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/F198278FX021

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/F198278BX023

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

JOURNAL OF THE CHEMICAL SOCIETY FARADAY TRANSACTIONS, PARTS I A N D I 1 The Journal of The Chemical Society 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, 11 Journal of The Chemical Society, Perkin Transactions, I Journal of The Chemical Society, Perkin Transactions, 11 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 II 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-phase kinetics, electrochemistry (other 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. Journal of The Chemical Society, Chemical Communications Journal of The Chemical Society, Dalton Transactions Journal of The Chemical Society, Perkin Transactions, I and 11 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 Transact ions. 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 shoula be sent to the Director of Publications, The Royal Society of Chemistry, Burlington House, Piccadilly, London W 1 V 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 OF THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE I T A L I A N A D I C H I M I C A F l S l C A S O C l i T i DE C H l M l E 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 ha1.e 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. Kell. 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 OF 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: M r s Y. A. Fish, The Royal Society of Chemistry Burlington House, London W1V OBN ... 111THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO. 75 Intramolecular Kinetics University of Warwick, 1&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; (b) 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 W1V OBiJ THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO.76 Concentrated Colloidal Dispersions Loughborough University of Technology, 1C16 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 particleparticle 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 31st August 1982 to: Professor R. H. Ottewill, School of Chemistry, University of Bristol, Cantock's Close, Bristol BS8 1TSTHE 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 R. A. Pethrick Dr P. Richmond Dr D. A. Young (Editor) ~ Dr M. La1 1 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 vol urne.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 Surface Reactivity and Catalysis Group Preparation, Performance and Characterisation of Supports in Heterogeneous Catalysis To be held at AERE, Harwell on 30 June 1982 Further information from: Mr L. Evans, Education and Training Centre, AERE Harwell, Didcot, Oxon OX11 ORA 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 Ma\ns Road, Edinburgh EH9 3JJ Industrial Physical 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 VFARADAY DIVISION INFORMAL AND GROUP MEETINGS 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 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 Rheology 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 1 982 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 8QX 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 Physics 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 V iPublications from The Royal Society of Chemistry SPECIALIST PERIODICAL REPORTS Catalysis Vol. 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 oi 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 0 851 86 554 2. Price f29.00 ($62.00). RSC Members f 17.50 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 f25.00 Mass Spectrometry Vol. 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, Ion - 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 f23.00 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, Herts SG6 1 HN.The Royal Society of Chemistry Burlington House Piccad i I I y London W I V 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 4 1 . Also available as a soft-cover booklet from Pergamon Press, Oxford.) Surface Chemistry. ‘ Definitions, Terminology, and Symbols in Colloid and Surface Chemistry - 1.’ (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 and 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 Purv 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’, 197 1 . pp. 71-90.) A complete listing of all IUPAC nomenclature publications appears in the 198 1 Index issues of J. Chem. SOC. viii

