摘要:

Ordinary Members PROFESSOR R. J. DONOVAN 1983 PROFESSOR M. C. R. SYMONS 1983 DR G. J. HILLS 1984 PROFESSOR J. M. THOMAS 1983 PROFESSOR A. J. LEADBETTER 1984 DR J. ULSTRUP 1985 DR I . W. M. SMITH 1985 PROFESSOR G. WILLIAMS 1985 PROFESSOR F. L. SWINTON 1983 DR D. A. YOUNG 1984 Honorarj, Secretarj-: DR G. J. HILLS Honorarj- Treasurer : PROFESSOR P. GRAY The President thanked the retiring members of Council, Vice-presidents Professor Sheppard and Professor Wagner, and Ordinary Members Professor King and Professor Purnell, for their services. 5. Reriew of Futurr Acfirifies A programme of future activities of the Division had been tabled and the President drew attention to the forthcoming General Discussions and Symposia. xiOrdinary Members PROFESSOR R. J. DONOVAN 1983 PROFESSOR M. C. R. SYMONS 1983 DR G. J. HILLS 1984 PROFESSOR J. M. THOMAS 1983 PROFESSOR A. J. LEADBETTER 1984 DR J. ULSTRUP 1985 DR I . W. M. SMITH 1985 PROFESSOR G. WILLIAMS 1985 PROFESSOR F. L. SWINTON 1983 DR D. A. YOUNG 1984 Honorarj, Secretarj-: DR G. J. HILLS Honorarj- Treasurer : PROFESSOR P. GRAY The President thanked the retiring members of Council, Vice-presidents Professor Sheppard and Professor Wagner, and Ordinary Members Professor King and Professor Purnell, for their services. 5. Reriew of Futurr Acfirifies A programme of future activities of the Division had been tabled and the President drew attention to the forthcoming General Discussions and Symposia. xi

ISSN:0300-9599    

DOI:10.1039/F198278FX025

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

3 708 REVIEW OF BOOKS is the absence of any reference to possible new and potentially significant applications for polymer latices. Novel applications may well be found in at least two directions, namely, those which exploit the large polymer-aqueous-phase specific surface area of latices, and those which exploit the electrical dissymmetry which is present at the interface between polymer and aqueous phase in the case of electrostatically stabilised latices. No reference is made in this book to the efforts which have so far been made to exploit for medical purposes the adsorptive and binding potentialities of the large area of polymer-aqueous-phase interface in latices. Nor is there any mention of possible catalytic applications of this large interfacial area. So far, catalytic applictions have been confined to those which rely essentially upon enhancement of the counter-ion concentration in regions of the electrical double layer which are near to the polymer surface.However, it is at least possible that the adsorptive capacity of the interface may also be useful in catalytic applications. Some discussion of possibilities such as these would have been welcome. D. C. BLACKLEY Received 14th April, 1982 Shock Waves in Chemistry. Ed. by ASSA LIFSHITZ. (Marcel Dekker, New York, 1981). Pp. ix + 390. Price SFr 182. After a somewhat hesitant start, the use of shock waves to study chemical and physical processes at high temperatures has become an accepted technique and reliable kinetic data can be obtained in this way. Several books have been written, notably by Bradley and by Gaydon and Hurle, which describe not only the underlying principles and the experimental procedures but also give some account of the early results obtained using shock waves to provide high temperatures for short, well defined times in the reactant gases.Inevitably, these books have become rather dated. This new book, edited by Lifshitz, is rather different. It is a collection of self-contained review articles on various aspects of shock waves. The first (by Khandelwal and Skinner) is concerned with hydrocarbon oxidation, and the next (by Tsang) describes the results obtained using the comparative rate technique which he has pioneered. Both these articles include extensive lists of references and represent useful summaries of the present situation.Boyd and Burns have contributed a chapter on dissociation-recombination reactions, while Kiefer describes the laser-schlieren method which he has done so much to develop. There is another chapter by an acknowledged expert, Just, on atomic resonance absorption spectrometry. Under shock-tube conditions it is very seldom that the concentrations of radicals and other species reach a steady state, and so the classical Bodenstein steady-state approximation cannot be used. Instead, it is necessary to integrate the differential equations describing the time-variation of species concentration, and Gardiner, Walker and Wakefield have provided a useful guide to the computational procedures available in this and other aspects of shock-tube work.In addition to these contributions there is another by Bar-Nun on Chemical Aspects of Shock Waves in Planetary Atmospheres which, although interesting in itself, fits rather uneasily with its companions. As is inevitable in a book of this type the standard and style of the chapters varies and there is some overlapping material; none of this, however. represents a serious drawback. What is more difficult to understand is the audience for whom the book is intended. Each chapter is a useful and interesting review which will appeal to a fairly restricted readership, but, in the opinion of this reviewer, the whole volume lacks coherence. The time-honoured phrase ‘should be on the shelves of every library’ probably applies, though the price, over &50 at the current exchange rate, must cause all university librarians to flinch in these days of U.G.C.cuts. There is still room for the definitive up-to-date book to be written on shock waves in chemistry. J. A. BARNARD Received 5th April, 19823 708 REVIEW OF BOOKS is the absence of any reference to possible new and potentially significant applications for polymer latices. Novel applications may well be found in at least two directions, namely, those which exploit the large polymer-aqueous-phase specific surface area of latices, and those which exploit the electrical dissymmetry which is present at the interface between polymer and aqueous phase in the case of electrostatically stabilised latices. No reference is made in this book to the efforts which have so far been made to exploit for medical purposes the adsorptive and binding potentialities of the large area of polymer-aqueous-phase interface in latices.Nor is there any mention of possible catalytic applications of this large interfacial area. So far, catalytic applictions have been confined to those which rely essentially upon enhancement of the counter-ion concentration in regions of the electrical double layer which are near to the polymer surface. However, it is at least possible that the adsorptive capacity of the interface may also be useful in catalytic applications. Some discussion of possibilities such as these would have been welcome. D. C. BLACKLEY Received 14th April, 1982 Shock Waves in Chemistry. Ed. by ASSA LIFSHITZ. (Marcel Dekker, New York, 1981). Pp. ix + 390.Price SFr 182. After a somewhat hesitant start, the use of shock waves to study chemical and physical processes at high temperatures has become an accepted technique and reliable kinetic data can be obtained in this way. Several books have been written, notably by Bradley and by Gaydon and Hurle, which describe not only the underlying principles and the experimental procedures but also give some account of the early results obtained using shock waves to provide high temperatures for short, well defined times in the reactant gases. Inevitably, these books have become rather dated. This new book, edited by Lifshitz, is rather different. It is a collection of self-contained review articles on various aspects of shock waves. The first (by Khandelwal and Skinner) is concerned with hydrocarbon oxidation, and the next (by Tsang) describes the results obtained using the comparative rate technique which he has pioneered.Both these articles include extensive lists of references and represent useful summaries of the present situation. Boyd and Burns have contributed a chapter on dissociation-recombination reactions, while Kiefer describes the laser-schlieren method which he has done so much to develop. There is another chapter by an acknowledged expert, Just, on atomic resonance absorption spectrometry. Under shock-tube conditions it is very seldom that the concentrations of radicals and other species reach a steady state, and so the classical Bodenstein steady-state approximation cannot be used. Instead, it is necessary to integrate the differential equations describing the time-variation of species concentration, and Gardiner, Walker and Wakefield have provided a useful guide to the computational procedures available in this and other aspects of shock-tube work.In addition to these contributions there is another by Bar-Nun on Chemical Aspects of Shock Waves in Planetary Atmospheres which, although interesting in itself, fits rather uneasily with its companions. As is inevitable in a book of this type the standard and style of the chapters varies and there is some overlapping material; none of this, however. represents a serious drawback. What is more difficult to understand is the audience for whom the book is intended. Each chapter is a useful and interesting review which will appeal to a fairly restricted readership, but, in the opinion of this reviewer, the whole volume lacks coherence. The time-honoured phrase ‘should be on the shelves of every library’ probably applies, though the price, over &50 at the current exchange rate, must cause all university librarians to flinch in these days of U.G.C. cuts. There is still room for the definitive up-to-date book to be written on shock waves in chemistry. J. A. BARNARD Received 5th April, 1982

ISSN:0300-9599    

DOI:10.1039/F198278BX027

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

JOURNAL OF THE CHEMICAL SOCIETY FARADAY TRANSACTIONS, PARTS I AND I 1 The Journal of The Chemical Societ?? is issued in six sections: Journal of The Chemical Society, Chemical Communications Journal of The Chemical Society, Dalton Transactions Journal of The Chemical Society, Faraday Transactions, I Journal of The Chemical Society, Faraday Transactions, I I Journal of The Chemical Society, Perkin Transactions, I Journal of The Chemical Society, Perkin Transactions, I I Thus, five of the sections are directly associated with three of the Divisions of The Royal Society of Chemistry: the sixth is Chemical Communications. This continues to be the medium for the publication of urgent, novel results from all branches of chemistry. Communications should not normally exceed one printed page in length and authors are required to submit three copies of the typescript and two copies of a statement of the reasons and justification for seeking urgent publication of the work.This Section is intended to be essentially a journal for inorganic chemists containing papers on the structure and reactions of inorganic compounds and the application of physical chemistry techniques to, e.g. the study of inorganic and organometallic compounds and problems, including work on the kinetics and mechanisms of inorganic reactions and equilibria, and spectroscopic and crystallographic studies of inorganic compounds. Journal of the Chemical Society, Faraday Transactions, I and I I These are, respectively, physical chemistry and chemical physics journals. P A R T I (physical chemistry) includes papers on such topics as radiation chemistry, gas- p hase kinetics, elect rochem istry (0 t her than preparative), surface and interfacial chemistry, heterogeneous catalysis, physical properties of polymers and their solutions and kinetics of polymerization, etc.P A R T I I (chemical physics) contains theoretical papers, especially those on valence and quantum theory, statistical mechanics, intermolecular forces, relaxation phenom- ena, spectroscopic studies (including i.r., e.s.r., n.m.r., and kinetic spectroscopy, etc.) leading to assignments of quantum states, and fundamental theory, and also studies of impurities in solid systems, etc. These are, respectively, the organic chemistry and the physical organic chemistry sections of the Journal. P A R T I (organic and bio-organic chemistry) is designed to contain papers on all aspects of synthetic, and natural product organic and bio-organic chemistry and to deal with aliphatic, alicyclic, aromatic, carboncyclic and heterocyclic compounds.Papers on organometallic topics are considered for either the Dalton or the Perkin Transactions. Journal of The Chemicul Society, Chemical Communications Journal of The Chemical Society, Dalton Transactions Journal of The Chemical Society, Perkin Transactions, I and II 1PART I I (physical organic chemistry) is for papers on reaction kinetics and mechanistic studies of organic systems and the use of physico-chemical, spectroscopic, and crystallographic techniques in the solution of organic problems.Notice to Authors ( 1 ) Although authors need not be members of the Royal Society of Chemistry it is hoped that they will be. (2) Authors must indicate the Part of the Journal they wish their paper to appear in. This preference will be respected unless it is obviously erroneous in terms of the scientific content of the paper. (3) Since all papers will be subjected to refereeing, in parallel, by two independent referees, the original typescript (quarto or A4 size) and two good-quality copies should be provided. (4) All papers should be sent to the Director of Publications, The Royal Society of Chemistry, Burlington House, Piccadilly, London W 1V OBN. ( 5 ) For details of manuscript preparation, preferred usages, etc. the Instructions to Authors, previously available from the Faraday Society, and now obtainable from The Royal Society of Chemistry, should be consulted. (6) The Society will adopt the following abbreviations for the new journals in all its publications.J . Chem. SOC., Chem. Commun. J . Chem. SOC., Dalton Trans. J . Chem. SOC., Faraday Trans. I J . Chem. SOC., Faraday Trans. 2 J . Chem. Soc., Perkin Trans. I J . Chem. Soc., Perkin Trans. 2 * The author to whom correspondence should be addressed is indicated by an asterisk after his name in the heading of each paper. 11FARADAY DIVISION O F THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE ITALIANA D I CHIMICA FlSlCA S O C l i T i DE CHlMlE PHYSIQUE DEUTSCHE BUNSEN G E S E L L S C H A F T F U R P H Y S I K A L I S C H E C H E M I E FARADAY DISCUSSION NO. 7 4 Electron and Proton Transfer University of Southampton, 14-1 6 September 1982 This meeting will be concerned with fundamental aspects of the chemical kinetics of electron and proton transfer reactions in solution and with particular reference to well defined biological systems.Attention will be focused on (i) the theory of charge transfer, (ii) critical experiments designed to test those theories and (iii) their application to the understanding of charge transfer reactions in molecules of biological interest. The meeting will encompass well characterised reactions in solution, redox reactions and elementary biochemical reactions; particular attention will be paid to isotope effects, to electron and proton tunnelling, to intermolecular and intramolecular transfers and to related questions concerning the organisation of biological systems.Among those who have agreed to take part are R. A. Marcus, R. R. Dogonadze, H. Gerischer, J. Jortner, R. M. Kuznetsov, N. Sutin, R. J. P. Williams, H. L. Friedman, J. M. Saveant, J. F. Holzwarth, F. Willig, J. C. Mialocq, M. Kosower, L. I. Krishtalik, E. F. Caldin, H. H. Limbach, W. J. Albery, M. M. Kreevoy, J. J. Hopfield, P. Rich, H. A. 0. Hill, K. Heremans, C. Gavach and D. B. Kelt. The final programme and application form may be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry Burlington House, London W1V OBN FARADAY DIVISION O F THE ROYAL SOCIETY O F CHEMISTRY SYMPOSIUM NO. 1 7 The Hydrophobic Interaction University of Reading, 15-16 December 1982 This term refers to interactions between chemically inert residues arising from perturbations in the unique spatial and orientational correlations in liquid water.These effects provide a major contribution to many of the non-covalently bonded structures that form the basis of life processes. Current advances in the statistical mechanics of polar fluids, intermolecular forces, computer simulation, and membrane physics are providing a new basis for the re-examination of various aspects of hydrophobic effects, their origin and their quantitative description. Such theoretical treatments will be confronted with recent experimental work on simple model systems which, it is hoped, will lead to a better understanding of hydrophobic interactions in more complex processes. The following have provisionally agreed to contribute to the symposium : A.Ben-Naim, H. J. C. Berendsen, D. L. Beveridge, S. D. Christian, L. Cordone, D. Eagland, D. Eisenberg, R. Lumry, P. J. Rossky, M. C. R. Symons, H. Weingartner, M. D. Zeidler The preliminary programme may be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry Burlington House, London W1 V OBN ... 111THE FARADAY DIVISION OF THE ROYAL SOCIETY O F CHEMISTRY GENERAL DISCUSSION NO. 75 Intramolecular Kinetics University of Warwick, 18-20 April 1983 Organising Committee Professor J. P. Simons (Chairman) Dr M. S. Child Professor R . J. Donovan Dr G. Hancock Dr D. M. Hirst Professor K. R. Jennings Dr R. Walsh Experimental and theoretical interest in the time-dependent behaviour of isolated molecules, radicals or ions is strong and increasing.The Discussion will be concerned with the kinetics of processes which occur in isolated species following their preparation in states with non-equilibrium energy distributions (e.g. by photon absorption or collisional activation), Topics covered will include: (a) theoretical and experimental studies of energy redistribution in isolated species; (6) observation and theoretical modelling of the competition between intramolecular energy redistribution and radiative decay or radiationless processes (e.g. internal conversion, fragmentation, isomerisation). The preliminary programme 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 GENERAL DISCUSSION NO.76 Concentrated Colloidal Dispersions Loughborough University o f Technology, 14-1 6 September 1983 The meeting will discuss the experimental investigation and the theoretical description of the properties of concentrated colloidal dispersions, i.e. those systems in which the particle-particle interactions are strong enough to cause significant deviations from ideal behaviour. Both the structural and dynamic features of concentrated systems as determined by scattering, rheological and other techniques will be considered. It is anticipated that a range of dispersion types will be discussed. These will include both 'model' systems and dispersions of importance to industry provided that the data from the measurements can be interpreted. Contributions for consideration by the organising committee are invited and abstracts of about 300 words should be sent by 31 st August 1982 to: Professor R.H. Ottewill, School of Chemistry, University of Bristol, Cantock's Close, Bristol BS8 1TS IVTHE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM NO. 18 Molecular and Microstructural Basis of Viscoelasticity and Related Phenomena Robinson College, Cambridge, 8-9 December 1983 Organising Committee Sir Geoffrey Allen (Chairman) Professor Sir Sam Edwards Dr M. La1 Dr R. A. Pethrick Dr P. Richmond Dr D. A. Young (Editor) The past few years have witnessed the development of new concepts which provide a deeper understanding of the relationship between molecular dynamic and microstructural features of systems and their viscoelastic behaviour.This Symposium is designed to bring together original contributions involving theoretical, computational and experimental studies which represent significant advances in this important field of current activity. It is hoped that such contributions, together with the discussion that they will generate, will lead to new insights into the molecular mechanisms underlying the viscoelastic/rheological behaviour of, for example, flexible and rigid rod-like polymer molecules, liquid crystals and composites. In addition to three oral sessions (at which the main papers will be presented and discussed), the Symposium may include a poster session. Such poster papers will not be published in the Symposium volume. Contributions for consideration by the organising committee are invited.Abstracts of ca. 300 words should be sent to: Dr M. Lal, Unilever Research, Port Sunlight Laboratory, Bebington, Wirral L63 3JW not later than 29 October 1982. Full papers for publication in the Symposium volume will be required by 19 August 1983. FARADAY DIVISION INFORMAL AND GROUP MEETINGS Gas Kinetics Group Seventh International Symposium on Gas Kinetics To be held at the University of Gottingen, West Germany on 23-27 August 1982 Further information from Dr R. Walsh, Department of Chemistry, University of Reading, Whiteknights, Reading RG6 2AD Colloid and Interface Science Group with the Colloid and Surface Chemistry Group of the SCI Adsorption from Solution To be held at the University of Bristol on 8- 10 September 1982 Further information from Dr W D Cooper, Department of Chemistry University of Edinburgh, West Mains Road, Edinburgh EH9 3JJ Industrial Ph ysicaf Chemistry Group Supercritical Fluids: Their Chemistry and Application To be held at Girton College, Cambridge on 13-1 5 September 1982 Further information from Dr W.R. Ladner, National Coal Board, Coal Research Establishment, Stoke Orchard, Cheltenham GL52 4RZ Neutron Scattering Group and Polymer Physics Group with the Institute of Physics The Neutron and its Applications To be held in Cambridge on 13-1 7 September 1982 Further information from The Meetings Officer, Institute of Physics, 47 Belgrave Square, London SW1 X 8QX VFARADAY DIVISION INFORMAL AND GROUP MEETINGS Theoretical Chemistry Group Molecular Electron Structure Theory and Potential Energy Surface To be held at the University of Bristol on 15-16 September 1982 Further information from Dr G.G. Balint-Kurti, School of Chemistry, University of Bristol, Cantock's Close, Bristol BS8 1TS Molecular Beams Group Molecular Beams and Molecular Structure To be held at the University of Bristol on 16-1 7 September 1982 Further information from Dr J. C. Whitehead, Department of Chemistry, University of Manchester, Manchester M13 9PL Surface Reactivity and Catalysis Group The Characterisation of Surface Layers in Chemisorption and Catalysis To be held at the University of East Anglia on 20-21 September 1982 Further information from: Dr M. A. Chesters, School of Chemical Sciences, University of East Anglia, Norwich NR4 7TJ Division Autumn Meeting: Energy and chemistry To be held at Heriot-Watt University, Edinburgh on 21 -23 September 1982 Further information from Dr J.F. Gibson, The Royal Society of Chemistry, Burlington House, London W1 V OBN Statistical Mechanics and Thermodynamics Group with the British Society of Rheolog y Microstructure and Rheology To be held at Trinity Hall, Cambridge on 21-24 September 1982 Further information from Dr P. Richmond, Unilever Research, Port Sunlight, Wirral, Merseyside L62 3JW High Resolution Spectroscopy Group High Resolution Fourier Transform, Laser Infrared and Electronic Spectroscopy To be held at the University of Newcastle-upon-Tyne on 22-24 September 1982 Further information from Dr P. J. Sarre, Department of Chemistry, University of Nottingham, Nottingham NG7 2RD Polymer Physics Group Polymer Electronics To be held in London on 20 October 1982 Further information from the Meetings Officer, The Institute of Physics, 47 Belgrave Square, London SWlX 8oX Electrochemistry Group Spectroscopic Studies of Electrode Surfaces To be held in Oxford on 13-1 4 December 1982 Further information from Professor W.J. Albery, Department of Chemistry, Imperial College, London SW7 2AZ Colloid and Interface Science Group Physical and Biological Aspects of Insoluble Monolayers and Multilayers To be held at the Scientific Societies Lecture Theatre, London on 14 December 1982 Further information from: Dr R. Aveyard, Department of Chemistry, The University, Hull HU6 7RX Division with Polymer Ph ysics Group and Macrogroup UK Annual Chemical Congress: Copolymers To be held at the University of Lancaster on 11-1 3 April 1983 Further information from Dr J.F. Gibson, The Royal Society of Chemistry, Burlington House, London W1 V OBN Polymer Physics Group, Macrogroup UK and the Plastics and Rubber Institute Polyethylenes To be held in London on 8-10 June 1983 Further information from The Plastics and Rubber Institute, 11 Hobart Place, London SW1 W OHZ viPublications from The Royal Society of Chemistry SPECIALIST PERIODICAL REPORTS Catalysis VOlm 4 Senior Reporters: C. Kemball and D. A. Dowden This volume reviews the recent literature published up to mid 1980. Brief Contents: The Design and Preparation of Supported Catalysts: Aspects of Characterization and Activity of Supported Metal and Bimetallic Catalysts; Metal Clusters and Cluster Catalysis; Olefin Metathesis; Superbasic Heterogeneous Catalysts; Hydration and Dehydration by Heterogeneous Catalysts; Sulphide Catalysts: Characterization and Reactions including Hydrodesulphurization; Carbon as a Catalyst and Reactions of Carbon.Hardcover 266pp 0 85186 554 2. Price f29.00 ($62.00). RSC Members f17.50 Gas Kinetics and Energy Transfer Vol. 4 Senior Reporters: P. G. Ashmore and R. J. Donovan A review of the literature published up to early 1980. Brief Contents: Reactions Studied by Molecular Beam Techniques; Reorientation by Elastic and Rotationally Inelastic Transitions; Infrared Multiple Photon Excitation and Dissociation: Reaction Kinetics and Radical Formation; Ultraviolet Multiphoton Excitation: Formation and Kinetic Studies of Electronically Excited Atoms and Free Radicals; Gas Phase Reactions of Hydroxyl Radicals; Gas Phase Chemistry of the Minor Constituents of the Troposphere.Hardcover 252pp 0 85186 786 3. Price f45.00 ($96.00). RSC Members a5.00 Mass Spectrometry VOlm 6 Senior Reporters: R. A. W. Johnstone This volume reviews the literature published between July 1978 and June 1980. Brief Contents: Theory and Energetics of Mass Spectrometry; Structures and Reactions of Gas-phase Organic Ions; Gas-phase Ion Mobilities, lon- Molecule Reactions, and Interaction Potentials; Interaction of Electromagnetic Radiation with Gas-phase Ions; Aspects of Secondary Ion Emission; Development and Trends in Instrumentation in Mass Spectrometry; Applications of Computers and Microprocessors in Mass Spectrometry; Gas Chromatography- Mass Spectrometry and High- performance Liquid Chromatography- Mass Spectrometry; Reactions of Negative Ions in the Gas Phase; Natural Products; The Use of Mass Spectrometry in Pharmacokinetic and Drug Metabolism Studies; Organometallic, Co-ordination, and Inorganic compounds Investigated by Mass Spectrometry.Hardcover 368pp 0 85186 308 6. Price f39.50 ($88.00). RSC Members €23.00 ORDERING RSC Members should send their orders to: The Royal Society of Chemistry, The Membership Officer, 30 Russell Square, London WCl B 5DT. Non-RSC Members should send their orders to: The Royal Society of Chemistry, Distribution Centre, Blackhorse Road, Letchworth, Hens SG6 1 HN.The Royal Society of Chemistry Burlington House Piccad i I I y London W1V OBN viiNOTES I t has always been the policy of the Faraday Transactions that brevity should not be a factor influencing acceptability for publication. In addition however to full papers both sections carry at the end of each issue a section headed “Notes”, which are short self-contained accounts of experimental observations, results, or theory that will not require enlargement- into “full” papers. The “Notes” section is not used for preliminary communications. The layout of a “Note” is the same as that of a paper. Short summaries are required. The procedure for submission, administration, refereeing, editing and publication of “Notes” is the same as for “full” papers. However, “Notes” are published more quickly than papers since their brevity facilitates processing at all stages.The Editors endeavour to meet authors’ wishes as to whether 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 I500 words or word-equivalents. NOMENCLATURE AND SYMBOLISM 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 rules themselves and guidance on their use are given.Physicochemical Quantities and Units. Manual of Symbols and Terminology for Physicochemical Quantities and Units. (Pure and Appl. Chem., Vol. 51, No. 1, 1979, pp. 1-41. Also available as a soft-cover booklet from Pergamon Press, Oxford.) Surface Chemistry. ‘ Definitions, Terminology, and Symbols in Colloid and Surface Chemistry - I . ’ (Pure and Appl. Chem., Vol. 31, No. 4, 1972, pp. 577-638.) ‘ Definitions, Terminology, and Symbols in Colloid and Surface Chemistry - 11. Heterogenous Catalysis. ’ (Pure und Appl. Chem., Vol. 46, No. I , 1976, In addition, the terminology and symbols for the following subject areas are available either in the form of soft-cover booklets from Pergamon Press (denoted by *) or have been the subject of articles in Pure and Applied Chemisrry in recent years: activities;* chromatography; electrochemistry; electron spectroscopy; equilibria, fluid flow; ion exchange; liquid-liquid distribution ; molecular force constants; Mossbauer spectra; nuclear chemistry; pH; polymers; quantum chemistry; radiation;* Raman spectra; reference materials (recommended reference materials for the realization of physico- chemical properties: general introduction, enthalpy, optical rotation, surface tension, optical refraction, molecular weight, absorbance and wavelength, pressure-volume- temperature relationships, reflectance, potentiometric ion activities, testing distillation columns); solution chemistry; spectrochemical analysis; surface chemistry; thermo- dynamics, and zeolites. Finally, the rules for the naming of organic and inorganic compounds are dealt with in the following publications from Pergamon Press: ‘Nomenclature of Organic Chemistry, Sections A, B, C , D, E, F, and H’, 1979. ‘Nomenclature of Inorganic Chemistry’, 1971. A complete listing of all IUPAC nomenclature publications appears in the 198 1 Index issues of J. Chem. SOC. pp. 71-90.) viii

