows, at pH 3 AS1 is very positive, suggesting that the solvent exchange in the U* state proceeds via a dissociative mechanism. This can be a D-type mechanism since AH1 is also high.g In contrast, the solvent exchange mechanism for the X* state has a positive AS: value and the AH: is also smaller than AH!. This suggests a weakly dissociative mechanism of the type In any case, the X* state has a more associative or assisted character and the U* state a more strongly dissociative character for solvent exchange.NATURE OF THE ELECTRONIC STATES OF THE EXCITED URANYL ION At all the pH values used the states U* and X* appear to exchange coordinated water by different mechanisms. For transition-metal complexes solvent exchange is currently interpreted in terms of crystal-field effects.*' Reasoning by analogy we can expect that for UO$+ and (UOE+)* crystal-field effects on the U atomic orbitals would play a role in the solvent-exchange processes. The small differences in the rates ki andS. J. FORMOSINHO AND M. DA G. M. MIGUEL 1751 k, suggest that the atomic orbitals involved in these processes have anf nature; d orbitals have crystal-field effects larger thanforbitals and affect solvent-exchange rates of transition-metal ions by several orders of magnitude.The involvement off orbitals is also in agreement with the studies of JerrgensenlO and Denning et aZ.,ll which consider the lowest vacant orbitals in UOi+ to be the Sforbitals 4, (f,,) and 6, (f,,). The fo andf,, orbitals are considered to be strongly involved in the U=O bonds. Whilst controversy exists over the actual hydration number of the uranyl ion in aqueous solution, and between 4 and 6 coordinated waters have been l3 it is generally agreed that the linear UOi+ ions are arranged at right angles to the water channels, the H,O molecules being in equatorial positions with respect to the linear In an excited hydrated (UOi+)* ion the approach of an H 2 0 molecule in an associative mechanism can occur on the plane in which the low-energy atomic orbitals are empty.'* The 4,(5f) orbitals have lobes pointing in an equatorial direction whereas the 6,(5f) orbitals have lobes pointing in directions making an angle of 45" with the equatorial plane.Consequently if in the excited state the 4, orbital is empty an associative mechanism for solvent exchange can be expected. However, the same cannot be said when the 6, orbital is empty but the 4, orbital is occupied by an un- paired electron, because repulsion between the lone-pair electrons of H 2 0 and the unpaired electron in 4; would hinder the approach of H,O molecules in the equatorial plane. This argument suggests that the state X*, with the more associative character, has the excited electron in a SL(5f) orbital and the state U* has the excited electron in a 4;(5f) orbital.In (UOi+)aq solvent-exchange rates are ca. 10 times higher than in (UOi+)&.6 This would not support the electronic transition 0, -+ 6, proposed by Denning et a1.l1 for the two lowest excited states of the uranyl ion. In fact, the 0, orbital has an electron density in an axial direction [p,(O) and fo(U) atomic orbitals] and consequently the 0, .+ 6, transition brings electron density from an axial position towards the equatorial plane. This will increase the repulsion between the lone pairs of H 2 0 and the U atom. The effect is equivalent to a decrease in the charge of the U atom, an effect which is known to increase the rate of solvent exchange.8v9 Such an effect was not expected by Jerrgensen1O for a n, + 6, or 4, transition because the n, orbital has a large electron density in the equatorial plane, mainly from the p + , atomic orbitals of the 0 atoms.However, the ground-state values were obtained b p H n.m.r., a technique which leads to considerable difficulties in the extraction of absolute solvent-exchange rates.g So the possibility that the electronic configurations of the lowest excited electronic states of (UO;') are n3,& (X*) and nt 4; (U*) needs to be treated with some care. The total angular momentum (0) components can be found for the possible electronic configurations of the excited states:l5$ l6 n3, 6; ,lI(i2 = 0 , l and 2) llI(i2 = 1 ) n3, 4; 3A(R = 1,2 and 3) lA(R = 2) ,r(R = 3,4 and 5 ) T(C2 = 4) 0; 6; ,A(Q = 1,2 and 3) lA(Q = 2) a; 4; "(0 = 2,3 and 4) l4(R = 3). For the first excited state spectroscopic data reveal R = 1 .ll In UO;+, because of large spin-orbit coupling, spin is not a good quantum number. If the spin states are combined, conserving the symmetry and the total angular momentum, the relevant states are x* 7Cn3,a; n(n = 1 ) U* nn3,4; A(i2 = 2) or r(Q =4).1752 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS ACIDITY EFFECTS The effect of pH on the fluorescence quantum yield and decay of excited uranyl has been interpreted in terms of the presence of different hydroxo and polymer ionic species that vary in concentration with pH., Each of these species has different rates of solvent exchange and the variation of the observed rates ki and k , with pH reflect such complexity.As table 1 shows, in spite of uncertainty in the coordination number of the water molecules, for all pH values the U* state always has a more dissociative solvent-exchange character than X* since AS! > AS$.The difference in the activation entropies varies between 60 and 29 J K-l mol-l, but is sufficiently large to be attributed to a difference in mechanism and not to a difference in the number of coordinated waters around the two states. Even for the more extreme situation between [UO,(H,O):+]* and [UO,(H,O)?+]* the difference in the entropies of activation would be only of the order of 14.5 J K-l mol-l. Consequently we will discuss the results of table 1 in terms of a mechanism of solvent exchange and consider that the pH dependence of AHt and AS? reflects the solvent-exchange mechanism of the dominant ionic species at each pH.Those species are considered to be at pH 1 [U02(H,0),2+]*, pH 2 [UO,(OH)(H,O)i]*, pH 3 [UO,(OH),(H,O),]* and pH 4 [(u02)2(0H),1**2 The guidance used to establish the mechanism of solvent exchange comes mainly from the activation parameters. Although the activation enthalpy of one particular cation has no diagnostic value, individual activation entropies may have. A broad distinction is generally made between markedly positive AS? for a dissociative process and markedly negative values, indicating an associative proces~.