ISSN:0300-9599    

DOI:10.1039/F198278FP041

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 1689-1704 Linear Solvation Energy Relationships Part 9.-Correlations of Gas/Liquid Partition Coefficients with the Solvatochromic Parameters, n*, a and D BY MORTIMER J. KAMLET* Naval Surface Weapons Center, White Oak Laboratory, Silver Spring, Maryland 20910, U.S.A. AND ROBERT W. TAFT* Department of Chemistry, University of California, Irvine, California 92717, U.S.A. AND PETER W. CARR* Department of Chemistry, University of Minnesota, Minneapolis, Minnesota 55455, U.S.A. AND MICHAEL H. ABRAHAM* Department of Chemistry, University of Surrey, Guildford, Surrey GU2 5XH Received 3rd December, 1980 When referenced to an alkane solute of similar molecular volume, i.e. log (Ksolute/Kalkane), gas/liquid partition coefficients for nitromethane, methylethylketone, dioxan, toluene and ethanol solutes are well correlated by the solvatochromic parameters, n*, 6, a and B.For nitromethane and dioxan solutes the partition coefficients depend only on the dipolarity/polarizability parameters (n*, 6). For methylethylketone, there is also a measurable dependence on hydrogen-bond donor acidity (a) of protic solvents. For ethanol in aliphatic non-hydrogen-bond donor solvents the dominant effects are solvent dipolarity/polarizability and solvent hydrogen-bond acceptor basicity (B>. For toluene solute there appears to be a negative or desolvation effect in protic solvents which is attributed to a ' hydroxyphobic effect' of weakly basic solutes in alcohol solvents. Earlier papers in this series' have dealt with the formulation of three scales of intrinsic solvent properties (solvatochromic parameters) by means of a ' solvatochromic comparison method', and with the use of these scales and the method to unravel and quantify medium effects on many types of physiocochemical properties and reactivity parameters.A scale of solvent dipolarities,? labelled n*, serves as an index of the ability of the medium to stabilize a charge or a dipole by virtue of its dielectric e f f e ~ t . ~ - ~ For ' select' solvents (non-protic, non-chlorinated aliphatic solvents with a single dominant bond dipole) n* values have been shown5 to be very nearly proportional to molecular dipole moments, p. Abboud and Taft6 have also demonstrated good linearity between the n* values for these select solvents and the dielectric constant function, 8, of Block and Walker.' More recently, Kollings has pointed out that n* values of 32 non-protic aliphatic, polychlorinated aliphatic, and aromatic solvents are reasonably well correlated by a McRae-type of equation which involves a combination t The term solvent dipolarity is intended as a more specific description than the often misused solvent polarity, which has frequently also included the effects of hydrogen-bonding interactions in varying combinations with the dipole/dipole effects.16891690 LINEAR SOLVATION ENERGY RELATIONSHIPS of functions of solvent dielectric constant, E , and solvent refractive index, n. In a similar vein, Bekarek9~lo has shown that there is a very good linear correlation of values of n* for select solvents with the Buckingham cross-term ( E - 1) (n2 - 1)/(2~ + 1) (2n2 + 1). Thus although the solvent n* values have been determined empirically, they can be related to fundamental properties of the solvent medium.An a scale of solvent h.b.d. (hydrogen-bond donor) acidities provides a measure of the ability of protic solvents to donate a proton (accept an electron pair) in a type-A (solvent to solute) hydrogen bond.2* 11, l2 AB scale of h. b.a. (hydrogen-bond acceptor) basicities quantifies the solvent’s ability to accept a proton in a type-B (solute to solvent) hydrogen Rather than being based on the solvatochromic behaviour of single indicators, as has been the case for most earlier solvent property scales,2f16 the solvatochromic parameters were determined by averaging normalized solvent effects on a variety of properties involving many types of indicators.We are now fairly well satisfied with the current values of n* and p for most h.b.a. solvents (i.e. additional results are not likely to materially change the averages), but a, and n* values of several amphiprotic solvents (most h.b.d. solvents are amphiprotic) are still uncertain and remain subject to change. The solvatochromic parameters were intended for use in solvatochromic equations of the following general types. When hydrogen-bonding effects are excluded, as when neither solvents nor solutes are hydrogen-bond donors, correlations of medium effects may take either of two forms: (a) For p + n* or n --* n* electronic spectral transitions with all solvents considered together, and for other properties when families of solvents with similar polarizability characteristics are treated separately (e.g.only non-chlorinated aliphatic solvents, only polychlorinated aliphatics or only aromatic solvent~),~-~ the linear solvation energy relationships take the form bond.2,13-15 XYZ = (XYZ),+sn* where s is a measure of the response of XYZ to changing solvent dipolarity. (b) For other values of XYZ, with all solvents considered together, the preferred linear solvation energy relationship becomes (1 b) XYZ = (XYZ), + s(n* + dd) where 6, a ‘ polarizability correction term = 0.0 for non-chlorinated aliphatic solvents, 0.5 for polychlorinated aliphatics, and 1 .O for aromatic s01vents.l~ We have reported correlations, now numbering in the hundreds, wherein the XYZ term in eqn (1 a) and (1 b) has been a position or intensity of maximal absorption in an ir., u.v.-visible, n.m.r.or e.s.r. spectrum, an n.m.r. coupling constant, a free energy or heat of solution or of transfer between solvents, or the logarithm of a rate or equilibrium constant or of a fluorescence lifetime. It is pertinent to subsequent discussions that correlations of solvent effects on properties other than thep --+ n* transitions by means of eqn (I a) have been best when restricted to a select set of non-protic, non-chlorinated aliphatic solvents, which contain a single dominant bond dipole and for which values of n* are generally proportional to molecular dipole moments. Abboud and coworkers5 have pointed out that most earlier ‘solvent polarity’ scales show good linear correlation with one another and with the n* scale when consideration is restricted to these select solvents.When solutes but not solvents are hydrogen-bond donors, contributions of type-B hydrogen bonding must be included in the solvent effects. In these instances, totalKAMLET, TAFT, CARR A N D ABRAHAM 1691 solvatochromic equations also include dependences on the solvent h.b.a. basicity parameter, B, i.e. and X Y Z = (XYZ),+SZ*+~B X Y Z = (XYZ), + S(Z* + dd) + bp. In the converse case, when solvents but not solutes are hydrogen-bond donors, type-A hydrogen-bonding interactions must be taken into account. Here the multiple solvent effect dependences are on n* and the solvent h.b.d. acidity parameter, a, i.e.X Y Z = (XYZ),+sn*+aa (3 a ) and X Y Z = (XYZ),+s(n*+dd)+aa. In situations where both solutes and solvents have been hydrogen-bond donors it has proved quite difficult to unravel solvent polarity, type-B hydrogen bonding and variable solvent self-association effects29 l1 from (usually multiple) type-A hydrogen- bonding interactions. For this reason we have demonstrated total solvatochromic equations with unambiguously distinct dependences on all three of the solvent parameters, n*, a and a, in only a few instances.l8$ l9 RESULTS AND DISCUSSION One of the present authors2, has recently pointed out that good linear regression is observed for the results in select solvents (for which n* z k p ) when logarithms of mole fraction gas/liquid partition coefficients reported by Rohrschneider21 for nitromethane (NM) solute are plotted against solvent n* values.The least-squares linear regression equation had a correlation coefficient, r, of 0.955. Such a good linear regression with n* was not observed with log Keg values for other solutes reported by Rohrschneider, however. Thus the correlation of log Keq for methylethylketone (MEK) solute in the same solvents led to r < 0.70. Further, the correlation coefficient was only 0.666 when log KMEK was plotted against log KNM. We now report, however, that when (in a procedure akin to that reported by Snyder)22 the partition coefficients for these solutes are referenced to corresponding partition coefficients for alkane solutes of the like molar volume, V, i.e.log (Ksolute/Kalkane), excellent correlations with n* values of select solvents are obtained for MEK as well as NM. Further, satisfactory correlations with n* are also obtained for logarithms of partition coefficients of dioxan, toluene and ethanol solutes, similarly referenced to alkane solutes of like V. We shall refer to Ksolute/Kalkane in a given solvent as the solute/alkane selectivity factor, or S (solute/alkane). REFERENCE SOLUTES OF LIKE MOLAR VOLUME We can justify the use of log K values referenced to alkane solutes on a number of grounds. If the activity coefficient of a solute at infinite dilution, yi, can be separated into a polar and a non-polar contribution such that Yi = Y? YB the non-polar contribution may be given23 by the expression RT In y? = V(Sf - St)21692 LINEAR SOLVATION ENERGY RELATIONSHIPS where df and &! are the non-polar or dispersive parts of the solubility parameters of solute i and solvent s.Thus in order to separate out this non-polar contribution, the yi (or K ) values should be referenced to a non-polar solute of comparable volume and of similar hd to the solute i. In the present case; when an alkane is the reference solute, there is at least a partial cancellation of the non-polar contribution, even though in general dd for the alkane is not exactly equal to df. Although the present method can be justified as above, in no sense can it be regarded as an extension of solubility parameter theory. The latter requires as input knowledge of certain bulk properties of the solute and solvent (i.e.molar volumes and enthalpies of vaporisation), whereas in our method no property of the solute at all is used as input. The intercepts and slopes of the various regression equations are, of course, characteristic of the solute, but these are the output from the procedure. In any case, these output solute characteristics do not refer to the bulk solute but refer to the (monomeric) solute at high dilution in the gas phase and in solution. Another approach is to regard the solution process as consisting of a number of distinct steps: (i) creation of a cavity of the same size as the solute in a given solvent, (ii) reorganization of the solvent molecules around the cavity, (iii) non-polar dispersive solute/solvent interactions and (iv) polar solute/solvent interactions through dipole/ dipole and dipole/induced-dipole effects.In order to separate out effect (i) and, at least partly, effect (iii), solution parameters have often been referenced to a non-polar solute of the same molar volume as the solute in question.24 Once again, the requirements for the reference solute are that it be non-polar and of the same Vas the indicator solute. For the solutes investigated by Rohrschneider, ethane (v = 55 cm3 mol-l) would be a suitable reference for NM (V = 54 cm3 mol-l) and ethanol ( V = 59 cm3 mol-l); propane ( V = 89 cm3 mol-l) is of almost the same molar volume as dioxan ( V = 86 cm3 mol-l) and MEK ( V = 90 cm3 mol-l); and n-butane ( V = 101 cm3 mol-l) would be a suitable reference for toluene ( V = 107 cm3 m ~ l - ' ) .~ ~ We have therefore collected gas/liquid partition coefficients for these and have calculated the ratios S(NM/ethane), S(EtOH/ethane), S(dioxan/propane), S(MEK/propane) and S(to1uenefbutane) in 41 solvents for which 7z* values are known. Logarithms of these selectivity factors are assembled in table 1, together with solvent 7t*, a and j? values.? We shall discuss seriatim the analysis by means of the solvatochromic comparison method of the solvent sensitivity factors for each of the solutes mentioned above. To unravel the effects of solvent dipolarity, polarizability and hydrogen bonding, we have carried out single and multiple parameter correlations involving various combinations of the 41 solvents. In a number of instances, to ascertain whether purely statistical considerations warrant the inclusion of a dependence on a second parameter, we have determined confidence limits, a%, for the individual variables shown using either the F-test or t-test.In all single and multiple parameter correlations a(n*) has been > 99.99%. We have also determined the confidence limit, CL%, that the single regression may be rejected in favour of the double regression according to the statistical test set forth -f Along conceptually similar lines, Snyder" has estimated K values for hypothetical reference alkane solutes of the same molar volumes as the indicator solutes from Rohrschneider's data for n-octane solute using the equation The present approach involves less of an approximation since it relies on more closely matched experimental information, but Snyder's approach has the advantage that all data are from the same source and determined under similar experimental conditions.log Kalkane = ( Vsolute/ 163) (1% KoctanehKAMLET, TAFT, CARR AND ABRAHAM 1693 TABLE 1 .-SOLVENT SENSITIVITES (PARTITION COEFFICIENT RATIOS) OF VARIOUS SOLUTEP RELATIVE TO ALKANE SOLUTES OF SIMILAR MOLECULAR VOLUME solventb select aliphatic solvents (non-protic) 2 cyclohexane 3 triethylamine 7 diethyl ether 1 hexane 4 di-isopropyl ether 13 tetrahydrofuran 5 di-n-butyl ether 11 ethyl acetate 16 methylethylketone 4 1 cyclohexanone 18 acetone 23 NN-dimethylacetamide 25 dimethylformamide 28 N-methylpyrrolidone 27 butyrolactone 29 dimethyl sulphoxide aliphatic non-select solvents 137 propylene carbonate 9 dioxan 26 hexamethylphosphoramide aromatic solvents 8 toluene 14 benzene 35 p-xylene 15 chlorobenzene 33 bromobenzene 97 fluorobenzene 62 iodobenzene 17 anisole 24 pyridine 59 diphenyl ether 37 benzonitrile 46 dibenzyl ether 58 acetophenone 33 nitrobenzene hydrogen-bond-donor solvents 101 t-butyl alcohol 103 n-butyl alcohol 1 12 n-propyl alcohol 104 ethanol 105 methanol 50 acetonitrile 32 nitromethane 102 propan-2-01 0.00 0.00 0.14 0.71 0.27 0.47 0.27 0.49 0.58 0.55 0.24 0.46 0.55 0.45 0.67 0.48 0.76 0.53 0.73 0.48 0.88 0.76 0.88 0.69 0.92 0.77 0.87 0.49 1.00 0.76 -0.08 0.00 1.649 1.815 1.843 1.680 1.786 2.084 1.819 2.134 2.038 2.062 2.107 2.423 2.433 2.463 2.397 2.518 0.55 0.37 2.290 0.81 n.y.d.c 2.370 0.87 1.05 2.289 0.54 0.11 0.59 0.10 0.43 n.y.d." 0.71 0.07 0.79 0.06 0.62 n.y.d.c 0.81 n.y.d.c 0.73 0.22 0.87 0.64 0.66 0.13 0.90 0.41 0.80 0.41 0.90 0.49 1.01 0.39 a - 2.144 1.918 2.071 2.221 2.135 1.980 2.178 2.296 2.276 2.591 2.319 2.218 2.337 0.41 0.46 0.51 0.54 0.60 0.85 0.85 0.46 0.67 0.79 0.80 0.86 0.98 0.15 0.23 0.78 1.596 1.818 1.836 1.936 1.903 2.3 14 2.450 1.763 1.730 2.01 1 2.018 1.672 1.969 2.404 1.995 2.662 2.502 2.471 2.702 2.952 3.020 2.994 3.106 3.326 - 3.217 2.714 2.373 2.642 2.349 2.513 2.750 2.629 2.509 2.815 2.852 2.8 17 2.819 2.895 2.812 2.997 2.38 1 2.444 2.533 2.796 3.170 3.554 2.41 5 - 1.224 1,614 1.601 1.234 1.627 1.946 1.567 - - 2.086 2.436 2.589 2.661 2.577 2.679 2.842 2.35 1 2.757 2.396 1.903 2.098 1.888 2.073 2.272 2.194 1.980 2.360 2.460 2.265 2.409 2.390 2.334 2.503 - 2.029 2.098 2.275 2.560 2.730 3.018 2.194 1.375 1.178 2.045 2.223 2.203 2.460 1.371 1.155 2.146 2.146 2.852 2.457 2.042 1.918 3.174 2.567 3.112 2.541 3.023 2.520 3.834 3.667 3.897 3.582 3.867 3.633 3.850 3.209 4.199 3.902 3.345 2.871 3.848 2.987 3.813 4.202 2.613 1.887 2.688 1.893 2.394 1.766 2.621 1.877 2.770 2.041 2.707 1.842 2.564 1.885 3.045 2.241 3.304 3.280 2.845 2.135 3.214 2.483 3.154 2.435 3.179 2.588 3.323 2.361 3.440 - 2.306 - 2.427 - 2.510 - 2.740 - 3.081 - 3.784 - - - 2.513 - Data of Rohrschneider, ref.(21); solvent numbering is the same in all papers of this series; p values for these solvents have not yet been determined.1694 LINEAR SOLVATION ENERGY RELATIONSHIPS by E h r e n s ~ n . ~ ~ Only when CL and a for the second variable are > 90% is the multiple parameter correlation statistically justified.We use the test of E h r e n s ~ n , ~ ~ and the similar procedure based on the t-test or F-test, only in the sense recommended by Shorter,28 namely that these are simple and convenient methods of assessing the relative importance of explanatory variables in linear regression analyses. We emphasize, however, that sometimes purely chemical considerations, as for example findings from a closely related solvatochromic comparison study, can strongly suggest a dependence on a second parameter, even though the purely statistical factors do not so indicate. NITROMETHANE SOLUTE NM is a moderately polar solute, n* = 0.85 and dipole moment, p = 3.5 D; a weak h.b.d. acid, a = 0.23; and a very weak h.b.a. base, D < 0.25. On this basis, it seems to be appropriate to carry out correlations with n*, (n*, a) and (n*, j?) to assess the FIG.1.-Log (KNM/Kethane) plotted against (a) n* and (b) n*-0.206. (a) 0, Select solvents, r = 0.984; A, non-select aliphatic solvents; 0, h.b.d. solvents; 0, aromatic solvents, r = 0.968 (excluding 15, 33 and 62). (b) 0 , 0, r = 0.988 (excluding 15, 33 and 62). contributions of solvent dipolarity and hydrogen bonding to the total solvatochromic effects. Plots of log S(NM/ethane) against n* and against (n* +dd) are shown in fig. 1 ( a ) and (b). The solid line in fig. 1 ( a ) represents the correlation equation with the n* values of 16 select solvents, which is log S(NM/ethane) = 1.50 + 2.60n* (4 a) r = 0.984, o = 0.17KAMLET, TAFT, CARR A N D ABRAHAM 1695 where CT is the standard deviation.The r value of 0.984 for eqn (4a) compares with r = 0.955 reported by Carr20 for the correlation of the same data not referenced to an appropriate alkane. If the results in the non-select aliphatic solvents, dioxan, propylene carbonate and hexamethylphosphoramide, are also included in the least- squares fit, the r value decreases slightly to 0.978. The dual regression equation with (n*, a), eqn (4b), represents an attempt to ascertain whether type-B hydrogen bonding by WM to h.b.a. solvents influences the partition coefficients. Again the correlation involves the 16 select solvents : logS(NM/ethane) = 1.45+2.50n*+0.21/? (4 b) r = 0.984, CT = 0.18, a(D> = 60%, CL = 50%. Eqn (4b) shows no improvement in r or the standard deviation over eqn (4a), and both the a and CL criteria indicate that the dependence of log S(NM/ethane) on the solvent h.b.a.basicities is insignificant. Comparison of a correlation with n* for a data set including results in h.b.d. solvents with a multiple parameter correlation with (n*, ayprovides information as to whether type-A hydrogen bonding by h.b.d. solvents to the nitro oxygens influences the partition coefficients. Here the appropriate data set includes the 16 select solvents and 7 aliphatic h.b.d. solvents in table 1 . The regression equation with n* is log S(NM/ethane) = 1.42 + 2.65n* n = 23, r = 0.978, CT = 0.17 (4 c) and the corresponding multiple linear regression equation is log S(NM/ethane) = 1.47+2.63n* -0.21a ( 4 4 r = 0.982, CT = 0.16, a(a) = 95%, CL = 99%.Here the a(a) and CL indicators suggest that the dependence of S(NM/ethane) on the solvent a values is significant (but barely so). However, it would be difficult to rationalize a solute/solvent hydrogen-bonding effect which led to a negative dependence on a, and it may be that we are seeing here a manifestation of the ‘hydroxyphobic effect ’, which is discussed below in connection with toluene solute. Next to be considered are the S(NM/ethane) values in aromatic solvents. For some reason, the results for iodobenzene (solvent 62 in table l), and sometimes chlorobenzene (solvent 15) and bromobenzene (solvent 33) have fallen out of line in several of the correlations reported here, behaving in effect as if their n* values were ca. 0.2 units lower than had been determined on the basis of many earlier correlations [as can be seen in fig.1 (a)]. If these data are included in the correlation for aromatic solvents only, the r value is 0.824; the correlation equation excluding these results, represented by the dashed regression line in fig. l(a), is log S(NM/ethane) = 1.70 + 1.71 n* n = 1 1 , r = 0.975, rs = 0.07. (4 4 The method used for the calculation of the d term in eqn (1 b) is given by eqn ( 5 ) 2A(XYZ) s(a1) + s(ar) d = where A(XYZ) is the difference between the values calculated by means of the aromatic and aliphatic solvent regression equations at n* = 0.7 and s(ar) and s(a1) are the slopes1696 LINEAR SOLVATION ENERGY RELATIONSHIPS of those regression equations.? From eqn (4a) and (4e), the d value for log S(NM/ethane) is calculated to be -0.20.A plot of log S(NM/ethane) against (n*-0.206) is shown in fig. l(b). The correlation equation for 16 select solvents and 1 1 aromatic solvents (excluding 15, 33 and 62), represented by the regression line in fig. 1, is log S(NM/ethane) = 1.59 + 2.48 (n* - 0.206) (4f) n = 27, r = 0.988, o = 0.12. If the results for the non-select and protic aliphatic solvents and the halogenoaromatic solvents 15, 33 and 62 are included, the r value drops to 0.955. In summary we conclude that solvent dipolarity and polarizability are the main factors influencing the distribution of NM between gas and solvent phases, and that effects of hydrogen bonding by NM to h.b.a. solvents or by h.b.d. solvents to NM are negligible. There are significant, as yet unexplained, aberrations involving chloro- benzene, bromobenzene and iodobenzene solvents, and a possible ‘ hydroxyphobic effect’ in alcohol solvents.METHYLETHYLKETONE SOLUTE With n* = 0.67 and p = 2.9 D, MEK is a slightly less polar solute than NM; with /3 = 0.48, it is a significantly stronger h.b.a. base; and with a = 0.03, its h.b.d. acidity is barely measurable, even in strong h.b.a. base solvents. It therefore seems reasonable to carry out correlations of log S(MEK/propane) with n* and with (n*, a). Again, the first correlation involves the data for the select solvents (table l), the regression equation being log S(MEK/propane) = 1.26 + 1.48n* n = 14, r = 0.976, o = 0.13. As before, the r value falls off (to 0.960) when the results for the non-select aliphatic solvents are included, and improves (to 0.988) if the cyclohexanone result is excluded.The lower s value (sensitivity to solvent dipolarity/polarizability) of 1.48 for MEK in eqn (6a), compared with 2.60 for NM in the corresponding eqn (4a), is consistent with the lower dipolarity of MEK solute. We next ascertain the effect of type-A hydrogen bonding by protic solvents to the carbonyl oxygen of MEK by comparing the correlations with n* and with (n*, a) for a data set including the 14 select solvents and 6 aliphatic h.b.d. solvents. For reasons which will be discussed below, the results for MEK solute in nitromethane solvent are very much out of line on the high side (see fig. 2) and are excluded from the correlations leading to eqn (6b) and (6c). The correlation equation1 with n* is log S(MEK/propane) = 1.31 + 1.51n* ~1 = 20, r = 0.952, o = 0.16 and the multiple linear regression equation with (n*, a) is log S(MEK/propane) = 1.25 + 1.52n* +0.26a (6 4 r = 0.971, o = 0.12, a(a) =- 99%, CL = 99.5%.t We have shown in a recent paper” that the d term in eqn ( I b) can be related quantitatively to the dipolarity/polarizability blend in the solvent effect on XYZ (expressed in terms of functions of the solvent refractive index and either the dipole moment or the dielectric constant), the polarizability term contributing increasingly more as the d value becomes less negative. 