ISSN:0300-9599    

DOI:10.1039/F198278FP049

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraduy Trans. 1, 1982, 78, 2033-2040 Fluorescence and Absorption Spectral Study of the Interaction between Cetylpyridinium and 2-Naphtholate Ions in Aqueous Solution BY 'SOJI A. AMIRE AND HUGH D. BURROWS* Chemistry Department, University of Ife, Ile-Ife, Nigeria Received 30th January, 198 1 Absorption spectral studies on aqueous solutions of cetylpyridinium chloride and sodium 2-naphtholate indicate the formation of cetylpyridinium naphtholate complexes. The complex is shown to be non- fluorescent, and the main cause for the decrease in 2-naphtholate fluorescence intensities with increasing cetylpyridinium concentrations below the critical micelle concentration of the surfactant involves complex formation. At higher surfactant concentrations, quenching of fluorescence of 2-naphtholate ions is also suggested to occur via a dynamic interaction between these ions and cetylpyridinium micelles.Photophysical studies of micelle-forming surfactant solutions have been shown to yield valuable information on both the structure and dynamics of such systems.l Fluorescence quenching has been shown to be particularly valuable in this respect.l? Pyridinium ions are known to be efficient quenchers of the fluorescence of aromatic hydrocarbon^,^ and it is of interest to see whether, by using a suitable long-chain alkylpyridinium compound, information can be obtained on the interaction between such a system and a fluorescent solute molecule in aqueous solution. In the present study we have looked at the interaction between cetylpyridinium cations and 2-nap ht hola te anions.EXPERIMENTAL 2-Naphthol was purified by boiling with decolourising charcoal and recrystallizing from hot water. The white crystals so obtained were washed with water and dried in an oven. Cetylpyridinium chloride (CPCI, Koch-Light pure grade) was used without further purification. 2-Naphthol was dissolved in water with sufficient 0.1 mol dm-3 sodium hydroxide to ensure complete neutralization. Stock solutions of CPCl (0.1 mol dm-3) were prepared in water. To avoid potential problems with ageing of micellar systems,2 all experiments were carried out with freshly prepared solutions. The surfactant-naphtholate systems used in this study were prepared either by mixing appropriate volumes of the stock solutions in volumetric flasks or by adding small volumes of the CPCl solution directly with a microsyringe to 2-naphtholate solutions in cuvettes and correcting for the slight volume change this introduced.Identical results were obtained with the two procedures. The pH of the solution was adjusted by adding dilute potassium hydroxide or hydrochloric acid solutions. Buffers were not used because of problems of precipitation and the potential of common buffers to complex with the cetylpyridinium system. Fluorescence spectra were recorded using Aminco-Bowman and EEL 244 fluorimeters. Absorption spectra were measured on Pye-Unicam SP6-400 and SP6-500 spectrophotometers. For absorption spectra, solutions were thermostatted ( & 0.1 "C) with a water-circulating bath. Apparent equilibrium constants were obtained from absorption spectra using a modified Hill Whilst this is essentially an empirical method, it has been shown to be valuable in systems such as the present one where multiple equilibria are pos~ible.~ The method involves 20332034 CETY LP Y R I D I N I U M-2-NAP HTHOL A TE COMPLEXES plots of log [(A -A,,)/(Am - A ) ] against log [CPCI],,,,, where A , is the absorbance in the absence of CPCl, A is that in the presence of CPC1, and A , is the absorbance of the solution at complete complex formation (saturation).Free CPCl concentrations were calculated from initial concentrations and absorption spectra. It can be simply shown that at 50% saturation the intercept on the x axis is -log K, whilst the slope of the plot, n, gives a measure of the extent of aggregation5 As defined in t h s way, the apparent equilibrium constant, K , takes the units of concentration.The K values were corrected for the parallel reaction between 2-naphtholate and hydrogen ions as de~cribed.~ In all cases studied, the Hill plots were linear over the range 15-85 % complex formation. RESULTS On addition of CPCl to 2-naphthol it was noted that the solution became pale yellow in colour. This colour formation was seen to be reversible upon dilution. The absorption spectra were run of solutions of sodium 2-naphtholate ( lop4 to lop3 mol dm-3) alone, and in the presence of varying concentrations of CPCl. Typical spectra are shown in fig. 1. The lowest energy absorption band of the 1.5 1 .o aJ 5 e 2 I) m 0.5 0 \ \ \ \ \ w avele ngt h/n rn FIG.1 .-Electronic spectra of aqueous solutions of sodium 2-naphtholate ( mol dm-3, pH 1 1.85): solid line, alone; dashed line, in the presence of 5 x mol dm-3 cetylpyridinium chloride. 2-naphtholate ion shifted by ca. 10 nm to longer wavelengths, suggesting complex formation. Addition of a similar concentration of sodium chloride to 2-naphtholate solutions caused no such red shift, indicating that complex formation was between cetylpyridinium and naphtholate ions. Isosbestic points were observed around 325 and 335 nm, suggesting that there was only one equilibrium process involved. Plots of fractional saturation, a, determined at a particular wavelength against free cetylpyridinium concentration were observed to be sigmoidal. A typical plot is shown in fig.2. For a particular 2-naphthol concentration and pH, plots of absorbance at 355 nm against absorbance at 370 or 400 nm at various CPCl concentrations were linear. Apparent equilibrium constants were determined by studying the absorbance at 355nm, the absorption maximum of the complex, as a function of CPCl concentration. A typical Hill plot obtained is shown in fig. 3. The values obtained for apparent equilibrium constants are comparable to those obtained in other studiesS. A. AMIRE A N D H. D . BURROWS 2035 1 .o 0 . 8 h 0.6 a s \ h T I s 0.4 II tt 0.; 0 2 4 6 [ CPCl] free / 1 0-4 mol dm-3 FIG. 2. FIG. 3, FIG. 2.-Plot of extent of complexing against [CPCl]f,ee for solutions of 2-naphthol (1.04 x lop4 mol dmp3, pH 8.93, 27.0 "C) in the presence of cetylpyridinium chloride studied at the absorption maximum of the complex (355 nm).FIG. 3.-Hill plot for the system 2-naphthol (1.0 x rnol dm-3, pH 9.64, L = 355 nm) in the presence of cetylpyridinium chloride at 20.6 OC. TABLE 1 .-APPARENT EQUILIBRIUM CONSTANTS AS A FUNCTION OF pH AND 2-NAPHTHOL CONCEN- TRATION FOR THE SYSTEM 2-NAPHTHOL/CETYLPYRIDINIUM CHLORIDE IN AQUEOUS SOLUTION AT 27.0 *C [2-naphthol]/ mol dmP3 K / lo4 mol dm-3 PH 9.64 8.93 9.45 11.18 9.32 1 .oo 1.04 1.96 1.90 3.02 1.09 3.08 1.99 1.59 2.68 using this m e t h ~ d , ~ and are in a range where the Hill plot can be accurately applied to spectrophotometric data. For these studies, the concentration of CPCl was always less than its critical micelle concentration (c.m.c. of cetylpyridinium chloride 9.0 x mol dm-3).6 Nevertheless, the slopes of the Hill plots were > 1 in all cases, having a value of 1.5kO.1 at pH 3 9.6 (the pK, of 2-naphthol),' but increasing to 2.6k0.3 at lower pH values.This was independent of temperature over the range studied (293-307 K), and suggests that the surfactant molecules are starting to aggregate below the c.m.c. Apparent equilibrium constants were determined from the Hill plots, and were found to increase with decreasing pH; they also appeared to show some dependence on 2-naphthol concentration (table 1). However, this apparent2036 CETY L P Y R I D I N I U M-2-N A P H T HO L A TE COMPLEXES dependence in 2-naphthol may result from pH changes. Further studies are in progress to elucidate this factor. Apparent equilibrium constants were determined at constant pH and 2-naphthol concentration as a function of temperature.A semilogarithmic plot of formation constant against inverse temperature was linear (fig. 4). The following thermodynamic parameters for complex formation were determined* from this: A H 0 = -26.0k4.4 kJ mol-1 A S 0 = - 14.5+ 11.3 J K-l mol-l. The enthalpy falls within the range of values reported for organic charge-transfer complexes between n-electron donors and acceptors.8 KIT FIG. 4.-Semilogarithmic plot of the apparent equilibrium constant for the 2-naphtholate-cetylpyridinium complex against inverse temperature. To gain further insight into the interaction occurring between the two ions, the fluorescence of the 2-naphtholate ion mol dm-3) was studied in the presence of varying concentrations of CPCl.Whilst the fluorescence emission and excitation spectra were found to be essentially independent of CPCl concentration, the intensity of the fluorescence was seen to decrease dramatically upon increasing the surfactant concentration (fig. 5). Addition of similar concentrations of chloride caused no such decrease in fluorescence, suggesting that the behaviour was due to interaction of cetylpyridinium and 2-naphtholate ions. This conforms with the known ability of alkylpyridinium compounds to quench the fluorescence of aromatic compo~nds.~ This decrease in intensity can arise from two causes, either static quenching by formation of a non-fluorescent complex, reaction (3), or dynamic quenching involving interaction between excited 2-naphtholate anion (BN-) and cetylpyridinium cation (CP+) (4).BN- + Av -+ lBN* monomer excitation (1) lBN* -+ BN- + kv monomer fluorescence (2) BN- + CP+ --+ (BN - CP) formation of non-fluorescent complex (3) 'BN* + CP+ -+ BN- + CP+ dynamic quenching (4) The fact that no shift in the emission spectrum was observed on complexation suggestsS. A. AMIRE AND H. D. BURROWS 2037 0 2.5 5.0 7.5 [ CPCll / 1 0-4 mol dm-3 FIG. 5.--Intensity of fluorescence of 2-naphtholate ion ( mol dmP3, ,Iexcitation = 350 nm, Aemission = 420 nm) as a function of CPCl concentration : circles, experimental data; solid line, theoretical curve from eqn (1 2). strongly that the complex was non-fluorescent. If only free naphtholate ion is fluorescing, it is possible to derive an equation relating the fluorescence intensity to its concentration and to the fraction of light absorbed at the excitation wavelength by fluorescer and complex. Using these assumptions, the intensity of light at this wavelength at any position in the cell, I, will be3 ( 5 ) I’ = I,, X 1 O-(‘BN ‘BN+&comp ‘comp) where E and c are the molar absorptivities and concentrations, respectively, of the absorbing species.Cetylpyridinium is not included in this as it does not absorb at the excitation wavelength. The amount of light absorbed by free 2-naphtholate ion in a length dl will be d I = (In 10) I,, X 1 O-(&BN ‘BN+&comp 'camp) EBN cBK dl. (6) This can be integrated to give the intensity of light absorbed over the length, x, from which emission is observed (7) (In 10) I,, EBN cBN( 1 - 1 O-(&BN CBK+EcomP ccomp) ”) (EBN CBN + Ecomp Ccomp) x I = The intensity of fluorescence, If, will be proportional to this K ’ c B N ( 1 - lO-(&BN CBNSEcompCcomp) ”) If = EBN CBN +Ecomp Ccomp where K’ is a constant for a particular fluorescer and fluorimeter geometry.If the total BN concentration, cT, remains constant, then Ccomp = CT-CBN K’cB,( 1 - 10-[CBN(EBN-Ecomp)+Ecomp ‘TI ”) and If = CBN(EBN - Ecomp) + Ecomp CT If we make the further assumption that for the denominator2038 CETYLPYRIDINIUM-2-NAPHTHOLATE COMPLEXES which is valid, except at very low CPCl concentrations, then where A and D are constant for a particular total BN concentration, excitation wavelength and fluorimeter, and B is a constant involving only the difference in molar absorptivities of BN and the complex, and x.By using values of cBN, either from absorption spectra or from apparent equilibrium constants, it is possible to construct a theoretical plot of fluorescence intensity against CPCl concentration. Good agreement was observed between theory and experiment under the conditions employed ([BN] = mol dm-3, pH 9.64, t = 25 OC, Aexcitation = 350 nm, Lemission = 420 nm) using A = 2.1 1 x lo8, B = 1000, and D = 0.018 (fig. 5). Interpret- ation of A and D is complex, as these terms involve intensity of exciting light, cT, etc. However, B involves x and the difference between the molar absorptivities at the excitation wavelength, and, given the fact that both complex and 2-naphtholate absorptions are changing very rapidly with wavelength in this region, the best-fit value is certainly of the correct order of magnitude.The results strongly support the suggestion that the decrease in fluorescence intensity at these CPCl concentrations arises primarily from the formation of a non-fluorescent complex. At concentrations around or above the c.m.c., an additional mechanism of reduction of fluorescence intensity appears to be operating. Whilst increasing the total CPCl concentration from 5 x to 10 x lov4 mol dm-3 only decreases the free 2-naphtholate concentration by a factor of 1.7, the fluorescence intensity over the same range decreases by a factor of five. The effect shows up most clearly when the fluorescence data are replotted in the form of a Stern-Volmer plot (fig.6). At low CPCl concentrations the plot is non-linear, as expected from the formation of a non-fluorescent complex absorbing strongly at the excitation wavelength. Above ca. 5 x mol dm-3 CPCl; however, the plot* is linear, with slope 9.32 x lo4 dm3 mol-l. Quantitative interpretation of the behaviour is difficult, as the data must be corrected for light absorption by the non-fluorescent complex, and for cases where more light is absorbed by the complex than by the fluorescer errors are likely to be large.3 A rough correction for this may be made by using experimental values from absorption spectral data at the excitation wavelength. From these, the slope of the linear region is found to be ca. lo4 dm3 mol-l. The high value for this quenching constant, and the fact that the linear region starts at CPCl concentrations close to the c.m.c., suggests that this probably is due to dynamic quenching of excited 2-naphtholate ions by cetylpyridinium micelles. DISCUSSION The absorption spectral changes observed when CPCl is added to aqueous solutions of 2-naphtholate ions provide clear evidence for complex formation. That this complex has charge-transfer character is suggested by the red-shift of the lowest energy absorption, and by the known tendency of alkylpyridinium ions to form charge-transfer complexe~.~~ lo The ionisation potential of 2-naphtholate ion? is certainly favourable for the formation of such a complex. Studies of the association equilibrium suggest that more than one cetylpyridinium monomer unit is present in each complex.Whether the anion is assisting aggregation or whether there are already submicellar * Because of the wide variation in fluorescence intensities in this study, the concentration range which can be studied at a particular instrument sensitivity is limited. However, studies on different sensitivity settings indicate that the Stern-Volmer plot is linear to higher concentrations than those indicated in fig. 6. t A value of 7.83 eV is reported for 2-naphtho1," quoted in ref. (12). The ionisation potential of the anion should be even lower.S . A. AMIRE AND H. D. BURROWS 2039 80 60 * 1 +o 40 20 01 0 0 o o I I 5 10 1 0-4 [ CPCl] /mol dm-3 FIG. 6.-Stern-Volmer plot for the quenching of fluorescence of 2-naphtholate ion ( mol dm-3) by CPCl at room temperature.aggregates of cetylpyridinium ions at CPCl concentrations corresponding to the onset of complex formation is not clear. Submicellar aggregates of surfactants have been reported with carboxylic acid soaps,13 but the existence of such aggregation remains contro~ersial.~~ The increase in the slope of the Hill plots, n,. when the pH is below the pK, of 2-naphthol suggests that there is an increase in the extent of aggregation in this case. This is supported by an increase in the apparent equilibrium constant with decreasing pH. The nature of the multiple binding in these complexes is not clear. However, both the sigmoidal shape of the plot of fractional saturation against concentration (fig. 2), and the fact that n > 1 strongly suggest15 that the binding sites in the aggregates are not independent, and that binding of one ligand increases the affinity of the aggregate for binding a second ligand.The high values observed for the apparent equilibrium constants, which are comparable to those found for the binding of various ligands to haem~proteins,~ may indicate that the naphtholate ions are binding in a region where the dielectric constant is lower than that of the bulk water. Studies of fluorescence intensity as a function of surfactant concentration indicate that the complex is non-fluorescent. In 1 -(9-anthrylmethyl)pyridinium chloride, the pyridinium ring is seen to quench the anthracene fluorescence efficiently by an intramolecular me~hanism.~ For such efficient quenching, however, the pyridinium and aromatic species must be close to each other.lH n.m.r. studies of aqueous solutions of various phenols in the presence of surfactant systems strongly suggest that they are solubilised in the palisade layer of the micelle, with the hydroxy group pointing towards the bulk aqueous so1ution.l69 l7 Specific charge-transfer interactions in the present case favour the 2-naphtholate and cetylpyridinium ions being adjacent, even in the absence of full micellar aggregation.2040 C E T Y L P Y R I D I N I U M-2-N A P H T HO L A TE C 0 M P L E X E S At higher CPCl concentrations the Stern-Volmer plot of the quenching of 2-naphtholate fluorescence is linear. This appears to be due to dynamic quenching of 2-naphtholate excited state by cetylpyridinium micelles. Using the corrected Stern-Volmer quenching constant, and the lifetime of the 2-naphtholate fluorescence (8.1 ns),18 an apparent second-order rate constant of ca.10l2 dm3 mol-l s-l was obtained for this quenching. A ‘true’ rate constant ( i e . in terms of concentration of micelles) can be estimated by dividing this apparent rate constant by the micellar aggregation number ( n = 95 for cetylpyridinium chloride).lg This gives a value ca. 1O1O dm3 (mole of micelle)-l s-l. This is comparable to the diffusion-controlled rate in aqueous solution, and can be contrasted with the rate constant for reaction between hydrated electrons and cetylpyridinium micelles [ca. 5 x 10l2 dm3 (mole of micelle)-l S - ~ ] . ~ O The extremely high rate constant in this latter case is suggested to arise from the high electrical potential of the micelle double layer.In the present case, the surface charge may be partially neutralised by naphtholate ions already bound to the micelle. K. Kalyanasundaram, Chem. SOC. Rev., 1978, 7, 453. * See, for example, H. D. Burrows, S. J. Formosinho, M. F. J. R. Paiva and E. J. Rasburn, J. Chem. SOC., Faraday Trans. 2, 1980, 76, 685, and references therein. R. A. Hanna, D. R. Rosseinsky and T. P. White, J. Chem. SOC., Faraday Trans. 2, 1974, 70, 1522. A. V. Hill, J. Physiol. (London), 1912, 40, 4. A. C. Anusiem, J. G. Beetlestone and D. H. Irvine, J. Chem. SOC. A , 1968, 960. E. J. Fendler and J. H. Fendler, Advances in Physical Organic Chemistry, ed. V. Gold (Academic Press, London, 1970), vol. 8, p. 276. A. Albert and E. P. Serjeant, The Determination of Ionization Constants (Chapman and Hall, London, 2nd edn, 1971), p. 87. R. S. Mulliken and W. B. Person, Molecular Complexes (Wiley, New York, 1969), p. 89. E. M. Kosower, Progress in Physical Organic Chemistry, ed. A. Streitwieser and R. W. Taft (Intersci- ence, New York, 1965), vol. 3, p. 81. lo R. Foster, Organic Charge-transfer Complexes (Academic, London, 1969). l 1 K. Beakers and A. Szent-Gyorgyi, Red. Trav. Chim. Pays-Bas, 1962, 81, 255. l2 J. B. Birks, Photophysics of Aromatic Molecules (Wiley-Interscience, London, 1970), p. 459. l3 D. Eagland and F. Franks, Trans. Faraday SOC., 1965, 61, 2468. l4 L. R. Fisher and D. G. Oakenfull, Chem. SOC. Rev., 1977, 6, 25, and references therein. l5 E. Antonini, Science, 1967, 158, 1417. l6 J. J. Jacobs, R. A. Anderson and T. R. Watson, J. Pharm. Pharmacol., 1971, 23, 148. li F. Tokiwa and K. Aigami, Kolloid Z. 2. Polym., 1971, 246, 688. A. Weller, 2. Phys. Chem. (Frankfurt am Main), 1958, 17, 224. l9 E. W. Anacker, J , Phys. Chem., 1958, 62, 41. 2o M. Gratzel, J. K. Thomas and L. K. Patterson, Chem. Phys. Lett., 1974, 29, 393. (PAPER 1/144)