~ At pH 1 and 3 the solvent exchange seems to be of a dissociative character in both states although, as previously stated, with stronger character in the U* state. Support for the dissociative character of these processes comes also from the activation-enthalpy values which are quite high, revealing that in the breaking of the H,O-UOi+ bond there is virtually no contribution from the incoming solvent molecule to such breaking.The limiting rate attained at high temperatures (30-40 "C) also supports this view and should correspond to a situation where the breaking of the H,O-UOi+ bond is no longer the rate-determining step. At pH 2 and 4 AH+ values are smallest and AS+ are negative, suggesting an associative mechanism, at least for the X* state. For U* and at least at pH 2 for solvent exchange with a dissociative character, a strong interaction with the incoming ligand exists (I, mechanism). The changes in mechanism can be attributed to the nature of the ligands in the different ionic species.At pH 1 in [UO,(H,O)i+]* the presence of the electron-supplying ligand H,O favours a dissociation mechanism.* At pH 2 the effect of the negative charge of OH- in [UO,(OH)(H,O)$]* seems to be the dominant factor and favours a more associative mechanism. At pH 3 in the neutral species UO,(OH),(H,O), the effect of charge is less significant and dominates the electron-withdrawing character of the two OH- groups; such an effect favours a dissociative mechanism. Finally, at pH 4 the dominant species is a polynuclear cation where the strong bonding of the ligands, particularly those in the bridges between the U atoms, makes dissociation of the ligand molecules difficult and therefore favours an associative mechanism. EFFECT OF D20 The decay of the excited uranyl ion was studied in D,O at pH 3 and analysed in terms of reversible crossing between U* and X*.For both states there is much stronger interaction of the incoming D20 molecules in the transition state since the enthalpies of activation are ca. 20 kJ mol-1 lower and the AS? are ca. 75 J mol-1 K-' lower inS. J. FORMOSINHO AND M. DA G. M. MIGUEL 1753 D,O than in H,O. Even for cations in the ground state very little is known about D,O effects. In V02+ at pH 7 D,O slightly increases the dissociative character of the exchange process.17 However, V has a much smaller radius than U and consequently for the latter an associative mechanism is favoured. D,O is an enhancer of the struc- ture-making or -breaking of the different cations and a similar effect may also be caused in UOi+ through enhancement of the associative character for solvent exchange.lo Kemp et a1.ls have studied the decay of the uranyl excited state at various temperatures in both H,O and D,O, and although they noted a strong isotope effect, the results are not strictly comparable, as under their conditions only single-exponential decay was observed. This may have been a result of either increased excitation intensity or a different excitation wavelength. AUTOQUENCHING AND LASER-INTENSITY EFFECTS Uranyl concentration has also a very marked effect on the characteristics of the fluorescence decay of urany1.l As far as the solvent-exchange rates are concerned an increase in [UOi+] increases the quenching rate k, but leaves k, virtually unaffected at all pH values studied, with the exception of pH 1 where there is no influence on ki., In the U* state with a n3,#; configuration the half-filled fk3 orbital has all the lobes pointing in the equatorial direction and can strongly overlap the same type of empty orbitals of a ground-state UOi+ molecule.This one-electron bond is the main bond for the excimer formed with the two O=U=O axes parallel, since no significant binding can occur with the nu orbitals owing to the fact that they are completely filled in the UO:+ ground-state species. However, in the n3, S; configuration of the X* state, the half-filled SL(f+,) orbital makes an angle of 45" with the equatorial ~ 1 a n e . l ~ The overlap of these orbitals for the formation of the uranyl excimer is much weaker than in the U* state and consequently the possible excimer would have very low binding energy.It is consequently no surprise that uranyl affects only the rates of decay for the U* state (n3,4L) and the solvent-exchange rate ki. At low pH (ca. 1 .O) the high charge on excited and non-excited uranyl ions hinders the formation of the excimer. The temperature studies of the autoquenching rate for the exchange rate ki at pH ca. 3 gives AH? = 46 kJ mol-1 and AS? = 21 J mol-l K-l. When such values are compared with those in table 1 it can be seen that uranyl increases ki by increasing the associative character of the solvent exchange. This effect can be attributed to an increase in the radii of the central atom (bicentre) in the excimers and of the charge of the central atoms; these two factors favour a stronger interaction of the incoming solvent molecule in the transition state.8p9 For pH values of 3.0 and 4.0 an increase in the laser intensity decreases both rates of solvent exchange, ki and kr.lY2 At pH 1 these rates pass through a minimum., (UOi+)* absorbs at 337 nm, similar to the ground state.This may be because of promotion of a second electron nu -+ 4,, S,, as these orbitals are degenerate. Such a transition will strengthen the binding H,O-U, since the empty bonding orbital nu can accept the 0 lone pairs from H,O. Other energy transitions that have electron-transfer character from U to 0 in uranyl will have a similar effect. Absorption of the excited uranyl ion within the laser pulse promotes (UOi+)* to a higher excited state (UOi+)** where