1 If the result for nitromethane solvent were included, r would be 0.943 for eqn (6b) and 0.960 for eqn (6 4-KAMLET, TAFT, CARR A N D ABRAHAM 1697 From the coefficients of n* and a in eqn (6c) and from the a(a) and CL terms, solvent h.b.d.acidity appears to play a small but real part in increasing the gas/liquid partition coefficient for MEK solute in protic solvents. A plot of log S(MEK/propane) against log S(MEK/propane),,,,, eqn (6c), is shown in fig. 2, where the out-of-line behaviour of nitromethane (solvent 32) and cyclohexanone (solvent 41) is particularly noteworthy. FIG. 2.-Log (KMEK/Kpropane) as a combined function of solvent n* and a values. 0, Select solvents, 0, aliphatic h.b.d. solvents, r = 0.971 (excluding 32). As was observed in the earlier instance, the relatively poor correlation of log S(MEK/propane) with n* for all aromatic solvents (r = 0.797) is improved significantly if the result for iodobenzene (solvent 62) is excluded.The regression equation with n* for 13 aromatic solvents is log S(MEK/propane) = 1.46+ 1.09n* Y = 0.905, 0 = 0.09. From eqn (9, (6c) and (64, the d term is calculated to be -0.09, and multiple parameter least-squares correlation of the results for 14 select solvents, 6 aliphatic h.b.d. solvents (excluding 32) and 13 aromatic solvents (excluding 62) with (n* - 0.096) and a leads to the total solvatochromic equation log S(MEK/propane) = 1.28 + 1.48 (n* - 0.096) + 0.24a (6 4 n = 33, r = 0.959, 0 = 0.14. If the results for the non-select aliphatic solvents are also included, r = 0.947. Summarizing for MEK solute, we conclude that the gas/liquid partition coefficients are influenced by solvent dipolarity/polarizability to a lesser extent than with NM solute, and that type-A hydrogen bonding leads to a small but measureable dependence1698 LINEAR SOLVATION ENERGY RELATIONSHIPS on solvent a values.Seemingly, from fig. 2, nitromethane solvent (whose a value is 0.23) has a greater hydrogen-bonding effect on S(MEK/propane) than the stronger h.b.d. alkanol solvents (whose a values are 0.79-0.98). This is an interesting observation, which we shall discuss further below. DIOXAN SOLUTE The n* value of dioxan is 0.55, the D value is 0.37; thus dioxan is a less polar solute than MEK and a slightly weaker h.b.a. base. On this basis, the logical correlations are again with (n*, a). The correlation equation of log S(dioxan/propane) with the n* values of the 16 select solvents of table 1 is log S(dioxan/propane) = 1.69 + 1.43n* r = 0.965, G = 0.14.As in the patterns of behaviour observed earlier, r decreases (to 0.949) if the results for the non-select aliphatic solvents are included and increases (to 0.978) if the cyclohexanone result is excluded. Despite the 18% lower solute n* value, s = 1.43 in eqn (7a) is quite close to s = 1.48 in corresponding eqn (6a) for MEK solute. We rationalize this larger than expected response to solvent dipolarity/polarizability as follows. In comparing the polarity of dioxan with NM and MEK, note that the gas-phase dipole moment of dioxan is nil. The appreciable n* value of 0.55 probably results from one or both of the following considerations. (a) When a dioxan solvent molecule is near a given solute molecule (or vice versa), the other molecule ‘sees’ only part of the dioxan molecule and hence regards dioxan as having a dipole.Consistent with this, the dioxan n* is very near the values of 0.58 for tetrahydrofuran and 0.51 for tetrahydropyran.2 (b) Dioxan can assume multiple conformations. The dominant form in the gas phase and in pure dioxan is the chair conformer, which has no dipole moment, but dioxan solvent molecules surrounding a polar solute (or dioxan solute molecules surrounded by a polar solvent) may very well assume the boat conformation, which has an appreciable dipole moment. In the present case, a change from predominantly the less polar chair form in cyclohexane solvent to predominantly the more polar boat form in DMSO solvent would be expected to lead to a larger value AS(dioxan/propane) between these two solvents than would be the case if it were all-chair or all-boat in both solvents.This could explain the higher than expected slope in eqn (7a). We have noted earlier examples where changing solute conformations (from sp3 to sp2 hybridization on aromatic amine nitrogens) have led to larger than expected solvatochromic effects. lib To assess the possible effects on the distribution coefficients of type-A hydrogen bonding by protic solvents to the dioxan oxygen atoms, we again compare the correlations with R* and with (n*, a) for the data set including the select solvents and the aliphatic h.b.d. solvents, but again excluding nitromethane. The correlation equation with n* is log S(dioxan/propane) = 1.71 + 1.47n* n = 22, r = 0.956, G = 0.14 and the mu!tiple linear regression with (n*, a) is log S(dioxan/propane) = 1.68 + 1.47n* + 0 .1 2 ~ ~ r = 0 . 9 6 1 , ~ = 0.14, a(a) = SO%, CL = 75%.KAMLET, TAFT, CARR A N D ABRAHAM 1699 The data point for nitromethane is out of line on the high side by 0.60 log units relative to either of the above equations (i.e. the distribution coefficient of dioxan solute into nitromethane solvent is 300% higher than calculated). From the a / s ratio of 0.075 in eqn (7c) and from the low values of the a(a) and CL terms, it appears that hydrogen bonding by most protic solvents has little effect on S(dioxan/propane). This was initially surprising to us, because the alcohol sol- vents in table 1 are relatively strong hydrogen-bond donors, and the hydrogen- bond acceptor properties of dioxan have been abundantly demonstrated, often with weak h.b.d.indicators than the alkanols (as will be discussed further below). Considering next the results in 13 aromatic solvents (excluding iodobenzene, 62), the correlation equation with n* is log S(dioxan/propane) = 1.95 + 1.03n* ( 7 4 r = 0.857, CJ = 0.1 1. 32 0 00 0 0.2 0.4 0.6 0.8 1.0 IT* FIG. 3.-Log (Kdioxan/Kpropane) plotted against solvent 1z* values. 0, Select solvents, a, aromatic solvents, 0, aliphatic h.b.d. solvents, r = 0.943 (excluding 32 and 62). From eqn (9, (7a) and (74, the d term in eqn (1 b) is calculated to be -0.02, or essentially nil. We have therefore carried out a correlation with n* for the data set including the 16 select solvents, the 6 aliphatic h.b.d.solvents (excluding 32), and the 13 aromatic solvents (excluding 62), to obtain the regression equation log S(dioxan/propane) = 1.74 + 1.38n* n = 35, r = 0.943, CJ = 0.14. The data are plotted in fig. 3; again the positions of the out-of-line data points are particularly to be noted.1700 LINEAR SOLVATION ENERGY RELATIONSHIPS Summarizing the findings for dioxan solute, we conclude that the gas/liquid partition coefficients are influenced by solvent dipolarity/polarizability to a slightly lesser extent than for MEK solute but, surprisingly, not by the h.b.d. acidity of most protic solvents. Again, however, there is a large out-of-line effect in nitromethane solvent. 2.5 h e, E 3 c 0 Y . 6 - v 2 2.0- r 4 w - TOLUENE SOLVENT As in the case of other less polar aromatic compounds, we see that the polarizability term contributes more and the dipolarity term contributes less to the n* value of toluene than is the case with most aliphatic compounds.(We have discussed this subject briefly,17 and will report on it in greater detail in a future paper.) Thus, although n* for toluene is 0.54, the dipole moment, p, is only 0.4 D. For comparison, ethyl acetate, a typical select solvent, has n* = 0.55 and p = 2.7 D. The value of toluene is 0.1 1, it being a weak h.b.a. n-base. - 0 # 1 1 I I I I I . 50 %y 632 0 t - BuOH FIG. 4.-Log (Ktoluene/Kbutane) plotted against solvent n* values. 0, Aliphatic non-h.b.d. solvents, r = 0.952; 0 , alkanol solvents; 0, other h.b.d. solvents. A plot of log S(toluene/butane) against n* of 16 select solvents is shown in fig.4. The correlation equation, represented by the regression line in the figure is log S(toluene/butane) = 1.65 + 0.80 n* i s a) Y = 0.952, t~ = 0.09 (excluding 41, Y = 0.968, t~ = 0.08). An interesting new effect becomes evident when we compare the results for protic solvents with the regression eqn (8a). These data are also plotted in fig. 4, different symbols being assigned to the aliphatic alcohols and the other h.b.d. solvents. The results for nitromethane (solvent 32) and acetonitrile (solvent 50) solvents conform with eqn (8a), but the data points for the alkanols are displaced to lower values of S(toluene/butane) than called for by the regression equation. Further, the negative displacements are larger the bulkier the R groups in R-OH (which means that the displacements are smaller the greater the solvent h.b.d.acidities).KAMLET, TAFT, CARR A N D ABRAHAM 1701 We have not earlier observed such an effect in our solvatochromic comparison studies, and it obviously does not lead to a simple linear solvation energy relationship between S(toluene/butane) and (n*, a). We believe that the negative displacements may be a manifestion of what might properly be called a ‘ hydroxyphobic effect’ of toluene or other very weakly basic solutes. The self-associating alkanol molecules in the solvent shell may orient themselves so that their more polar hydroxy groups are facing towards one another to form cyclic hydrogen-bonded trimer or tetramer clusters;29 thus the less dipolar alkyl groups would be pointing towards the solute molecules. Effectively, then, the cybotactic environments which the toluene molecules would be ‘ sensing’ would be less dipolar than the bulk solvent and, as a consequence, the distribution coefficients of toluene solutes into R-OH solvents would behave as if the alkanols had significantly lower n* values.If this is indeed the proper explanation for the downward displacements of the R-OH data points from the regression line in fig. 4, the ‘ hydroxyphobic effect’ must be different in origin from the ‘hydrophobic effect’ discussed recently by one of the present authors,26$ 30 as water has no non-polar terminus to point towards non-polar solutes. We prefer the above explanation, but it should also be noted that the results in fig.4 do not exclude an explanation based on a different dipolarity/polarizability mix for usual alkanol solvents compared with usual select solvents, coupled with an unusually large susceptibility of the free energy of solution of toluene solute to solvent polarizability (oide infra). We have also correlated log S(toluene/butane) with the n* values of 12 aromatic solvents (again excluding 62). The regression equation is log S(toluene/butane) = 1.64+ 0.77n* r = 0.774, o = 0.11. From eqn (5), (8a) and (8b), the d value in eqn (1 b) is again calculated to be -0.02 (i.e. again essentially nil). We have therefore carried out a correlation with n* for the data set including the 16 select solvents and the 12 aromatic solvents (excluding 62) (8 4 to obtain n = 26, r = 0.926, o = 0.1 1.Excluding solvent 41, r becomes 0.936; including the non-select aliphatic solvents, r becomes 0.914. We have mentioned that we have related the d term in eqn (lb) to the dipolarity/polarizability blend in the solvent effect on XYZ.” The low negative d values for dioxan and toluene indicate that the free energies of solution of these solutes are more strongly influenced by solvent polarizability than is the case for NM or MEK solutes. Summarizing, we conclude that the distribution of toluene solute between gas and solvent phases is influenced by solvent dipolarity and polarizability, the part played by the latter possibly being greater than for NM or MEK solutes. Further, there appears to be an effect involving alkanol solvents, which leads to lower partition coefficients than would normally be expected based on the alcohol n* values.log S(toluene/butane) = 1.65 + 0.78n* ETHANOL SOLUTE With n* = 0.54 andp = 0.77, EtOH is a moderately polar solute and a strong h.b.a. base. As concerns its a value of 0.86, however, note that this value is for strongly self-associated (polymeric) bulk ethanol solvent, whereas ethanol solute at high dilution is probably mainly monomeric and would be expected to have a lower a value.1702 LINEAR SOLVATION ENERGY RELATIONSHIPS This expectation is based on Huyskens’ obser~ation~l that, when an amphiprotic molecule acts simultaneously as a hydrogen-bond acceptor and donor at the same site, both the donor and acceptor strengths are enhanced substantially relative to the same species acting only as donor or only as acceptor.As was mentioned earlier, multiple type-A and type-B interactions strongly complicate the analysis of the solvatochromic behaviour of amphiprotic solutes in amphiprotic solvents. For this reason we have considered only the S(EtOH/ethane) values in aliphatic non-h.b.d. solvents, and carried out correlations only with n* and with (n*, 8). The correlation equation with the n* values of 15 select solvents is log S(EtOH/ethane) = 1.44 + 2.20n* r = 0.933, r~ = 0.32 and the corresponding correlation equation with (n*, 8) is log S(EtOH/ethane) = 1.11 + 1.54n* + 1.358 (9 b) r = 0.967, CJ = 0.24, a(D) = 99.5%, CL = 99.5%. From the b/s ratio of 0.88, and from the a(/?) and CL values, it is evident that type-B hydrogen bonding by ethanol solute to h.b.a.solvents contributes appreciably to the free energies of solution of ethanol in those solvents, and hence to the gas/liquid par tit ion coefficients. COMPETITION BETWEEN ALCOHOL SOLVENT SELF-ASSOCIATION A N D SOLVENT-SOLUTE HYDROGEN BONDING It is instructive to compare the magnitudes of the enhanced partition coefficients into the solvent phase caused (a) by type-A hydrogen bonding by ethanol solvent to dioxan and MEK solutes, and (b) by type-B hydrogen bonding by ethanol solute to dioxan and MEK solvents. Considering first MEK, (a) the effect of hydrogen bonding by ethanol solvent to MEK solute is given by an aa term in eqn (6c); a = 0.26, a = 0.86, aa = 0.22; (b) the effect of hydrogen bonding by ethanol solute to MEK solvent is given by the b8 term in eqn (9b); b = 1.35, 8 = 0.48, b/? = 0.65.Thus, ethanol-solvent/MEK-solute hydrogen bonding increases partition into the solvent phase by 80 %, whereas ethanol-solute/MEK-solvent hydrogen bonding increases partition into the solvent phase by 350%. A similar analysis for the ethanol/dioxan pair indicates that the ethanol-solvent/dioxan-solute effect is nil, whereas the ethanol- solute/dioxan-solvent effect amounts to 0.50 log units or a 200% enhanced solute partition into the solvent phase. It remains to be explained why ethanol solute, although a weaker h.b.d. acid than ethanol solvent, associates with dioxan or MEK to the greater extent. The answer, of course, lies in the competitive self-association of ethanol solvent. Ethanol (/? = 0.77) being a stronger base than MEK ( 8 = 0.48) or dioxan ( p = 0.37), an ethanol solvent molecule achieves greater stability by associating with another solvent molecular rather than with a dioxan or MEK solute molecule.Only when the ethanol molecule is a solute, and the dioxan or MEK solvent molecules are present in appreciably greater concentrations, is there significant solvent/solute association. This rationale also explains the out-of-line behaviour of nitromethane solvent with dioxan and MEK solutes. Nitromethane ( p < 0.25) is an appreciably weaker h.b.a. base than MEK or dioxan. Hence, with the latter as solute molecules in nitromethane solvent, there is significant type-A (solvent to solute) hydrogen bonding rather than solvent self-association.From the upward displacements of the nitromethane-solventKAMLET, TAFT, CARR AND ABRAHAM 1703 data points from the regression lines in fig. 2 and 3, the nitromethane-solvent/MEK- solute hydrogen-bonding effect is estimated to be 0.5 log units, and the nitromethane- solvent/dioxan-solute association effect is estimated to be 0.6 log unit (that the effect is greater for dioxan than for MEK suggests that dioxan acts as a hydrogen-bond acceptor at both oxygens). On close inspection of fig. 2 and 3, it is seen that similar but smaller effects probably apply for acetonitrile solvent (50) with dioxan and MEK solutes. Thus we have the seeming anomaly that a weaker non-self-associating h.b.d. acid (NM, a = 0.23 ; CH,CN, a = 0.15) participates more effectively in a hydrogen-bonding interaction than a stronger but self-associating h.b.d.acid (EtOH, a = 0.86). This problem of seemingly inverted h.b.d. acidity orders caused by competitive self- association of the alkanols has hindered many of our attempts at correlations with (n*, a).? Only when the indicators have been sufficiently stronger h.b.a. bases than the alkanols to overcome their self-association have we observed proper orders of t ype-A hydrogen- bonding effects. Our analysis of the h.b.a. basic properties or h.b.d. acidic prqerties of solutes that undergo self-association highlights the comments made in the introduction : in order to discuss solute-solvent interactions that involve (as in the present case) the solute effectively at infinite dilution in the solvent, it is of limited use to do so in terms of bulk solute properties such as the solute solubility parameter.It is more fruitful to use solute characteristics such as the output a, b and s values in eqn (2) and (3) that refer to the (monomeric) solute at infinite dilution. The work by R. W.T. was supported in part by a grant from the Public Health Service. The work by M.J.K. was done under Naval Surface Weapons Center Independent Research Task IR-201. Part 8.-R. W. Taft and M. J. Kamlet, Org. Magn. Reson., 1980, 14, 485. M. J. Kamlet, J. L. M. Abboud and R. W. Taft, Prog. Phys. Org. Chem., 1981, 13, 485. M. J. Kamlet, J. L. M. Abboud and R. W. Taft, J. Am. Chem. SOC., 1977,99, 6027. M. J. Kamlet, T. N. Hall, J. Boykin and R. W. Taft, J. Org. Chem., 1979, 44, 2599. J. L. M. Abboud, M. J. Kamlet and R. W. Taft, J. Am. Chem. SOC., 1977,99, 8327. J. L. M. Abboud and R. W. Taft, J. Phys. Chem., 1979, 83, 412. H. Block and S. Walker, Chem. Phys. Lett., 1973, 19, 363. 0. Kolling, Trans. Kansas Acad. Sci., 1981, 84, 32. V. Bekarek, CON. Czech. Chem. Commun., 1980, 45, 2063. lo V. Bekarek, J. Phys. Chem., 1981, 85, 722. l 1 (a) R. W. Taft and M. J. Kamlet, J. Chem. SOC., Perkin Trans. 2, 1979, 1723; (b) M. J. Kamlet and R. W. Taft, J. Chem. SOC., Perkin Trans. 2,1979,349; ( c ) R. W. Taft and M. J. Kamlet, J. Am. Chem. Soc., 1976, 98, 2886. l 2 R. W. Taft et al., paper on a scale in preparation. l3 M. J. Kamlet and R. W. Taft, J. Am. Chem. SOC., 1976, 98, 377. l4 M. J. Kamlet, J. L. M. Abboud, M. E. Jones and R. W. Taft, J. Chem. SOC., Perkin Trans. 2, 1979, l 5 M. J. Kamlet, A. Solomonovici and R. W. Taft, J. Am. Chem. SOC., 1979, 101, 3734. l6 C. Reichardt, Solvent Effects in Organic Chemistry (Verlag Chemie, Weinheim, 1979). l7 R. W. Taft, J. L. M. Abboud and M. J. Kamlet, J. Am. Chem. SOC., 1981, 103, 1080. l9 M. J. Kamlet, C. Dickinson and R. W. Taft, J. Chem. SOC., Perkin Trans. 2, 1981, 353. 2o P. W. Carr, J. Chromatogr., 1980, 194, 105. 21 L. Rohrschneider, Anal. Chem., 1973, 45, 1241. 22 L. R. Snyder, J. Chromatogr. Sci., 1978, 16, 223. t Unfortunately, complications caused by this effect have led us to revise published a values more 342. M. J. Kamlet, C. Dickinson and R. W. Taft, Chem. Phys. Lett., 1981, 77, 69. frequently than we would care to admit.1704 LINEAR SOLVATION ENERGY RELATIONSHIPS 23 J. H. Hildebrand and R. L. Scott, The Solubility of Nonelectrolytes (Dover Publications, New York, 3rd edn, 1964); J. H. Hildebrand and R. L. Scott, Regular Solutions (Prentice Hall, Englewood Cliffs, N.J., 1962). 24 M. H. Abraham, A. Nasezadeh, J. J. Moura Ramos and J. Reisse, J. Chem. SOC., Perkin Trans. 2, 1980, 854. 25 M. H. Abraham and G. F. Johnson, J. Chem. SOC. A, 1971, 1610. 26 M. H. Abraham, J. Am. Chem. SOC., 1979, 101, 5477. 27 S. Ehrenson, J. Org. Chem., 1979, 44, 1793. 28 J. Shorter, Correlation Analysis in Organic Chemistry (Clarendon Press, Oxford, 1973). 29 C. Duboc, Spectrochim. Acta, Part A, 1974, 30, 431, 440. 30 M. H. Abraham, J. Am. Chem. SOC., 1980, 102, 5910. 31 P. L. Huyskens, J. Am. Chem. SOC., 1977, 99, 2579. (PAPER O / 1858)