ISSN:0300-9599    

DOI:10.1039/F19827802033

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 2041-2049 Thermodynamic Study of the Influence of Complexation on Exchange Equilibria in Wyoming Bentonite Clay BY ANDRE MAES,* EDDY RASQUIN AND ADRIEN CREMERS Centrum voor Oppervlaktescheikunde en Collo'idale Scheikunde, Katholieke Universiteit Leuven, De Croylaan 42, B-3030 Leuven (Heverlee), Belgium Received 9th February, 198 I In Wyoming Bentonite the ion-exchange reactions for the homovalent pairs Ag(en);-Cs+ and Cu(en)i+-Ca2+ and for the heterovalent pair Ag(en)$-Cu(en)z+ are characterized by A G e values of exchange which are very much larger than those for the aqueous metal ions. These effects, which emphasize the importance of ion-ligand uersus ion-solvent (water) interactions, are discussed in terms of the relative contributions of the interactions occurring in the surface and bulk solution phase. A significant decrease is found in bivalent ion selectivity with increasing loading in the Ag(en)l-Cu(en)i+ case (two orders of magnitude difference in K , at both ends of the composition scale).This is opposite to what is commonly found in mono-divalent ion-exchange equilibria in montmorillonite. This effect is discussed in terms of heterogeneity in charge distribution in the exchanger and variations in electrostatic interaction. The exchange selectivities of the aqueous transition-metal ions Zn2+, Ni2+, Cu2+, Co2+ and Cd2+ versus Na+ in montmorillonite are quite similar,l the overall AG* value varying around - 0.75 kJ equiv.-l. In contrast, ethylenediamine complexes of Zn2+, Ni2+, Cu2+ and Ca2+ are very selectively exchanged,' the overall AG* values for the Ca2+ displacement ranging from - 7 (Zn2+) to - 10.25 (Ca2+) kJ equiv.-l.The present paper reports on the exchange of two complexes, Ag(en)l (versus Cs+) and Cu(en)i+ (versus Ca2+) and their aqueous counterparts in Wyoming Bentonite (W.B.) clay. The purpose is twofold: (1) to gather more information on the role of the ion-solvent interaction on the magnitude of the exchange selectivity by modi- fying the ion environment using a complexation with neutral ligands; (2) to verify whether the characteristic selectivity rise with bivalent ion loading which is observed in heterovalent exchange reactions among hydrated cations3* is also observed in heterovalent exchanges among complexes. EXPERIMENTAL The fraction of Wyoming Bentoniteclay smaller than 0.5 pm was separated usingconventional techniques, and was stored as a 3% Na+ suspension in 1 mol dm-3 NaNO, at 5 OC in the dark.Appropriate quantities of W.B. clay were dialysed to 0.01 mol dmP3 Na+ prior to use. The Cu(ethylenediamine)i+-Ca2+ selectivity was measured by overnight equilibrium of 10 cm3 of a 1 % Ca2+ suspension of W.B. with 20 cm3 of mixed Cu(en)i+-Ca2+ solution of 0.01 total normality in Cu2+ + Ca2+. The formation of the fully coordinated Cu(en)i+ complex was ensured by adding an excess of 2 x mol dm-3 ethylenediamine (en). The pH of the starting solutions was brought to 8.5 with HNO,. The Ag(en),f-Cs+ equilibrium in W.B. was measured by mixing 5 cm3 Na+ suspension of the clay (1 %) at 0.01 mol dm-3 NaNO, with 20 cm3 mixed solutions of Ag(en),f and Cs+ at 0.01 total normality.From the stability constants of the Ag(en)z complex (Kl = lo4.', K2 = 103.0)5 we calculated that a free ligand concentration of 0.1 mol dm-, (en) at pH 11 was necessary to form the biscomplex. The solutions were therefore 0.125 mol dmP3 in (en) and at pH ca. 11.5. 204 12042 EXCHANGE EQUILIBRIA IN WYOMING BENTONITE The selectivity between the Ag(en)l and Cu(en)i+ complexes was measured at 25 and 5 OC by equilibrating 10 cm3 1 % Na+ suspension of W.B. with 40 cm3 of mixed Cu(en);+-Ag(en)$ solutions of 0.0 1 total normality in Cu2+ + Ag+. The (en) concentration was 0.125 mol dmP3 and the pH was ca. 11.5. The binary Na+-Ag+ and Na+-Cs+ exchange reactions were also studied at 0.01 total normality using an Na+ suspension of W.B.at 0.01 mol dm-, NaNO, and mixed Na+-Ag+ and Na+-Cs+ solutions, respectively. All equilibrations were performed in a thermostatted environment at 25 OC by overnight end-over-end shaking, followed by phase separation using a thermostatted supercentrifuge. Concentrations were determined by radiotracer methods using 22Na, lloAgm and 13'Cs. Cu2+ was analysed by flameless and conventional atomic absorption techniques (varian Techtron AAG). d(OO1) spacings of wet, fully exchanged Cu(en)i+ at 5 x loP4 and 0.1 mol dm-, free (en) and Ag(en)l at 0.1 mol dmP3 free (en) were measured using the Debye-Scherrer technique. RESULTS The natural logarithm of the selectivity coefficients for the Cs+-Ag(en)$, Ca2+-Cu(en)i+ and Cu(en)i+-Ag(en)$ equilibria in W.B.are shown in fig. 1, 2 and 3 as a function of the equivalent fraction of Ag(en)i, Cu(en)i+ and Ag(en)z, respectively. ZAg FIG. 1 .-Dependence of In K,[Cs+-Ag(en)z] on silver loading. In K,(Cs+-Ag+), shown by the broken line, is calculated (see text). The sum of Ca2+ +Cu(en)g+ adsorbed ranges from 0.84 to 1 .O mequiv. g-l at the Ca2+- and Cu(en)i+-rich ends of the isotherm. The equilibrium pH is 7. In the case of the Cs+-Ag(en)l equilibrium the sum total ranges from 0.96 mequiv . g-l (Cs+-rich end) to 1.08 mequiv. g-' [Ag(en)l-rich end], and the equilibrium pH is 1 1.5. The use of a well-dispersed Na+ clay as starting material in the Cs+-Ag(en)$ and Cu(en)i+-Ag(en)l equilibria is preferred over using the flocculated Ag(en)i, Cs+ or Cu(en);+ clays because the former allows easier handling and better reproducibility in pipetting.The levels of Na+ displaced in the liquid phase are sufficiently lowA. MAES, E. RASQUIN A N D A. CREMERS FIG. 2.-Dependence of In Kc[Ca2+-Cu(en),2+] on copper loading. In Kc(Ca2+-Cu2+), shown by the broken line. is taken from the literature.s 2043 I ' I I I I I 1 I I 0.2 0.4 0.6 0.8 ZA, ZAg FIG. 3.-Dependence of In K,[Cu(en)i+-Ag(en);] on silver loading at 5OC (A) and 25OC (0). In K,(Cu2+-Ag+), shown by the broken line, is calculated (see text). Corrections for solution-phase activity ,.-dX,.:-..tA n,-o :.,,-.l-.A-A bUGlIIblG1lLJ a l G 1 1 1 ~ 1 u u G u . (ca. 4 x I 0-3 mol to justify an analysis in terms of a purely binary system. Indeed, using K , (Na+-Cs+) = 10.25 (see later) one estimates the maximum Na+ occupancy to be ca.5% of the Cs+ content. The amount of Na+ adsorbed therefore equals 0.05 mequiv. g-l at high Cs+ contents and vanishes towards the Ag(en)i-rich end. In the case of the Cu(en)i+-Ag(en)i exchange, the selectivities of both cations versus Na+ are high enough to displace Na+ from the exchange complex entirely.2044 EXCHANGE EQUILIBRIA I N WYOMING BENTONITE Z A g FIG. 4.-Dependence of In K,:(Na+-Ag+) on silver loading. I I I I I 1 3 t 4 ZCS FIG. 5.-Dependence of In K,(Na+-Cs+) on caesium loading. Kielland plots for the Na+-Ag+ and Na+-Cs+ exchanges are given in fig. 4 and 5. The total ion occupancy, i.e. Ag+ + Na+, remains constant at 0.85 f 0.02 mequiv. g-l. Beyond an equivalent fraction of Ag+ of 0.7-0.8 an oxidoreduction phenomenon occurs, as evidenced by the blackening of the sample.The selectivities in this region are therefore extrapolations. The sum of Na+ + Cs+ increases gradually from the Na+ CEC of 0.82 mequiv. g-l to 0.98 mequiv. g-l at the Cs+-rich end. All selectivity data were obtained at low total normality (0.01), which allows us to identify the activity ratio with the concentration ratio in solution in the cases of homovalent exchange, i.e. the Cs+-Ag(en)l, Ca2+-Cu(en)i+, Na+-Ag+ and Na+-Cs+ exchanges. In the case of the heterovalent Ag(en)i-Cu(en)i+ exchange the activity coefficient ratio equals y"g(en)l - rf Ag(en), NO, rCu(en>f+ 7% Cu(en),(NO,), which may differ from unity. The Davies equation6 fits the activity coefficients of all electrolytes up to I = 0.2 and was used to calculate the solution-phase activity coefficient ratio.The ionic strength was obtained by accounting for Na+, Cu(en):+, Ag(en)i and the monoprotonated form of (en). The correction, expressed in terms of In K, units, then varies from 0.266 to 0.249 over the experimental composition range and is included in the data in fig. 3. The selectivities for the exchange reactions of aqueous ions were determined or taken from literature data and are represented by the broken lines in fig. 1-3. In K(Cs+-Ag+) = -2.05 and is calculated from In K(Na+-Cs+) = 2.45 andA. MAES, E. RASQUIN A N D A. CREMERS 2045 In K(Na+-Ag+) = 0.4, which were obtained by graphical integration of the experi- mental data in fig. 4 and 5; In K(Ca2+-Cu2+) = - 0.044 is as given by El Sayed et af.;8 In K(Ag+-Cu2+) = - 0.784 is obtained from the former Ca2+-Cu2+ and Ag+-Na+ data and In K(Na+-Ca2+) = 0.06;9 In K,(Ag+-Cu2+) is then -0.784+ 1 = 0.216. All thermodynamic data, obtained by using the Gaines and Thomas integrati~n,~ are given in fig. 6 in terms of AGe of exchange and are combined with literature data. I \ \ A I E! d h 0 d I Ag+ *-0.986 N a+ \ c s+ FIG. 6.-Combination of thermodynamic data given in kJ equiv.-' of the indicated binary equilibria in Wyoming Bentonite clay. Full and broken lines represent experimentally obtained and calculated exchanges, respectively. Literature data are indicated where used. TABLE l.-AG* VALUES (kJ mol-l) OBTAINED AT 25 OC FOR THE INDICATED HYPOTHETICAL The overall stability constants corresponding to the formation of Cu(en)i+ and Ag(en)l in W.B.v2) and in the solution phase ( p2) are also compared. EXCHANGES Cu2+-Cu(en)i+ - 14.41 22.57 20.03 Ag+-Ag(en); - 16.40 10.59 7.7 AGe(Na+-Cs+) = -6.04 kJ equiv.-l differs markedly from the value obtained by Gast et a(.', in another bentonite, but fits perfectly into the charge-density relationship for the Na+-Cs+ equilibrium in octahedrally substituted montmorillonites.ll Fig. 1 and 2 clearly demonstrate that complexing of Ag and Cu with (en) results in high selectivities for both cations. In K[Ca2+-Cu(en)i+] = 5.80, whereas In K(Ca2+-Cu2+) is only -0.044. These data fit into the charge-density relationship found in other montmori1lonites.l2 Complexing of Ag+ enhances its selectivity uersus Cs+ from In K(Cs+-Ag+) = - 2.05 to In K[Cs+-Ag(en)z] = 4.60.Complexation has a similarly dramatic effect on exchanges among complexes, as seen in fig. 3 ; the In K value for the Ag(en)i-Cu(en)i+ is -7.84 and 0.47 for the aqueous ions. The free-energy changes for the hypothetical exchange reactions of the uncomplexed by the complexed ions are easily calculated from these data and are given in table 1. They correspond to the extra stabilities gained by the complexes upon adsorption in the interface.13 These free-energy changes were shown13-14 to be related to the overall complex formation constants in the surface g2) and solution phase v2) by Pz2046 EXCHANGE EQUILIBRIA I N WYOMING BENTONITE in which M and MLn stand for the aqueous metal ion and its coordinatively saturated complex.The data in table 1 show an increase in the logarithm of the overall stability constant of Ag(en); and Cu(en);+ complexes in W.B. by, respectively, 2.89 and 2.54 units with respect to their value in bulk solution. The foregoing data are useful in assessing the reversibility of the Cu(en)i+-Ag(en)t exchange reaction. From the cycle in fig. 6 one calculates a free-energy change of - 10.17 kJ equiv.--l in favour of the Ag(en)l complex. This is in good agreement with the experimentally obtained value (- 9.67 kJ equiv.-l), proving the reversibility of exchange reactions of complexed ion species. The good correspondence between calculated and experimental AG*[Ag(en)i- Cu(en)X+] values is remarkable in view of the fact that equilibria involving Ag(en)t were obtained at high pH, whereas the Na+-Ca2+, Na+-Cs+ and Ca2+- Cu(en)t+ equilibria were studied at neutral pH values.The formation of a Cu(en)i+ complex at high (en) concentration and high pH in solution and/or at the surface might invalidate the use of the Cu(en)i+-Ca2+ equilibrium at pH 7 in the cycle (fig. 6) to calculate the Ag(en)i-Cu(en)t+ equilibrium. However, the formation of Cu(en)i+ does not occur in solution under the present experimental conditions. Indeed, the U.V. spectrum of a solution containing 0.25 mol dmP3 Cu2+ and 1 mol dm-3 (en), which should be sufficient to make the triscomplex if K3 were 1 ,5 only shows the 18 000 cm-l band of the biscomplex Cu(en)i+. The observed d(OO1) spacing of 1.34 nm on Cu(en)t+-W.B. suspension at 0.1 mol dm-3 (en) and pH 11.5 also points to the formation of a biscomplex (see table 2).In Camp Berteau Montmorillonite, Cu(en)i+ species can be generated under anhydrous conditions from Cu2+-C.B. and gaseous (en). However, on contact with air, Cu(en)i+ is immediately transformed into Cu(en)i+. l5 The involvement of Cu(en)i+ in W.B. is therefore believed to be minimal. TABLE 2.--d(OO 1) SPACING (nm) OF FULLY EXCHANGED Ag(en)l AND Cu(en),"+ WYOMING BENTONITE CLAY AT LOW AND HIGH FREE (en) CONCENTRATION 2 x 10-4 no complex formed 1.5 1 x 10-1 1.34 1.34 DISCUSSION 'THERMODYNAMICS OF EXCHANGE OF COMPLEXES A characteristic feature of ion exchange involving complexes in montmorillonites is the overall increase of AG* of exchange, as compared with the value for the aqueous counterparts.Important changes occur in AG* when one of the exchanging ions is complexed with a ligand such as (en), as exemplified in the pairs Cu(en)i+-Ca2+, Ag(en)t-Cs+ and the Cu(en)i+-Cu2+ and Ag(en)t-Ag+ systems (see fig. 6 and table 1). Considering exchanges among complex ions one finds AG* changes of ca. 2 kJ equiv.-l for the exchanges of the (en) complexes of Ni2+ and Zn2+, Ni2+ and Cd2+ in Camp Berteau Montmorillonite.2 The selectivity among the aqueous ions, however, is very sma1l.l AG* [Ag(en):-Cu(en)i+] is similarly high (9.67 kJ equiv.-l) compared with AG* (Ag+-Cu2+) of the aqueous ions (0.967 kJ equiv.-l). TheseA. MAES, E. RASQUIN A N D A. CREMERS 2047 observations stress the importance of the ion-ligand and/or ion-water entities in determining the magnitude of the selectivity. The present results corroborate the observations made for alkali-metal16 and alkaline-earth17 ion exchanges in montmorillonite, in the sense that homovalent exchange equilibria involving complexes in montmorillonite are also exergonic when the most polarizable cation [Ag(en)i or Cu(en)i+] exchanges for a less polarizable one (Cs+ or Ca2+).This relation holds in cases where the polarizability differences are high (complex uersus aqueous ions) or for homologous series of ions as the alkali metals. Any change in the thermodynamic state functions of exchange relates to a difference of such functions, for both ions concerned, at the surface and in the solution phase. In terms of the Eisenmann18 model, exchanges in montmorillonite follow the low field-strength pattern, i.e.the sign of the AG, AH and TAS terms is governed by the solution properties of the ions. In the case of alkali-metal and alkaline-earth ion exchange the sign of all three thermodynamic functions is predicted. Those exchange reactions are always exergonic, exothermic and occur with a negative entropy change when an ion is replaced by one with smaller AG, AH and A S terms of hydration. The parallel with the polarizability parameter is self-evident. The influence of complexation on exchange becomes apparent by comparing the thermodynamic functions in the presence and absence of (en). Taking the Cu(en)i+-Ca2+ case as an example, the equations take the form - AG,en) = (Gcucen), - G d - (Gcucen,, - Gca) in which G stands for the free-energy content.The bar denotes the surface phase. The change in the overall AG of exchange on complexing corresponds in sign to the variation in the solution terms. Indeed, smaller AG values of hydration of the complex compared with the uncomplexed cation are expected. In the case of exchange among complexes the following relevant equations are written for the Ag+-Cu2+ system in the presence and absence of (en): AG = (Ccu - 2 G A g ) - (GCu - 2G,,) AG(en) = (Gcuccn), - 2-dAg(en),) - (Gcu(en), - 2G,g(en),)* Here again the variation in the overall AG* on complexing parallels in sign the variation in the solution term. The exact magnitude of the changes in hydration energy on complexing may be expected to be high, but are difficult to assess, and consequently it is not possible to deduce the separate contributions of both surface and solution terms to the overall AGe change.The importance of the surface term can be judged from the surface charge-density dependence of the Ca2+-Cu(en)i+ and Ca2+-Cu2+ exchange equilibria in a series of montmorillonites of variable charge density.12 AG(Ca2+-Cu2+) is almost independent of charge density, whereas AG[Ca2+-Cu(en)i+] decreases linearly from ca. - 10 kJ equiv.-l at 1.5 x mequiv. cm-2 to zero at vanishing charge density. (The charge in W.B. corresponds to 1 x lop7 mequiv. cmP2.) CHANGES I N SELECTIVITY WITH SURFACE COMPOSITION Kc[Ag(en),'-Cu(en)i+] decreases by more than two orders of magnitude with increasing bivalent ion loading (see fig. 3), in contrast to the commonly observed opposite behaviour in heterovalent exchange reactions among hydrated cati0ns.l.3, Two alternatives have been proposed to explain the general observation of increasing K,(Mn-M1+n) with M1+n ion loading in the case of heterovalent exchanges2048 EXCHANGE EQUILIBRIA IN WYOMING BENTONITE among aqueous metal ions in montmorillonites. First, the increase in K,(Na+-Ca2+) with Ca2+ content is explained as an autocatalytic effect provoked by the progressive collapse of double layers as exchange proceeds and by the fact that the selectivity for planar sites exceeds that for edge sites.19 Secondly, configurational entropy terms arising from the replacement of a monovalent by a divalent cation have been invoked3 to explain the overall entropy gain, because of the larger number of possibilities of arraying bivalent versus monovalent cations and also to explain the K,(M+-M2+) increase with bivalent ion loading.Heterovalent exchange among the complexes considered occurs in collapsed layers 4 A thick, as indicated by the d(OO1) spacings of the end members (table 2), in contrast to the more swollen state in the Na+-Ca2+ case. The selectivity changes with loading can therefore not be explained by a progressive interlayer collapse. Using a hypothetical charging-discharging process it has been predicted20 that the overall free energy of heterovalent exchange reactions increases with charge-density increase. This was verified for the Na+-Ca2+ exchange in montmorillonites of variable charge density.21 The mean charge density of W.B.clay itself is a composite of different charge densities.22 One may therefore expect to observe the highest affinity for the Cu(en)$+ ion at the small bivalent ion loadings corresponding to the exchange occurring in the highly charged regions of the clay. This is experimentally verified by the observation that the relative selectivity for Cu(en)i+ is highest at the low Cu(en)i+ constant side of the Kielland plot in fig. 3. The observed changes in Cu(en)i+-Ag(en)i selectivity with loading could then merely result from Coulombic interaction terms, and are thus of energetic origin. This is verified by two observations. (1) From the temperature dependence in fig. 3 it is calculated that the overall Ag(en)l-Cu(en)i+ exchange reaction itself is governed by enthalpic factors (AH x 12.5 kJ equiv.-l).( 2 ) Unpublished results on the charge-density dependence of the Ag(en)l - Cu(en)i+ exchange agree with the predicted increase in AG*[Ag(en)i-Cu(en)i+] with charge- density increase. The conclusion is therefore that configurational entropy terms may contribute to the same extent to the selectivity changes with loading in both heterovalent exchanges between complexed and uncomplexed cations. However, in the case of exchange among complexes energetic terms exceed the entropic contribution and are responsible for the observed selectivity changes with loading. We thank the Belgian Government (Programmatie van het Wetenschapsbeleid), the Fonds voor Kollektief Fundamenteel Onderzoek (F.K.F.O.) and the K. U. Leuven (3e cyclus Fonds) for financial support.A. Maes, P. Peigneur and A. Cremers, Proc. Znt. Cfay Con$ (Applied Publ. Ltd, 1976), p. 319. P. Peigneur, A. Maes and A. Cremers, Proc. Znt. Cfay Con$ (Elsevier, Amsterdam, 1979), p. 207. M. B. McBride, Clays Cfay Min., 1980, 28, 255. R. Van Bladel, G. Gaviria and H. Laudelout, Proc. 4th Inf. Cfay Con$ (Pergamon Press, Oxford, 1972), vol. 2. p. 15. G. L. Sillen and A. E. Martell, Stability Constants of Metal-ion Complexes (The Chemical Society, London, 1964-1971). C. N. Davies, Zon Association (Butterworths, London, 1962). G. L. Gaines and H. C. Thomas, J . Chem. Phys., 1953, 21, 714. M. M. El Sayed, R. G. Burau and K. L. Babcock, Soil Sci. SOC. Am., Proc., 1970, 34, 397. J. Dufey, Ph.D. Thesis (Universite Catholique de Louvain, 1974). lo R. G. Gast, R. Van Bladel and K. B. Deshpande, Soil Sci. SOC. Am., Proc., 1969, 33, 661. l1 A. Maes and A. Cremers, J . Chem. SOC., Faraday Trans. 1, 1978, 74, 1234. l2 A. Maes and A. Cremers, J . Chem. SOC., Faraday Trans. I , 1979, 75, 513. l 3 A. Maes, P. Marynen and A. Cremers, J . Chem. SOC., Faraday Trans. I , 1977, 73, 1297. l4 A. Maes, P. Peigneur and A. Cremers, J . Chem. SOC., Faraday Trans. I , 1978, 74, 182.A. MAES, E. R A S Q U I N A N D A. CREMERS 2049 l5 F. Velghe, R. A. Schoonheydt, J. B. Uytterhoeven, P. Peigneur and J. H. Lunsford, J . Phys. Chem., l6 H. Martin and H. Laudelout, J . Chim. Phys., 1963, 60, 1086. l7 H. Laudelout, R. Van Bladel, G. H. Bolt and A. L. Page, Trans. Faraday SOC., 1968, 64, 1477. l8 G. Eisenmann, Biophys. J., 1968, 2, 259. lS I. Shainberg and W. D. Kemper, Clays Clay Min., 1966, 14, 117. *O G. H. Bolt and C. J. Winkelmolen, Zsr. J. Chem., 1968, 6, 175. *l A. Maes and A. Cremers, J . Chem. SOC., Faraday Trans. I , 1977, 73, 1807. 22 M. S. Stul and W. J. Mortier, Clays Clay Min., 1974, 22, 391. 1977, 81, 1187. (PAPER 1 /211)