the solvation shell acquires a configuration with a slower exchange rate.Although (UOg+)** rapidly relaxes to the lowest electronic states (UOg+)*, the same might not happen with a hydration shell that remembers the higher excited electronic configuration during the (UOE+)* fluorescence decay. At pH 1 the slight increase in ki and k , at high laser intensities could be because of some more significant excitation of a double excited state involving the promotion of a ou electron.1754 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS HYDROGEN ABSTRACTION FROM WATER EFFECT OF TEMPERATURE AND DEUTERIUM ISOTOPE The decay process characterized by the rates k, and k , at pH 3 can clearly be analysed in terms of the current tunnelling theories of radiationless transitions.According to the model of Jortner,20 when the physical process is characterized in terms of single-mode formalism the temperature dependence of a non-radiative transition has an activationless region at low temperatures and an activation region at high temperatures. The transition temperature, T,, between the two regions is given by P-Pl 4s kTa = where cu is the frequency of the modes involved in the radiationless process, p = A E / h and S = A2/2; AE is the difference in energy of the vibrational levels involved in the transition and A is the horizontal reduced displacement of their minima. The activation energy is given by ?iu)(S-p)2 E, = 4s and the ratio of the frequency factor at high temperatures and the rate in the activationless region is A , exp (- s- YP) with y = In (p/S) - 1 .Good agreement between theory and experiment is obtained for k, and k, (pH 3) throughout the temperature range studied with Tim = 0.1 eV, S = 9.0 and p = 1.2. These are reasonable parameters.20 The most striking feature of excited uranyl decay is the fact that, at low [NO;], k , and k , have a very high transition temperature, T, z 298 K, whereas in organic molecules this temperature is ca. 100 K. In an organic molecule normally only one XH local mode is effective in the radiationless process, but for hydrated [UO,(OH),(H,O),]* w x 5000 cm-l. This shows that in UO:+ the high-frequency local OH modes (v z 3800 cm-l) and also the UO stretching frequency ( v = 745 cm-l) are the most relevant vibrations in these radiationless transitions.Support for the involvement of these modes comes from the heavy-water effect in k,. According to eqn (1) and (2) a decrease in the frequency of the OD mode (v = 2700 cm-l) should decrease T, and E,. The decrease in Ea of 1.33 would correspond to u) = 3770 cm-l if S and p remain the same. Such a value agrees well with the involvement of v(OD)+ v(U0) x 3500 cm-l. T, should also decrease to 225 K and consequently it is not observed under the present experimental conditions. However, for k , there is no clear definition between activation and activationless regions, a situation that is possible for X-D vibrations.20 Further support for the involvement of 0-H and U=O modes also comes from the deuterium isotope effect, which at 50 "C is ca. 2.0 for k, and k,.In contrast for aquo-ions such as Sm1I1 and EulI1, which are de-excited by OH vibrations, D,O decreases the fluorescence lifetimes by 20-50 times4 Theories of radiationless transitions5 show that deuterium isotope effects are quite small if the radiationless reaction coordinate includes vibrations of heavy atoms together with the UO modes.S. J. FORMOSINHO AND M. DA G . M. MIGUEL 1755 ACIDITY EFFECTS Several physical and chemical processes can be conceived for the irreversible decay in the excited uranyl i0n.~9~ One of them is electron transfer from the ligands to the uranium. Such processes can be described in terms of radiationless transition theories, but involve low-frequency phonon modes and, consequently, have very low activation energies,2o contrary to what happens here.The possibility of a purely physical non-radiative transition cannot be excluded under conditions that will be discussed later, but again such processes are expected to have low activation energies (< 15 kJ mol-l) in the temperature range of this study. The most favoured interpretati~n~-~q 21 is hydrogen abstraction from H 2 0 with significant charge-transfer character. The remaining OH radical recombines with the uranium(v) intermediate, producing no overall photochemical reaction. Theoretical support for such a mecha- nism has also been prod~ced.~ Hydrogen abstraction can occur from the more strongly bonded water in the equatorial position or from loose water molecules that approach towards an axial direction. The former situation can be envisaged and 0 II u-0, II I €4 O H in the equatorial field assumes pyramidal geometry that allows a closer approximation of the hydrogen atoms from the U=O bond.For this situation the activation energy should be high, because the distance of approach O-H.--O=U is not very small; in contrast the activation frequency factor can be of the order of the frequency of vibrations (ca. 1013 s-l). Although we ignore completely the detailed structure of the water molecules and the hydrogen bonds around the aquo-ions, it seems that this situation satisfies the data at pH 2. At pH 3 the rates k , and k , attain a minimum; the activation energies are smaller than at pH 2 (ca. 25 kJ mol-l) revealing that the H,O molecules get very close to the U=O bond and approach it with the appropriate geometry (linear approach).Those are probably water molecules coming from outside the first solvent shell towards an axial position. However, the activation frequency should now be much smaller. For the U* state the situation is intermediate between these two extremes at pH 1 and pH 4. However, for the X* state the acti,vation energies are very low (particularly at pH 1, Ea z 13 kJ mol-l) and the pre-exponential factors are also very low. It may be that this process is a physical non-radiative transition. The structure of the water molecules and OH- ions in a more rigid equatorial position may hinder the