ISSN:0300-9599    

DOI:10.1039/F19827801689

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1705-1715 Electron-donor Sites of Oxides as Investigated on Oxygen- 17 Exchanged CaO Surfaces BY VALERIO INDOVINA* AND DANTE CORDISCHI Centro di Studio su 'Struttura ed Attivita Catalitica di Sistemi di Ossidi', Istituto Chimica Generale ed Inorganica, Universita di Roma, Rome, Italy Received 22nd April, 198 1 The electron-donor properties of various oxides (BaO, SrO, CaO, MgO, y-Al,O, and Si0,-Al,O,), activated under vacuum in the temperature range 1000-1273 K, have been investigated by adsorption of 9,IO-dimethylanthracene (DMAN), chlorine and oxygen (90% "0-enriched or non-enriched). Upon adsorption of DMAN, radical anions are formed on CaO, whereas radical cations are formed on Si0,-AI,O,. On oxides of intermediate base strength (MgO and y-Al,O,) very low radical concentration was obtained.The data allow for a further specification of the correlation between electron acceptordonor properties of oxides and their acid-base strength. Upon adsorption of Cl,, a two g-value e x . signal (8, = 2.002, g, = 2.013) was observed on alkaline-earth oxides. A nearly identical signal was also observed on adsorption of 0, on the same oxides. Upon adsorption of 0, (non-enriched) on CaO samples, previously exchanged with 1702, the above signal showed some relevant modifications, suggesting that the paramagnetic species contains oxygen atoms. The signal is tentatively assigned to a 0;- surface species. This species is thought to be formed on the surface of alkaline-earth oxides in the electron-donation process from low-coordination oxidic sites (Oiks) toward acceptor molecules.The 0$ks sites are therefore identified as the electron-donor sites. The work also includes a volumetric study of 0, adsorption on the various oxides. By comparing volumetric adsorption data with the concentration of radicals (0; and, O;), it is possible to provide a description of the 0, adsorption mechanism on alkaline-earth oxides. The formation of anion radicals upon adsorption of various acceptor molecules (A) on the surface of alkaline-earth oxides, previously activated under vacuum at high temperature, is a well documented phenomenon. Some typical examples can be found in ref. (1)-(6). The nature of the electron-donor sites has attracted much attention. Most authors1? 4 7 5 7 appear to agree on two main points: (a) the donor centres are surface 0,- ions in low-coordination sites and (b) the mechanism of the electron ., donation is O&+A -+ O-+A'-. A point of some concern with respect to the above mechanism is that whereas the A'- radicals have been detected in several cases, there is as yet no experimental evidence for the simultaneous formation of 0- species. The arguments put forward to explain this result are not completely convincing since 0- species are known to be formed on the surface of alkaline-earth oxides and to possess a stability which allows their detection by e.s.r.* Moreover, Garrone et al. have recently proposed a mechanism for anion-radical formation which does not require electron transfer from the s01id.~ Further experimental work is therefore needed to support these suggestions as to the nature of the electron-donor sites and the electron-donation process.In the present work, we have studied the adsorption of 9,1O-dimethylanthracene, C1, and 0, (90% "0-enriched or non-enriched) on alkaline-earth oxides, Al,O, and Si0,-Al,O,. Relying also on 0, adsorption experiments on l'o-exchanged CaO 17051706 ELECTRON-DONOR SITES OF OXIDES surfaces, we have been able to show that: (a) surface oxygen ions of the oxide participate directly in the electron-donation step and (b) 0- species are formed in this step. EXPERIMENTAL MATERIALS High surface area MgO, y-Al,O, and SO,-Al,O, (alumina content 25%) were prepared as previously described.'? lo BaO, SrO and CaO were prepared by decomposition of the carbonates (Erba, RP) at 1173 K (CaO) or at 1273 K (BaO and SrO) under vacuum.B.E.T. surface area values (m2 g-l), determined by Kr adsorption at 77 K, were: MgO (200), CaO (120), SrO (3.6), BaO (0.8), Al,O, (120) and Si0,-Al,O, (120). 9,lO-dimethylanthracene (DMAN) and n-hexane were of Reagent grade and were further purified by distillation (n-hexane) or by recrystallization (DMAN). 0, and C1, were dried before admission to the adsorption chamber. High-purity non-enriched 0, ('Air Liquide ' 99.95 %), 90 % 170-enriched 0, (Yeda, Israel) and C1, (Matheson) were used without further purification. APPARATUS AND PROCEDURE A weighed amount of specimen (ca. 0.1 g) was placed in a silica bulb equipped with a side e.s.r. tube. Samples were activated under vacuum for 5 h at 1173 K (MgO and CaO), at 1273 K (SrO and BaO) or at 1000 K (y-Al,O, and SO,-Al,O,) before exposure to 0,, C1, or DMAN, generally at 298 K.Volumetric determinations of oxygen adsorption were performed by contacting the activated samples with 0, at a pressure of ca. 50 Pa. Pressure readings were made with a differential pressure transducer (MKS, Baratron) capable of detecting variations of 0.1 Pa. The adsorption was considered complete when two successive readings at 5 min intervals did not differ by > 0.5 Pa. Total adsorption and irreversible adsorption (molecule m-,) were determined as described below. Solutions of DMAN in n-hexane ( lop2 mol drn-,) were evacuated at room temperature and successively contacted with the activated samples by means of a break-seal system.170-exchanged CaO samples were prepared by heating CaO samples, previously activated under vacuum, in the presence of 90% "0-enriched oxygen. The extent of isotopic exchange was monitored by gas-phase analysis on a mass spectrometer (VG, Micromass 601). The e.s.r. spectra were recorded at X-band frequencies on a Varian E-9 spectrometer. The absolute number of spins was determined from electronically integrated spectra using Varian 'strong pitch' (3 x 1017 spin m-l) as a standard. RESULTS ADSORPTION OF 9,lO-D I METHY L ANTHR A C ENE (DMAN) Upon adsorption from a solution of n-hexane, an e.s.r. signal consisting of seven main lines with a splitting of ca. 7.6 G* is observed on Si0,-Al,O,, previously activated under vacuum at 1000 K [fig.1, spectrum (a)]. Each line of the spectrum is further resolved into a multiplet with a splitting of ca. 3 G. The signal corresponds to ca. 4 x 1015 spin m-,. CaO, activated under vacuum at 1173 K, gives, upon adsorption from the same solution as above, an e.s.r. signal [fig. 1, spectrum (b)] showing a poorly resolved sequence of lines with a splitting of ca. 2 G. The signal intensity corresponds to 1.6 x 1015 spin m-,. y-Al,O,, activated at 1000 K, and MgO, activated at 1173 K, give weak and structureless e.s.r. signals. In table 1 the hyperfine proton-coupling constants of DMAN radicals chemisorbed on the surface of CaO and Si0,-Al,O, are compared with those of the radical cation and anion in solution. The comparison shows that an anion radical forms on the surface of CaO and a cation radical on the surface of SiO,-Al,O,.* 1 G = 10-4 T.V. INDOVINA AND D. CORDISCHI 1707 FIG. I.-E.s.r. spectra recorded at 77 K on (a) Si0,-Al,O, and (b) CaO after contacting the samples at 298 K with a solution of 9,lO-dimethylanthracene in n-hexane. TABLE 1 .-HYPERFINE PROTON-COUPLING CONSTANTS OF 9,l 0-DIMETHYLANTHRACENE RADICALS CHEMISORBED ON OXIDE SURFACES AND IN SOLUTIONa CH3 H H H radical 9-10 1-4 2-3 1, 2, 3, 4 ref. 11 cation in solution 8.00 2.54 1.19 - anion in solution 3.88 2.90 1.52 11 chemisorbed on Si02-A1203 7.6 3.0 this work chemisorbed on CaO (4.0) - - 2.0 this work - - - a Hyperfine constants are expressed in gauss (1 G = T). Estimated error: kO.1 G ADSORPTION OF CHLORINE Upon adsorption of C1, at 298 K on alkaline-earth oxides, previously activated under vacuum at 1173 K, a two g-value e.s.r.signal is observed (8, = 2.002 and g, = 2.010). The e.s.r. spectra recorded at 77 K are collected in fig. 2: MgO [spectrum (a)], CaO [spectrum (b)] and SrO [spectrum (c)]. A nearly identical signal has been1708 ELECTRON-DONOR SITES OF OXIDES v FIG. 2.-E.s.r. spectra recorded at 77 K on (a) activated MgO, (b) CaO and (c) SrO, after exposure to chlorine at 298 K. observed by Kibblewhite and Tenchl, upon adsorption of C1, on MgO. The species will be hereafter referred to as 0;-, the assignment being discussed below. The presence of an additional set of low-intensity lines in the low-field side of the spectra, clearly visible on the CaO sample [fig. 2, spectrum (b)], suggests the formation of a second paramagnetic species which probably contains chlorine atoms.Although an identification of this species is not possible, the formation of the C1; radical can be ruled Adsorption of Cl, at a lower temperature (146 K) also fails to give Cl;. OXYGEN RADICALS ON CaO Upon adsorption of 0, (non-enriched) at 298 K on CaO activated under vacuum at 1173 K, a two g-value signal ( g , = 2.0020, g , = 2.013) was observed in the e.s.r. spectrum recorded at 77 K in the presence of 0.1 kPa of 0, in the gas phase [fig. 3, spectrum (a)]. The signal is the same as that observed after C1, adsorption on the alkaline-earth oxides and will be therefore designated as 0:-. The spectrum of this species is broadened at higher oxygen pressure. In particular, the half-height linewidth of the sharp component at g , = 2.0020 doubles when the oxygen pressure is increasedV.INDOVINA AND D. CORDISCHI 1709 f f l x 2.5 FIG. 3.-E.s.r. spectra recorded at 77 K on CaO samples, previously activated under vacuum at 1173 K, after exposure to oxygen at 298 K. Adsorption of non-enriched 0,: (a) in the presence of 0, in the gas phase and (b) after evacuation at 298 K. Adsorption of 90% "0-enriched oxygen: (c) in the presence of 0, in the gas phase and ( d ) after evacuation at 298 K. Adsorption of non-enriched 0, on "0-exchanged CaO: (e) in the presence of 0, in the gas phase and (f) after evacuation at 298 K. from 0.1 to 1 kPa. After removal of the gas-phase 0, by evacuation at 298 K, the e.s.r. signal becomes much more complex.The spectrum observed in these conditions was previously shown to arise from three different surface species: OF, 0; and O;-.14 The three g-value signal of the 0; species (gl = 2.0023, g, = 2.0095 and g, = 2.0185) is shown in fig. 3 [spectrum (b)]. As previously reported, the 0; and 0; species are not detected in the presence of oxygen in the gas phase because of the presence of well known strong broadening effects. Upon adsorption of 0, (90% 170-enriched) on a CaO sample, thermally activated as described above, a spectrum identical to that obtained after adsorption of non-enriched 0, was observed in the presence of 0, in the gas phase [fig. 3, spectrum (c)]. By contrast, after removal of 0, at 298 K, the spectrum appears to be drastically modified. In particular, in the central region of the spectrum [fig.3, spectrum (43 the three components of the 0; species, which dominated the spectrum after adsorption of non-enriched 0, [spectrum (b)], are now absent. Thus, with 170-enriched 0,, the central parts of the spectra recorded in the presence or in the absence of oxygen differ very little: namely, only the lines of the 0;- species are clearly visible [compare spectra (c) and (d)]. This result demonstrates that the species 0;- is also present after removal of oxygen at 298 K. The 0;- signal is not observed in spectrum (b), being obscured by the more intense signal of the 0; and 0; species. An important difference with respect to spectrum (b) is the presence in spectrum (d) of a set of six low-intensity lines with a splitting of ca.5 G. These lines are most probably some of the hyperfine components of the 0; species. Indeed, according to the analysis made by Tench,15 the species (1s0,-170,-170,)-, where 0, is a lattice oxygen of the surface, should give a spectrum consisting of 36 lines centred around g, with At, = 108 G and A: = 70 G.1710 ELECTRON-DONOR SITES OF OXIDES 200 G % FIG. 4.-E.s.r. spectrum recorded at 77 K on CaO after exposure to 90% "0-enriched oxygen. The details of the central part of the spectrum (out of the scale in this figure) are shown in fig. 3, spectrum ( d ) . These lines could not be detected because of their low intensity. However, as A% = A2 = 15 G and A: = A: = 10 G around g, and g,, two sets of 24 lines should be obtained with a splitting of ca.5 G. In the spectrum recorded under vacuum, the simultaneous presence of the species 0; is shown by the persistence of the line at g = 2.10 (g,-component) and by the fact that, with higher receiver gain, 11 lines centred around g, are observed (fig. 4). The &-value (75 G) is in good agreement with that reported (76 G) by Che et al. for 0; on pyridine-promoted CaO. l6 ADSORPTION OF 0, ON 170-EXCHANGED CaO SAMPLES Prior to oxygen adsorption, a CaO sample was exchanged with 90% 170-enriched oxygen. Three successive portions of enriched 0, (2.4 x mol) were contacted at 1173 K for 2 h with a CaO sample (oxygen content 4.9 x mol, expressed as 0,) previously activated under vacuum for 5 h at 1 173 K. The final extent of exchange was 26% of the total oxygen content (surface and bulk) in the CaO sample.The CaO was further activated under vacuum at 1 173 K for 1 h and then exposed to non-enriched 0, at 298 K. The spectra recorded at 77 K in the presence of 0, [spectrum (e)] and after removal of 0, [spectrum cf)] are reported in fig. 3. Spectrum (e) (species OE-), when compared with spectra (a) and (c), shows three main differences: (i) a marked broadening of the line at g, = 2.0020, (ii) the appearance of a new line at g = 2.028 and (iii) a decrease by a factor 2 of the signal intensity (from integrated spectra). The relative intensities of the three components of spectrum (e) are unaffected by varying the microwave power.* In the spectrum recorded under vacuum [spectrum (f)] the 0; signal is again visible and the component of the 0;- species at g = 2.028 is still present.* An expanded spectrum recorded with higher receiver gain does not show additional lines.V. INDOVINA A N D D . CORDISCHI 171 1 The 170-exchanged CaO sample was successively contacted with a large excess of non-enriched 0, at 1173 K for 6 h, evacuated at 1173 K for 1 h and then exposed to non-enriched 0, at 298 K. Spectra (a) and (b) were obtained. In particular, the line at g = 2.028 disappeared and the intensity of the 0;- species was restored. OXYGEN RADICAL ON OTHER OXIDES The formation of oxygen radicals on MgO, thermally activated under vacuum, has been investigated previ0us1y.l~ The main results can be summarized as follows. After 0, adsorption, MgO samples, activated at 1173 K, show a weak signal from the 0;- species.The formation of 0; and 0; observed in some cases is strongly dependent on the previous history of the sample: pre-exposure to H,17-19 and different thermal treatments in air, such as quenching or annealing of the samples after heating at high temperature. 