ISSN:0300-9599    

DOI:10.1039/F19827802041

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 2051-2058 Structural Analysis of Molten Na,BeF, and NaBeF, by X-ray Diffraction BY NORIMASA UMESAKI AND NOBUYA IWAMOTO Welding Research Institute, Osaka University, Yamada-kami, Suita, Osaka 565, Japan AND HIDEO OHNO* AND KAZUO FURUKAWA Molten Material Laboratory, Division of Nuclear Fuel Research, Japan Atomic Energy Research Institute, Tokai-mura, Ibaraki 3 19-1 1 , Japan Received 30th April, 1981 The structures of molten Na,BeF, (923 & 10 K) and NaBeF, (743 f 5 K) have been investigated by X-ray diffraction analysis. BeF, tetrahedra existing in the crystalline state persist as the fundamental structural unit in the molten state. Molten Na,BeF, contains mainly monomeric BeF,2- with four unshared F- comers. A configuration for four Na+ cations around a BeF, tetrahedron is suggested for molten Na,BeF,. In molten NaBeF,, short-chain anions, such as the common F- comer sharing dimeric Be$- and/or trimeric Be,F$j, mainly occur.The physical properties, such as self-diff~sion,l-~ electrical cond~ctance,~ viscosity5 etc., of systems containing molten alkali fluoroberyllates, XF(X = Li, Na or K) and BeF, are very different from those of molten alkali halides owing to the three- dimensional network character of BeF,. Moreover, the corresponding-states principle6 may be applied between the following pairs, (1) LiF-BeF, and MgO-SiO,, (2) NaF-BeF, and CaO-SiO, and (3) KF-BeF, and BaO-SiO,, in their intermediate composition regions under a reduced absolute temperature scale ; the phase diagrams and physical properties, such as ionic packing density, self-diffusion coefficient, viscosity coefficient and equivalent conductivity, in these pairs are in excellent agreement.We have measured the self-diffusion coefficients of lithium3 and fluorine'? in molten Li,BeF4 and LiBeF,, respectively. Fluorine is exchanged between neigh- bouring beryllate units with the rotation of beryllate anions, and there is a strong similarity between fluorine diffusion in molten fluoroberyllates and oxygen diffusion in molten silicates. Consequently, a direct comparison of the structures of molten fluoroberyllates with those of molten silicates is essential for elucidating similarities in the physical properties of these two systems. A structural investigation of the LiF-BeF, system using X-ray diffraction methods has been carried out by Vaslow and Narten.' They interpreted their results in the following manner: (1) the predominant structural unit is the BeFi- tetrahedron; (2) at low LiF concentrations these are linked together with each F- common to two tetrahedral units; (3) with increasing LiF concentration the linkage breaks down, but each Be2+ is maintained in a tetrahedral F- environment; (4) the Li+ ions are located outside the tetrahedral unit in locally disordered F- environments, the degree of disorder increasing with LiF concentration ; (5) the changes in interatomic spacing and coordination number that accompany fusion are consistent with the changes in density.Nevertheless, the configuration of the Li+ ions located around the BeF tetrahedral units is unknown.On the other hand, 205 12052 X-RAY STUDIES OF MOLTEN Na,BeF, AND NaBeF, Waseda and Toguris have reported an X-ray study of the molten systems MgO-SiO, and CaO-SO,. Unfortunately, no discussion of the structural similarities between these molten fluoroberyllates and silicates has been attempted. In this paper we report an X-ray diffraction analysis of molten Na,BeF, and NaBeF, and discuss the structural similarities between molten NaBeF, and CaSi0,.8 EXPERIMENTAL Prior to the X-ray diffraction experiment the samples were prepared as follows. Weighed amounts of NaF (analytical reagent grade, Merck) and BeF, (Rare Metallic Co) were thoroughly mixed in a dry glove box and melted in a Pt crucible under a He atmosphere. After being melted for ca.1 h the samples were cooled and crushed. The prepared samples were placed on a flat Pt tray (35 x 25 x 3 mm3) and heated in a small electric furnace made of Pt wire. The sample-heater assembly was enclosed under a He atmosphere by putting it in an air-tight chamber with a window of A1 foil of thickness 10 pm to allow passage of the X-ray beam. The temperature was controlled to within a maximum error of 10 K throughout the measurement. The X-ray diffraction experiment was carried out with the use of a 8-8 diffractometer with parafocusing reflection geometry and Mo K, (A = 0.7107 A) radiation monochromatized by a curved graphite monochromator mounted in the path of the diffracted beam. Slit systems of pf" and lo-lo were employed for the low (3 < 8" < 10) and high (8 < 8" < 50) scattering angles, respectively (0 is the scattering angle).The X-ray scattering intensities were measured at 0.25' intervals over the whole range of scattering angle using the step-scanning technique. Several runs were made in order to accumulate 2 x lo4 counts per datum point for low-8 and (3.4-4.0) x lo4 counts per datum point for high-0. After the X-ray intensities measured uia the two different slit systems had been normalized for the overlapping section they were corrected for background, polarization and Compton scattering, and then were scaled by means of the Krough-Moe-Norman method to the theoretical intensities arising from independent atoms contained in the stoichiometric unit. The radial distribution function D(r), the correlation function G(r) and the reduced intensity function S .i ( S ) are given as m m D(r) = 4nr2p0 X K + (QZ 2r/n fSmax S . i ( S ) sin (Sr) d S (1) i-1 i-1 J o (2) i-1 / i m \ where m is the number of atoms contained in the stoichiometric unit, Ki the effective electron number of atom i, po the mean electron density,h(S) the independent atomic scattering factor of atom i corrected for anomalous dispersion, cEh ( S ) the total coherent scattering intensity and S,,, the maximum value of S (= 4nsin8/3,) reached in the diffraction experiment. The constants used in the calculations of eqn (1)-(3) are given in table 1. RESULTS AND DISCUSSION Fig. I shows the observed reduced intensity functions S . i(S) of molten Na,BeF, (923 f 10 K) and NaBeF, (743 Ifr 5 K).The radial distribution functions D(r), the function D(r)/r and the correlation functions G(r) are shown in fig. 2 and 3. Ghosts originating from experimental errors and/or incorrect treatments are not recognized in these curves. As shown in fig. 2 and 3, these curves indicate peaks at 1.60-1.65 A, 2.22-2.35 A, 2.60-2.65 Aetc. Table 2 shows thedistances (rii 0.01 A)andcoordinationN. UMESAKI, N . IWAMOTO, H. O H N O A N D K. FURUKAWA 2053 TABLE PARAMETERS USED IN THE CALCULATIONS OF EQN (1)-(3) parameter Na,BeF, NaBeF, temperature / K 923 _+ 10 743 f 5 densityZ0/g CM-, 2.093 2.1 13 molar weight 130.997 88.991 effective electron number 11.84 11.98 3.73 3.77 8.65 8.75 13.5 13.0 5 ._ oj 0 2 4 8 10 12 14 S I P FIG. 1 .-Reduced intensity functions S .i ( S ) ; (a) molten Na,BeF, (923 + 10 K), observed (full line) and calculated (dotted line); (b) molten NaBeF, (743 f 5 K). numbers (nilj 0.1 atoms) of the nearest-neighbour ionic pairs in molten Na,BeF, and NaBeF, derived from the functions D(r)/r by assuming a Gaussian distrib~tion.~ In the known crystalline forms of some alkali fluoroberyllates, such as Na2BeF,10 and Li,BeF,,ll each Be2+ cation is tetrahedrally surrounded by four F- anions with a Be-F distance of 1.54-1.5 A and a F-F distance of 2.50-2.56 A. Therefore, the first peaks at 1.60-1.65 A and the third peaks at 2.60-2.65 A of the observed curves are due to Be-F and F-F pairs in the BeF, tetrahedra. In addition, the observed coordination number of the nearest-neighbour F- anions around Be2+, nBe,F, is ca.3.8-4.0. These results indicate that BeF, tetrahedra existing in the crystalline state persist as the fundamental structural unit in the molten state, and the mean distances between the ions in a BeF, tetrahedron become slightly longer in the melt than in the solid. The second peaks at 2.22-2.35 A are due to the nearest-neighbour Na-F pairs; the distance between these ions is nearly the sum of the ionic radii12 of Na+ (0.95 A) and F- (1.36 A). Our results are similar to those for the molten LiF-BeF, system reported previously by Vaslow and Narten.72054 X-RAY STUDIES OF MOLTEN Na,BeF, AND NaBeF, c 4 2 *G I 3 2 1 0 ..i 0 2 4 6 8 1 0 0 2 4 6 8 1 0 FIG. 2. FIG. 3. r l A rlA FIG. 2.-Radial distribution function D(r), function D(r)/r and correlation function G(r) of molten Na,BeF, (923 f 10 K).FIG. 3.-Radial distribution function D(r), function D(r)/r and correlation function G(r) of molten NaBeF, (743 f. 5 K). TABLE 2 . a B S E R V E D DISTANCES AND COORDINATION NUMBERS OF THE IONIC PAIRS IN MOLTEN Na,BeF, AND NaBeF,. INCLUDED FOR COMPARISON ARE THE DATA FOR MOLTEN CaO-SiO, REPORTED BY WASEDA AND TORGURI* liquid crystal liquid Na,BeF, NaBeF, Na,BeF, ionic CaO-SiO, (1873 K) sum of (923 & 10 K) (743 & 5 K) (DeganellolO) radii (Waseda et aL8) Be-F rBe-F/A 1.65 1.60 1.54-1.56 1.67 si-o r,-o/A 1.63 Na-F rNa-F/A 2.22 2.35 2.27-2.54 2.31 Ca-o rCa.O/A 2.41 nBe/F 4.0 3.8 nsi/o 3-8 nNa/F 2.8 1.8 nca/o 6.9 2.60 2.50-2.56 2.72 o-o ro-o/A 2.66 3.5 noio 5.8N. UMESAKI, N. IWAMOTO, H. OHNO AND K. FURUKAWA 2055 MOLTEN Na,BeF, In molten Na,BeF, the observed coordination number of the F-F pair in BeF, tetrahedra, n,/,, is 3.0.This suggests that BeF, tetrahedra exist mainly in an isolated form with four unshared F- corners, monomeric BeFi-, in molten Na,BeF,. Quist et a1.13 showed, using Raman spectroscopy, that monomeric BeFi- was the predominant beryllium-containing species in molten Na,BeF, and Li,BeF,. A similar result was obtained from e.m.f. measurement of the molten LiF-BeF, system by Holm and Kleppa.', They reported that, up to X(BeF,) = 0.33, corresponding to the composition Li,BeF,, the partial molar excess entropy of beryllium fluoride was positive and changed little with composition, owing to complete depolymerization of the three- dimensional network structure of beryllium fluoride ; i.e.the formation of monomeric BeFi- anions bonded by Coulombic forces. Most of the Naf ions might occupy the various stable positions around the monomeric BeFi- for some time and then migrate to other positions through voids in the melt. ( b ) FIG. 4.-Schematic illustration of typical configurations of Na+ ions around a BeF, tetrahedron : (a) comer-site position, (6) edge-site position and (c) face-site position. In order to advance understanding of the transport properties of molten Na,BeF,, the short-range arrangements were examined with reference to the three typical configurations of sodium atoms around the BeF, tetrahedron, as shown in fig. 4. In one configuration the sodium atom occupies a corner site in the BeF, tetrahedron, and Na-F and Na-Be distances are 2.31 and 3.96 A, respectively.In another configuration the sodium atom occupies an edge site and the Na-Be distance is 2.87 A. In the third configuration the sodium atom occupies a face-site and the Na-Be distance is 2.35 A. Various models of the short-range arrangement were constructed by combining the three typical configurations with one another in a stoichiometric unit and were examined by trial and error using the following Debye scattering equation15 r m 1 m m S i(W E fi ( S ) = I: C ni/jfi2(Slf3(S) exp (- bijS2) sin (Srii)/rii (4) li-1 1 i-lj-12056 X-RAY STUDIES OF MOLTEN Na,BeF, A N D NaBeF, wheref,(S) andf,(S) are the independent atomic scattering factors of atoms i and j , rij the distance between atoms i and j and bij the temperature factor, i.e.half the mean-square variation in rij. To check the ‘goodness of fit’ of the calculated reduced intensity to the observed one, the quantity R was introduced, where Of the possible models, the one giving the best fit was as follows: four sodium atoms surrounding the BeF, tetrahedron, two of these occupying corner-sites while the others occupy edge-sites. The calculated reduced intensity curve was refined by the least- squares method in the range 4.0 < S/A-l < 13.5 in order to achieve a better agreement with the observed curve, starting with parameters obtained from typical arrangements of the four sodium atoms around the BeF, tetrahedron as mentioned above. Here the interactions of atomic pairs beyond values of r = 6 A were neglected because peaks in the G(r) curve in this range were insignificant, and the numbers of atomic pairs were fixed in order to maintain the stoichiometric unit.The final parameters of the most probable short-range arrangement in molten Na,BeF, are listed in table 3. The final value of R for this model is 0.326, which was obtained by using the S . i ( S ) values above S = 4.0 A-1 with interval A S = 0.05 A-l. The calculated reduced intensity curve S. i(S) for this model is shown in fig. 1 (dotted line) and is in good agreement with the observed curve for S > A-l. However, agreement was poor for S < 4 A-l. This disagreement is due to the contribution from the long-range arrangement of molten Na,BeF,. TABLE 3.-PARAMETERS USED IN CALCULATING THE MOST PROBABLE STRUCTURAL MODEL BY EQN (4). ni,j, rij AND < Arij2 > ARE THE COORDINATION NUMBERS OF^ IONS AROUND ANY ORIGIN i ION, THE DISTANCE AND THE ASSOCIATED ROOT-MEAN-SQUARE DISPLACEMENT BETWEEN IONS i AND J, RESPECTIVELY m m Be F F F Na Be Na Be Na F Na F Na F Na Na Na Na 8.0 1 2 .0 2.0 2.0 6.0 4.0 6.0 0.7 1.3 1 . 6 5 2.62 3.35 4.24 2.22 3 . 9 9 4.72 3 . 3 0 5.51 ~ 0.100 0 . 1 4 2 0.085 0.095 0.098 0 . 1 3 1 0 . 1 8 4 0.245 0.283 a rij+O.O1 A: < Arij2 > 1+0.005 A. The coordination number of the nearest-neighbour Na-F pairs, nNa/F, in this model is 3.0, which is almost equal to the observed value (2.8). From measurements of the spin-lattice relaxation time for molten NaBeF, by means of nuclear magnetic resonance, Matsuo and Suzuki16 have pointed out the possibility of an edge-site position for the Na+ ion around the BeF, tetrahedron.These results are not inconsistent with those obtained by X-ray diffraction.N. UMESAKI, N. IWAMOTO, H. OHNO A N D K. FURUKAWA 2057 MOLTEN NaBeF, Up to the composition range of Na,BeF,, mixtures rich in NaF will chiefly contain Na+ cations, and F- and monomeric BeF:- anions. As the BeF, concentration increases, the polymerization process proceeds and pure molten BeF, forms a three-dimensional network structure of BeF, tetrahedra. NaBeF, is the intermediate phase between these composition extremes. From viscosity5 and thermodynamic data1, we can estimate that a mixture with the composition of NaBeF, contains a high percentage of polymeric anions of small chain size, such as (BenF3n+l)(n+1)- (n = 2, 3,4,.. . ), or closed rings, such as (BenF3n)n- (n = 3,4, 5,. . .). In molten NaBeF, the observed coordination number nF/F was ca. 3.5. The calculated average nF,F of chain anions such as (BezF7),-, (Be3FlJ4-, (Be4F13)5- and (BeF,)g- are 3.4, 3.6, 3.7 and 4.0, respectively. On the other hand, the value of +/F for closed-ring anions is always 4.0. Therefore, the possibility of the formation of longer-chain polymers or large-ring anions is small, and the probability of small-chain anions such as (Be2F7),- and (Be3Flo),- is considered to be large in molten NaBeF,. The existence of (Be2F7),- anion in molten Na,LiBe,F, has been shown by Raman spectroscopy . As noted in the introduction, a molten NaF-BeF, system should have essentially the same structure as molten CaO-SiO, system by the corresponding-states principle.6 X-ray structural analysis of molten CaO-SiO, (50-50 m/o) was carried out by Waseda and Toguri.s We examined the similarity between the structures of molten NaBeF, (measured temperature, T = 743 K; melting point, T, = 645 K; reduced absolute temperature scale, T/T, = 1.15) and CaO-SiO, ( T = 1873 K, T, = 1813 K and T/T, = 1.03).As shown in table 2, the coordination number of nearest-neighbour 0-0 pairs in molten CaO-SiO,, nolo (= 5.8), is different from nF/F (= 3.5) in molten NaBeF,. Waseda and Toguri8 also reported that the addition of CaO up to x(Ca0) = 0.57 had no effect on no,o in molten CaO-SiO,. This result indicates that the three-dimensional network structure is present at all CaO-SiO, compositions, in conflict with physical properties such as viscosity,ls electrical conductancelg etc., of this molten system.When CaO is added to molten SiO,, the viscosities of the resulting mixtures drop dramatically, owing to the rupture of three-dimensional network structure of the SiO, tetrahedra. In contrast Vaslow and Narten7 showed that nF/F in the molten LiF-BeF, system decreased from a value of 6(BeF,) to 2(LiBeF,) on the addition of LiF. Their results are in good agreement with ours. From the discussion above, we consider that there is room for reconsideration with respect to the X-ray study by Waseda and Toguri.8 CONCLUSIONS We can summarise the results obtained as follows. (1) In molten Na,BeF, and NaBeF,, BeF, tetrahedra exist as the fundamental structural unit.The mean distances between the atomic pairs in a BeF, tetrahedron become slightly longer in the melt than those found in the solid. The mean distances of the nearest-neighbour Na-F pairs are close to the sum of the ionic radii of Na+ and F-. (2) Molten Na,BeF, contains mainly monomeric BeFi- ions with four unshared F- corners. This result is supported by the results from Raman spectroscopy and e.m.f. measurements. Furthermore, four Na+ ions are situated in the following configurations around a BeF, tetrahedron : two occupy corner-site positions and the others occupy edge-site positions. (3) The probability of small-chain anions such as (BezF7),- and (Be,Flo),- is large in molten NaBeF,. 67 FAR 12058 X-RAY STUDIES OF MOLTEN Na,BeF, AND NaBeF, We thank Mr Y.Takagi (Research Laboratory of Engineering Materials, Tokyo Institute of Technology) for useful suggestions during this work. H. 0. and K. F. also express their appreciation of the encouragement given by Dr J. Shimokawa (Divisional Director of Nuclear Fuel Research, Japan Atomic Energy Research Institute). T. Ohmichi, H. Ohno and K. Furukawa, J. Phys. Chem., 1976, 80, 1628. * H. Ohno, Y. Tsunawaki, N. Umesaki, K. Furukawa and N. Iwamoto, J. Chem. Res. (M), 1978,2948. N. Iwamoto, Y. Tsunawaki, N. Umesaki, H. Ohno and K. Furukawa, J. Chem. Soc., Faraday Trans. 2, 1979, 75, 1271. G. D. Robbins and J. Braunstein, MSR Program Semiannual Prog. Rept. ORNL-4548, Oak Ridge National Laboratory, 1970, p. 156. S. Cantors, W. T. Ward and C. T. Moynihan, J. Chem. Phys., 1969, 50, 2874. K. Furukawa and H. Ohno, Trans. Jpn Znst. Metall., 1978, 19, 553. F. Vaslow and A. H. Narten, J. Chem. Phys., 1973, 59, 4949. Y. Waseda and J. M. Toguri, Metall. Trans., 1977, 8B, 563. C. A. Coulson and G. S. Roushbrooke, Phys. Rev., 1939, 56, 1216. lo S. Deganello, Acta Crystallogr., Sect. B, 1973, 29, 2593. l1 J. H. Bums and E. K. Gordon, Acta Crystallogr., 1966, 20, 135. l2 L. Pauling, Nature of The Chemical Bond (Cornell University Press, Ithaca, 3rd edn, 1960). l 3 A. S. Quist, J. B. Bates and G. E. Boyd, J. Phys. Chem., 1972, 76, 78. l4 J. L. Holm and J. 0. Kleppa, Znorg. Chem., 1969, 8, 207. l5 H. A. Levy, M. D. Danfold and A. H. Narten, ORNL-3960, Oak Ridge National Laboratory, 1966, l6 Y. Matsuo and H. Suzuki, personal communication. l7 L. M. Toth, J. B. Bates and G. E. Boyd, J. Phys. Chem., 1973, 77, 216. l8 J. O’M. Bockris and D. C. Lowe, Proc. R. Soc. London, Ser. A, 1954, 226, 423. lo J. O M . Bockris, J. A. Kitcherner, S. Ingatowicz and J. W. Tomlinson, Trans. Faraday Soc., 1952,48, *O B. C. Blanke, K. W. Foster, L. V. Jones, K. C. Jordan, R. W. Joyner and E. L. Murphy, MLM-1079, p. 1. 75. Mound Laboratory, 1958, p. 1. (PAPER 1 /695)