hydrogen-abstraction reactions, but favour the non-radiative physical transitions with stronger cation-solvent vibrations and higher matrix elements. The uranyl ion affects the rates of decay of the U* state more strongly and consequently also affects the rate k,.The temperature dependence of the autoquenching rate for k , studied at pH 3 gives Ea = 26 kJ mb1-l and A = 4 x 1O1O s-l. This reveals that excimer formation helps hydrogen abstraction from H,O in the U* state and in fact the observed E, and A values approach the values of the same parameters in the polynuclear cations at pH 4. The intensity of the laser leaves the rates k , and k , unaffected at all the pH values studied. This reveals that radiationless conversion from higher excited states is faster than the hydrogen-abstraction reaction and is mainly to the lowest U* and X* excited states.1756 PHOTOPHYSICS OF URANYL ION IN AQUEOUS SOLUTIONS CONCLUSIONS Reversible crossings between two excited states in condensed media are rare, but the delicate balance between all the rates under different conditions of pH, temperature, coordinated anions, uranyl concentration, ionic strength, nature of the solvent, exciting intensity etc.makes the excited uranyl ion an unique case. The quasi-equili- brium between the two lowest states, revealed by the biexponential decay, can be destroyed by several factors, but possibly the most significant of these is the nature of the coordinated ligands. Ligands can stabilize the 7zL& and n",sl, states by increasing or decreasing the energy gap, but are expected to increase the uranyl lifetimes by decreasing hydrogen abstraction from the coordinated waters and quench the reversible decay by decreasing the solvent-exchange rates.Under such conditions it is no surprise that the uranyl ion behaves as a molecular chameleon by presenting so great a variety of contradictory factors. In spite of these results, uranyl will continue to present new and unexpected features, particularly as it has > 20 different Q states in a very small energy region. The present studies reveal that some of these states can play a significant role on the photophysical processes in condensed phases and it is really a matter of luck that a great number of these features are amenable to analysis in terms of only two of those states, although the electronic configurations may be the determining factor. However, it is still not certain how many different low-energy configurations are important.We thank Prof. J. S. Redinha and Prof. H. D. Burrows for many helpful discussions and one of the referees for useful suggestions. We acknowledge also the financial support of INIC. S. J. Formosinho, M. G. Miguel and H. D. Burrows, J. Chem. SOC., Faraday Trans. I , 1984, 80, 1717. M. G. Miguel, S. J. Formosinho, A. C. Cardoso and H. D. Burrows, J. Chem. Soc., Faraday Trans. I , 1984,80, 1735. H. D. Burrows and T. J. Kemp, Chem. Soc. Rev., 1974,3, 139. C. K. Jrargensen and R. Reisfeld, Struct. Bonding (Berlin), 1982, 50, 121. H. D. Burrows and S. J. Formosinho, J. Chem. Soc., Faraday Trans. 2, 1977, 73, 201. Y. Ikeda, S. Soya, H. Fukutomi and H. Tomiyasu, J. Inorg. Nucl. Chem., 1979, 41, 1333. F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions. A Study of Metal Complexes in Solution (John Wiley, New York, 2nd edn, 1967), pp. 137 and 172. J. Burgess, Metal Ions in Solution (Ellis Horwood, Chichester, 1978), chap. 11 and p. 198. ' H. P. Bennett0 and E. F. Caldin, J. Chem. SOC., A , 1971, 2198. lo C. K. Jrargensen, J. Lumin., 1979, 18/19, 63. l 1 R. G. Denning, T. R. Snellgrove and D. R. Woodwark, Mol. Phys., 1979, 37, 1109. l 2 0. Ya. Samoilov, Structure of Aqueous Electrolyte Solutions and the Hydration of Ions (Consultants l3 V. A. Shcherbakov and L. L. Shcherbakova, Radiokhimiya, 1976, 18, 207; C. Gorller-Walrand and l4 L. E. Orgel, J. Chem. SOC., 1952, 4756. l5 G. Herzberg, Spectra of Diatomic Molecules (Van Nostrand, Princeton, 1967), p. 214. l6 C. Gorller-Walrand and L. G. Vanquickenborne, J. Chem. Phys., 1972, 57, 1436. l8 A. Cox, T. J. Kemp and W. J. Reed and 0. Traverso, J. Chem. Soc., Faraday Trans. 1,1980,76,804. l9 C . Becker, J . Chem. Educ., 1964,41, 358. 2o J. Jortner, J. Chem. Phys., 1976, 64, 4860; E. Buhks and J. Jortner, J. Phys. Chem., 1980, 84, 3370. 21 M. D. Marcantonatos, J. Chem. Soc., Faraday Trans. I , 1980,76, 1093. Bureau, New York, 1965). W. Colen, Chem. Phys. Lett., 1982, 93, 82. J. Reuben and D. Fiat, J. Am. Chem. SOC., 1969, 91, 4652. (PAPER 3/ 1284)

ISSN:0300-9599    

DOI:10.1039/F19848001745

出版商:RSC    

年代:1984

数据来源: RSC

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J. Chem. Soc., Faraday Trans. 1, 1984,80, 1757-1768 pH Dependence of Hydrogen Bonding in Complexes between Trimethyl-N-Oxide and Pentachlorophenol and Trifluoroacetic Acid in Acetonitrile B Y ZENON PAWLAK*t Chemistry Department, The University of Texas at Austin, Austin, Texas 78712, U.S.A. AND MARIA TUSK AND STEFAN KUNA Institute of Chemistry, University of Gdansk, 80-952 Gdansk, Poland AND FRANK STROHBUSCH Institut fur Physicalische Chemie der Universitat Freiburg, D-78 Freiburg, Federal Republic of Germany AND MALCOLM F. Fox School of Chemistry, Leicester Polytechnic, Leicester LE 1 9BH Receitled 1st August, 1983 Electrometric properties of the homocomplexes (CF,COO),H- and (C,Cl,O),H- and the heterocomplexes Me,NOH+A-, Me,NOH+AHA- and (Me,NO),H+A- (where HA is trifluoracetic acid or pentachlorophenol) have been investigated by the electrometric titration method (e.m.f.) in acetonitrile (AN).The formation constants, K,, of the above complexes have been determined. In the homocomplexes log K,[(CF,COO),H-] = 4.18 and log K,[(C,Cl,O),H-] = 3.5 in AN; for the heterocomplexes Me,NOH+A-, Me,NOH+AHA- and (Me,NO),H+A-, A- = CF,COO-, log Kf = 3.92, 5.52 and 6.82, and when A- = C,Cl,O-, log Kf = 4.2, 3.87 and 5.22, respectively. The acid pKtN was found to be 13.03 for CF,COOH and 16.46 for C,Cl,OH. For hydrogen-bonded complexes a gradual change in proton activity, paH, for a series