2o On adsorption of 0, at 298 K on SrO, activated in vacuo as for the other alkaline- earth oxides, no e.s.r. signals are observed. However, if adsorption is carried out at 146 K, the signals of 0; and of 0; (the latter in trace amounts) appear. On leaving the sample for a few hours at room te.mperature the signals disappear, indicating a lower stability of these species on the surface of SrO as compared with MgO and CaO. No e.s.r. signals are observed when 0, is adsorbed on BaO, y-Al,O, and Si0,-Al,O,.VOLUMETRIC MEASUREMENTS OF 0, ADSORPTION A measurement of more strongly bound oxygen (irreversible oxygen) was taken as follows. 0, was first adsorbed at 298 K on samples activated under vacuum at 1023 K (7-Al,O, and SiO,-Al,O,), at 1173 K (MgO and CaO) or at 1273 K (BaO and SrO). TABLE 2.-oXYGEN ADSORPTION AT 298 K irreversible oxygena oxygen radicals / 1 Owl5 molecule mP2 / 1 0-15 spin mP2 sample BaO SrO CaO MgO Si02-A1203 r-A1203 1600 170 14 0.5 0.2 0.2 0.0 7c O.Od 0.0 0.0 1 b . c a Amount of oxygen not desorbed by evacuation for 10 min at 298 K. Concentration of radicals obtained by adsorption at 146 K. No radicals were formed at 298 K. Simultaneous formation of 0; and 0; (0; = 5.5 x 1015 and 0; = 1.5 x 1015) on CaO. On SrO the species 0; is present in trace amounts.In some cases 0; and 0; are also formed on MgO (for details see the references quoted in the text). The amount of 0, determined in this way will be called total oxygen. Subsequently, the samples were evacuated for 10 min at 298 K, and a seond portion of oxygen was adsorbed, always at 298 K. The amount of oxygen adsorbed in this last experiment will be called reversible oxygen. The difference between the total adsorption and the reversible amount is the irreversible oxygen reported in table 2. Table 2 also lists the concentration of oxygen radicals as determined by e.s.r.1712 ELECTRON-DONOR SITES OF OXIDES DISCUSSION E L E C T R O N A C C E P TO R-D 0 NOR PROPERTIES OF 0 XI D E S The formation of the radical cation of DMAN on the surface of SiO,-A1,0, and that of the radical anion on the surface of CaO provide new and definitive evidence for the existence of a correlation between electron acceptor-donor properties of oxides and their acid-base properties. In fact, previous work from our group has shown that the concentration of the negative radical of nitrobenzene (NB), formed by adsorption of NB on the surface of oxides, monotonically decreases when the oxides are taken in the order of their decreasing basicity (BaO > SrO > CaO > MgO > Al,O, > Si02-A1,0,).10~ 21 Moreover, the concentration of positive radicals of hexamethylben- zene (HMB) increases when the oxides are taken in the same order of basicity given above.21 Namely, the higher the basicity of an oxide, the higher its electron-donor properties on the one hand, and the lower its electron-acceptor properties on the other.Conversely, when the concentration of the radicals of perylene (PE), anthracene (AN) and naphthalene (NA) is considered for the same oxides (again listed in the order given above), there is a minimum in the radical concentration on oxides of intermediate basicity (MgO and A120,).21 This behaviour may be explained using the assumption that radical anions of PE, AN and NA are formed on the surface of strongly basic oxides (BaO, SrO and CaO) whereas radical cations are formed on strongly acidic oxides (Si02-A1,0,). However, for alternant hydrocarbons, such as PE, AN and NA, the spectroscopic differences between their anion and cation radicals are small, both in the e.s.r. and electronic spectra.,, Therefore, in view of the poor resolution of powder spectra, an unambiguous assignment of the radicals cannot generally be made on the basis of spectroscopic evidence alone.The above considerations account for some difficulties encountered in the assignment of the e.s.r. spectrum of PE chemisorbed on the surface of A1,0,.23724 In particular it is rather difficult to decide whether the radical anion or the radical cation of PE is formed. The situation is much simpler with DMAN in view of the large difference (a factor of ca. 2) in the hyperfine coupling constants of the methyl protons for the radical anion as compared with the radical cation (table 1). Thus (i) the formation of the radical cation of DMAN on SiO,-Al,O,, (ii) the formation of the radical anion on CaO and (iii) the lack of formation of such radicals on MgO and A1,0, are in agreement with the trend observed with PE, AN and NA on the same oxide surfaces.21 The results also provide evidence for the general statement that radical anions are formed by adsorption of PE, AN and NA on basic oxides and radical cations on acidic surfaces.NATURE OF THE ELECTRON-DONOR SITE The simplest way to visualize the formation of negative radicals on the surface of oxides is to invoke a direct electron-transfer process from a surface site toward a given acceptor molecule (02, Cl,, nitrobenzene, etc.). As far as the nature of the electron-donor sites on alkaline-earth oxides is concerned, most authors agree that these consist of surface 02- ions in low-coordination sites (O&J.l9 4 7 5 7 Recently, Garrone et al.have proposed a different mechanism to explain the formation of anion radicals on the surface of alkaline-earth ~ x i d e s . ~ According to these authors, the anion radicals can be formed without any electron transfer from the solid. In fact, an XH molecule (e.g. a hydrocarbon) with a large enough electron affinity can be heterolytically chemisorbed on a surface (Me2+ 02-) site leading to species Me2+ X- and OH;. Subsequently, the carbanion can transfer an electron to a second XH molecule leading to the anion radical XH’-. Note that both mechanisms require the participation of OEts sites. These surface centres are Lewis base sites, althoughV. INDOVINA AND D. CORDISCHI 1713 they act as a source of one electron only, in the first case, and Bronsted base sites, in the second case.Accordingly, a correlation between electron-donor properties and basicity of the surface is expected in both cases. In our opinion, the mechanism proposed by Garrone et aL9 might well be operating with specific molecules and satisfactorily explain the sensitizing effect of pyridine and other molecules in promoting the formation of 0; on alkaline-earth oxides, but the participation of 065, in the electron-donor process emerges in the present study, as will be illustrated below. In particular, the species 0:- and 0; will be shown to originate from 0- species which are the first product in the oxidation of the electron-donor centre O&,. It is therefore convenient to consider first the main features of the 0:- species: 1.The centre is on the surface (or very near to it), as shown by the fact that its e.s.r. signal is substantially broadened on increasing the 0, pressure. 2. The centre contains oxygen atoms since the e.s.r. signal is different on the 'natural' CaO as compared with the 170-exchanged sample. On this latter sample, the broadening of the signal, the appearance of a new, broad component and the decrease of intensity are thought to arise from hyperfine interactions with 170. 3. The centre originates from an oxygen species already present on the CaO sample after activation. The same signal is in fact obtained whether 0, or C1, is adsorbed. An identical signal is obtained upon lS0, or 1702 chemisorption. In principle, three possible species could account for the above features: 0-, 0;- and O!-.Species 0;- with n odd and > 3, also possible in principle, do not appear likely. The 0- species can be ruled out in view of the following two main points. First, the e.s.r. spectrum of the species 0-, previously detected on alkaline-earth 25 shows features which do not agree with those of the 0;- signal. In particular, the g-values are substantially different: 2.047 and 2.0014 for 0- as compared with 2.01 and 2.0020 for 0:-. Secondly, the 0- species readily reacts with 0, leading to the species 0; and, therefore, cannot be observed in the presence of 0,. The species 0;- can also be ruled out. Two possibilities must be considered: (a) the species is the diamagnetic peroxy ion and (b) the species contains two unpaired electrons (0-. - SO-).In the first case no e.s.r. signal should be detected, whereas in the second case the molecule is in a triplet state. However, the e.s.r. spectra of the species 0- - * 0- (Vo centres), previously observed in MgO single crystals,26 are rather different from the spectrum we observe for 0:-. In particular, due to a very large D-term (> 200 G), separated e.s.r. lines are observed for the Vo centre.26 In the light of the above arguments, the 0;- species appears to be the most suitable model for the centre under discussion. 0:- centres, consisting of a triangular array of 0- species, are thought to be placed either on (1 11) surface micro-planes, which are formed during the activation of alkaline-earth oxides under vacuum at high temperature, or more simply on the corner of the oxide p a r t i ~ l e .~ ? ~ ~ A mechanism for 0, adsorption on alkaline-earth oxides can now be considered. The following sequence of surface reactions is proposed : 2 0;- + O,(g) + 2 02,- 2 o,,, + 0;- 3 o,,, -+ 0;- (2 4 (3) 0;- + 0;- + 2 02- +O,(g). 56 FAR 11714 ELECTRON-DONOR SITES OF OXIDES The mechanism is based upon the detection by the e.s.r. technique of 0; [step (1 a)], 0; [step (2a)l and 0;- species [step (2c)l. The formation of diamagnetic species, 0;- [steps (I b) and (2b)l and 0,- [step (3)] is inferred from a comparison of the adsorption data, as determined by the volumetric method, with the concentration of radicals (table 2). The data of table 2 also suggest that reaction (1 b) is prevalent with respect to reaction (1 a) on the surface of BaO and SrO, in agreement with the more pronounced tendency of these materials to give peroxides.Accordingly, 02- surface ions taking part in reactions (1 a) and (1 b) are designated by different symbols, respectively O&,s and OE-, to underline the fact that whereas a large fraction of them can participate to step (1 b) (such as, for instance, on BaO), only those in low- coordination sites are active in step (1 a). Following step (1 a), the electron-donation step, the 0- ions formed undergo surface reactions (2a), (2b) and ( 2 c ) . The occurrence of these reactions accounts for the lack of 0- detection by e.s.r. However, the detection of 0; and 0;- species strongly supports the suggestion that 0- ions are formed first.Moreover, in the case of CaO, which may be studied in a more quantitative way, the concentrations of OF, 0; and 0;- are consistent with the stoichiometry of the adsorption scheme. In particular, from step (1 a) [O;] = [O-] and from steps (2a) to ( 2 c ) , [O-] = [O;] + 2[0:-] + 3 [ 0 ; - ] (5.5 x 1015 > 1.5 x 1015 + 3 x 0.6 x 1015). Finally, step ( 3 ) accounts for possible formation of 0,- species by surface migration of peroxy ions and 0, desorption from particular sites (near to kinks, edges or corners). TABLE 3.-cONCENTRATION OF SPECIES 0;- AFTER 0, OR c1, ADSORPTION AT 298 K C/ 1 O-I5 spin rnU2 sample 0 2 c12 10 SrO - CaO 0.6 2 MgO 0.2 0.8 A1203 0.0 0.0 Finally, we briefly consider the concentration of 0;- species formed on the surface of oxides after adsorption of C1, or 0, (table 3).As expected from the electron-affinity values, the concentrations obtained upon C1, adsorption are higher than those obtained with 0,. Furthermore, the concentration of 0;- is found to decrease mono- tonically in passing from SrO to y-A1203. This finding is relevant in view of the fact that the 0;- concentration is expected, on the basis of the mechanism proposed here, to be proportional to the concentration of O t i s donor sites. Therefore, the concentration of 0;- can be regarded as a rough measure of the concentration of the electron-donor sites. Thus, the correlation between electron-donor properties and surface basicity, previously found to rely on the concentration of anion radicals formed,l09 21 is now confirmed more directly uia the concentration of electron-donor sites. A.J. Tench and R. L. Nelson, Trans. Faraday SOC., 1967, 63, 2254. B. D. Flockhart, L. McLoughlin and R. C. Pink, J . Cafal., 1972, 25, 305. M. Che, C. Naccache and B. Imelik, J. Catal., 1972, 24, 326. D. Cordischi, V. Indovina and A. Cimino, J. Chem. SOC., Faraday Trans. 1, 1974, 70, 2189. * B. D. Flockhart, 1. R. Leith and R. C. Pink, Trans. Faraday SOC., 1969, 65, 542; 1970, 66, 469.V. INDOVINA A N D D. CORDISCHI 1715 B. D. Flockhart, P. A. F. Mollan and R. C. Pink, J. Chem. Soc., Faraday Trans. I , 1975, 71, 1192; B. D. Flockhart and R. C. Pink, J . Catal., 1980, 61, 291. S. Coluccia, A. J. Tench and R. L. Segall, J . Chem. SOC., Faraday Trans. 1, 1979, 75, 1769. A. J. Tench, T. Lawsan and J. F. J. Kibblewhite, J. Chem. Soc., Faraday Trans. I , 1972, 68, 1169. @ E. Garrone, A. Zecchina and F. S. Stone, J . Catal., 1980, 62, 396. lo D. Cordischi and V. Indovina, J . Chem. SOC., Faraday Trans. 1, 1976, 72, 2341. l 1 J. R. Bolton, A. Carrington and A. D. McLachlan, Mol. Phys., 1962, 5, 31. l 2 J. F. J. Kibblewhite and A. J. Tench, J . Chem. SOC., Faraday Trans. I , 1974, 70, 72. l3 J. Roncin, Chem. Phys. Lett., 1968, 49, 2876; D. L. Griscom, P. C. Taylor and P. J. Bray, J . Chem. l4 D. Cordischi, V. Indovina and M. Occhiuzzi, J. Chem. SOC., Faraday Trans. I , 1978, 74, 883. l5 A. J. Tench, J. Chem. SOC., Faraday Trans. 1, 1972, 68, 1181. l6 M. Che, A. J. Tench, S. Coluccia and A. Zecchina, J . Chem. SOC., Faraday Trans. I , 1976, 72, 1553. l7 D. Cordischi, V. Indovina and M. Occhiuzzi, J. Chem. Sac., Faraday Trans. 1, 1978, 74, 456. IH E. G. Derouane and V. Indovina, Chem. Phys. Lett., 1972, 14, 455. lo V. Indovina and D. Cordischi, Chem. Phys. Lett., 1976, 43, 485. 2o C. Angeletti, A. Cimino, V. Indovina, F. Pepe and M. Schiavello, Z . Phys. Chem., (N.F.), 1980, 122, 21 D. Cordischi, V. Indovina and M. Occhiuzzi, in Proc. IInd Symp. Magnetic Resonance in Colloid and 2 2 D. Distler and G. Hohlneicher, Ber. Bunsenges Phys. Chem., 1970, 74, 960. 23 G. M. Muha, J. Catal., 1979, 58, 470; 1980, 61, 293. 24 B. D. Flockhart and R. C. Pink, J . Catal., 1980, 61, 291. 25 N. B. Wong and J. H. Lunsford, J . Chem. Phys., 1971, 55, 3007. 26 B. H. Rose and L. E. Hallinburton, Solid State Phys., 1974, 7, 398 1. Phys., 1969, 50, 977. 237. Interface Science, ed. J. P. Fraissard and H. A. Resing (D. Reidel, Dordrecht, 1980), p. 461. (PAPER 1 /640) 56-2