ISSN:0300-9599    

DOI:10.1039/F19827802051

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 2059-2072 Silanol and Water on Silica Studied by the CNDO MO Method BY KATSUYUKI TAKAHASHI Department of Mineral Science and Technology, Faculty of Engineering, Kyoto University, Sakyo, Kyoto 606, Japan Received 13th May, 1981 The CND0/2 molecular-orbital method has been applied to silanol or water on silica. Silanols on silica bind strongly to the surface, though they are more acidic than alcohols. Hydrogen-bonding hydroxyls have a lower bond strength than the surface hydroxyls. On the specific adsorption of water or water dimers on silica, two types of electron flow are observed: H,O (donor) -+ SiOH (acceptor) and SiOH (donor) + H,O (acceptor); in the former the oxygen of the water interacts with the hydrogen of SiOH and in the latter the hydrogen of the water interacts with the oxygen of SOH.The empirical shift of the hydroxyl absorption spectra of silanol and water on silica corresponds to the difference of calculated hydroxyl bond energies in the adsorption system and the isolated system. It is also theoretically explained that hydroxyls on silica tend to be centres for water adsorption and that it is difficult for the silica to release protons from its surface. On siloxane formation, the SiOSi bond angle may decrease in order to form a ring, which is the preferential shape for the Kiselev process. Following some recent papers, our understanding of silanol and water molecules adsorbed on a silica surface on a molecular basis has improved. The physicochemical studies of silanol or water sorption on silica surfaces, such as vapour sorption, calorimetry, i.r.and n.i.r., n.m.r., X-ray scattering, LEED or Auger spectroscopy, etc., are explained in several b00ks.l.~ Those studies indicated that the hydroxyls are centres for the adsorption of water on silica. Klier et al.4 looked into the mechanism of water adsorption and agglomeration on silica and proposed a model for the interaction of water with hydroxyls on silica using information obtained from the position and structure of fundamental and overtone H,O absorption bands. However, neither the binding state of silanol on silica nor the interaction of water molecules with hydroxyl groups on silica have been completely elucidated, This paper is an attempt, using the molecular-orbital method, to throw light on the electronic state and the binding state of both silanol and water molecules adsorbed on silica.The LCAO MO method used here is the completely neglect differential overlap (CND0/2) method, originaly developed by Pople et al.5-9 In quantum-chemical calculations of molecular interactions the molecular complex or cluster in question is considered as a supermolecule, for which the CND0/2 calculation in this paper is carried out. The interaction energy (AE) is determined as the difference between the energy of the supermolecule (Ep~Q) and the sum of the energy of the isolated constituent molecules ( E p , EQ) in their equilibrium geometry, eqn (1): AE = Ep*Q - (Ep + EQ). (1) The energy of the supermolecule or the single molecule including nuclear repulsion 67-2 20592060 SILANOL A N D WATER O N SILICA is divided into two parts, the one-centre energy part ( E A ) and the two-centre energy part (EAB), eqn (2): where A and B represent atoms.RESULTS AND DISCUSSION Silica exists in numerous crystalline modifications, the most important forms being quartz, cristobalite and tridymite. Their basic crystal unit is a tetrahedral structure of Si-0, as shown in the molecular diagram of SiOj- (Si-0 = 1.61 A*) in fig. 1. The charges of Si and 0 become ca. 0.5 (3.502 in electron densities) and - 1.1 (7.124 in electron densities) electrons, respectively, and the bonding strength of Si-0 is 01 - 0.409 ii2\, '\ 0. % \ \ \ 0, 0 4 FIG. 1 .-Charge distribution (electron units) and bond energy (atomic units) in SiOf. 08 P9 \ / Si, - 07 07- Si, \ 142.0' 141.06* ( p 5 1 .6 l i ,! 1.61 O1-J Si '\\ 2 \ \ 03 04 FIG. 2.-Schematic diagrams of quartz and symmetrical silica with their numbering. almost twice the antibonding strength of 0-0. Fig. 2 shows schematic structures of quartz and symmetrical silica with their numbering. Calculations were done for simplified models of silica, which were assumed to be Si,O;- consisting of two SiO, tetrahedra sharing a bridging oxygen atom, with the distance SiO = 1.61 A, the angle OSiO = 109.47' in a tetrahedral structure, 142.0' for quartz, 180.0' for cristobalite and tridymite and 141.06' for symmetrical silica. The bond length and bond angle we have taken for the equilibrium geometry compare well with 1.617 A and 143.3' calculated from the ab initio methodlo and with the experimental values 1.634 A and 144.1O l1 in disiloxane (H,SiOSiH,) 1.61 A and 143.7' in low quartz.12 Table 1 shows * 1 A E 10-10 m = 10-1 nm.K.TAKAHASHI 206 1 TABLE 1 .-VARIOUS VALUES FOR QUARTZ, CRYSTOBALITE, TRIDYMITE AND SYMMETRICAL SILICA quartz cristobalite tridymite symmetrical charge distribution 0, 0 5 0,-Si, Si,-0, Si , bond energy/a.u. total energy/a.u. - energy gap/eV transition energy/eV - 1.139 0.704 -0.816 - 0.490 - 0.347 129.78 12.81 6.64 - 1.097 0.695 - 0.808 - 0.486 -0.359 - - 129.83 12.51 6.28 - 1.096 0.693 - 0.808 -0.485 -0.358 .129.82 - 12.51 6.28 - 1.041 0.695 -0.817 - 0.478 - 0.344 129.77 12.73 6.48 charge distributions, bond energies, energy gaps and transition energies of the four crystalline models. Unbroken Si2-0, and Si,-0, bonds are weaker than other broken Si-0 bonds with the result that the charge of 0, between Si, and Si, becomes smaller than that of other oxygen atoms binding one silicon. As in the formation of a Si-0-Si bridge, electrons from oxgyen are attracted to two silicon atoms, the electron densities between unbroken Si and 0 became smaller than between broken Si and 0.Various values for the four structures of Si,O!- are approximately the same in table 1, and the value of the energy gap, ca. 12 eV, is similar to a value of 11 eV for a-quartz.13 The CND0/2 results for perfect a-quartzl* and Si20,H,15 and the ab initio STO-3G result for Si,07H,14 are -0.6, -0.7 and -0.6, respectively for the formal charge of the central oxygen atom. Our calculations show that the formal charge of the central oxygen is -0.8 (for si20;-, table 1) and -0.7 (for Si207H,, see later, fig.3), close to the above values and to the oxygen charge in quartz of -0.7 inferred from the shift of the Si Kg X-ray emission spectral6 and the charge of -0.7 calculated from the ab initio SCF method for molecular Si02.17 After the supposition18 that the valences of silicon atoms on surfaces must be saturated with silanol groups, Carman19 visualized the structure of a particle of colloidal silica as consisting of a network of interlinked SiO, tetrahedra with hydroxyl groups attracted to the surface due to the tendency of silicon to complete tetrahedral coordination. Boehrnz0 then considered each particle of silica to be a macromolecule of polysilicic acid.Taking symmetrical silica as the basic structure of the various silicates in order to study the electronic state of silanol or water on silica (see later), we first calculated the total energy of the Si,O;- + 2H+ system as a function of angles from the Si-0 axis and fixed the 0-H bond length as 0.96 A. The total energy of the system falls as the angle from the Si-0 axis increases, attaining its lowest total energy, i.e. the most stable state, at an angle SiOH = 137.5', and then rises up again. This angle is different from the angle of ca. 113' for the SiOH bond angle as found by Peri from i.r. spectroscopy.21 However, the difference in the electron densities and bond energies in both systems is only ca. orders of magnitude in the calculation. The total energy (- 133.398 a.u.) in the geminal configuration2 connecting two hydroxyls to one silicon atom is higher than that (-133.501 a.u.) in the vicinal configuration2 connecting one hydroxyl to one silicon atom, because of higher E A B than E A in eqn (2).We next changed the bond length of 0-H and fixed the angle SiOH as 137.5'. The most stable state appears at 0-H = 1.06 A, slightly longer than the normal bond length of 0.96 A. On addition of the six protons to Si207 ,- , five configurations were considered (table 2) and their total energies calculated with the2062 SILANOL AND WATER ON SILICA 0.881 ‘07 Si, 2.383 -0.583 6.536 \ 0, 6.728 -0.596 1 0.884 H1o Si2 2.389 01 ~ -0.76\;\ -0.476 6.551 I j._i ~ o . 4 7 9 0 4 / ,’ ,’ * $:. 7 63 H13 H 12 0.884 FIG. 3.-Electron density (electron units) and bond energy (atomic units) in the adsorption system Si,O:- + 6H+.TABLE 2.-TOTAL ENERGY OF FIVE CONFIGURATIONS OF THE ADSORPTION SYSTEM si,o!- -k 6H+ configuration 1-2 cis 1-2 trans 5-2 cis 5-2 trans 5-2 trans 7-6 cis 7-6 trans 5-6 cis 5-6 trans 5-6 cis total energy/a.u. - 1 3 8 . 1 0 4 - 1 3 8 . 1 0 6 - 1 3 8 . 1 0 4 - 1 3 8 . 1 0 6 - 1 3 8 . 1 0 7 fixed angle SiOH = 1 37S0 and the fixed 0-H bond length of 1.06 A. With reference to fig. 3, 5-2 trans, 5-6 cis in table 2 means that H,, (or H13) locates itself at a trans position against 0, on the Si2-03 (or Si2-04) axis, and Hi4 (or HIS) locates itself at a cis position against 0, on the Si,-0, (or Si,-0,) axis. The other expressions have similar meanings. The optimal state among the five Si,O$- + 6H+ systems could be obtained for 5-2 trans, 5-6 cis, whose molecular diagram is illustrated in fig.3. When six protons add to Si,O;-, hydrogen has a slight positive charge, +0.12 (0.88 in electron densities), due to electron flow from silicate to hydrogen. Consequently, the positive charge of Si increases as the negative charge of 0 decreases due to the decreasing electron densities of both Si and 0. The large value of the bond energy of hydroxyl groups, e.g. 0,-H,, or 07-H11, means that it is difficult to remove silanol from the silica surface. Note, however, the weak hydrogen-bonding strength of O7-HlO (- 0.01 6 a.u.), which is approximately equivalent to one-fiftieth that of O,--H,o.K. TAKAHASHI 2063 When the silanol groups of silica are compared with the hydroxyl group of alcohol, the former are considered more acidic than the latter when measuring the shifts of the infrared 0-H stretching bands22 upon admixture with the bases ether and mesitylene and the acid phenol, respectively.The energetic calculations (AE), using a configuration given by ref. (23), for reactions to release a proton, eqn ( 3 ) : difference bond energy of energy of leaving H CH,OH -+ CH,O-+ H+ AE = 0.898 -0.741 a.u. Si20,6H + Si20,5H-+ H+ AE = 0.853 - 0.736 a.u. } ( 3 ) result in silanol on silica being more acidic than the hydroxyl group in alcohol. Furthermore, the bond energy of the hydroxyl groups is lower in silica than in alcohol, so the silanol on silica may be more acidic. The CND0/2 calculations for a series of compounds containing the Si-0 bond indicated that the experimental oxygen bond angles in these compounds are reproduced reasonably well without the use of Si ( 3 4 functions.24 However, we will investigate the effect of the Si 3d orbital on the electron densities of each atom as well as on the bond energies of Si-0 and 0-H or the total energy of the Si20!-+6H+ system, making use of the work of Santry and SegaP on the effect of 3d functions of various second-row elements on the charge distributions, bond angles and dipole moments TABLE 3.-EFFECT OF si d ORBITALS ON ELECTRON DENSITY, BOND ENERGY AND TOTAL ENERGY OF THE ADSORPTION SYSTEM, Si,O!- + 6H+ bonding electron density 0, Si, 0 5 HI0 bond energy/a.u. 0,-Si, Si,-0, O,--H,, E'4 E m total energy/a.u.total 6.551 2.389 6.728 0.884 0.476 - 0.596 - 0.763 - 138.107 - 130.509 - 7.599 6.326 3.418 1.861 (sp) 1.557 ( d ) 6.267 0.828 -0.849 -0.817 -0.746 - 140.883 - 129.750 - 11.133 of molecules.The results of the sp and spd methodsa are shown in table 3. The Si 3d orbital causes not only a large increase in the electron densities of Si (the positive charge of Si decreases), but also a decrease in the electron densities of 0 (negative charge of 0 decreases) and H (positive charge of H increases). Note that the Si 3d orbitals are to play a great role in electron populations on Si. Regarding the bonds, the Si-0 bond strength increases as the 0-H bond strength decreases because of the Si 3d orbital. The total energy of the system becomes lower, i.e. the system is stabilized due to mutual interaction energies (,TAB), including the Si 3d orbitals, between different atoms.2064 SILANOL A N D WATER ON SILICA The structure and physicochemical properties of water are explained in a few book^.^^.^^ As far as water clusters are concerned, recent quantum-mechanical calculations were carried out from dimer to hexamer or with water-ion interactions.The dimers of water and their configurations, stabilization and hydrogen bonding have been discussed using ab initio method^.^^-^^ Dimeric water was here considered as only a linear dimer in its most stable state, though other bifurcated and cyclic dimers may also be present. Fig. 4 shows molecular diagrams of the water molecule and of the 0.142 0.155 H2 "@qJ -0.060 -0.737 -0.326 H3- 02 Y.755 A 0 .7 5 5 v.756 01-0.2 84 ,Qpoo;;----- 0.1 76 H1 U . H, O . l i 2 0.155 0.;; 6 (a) ( b ) FIG. 4.-Charge distribution (electron units) and bond energy (atomic units) in water and a linear dimer of water. linear dimer of water in the most stable configuration obtained using the CND0/2 calculation, with the bond length of 0,-H, (or 0,-H,) taken as 0.9572 A and the angle H,O,H, as 104*52', in common with the values for water. In this water dimer the bond length of 0,-0, for the lowest energy was determined to be 2.53 A, different from the value of 3.00 A obtained in an ab initio calculation by Popkie et ~ 1 . ~ ~ On going from water dimer to the water molecule, the gross electron densities of H,O,H, decrease slightly as those of H,O,H, increase slightly, so that the linear dimer may form a molecular complex as H,O,H, becomes a donor and H,O,H, an acceptor.The new hydrogen bond in the water dimer is calculated to have a strength of -0.06845 a.u. (-42.95 kcal mol-l), one-tenth the strength of the 0-H bond of water, resulting in weakening of its nearest-neighbour bond. If the 0-H binding energy in isolated water molecule is estimated to be 109 kcal mol-l, the H,-O, bond energy is ca. 10 kcal mol-1 in the range of empirical hydrogen-bond energies. In water oligomers and tetrahedral clusters the hydrogen-bond energies obtained using the CND0/2 method (8.7-10.1 kcal m01-l)~~ are in fairly good agreement with those obtained using the ab initio method (6.1-8.2 kcal m ~ l - l ) . ~ ~ The stabilization energy for water dimer formation was calculated to be 6.5 kcal mol-l, larger than the value of 5.1 kcal mol-1 when including its correlation energy,35 and close to the calorimetric heat of a single hydroxyl-water bond, 6 kcal m01-l.~~ In addition, the values of dissociation energy of the linear water dimer obtained by different methods are similar, e.g.CND0/2 (5.0-8.7 kcal mol-l), INDO (4.0-6.9 kcal mol-l), STO-3G (6.0 kcal mol-l), 6-31G (5.6 kcal mol-l) and extended +polarization (4.7-5.1 kcal m ~ l - l ) . ~ ' 39 Therefore, the CND0/2 calculation is reliable for hydrogen-bonded systems. Five schematic models of water adsorbed on silica using the sp method are illustrated in fig. 5. In structure (1) oxygen (Ole) of water strongly interacts with hydrogen (H,,,) of silanol, while in structure (11) hydrogen (H,,) of water strongly interacts with oxygen (0,) of silanol.Structure (111) shows that another water molecule behaves like the water molecule in structure (I). In both structure (IV) and structure (V) another water molecule adsorbs and forms a linear water dimer, analogously to structures (I) and (11),K. TAKAHASHI 2065 (IV) (V) FIG. 5.-Five schematic models of water adsorbed on silica and the change of electron density and bond energy of the adsorption system as compared with the isolated system. AElkcal mo1-l. (I) -7.84, (11) -8.10, (111) - 16.04, (IV) - 13.79, (V) -7.35. respectively. The most stable configuration of water on silica was first determined for structure (I). Fixing 0,-0,, = 2.53 A on an extension of the 0,-H,, line, we rotate water H,,O,,H,,, taking O,, as the centre and angles XO,,H,, as 8, where X means a horizontal axis through O,,.The total energy of the system (Si20,6H + H,O) as a function of angles (0) indicates the most stable state at an angle XO,,H,, = 25O, a bisector of H,,O,,H,, being located at 1 1 . 2 4 O anticlockwise from the line of O,H,,O,,. Secondly, by changing the 0,-0,, distance whilst fixing the angle XO,,H,, as 2 5 O , we find the total energy of this system indicates the system is at its most stable when 0,-0,, = 2.53 A. For structure (11) we changed the angles XO,H,, = 8, where X means a horizontal axis through 0,, keeping O,H,,O,, in a straight line and fixing 0,-O,, as 2.53 A. The lowest energy for this system was obtained at an2066 SILANOL A N D WATER ON SILICA angle XO1Hl7 = 60'.By changing the distance between 0, and 0,, at the definite angle XO1Hl7 = 60Othe most stable state exists at 0,-O,, = 2.53 A. In the next step, the water molecule H1701,H,, was rotated around OI6, in the plane of the paper. The total energy of the system is found as a function of angles, ZO,,H,, = 8, where Z means a vertical axis through O16. The optimal angle becomes ZO,,H,, = 26O, O,,H,, deviating slightly from the 0,01, line by 4' in a counterclockwise direction. The interaction energy (AE) for adsorption of a water molecule on disiloxane (H,SiOSiH,) at its equilibrium geometry is -5.43 kcal mol-l by the STO-3G method4, and -5.35 kcal mol-l by the 4-31G method.40 Also, the adsorption energy of a water molecule on a surface cluster of silica was calculated to be -9.95 kcal mot1 by the CNDO MO method,41 while the experimental enthalpy of a water molecule on silica is ca.6 kcal m01-l.~~ Our calculations for :he adsorption of a water molecule on silica (Si2O76H), using a different model from those mentioned above for disiloxane and silica, give -7.8 kcal mol-l for structure (I) and -8.1 kcal mol-l for structure (11), in good agreement with the above values. Fig. 5 illustrates that the plus and the minus signs for the electron density and bond strength are increased and decreased on adsorption of water on silica, in comparison with those in the adsorption and isolated systems. When water adsorbs on silica in structure (I), the electron densities of 0,, in adsorbed water and of 0, in silanol increase as those of H,, (or H18) in water and H,, in silanol decrease.Electrons flow from the adsorbed water molecules to silica, so that the gross electron densities of water may decrease by -0.039, as shown in table 4. On the contrary, for structure (11) electrons flow from silica to the adsorbed water, the gross electron densities of water increasing by +0.032. For structures (111) and (IV), electrons flow from two water molecules to silica, as in the case with structure (I), decreasing the gross electron densities of dimeric water by - 0.044 and - 0.029, respectively. On the other hand, for structure (V) electrons flow from silica to water Hl70,,Hl,, as in the case with structure (II), so the gross electron densities of this water molecule increase by + 0.027. Another water molecule H,OO,gH,, scarcely changes in its gross electron densities. In short, we can classify two types of electron movement: one is H,O (a donor) + silica (an acceptor), for structures (I), (111) and (IV), and the other is silica (a donor) -, H,O (an acceptor), for structures (11) and (V).As for bonding states, if a new hydrogen bond is formed between 0,, and Hlo, as shown for structure (I), the bond strength of the nearest-neighbour bond H,,-0, decreases and the next nearest neighbour bond 0,-Si, increases to a lesser extent. The successive alteration of bond strength which occurs for silica and water also occurs when a new hydrogen bond is formed. We note that the frequency of the hydrogen- bonded OH should be lowered. It is already known that the frequency of hydrogen- bonded OH is lower than the frequency of the surface OH43344 and also that the frequency of the surface hydroxyl absorption band is lowered due to the interaction between surface OH and adsorbed molecules such as 0,, N2,45 acetone, ammonia,46* 47 toluene, ethylben~ene,~, e t ~ .~ ~ Klier et aL4 classified the interaction of partially hydrophobic silicas with water and discussed the structures and frequency shift in SOH-H,O on a donor-acceptor basis. As shown in table 4, for structures (I), (11) and (111) the shift of OH bond energies of silanol and water, determined by the difference of OH bond energies both in the adsorption system and in the isolated system, is likely to correspond to the spectroscopic shift of H20 and SOH. The shift of OH bond energies for structures (IV) and (V) tends to be similar to the shift for structures (I) and (11), respectively.The new hydrogen bond formed between water and silanol was calculated to be at ca. -0.09 a.u., an actual hydrogen-bonding value being ca. 13 kcal mol-1 by analogy with the 0-H binding energy in isolated water. The bond strength (-0.059 a.u.) between water far from silica and silanol isK. TAKAHASHI 2067 TABLE 4.-vARIOUS CHANGES IN THE FIVE ADSORPTION SYSTEMS AS COMPARED WITH EACH ISOLATED SYSTEM configuration difference of electron density of H,O electron flow difference of bond energy of OH in silica/a.u. shift difference of bond energy of OH in H20/a.u . shift new bond of OH/a.u. - 0.039 H20 -+ silica -0.019 large - 0.006 (O16H17) - 0.01 1 (O16H18) small - 0.096 + 0.032 silica -+ H20 -0.001 small - 0.023 (O16H17) + 0.001 (O16H18) large - 0.090 H,O + silica -0.019 (OIHlo) - 0.024 (O,Hll) large - 0.006 (O16H17) -0.014 (O16H18) - 0.006 (019H20) - 0.005 (O19Hzl) small - 0.095 (O16H10) - 0.092 (O19Hll) -0.010 (O16H11) configuration difference of electron density of H,O electron flow difference of bond energy of OH in silica/a.u.shift difference of bond energy of OH in H,O/a.u. shift new bond of 0Hla.u. -0*04 (H17016H18) -o*029 (H20019H21) H,O -+ silica -0.020 (OIHlo) - 0.01 7 (07H11) large - 0.005 (O16H17) -0.010 (O16H18) small - 0.102 (H10016) - 0.059 (Hl1019) - 0.03 1 (0,H18) ~ ~~ +o'027 (H17016H18) +O*Oo2 (H20019H21) silica + H,O 0.000 (07Hll) - 0.002 (OIHlo) small -0.019 (O16H17) + 0.004 (O16H18) large - 0.084 (0,H17) -0.016 (OlH18) - 0.005 (HloO19) approximately half the bond strength (-0.102 a.u.) between water near silica and silanol with structure (IV) and with structure (V) the difference is of the order of a fifth ( - 0.0 16 and - 0.084 a.u.). We will now see how an oxygen atom removing a proton and a silanol group com- pare as an adsorption centre from a molecular point of view.Fig. 6 shows the main bond energies in the adsorption system of structure of fig. 5 without the H,, and H,, hydrogen atoms. The oxygen-oxygen bonds, especially the 0,-0,, bond, have relatively large antibonding values which are different from the relatively large bonding values of the Hlo-Ol6 bond of structure (I). The hydrogen-oxygen bonds2068 SILANOL A N D WATER ON SILICA F15 -0.004 0 " 1 4 ; H FIG.6.-Main bond energies (atomic units) for structure (I) of fig. 5 without the H,, and H,, hydrogen atoms. also have comparatively smaller bond energies. The energy difference of this water adsorption was calculated to be AE = 2.53 kcal mol-l, so that the adsorption system may not be theoretically stabilized by water adsorbed on silica. Although there is the possibility that hydrogen Hi7 or H,, interacts strongly with 0, or 07, we shall not discuss that here as it is similar to structure (11) of fig. 5 . On the other hand, the five structures of fig. 5 are stabilized by water adsorption on silica, as indicated by the values of AE. The isolated hydroxyl groups were far more resistant to removal from the surface during evacuation as the hydroxyl groups already involved in hydrogen bonding were more favourably situated for This is supported by the fact that although a narrow absorption band at 3770 cm-l due to free hydroxyl groups and a broad band at 3450 cm-1 due to hydrogen-bonded hydroxyls were observed in the i.r.spectrum of silica gel evacuated at 350 O C , the low-frequency band disappeared after evacuation at 800 O C while the band at 3770 cm-l remained.51 Our calculations also showed that hydrogen-bonded OH is less strongly bound than free OH. Little et al.43 discussed the removal of water from the neighbourhood of the hydroxyl groups, and Young52 has pointed out that this reaction is reversible at low temperatures and is irreversible at high temperatures. The CND0/2 calculation results in AE = 151.46 kcal mol-1 as the energy difference for eqn (4) from which we deduce that we need to heat silica to a very high temperature in order to obtain siloxane by removal of H20.Fig. 7 is a molecular diagram of siloxane with 0, located at the symmetrical position against 0, on the Si,-Si, line. In comparison with fig. 3 the Si,-O, and Si,-0, bonds become weaker. Furthermore, the 01-Si2 bond or the 0,-Si, bond is antibonding while the 0,-0, bond is strongly bonding,K. TAKAHASHI 2069 1.069 H12 / I 6.539 / / 6.418 O j O5 6.251 0 , 1 3 4 /0.007 l,,?i\3. 466 -0.109 I / FIG. 7. 1.051 1.034 FIG. 8. 66 0 FIG. 7.-Electron density (electron units) and bond energy (atomic units) in the Si,O:-+ 4H+ system-for the formation of siloxane. FIG. 8.-Electron density (electron units) and bond energy (atomic units) in the Si,OE-+ 4H+ system without formation of siloxane.-0.418 -0.418 0.108 0 7 6.785 I I 6.785 0, eqn (6) (the lower part is symmetrical). FIG. 9.-Electron density (electron units) and bond energy (atomic units) of the upper part of Si,Oy; from2070 SILANOL A N D WATER ON SILICA the opposite to the case of the 0,-0, bond in fig. 3. We now consider the case where one silicon atom is missing a tetrahedral bond while the other SiO, has an exact tetrahedral structure, eqn ( 5 ) : A calculated result is shown in fig. 8. Looking at both fig. 7 and 8, we see that the 0,-Si, and 0,-0, bond energies are different: in fig. 8 the 0,-Si, bond is weakly bonding and the 0,-0, bond strongly antibonding, while in fig.7 the opposite is the case. These calculations result in the formation of a strained Si-0-Si bridge, similar to that found by Hocky et af.53 on the irreversible removal of hydroxyls from silica. In fact, the angle of Si,O,Si, will decrease on the formation of siloxane. Another calculation for siloxane formation using Kiselev’s prop~sition~~ was done for eqn (6) : - 0 * +H,O 0 0 I / o O,,\ 1 0 kSi\,/ k0 Si ‘ The molecular diagram of the product compound is given in fig. 9. The bonds between silicon and oxygen become strongly bonding to form a ring, in spite of their antibonding in fig. 7. This indicates that the ring formation of siloxane in fig. 9 is rather easier than that in fig. 7 : the Kiselev process tends to be preferential for ring formation of siloxane.CONCLUSION CNDO/2 calculations show that silanols bound to silica need to be heated to high temperatures in order to remove them from the silica surface, although they are more acidic than alcohol groups. The bond energies of hydrogen-bonding hydroxyls have a lower frequency than those of the surface hydroxyls, in agreement with the empirical results. On specific adsorption of water or dimeric water on silica, two types of electron movement occur: one is H,O (a donor) -+ silica (an acceptor) and the other is silica (a donor) -+ H20 (an acceptor). The empirical shift of the OH absorption spectra of silanol and water due to their mutual interactions on silica corresponds to the difference of their hydroxyl bond energies in the adsorption system and in the isolated system.Hydroxyls on silica are, theoretically, centres for water adsorption, while silica which is releasing protons from its surface is not. In siloxane formation, the Si-0-Si bridge leads to a strained bridge and the Kiselev process becomes the preferential one.K. TAKAHASHI 207 1 I am grateful for permission to use the FACOM 230-75 computer of the Data Processing Center, Kyoto University. I also thank a referee for helpful comments. J. Texter, K. Klier and A. C. Zettlemoyer, in Progress in Surface and Membrane Science, ed. D. A. Cadenhead and J. F. Danielli (Academic Press, New York, 1978), vol. 12, p. 327. V. L. Snoeyink and W. J. Weber, in Progress in Surface and Membrane Science, ed. J. F. Danielli, M. D. Rosenberg and D. A. Cadenhead (Academic Press, New York, 1972), vol.5 , p. 63. L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, New York, 1966), p. 228. K. Klier, J. H. Shen and A. C. Zettlemoyer, J . Phys. Chem., 1973, 77, 1458. J. A. Pople, D. P. Santry and G. A. Segal, J . Chem. Phys., 1965, 43, S129. ti J. A. Pople and G. A. Segal, J. Chem. Phys., 1965, 43, S136. J. A. Pople and G. A. Segal, J . Chem. Phys., 1966, 44, 3289. D. P. Santry and G. A. Segal, J . Chem. Phys., 1967, 47, 158. D. P. Santry, J. Am. Chem. Soc., 1968, 90, 3309. l o H. Oberhammer and J. E. Boggs, J . Am. Chem. SOC., 1980, 102, 7241. l 1 A. Almenningen, 0. Bastiansen, V. Ewing, K. Hedberg and M. Traetteberg, Acta Chem. Scand., 1963, 17, 2455. l 2 Y. Le Page and G. Donnay, Acta Crystallogr., Sect. B, 1976, 32, 2456.0. V. Krylov, Catalysis by Nonmetals (Academic Press, New York, 1970), p. 256. l 4 H. Haberlandt and F. Ritschl, Phys. Stat. Solidi B, 1980, 100, 503. l 5 B. H. W. S. de Jong and G. E. Brown Jr, Geochim. Cosmochim. Acta, 1980, 44, 491. l6 V. S. Urusov, Geochem. Int., 1970, 7, 143. l7 J. Pacansky and K. Hermann, J. Chem. Phys., 1978, 69, 963. U. Hofmann, D. Endell and D. Wilm, Angew. Chem., 1934, 47, 539. P. C. Carman, Trans. Faraday Soc., 1940, 35, 964. 2o H. P. Boem, A h . Catal., 1966, 16, 179. J. B. Pen, J . Phys. Chem., 1966, 70, 2937. 22 R. West and R. H. Baney, J . Am. Chem. Soc., 1959, 81, 6145. 23 L. E. Sutton, Tables of Interatomic Distances and Configurations in Molecules and Ions (The Chemical Society, London, 1958) Special Publication no. 11, M 114.24 J. A. Tossell and G. V. Gibbs, Acta Crystallogr., Sect. A, 1978, 34, 463. 25 D. Eisenberg and W. Kauzmann, The Structure and Properties of Water (Oxford University Press, 26 G. J. Safford and P. S. Lering, in Progress in Surface and Membrane Science, ed. J. F. Danielli, M. D. *’ A. D. Buckingham, in Intermolecular Interactions: From Diatomics to Biopolymers, ed. B. Pullman London, 1969), p. 1. Rosenberg and D. A. Cadenhead (Academic Press, New York, 1973), vol. 7, p. 231. (John Wiley, New York, 1978), p. 1. K. Morokuma and L. Pedersen, J . Chem. Phys., 1968, 48, 3275. 29 G. H. F. Diercksen, W. P. Kraemer and B. 0. Ross, Theor. Chim. Acta, 1975, 36, 249. 3n 0. Matsuoka, E. Clementi and M. Yoshimine, J. Chem. Phys., 1976, 64, 1351. 31 R. C . Kerns and L. C . Allen, J . Am. Chem. SOC., 1978, 100, 6587. 32 H. Popkie, H. Kistenmacher and E. Clementi, J . Chem. Phys., 1973, 59, 1325. 33 J. R. Hoyland and L. B. Kier, Theor. Chim. Acta, 1969, 15, 1. 34 J. E. Del Bene and J. A. Pople, J . Chem. Phys., 1970, 52, 4858. 35 P. A. Kollman, J . Am. Chem. Soc., 1977, 99, 4875. 36 K. Klier and A. C. Zettlemoyer, J. Colloid Interface Sci., 1977, 58, 217. 37 A. S. N. Murthy and C. N. R. Rao, J . Mol. Struct., 1970, 6, 253. 3R P. A. Kollman and L. C. Allen, Chem. Rev., 1972, 72, 283. P. Schuster, in The Hydrogen Bond, ed. P. Schuster, G. Zundel and C. Sandorfy (North Holland, Amsterdam, 1976), vol. 1, p. 25. 40 J. Sauer, P. Hobza and R. Zahradnik, J. Phys. Chem., 1980, 84, 3318. 41 D. Mikheikin, I. A. Abronin, G. M. Zhidomirov and V. B. Kazansky, J . Mol. Catal., 1977, 3, 435. 4 2 E. M. Flanigen, J. M. Bennet, R. W. Grose, J. P. Cohen, R. L. Patton, R. M. Kirchner and J. V. O3 L. H. Little and M. V. Mathieu, Proc. 2ndZnt. Cong. Catalysis (Edition Technip, Pans, 1961), p. 771. 44 J. B. Pen and R. B. Hannan, J . Phys. Chem., 1960, 64, 1526. 45 R. S. McDonal, J . Am. Chem. Soc., 1957,79, 850. 4 6 M. Folman and D. J. C. Yates, Proc. R. SOC. London, Ser. A, 1958, 246, 32. 4 7 N. W. Cant and L. H. Little, Can. J. Chem., 1964, 42, 802. 49 A. V. Kiselev and V. I. Lygin, Surf. Sci., 1964, 2, 236. Smith, Nature (London), 1978, 271, 512. A. N. Sidorov, Opt. Spectrosc. (USSR), 1960, 8, 424.2072 SILANOL A N D WATER O N SILICA R. S. McDonald, J . Phys. Chem., 1958, 62, 1168. 51 A. Chevet, J . Phys. Radium, 1953, 14, 493. 5 2 G. L. Young, J . Colloid Sci., 1958, 13, 699. 53 J. A. Hocky and B. A. Pethica, Trans. Faraday Sac., 1961, 57, 2247. j4 A. V. Kiselev and V. I. Lygin, Infrared Spectra of Surface Compounds (John Wiley, New York, 1975), p. 100. (PAPER 1 /767)