of species passes through a maximum or a sigmoidal curve when paH(C, = CHA) =f(ApKtN). The sigmoidal and lambda curves which are produced when the formation constant and other parameters of charged and molecular complexes are plotted against ApK, in water and non-aqueous solvent are discussed.Hydrogen-bond formation of the charged species AHA- and BHB+ and the uncharged species BHA, with and without proton transfer, have been followed and a method for distinguishing the various phenomena has been developed.' Hydrogen-bond strength has been related to the base and acid enthalpie~,~-~ the gradual disappearance of the stretching band v(C=O) and gradual appearance of the asymmetrical band v(C00-) in the i.r. spectra,'-ll and to changes in the pH,12-14 proton chemical shift5-*, l5 and rate constant, k.16 These parameters depend upon the acid-base strength of the partners [ApKtN = pK,AN(acceptor) - pK,AN(donor)]. Some relative properties of the complexes are presented graphically as a function ApKSaolv t Present address : Institute of Economics and Organization of Production, University of Gdansk, 81-824 Sopot, Poland.17571758 - - - - HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE 17 -16 5 a 1s 14 13 c t corn plexes without corn plexes with proton transfer , proton transfer I 1 I I I I I I I I 1 -4 -3 - 2 - 1 0 1 2 3 4 S ApKFlV = [ pK;Oh’ (acceptor) - pKi0* (donor)] Fig. 1. Sigmoidal [-AH, A p , v(C0, COO-), Kes, k] and lambda (paH, dOH, Kj) curves describing complexes with and without proton transfer based on experimental literature data. Inflection points describe centrosymmetric proton location (see discussion in the text). , Results from this work. Exponent of hydrogen-ion activity, paH (C,,,, = Cacid), plotted as a function of ApKtN of the partners in acetonitrile. (1) Bu,N+C,C1,0-+C,C150H, (2) Me,NO + C,Cl,OH, (3) Me,NOH+ClO, + Bu,N+C,Cl,O-, (4) C,C150H + Me,NO, ( 5 ) N-MeImH+ClO; + Bu,N+C,Cl,O-, (6) P,H+ClO, + Bu,N+C,Cl,O-, (7) Bu,N+CFCOO- + CF,COOH, (8) hle,NO + CF,COOH, (9) Me,NOH+C10, + Bu,N+CF,COO-, (10) CF,COOH +Me,NO.The titrant occupies second place in each system and the number identifies the systems listed in table 2. Table 1. Properties of complexes containing hydrogen bonds complex without complex after experimental method, proton transfer, proton transfer, evaluated properties ‘X’ ‘Y’ reference calorimetry, -AH n.m.r. proton chemical shift, 6,, potentiometry, paH i.r. spectroscopy, Av(C=O/COO-) dipole moment, A,u equilibrium constant, Keg formation constant, Kf kinetic rate constant, k PaHW N PWX) N v(C=O)(X) -9 v(CO0-)(Y)z.PAWLAK et al. 1759 (fig. 1); those which show increasing values (as a sigmoidal curve) are the enthalpy, -AH, dipole moment, Ap, bond strength v(C=O, COO-), and equilibrium constant, Keq. In the same figure other relative properties of the complexes are shown which follow a lambda curve: chemical shift, dOH, proton activity, paH, and formation constant, Kf. The importance of these properties with respect to hydrogen bonding is summarized in table 1. In our previous ~ t u d i e s ~ * ? ~ ~ a series of reactions of Me,NOH++B (where B = Me,NO or an amine) in acetonitrile produced the following conclusions. (1) Protonated N-oxides and N-bases form very stable complexes in reactions with Me,NO [(Me,NO),H+, log Kf = 5.51 and (Me,NOHB)+, log Kf z 4 when pKtN(BH+) pKtN(Me,NOH+)].( 2 ) The titration curve, paH = f(CMeaNO/ CBZesNOH+) in CH,CN exhibits a sigmoidal shape with an inflection point at the 1 : 1 ratio and a 5 unit change in paH. This is unusual behaviour for homoconjugation. In the present work some of the complexes of an Me,NO base with trifluoroacetic acid and pentachlorophenol (HA) were studied. The possible products of the interaction of Me,NO with HA may be represented by the following scheme: Me3N0 * . . HA Me3N0 (excess) (Me3N0)*H+ . . A- \ !/ Me?NOH+ . . . A- complex Y \A (excess) Me,NOH+. . . AHA- complex Z complex XY VII A-+ HA AHA- The complex X should form in reactions (111) and (IV) when the acidity of HA is much lower than that of Me,NOH+; however, in our experiments the species Me,NO.- .HA was not observed. In the formation of complex Y the acidity of Me,NOH+ is lower than that of HA, but when an excess of Me,HO is added the homocomplex Z , (Me,NO),H+, is formed [reaction (V)]. Reaction (VI) proceeds from the formation of complex Y: when an excess of HA is added the complex XY, Me,NOH+AHA-, is formed. The anionic homocomplexes (AHA-) formed in reaction (VII), (CF,COO),H- and (C,Cl,O),H-, have been studied by means of the change in paH taking place during titration. Also, the reactions of C,Cl,O- with PyH+ and C,Cl,O- with N-MeImH+ were investigated. The paH values of the acid-base interreactions were then used to calculate formation constants (Kf) and proton-transfer constants (KpT).1760 HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE EXPERIMENTAL APPARATUS The e.m.f.was measured with a Precision potentiometer E353 (Metrohm Herisau). The reference half-cell was a calomel electrode filled with a 0.1 mol dm-3 solution of (C,H,),N+Cl- in acetonitrile and the salt bridge was filled with a 0.1 mol dm-3 solution of (C,H,),N+ClO; in acetonitrile. All measurements were carried out at 298 & 0.05 K. The titration cell initially contained 30 cm3 of ammonium perchlorate and the acid (HA) or salt (R,N+A-) and was thermostatted. The concentration of the titrant was 0.025 mol dmU3 and that of the solution placed in the cell was 0.001 mol dm-3. REAGENTS Acetonitrile was purified and vigorously dried.,, The liquid amines were dried over solid KOH and then distilled. Pentachlorophenol, picric acid and trifluoroacetic acid were purified by crystallization.The tetra-alkylammonium salts, perchlorates of N-bases and trimethylamine- N-oxides were prepared as described elsewhere.l2I l4 CALIBRATION OF THE GLASS ELECTRODE The reversibility of the glass electrode was checked by e.m.f. measurements in buffer solution containing C(C2H5)4NPI = 2.5 x (as a stock solution). The paH