ISSN:0300-9599    

DOI:10.1039/F19827801705

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Faladay Trans. I , 1982, 78. 1717-1732 Chloride-ion Effects on the Reversible and Irreversible Surface Oxidation Processes at Pt Electrodes, and on the Growth of Monolayer Oxide Films at Pt BY B. E. CONWAY* AND J. MOZOTA Chemistry Department, University of Ottawa, Ottawa, Canada Received 14th May, 1981 By means of cyclic-voltammetry experiments conducted at Pt electrodes in 6 mol dmP3 H,SO, and HClO, at low (21 3 K) temperature, the effects of C1- on the resolvable reversible and irreversible stages of Pt surface oxide formation and reduction have been evaluated in relation to the competition between anodic deposition of OH and 0 species and specific adsorption of C1- at Pt surfaces at potential > ca. 0.8 V EH. C1- effects on the logarithmic growth of the monolayer surface oxide film with time have also been investigated.Both types of effects are discussed in terms of the tendency of adsorbed C1- ion to promote place-exchange reconstruction of the OH/O monolayer on Pt-electrode surfaces. The observed effects of C1- are of interest in the kinetic behaviour of the anodic C1, evolution reaction at oxidized Pt-electrode surfaces where the electrocatalytic properties of the interface depend on the state of the oxide film with co-adsorbed C1- ion present in it. The kinetics of anodic C1, evolution at Pt and Ir have been shown to depend in important ways on (a) the state of surface oxidation of the noble metal s ~ r f a c e l - ~ and (b) the co-adsorption of C1- in the sub-monolayer oxide film394 that is generated at Pt above ca.0.8 V E , and at Ir above ca. 0.6 V E,. The state of surface oxidation and the extent of C1- ion adsorption are not independent factors, since it has been shown by Bagotzky et al.,5 by Breiter6 and in work from this laboratory3> that C1- interferes competitively with surface oxidation process at Pt and Ir, depending on electrode potential and C1- concentration. The competitive adsorption behaviour was shown4 to follow a Temkin-type isotherm, linear in log [Cl-] over 4.5 decades of [Cl-1. A comprehensive review on the relation between electrochemical C1, evolution kinetics and the state and band structure of catalytic conducting oxide films has recently been prepared’ and covers the relevant literature over the past 12-15 years. Also, three recent paper^^-^ from this laboratory on C1, evolution kinetics and C1- adsorption have contained bibliographies of earlier work, so here we shall not reference in any detail the prior research on these topics.drew attention to the C1- dependence of the state of surface oxidation of noble metals, and showed that it determined the kinetics of anodic C1, evolution taking place in aqueous solutions at Pt electrodes. In the present paper we report studies on the effect of C1- on the resolution of the almost reversible from the irreversible components of surface oxide on Pt and the effects of low concentrations of C1- on the potentiostatic growth of sub-monolayer quantities of oxide species at Pt electrodes. Our previous 17171718 SURFACE OXIDATION A T Pt ELECTRODES EXPERIMENTAL 1.GENERAL AIMS The purpose of the experiments was to apply a technique previously developed in this laboratory8 for resolution of the reversible and irreversible stages of Pt surface oxide formation and reduction to evaluate the influence of adsorbed C1- on these processes. Low temperatures, down to 213 K, are employed in fast-sweep cyclic voltammetry experiments to achieve resolution of the reversible from the irreversible processes in surface oxidation and place exchanges, 10 at sub-monolayer levels of OH/O coverage at Pt. Potentiostatic oxide-film growth experiments were also conducted at the monolayer and sub-monolayer stages of Pt surface oxidation in the presence of various low concentrations of Cl- in order to examine how the known3~ 4 coadsorption of C1- in the developing electrodeposited oxide film affects the kinetics of extension of the film. 2.EXPERIMENTAL PROCEDURES The general procedures for high-purity electrochemical studies of surface processes at Pt electrodes were followed, as described in recent previous papers.3* 4~ l1 A small, three-electrode, all-glass cell was used with a H, reference electrode, as employed previo~sly.~-~ A simple cell design is desirable in order to facilitate cleaning. All solutions were made up in water twice distilled, followed by pyrodistillation, as described previously.ll This procedure was found, as in earlier work,R to give very satisfactory clean solutions as determined by the various criteria described in ref. (1 1). For the experiments at low temperatures, the cell was mounted in a large Dewar flask containing an alcohol + solid-CO, mixture.The temperature was measured by means of a pentane thermometer. Temperatures constant to within 1 K could be maintained during the 1-2 h duration of experiments. Solutions were outgassed by 0,-free, purified N,. All experiments were conducted at 298 1 and 21 3 1 K in solutions stirred by bubbling purified N, into them. 3. ELECTRODES A Pt wire electrode, degreased in refluxed acetone, was sealed in a soft-glass tube. The mounted electrode was washed initially in 98% H,S04 for 48 h and then rinsed many times with hot pyrodistilled water before being immediately set in the cell. Such electrodes gave ‘clean’ cyclic-voltammetry behaviour,? l1 after 5-10 anodic/cathodic cycles between 0.05 and 1.4 V E,.4. REFERENCE ELECTRODE All experiments were conducted with a Pt/H, reference electrode in the same solution. Potentials referred to this electrode are denoted by V EH. 5. ACIDS B.D.H. AristaR grade HClO, and H,SO, were used for preparing the electrolyte solutions. As noted previously,ll this grade of the acids is found to be free from any significant traces of heavy metals or from other oxidizable or reducible impurities that might give rise to spurious diffusion-controlled currents in cyclic-voltammetry experiments. 6. MICROMETER TITRATION PROCEDURE FOR C1- ADDITION In order to obtain information on the progression of the competitive anion adsorption effects, from initially very low concentrations mol dm-3), on the Pt surface oxidation process, a dilute solution of KC1 in the acid supporting electrolyte was titrated, in successive aliquots, into the electrochemical cell by means of an accurate micrometer syringe.The solution was delivered through 1 mm Teflon ‘spaghetti’ tubing, terminating in a fine glass jet. The extent of anion blocking of electrodeposition of OH and 0 species at the Pt electrode was derived from the diminution of surface oxide reduction charge in a cyclic-voltammetry i against V profile, as described previou~ly.~, i against Vprofiles were recorded on separate sheets of graph paper after each addition of CI-. The extents of blocking of surface oxidation up to a givenB. E. CONWAY AND J. MOZOTA 1719 potential (see below) by C1- could be evaluated with an accuracy of 2-10%, depending on C1- concentration.7. POTENTIODYNAMIC EXPERIMENTS AND REVERSIBILITY The potentiodynamic experiments in the cyclic mode give the current densities, i, for the surface processes which take place on the Pt electrode in response to changing electrode potential, V. In order to display the reversible and irreversible components of Pt surface oxide formation and reduction at progressively increasing potentials, a potential-sweep programme was employed in which the reversal of the anodic sweep was made at successively higher positive potentials on each successive sweep. Then the reversible and irreversible components are displayed on the current-potential record (on an oscilloscope, in the present work) as easily distinguishable regions. 8. POTENTIAL-TIME PROGRAMME FOR THE STUDY OF Pt OXIDE GROWTH I N C1- From the experimental point of view, the study of the effect of C1- ion on the growth of Pt oxide at a given constant anodic potential, E,, presents several difficulties, one of which cannot be completely overcome, as will be discussed below.Thus, a special potential-time programme had to be devised to ensure the accuracy and reproducibility of the results. The problem arises as follows: when studying the growth of oxide at a noble metal surface at constant potential, E,, by the potentiodynamic method, it is important to realize that the time spent, z, in the sweep to reach the required growth potential can influence the results if it is of comparable magnitude to the growth time, i.e. the holding time, zh, when t h is small.In order to minimize z, in relation to zh for short 7h, a potential-stepping programme is required to reach the holding or growth potential, E,, in order to avoid erroneous measurements which arise from H ionization taking place during the growth time if the step to the growth potential is initiated from, say, 0.0 V EH. While the first effect regarding z, can be minimized in a controlled way, the second experimental difficulty proves to be the limiting one in terms of accuracy of the technique at short times. A problem also arises from the potential-dependence of the Cl- coverage, especially because in the anodic direction, which spans the range 0.0-1.4 V EH, the p.2.c. of Pt is crossed in the potential step. Thus, if a step is applied to reach the growth potential immediately after completion of ionization of the adsorbed H species (i.e.at ca. 0.45 V EH), the coverage of adsorbed C1- will be quite low, and for short holding times (zh < lo-* s) and/or for low C1- concentrations ( < mol dm-3) the equilibrium coverage is not immediately established. This can result in a lower degree of C1- blocking than expected, as is seen experimentally (see below). In order to circumvent this problem to a substantial extent, and the other difficulty mentioned above, the potential-time programme shown in fig. 1 was used. It can be described as follows. (a) Following a period of initial anodic/cathodic cycling to 1.4 V E,, a preconditioning step of formation of oxide at a given potential, E,, is made for a short programmed period of time.(b) Reduction of this oxide and deposition of H is then made to a potential E, in the H region at a sweep-rate s = 20 V s-l. (c) Ionization of H at the same s = 20 V s-l is made, followed by continuation of the potential sweep at the same s up to 0.75 V E,, since no Pt oxide is formed up to that potential in this step of the programme; the fig. 1 programme starts at this step on the left-hand side. There then follow ( d ) a holding period of 10 s at 0.75 V E, to allow C1- to establish an equilibrium coverage at that potential; ( e ) a 'step', at s > 200 V s-l, from 0.75 V EH to the desired constant oxide growth potential, E,; cf) growth of the oxide at the experimental growth potential, E,, for a controlled period of time zh, as in step (a).Steps (6)-cf) are repeated for a series of growth times, zh, or E, can be varied. If stage (c) is eliminated from the programme, the results deteriorate to the point that it is difficult to obtain satisfactory behaviour at small growth times, zh. Although an equilibrium coverage of C1- ion can be established at the potential of 0.75 V EH, this does not guarantee that equilibrium C1- coverage will be re-established at the E, potential chosen for oxide growth, but the procedure using stage (c) does appear to help minimize that problem, i.e. it extends the lower limit of the range of zh and of [Cl-] that can be accurately studied. SOLUTIONS1720 SURFACE OXIDATION AT Pt ELECTRODES variable T,, ,10-3-100s I 0.85-1.4V ( I----\ S-’ v s-’ I time/s FIG. 1.-Potential-time programme for the study of C1- effects on the kinetics of growth of Pt oxide at constant potential, E,, for programmed periods of time, th.The programme described above proved to be the best, but significant remaining experimental limitations which cannot be avoided are evident in the results obtained at th < s and with [Cl-] -= mol dm-3; the extent of the limitations depends on E,. However, under most other conditions the results appeared satisfactory, and the deviations observed for conditions below the above limits were consistent with the trends discussed above. For th > s, there are no serious problems. RESULTS AND DISCUSSION The results obtained will be dealt with in three parts: first the C1- competitive adsorption behaviour at low temperatures ; secondly, the resolution of reversible from irreversible processes in Pt surface oxide formation and reduction; and finally the growth or extension of Pt oxide monolayer films in the presence of Cl-.1. COMPETITIVE ADSORPTION OF c1- I N Pt SURFACE OXIDE FORMATION AT LOW TEMPERATURES The effects of C1- on surface oxide formation at Pt were recently studied by Conway and Novak37 in 0.1 mol dm-3 H2S04 and 0.1 mol dm-3 HClO, with regard to competitive adsorption behaviour. In order to be able to carry out the experiments on evaluation of the reversible component at the temperature of interest, viz. 213 K (alcohol + dry-ice bath), required in the present work, it was necessary to increase the acid concentration to 6 mol dmP3 in H2S04, since more dilute solutions freeze at these low temperatures. The low temperature is necessary in order to distinguish satisfactorily the reversible and irreversible surface processes, in particular to slow down the irreversible process, which becomes more predominant at higher anodic potentials,* sufficiently to distinguish it from the faster, almost reversible process. The required high acid concentration, however, brings about an unavoidable increase in the blocking of surface oxide formation at Pt due to the adsorption of HSO, ions.Normally HSO; adsorption is much weaker than that of C1- but becomes substantial at the high (6 mol dm-3) acid concentrations and the lower temperature.B. E. CONWAY AND J. MOZOTA 1721 This ‘extra’ blocking by the HSO; appears to be similar in significance to the initial effect of C1- observed at low concentration^,^ viz.10-6-10-5 mol dm-3, at 298 K. Thus, it blocks and displaces the H peaks, and selectively blocks the initial stages of surface oxide formation at Pt, indicating that its electrosorption valency4 is, as with Cl-,5* l2 close to one. The effect of C1- was first determined at a sweep-rate of 0.100 V s-l, with a growth potential, E,, fixed at 1.4 V E,. The amount of oxide deposited in the anodic sweep was measured by integration of the i against V profile of the cathodic sweep, i.e. over the oxide reduction peak, in the usual way. -6 -5 -4 -3 - 2 log ( [ KCI] /mol dm-3 ) FIG. 2.-Percentage of surface oxide blocked by CI- as a function of log [Cl-] at: (a) Pt in 6 mol dmP3 H,SO, in potential sweeps taken to 1.4 V E , at 0.10 V s-’ at 21 3 K and (6) Pt in 0.1 mol dmV3 H,SO, in potential sweeps taken to 1.4 V EH at 0.10 V s-’ at 298 K [ref.(4)]. The results shown in fig. 2(a) are; in principle, similar to those observed previously by Conway and Novak under the same potentiodynamic conditions but at 298 K, as shown in fig. 2(b). The first competitive effects of Cl- adsorption on the surface oxidation of Pt at 213 K become discerned at ca. mol dmP3 compared with 2 x lop6 mol dm-3 at 298 The difference can be understood on the basis of blocking by HSO; adsorption which already competes with C1- adsorption at low concentrations of the latter ion, since the percentage blocking is measured relative to the amount of oxide formed at a given Ea in the absence of C1- (in this case, in the presence of the HSO, in the supporting electrolyte).Once the blocking effect at 213 K becomes measurable when [Cl-] > lop5, the relation between percentage oxide blocked and log[Cl-] is found to be a linear one for the range 10-5-10-3 mol dmp3, as seen in fig. 2(a). As [Cl-] > mol dm-3, the blocking is already > 80% and tends to approach a limit of 100% (note that determination of such small amounts of oxide, as 100% blocking is approached, becomes unavoidably rather inaccurate). The initial S-shaped region found at 298 K,374 which is associated with selective blocking of the OH species at Pt up to ca. 1.1 V, is absent in the results for 213 K in 6 mol dmP3 H,S04. This is almost certainly due to the fact that HSO, at high concentration and at the low temperature already itself has some competitive effect with respect to OH deposition.1722 SURFACE OXIDATION A T Pt ELECTRODES The study of C1- adsorption effects was extended to higher sweep-rates and higher oxide formation potentials, Ea.Thus, for the case of s = 10 V s-l and Ea = 1.6 V E,, the effect of Cl- on surface oxide formation at Pt is shown in fig. 3 (a). As can be seen, the linear dependence of percentage blocking on log[Cl-] is observed from 10-5-10-1 mol dm-3, as at 298 K. The percentage blocking at 10-1 mol dm-3 was ca. -6 -5 - 4 - 3 - 2 - 1 log ([KCII /mol dm-3) FIG. 3.-Percentage of surface oxide blocked by C1- as a function of log[Cl-] at Pt in 6 mol dm-3 H,SO, in potential sweeps taken to 1.6 V EH at 10 V s-' at 213 K: (a) percentage blocked based on total oxide charge measured up to 1.6 V EH; (b) blocking of reversible component in Pt surface oxide formation and reduction from relative charge measurements [fig.4(a)-(f)]; (c) percentage of total surface oxide that is irreversibly reduced. 70%, compared with 75% at 298 K, and the slopes for both temperatures were similar, viz. 19% blocking decade-l of [Cl-] at 213 K, compared with 15% at 298 K. The difference between the extents of oxide blocking as a function of log[Cl-] at 21 3 and 298 K (measured by integration of cathodic i against V profiles taken from E, = 1.4 V E,) is due to the smaller oxide coverage attained at a given Ea at low T. The results of fig. 3(a), obtained by integrations from 1.6 V EH, confirm this conclusion, since under the latter conditions the oxide coverage is increased (relative to that attained at 1.4 V) and the behaviour becomes closer to that observed at 298 K for a similar oxide coverage.This stresses the importance of comparing results on the basis of total oxide coverage initially present at the surface. Competitive adsorption effects were also investigated at higher sweep-rates, ca. 10 V s-l, in order to determine the influence of C1- on the reversibility of the early stages of Pt surface oxide formation, as will be described in the next section. 