ISSN:0300-9599    

DOI:10.1039/F19827802059

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1982, 78, 2073-2084 Coupling between Binding-induced Conformational Phenomena and Stereospecific Effects in Asymmetric Reactions BY MARIO BARTERI AND BASILIO PISPISAT Istituto di Chimica Fisica, Universita di Roma, 00185 Roma, Italy Received 27th May, 198 1 The oxidation by H,O, of L-( + )-ascorbate anion in the presence of 2,2’,2”,2”’-tetrapyridineiron(111) complex ions anchored to poly(L-glutamate) (FeL) or poly(D-glutamate) (FeD) has been studied at pH 7.0 and varying complex-to-polymer-residue ratio [C]/[P]. The reaction follows two parallel routes; one corresponds to an electron-transfer process within a substrate-catalyst adduct and the other refers to an uncatalysed pathway to the dehydroascorbic acid. Unusual phenomena are observed in the catalysis in the sense that only the conformational dissymmetry of the active sites, arising from the binding-induced coil-to-a-helix transition of polypeptide matrices by Fe”’ complex counter-ions, is able to impart stereospecific effects in the reaction.Evidence is produced to show that stereoselectivity is driven by activation entropy. The effect probably arises from the formation of a rather rigid precursor complex, with the optically active substrate molecules bound to the chiral residues of the ordered polymer surrounding the active sites. The stereochemistry of such an intermediate allows the reaction to proceed only by a remote electron-transfer pathway, through the quaterpyridine ligand of the metal chelate. Evidence suggests that the asymmetric [Fe(tetpy)(OH),]+- polyelectrolyte systems also behave as environmental controllers of the uncatalysed oxidation of the L-(+)-axorbate anion.This effect is briefly discussed in terms of the role played by macroions in ionic reactions in solution. Active molecules anchored to polymeric matrices are useful systems for the investigation of catalytic phenomena. For example, transition metal ions or complex ions bound to macromolecules have been used as enzymatic models’ or active catalysts for a number of reactions in s ~ l u t i o n . ~ ~ We have recently reported that haemin-like 2,2’,2”,2”’-tetrapyridyliron(111) complex ions anchored to sodium poly(L-glutamate) (PLG) or poly(D-glutamate) (PDG) are efficient catalysts for the oxidation of L-( +)-ascorbic acid,6 according to the reaction O=C- C - OH O=C-c=o I II catalyst I t 0 C-OH + H202 ---+ 0 C=O + H,O ‘C! ‘C! I I CHOH CH ,OH I CHOH CH,OH Moreover, when the polypeptide matrices are in an a-helical conformation oxidation of the optically active substrate by the enantiomeric catalysts [Fe(tetpy)(OH),]+-PLG (FeL) and [Fe(tetpy)(OH),]+-PDG (FeD) becomes stereoselective.6 Since an ordered structure in PLG or PDG can be achieved by either lowering the pH or increasing the amount of bound counter-ions,’ the effect of both these parameters on the kinetics of the reaction has been extensively investigated.We report here the results obtained t Permanent address: Istituto Chimico, Universita di Napoli, 801 34 Napoli, Italy. 20732074 STEREOSPECIFIC EFFECTS I N ASYMMETRIC REACTIONS in a study of the peroxidase-like activitys of the aforementioned catalysts with L-( +)-ascorbic acid at fixed pH (7.0) and varying complex-to-polymer-residue ratio. The aims of the work were as follows: first to elucidate the mechanism of electron transfer catalysed by systems which also decompose hydrogen peroxide effi~iently,~ and secondly, to investigate the relationship between binding-induced conformational dissymmetry of the active sites and stereoselectivity in the catalysis.EXPERIMENTAL MATERIALS Poly(L-glutamic) acid and poly(D-glutamic) acid were purchased from Miles-Yeda (m.w. = 30000). They were converted into the sodium salts using 0.1 mol dm-, NaOH. Stock solutions were exhaustively dialysed against water to eliminate excess sodium ions and the materials recovered by freeze-drying.Concentrations of the polymers were determined by U.V. absorption at 200 nm (F = 5500). The pseudo-octahedral trans-iron(m) derivative is an 0x0-bridged dimeric compound in the solid state.loa In solution, at the pH values and concentrations with which we are concerned it is almost entirely in the mononuclear form [Fe(tetpy)(OH),]f.lo Concentrations of the iron(m) chelate were determined at 352 or 300 nm (E = 12 190 and 13 430, respectively; pH 7.0). Tris(hydroxymethy1)aminomethane (Sigma Chemicals) was used as buffer in the chloride form (tris buffer) at a concentration of 0.05 mol dm-, (pH 7.00 k 0.03). Interaction between polymer or complex ions and buffer was ruled out on the basis of preliminary optical measurements.Under the experimental conditions used the degree of association of [Fe(tetpy)(OH),]+ counter-ions by polypeptides is > 93 %, according to equilibrium dialysis experiments.' Both L-( + )-ascorbic acid (Merck) and stabilizer-free H,02 (C. Erba) were analytical-grade reagents. All measurements were performed on freshly prepared solutions, using doubly distilled water with conductivity .c 2 x f2-l cm-l (20 "C). METHODS A N D A P P A R A T U S Kinetic experiments were carried out spectrophotometrically under pseudo-first-order conditions, measuring the disappearance of ascorbic acid at 265 nm.ll A typical run con- sisted of adding hydrogen peroxide uia a microsyringe into the 1 cm optical cell contain- ing 2 cm3 of catalyst and substrate mixture, both systems being thermostatted at 25.9k0.1 OC.The experimental conditions normally used were [AH-] = 1 x lo-, mol dm-3, [H,O,] = 1 x mol dm-3 and [C]/[P] = 0.01-0.20, AH-, C and P denoting L-( +)-ascorbate anion, complex and polymer, respectively. [The pK,, value of ascorbic acid is reported to be 4.04 ( p = 0.1 rnol dmP3 KNO,, 25 "C)lZa or 4.03 ( p = 1 .O mol dm-, NaClO,, 20 "C)lzb]. Polymer concentration [PI is referred to the monomer, i.e. monomol drn-,. Measurements were also carried out with different initial concentrations of H202 and at another temperature (16.0 "C). Plots of log ( A , - A , ) as a function of time were linear over several half-lives. The observed pseudo-first-order rate constants k (s-l) were obtained from the slopes of the straight lines. At least four kinetic measurements were performed from each run to obtain consistent results.A slight oxidation of the substrate by hydrogen peroxide was detected in the presence of the sole polypeptide matrices. Its rate amounted to at most ca. 20% of the total rate. All first-order rate constants were properly corrected for this effect. Plots of k against complex concentration [C], at fixed complex- to-polymer-residue ratio, [C]/[P], always gave straight lines and the observed second-order rate constants kFeD and kFeL (dm3 mol-l s-l) were obtained from the slopes. The stereoselectivity factor of the catalysis is given by kPeD/kFeL. Absorption spectra were determined on a Berkman DBGT or Cary 219 spectrophotometer. The pH measurements were made with a Radiometer 26 pH-meter using standard semimicroelectrodes.mol dmP3, [C] = (0.5-7) xM. BARTER1 AND B. PISPISA 2075 RESULTS The pseudo-first-order rate constants k of the H,O, oxidation of L-( +)-ascorbate anion (pH 7.0, tris buffer 5 x lo-, mol dm-3) as a function of complex concentration [C], at fixed [C]/[P] values of 0.01 and 0.10 and at two temperatures, are reported in fig. 1 and 2. Within the range of complex concentration explored, the specific rates for the electron transfer increase with increasing [C]. The linear variation of rate with the concentration of (polymer-supported) complex ions at 25.9 and 16.0 OC indi- cates true catalytic behaviour for the iron(m) chelate. However, at [C]/[P] > ca. 0.01 extrapolation of the straight lines for [C] -+ 0 gives intercept values on the specific 60 40 I v1 m I 2 --- Y 2c 0 [Cl/10-5 rnol d ~ n - ~ FIG.1 .-Catalytic effect for the oxidation of L-( +)-axorbate anion in the presence of iron(m) complex ions bound to poly(L-glutamate) (empty symbols) or poly(D-glutamate) (full symbols), at [C]/[P] = 0.01 and two temperatures (25.9 OC, full line; 16.0 OC, broken line); pH 7.0 (tris buffer 5 x lo-* rnol dmP3). k = difference between first-order rate constants in the presence and absence of bound complex coun ter-ions. rate ordinate which are not zero (fig. 2). Thus k is a composite rate constant reflecting contributions from parallel pathways (k = k , + kcat [C]), of which one (kcat/dm3 mol-1 s-l) corresponds to the catalytic process and the other (k0/ss1) refers to an uncatalysed route for electron transfer, becoming negligible at very low complex-to-polymer-residue ratios.The slopes and intercepts of the plots of k against [C], corresponding to kcat and k,, respectively, are reported in table 1. (Subscripts FeD and FeL denote the enantiomeric catalytic systems whereas D and L denote the asymmetric polymeric matrices, see below.) Table 1 shows that the complex-to- polymer-residue ratio markedly affects both parallel pathways for the oxidation of2076 STEREOSPECIFIC EFFECTS IN ASYMMETRIC REACTIONS 100, c 0 0 1 2 3 4 5 6 7 8 [ C ] / 1 O-s rnol ~ i r n - ~ FIG. 2.-Catalytic effect for the oxidation of L-( + ) -axorbate anion in the presence of FeL (empty symbols) and FeD (full symbols) enantiomeric systems, at [C]/[P] = 0.10 and two temperatures: 25.9 (0, 0 ) and 16.0 O C (A, A); pH 7.0 (tris buffer 5 x lo-* mol dm-3).k = difference between first-order rate constants in the presence and absence of bound counter-ions. (ko, u, ko, L/s-l) OXIDATION OF L-( + )-AXORBATE ANION AT DIFFERENT [C]/[P] VALUES TABLE 1 .-RATE CONSTANTS FOR THE CATALYTIC (kFeD, kFeL/dm3 mOl-' S-l) AND NON-CATALYTIC T = 25.9 O C , pH 7.0, tris buffer 5 x lop2 mol dmP3, [H,O,], = 1 x [AH-], = 1 x mol dm-3, mol dm-3. ~~~~~ 0.0 1 3705 f 116 3608 k 1 19 1 .o ca. 0.6 ca. 0.6 1 .o 0.02 1144f62 1160f53 1 .o f 0 . 1 22.1 20.0 1 . 1 0.04 485.5k40.3 317.7k28.0 1.5f0.2 41 .O 32.7 1.2 0.06 409.3f39.8 208.7+ 17.9 2.0k0.3 46.8 34.9 1.3 0.10 442.3k42.9 153.6+ 12.6 2.9k0.5 59.6 38.7 1.5 0.20 414.0 k 33.4 106.4 f9.2 4.0 f 0.5 70.9 39.9 1.8 a Stereoselectivity factors of the catalytic reaction.ascorbate anion. At [C]/[P] = 0.01, not only is the rate of the non-catalytic process negligibly small compared with that of the catalytic one, but also the catalysis exhibits no stereoselectivity (kFeD = kFeL = 3.66 x lo3 dm3 mol-1 s-l). However, the observed second-order rate constant compares very favourably with those obtained using other Fe"' complexes as catalysts for the oxidation of the same substrate, although at lower pH.l3- 14a When [C]/[P] is increased the reaction becomes stereoselective, kFeD being definitely larger than kFeL. In addition, the higher is the complex-to-polymer ratio the greater is the stereoselectivity factor kFeD/kFeL. At the same time, the rate of the parallel uncatalysed reaction increases and also shows some stereospecificity.M.BARTER1 A N D B. PISPISA 2077 On the basis of these results, the following empirical rate expression may be formulated : rate = k , [AH-] + kcat [AH-] [C] (1) where AH- denotes the L-( +)-ascorbate anion1, and k , is an apparent first-order rate constant, which vanishes as [C]/[P] + 0.15 The oxidation of ascorbic acid was also studied as a function of the initial concentration of hydrogen peroxide in the range [H,O,],/[AH-], = 0.5-100. The second-order rate constants for the catalytic process were found to be independent of [H,O,] for all values of [C]/[P] (fig. 3 and table 2). This result agrees with that 40-----" 30 _ _ I I I I I I 0 1 2 3 4 5 [C] rnol dm-3 FIG. 3.-Catalytic effect for the oxidation of L-( +)-axorbate anion in the presence of FeL system at dif- ferent initial concentrations of hydrogen peroxide (see table 2): A, 1 .1 x 7, 1.1-x 0 , 5 . 4 x lop5 mol dm-3 [C]/[P] = 0.10, T = 25.9 OC, pH 7.0 (tris buffer 5 x lop2 mol drnp3). k = difference between first-order rate constants in the presence and absence of bound complex counter-ions. TABLE 2.-RATE CONSTANTS (dm3 m01-l S-') FOR THE CATALYTIC OXIDATION OF L-( +)-AXORBATE ANION AS A FUNCTION OF THE INITIAL CONCENTRATION OF HYDROGEN PEROXIDE (mol dm-3) AT TWO [c]/[P] VALUES T = 25.9 O C , pH 7.0, tris buffer 5 x lop2 mol dm-3, [AH-], = 1 x mol dm-3. 1.0 x 10-2 3705 3608 1.ox 10-2 442.3 153.6 1.2x 10-3 3922 3794 1 . 1 x 10-3 408.8 129.2 1.2 x 10-4 3372 347 1 1 . 1 x 10-4 405.1 156.1 4.9 x 10-5 366 1 3716 5.4x 10-5 471.0 140.7 av .3665 196 3647+ 121 av . 431.8 & 30.6 144.9 & 10.8 obtained by Grinstead on the H,O, oxidation of the same substrate on EDTA-Fe"' complex13 and is similar to those obtained by Martell et aZ.,14" who used 0, as an oxidant and diverse ferric chelates as catalysts. In contrast, the parallel uncatalysed reaction depends on [H,O,] (fig. 3 and 4), so that eqn (1) can be written as follows:2078 STEREOSPECIFIC EFFECTS IN ASYMMETRIC REACTIONS where ko, ?,,/dm3 mo1-1 s-' is still a complicated function of the complex to polymer ratio, vanishing as [C]/[P] -+ 0. From the slopes of the straight lines of fig. 4 one obtains ko, D, = 230.6 and ko, L, = 157.0 dm3 mol-l s-l. Using these par- ameters and kFer, = 431.8 dm3 mol-' s-l and kFeL = 144.9 dm3 mol-' s-' (table 2), the reaction velocities were calculated at [C]/[P] = 0.10 and [H,O,],/[AH-1, z 1.Comparison with the experimental rates shows a rather good agreement (table 3). [ H,O,], / 1 O-' rnol dm-3 FIG. 4.-Dependence of pseudo-first-order rate constants of the uncatalysed oxidation of L-( + )-axorbate anion on the initial concentration of H,O, {empty symbols, FeL; full symbols, FeD system ([C]/[P] = 0.10)). T = 25.9 OC, pH 7.0 (tris buffer 5 x lo-, mol dm-3); see text., TABLE 3.-cOMPARISON BETWEEN EXPERIMENTAL AND CALCULATED [EQN (2)] REACTION VELOCITIES ( lop6 mol dm-3 s-l) OF THE OXIDATION OF L-( +)-ASCORBATE ANION IN THE PRESENCE OF THE ENANTIOMERIC SYSTEMS F e D AND F e L , AT [c]/[P] = 0.10 T = 25.9 OC, pH 7.0, tris buffer 5 x mol dm-3, [C] = 5 x lop5 mol dm-3; see text.e x p t . calc. [AH-], / 1 0-4 mol dm-3 [H,O,],/[AH-], VF,D V F ~ L VF,D VF,L 0.6 0.9 2.43 1.15 2.04 0.94 0.6 1.8 3.35 1.21 2.82 1.47 1 . 1 0.5 3.26 2.02 3.74 1.73 1.1 1 .o 6.04 3.32 5.16 2.70 From the empirical rate law [eqn (2)] it appears that electron transfer in the non-catalytic reaction takes place between ascorbic acid and H,O,, whilst in the catalytic process it occurs between substrate and the transition metal complex, possibly within a Michaelis adduct. In the latter case, oxidation of both the lower valence metal chelate and A * - (or AH.) species by H,O, takes place in subsequent fast steps. A 0 - may also disproportionate very rapidly.16 Two further findings deserve a few comments. First, when the stereoselectivity factors kFeD/kFeL are plotted as a function of [C]/[P] ratio a trend similar to that followed by the a-helical fraction of the polymeric support (x,) on binding of [Fe(tetpy)(OH),]+ ions is observed (fig.5). The relationship between the parameter x, and the amount of bound complex counter-ions, under experimental conditions of binding equilibrium similar to those of this investigation, has been already reported by It was essentially based on a statistical treatment by a two-state model for the polypeptide17 and a preferential association of complex ions to the helical conformation of the polyelectrolyte, in agreement with the cooperative behaviour ofM. BARTER1 AND B. PISPISA 2079 the binding isotherm.' The results illustrated in fig. 5 clearly indicate that the conformational dissymmetry of the active sites is chiefly responsible for stereospecificity in the catalysis.Secondly, when the activation energies of the catalytic reaction are reported as a function of [C]/[P] an S-shaped curve is obtained, similar to that shown by the stereoselectivity factors (see above). Furthermore, at each value of [C]/[P] the activation energies for both enantiomeric catalysts are equal within experimental error (fig. 6). Finally, according to the theory of absolute reaction rates, the second-order 0 s, 1 1 I 0.1 0 0.20 [ C I / [ P I FIG. 5. I 2 3 ; $ -x . 3 5 5 I I I I 0 010 0.20 [CI /[PI FIG. 6. FIG. 5.-Variation of the stereoselectivity factor, kPeD/kFeL (O), of the catalytic oxidation of L-( + )-ascorbate anion at 25.9 OC and of the a-helical fraction of polypeptide matrix, x, (solid line), induced by the binding of 1rans-Fe"' molecules, as a function of [C]/[P] (degree of association > 93%); see text and ref.(7b). FIG. 6.-Variation of the activation energies of the catalytic oxidation of L-( + )-ascorbate anion, measured between 16 and 26 OC, as a function of [C]/[P]. The different symbols refer to the enantiomeric catalysts used; pH 7.0 (tris buffer 5 x lo-' mol dmP3). rate constants of the reaction between the FeD or FeL isomers of the asymmetric catalyst and the L isomer of the substrate are related by k,, = exp[-(AG,Z,-AG,f,)/RT] (3) k L L where AG,Z, and AG,Z, are the standard free energies of activation, and the notation DL and LL are used instead of FeD and FeL, respectively, to emphasize the diastereomeric character of the reacting systems under consideration.Assuming ideal behaviour, owing to the very low concentration of reagents, eqn (3) becomes K-LL where G,Z, and G,Z, are the standard free energies of the diastereomeric transition states. Eqn (4) was originally used by Prelog,18 who assumed that the difference (G,Z, - G,Z,) reflects the difference in steric hindrance between the two diastereomeric transition states, although this interpretation has recently been que~ti0ned.l~ In our case, the difference is linearly dependent on the a-helical fraction of the polymeric support (xa), as shown in fig. 7. This finding suggests that the Michaelis complex also2080 STEREOSPECIFIC EFFECTS I N ASYMMETRIC REACTIONS involves the chiral polymer residues surrounding the active sites.On increasing the amount of ordered polypeptide, the local stereochemistry would be then such that the LL diastereoisomer experiences larger steric hindrances than does the DL diastereo- isomer, which thus produces a less efficient electron-transfer process. The overall results lead to the conclusion that: (i) stereoselectivity is coupled with the amount of a-helix in the polymer support, which in turn depends on [C]/[P] (fig. 5 and 7 ) ; (ii) the effect is remarkable although it occurs at the expense of the catalytic efficiency, because the rate constants decrease by ca. one order of magnitude with respect to that of the non-stereospecific reaction (fig. 8); (iii) the diverse efficiency in the electron-transfer process between the FeD-L-( + )-ascorbate (DL) and FeL-L- ( + )-ascorbate (LL) diastereomeric systems is ascribable to a difference in entropies of activation, since at each [C]/[P] AH& = AHZ, within experimental error (fig.6); (iv) besides their catalytic role, the asymmetric [Fe(tetpy)(OH),]+-polyelectrolyte materials play the unusual role of environmental controller of the uncatalysed oxidation of L-(+)-ascorbate anion in terms of specific rate as well as stereospecific effects (table 1 and fig. 4), the latter being much smaller than those observed in the catalytic 3 FIG. 7.-Dependence of the difference between the standard free energies of the diastereomeric transition states (G,f, - GgL) on the a-helical fraction of polypeptide matrices; see text. DISCUSSION Evidence is produced that the H,O, oxidation of L-(+)-ascorbate anion in the presence of [Fe(tetpy)(OH),]+-polyglutamate systems follows parallel routes, one of which corresponds to an electron transfer process between substrate and central metal ion and the other refers to an uncatalysed pathway to products [eqn (2)].Furthermore, the observation that even the uncatalysed process exhibits stereoselective effects, although smaller than those of the catalytic reaction, indicates that complexes between counter-ions and polyelectrolytes play a role additional to that as catalyst. For sake of clarity, we discuss the data of the two parallel reactions separately. CATALYTIC OXIDATION The relevance of the results described here is that the enantiomeric systems [Fe(tetpy)(OH),]+-PLG (FeL) and [Fe(tetpy)(OH),]+-PDG (FeD) act as stereospecific catalysts in the oxidation of L-( +)-ascorbate anion only when the polypeptideM.BARTER1 A N D B. PISPISA 208 1 supports are in a-helical conformation, a situation that at fixed pH (7.0) can be achieved by increasing the amount of bound co~nter-ions.~ As [C]/[P] was increased both the stereoselectivity factor kFeD/kFeL and the activation energy for catalysis were found to increase following a sigmoid trend, clearly reflecting the helix-coil transition of the polymeric support (fig. 5 and 6). These observations indicate that at least two different factors are involved in determining the catalytic reaction rate; one is favoured by a decrease and the other by an increase in [C]/[P].It is thus reasonable that these factors are the amounts of random coil and a-helix in the polypeptide matrices. At very low [C]/[P] the (negatively charged) random-coil form of the support predominates. Under these conditions the anionic substrate probably interacts with the catalyst only through the bound, positively charged, iron(r1r) molecules, possibly forming a mixed-ligand chelate c ~ m p l e x . ~ ~ ~ 2o By analogy with simpler redox systems, electron transfer may be regarded as taking place between the two species directly. The direct attack mechanism of the substrate molecule on the central metal ion can account for the observed low activation energy of the reaction14121 (fig. 6), and is consistent with the idea that the precursor complex is conformationally mobile, owing to the size of the substituent in the ascorbate anion.According to this hypothesis, some of the conformers may favour the LL and others the DL reaction so that, on average, low stereoselectivity should be observed. This is indeed the case, as shown in table 1 and fig. 5. At high [C]/[P], the fixed charges in the polymeric matrix are effectively shielded and the support acquires rigidity, now being predominantly in the a-helical conf~rmation.~ Both these effects allow extensive interactions between substrate molecules and the catalyst, probably involving the chiral polymer residues surrounding the active sites. In this case a rigid intermediate would form because of hydrogen bonding between hydroxy groups of the ascorbate anion substituent and y-carboxy groups of the side-chains of the polypeptide.Conformational rigidity of the reacting species should enhance stereospecific effects in asymmetric processes,lg so that different rate constants for the oxidation of the optically active substrate on enantiomeric catalysts may be expected. This hypothesis is consistent with the experimental results illustrated in fig. 5 and 7. Furthermore, according to molecular models the stereochemistry of such an adduct allows electron transfer to take place only through the peripheral quaterpyridine ligand of the active sites, a mechanism for this type of reaction already found in metalloproteins and ferriporphyrin ~ y s t e m s . ~ ~ - ~ ~ This is probably a route with an activation energy higher than that of the direct attack mechanism (see above), as experimentally observed (fig.6). We may therefore conclude that the stereospecific electron transfer catalysis does not involve substitution on the iron centre but that the reaction proceeds by remote attack, making use of an electron-transfer site far from the central metal ion, possibly the edge of tetrapyridyl group. Further support for the above conclusions is provided by the following observation. The dependence of the second-order rate constants of the overall catalytic process (kea+) on [CJ/[P] (hereafter denoted by X ) was found to be satisfactorily described by where k, is the specific rate constant at some low value of X (around 0.01), k , is the specific rate constant at X z 0.20 (where kcat exhibits an almost asymptotic value and stereoselectivity approaches the maximum value), and b and b’ are adjustable constants.The data of fig. 8 were fitted with an average deviation of 6%, which2082 STEREOSPECIFIC EFFECTS IN ASYMMETRIC REACTIONS I 1 I 1 010 020 x = [Cl /[PI FIG. 8.-Experimental second-order rate constants for the oxidation of L-( + )-axorbate anion catalysed by FeL (empty symbols) and FeD (fun symbols) systems and calculated curves according to eqn (5) (solid lines). T = 25.9 OC, pH 7.0 (tns buffer 5 x !O-2 mol dm-3); see text. is of the same order of magnitude of the standard deviation of kcat, using k, = 4.02 x lo3 dm3 mol-l s-l for both catalysts and k,, = 100 and k,, = 420 dm3 mol-l s-l (solid lines in the fig. 8). [The constants b and b’ were - 108 and 1.20 x lo4, respectively.In fitting the data obtained on the FeD system, (bX+ b’X2) was replaced by (cX+ c’P + c’’X3) in order to get the same agreement, with c = - 59.4, c’ = 3.57 x lo3 and C” = 4.85 x lo5.] According to eqn (9, the reactant molecules can be divided into two groups which follow parallel pathways. One group would contain those particles which are oxidised with rate constant k,, irrespective of the catalyst used. This fraction, designed by (1 +bX+b’P)-l, is obviously the larger the lower is A’, so that it predominates when the support is highly disordered. The other group would contain those substrate molecules which are oxidised with rate constant k,, or k2L, depending on the catalyst employed. This fraction, designed by (bX+b’X2)- (1 + bX+ b’P)-l, predominates at high values of X , i.e.when the amount of a-helix in the polypeptide matrices is large. In agreement with the above-mentioned data, the stereoselectivity increases at the expense of catalytic efficiency because k,, and k,, are much smaller than k,. Furthermore, the maximum stereoselectivity is found to be k,,/k,, = 4.2 & 0.5. In conclusion, [Fe(tetpy)(OH),]+-polypeptide systems may be considered as a simple model of the allosteric effects attributed to enzymes. The progressive binding of complex counter-ions determines the coil-to-a-helix transition in the charged polymeric matrices. This in turn brings about a change in the mechanism of oxidation of the L-( +)-ascorbate anion, leading to marked stereospecific effects in the catalysis. At very low values of [C]/[P] the reaction is scarcely stereospecific and proceeds predominantly by a route involving direct attack of the ascorbate anion on the iron(II1).On the other hand, at high [C]/[P] the electron-transfer process is slower but exhibits stereoselectivity, which is greater the larger the amount of a-helix in the polypeptide matrices. In this case the precursor complex probably sees the substrate molecules bound to the chiral polymer residues surrounding the active sites and theM. BARTER1 AND B. PISPISA 2083 reaction proceeds predominantly by a remote pathway, through the quaterpyridine ligand. In the sense that the complex-polymer system contains both binding sites and catalytic sites it may be considered to be an enzymic model. NON-CATALYTIC OXIDATION The data reported in table 1 and fig, 4 clearly indicate that the uncatalysed H202 oxidation of ascorbate anion AH-+H,O, -+ A+OH-+H,O takes place in the domains of the complex-polyelectrolyte systems.The rate constant of the reaction varies with both [C]/[P] and the chirality of the polypeptide matrix used. Enhancement of the specific rates of ionic reactions carried out in the presence of macroions bearing opposite charges to those of the reacting particles has been widely reported. The phenomenon, known as ‘ polyelectrolyte catalysis’,26-28 is ascribable to electrostatic interactions between the species in solution. According to this view our results could be rationalised if the [Fe(tetpy)(OH,)]+-polyglutamate system were to behave as a cationic polyelectrolyte. This is in fact the case, since the y-carboxylate anions of the side-chains of the polymer, acting as unidentate ligands, are coordinated by the trans-iron(II1) derivative through an apical site.’ As a result, the negative charges on the polypeptide chains are effectively shielded by the bulky complex counter-ions, but a polycationic material simultaneously forms in situ, since the bound molecules preserve on coordination their formal univalent positive charge.This hypothesis can account for the dependence of k, on [C]/[P], owing to the fact that the charge density of the ‘ new ’ polyelectrolyte increases with increasing number of bound co~nter-ions,~~ as well as the finding that stereoselectivity of the uncatalysed reaction is lower than that of the true catalytic process (table 1).In the latter case the formation of a substratesatalyst complex allows closer contact between the asymmetric which is reflected in a much larger difference between the standard free energies of the diastereomeric transition states. The true rate constants of complicated reactions, such as those involving enzymes, can be evaluated only when the overall kinetic data are corrected from the contribution of the primary salt effect due to the presence of charged macromolecules.28b As far as we know, the results reported here represent the first example in which the contribution of the so-called ‘catalytic’ effects of the polyelectrolyte is clearly discriminated from that of the real active sites. Further experiments are, however, needed to clarify this phenomenon.We thank the Italian National Research Council (C.N.R.) for partial financial support. J. H. Wang and W. S. Brinigar, Proc. Natl Acad. Sci. USA, 1960, 46, 958; J. H. Wang, Ace. Chem. Res., 1970, 3, 90. A. Levitzki, I. Pecht and M. Anbar, Nature (London), 1965,207, 1386; J . Am. Chem. Soc., 1967,89, 1587. T. Nozawa, Y. Akimoto and M. Hatano, Makromol. Chem., 1972, 158,21; M. Hatano, T. Nozawa, S. Ikeda and T. Yamamoto, Mukromol. Chem., 1971, 141, 1 1 . G. Menecke and P. Reuter, Pure Appl. Chem., 1979, 51, 2313. R. H. Grubbs, Polymer-attached Homogeneous Catalysis, in Enzymic and Non-enzymic Catalysis, ed. P. Dunnill, A. Wiseman and N. Blakebrough (Ellis Horwood, Chichester, 1980), chap. 9. M. Barteri, B. Pispisa and M. V. Primiceri, Biopolymers, 1979, 18,3115; J .Znorg. Biochem., 1980, 12, 167. ’ (a) M. Branca, M. E. Marini and B. Pispisa, Biopolymers, 1976, 15, 2219; (b) M. Branca and B. Pispisa, J . Chem. Soc., Faraday Trans. I , 1977, 73, 213.2084 STEREOSPECIFIC EFFECTS I N ASYMMETRIC REACTIONS B. C. Saunders, Peroxidases and Catalases in Inorganic Biochemistry, ed. G. L. Eichhorn (Elsevier, Amsterdam, 1973), vol. 2, p. 988; G. A. Hamilton, Adu. Enzymol., 1969, 32, 55. M. Barteri, M. Farinella and B. Pispisa, Biopolymers, 1977, 16, 2569; J. Inorg. Nucl. Chem., 1978, 40, 1277; M. Barteri, M. Farinella, B. Pispisa and L. Splendorini, J. Chem. SOC., Faraday Trans. I , 1978, 74, 288. lo (a) M. Branca, B. Pispisa and C. Aurisicchio, J. Chem. SOC., Dalton Trans., 1976, 1543; (b) M.Cerdonio, F. Mogno, B. Pispisa and S. Vitale, Inorg. Chem., 1977, 16, 400. l 1 E. Racker, Biochim. Biophys. Acta, 1952, 9, 577. l 2 (a) M. M. T . Khan and A. E. Martell, J. Am. Chem. Soc., 1967,89,4176; (b) E. Pelizzetti, E. Mentasti and E. Pramauro, Inorg. Chem., 1976, 15, 1898. l 3 R. R. Grinstead, J . Am. Chem. Soc., 1960, 82, 3464. l4 (a) M. M. T. Khan and A. E. Martell, J . Am. Chem. Soc., 1967, 89, 7104; (b) A. E. Martell and M. M. T. Khan, Metal-ion Catalysis of Reactions of Molecular Oxygen, in Inorganic Biochemistry, ed. G. L. Eichhorn (Elsevier, Amsterdam, 1973), vol. 2, p. 654. l 5 K. J. Laidler and P. S. Bunting, The Chemical Kinetics of Enzyme Action (Clarendon Press, Oxford, 1973), chap. 11. l6 B. H. J. Bielski and H. W. kchter, Ann. N. Y. Acad. Sci., 1975, 258, 231. l7 J. Monod, J. Wyman and J. P. Changeux, J . Mol. Biol., 1965, 12, 88. l9 F. J. Hwang, L. C. De Bolt and H. Morawetz, J . Am. Chem. Soc., 1976, 98, 5890. 2o G. S. Laurence and K. J. Ellis, J. Chem. Soc., Dalton Trans., 1972, 1667. 21 N. Sutin, Oxidation-reduction in Coordination Compounds, in Inorganic Biochemistry, ed. G. L. Eich- 22 C. E. Castro and H. F. Davis, J . Am. Chem. Soc., 1969, 91, 5405. 23 J. K. Yandell, D. P. Fay and N. Sutin, J. Am. Chem. Soc., 1973, 95, 1131. 24 D. L. Toppen, J . Am. Chem. Soc., 1976,98,4023; F. L. Harris and D. L. Toppen, Inorg. Chem., 1978, 25 L. K. Hanson, C. K. Chang, M. S. Davis and J. Fajer, J. Am. Chem. Soc., 1981, 103, 663. 26 C. G. Overberger and J. C. Salamone, Ace. Chem. Res., 1969, 2, 217. 27 H. Morawetz, Ace. Chem. Res., 1970,3,354; Macromolecules in Solution (Wiley, New York, 2nd edn, 1975), chap. 9; J . Macromol. Sci., Macromol. Chem., 1979, A 13, 31 1. 2n (a) N. Ise, in Polyelectrolytes and their Applications, ed. A. Rembaum and E. Selegny (Reidel, Dordrecht, 1975), p. 71; J. Polym. Sci., Polym. Symp., 1978, 62, 205; N. Ise and T. Okubo, Macromolecules, 1978, 11, 439; (b) T. Okubo and N. Ise, Ado. Polym. Sci., 1977, 25, 136. 29 K. Mita, T. Okubo and N. Ise, J. Chem. Soc., Furaduy Trans. I , 1976,72,1033; A. Enokida, T. Okubo and N. Ise, Macromolecules, 1980, 13, 49. 30 C. Mavroyanis and M. J. Stephen, Mol. Phys., 1962, 5, 629. V. Prelog, He$. Chim. Acta, 1953, 36, 308. horn (Elsevier, Amsterdam, 1973), vol. 2, p. 61 1. 17, 74. (PAPER 1 /854)