values of these solutions were calculated assuming complete dissociation of (C,H,),NPi in the dilute solution, pK$& = 1 1 .0.22 The activity coefficient was calculated from the expression -logf= 1.51 1:. On calibrating the glass sensor electrode in the paH region 7.0-12.0 (with 12 points) a linear relationship was obtained, with a slope of 75 mV (paH)-l. The Nernst slope of the glass electrode varies in non-aqueous indicating some irreversibility of the electrode process.mol dm-3 and picric acid, CHPi = 1.0 x For our electrodes we obtained the following relation: paH = (Eh-E)/W= (1266-E)/75 where W is the slope and Ei and E are the apparent potential of the reference electrode and the measured potential, respectively, in mV. RESULTS AND DISCUSSION DETERMINATION OF THE FORMATION CONSTANTS, Kf, AND PROTON-TRANSFER CONSTANTS, K ~ T In a solution containing a proton donor (HA, BH+, Me,NOH+) and a proton acceptor (A-, B, Me,NO) the following equilibria occur: K f Me,NOH+ + A- Me,NOH+A- HA + Me,NO & Me,NOH+A- Kuncor (Me,NO),H+A- or Me,NOH+AHA- (V, VI) where Kf and Qncor are the corrected and uncorrected formation equilibrium constants.The equations relating the hydrogen-ion activity, aH, the total analytical concen- trations of the base, Cb (Me,NO, B or A-), and acid, C, (Me,NOH+, BH+ or HA),Table 2. Formation constants, K,, proton-transfer constants, KpT, and acid dissociation constants, KtN, in acetonitrile at 298 K (Me,NOHfA-) (Me,NO.H .ONMe,)+A- (Me,NOH+AHA-) no. system (pKtN)” log KPT 1 2 3 4 5 6 7 8 9 10 A-+HA (16.46y Me,NO + HA Me,NOH++ A- HA+Me,NO (16.93)e N-MeImH’ (1 4.30)f + A- P,H+ (12.52)f+A- A- + HA (1 3.03)“ Me,NO + HA Me,NOH++ A- HA + Me,NO HA = C6C1,0H, A- = C6C1,0- 0 - - (AHA-): 3.50 (0.30) d 4.20 (0.42) - - +2.12 1.74, [3.86(0.26)]9 - - + 3.94 1. 16, [5.10(0.47)]g - - d - 5.98 (0.38) 3.80 (0.10) 3.87 (0.25) d - 5.22 (0.27) HA = CF,COOH; A- = CF,COO- 0 - - (AHA-): 4.18 (0.20) d - 6.40 (0.50) 4.20 (0.15) d 3.92 (0.32) - - d 6.82 (0.35) 5.52 (0.30) - a Titrant equals second species throughout; +6 is the standard deviation; this study; log KPT < 0; values from ref.(14); f values from ref. (21); g q n c o r .1762 HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE and the acid dissociation constant, Ka, are derived below: Kt(EA) =aHIA-lfA/[HAl e&+) = aH[BI/PH+lfRH Ca = [donor] + [complex] c b = [acceptor] + [complex] (3) (4) where the donor is HA, BH+ or Me,NOH+, the acceptor is A-, B or Me,NO and the complex is AHA-, BH+A-, (Me,NO),H+A-, Me,NOH+AHA- or Me,NOH+A-. Kf = [complex]/( Ca - [complex]) (Cb - [complex]). LAHA-] = (KkNCa-aH fA Cb)/(KkN -aH fA) ( 5 ) (6) To obtain [complex] we have [Me,NOH+A-l = (KtNCa fMe,NOH Cb)/(KkNfMe,NOH (7) If charge delocalization does occur (the proton-transfer equilibrium constants KPT are near to or greater than unity), then Kf requires a correction of the form log Kf = log encor - log K p T .(8) The Qncor values were also calculated from eqn ( 5 ) . The proton-transfer equilibrium constant can be calculated from KPT = K t (donor)/ K t (accep tor). (9) Values of Kf and KPT are given in table 2, and a plot of paH(C, = Ca) against ApK,AN is shown in fig. 1. The dominant factors affecting paH changes occurring during acid-base titration are the acidity of the interacting substances and the value of the formation ~0nstants.l~ Acetonitrile [pKtN(Me,NOH+) = 16.931 is the weakest acid studied here (see table 2), with a tendency to form very stable complexes, log K,[(Me,NO),H+] = 5.51.14 THE! HETEROCOMPLEX Me,NO + C,Cl,OH (OR CF,COOH) The paH of trimethylamine-N-oxide on titration with C,CI,OH or CF,COOH in CH,CN (fig.2, curves 2 and 8) exhibits a decrease with an inflection point at the 1 : 1 ratio: Me,NO,,,,,, + C,Cl,OH + (Me,NO),H+C,C1,0- (V) C,Cl,OH 11 (excess) Me,NOH+(C,Cl,O),H-. (VI) After the equivalent point an excess of titrant causes a smaller decrease in paH in both cases. However, the change is sufficiently pronounced to indicate interaction. Also, after the inflection point curves 2 and 8 have identical slopes, corresponding to the formation of the homocomplexes (C,Cl,O),H- and (CF,COO),H- (curves 1 and 7). The reverse titrations of pentachlorophenol and trifluoroacetic acid with Me,NO as a titrant (fig.2, curves 4 and 10) show a small initial rise in paH up to the 1 : 1 ratio, reaction (11), and a sharp increase in paH with an excess of Me,NO, reaction (V), as a result of formation of the very stable complex (Me,NO),H+.14z. PAWLAK et al. 1763 20 18 2 16 14 12 -0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8 log (Cbase/Cacid (curves 4.9 and 10); log (Cacid/Cbase) (curves I , 2, 7 and 8) Fig. 2. Relationship between the exponent of hydrogen-ion activity, paH, as a function of log (Cbase/Cacid) [or log (Cacid/Cbase)] in acetonitrile. (1) Bu,N+C,Cl,O- + C,Cl,O,H, (2) Me,NO + C,Cl,OH, (3) Me,NOH+ClOh + Bu,N+C,C1,0- [shape of curve as in (9)], (4) C,Cl,OH + Me,NO, (5) N-MeImH+ClO; + Bu,N+C,Cl,O- [shape of curve as in (9)], (6) P,H+ClO, + Bu,N+C,Cl,O- [shape of curve as in (9)], (7) Bu,N+CF,COO- + CF,COOH, (8) Me,NO + CF,COOH, (9) Me,NOH+ClO; + Bu,N+CF,COO-, (10) CF,COOH + Me,NO.The titrant occupies second place in each system and the number identifies the systems listed in table 2. THE HETEROCOMPLEX Me,NOH+ + C,c1,0- (OR CF,COO-) A plot of paH against log (C,-/CMesNOH+) shows the shape of a sigmoidal titration curve when Me,NOH+ClO, solution is titrated with (C,H,),N+C,Cl,O- or (C,H,),N+ CF,COO-, (fig. 2, curves 3 and 9). At ca. 2 paH units for the 1 : 1 acid/base ratio [reaction (I)] a point of inflection is observed: Me,NOH+ + c,c~,o- + Me,NOH+C,Cl,O-. (1) For two of the systems studied the ApK,AN values are negative (-0.47 and - 3.90). The sigmoidal titration curves indicate the existence of very stable complexes.These have never been observed for the formation of molecular complexes from charged species.l THE HOMOCOMPLEXES (C6C1,0),H- AND (CF,COO),H- The quaternary salts (C,H,),N+C,Cl,O- and (C,H,),N+CF,COO- exhibit a paH change with a linear relationship [reaction (VII)] on titration with their parent acid. 58-21764 HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE The much greater change in proton activity occurring during complex formation (CF,COO),H- can be explained in terms of the formation of a species that is more stable than (C,Cl,O),H-. Additionally, the following reactions have been studied : N-methylimidazole H+CIO; + (C4Hs),N+c,Cl,0-, ApKfN = 2.16 pyridine H+ClO; + (C4Hs)4N+c,Cl,0-, ApK,AN = 3.94.The paH curves show a sigmoidal shape (fig. 2, curves 5 and 6), but for both reactions ApK,AN is positive; thus proton transfer is indicated. INTERPRETATION OF SIGMOIDAL AND LAMBDA PROTON-TRANSFER CURVES Several experimental methods have supplied evidence for the existence of separate X and Y species. When graphically presented, the measured properties of the complexes {So,, AH, v(C=O, COO-), Ap, paH, Kf and Keq asfTApK,AN)[pK,AN (acceptor) -pK,AN (donor)]] can be described by one of two types of curves : lambda 'A'5-8912-14 or sigmoidal (fig.1). The results of our studies of changes in paH(Cdono, = Cacceptor) are presented in fig. 1 and are described by a lambda curve. The inflection points of the lambda and sigmoidal curves in fig. 1 are observed at ApKSa0lv(critical) = 0 for the measured properties of these complexes when they are plotted as a function of acidity in a non-aqueous medium.This has been recognized by us in previous papers.'? 8 * 12-14 From experimental data it is known that no simple quantitative correlation exists between the pK, of acids in organic polar solvents and in water.23 However, for the same class of acids, ApK, is approximately constant [ApK, = pKs,olv - pKFzo]. For example, ApK, values of ammonium ions are ca. 8 in acetonitrile21 and ca. 3 in acetone;13 their values for phenols are ca. 13.0 in CH3CN,23 ca. 12.5 in acetone12 and ca. 12 in propylene ~arbonate;,~ the ApK, values of aromatic acids in acetonitrile range from 14.1 to 16.5.23 For most slightly polar solvents such as eel,, C,H, or C,H,Cl acidity cannot be defined in the conventional manner because of the non-existence of ionic equilibria.The decrease in acidity is much larger for molecular acids than for cationic acids. The contributing factor in the deviation for some ortho-substituted acids is that the steric effect is often greater in organic solvents than in water. Data presented for water show ApK,HzO(critical) in the range - 2 to 8,lv 25 the value of ApK,(critical) indicating centrosymmetric hydrogen-bonded complexes. The values are higher in less polar solvents than those which were determined in polar aprotic solvents or in the solid phase. ApK,HzO(critical) values are collected in table 3. In less polar solvents ( E < 10) values of ApK,HZ0(critical) are as follows : RCOOH +aliphatic N-bases, 7.7-5.8; RCOOH +aromatic N-bases, 5.2-3.5; RCOOH + aromatic amine- N-oxides, 1.6; ArOH +aromatic N-bases, 5.0-4.0.In polar aprotic media such as acetone, acetonitrile, nitrobenzene or propylene carbonate values of the ApK,Hz0 are as follows: RCOOH+aliphatic N-bases, 3; ArCOOH + aromatic N-bases, 3.25; RCOOH + ArO-, 2.5-3.5 ; ArCOOH + ArO-, 2.0-2.5; Me3NOH + aliphatic N-bases, 4.1. The values of ApPZlv(critical) reach zero for the systems RCOOH + RCOO-, ArCOOH + ArCOO-, ArOH + ArO- and PyH+ +aliphatic N-bases. What conclusions can we draw from the ApK,HZ0(critical) data? In less polar solvents ( E < 10) aliphatic carboxylic acids are ca. 10, weaker than protonated N-bases, but only ca. lo2 weaker than pyridine-N-oxide. Aromatic carboxylic acids are ca.lo4Table 3. Values of ApKFzO [ = pK,Hzo (acceptor) - pK?zo (donor)] corresponding to centrosymmetric hydrogen-bonded complexes for various solvent systems with solvent acid base less polar more polar without solvent RCOOH aliphalic N-bases aromatic N-bases aromatic amine-N-oxides RCOO- ArO- ArO- ArCOO- ArCOOH aliphatic N-bases ArOH aromatic N-bases ArO- Me,NOH+ aliphatic N-bases PyH+ aliphatic N-bases aromatic N- aromatic N-bases base H+ 7.7a(B); 7.5b(B); 5.8,(F) 5.2'(B) ; 4. sd(B) ; 3. Y(F) ; 1.62;f 1.689 (B,F,G,H) Oh(S); 2.3i(Q); 3.743(4) ca. Ok(M); ca. Oz(K); 3.2 5 O ( K) 2.0m(L); 2.5%(K); CU. O m 7 n(K, L)* ca. Ok(M)* 2.P(L); 2.Y(K); 3Sn(I) 5.0Q(C, D); 4.OQ(E); 5.0r(B); 1.3P(S) 4.5p(A) ca. O"(K); ca. Ok(M)* 4.1t(K); ca. Ot(K)* CQ.OU(K)* ca. OW(K)* N (A) Carbon tetrachloride (E = 2.2); (B) benzene (2.3); (C) cyclohexane (2.0); (D) toluene (2.4); (E) trichloroethylene (3.5); (F) chloroform (4.8); (G) =hlorobenzene (5.6); (H) 1,2-dichloroethane (8.9); (I) acetonitrile + 1,2-dichloroethane (1 : 1, v/v); (J) acetone (20.7); (K) acetonitrile (36); (L) nitrobenzene (34.8); (M) propylene carbonate (65); (S) solid; (Q) liquid. aT. Duda and M. Szafran, Bull. Acad. Pol., Ser. Sci. Chim., 1978, 26, 207; L. Sobczyk, Z. Pawe€ka, Roczn. Chem., 1973, 47, 1523; ref. :15); G. M. Barrow, J . Am. Chem. SOC., 1956,78, 5802; f ref. (6a); Q ref. (6b); J. Pietrzak, B. Nogaj, Z. Dega-Szafran and M. Szafran, Acta Phys. Pol.. Ser. A , 1971, 52, 779; ref. '7); ref. (1 1); O G. Zundel and A. Negyrevi, J. Phys. Chem., 1978, 82, 685; P ref.(19); Q J. Jadzyn and J. Malecki, Acta Phys. Pol., 1972, 5, 599; Z. Dega-Szafran and E. Dulewicz, Adu. Mol. Relax. Proc., 1981, 21, 207; ref. (3); ref. (16); R. Lindeman and A. Zundel, J . Chem. SOC., Faraday Trans. I , 1977, 73, 788; j ref. (9); ref. (17); ref. (9); ref. (14); ref. (10); ref. (21); * on a non-aqueous scale 1.1766 HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE weaker, but phenols are ca. lo5 weaker than protonated N-bases. We do not have an absolute scale of acidity in slightly polar solvents, but a knowledge of the differences among a group of acids is very helpful in planning experiments with these acid-base groups. For polar solvents, such as acetonitrile and nitrobenzene, aliphatic acids are weaker than phenols by a factor of 102-103, but protonated N-bases are ca.lo4 weaker than trimethylamine-N-oxide charged species. Also, phenols are ca. 10, weaker than benzoic acids but protonated-aliphatic