2. RESOLUTION OF THE REVERSIBLE COMPONENT OF Pt SURFACE OXIDE FORMATION A N D REDUCTION IN THE PRESENCE OF c1- As was shown by Angerstein-Kozlowska et aL8 in experiments performed at low T and high s, and also by Gottesfeld and Conway13 by means of reflectance measurements, the early stages of monolayer oxide formation and reduction at Pt, viz.OoH < 15%, are almost reversible. By operating at a sweep-rate of 10 V s-l, it was possible to distinguish clearly theB. E. CONWAY AND J. MOZOTA 1723 reversible component of reduction of the early stages of surface oxide formation at 213 K, as shown in fig. 4(a)-(f). As the concentration of Cl- was increased, the competitive adsorption effect of C1- on the total extent of surface oxide formation (fig. 2) is again seen. The effect of C1- on the resolved reversible component in oxide formation and reduction [fig. 4(a)-df)] was surprising since it was expected that C1- would tend to block this process already at low [Cl-] by favouring the formation of 8 - 4 - 4 E + 0 - 5 1 li - 4 - - a - I I I I I I I 0 0.8 1.6 0 0.8 1.6 potential/V E, potentiallv E, FIG.4.-Series of cyclic-voltammetry i against V profiles for Pt in 6 mol dmP3 H,SO, +x mol dmP3 KCl at 10 V s-l at 213 K: (a) 0; (b) 4 x ( c ) 8 x loP5; ( d ) 4 x low4; (e) 4 x and (f) 6 x lo-* mol dm-3. the irreversible component through increased rate of place exchange due to lateral anion-interaction effects14 on the Pt/OH or Pt/O surface dipoles. However, as can be seen in fig. 4(a)-(f) and as represented quantitatively in fig. 3(b), C1- appears initially to have only a small effect on the reversible component up to [Cl-] = lop3 mol dm-3; at higher concentrations ([Cl-] > ca. 3 x mol dm-3),1724 SURFACE OXIDATION A T Pt ELECTRODES this reversible component also becomes blocked over a relatively small range of increased concentration [fig.4(e) and (f)]. At 213 K it is to be concluded that the rate of place exchange, leading to formation of the irreversibly reduced species, is slow enough not to be influenced substantially by the presence of adsorbed C1- at low surface concentrations. Fig. 3 ( c ) shows the percentage of total oxide formed up to 1.6 V E , at various Cl- concentrations that is reduced irreversibly, allowing as best as possible for the double-layer charging component of the currents. This percentage approaches ca. 95 % as [Cl-] increases to 10-1 mol dm-3. Attempts were also made to study the reversibility of the early stage of oxidation 0.0 0.5 1 .o 1.5 potential/V E, FIG.5.-Cyclic-voltammetry i against V profiles for Pt in 6 mol dm-3 HCIO,; conditions as follows: (a) i, = S, = 10 V s - * , T = 298 K; (b) S, = S, = 10 V S - ' , T = 213 K ; (c) S, = 10 V s-', S, = 60 V S-', T = 213 K.B. E. CONWAY A N D J . MOZOTA 1725 in HClO,, since C10, adsorbs much less strongly than HSO, and thus should allow the reversible component to be observed much better. While this is the case, 6 mol dmP3 HClO, solutions tend to freeze at 213 K, making it difficult to carry out experiments over a long period of time. In fig. 5(a) and (b) the i against V profiles in 6 mol dm-3 HC10, at 10 V s-’ are compared for 298 and 213 K. There is less total oxide formed at 213 than at 298 K, but the reversible component is more evident at the lower temperature.By increasing the cathodic sweep-rate to 60 V s-l [fig. 5(c)], while keeping the anodic sweep-rate at 10 V s-l, the main (irreversible) oxide reduction peak can be seen [fig. 5(c)] to consist of two components, as is also found for the case of Au.14 In fig. 5(c) the value of Ea was made less positive (1.15 V EH) so as to increase the relative amount of the reversible component of the oxide and thus to make it more easily distinguishable. 3. GROWTH OF Pt OXIDE AT CONSTANT POTENTIAL I N THE PRESENCE OF C1- Measurements of the growth of surface oxide layers at constant potential yield interesting information on the growth mechanism and on the ‘ageing’ process in the film, i.e. with regard to transition between reversible and irreversible behaviour.*? l5 In the case of Pts the early stages (OoH < 15%) of surface oxide formation behave almost reversibly in terms of reduction of the deposited OH species.As O,, increases, however, a place-exchange process is considered8, 13, l6 to be responsible for the cooperative turnover of the Pt/OH layer, leading to irreversibility in the subsequent reduction of the restructured oxide film. When the oxide is formed at constant potential, the turnover process, which is relatively slow, is given time to take place. As the growth time, zh, increases, the amount of surface oxide which is irreversibly reduced becomes larger.8 Since one of the effects of adsorbed anions on surface oxide formation at Pt is to assist the time-dependent turnover process,16 presumably by introducing a lateral repulsive interaction between the adsorbed anion and neighbouring OH surface dipoles, it was considered to be of interest to evaluate the effects of C1- on the kinetics of growth of Pt oxide at constant potential as a function of the C1- concentration.First, however, we shall describe the kinetics of Pt surface oxide growth in the absence of C1-. As described earlier, a special potential-time programme was devised for this part of the work (fig. 1). 4. KINETICS OF Pt OXIDE GROWTH I N THE ABSENCE OF C1- The growth of Pt oxide at constant potential in 0.5 mol dm-3 H,SO,, in the absence of C1-, obeys a direct logarithmic law16 l 8 in growth time, zh, as shown in fig. 6 ( a ) and (b) for the range zh = 10-3-100 s, for various anodic growth potentials, Ea. The relation can be expressed by an equation of the form qt = A logt+q(t = 0) (1) where q(t = 0) is the charge for oxide formation measured in a multisweep experiment at highs and A is a constant.Both A and q(t = 0) depend on Ea, so that as Ea increases, A and q(t = 0) also increase. The dependence of A on Ea will be examined in more detail below. It can also be observed in a qualitative way [fig. 7(a)] that, as the growth time increases when Ea is low (ca. 0.9 V), it is the relative amount of the more irreversibly reduced component that increases with time. This also applies in the presence of C1- [fig. 7 m i .SURFACE OXIDATION A T Pt ELECTRODES N I E Y 2 s . P) EJV E , I I I I I -3 - 2 -1 0 1 2 log (time/s) FIG. 6. Oxide charge as a function of log (time) for growth of Pt oxide at nine constant potentials, E,, as indicated on the figure.A ‘direct’ logarithmic law such as eqn (1) for oxide film extension with time, t, is often found for growth of very thin films or sub-monolayers.16-19 Such a law can be deduced on the basis that the rate constant, k,, at a given potential is exponentially dependent on coverage, 8, through a Frumkin-type interaction parameter, g, [ke = ke-,, exp ( - g o ) ] . Then a growth-rate equation of the form (2) dI3/dt = k exp (-g8) integrates to the following logarithmic growth law: (3) 1 1 8, = -ln(t+c/k)+-lngk g g where c is an integration constant. Only an equation of the form of eqn (2) will do this, i.e. one not containing a pre-exponential term in 1-13 (see forms of integrals in integral tables).Alternatively, if the local potential drop across a sub-monolayer oxide film layer,B. E. CONWAY A N D J. MOZOTA 1727 1 .o 0.5 0 - 0 . 5 d f c -1.0 2 . * 3 u 1.0 0.5 0 -0.5 - 1 .O I 0.0 0.5 1 .o potential/V E , FIG. 7.-Cyclic-voltammetry i against V profiles for Pt in 0.5 mol dm-3 H,SO,, 298 K, after various rh: 0, lO-l, 1, 10 and 100 s, at E, = 0.9V E , in (a) 0.5 moldm-3 H,SO,; (6) 0.5 mol dmP3 H,SO, + 1 x loPs mol dmP3 KCl. within which place-exchange processes are occurring, falls with time due to change of the surface dipole potential difference, which will be approximately proportional to the coverage of oxide species in the place-exchanged state, a similar integrated rate relation, logarithmic in t, results. A more detailed treatment of the ‘direct’ logarithmic growth law will be given elsewhere in a forthcoming publication.5. KINETICS OF Pt OXIDE GROWTH I N THE PRESENCE OF c1- When C1- is added to the supporting electrolyte (here 0.5 mol dm-3 H,SO,) two effects appear immediatel~:~.~ H deposition is blocked and the distinguishable peaks in a cyclic-voltammetry experiment are shifted towards less positive potentials ; also surface oxide formation is blocked by the competitive adsorption of C1-. These are the usual effect^.^-^ In addition, however, as the oxide grows at constant potential in the presence of Cl-, it becomes qualitatively apparent, as shown in fig. 7(b), that C1- adsorption leads to preferential formation of the more irreversibly reduced component(s) [cf. fig. 7(a) and (b)], i.e. C1- increases the extent or rate of turnover of the deposited OH/O species.1728 SURFACE OXIDATION A T Pt ELECTRODES 1 I I 0 0.5 1.0 I I I 0 0.5 1 .Q potential/V E , FIG.8.-Series of cyclic-voltammetry i against Vprofiles for Pt in 0.5 mol dm-3 H,SO,, 298 K, for various th: 0, lo-*, (c) lo-,; ( d ) 1 , 10 and 100s for E, = 1 . 1 V E,, at various [Cl-1: (a) 0; (b) (e) lo-, and (f) 10-1 mol dm-3. Fig. 8 shows a series of i against V profiles for various increasing [Cl-] (0-10-l mol dmP3), for oxide growth at E, = 1.1 V E, for various rh (10-3-100 s), which clearly demonstrates the progressively increasing blocking of oxide formation as [Cl-] increases, together with favoured growth of the irreversible component. Thus the apparent cathodic peak potential, Ep, determined at the capacitance maximum for oxide reduction, is displaced to less positive values as shown in table 1 for zh = 100 s together with the oxide coverage evaluated in terms of charge for its reduction.This effect of Ep becoming less positive with diminishing qoH due to C1- adsorption is even more striking if it is noted that as [Cl-] increases, the coverage of OH species becomes substantially decreased. Contrarily, in the absence of C1- as discussedB. E. CONWAY AND J. MOZOTA I729 [Cl-]/mol dm-3 Ep/V E , 0 0.750 1.3 273 10-5 0.750 0.93 195 10-4 0.741 0.79 166 10-3 0.733 0.50 105 10-2 0.672 0.19 40 lo-' 0.568 0.08 17 a 6,, = 1 = 210 pC cm-2. [Cl-1 /mol dm-3 log (time/s) FIG. 9.-Oxide charge, q, as a function of log (time) for growth of Pt oxide at constant potential, E,, of 1.05 V EH (298 K) for various [CI-1, as indicated on the figure.previously,s the Ep normally appears at more positive potentials with decreasing surface oxide coverage, i.e. as reduction becomes less irreversible. The presence of C1- does not seem, however, to change the direct logarithmic dependence of extent of growth on time, as is shown in fig. 9; however, as [Cl-] increases, the slopes of the lines decrease. Notice that for low [Cl-] and small zh values the charge readings appear to be too high; this is for the experimental reason discussed in part 8 of the Experimental section. However, when this graphical approach is applied to the data obtained at all the E, values studied, viz. in the range 0.85-1.4 V E,, the slopes of the charge against log[Cl-] lines behave in an interesting way, as illustrated in fig.10: two types of behaviour are observed: first, for [Cl-] = 0, lo+, and 10-1 mol dm-3, the values of the slopes increase with increasing E,, but after a narmw S-shaped transition region the relation becomes linear. The curves for [Cl-] = and 10-1 mol dm-3 in fig. 10 are obviously displaced towards higher potentials due to the strong blocking effect of C1- at these two high concentrations, for which the onset of oxide formation has been displaced to 1.0 and 1.05 V EH, respectively. lo-* and Secondly, although it would be expected that the results for [Cl-] =1730 SURFACE OXIDATION AT Pt ELECTRODES Y 0 0.8 L Ea/V E, FIG. lO.--dq/d(log t ) from fig. 9 as a function of the growth potential, E,, for various [Cl-1, as indicated on the figure.mol dm-3 should fill the gap between the above two extreme curves and behave in a similar way with respect to increasing E,, this is only observed, to a certain extent, up to potentials E, d 1.05 V E,, quite close to the potential at which the monolayer of OH species is completed at Pt in the absence of C1-. When E, > 1.05 V E, however, the value of q,,/d(logt) is higher than expected, although the S-shaped region, followed by the linear region, is still maintained. It seems reasonable to interpret the abnormally high slopes observed when E, > 1.05 V E, and [Cl-] = mol dm-3 as being due to competition between oxide growth and oxide blocking by C1- adsorption. Thus, at low [Cl-1, viz. 10-5-10-3 mol dm-3 and with E, < 1.05 V E,, C1- blocks the oxide formation in the expected way.However, when E, > 1.1 V E,, then the oxide formed on the surface competes actively with the adsorbed C1-, probably reducing the coverage of the latter and thus increasing the growth rate, i.e. the dq/d(logt) value. It is evident from the charge against log t plots that at longer zh and high E, (> ca. 1.1 V E,) the lines corresponding to the above [Cl-] values tend to join with the one corresponding On the other hand, for high [Cl-] (10-2-10-1 mol dm-3), the effect of C1- is strong enough that the electrodeposition of oxide cannot compete effectively (at least at 1.4 V E,, the maximum E, attained) with the C1- blocking effect, and hence the curves behave as expected. The change in behaviour around 1.1 V EH was in a sense anticipated, since the previous work4 had indicated that C1- blocks selectively the formation of the OH monolayer, which is completed at ca.1. I V, before starting to block the formation of oxide in a higher state of oxidation. The transition is even more evident when the percentage oxide blocked is plotted as a function of log[Cl-1, for various E, and for two zh values, as shown in fig. I 1 (a) and (b). The behaviour shown in fig. 11 (a) is similar to that found in the multisweep studies of Conway and Novak4 and shows a linear-log region for the blocking effect, which starts to curve and reaches close to 100% blocking at high [Cl-] for E, values and [Cl-] = 0. < 1.1 VEh.B. E. CONWAY AND J. MOZOTA 1731 log ([KCll/mol dm-3) FIG.1 1 .-Percentage of surface oxide blocked by C1- at Pt (298 K) as a function of log [Cl-] for growth times, ?hr of (a) lop3 and (b) 10 s at various growth potentials, EJV EH: (a) x , 1.0; 0, 1.05; A, 1.1; 0, 1.2; V, 1.3; +, 1.4; (6) x , 0.85; 0, 0.9; A, 1.0; 0, 1.05; V, 1.1; +, 1.2; 0 , 1.3; A, 1.4. When the results for a longer growth time are plotted, as in the case of fig. 11 for zh = 10 s, then the linear region disappears for E, > 1.1 V E,, indicating again that some other component influences the behaviour. Notice that the curves for Ea = 1.2, 1.3 and 1.4 V E , are bent down in the range 10-5-10-3 mol drnp3, i.e. where the higher dq/d(log t ) values were found (fig. 10). This bending effect is consistent with the higher slopes observed in fig.10 and thus indicates less blocking by C1- than expected on the basis of comparison with the behaviour in the absence of C1- (fig. 6). CONCLUSIONS It is concluded that, despite some clearly understood limitations of the experimental procedure available for study of C1- effects on reversibility and surface oxide growth at short times, several definite and interesting new effects can be demonstrated; C1-, as expected, blocks the formation of Pt surface oxide. It also reduces the rate of growth at constant potential, except for conditions such as high anodic potentials and low [Cl-1, where it appears that the oxide being formed is able to displace C1- from the surface and thus an unexpectedly high rate of growth is observed. At the same time.1732 SURFACE OXIDATION A T Pt ELECTRODES C1- favours the growth of the resolvable irreversible component of surface oxide formation and reduction, probably by increasing the rate of turnover or place exchange.On the other hand, the blocking effect of C1- on surface oxide formation at Pt under cyclic conditions at 213 K is similar to that observed at 298 K in that a linear Temkin-type relation is found between percentage oxide blocked and log [Cl-1. However, it is clear that, for comparative purposes, results should be referred to the same initial oxide coverage. At 213 K, the behaviour of C1- at low oxide coverages ressmbles that of Br- and I- at higher oxide coverages at 298 K. The reversible component of surface oxide formation at Pt is clearly distinguishable at 213 K in 6 mol dm-3 H,SO, and even more in 6 mol dm-3 HClO,.Investigation of the effect of C1- on the charge for this resolved component showed effects smaller than those expected on the basis of anion/oxide interaction effects, relative to those on the irreversible component. The C1- competitive adsorption effect on monolayer oxide film formation is stronger and qualitatively different from that on Ir recently reported elsewhere.20 Support for this work from the Natural Sciences and Engineering Council of Canada is gratefully acknowledged. J. M. acknowledges the award of a CONICIT (Venezuela) post-graduate research scholarship. E. L. Littauer and L. L. Shrier, Electrochim. Acta, 1966, 11, 527. B. E. Conway and D. M. Novak, J. Electroanal. Chem., 1979, 99, 133. B. E. Conway and D. M. Novak, J. Chem. SOC., Faraday Trans. I, 1979, 75, 2454. D. M. Novak and B. E. Conway, J. Chem. SOC., Faraday Trans., I , 1981, 77, 2341. V. S. Bagotzky, T. B. Vasillyev, J. Weber and J. N. Pirtshkaleva, J. Electroanal. Chem., 1970, 27, 31. M. W. Breiter, Electrochim. Acta, 1963, 8, 925. ’ D. M. Novak, B. V. Tilak and B. E. Conway, in Modern Aspects of Electrochemistry, vol. 14, ed. J. O’M. Bockris, B. E. Conway and R. White (Plenum Press, New York, 1982), in press. H. Angerstein-Kozlowska, B. E. Conway and W. B. A. Sharp, J. Electroanal. Chem., 1973, 43, 9. M. A. H. Lanyon and B. M. W. Trapnell, Proc. R. SOC. London, Ser. A , 1955, 227, 387. lo N. Sat0 and M. Cohen, J. Electrochem. SOC., 1964, 111, 512. l1 B. E. Conway, H. Angerstein-Kozlowska, W. B. A. Sharp and E. E. Criddle, Anal. Chem., 1973,45, 1331. J. W. Schultze, Electrochim. Acta, 1977, 21, 327. l3 B. E. Conway and S. Gottesfeld, J. Chem. Soc., Faraday Trans. I , 1973, 69, 1090. l4 H. Angerstein-Kozlowska, B. Barnett and B. E. Conway, to be published. l5 M. E. Folquer, J. 0. Zerbino, N. R. de Tacconi and A. J. Arvia, J. Electrochem. SOC., 1979,126, 592. A. K. N. Reddy, M. A. Genshaw and J. O’M. Bockris, J. Chem. Phys., 1969,48, 671. D. Gilroy and B. E. Conway, Can. J . Chem., 1968, 46, 875. D. Gilroy, J . Electroanal. Chem., 1976, 71, 257. l9 P. T. Landsberg, J. Chem. Phys., 1955, 23, 1079. *O J. Mozota and B. E. Conway, J. Electrochem. SOC., 1981, 128, 2142. (PAPER 1 /774)