ISSN:0300-9599    

DOI:10.1039/F19827802073

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1982, 78, 2085-2094 Coupling between pH-induced Conformational Phenomena and Stereospecific Effects in Electron-transfer Reactions BY MARIO BARTERI AND BASILIO PISPISAT Istituto di Chimica Fisica, Universita di Roma, 00185 Roma, Italy Received 2 1 st September, 198 1 The oxidation by H,O, of L-( + )-ascorbate anion in the presence of 2,2’,2”,2”’-tetrapyridineiron(111) complex ions anchored to poly(L-glutamate) (FeL) or poly(D-glutamate) (FeD) has been studied at a complex-to-polymer-residue ratio of 0.10 and in the pH range 6-8. Evidence is produced that the reaction is a composite process reflecting contributions from parallel routes; one of these corresponds to a catalytic, [H,O,]-independent pathway and the other refers to an uncatalysed electron-transfer process between ascorbate anion and hydrogen peroxide. Stereospecific effects in the catalysis are observed with decreasing pH, accompanied by an increase in the amount of a-helix in the polypeptide supports &). Thus at pH 7.8 v;, z 0.13), k,,, = 1382.3 f 1 13.2 and kFeL = 1034.4 f 79.3 dm3 mol-l s-’ and the activation energy is 3.8f0.3 kcal mol-’ with both enantiomeric catalysts, whereas at pH 6.3 (fh x 0.84), kFeD = 70.9k5.5 and kFeL = 13.6f 1.1 dm3 mo1-ls-l and the activation energy is 18.0f 1.3 kcal mol-I in both cases.The results indicate that stereoselectivity is an entropy-controlled phenomenon. The effect is probably caused by conformational rigidity of the precursor complex, which arises from interactions between the optically active substrate molecules and the chiral residues of the ordered polymer surrounding the active centres.Effects of the stereochemical features of the substrate-catalyst adduct on the mechanism of electron transfer are discussed. The evidence suggests that the asymmetric [Fe(tetpy)(OH),]+-polyelectrolyte systems also behave as environmental controllers of the uncatalysed oxidation of ascorbate anion. In a previous paperf we showed that haemin-like 2,2’,2”,2”‘-tetrapyridineiron(111) complex ions anchored to sodium poly(L-glutamate) (PLG) or poly(D-glutamate) (PDG) are efficient catalysts for the oxidation of L-( +)-ascorbic acid, according to the reaction O=C-C-OH o=c-c=o I II catalyst I I 0 C-OH + HzOz 0 C=O + HZO \CL \CL I I CHZOH CHOH I I CHOH CH,OH It has been also reported that the conformational features of the polypeptide supports are primarily responsible for stereospecific effects in the catalysis.The specific rate constant of the reaction catalysed by [Fe(tetpy)(OH),]+-PDG (FeD) system was found to be higher than that obtained using the enantiomeric [Fe(tetpy)(OH),]+-PLG (FeL) material only when the amount of a-helix in the polymeric matrices was 1arge.l Since ordered structure in PLG or PDG can be achieved either by lowering the pH or increasing the amount of bound counter-ions,2 the effect of both these parameters on the kinetics of the reaction was extensively investigated. In the previous paper1 the t Permanent address: Istituto Chimico, Universita di Napoli, 80134 Napoli, Italy. 20852086 STEREOSPECIFIC EFFECTS IN ELECTRON TRANSFER results obtained at fixed pH (7.0) and varying complex-to-polymer-residue ratio [C]/[P] were illustrated, whilst the kinetic data obtained at fixed [C]/[P] (0.10) and varying pH are reported here.Evidence is produced that there is a close relationship between stereochemical features of the precursor complex and both mechanism and stereoselectivity of the electron-transfer reaction. The results are fully consistent with those previously reported1 and are discussed in the light of some general considerations concerning the structural characteristics of the catalytic systems. EXPERIMENTAL MATERIALS Sodium poly(L-glutamate) and poly(D-glutamate) were obtained as previously described.l Polymer concentration [PI, determined by U.V. absorption at 200 nm (E = 5500), is referred to the monomeric unit, i.e.monomol dmP3. [Fe(tetpy)(OH),]+ complex ions were prepared as previously reported.'* Under the experimental conditions used [pH 6.33 and 6.54 (phosphate buffer 4 x lo-, mol dm-3), 7.02, 7.48 and 7.76 (tris buffer 5 x lo-, mol dm-3)], the degree of association of complex counter-ions by polypeptides is ca. 84,89,94,98 and loo%, respectively, according to equilibrium dialysis measurements.2* Both ascorbic acid (Merck) and stabilizer-free H,O, (Erba) were analytical-grade reagents. All measurements were performed on freshly prepared solutions, using doubly distilled water with a conductivity < 2 x f2-l cm-l (20 "C). METHODS AND APPARATUS Kinetic experiments were carried out under pseudo-first-order conditions, measuring the disappearance of ascorbate anion (AH-) at 265 nm.5 (The pK, values of ascorbic acid are 4.04 and 1 1.34 at 25 0C.6) They were performed at two temperatures (25.9 & 0.1 and 16.0 & 0.1 "C) as already described.' At each pH the range of catalyst concentration explored was 0.5 < [C]/10-5 mol dm-3 < 5 ([C]/[P] = 0.10).In all cases, plots of log(A, -Am) as a function of time were linear over several half-lives. The observed pseudo-first-order rate constants k (s-l) were obtained from the slopes. Four kinetic measurements were performed for each run to obtain consistency in results. A slight oxidation of ascorbate anion by H,O, was detected in the presence of the sole polypeptide matrices,'?' its rate depending on both [PI and pH.All first-order rate constants were properly corrected for this effect. At each pH investigated, plots of k as a function of [C] always gave straight lines and the second-order rate constants kFeD and kFeL (dm3 mol-l s-l) were obtained from the slopes. Stereoselectivity of the catalysis is expressed as kFeD/kFeL. Absorption spectra were recorded on a Beckman DBGT or a Cary 219 spectrophotometer. Circular dichroism (c.d.) spectra were determined on a Cary 61 dichrograph, with appropriate quartz cells. The pH values were measured by a Radiometer 26 pH-meter using standard semimicroelectrodes. RESULTS The pseudo-first-order rate constants k of the oxidation reaction of L-( + )-ascorbate anion as a function of complex concentration [C], at fixed [C]/[P] of 0.10 and at pH 6.5 and 7.5, are reported in fig.1. As already observed for the same type of plot with the data obtained at fixed pH (7.0) and diverse [C]/[P] values,l the specific rate for electron transfer linearly increases with increasing [C]. This trend indicates true catalytic behaviour for the iron(m) chelate. Furthermore, within the whole range of pH explored (6.3-7.8) the straight lines exhibit intercept values which differ significantly from zero, i.e. k = k, + kcat [C]. This implies that the reaction is a composite process reflecting contribution from parallel pathways, one of which refers to the catalytic process (kcat/dm3 mol-l s-l) and the other corresponds to an uncatalysed route to dehydroascorbic acid (k0/s-l). The values of kcat and k, (25.9OC) at different pHM.BARTER1 A N D B. P I S P I S A 2087 100 - I m m 2 c 0 1 2 3 4 5 6 [C]/10-5 mol dm-3 FIG. 1 .-Catalytic effect of the oxidation of L-ascorbate anion in the presence of iron(w) complex ions bound to poly(L-glutamate) (empty symbols) or poly(D-glutamate) (filled symbols) at [C]/[P] = 0.10 and pH 6.5 (broken lines) and 7.5 (full lines), in 0.04 mol dmP3 phosphate buffer or 0.05 mol dm-3 tris buffer, respectively. T = 25.9 OC, [H,O,],/[AH-1, = 100. k is the difference between first-order rate constants in the presence and absence of bound complex ions. CATALYTIC (ko,D AND ko,L/s-l) OXIDATION OF L-ASCORBATE ANION AT DIFFERENT pH VALUES TABLE l.-RATE CONSTANTS FOR THE CATALYTIC (kFeD AND kFeL/dm3 mOl-' S-l) AND NON- T = 25.9 OC, [C]/[P] = 0.10, p = 0.04 dm3 mol-', [H,O,],/[AH-1, z 100.~~ PH fhU k F e D k F e L kFeD/kFeLb k O , D/ k O , L / k ~ , D / ~ o , L 7.76 0.13 1382.3k113.2 1034.4f79.3 1.3f0.2 110.3 106.1 1 .o 7.48 0.19 988.3f101.6 626.0k56.1 1.6f0.2 87.3 80.0 1.1 7.02 0.45 442.3f42.9 153.6k12.6 2.9f0.5 59.6 38.7 1.5 6.54 0.75 143.0f 13.1 40.2 k 3.7 3.7 k0.5 39.9 21.4 1.9 6.33 0.84 70.9f5.5 13.6 k 1.1 5.2 & 0.6 34.8 17.1 2.0 a-Helical fraction of polypeptide matrices at [C]/[PJ = 0.10 (see text); stereoselectivity factors of the catalytic reaction. values are reported in table 1, where subscripts FeD and FeL denote the enantiomeric catalysts, whereas D and L denote the asymmetric polymeric supports. In the same table the fraction of a-helix in the polypeptide matrices ( f h ) , at [C]/[P] = 0.10, is also shown.It was evaluated by the conventional expression A E o ~ s = f h AEh + ( -fh) AEc where AE = E~ - eR is the differential circular dichroism absorption (dm3 mol-l cm-l), using At+, = - 9.2 and A E ~ = 1.4, which represent the estimated values of ( E ~ - E ~ ) at2088 STEREOSPECIFIC EFFECTS IN ELECTRON TRANSFER 220 nm of helical and coil polypeptides at pH ca. 5 and ca. 9, respectively, in the presence of the complex ions.' The values of are those reported in fig. 2, where the change in ellipticity at 220 nm of poly(L-glumatate) solutions in the absence and presence of the iron(m) chelate molecules, at [C]/[P] = 0.10, is plotted as a function of pH. As already seen,' the binding of complex counter-ions determines a stabilization of the a-helical structure in the polypeptide in terms of an upward shift in the pH region of the conformational transition.Furthermore, the curve levels off at a value of ellipticity lower than that of the helical-polypeptide-free-complex solutions because of an induced c.d. band of opposite sign in the bound complex, within the same frequency 0 -10 I I I I 1 I 4 5 6 7 8 9 PH FIG. 2.-Variation of ellipticity at 220 nm of PLG solutions, in the absence (A) and presence (0) of chelate ions ([C]/[P] = 0.10) as a function of pH. Inspection of table 1 shows that the hydrogen-ion concentration markedly affects both parallel reactions in terms of specific rates as well as stereospecific effects. At around pH 8 a very small stereoselectivity is observed, despite the fact that the con- figurational dissymmetry2y8 of the active sites is relatively high, as indicated by the induced ellipticity of, say, the FeL system at 287 nm, where the chiral polymeric sup- port does not absorb ([el = - 32.0 x lo3 deg cm2 dmol-l, [C]/[P] = 0.10).With in- creasing [H+] the catalysis becomes stereoselective, kFeD being definitely higher than kFeL. In addition, the higher the hydrogen-ion concentration ( i e . the larger the amount of a-helix in the polypeptide matrices), the greater the stereoselectivity, which is seen to occur at the expense of catalytic efficiency because the rate constants decrease by ca. two orders of magnitude with respect to that of the non-stereospecific process. At the same time, the parallel uncatalysed reaction also shows some stereospecificity ; however it is much smaller than that of the catalytic process.When the stereoselectivity factors kFeD/kFeL are plotted as a function of pH, a trendM. BARTER1 A N D B. PISPISA 2089 similar to that followed by the a-helical fraction of the polypeptide supports (fh) is observed (fig. 3). This finding clearly indicates that the conformational dissymmetry2 of the active sites is primarily responsible for stereospecific effects in the catalysis. Similar conclusions were reached when studying the same reaction at pH 7.0 and varying [C]/[P].l In fact, an increase in the amount of bound counter-ions at fixed pH (7.0) has a similar (although less pronounced) effect on the coil-to-a-helix transition of the chiral polymeric supports to that of decreasing the pH at fixed [C]/[P] = 0.1 .o fh 0.5 0 6 5 i 3 * 2 1 PH FIG. 3.-Vanation of a-helical fraction in PLG, fh (a), and of stereoselectivity factor, kFeD/kFeL (vertical bars), as a function of pH. The parameterf, was evaiuated by c.d. measurements at 220 nm of complex-PLG solutions at [CJ/[P] = 0.10 (see text). TABLE 2.-RATE CONSTANTS (dm3 m0l-l S-l) FOR THE CATALYTIC OXIDATION OF L-ASCORBATE ANION AS A FUNCTION OF THE INITIAL CONCENTRATION OF HYDROGEN PEROXIDE AT TWO pH VALUES 7'= 25.9 O C , [C]/[P] = 0.10, [AH-] z 1 x mol dmP3. pH 7.0 pH 6.3 [H202],/mol dm-3 k F e o b e t . k F e D k F e L 1.ox 10-2 442.3 153.6 70.9 13.6 1.1 x 10-3 408.8 129.2 1.1 x 10-4 405.1 156.1 5 9 . 2 10.6 5.4x 10-5 47 1 .O 140.7 84.0 16.1 - - av . 431.8+30.6 144.9+ 10.8 71.4+ 10.1 13.4 2.0 The oxidation of L-ascorbate anion was also studied as a function of the initial concentration of hydrogen peroxide, in the range 0.5 < [H,O,],/[AH-], < 100.The results obtained at pH 6.3, which are entirely consistent with those observed at pH 7,' show that with increasing [H202] the slopes of the straight lines remain practically constant within experimental error (table 2), whilst the intercept values increase. This implies that the second-order rate constants of the catalytic oxidation of ascorbate anion is independent of [H202], at variance with the rate of the uncatalysed reaction. The dependence of k , on [H202],, at [H202],/[AH-], z 1 and pH 6.3, is illustrated in fig. 4. 68 FAR 12090 STEREOSPECIFIC EFFECTS I N ELECTRON TRANSFER On the basis of these results, the following empirical rate expression may be (1) where the second-order rate constant k,, app (dm3 mol-l s-l), which can be evaluated from the slope of the straight lines of plots of k , against [H2O2l0 (see fig.4), is a complicated function of [C]/[P]l as well as of [H+], as illustrated in table 3. formulated : d[AH-] dt --- - ko, app [AH-] [H202l+ kcat [AH-] [CI 30t ---- 10 0 5 10 [HzOzlo/ 10- mol dm-3 FIG. 4.-Dependence of pseudo-first-order rate constant k, of the uncatalysed oxidation of L-ascorbate anion on the initial concentration of hydrogen peroxide, in the presence of FeL (0) or FeD (A) systems ([C]/[P] = 0.10). [H,O,],/[AH-1, z 1, T = 25.9 O C , pH 6.3 (0.04 mol dmP3 phosphate buffer). TABLE 3 .-APPARENT SECOND-ORDER RATE CONSTANTS (dm3 m0l-l S-l) FOR THE UNCATALYSED OXIDATION OF L-ASCORBATE ANION.T = 25.9 O C , [H,O,],/[AH-1, z 1 . PH 6.3 0.10 178.3 93.3 7.0 0.10 230.6 157.0 7.0 0.20 287.8 156.0 Eqn (l), which is formally similar to that reported in our previous paper,l suggests that the rate-determining step of the uncatalysed reaction involves one molecule of substrate per molecule of hydrogen peroxide (which does not necessarily imply a two-electron step),g whereas that of the catalytic process occurs between one molecule of complex ion and one molecule of substrate. In the latter case, oxidation of both the lower-valence metal chelate and the ascorbate radical takes place in subsequent fast steps. A '- may also disproportionate very rapidly.1° These results are reminiscent of those obtained by other authors for the same reaction in acid solutions, catalysed by diverse iron(m) l1 Two further findings are worth mentioning.First, the activation energy of the catalysis is seen: (i) to vary as a function of pH following a sigmoid trend, like that shown by the stereoselectivity factor under the same experimental conditions, and (ii) to exhibit equal values with both enantiomeric catalysts, within experimental error (fig. 5). According to circular dichroism data, the helix content of poly(L-glutamate) in FelI1 chelate solutions ([C]/[P] = 0.10) at 16 O C does not practically differ from that evaluated at 26 OC, other experimental conditions being equa1.12 The variation of activation energies as a function offh may therefore be safely ascribed to someM.BARTER1 AND B. PISPISA 209 1 1250 1 0 5 6 7 0 PH FIG. 5 . 0 0.5 1.0 f h FIG. 6. FIG. 5.-Variation of the activation energies of the catalytic oxidation of L-ascorbate anion with pH. The different symbols refer to the enantiomeric catalysts ([C]/[P] = 0.10). FIG. 6.-Dependence of the difference in standard free energies of the diastereomeric transitions states on the amount of a-helix in the polypeptide matrices (see text); T = 25.9 O C . change in the electron-transfer mechanism with conformational features of the polypeptide matrices. In addition, stereoselectivi ty appears to be a process controlled by activation entropy, as already found in the study of the same reaction at fixed pH (7.0) and varying [C]/[P].l Secondly, the difference in the standard free energies of the diastereomeric transition states (G,Z, - GiL) is observed to increase with increasing a-helical fraction of the supports (fh), as shown in fig.6. The difference (GZ - GgL) was evaluated by the expression1 LL which relates the second-order rate constants of the electron-transfer reaction in FeD-L-ascorbate (DL) and FeL-L-ascorbate (LL) diastereomeric systems, on the assumption of ideal behaviour owing to the very low concentration of the reacting species. The data of fig. 6 indicate that stereoselectivity is closely connected with the ordered structure of the support in the sense that on increasing the amount of helical polypeptide the structural characteristics of the catalyst probably allow the chiral polymer residues to participate in the formation of the precursor complex.The local stereochemistry would be then such that the LL diastereoisomer experiences larger steric hindrances than does the DL diastereoisomer, which is responsible for a less efficient electron-transfer pro~ess.~’ 13* l4 Finally, according to the data reported in tables 1 and 3, the overall rate of the reaction varies as a function of pH in a rather complicated way. For sake of simplicity, we neglect here the uncatalysed process and focus attention on the variation of kcat with pH. 68-22092 STEREOSPECIFIC EFFECTS I N ELECTRON TRANSFER In principle, two concurrent factors may account for the observed dependence of k,,, on [H+], i.e. (i) the competition between non-stereospecific and stereospecific pathways in the overall catalytic cycle because two routes have different rate constants,l and (ii) the diverse reactivities of the ionic species of ascorbic acid toward the catalysts employed.6 By analogy, with the expression describing the dependence of kcat on [C]/[P],l the first effect is empirically describable as follows: kcst = k’f([H+I) + k”( 1 -f([H+I)) (3) where k’ is the specific rate constant at some low value of [H+] (say, ca.1 x 1 O-* mol dm-3), where the non-stereospecific catalysis predominates, k” is the rate constant at [H+] z 1 x lop6 mol dm-3, where kcst exhibits an asymptotic value and stereoselectivity approaches the maximum value, and f([H+]) is a function of hydrogen-ion concentration such that it tends to 1 or 0, respectively as [H+] approaches the former or latter value.On the other hand, since ascorbic acid (AH,) undergoes the following protolytic K , K2 equilibria : (4) and both the neutral and ionic forms can be oxidized, in principle, by the iron(rI1) derivative, the rate law of each catalytic pathway may be expressed as: AH, e AH- e A,- catalytic rate = [C] (k,[AH2] + k2[AH-] + k,[A2-]) ( 5 ) where k,, k, and k, are constants for the catalytic effect of the metal chelate on the un-ionised, monoionic and bi-ionic forms of ascorbic acid, respectively, on the assumption that the three substrate species react independently with the catalyst. Taking into account the equilibium constants in equilibrium (4), mass balance allows us to write eqn ( 5 ) as follows:1s [SI [Cl k, K, K, + k, K,[H+] + [H+I2 catalytic rate = K, K2 + K,[H+] + [€I+], where [S] denotes the total substrate concentration ([S] = [AH,] + [AH-] + [A2-]).Even simplifying eqn (6), by assuming that the concentration of AH, is negligibly small as compared with the others in the pH region l6 the concomitant occurrence of the effects described by eqn ( 3 ) and (6) emphasizes the difficulty in interpreting the change of kcat with pH. Additional complications arise not only from failing to have pK values of the substrate in the domains of the charged polymeric catalysts (owing to the fact that protolytic equilibria of ionic species in proximity to macroions are perturbed by the high electrostatic potential of the p o l y e l e ~ t r o l y t e ) ~ ~ - ~ ~ but also from having insufficient data within the range of pH explored.All these features make it possible, at present, to draw only qualitative conclusions on the dependence of the rate of oxidation of ascorbic acid on pH. The reaction is sensitive to [H+] in terms of rate constant as well as stereoselectivity, the former increasing and the latter decreasing as a function of pH. In addition, both effects were found to be more pronounced in the catalytic process than in the parallel uncatalysed reaction. Further study in this area is clearly required. DISCUSSION Evidence is produced that stereospecific effects in the catalytic oxidation of ascorbate anion are coupled with the amount of a-helix in the polypeptide matrices, which in turn depends on pH. At the same time, the activation energy is shown toM. BARTER1 A N D B.PISPISA 2093 increase with a sigmoid trend and to exhibit equal values with both enantiomeric catalysts. These results are suggestive of different routes to products, depending on the structural characteristics of the catalytic systems used. At or above pH 7, the negatively charged polymeric support repels the anionic substrate molecules, which may thus interact with the catalyst only through the bound, positively charged metal chelate. Under these conditions, the accessible axial position of the ligated complex ions would provide the electron-acceptor site,6- 16-21 and the transfer of one electron may occur directly between the two species through the bridge. A direct-attack mechanism by the substrate on the central metal ion has been proposed by us1 under conditions where similar structural features of the catalytic systems were matched, i.e.at pH 7 and very low [C]/[P]. Accordingly, this mechanism is expected to be coupled with a low stereoselectivity in that the degrees of rotational freedom of the diastereo- merically related transition states should be high, owing to the size of substituent in the substrate and the lack of assistance from the polypeptide in the formation of the intermediate. This is indeed the case, as experimentally observed (table 1 and In addition, a direct electron-transfer process between ascorbate and iron(1rr) ion within a mixed-ligand metal chelate should require a low activation 2 2 as observed in neutral or weakly alkaline solution (fig. 5). In contrast, the catalytic reaction in acid solution shows both high stereoselectivity and high activation energy.This suggests a different electron-transfer route to the products, probably owing to the diverse stereochemistry of the precursor complex. With increasing [H+] the negative charge density of the polypeptide matrix decreases, so that extensive interactions between substrate molecules and the ordered polymeric catalyst are expected to occur. For instance, hydrogen-bonding interactions between hydroxy groups of the substituent in ascorbate and y-carboxylate groups in the side-chains of the polymer may take place, leading to a much more rigid Michaelis adduct than that seen above. Such a loss of conformational mobility should enhance the difference in steric hindrances between the two diastereomeric transition states, which must be reflected in an increase of stereospecific effects.13 This is indeed the case, as shown in table 1 and fig.3 and 6. Furthermore, molecular models suggest that the stereochemistry of this type of adduct allows electron transfer from ascorbate to iron(m) ion to proceed only by a remote-attack mechanism, possibly through the peripheral quaterpyridine ligand of the active sites. This hypothesis, which is reminiscent of that proposed for the reduction of ferriporphyrins by ascorbic acid16 and for a number of redox reactions between metalloproteins and various reduc- tants,22 25 may account for the relatively high activation energy in acid solution. We may therefore conclude that there is a close relationship between stereochemical characteristics of the precursor complex and both mechanism and kinetic stereo- selectivity of the electron-transfer catalysis. All these features depend markedly on pH because it controls the structure of the chiral polymeric support, which appears to be primarily responsible for the observed phenomena.As far as the parallel uncatalysed oxidation of t-ascorbate anion is concerned, it was shown that with decreasing hydrogen-ion concentration the specific rate increases (tables 1 and 3 ) . By analogy with the catalytic reaction, this trend may be explained in terms of different rate constants for the oxidation of the ionic species of ascorbic acid (see above). Nevertheless, we are inclined to think that an additional, more subtle, factor connects the rate constant of the uncatalysed process with pH.It has been reportedly 2 * * that coordination of [Fe(tetpy)(OH),]+ ions by y-carboxylate groups of the side-chains of the polymer leads to the formation of positive ionic sites on the polypeptide because the ligated molecules keep the univalent positive charge of the free fig. 3 ) .2094 STEREOSPECIFIC EFFECTS I N ELECTRON TRANSFER chelate ions. Coulombic interactions between ascorbate molecules and the positively charged polymeric material are thus expected to take place, the effect being more pronounced the higher the pH because the degree of counter-ion association As a result, the specific rate of the uncatalysed oxidation of ascorbate anion is enhanced (table l), in agreement with the fact that the rate of ionic reactions is raised by polyions bearing fixed charges opposite in sign to those of the reacting 27 This hypothesis may also account for the fact that stereospecific effects in the uncatalysed reaction are definitely smaller than those in the true catalytic process.The formation of a Michaelis adduct as a preliminary step in the catalysis implies a more intimate contact between the asymmetric partners than that which one would predict on the basis of long-range Coulombic interactions alone. This should be reflected in a much larger difference between the standard free energies of the diastereomeric transition states, as experimentally observed (tables 1 and 3). In conclusion, although at present the pH dependence of the reaction can be described only qualitatively because of insufficient experimental data and quantitative knowledge of pK values of ascorbic acid in complex-polyelectrolyte solutions, the overall results are entirely consistent with a picture in which the complex-polypeptide system behaves not only as an efficient and stereospecific electron-transfer catalyst but also as an environmental controller of the uncatalysed oxidation of ascorbate anion.We thank the Italian National Research Council (C.N.R.) for partial financial support. M. Barteri and B. Pispisa, J. Chem. SOC., Faraday Trans. I , 1982, 78, in press. M. Branca and B. Pispisa, J. Chem. SOC., Faraday Trans. I , 1977, 73, 213. M. Branca, B. Pispisa and C. Aurisicchio, J. Chem. Soc., Dalton Trans., 1976, 1543; M. Cerdonio, F. Mogno, B. Pispisa and S.Vitale, Znorg. Chem., 1977, 16, 400. M. Barteri, B. Pispisa and M. V. Primiceri, Biopolymers, 1979, 18, 31 15. E. Racker, Biochim. Biophys. Acta, 1952, 9, 577. M. M. T. Khan and A. E. Martell, J. Am. Chem. SOC., 1967, 89, 7104. M. Branca, M. E. Marini and B. Pispisa, Biopolymers, 1976, 15, 2219. W. Weis, Ann. N. Y. Acad. Sci., 1975, 258, 190. lo B. H. J. Bielski and H. W. Richter, Ann. N . Y. Acad. Sci., 1975, 258, 231. l1 R. R. Grinstead, J. Am. Chem. SOC., 1960, 82, 3464. l 2 W. L. Mattice, R. W. McCord and P. M. Shippey, Biopolymers, 1979, 18, 723. l3 V. Prelog, Helv. Chim. Acta, 1953, 36, 308. l4 F. J . Hwang, L. D. De Bolt and H. Morawetz, J. Am. Chem. Soc., 1976, 98, 5890. l5 L. Michaelis and H. Davidson, Biochem. Z., 191 I , 35, 386; A. Cornish-Bowden, Fundamentals of Enzyme Kinetics (Butterworths, London, 1979), chap. 7. l6 F. L. Harris and D. L. Toppen, Znorg. Chem., 1978, 17, 74. l7 V. Crescenzi, F. Delben, F. Quadrifoglio and D. Dolar, J. Phys. Chem., 1973, 77, 539. ’ M. Barteri, B. Pispisa and M. V. Primiceri, J. Znorg. Biochem., 1980, 12, 167. M. Barteri, M. Farinella, B. Pispisa and L. Splendorini, J. Chem. Soc., Faraday Trans. 1, 1978, 74, 288. l9 M. Branca, P. Checconi and B. Pispisa, J. Chem. Soc., Dalton Trans., 1976, 481. 2o Y. Yano, S. Kawada and W. Tagaki, Bull Chem. SOC. Jpn, 1981, 54,493. 21 C. S. Laurence and K. J. Ellis, J. Chem. SOC., Dalton Trans., 1972, 1667. 22 N. Sutin, Oxidation-Reduction in Coordination Compounds in Inorganic Biochemistry, ed. G. L. 23 C. E. Castro and H. F. Davis, J. Am. Chem. SOC., 1969, 91, 5405. 24 N. Sutin and A. Forman, J. Am. Chem. SOC., 1971,93, 5274; J. K. Yandell, D. P. Fay and N. Sutin, 25 L. K. Hanson, C. K. Chang, M. S. Davis and J. Fajer, J. Am. Chem. SOC., 1981, 103, 663. 26 H. Morawetz, Macromolecules in Solution (Wiley, New York, 2nd edn, 1975), chap. 9; Pure Appl. 27 T. Okubo and N. Ise, Adv. Polym. Sci., 1977, 25, 136; N. Ise, J. Polym. Sci., Polym. Symp., 1978, Eichhorn (Elsevier, Amsterdam, 1973), vol. 2, p. 61 1 . J. Am. Chem. Soc., 1973, 95, 1131. Chem., 1979, 51, 2307. 62, 205; A. Enokida, T. Okubo and N. Ise, Macromolecules, 1980, 13, 49. (PAPER 1 / 1472)