N-bases are ca. lo3 weaker than benzoic acids. The ApK,HZO(critical) data give a better estimate of acidity and basicity on the molecular level than p e l v , since the latter function is much more influenced by electrostatic interactions with the solvent. From the results given in table 2 the formation constants, K,, may be calculated for three reactions: Me,NO + HA(titrant), Me,NOH+ClO,- + R4N+A-(titrant) and HA + Me,NO (titrant). The tendency for Me,NO and A- (C,Cl,O-, CF,COO-) to undergo homoconjugation [uiz. (Me,NO H ONMe,)+, log Kf = 5.51 ; (CF,COO),H-, log Kf = 4.18; (C,Cl,O),H-, log Kf = 3.101 has a big influence on the paH titration curve.In the reaction of Me,NO with C,Cl,OH and CF,COOH as titrants (fig. 2, curves 2 and 8) two types of complexes are formed before and after the equivalent point: before after (Me,NO - H . ONMe,)+C,CI,O-, Me,NOH+C1,0 * H * OC,Cl;, log Kf = 5.98 log Kf = 3.80 (Me,NO * H .ONMe,)+CF,COO-, Me,NOH+CF,COOH -OOCF;, log Kf = 6.40 log Kf = 4.20 In the reverse reaction of C6C1,0H and CF,COOH with Me,NO as a titrant (fig. 2, curves 4 and 10) the same complexes are formed in the reverse order. From these curves the formation constants log Kf = 5.22, 3.87, 6.82 and 5.52 are calculated (complexes listed in the same order as above). The difference in the values obtained from both kinds of titration is ascribed to the irreversible response of the glass electrode.before after Me,NOH+C,Cl,O - H * OCl,C;, log Kf = 3.87 Me,NOH+CF,COO . H - OOCF,, log Kf = 5.52 (Me,NO * H - ONMe,)+C,Cl,O-, log Kf = 5.22 (Me,NO * H . ONMe,)+CF,COO-, log K, = 6.82 Taking into account the influence of the basicity of the anion A- (C,CI,O- and CF,COO-) on the formation constant for Me,NOH+A-, some differences are expected. The more basic C,CI,O- forms a more stable complex (Me,NOH+C,C1,0-, log Kf = 4.2) than CF,COO-(Me,NOH+CF,COO-, log Kf = 3.92). For both reactions the sigmoidal titration curve (ApKkN is negative, no proton transfer occurs) shows the formation of stable complexes when two charged species are interacting [reaction (111 *z. PAWLAK et al. 1767 The influence of the basicity of C,CI,O- and CF,COO- was observed on the formation of (Me,NO),H+A- complexes.Table 2 shows that the values of Kf increase in the order: CF,COO-> C,Cl,O-. The same tendency is observed for the Me,NOH+ AHA- complexes. The reactions of N-MeImH+ClO; and PyH+ClO, with R,N+C,C1,0- as the titrant (with positive ApK,) show sigmoidal titration curves, and the following formation constants were obtained : log Kf(C,C1,OH-.-N-MeIm) = 1.70 log Kf(C,Cl,OH...Py) = 1.16. CONCLUSIONS It is expected that proton transfer of the form HA + B takes place when ApPaol" > 0 and a sigmoidal paH curve will result. This condition is not valid in an aprotic solvent when strong hydrogen bonding takes place. We have shown that complexes such as (Me,NO-H.ONMe,)+, for which Kf = 5.51, ApK,AN = 0 and which show high formation constants, demonstrate an inflection point with a change in pH of 5 units.14 Furthermore, we found systems with an inflection point when ApK,AN < 0, e.g.the reactions Me,NOH+ + C,CI,O- (ApK, = - 0.47) and Me3NOH+ + CF,COO- (ApK, = -3.90) with formation of Me,NOH+C,Cl,O- (log Kf = 4.20) and Me,NOH+CF,COO- (log Kf = 3.90). Anionic bases such as C,CI,O- and CF,COO- form slightly stronger complexes with Me3NOH+ than N-bases.14 We can predict the value of ApK,HzO(critical) from the expression14 We thank Prof. J. J. Lagowski, Ms Debora Bittaker and Mrs Katherine Mueller for their assistance in the preparation of this manuscript. 1 2 3 4 5 6 7 8 9 10 11 1 2 13 14 15 16 17 18 19 20 (a) Th. Zeeger-Huyskens and P. Huyskens, in Molecular Interactions, ed.H. Ratajczak and W. J. Orville-Thomas (Wiley, New York, 1981), vol. 2, pp. 1-97; (6) J. Emsley, Chem. Soc. Rev., 1981, 9, 91. L. Lamberts and P. Huyskens, Proc. 1st Int. Con$ Calorimetry Thermodynamics, Warsaw, 1949, p. 849. Z. Pawlak and R. G. Bates, J. Chem. Thermodyn., 1982, 14, 1035. Z. Pawlak, L. M. Mukherjee and R. G. Bates, J. Chem. Thermodyn., 1982, 14, 1041. M. Ilczyszyn, L. Le-Van, H-Ratajczak and I. M. Ladd, in Proton and Ions Involved in Fast Dynamic Phenomena, ed. P. Laszlo (Elsevier, Amsterdam, 1978), p. 257. B. Brycki, Z. Dega-Szafran and M. Szafran, (a) Adv. Mol. Relax. Proc., 1979,15,71; (6) Pol. J . Chem., 1980, 54, 221. Z. Pawlak, J. Magonski and T. Jasinski, J. Mol. Struct., 1978, 47, 329. J. Magonski and Z. Pawlak, J. Mol. Struct., 1982, 80, 243. S. L. Johnson and K. A. Rumor, J. Phys. Chem., 1965,69, 74. R. Lindeman and G. Zundel, J. Chem. SOC., Faraday Trans. 2, 1972,68, 979. Z. Pawlak and J. Magonski, J. Mol. Struct., 1980, 60, 179. Z. Pawlak, B. Nowak and M. F. Fox, J. Chem. SOC., Faraday Trans. 1, 1982, 78, 2157. S. Kuna, Z. Pawlak and M. Tusk, J. Chem. SOC., Faraday Trans. I , 1982, 78, 2685. Z. Pawlak and A. Wawrzynow, J. Chem. SOC., Faraday Trans. I , 1983, 79, 1523. N. S. Gotubow and S. F. Burejko, Adv. Mol. Relax. Proc., 1981, 24, 995. Z. Pawlak, M. F. Fox, M. Tusk and S. Kuna, J. Chem. SOC., Faraday Trans. I , 1983, 79, 1987. H. Ratajczak and L. Sobczyk, J. Chem. Phys., 1969, 50, 556. L. Sobczyk and 2. Pawelka, J. Chem. Soc., Faraday Trans. I , 1979, 70, 832. Z. Malarski, M. Rospenk, L. Sobczyk and E. Grech, J . Phys. Chem., 1982,86,401. R. Nouwen and P. Huyskens, J. Mol. Struct., 1973, 16, 459.1768 HYDROGEN BONDING OF COMPLEXES IN ACETONITRILE 21 Z. Pawlak, G. Zundel, J. Fritsch, A. Wawrzynow, S. Kuna and M. Tusk, Electrochim. Acta, g2 J . M. Kolthoff and M. K. Chantooni Jr, J , Am. Chem. SOC., 1965,87,4428. 23 J. F. Coetzee, in Progress in Physical Organic Chemistry, ed. A. Streitwieser Jr and R. W. Taft (Wiley, 24 Z. Pawlak and J. Magonski, J. Chem. Soc., Faraday Trans. I , 1982,78,2807. 25 S. N. Vinogradov, in Molecular Interactions, ed. H. Ratajczak and W. J. Orville-Thomas (Wiley, New 1984, 29, in press. New York, 1967), voi. 4, pp. 45-92. York, 1981), vol. 2, pp. 179-229. (PAPER 3/ 1348)

ISSN:0300-9599    

DOI:10.1039/F19848001757

出版商:RSC    

年代:1984

数据来源: RSC