ISSN:0300-9599    

DOI:10.1039/F19827802085

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1 , 1982, 78, 2095-2100 Heats of Dilution of Aqueous Solutions of Sodium Tetraphenylboron at 25 OC BY THELMA M. HERRINGTON* AND ELSPETH L. MOLE Department of Chemistry, University of Reading, Reading RG6 2AD Received 8th June, 1981 The heats of dilution of aqueous solutions of sodium tetraphenylboron are determined at 25 "C. The microcalorimeter was first tested by determining the heat of dilution of sodium chloride. The temperature dependence of the solute-solute attraction, after subtracting the Debye-Hiickel contribution, was compared with that of other electrolytes and non-electrolytes. In studies of the physico-chemical behaviour of the tetraphenylborate ion, it has been found that very different behaviour is observed compared with other quaternary ions of similar size (e.g.Bu,N+, Ph,P+); these include enthalpies of transfer,' conductivity measurements2 and n.m.r. ~tudies.~ The apparent molar volumes of aqueous sodium tetraphenylboron solutions have been determined between 0 and 60 OC by Millero4 and also by Jolicoeur and Philip;5 the latter authors also found that the apparent molar heat capacity of sodium tetraphenylboron showed a marked temperature dependence compared with Ph,P+ and Bu,N+, which are only weakly temperature dependent. The observed temperature dependence, however, was less than that found earlier by Subramanian and Ahluwalia.6 The relative magnitudes of solute-solute and solute-solvent interactions have been studied for non-electrolytes using rigorous statistical mechanical theorie~.~ In dilute solution up to 0.1 mol kg-l, if the Debye-Huckel electrostatic contribution is subtracted from a thermodynamic parameter, e.g.In y, then the remainder is linear in the molality as it would be for a non-electrolyte. It was decided to determine the heats of dilution and to compare these data with our previous treatment of non-electrolytes. * EXPERIMENTAL MATERIALS The sodium tetraphenylboron was purified by the method of Skinner and FUOSS;~ finally it was recrystallized three times from conductivity water, dried and stored under argon. All solutions were prepared using once distilled but previously deionized water, which had a conductance of < 1 x S cm-l. Solutions were made up by weight and buoyancy corrections applied to give a precision of & 0.1 mg.T H E CALORIMETER The microcalorimeter and the method of operation have been described previously.8 The aqueous solution of sodium tetraphenylboron, NaBPh,, was contained in a sealed pipette and diluted with water in the outer vessel. The temperature rise was monitored by a thermistor a.c. bridge. Heat was evolved on dilution and compensated for by the cooling effect of frigistors. Once experience had been gained of the heat liberated on dilution, compensation was arranged so as to minimise cooling corrections. The heat of opening of the pipette valves was found to 20952096 HEATS OF DILUTION OF AQUEOUS SOLUTIONS be negligible. The accuracy of the calorimeter was tested by determining the enthalpies of dilution of sodium chloride solutions and comparing the results with those of Gulbransen and Robin~on.~ Unfortunately, the number of dilutions that could be performed with sodium tetraphenylboron was limited by its solubility.RESULTS The results of both the sodium chloride and the sodium tetraphenylboron dilutions were analysed by first subtracting the Debye-Huckel contributions. The molar enthalpy of dilution is equal to the difference in the apparent molar enthalpy of the solution on dilution from an initial molality, mI, to a final molality, mF. Thus if H , is the value of the heat function for a quantity of solution containing one mole of solute at a molality m, then (1) AH$' = HF , -HI , = A4H2 = A4L2 (2) d d T where Hm = -2RT2-[lny+(l -#)I for a 1 : 1 electrolyte.* Let us define a standard enthalpy function by As the solutions are dilute (0.1 mol kg-l), the Debye-HuckellO values for ySt and bst (4) are used, then In yst = - arni/(l+ mi) where and a = (2nNp,)4 (e2/411&k~)t. (9) From Sylvester and PitzeP C is 1.947 kJ mol-l.The experimental data for sodium chloride were compared with those obtained by Gulbransen and Robin~on.~ Their results in the form of A H 2 l - A H z were plotted against Am; this plot is linear up to Am = 0.1 mol kg-l, so that a smoothed value of AH$," could be obtained to compare with our own values. Our experimental results are compared with those of Gulbransen and Robinson in table 1; the agreement is satisfactory . The heat of dilution of sodium tetraphenylboron solutions is much greater than that of sodium chloride between the same two molalities; heat is evolved on dilution.The results for sodium tetraphenylboron are given in table 2, as also are values for A H 2 and AH$' -AH:. Let the activity coefficient of the solute in this molality range be represented by In y = In ySt + mm then, if the plot of (AHZl-AHg) against Am is a straight line, the slope is -RT2(8m/aT). For sodium chloride the value obtained for au/aT is 2.04 x mol-1 kg K-l. For sodium tetraphenylboron the plot is also a straight line * See Glossary of Symbols on p. 2099.T. M. H E R R I N G T O N AND E. L. MOLE 2097 TABLE HEATS OF DILUTION OF AQUEOUS SODIUM CHLORIDE SOLUTIONS AT 25 O C 0.099 82 0.004 62 272 270 10 0.100 25 0.005 37 262 252k 15 0.101 38 0.005 11 267 270 & 8.2 0.100 03 0.004 33 276 264+ 10 0.099 25 0.006 38 247 243 & 7.2 From ref. (9).TABLE 2.-HEATS OF DILUTION OF AQUEOUS SODIUM TETRAPHENYLBORON SOLUTIONS AT 25 OC ~~ -(AH$' - AH$) rn'/mol kg-l rnF/moi kg-' -AP;/J mol-' -AHE1/J mol-1 /J mol-l 0.024 76 0.001 26 200 1464&33 1264f33 0.025 20 0.001 27 209 1477Ifr 13 1268f 13 0.025 14 0.001 24 209 1515+25 1306 & 25 0.020 03 0.000 96 190 1188f25 998 k 25 0.001 10 203 1381 +25 11 78 + 25 0.023 046 _~ __ ~ ~ _ _ _ ____ ____ and hence aw/aT is -7.22 x lo-, mol-l kg K-I. An interpretation of ao/aTmay be found from the theory of dilute solutions of non-electrolytes. DISCUSSION From theoretical considerations12 the Gibbs energy of a solution of mole ratio of solute to solvent m may be written _ _ - - G / N l k T = ~ ~ / k T + ~ ~ ~ / k T - m + m l n r n + ~ A , , ~ 2 f ~ B , 2 , r n 3 + .. . (1 1) where the coefficients A,, etc. are functions of temperature and pressure only, then the apparent molar enthalpy is given by 4H, = H P - RT2[~(2A2,/L1T)pMIrn+~(c?B,2,/c?T),M~m2.. .I. 4H2 = H P +xrn + yrn2 -+ . . . . (12) (13) For sodium tetraphenylboron, let us denote the non-electrolyte contribution to the apparent molar enthalpy by 4HP, then as it was found, as discussed above, that AHp/Am (where Am = mF-rnl) was a constant over the range of molalities investi- gated, the apparent relative heat content of the solutions due to the non-electrolyte (14) contribution is given by 4Lp =xm and the partial relative molar enthalpies of solute, Lp, and of solvent, Lp, are given Lp = 2xm (15) Lp = -Xrn2M1. (16) For simplicity eqn (13) will be written in the form by2098 Then HEATS OF DILUTION OF AQUEOUS SOLUTIONS "P/J mol-1 = 5.335 x lo3 (m,/mol kg-l) LF/J mol-l = 10.67 x lo3 (rn,/mol kg-l) Lp/J mol-1 = - 96.11 (m,/mol kg-l),. (17) (18) (19) Now from eqn (1 1) for a 1 : 1 electrolyte 21nyo = A , , ~ + B , 2 2 ~ 2 + .. . . (20) Thus from eqn (10) and (20), co = A2,M1/2 and the value of (aA,,/aT), is obtained from the heats of dilution since According to the theory of McMillan and Mayer13 for a solution of a solute in a solvent, the osmotic pressure, ll, is given by I l / k T = n+Bz,n2+B,*,,n3+ . . . (23) (24) where n is the number density of the solute. From eqn (16) and (19) of Garrod and Herringtonl, A,,v; = 2B,*,0-2~p+kT~ where B,*,O = - b:,. From KelP4 the molar volume of water at 25 OC is 18.07 cm3 mol-1 and (au;/aT), is 4.56 x cm3 mol-1 K-l.The critical compilation of Bradley and Pitzer15 was used for the compressibility of water 1O1lh-/(N m-2)-1 = 51.5-0.343(t/'C)+ 3.63 x (t/°C)2. The value of (aA2,/i3T>, is given by the heats of dilution. From determination of the osmotic coefficient16 at 25 OC, co = - 1.84 kg mol-l. Millero4 has determined the apparent molar volumes of aqueous solutions of sodium tetraphenylboron in the temperature range 0-60 "C; his data give (av$+/aT), at 25 OC as 0.33 cm3 mol-1 K-l. These figures give for the temperature dependence of B&!' for sodium tetraphenylboron N(aB:,O/aT), = - 72.6 cm3 mol-1 K-l. at 25 OC B:; can be considered to be composed of an attractive and a repulsive contribution from the intermolecular forces, thus B:: = S+q5A (26) where S is the repulsive and q5* the attractive contribution.If a hard-sphere model is assumed then the temperature dependence of B,*,O is that of the attractive contri- bution. In table 3 values of N(8q5A/ii?T)p are compared with those for tetrabutyl- ammonium chloride, sodium chloride, hexamethylenetetramine, sucrose and urea. For tetrabutylammonium chloride aw/aTwas taken as - 8.10 x mol-1 kg K-l,17 (avF/aT), is 0.275 cm3 mol-1 K-l14 and w = -0.032 kg mo1-l.18 For sodium chloride our own data for heat of dilution were used for aco/aT; from ref. (19) (o is 0.29 kg mol-1 and (c?vp/aT), is 0.073 cm3 mol-1 K-1.20 The values for hexa- methylenetetramine, sucrose and urea were taken from ref. (8). It can be seen thatT.M. HERRINGTON AND E. L. MOLE 2099 TABLE 3.-TEMPERATURE DEPENDENCE OF THE ATTRACTIVE CONTRIBUTION TO THE SOLUTE-SOLUTE INTERACTION COEFFICIENT, THE SOLUTE-SOLUTE VIRIAL COEFFICIENT B,*,O AND SOLUTE-SOLVENT VIRIAL COEFFICIENT B:: sodium tetraphenylboron 72.6 - 1570 275 tetra bu t ylammonium chloride 7 . 8 6 262 293 sucrose 0.56 285 210 urea -0.51 1 43 sodium chloride - 2.20 307 16 hexame th ylene te tramine 1 . 5 8 338 110 the solute-solute attraction increases with temperature for sodium tetraphenylboron, tetrabutylammonium chloride, hexamethylenetetramine and sucrose, but decreases for sodium chloride and urea. In sucrose and urea, attractive forces between the molecules may include hydrogen bonding. It has been suggested4 that the large tetraphenylboron anions behave like the large tetraalkylammonium cations in the solute-solute interactions, but like the chloride ion in their solute-solvent interactions.From eqn (24) the solute-solute virial coefficients, Btt, can be calculated. The apparent molar volume of sodium tetraphen- ylboron at infinite dilution at 25 O C is 276.4 cm3 m ~ l - l , ~ for tetrabutylammonium chloride 294.3 cm3 mol-l 21 and for sodium chloride 16.6 cm3 mo1-1.20 Values for NB,*,o are given in table 3 for these salts and also for hexamethylenetetramine, sucrose and urea.8 Sodium tetraphenylboron alone has a large negative value. Solute-solvent interaction can be calculated from apparent molar volume data. (27) From eqn (69) Of ref' (12) byl = -,,?++TK. Values for NB:: (where BT: = -by1) are also given in table 3 for the above solutes.It can be seen that sodium tetraphenylboron and tetrabutylammonium chloride have very similar values for BTF, which, if we assume comparable hard-sphere molar volumes, implies similar solute-solvent interaction. GLOSSARY OF SYMBOLS Superscript denotes the non-electrolyte contribution to the thermodynamic state function X cluster integral for two molecules of solute in pure solvent cluster integral for one molecule of solute and one of solvent in pure solvent Gibbs energy partial molar enthalpy of solute at infinite dilution apparent molar enthalpy Boltzmann's constant partial relative molar enthalpy of solvent partial relative molar enthalpy of solute apparent relative molar enthalpy molar ratio of solute to solvent ( N 2 / N l ) molality of solute - bE2 - by12100 Ml N , N2 n N R S T Lf U Q Y Er n P W & K 4 4* Pl PZ co ' C HEATS OF DILUTION OF AQUEOUS SOLUTIONS molar mass of solvent in kg mol-l number density of solute Avogadro's constant number of molecules of solvent number of molecules of solute gas constant repulsive contribution to the configuration integral absolute temperature molecular volume of pure solvent partial molecular volume of solute at infinite dilution activity coefficient of solute on molality scale rationalized permittivity relative permittivity isothermal compressibility of water osmotic pressure density of water in kg per unit volume osmotic coefficient attractive contribution to the configuration integral chemical potential of solvent chemical potential of solute non-electrolyte interaction coefficient V.Krishnan and H. L. Friedman, J . Phys. Chem., 1971, 75, 3606. J. F. Skinner and R. M. Fuoss, J . Phys. Chem., 1964, 68, 1882. J. F. Coetzee and W. R. Sharpe. J . Phys. Chem., 1971, 75, 3141. F. J. Millero, J. Chem. Eng. Dam, 1970, 15, 562. C. Jolicoeur and P. R. Philip, J . Solution Chem., 1975, 4, 3. S. Subramanian and J. C. Ahluwalia, J . Phys. Chem., 1968, 72, 2525. T. M. Herrington and E. L. Mole, J . Chem. Soc., Furuday Trans. I , 1982, 78, 213. J. E. Garrod and T. M. Herrington, J . Chem. SOC., Faraduy Trans. I , 1981, 77, 2559. E. A. Gulbransen and A. L. Robinson, J . Am. Chem. SOC., 1934, 56, 2637. L. F. Sylvester and K. S. Pitzer, J . Phys. Chem., 1977, 81, 1822. J. E. Garrod and T. M. Herrington, J . Phys. Chem., 1969, 73, 1877. l 3 W. G. McMillan and J. E. Mayer, J. Chem. Phys., 1945, 13, 276. G . S. Kell and E. Whalley, Philos. Trans. R . Soc. London, Ser. A , 1965, 258, 565. D. J. Bradley and K. S. Pitzer, J . Phys. Chem., 1979, 83, 1599. lo P. Debye and E. Hiickel, Phys. Z . , 1923, 24, 185. l 6 T. M. Herrington and C. Taylor, to be published. l 7 S. Lindenbaum, J . Phys. Chem., 1966, 70, 814. l9 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 1959). 2o L. A. Dunn, Truns. Furuday Soc., 1968, 64, 2951. I * B. E. Conway, R. E. Verrall and J. E. Desnoyers, Trans. Faraday Soc., 1966, 62, 2739. S. Lindenbaum and G. E. Boyd, J. Phys. Chem., 1964, 68, 91 1. (PAPER 1 /9 19)

ISSN:0300-9599    

DOI:10.1039/F19827802095

出版商:RSC    

年代:1982

数据来源: RSC