ISSN:0300-9599    

DOI:10.1039/F19827801717

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1733-1740 Effect of Electrolyte and pH on the Interaction of Sodium Carboxymethyl Cellulose on Barium Sulphate Particles BY PETER A. WILLIAMS, RAYMOND HARROP* AND GLYN 0. PHILLIPS North East Wales Institute of Higher Education, Connah's Quay, Deeside, Clwyd CH5 4BR AND GEOFFREY PASS Department of Chemistry and Applied Chemistry, University of Salford, Salford M5 4WT AND IAN D. ROBB Unilever Research, Port Sunlight Laboratories, Wirral, Merseyside L62 4XN Received 18th May, 1981 Adsorption isotherms of sodium carboxymethyl cellulose onto barium sulphate have been determined as a function of the degree of carboxylation of the polyelectrolyte, pH and added electrolyte concentration. At pH 6.5, high-affinity isotherms are obtained, the adsorption capacity increasing with decreasing charge density of the polyelectrolyte. The addition of simple electrolyte enhances the adsorption by an amount which depends on the nature of the ions present.At pH 2, the affinity for the surface decreases but the adsorption is enhanced for all charge-density types. These observations are discussed in conjunction with electrophoretic mobility and viscosity measurements, and it is demonstrated that electrostatic contributions largely dominate the adsorption process. Although a considerable amount of systematic work has been carried out over the last thirty years or so on the adsorption of polymers at the solid-liquid much of the work has involved non-ionic polymers, often in non-aqueous solvents. There has been relatively little work reported on polyelectrolyte adsorption from aqueous solution. This is perhaps surprising, since not only are polyelectrolytes widely used as flocculants, particularly in water clarification,,* soil aggregations and mineral processing,' and as stabilising agents in many food and pharmaceutical products,6-10 but more importantly this category of compound includes many biopolymers and a knowledge of their interfacial behaviour may eventually lead to a clearer understanding of many biological processes.The theory of non-ionic polymer adsorption is now well-establishedl19 l2 but this is not the case for polyelectrolytes. Hesselink,13 however, has recently put forward a theory for polyelectrolyte adsorption onto charged surfaces from aqueous electrolyte solutions where the adsorption is thought to resemble that of non-ionic polymers. The theory predicts that an increase in adsorption is brought about by an increase in ionic strength and a decrease in the degree of dissociation, which had been generally found experimentally.'* 14-16 The purpose of the present investigation was to provide a systematic study of the adsorption of a polyelectrolyte, namely sodium carboxymethyl cellulose of various charge densities, onto a crystal surface, namely that of BaSO,, at different degrees of ionisation under different conditions of ionic strength.17331734 SODIUM CARBOXYMETHYL CELLULOSE ON BaSO, EXPERIMENTAL MATERIALS Commercial-grade barium sulphate was washed by decantation five times with hot distilled water and was then dialysed against distilled water for one week.It was then freeze-dried, followed by heating at 378 K for 12 h. The surface area was found to be 2.9 m2 g-l (by nitrogen adsorption) and the average particle diameter 0.5 pm (by electron microscopy). ESCA showed the surface to contain the elements Ba, S, 0, Ca and P in the ratio 1: 1.05:4.34:0.03:0.05, respectively. Sodium carboxymethyl cellulose (SCMC) samples were supplied by Hercules Powder Co. Ltd, and are referred to as 4 M, 7 M and 12 M, respectively. There were found to have a degree of substitution of 0.45,0.78 and 1.27, respectively (determined by spectrofluorimetric titration”), and a reported molecular mass of 250 000. The samples were dialysed against distilled water for one week and then freeze-dried before use.ADSORPTION ISOTHERMS These were determined as follows: 20 cm3 of SCMC solution (10-500 ppm) were added to 1 g quantities of BaSO, contained in sealed tubes. The tubes were agitated on a laboratory shaker at room temperature for 1 h and left for a minimum of 24 h to ensure that adsorption had reached equilibrium. It was shown that the adsorption was 98% complete within 1 h irrespective of the experimental conditions. The samples were then centrifuged at 1300 g for 30 min and the amount of SCMC remaining in solution was determined using a spectro- photometric method based on anthrone.18 Hence the amount adsorbed could be found by difference. For isotherms determined in the presence of salts or acid, the latter were added to the SCMC solution before the solutions were added to the BaSO,.ELECTROPHORETIC MOBILITY The electrophoretic mobility of barium sulphate or barium sulphate with adsorbed SCMC was determined using a Rank Brothers particle microelectrophoresis apparatus (mark 11). A flat cell with platinized platinum electrodes was used and maintained at 298 K. The samples of coated BaSO, particles were first washed with distilled water to remove an excess of polymer. Dispersions containing ca. lo8 particles were then prepared in 0.02 mol dm-3 NaCl. The electrophoretic mobilities were determined at various pH values but at constant ionic strength by adjusting the solution pH with 0.02 mol dm-3 HC1 or NaOH. Zeta potentials were calculated using the Smoluchowski equation. VISCOSITY MEASUREMENTS The viscosities of the SCMC solutions were determined using an Ostwald U-tube viscometer at 298 K.All times were > 100 s, and hence kinetic-energy corrections could be neglected. RESULTS ADSORPTION ISOTHERMS Fig. 1 shows the amount of SCMC adsorbed on barium sulphate from water at pH 6.5 as a function of equilibrium concentration of solute in water, for the three samples of SCMC. In each case high-affinity isotherms are obtained, and the amount adsorbed reaches a constant value, the adsorption capacity, at low equilibrium concentrations. The capacity is shown to be greater the lower the degree of substi- tution. Thus the adsorption capacity for 4 M is greater than that for 7 M, although the latter is only slightly higher than that for 12 M. The adsorption capacity increases as the ionic strength of the solution is increased by the addition of simple electrolyte.Fig. 2 illustrates the effect of sodium chloride on the adsorption of 7 M at pH 6.5 as a function of ionic strength. Sodium sulphateP. A. WILLIAMS, R. HARROP, G . 0. PHILLIPS, G. PASS A N D I. D . ROBB 1735 ‘60 I 100 200 3 00 400 equilibrium concentration/mg dm-3 FIG. 1.-Adsorption isotherms of SCMC in water onto BaSO, at pH 6.5: 0,4 M SCMC; 0, 7 M SCMC; .,12 M SCMC. 2.5 -i 2.0 E 2 1.5 2 4 * 1.0 5 0.5 Do M \ 2 1 1 1 1 0 100 200 300 400 equilibrium concentration/mg dm-3 FIG. 2.-Adsorption isotherms of 7 m SCMC onto BaSO, at pH 6.5: a, from water; 8, from NaCl ( I = lop3 mol dm-3); 0, from NaCl ( I = lo-* mol dm-3); 0, from NaCl ( I = lo-’ mol dm-9. co ” 3 2 1 p [Il/mol dm-3 FIG.3.-Effect of ionic strength and nature of anion on adsorption capacity: 0, sodium chloride; 0, sodium sulphate.1736 SODIUM CARBOXYMETHYL CELLULOSE ON BaSO, 0 100 200 300 400 equilibrium concentration/mg dm-3 FIG. 4.-Adsorption isotherms of SCMC onto BaSO, at pH 2: A, 7 M from H,SO,; A, 12 M from H,SO,; A, 4 M from H,SO,; 0, 7 M from HCI; (>, 12 M from HC1; 0, 4 M from HCI. also gives an increase in adsorption capacity although lower than that for sodium chloride. This is illustrated in fig. 3. Fig. 4 shows the adsorption isotherms of the three SCMC samples when the pH of the solution had been adjusted to 2 using both hydrochloric and sulphuric acids. Although the adsorption capacity is not reached in the concentration range studied, note that the amount adsorbed is considerably higher than at pH 6.5.In addition the initial slope of the isotherms indicate low-affinity adsorption. When sulphuric acid is used to adjust the pH, the adsorption capacity is still greater that at pH 6.5 but lower than in the presence of HCl. The reversibility of the adsorption process was tested in two ways. In the first method, 80% of the supernatant above the BaSO, (with 7 M SCMC adsorbed to saturation coverage) was removed and replaced with an equal volume of the pure solvent and then left to equilibrate for 24 h. The procedure was carried out for adsorption from water, 0.1 mol dm-3 NaCl and 0.01 mol dm-3 HCl. In the second method, 7 M SCMC was adsorbed to saturation coverage from 0.1 mol dm-3 NaCl and 0.01 dma3 HCl and then the unadsorbed polymer removed by washing with the same solvent.The solvent was then poured off, replaced with water and the system equilibrated for 24 h. Both methods showed that there was no desorption. VISCOSITY In fig. 5 the reduced viscosity (vsP/c) is plotted against concentration for the three samples of SCMC at pH 7 and pH 2. At pH 2 the results accord with the behaviour of a neutral polymer and the limiting viscosity numbers may be obtained by extrapolating to zero concentration. At pH 7 SCMC exhibits the typical behaviour of a polyelectrolyte, for which the reduced viscosity is related to concentration by the empirical equation qsp/c = A / ( l +Be$)P. A. WILLIAMS, R. HARROP, G. 0. PHILLIPS, G. PASS A N D I. D. ROBB 1737 ,,i ,,i 30t * O t 0 0 0 - - w pH 7.0 ~ 0 0 02 0.04 0.06 0.0 8 0.10 concentration/g ~ r n - ~ FIG.5.-Reduced viscosity of SCMC as a function of concentration: 0, 4 M SCMC; 0, 7 M SCMC; 0 , 12M SCMC. m E! I / FIG. 10' 1 I I L I I I I 0 005 0.10 0.15 0.20 0.25 0.30 0-35 0.40 (concentration/ 102 g cm-3)' 6.-Reciprocal reduced viscosity as a function of the square root of concentration: 0 , 4 M SCMC; 0, 7 M SCMC; a, 12 M SCMC.1738 SODIUM CARBOXYMETHYL CELLULOSE ON BaSO, where A and B are constants.lg When the reciprocal of the reduced viscosity is plotted against ca linear plots are obtained (fig. 6), from which the limiting viscosity number may be found from the reciprocal of the intercept at zero concentration. ELECTROPHORETIC MOBILITY Fig. 7 shows the electrophoretic mobility and zeta potential of the barium sulphate particles as a function pH.At pH 6.5 the particles are negatively charged (zeta potential - 10.3 mV) and at pH 2 positively charged (zeta potential + 10.3 mV). The isoelectric point occurs at pH 3.8. FIG. 7.-Electrophoretic mobility and zeta potential of BaSO, particles as a function of pH: a, no adsorbate; 0, with adsorbed 4 M SCMC; 0, with adsorbed 7 M SCMC; (>, with adsorbed 12M SCMC. Fig. 7 also illustrated the electrophoretic mobility of the barium sulphate particles after adsorption of the three samples of SCMC at adsorption capacity. Between pH 6 and pH 9, when the polyanions are highly dissociated,20 the zeta potential reaches a steady value, for all three samples, related to the degree of substitution.At pH values < 5 the zeta potentials increase rapidly and converge to a common positive value below pH 3. DISCUSSION At pH 6.5 the three SCMC samples are extensively dissociated into sodium ions and polyanionslO and consequently exhibit typical polyelectrolyte behaviour as indicated by the viscosity data. The adsorption capacity is found to increase as the charge density of the polyanions decreases, and it is likely that the predominant factor giving rise to this observation is the build-up of charge near the surface due toP. A. WILLIAMS, R . HARROP, G. 0. PHILLIPS, G. PASS AND I. D. ROBB 1739 unadsorbed carboxylate ions. For 7 M SCMC in water at pH 6.5 the area available to each pyranose unit is 1.7 nm2 (as calculated for the adsorption capacity), but since each of these units has an area of only 0.57 nm2 (from molecular models) then only about one third of the surface can be occupied even if the polymer is assumed to adsorb with a completely flattened configuration.Similar results were obtained by Ishikawa,21 who studied the adsorption of quaternary poly(4-vinyl pyridine) onto negatively charged polyvinyl acetate lattices. The Gouy-Chapman relationship applies only when the diffuse region of the double layer is not significantly distorted. Although severe distortion may take place in the case of adsorbed non-ionic polymers5 owing to the large number of loops and tails extending into solution, it would not be significant for polyelectrolytes, which are thought to adsorb predominantly as trains such that the adsorbed layer thickness (d) has the dimensions of the polymer i t ~ e l f .~ ~ * ~ ~ Hence, assuming that d is between 0.75 and 1.5 nm, t+vs can be calculated from the above relationship, and if the relative permittivity of the electrical double layer is taken as 10, approximate values of o0 can be calculated. For 7 M SCMC, for example, the values will be somewhere between 0.01 17 and 0.00425 C m-2. From a knowledge of the number of carboxy groups adsorbed and the surface charge density it is evident that a high proportion of carboxy groups interact with the surface (between 70 and 90%), the effective charge on the particle being reduced by the desorption of SO:- ions.23 Although addition of electrolyte weakens the segment-surface binding energy, more polyelectrolyte is adsorbed at higher ionic strengths.This is likely to be a result of lower (electrostatic) repulsion between the segments of the adsorbed phase. More polymer generally adsorbs from poorer solvents and in the case of polyelectrolytes in water, the electrostatic repulsion between segments is the main contributor to the osmotic pressure. The free-energy change for a segment adsorbing at the surface will predominantly depend on interactions between segments in the adsorbed layer, where the segment density is high, rather than on those in the solution phase. For 7 M SCMC the surface area available for each segment is reduced to 0.54 nm2 in 10-1 mol dm-3 NaCl. The nature of the electrolyte added is also important, since the adsorption capacity was found to be greater in the presence of sodium chloride than sodium sulphate at the same ionic strength; this was thought to be due to the fact that sulphate ions compete to a greater extent with the polymer segments for surface sites.So far only the electrostatic effects arising from the negative charge on the polyanions have been considered. However, at pH 6.5 the electrophoretic mobility measurements show that the barium sulphate itself is also negatively charged, although the charge density is too low for the repulsion between the surface and polyanion to prevent adsorption. The lowering of the solution pH from 6.5 to 2 will have a significant effect on the properties of both the SCMC and the BaSO,. The viscosity measurements (fig. 5 ) indicate that the three SCMC samples behave as neutral polymers; this is expected, since at pH 2 virtually all the carboxy groups would be undissociated.This is reflected by the electrophoretic mobility results, which show that both coated and uncoated BaSO, particles are positively charged. The increased adsorption capacity at this lower pH is thought to result from the decrease in the intermolecular repulsion between segments in the adsorbed phase, which was thought to be the limiting factor at pH 6.5.1740 SODIUM CARBOXYMETHYL CELLULOSE ON BaSO, In addition the reduction in the solvency of the polymer molecules must now play an important role. The area available for an adsorbed segment at pH 2 (in HC1) is now only 0.41 nm2, and therefore the molecules must adsorb with a ‘loopier’ configuration, with loops and tails protruding out into the solution, which is typical of non-ionic polymers.The decreased adsorption of the SCMC samples in the presence of SO:- ions as compared to CI- ions at pH 2 suggests that there is still competition between SO:- ions and polymer segments for the surface sites. The low-affinity isotherms obtained at pH 2 may be due to fra~tionation,~~ although further experiments are necessary to elucidate this point. Yu S. Lipatov and L. M. Sergeeva, Adsorption of Polymers (Wiley, New York, 1974). T. Sat0 and R. Ruch, Stabilization of Colloidal Dispersions by Polymer Adsorption (Marcel Dekker, New York, 1980). S. G. Ash, Polymer Adsorption at the SolidlLiquid Interface, in Colloid Science (Specialist Periodical Report, The Chemical Society, London, 1973), vol.1, chap. 3. P. F. Wilds and R. Dexter, Br. Polym. J., 1972, 4, 239. J. A. Kitchener, Br. Polym. J., 1972, 4, 217. R. W. Slater and J. A. Kitchener, Discuss. Faraday Soc., 1966, 42, 267. R. L. Whistler, Industrial Gums (Academic Press, New York, 1973). M. Glicksman, Gum Technology in the Food Industry (Academic Press, New York, 1969). lo A. M. James and G. H. Goddard, Pharm. Acta Helv., 1972, 47, 244. A. Silberberg, J. Chem. Phys., 1968, 48, 2835. l2 J. M. H. M. Scheutjens and G. J. Fleer, J. Phys. Chem., 1979, 83, 12; 1619. l3 F. T. Hesselink, J. Colloid Interface Sci., 1977, 60, 3 ; 448. l4 B. J. Fontana, in Adsorption of Biological Analog Molecules on Non-biological Surfaces in the l5 A. Takahashi, M. Kawaguchi and T. Kato, Polymer Science and Technology, ed. L. H. Lee (Pergamon l6 R. Buscall, J. Chem. SOC., Faraday Trans. I , 1981, 77, 909. l7 G. P. Diakun, H. E. Edwards, J. C. Allen, G. 0. Phillips and R. B. Cundall, Anal. Biochem., 1979, l8 H. D. Graham and G. Mitchell, J. Food Sci., 1963, 28, 546. 2o R. Sarkar, M.Sc. Thesis (University of Salford, 1974). *’ M. Ishikawa, J. Colloid Interface Sci., 1976, 56, 3; 506. 22 F. R. Eirich, J. Colloid Interface Sci., 1977, 58, 4; 432. 23 M. C. Cafe and I. D. Robb, J. Colloid Interface Sci., in press. 24 R. E. Fetter, J. Polym. Sci. Lett., 1974, 12, 147. * R. A. Ruehrwein and D. W. Ward, Soil Sci., 1952, 73, 485. Chemistry of Biosurfaces, ed. M. L. Hair (Marcel Dekker, New York, 1971). Press, New York, 1980, vol. 12, part B. 94, 378. R. M. Fuoss and U. P. Straws, J. Polym. Sci., 1948, 3, 246; 602. (PAPER 1 /797)

ISSN:0300-9599    

DOI:10.1039/F19827801733

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 1741-1753 Exchange of Ammonium and Sodium Ions in Synthetic Faujasites BY PHILIP FLETCHER AND RODNEY P. TOWNSEND* Department of Chemistry, The City University, Northampton Square, London EClV OHB Received 19th May, 1981 Ion-exchange isotherms for the Na $ NH, equilibrium in zeolites X and Y were constructed at 298 K and a total solution normality of 0.1 g-equiv . dmP3. Values for the standard free energy of exchange were calculated and shown to be consistent with the affinity sequence predicted using dielectric theory. Comparisons are made between previously reported studies on the NaGNH, exchange in X and Y and this current work, and the differences in experimental data are discussed. Statistical thermodynamic considerations provide a greater understanding of the Na s NH, exchange equilibrium in X and Y, and collectively the evidence suggests that high framework charge and ion-sieving factors in X are the dominant properties determining exchange selectivity and maximum exchange levels in this zeolite.In Y, the ion-sieve effect and repulsion between the entering ammonium ions are the dominant factors. Studies of the exchange of ammonium ions into the sodium forms of the synthetic zeolites X and Y have been undertaken by various workers in the past, ever since Barrer et a!.' reported a 92 % replacement of sodium by ammonium ion in X using 2 mol dm-3 ammonium chloride solution. Sherry2 first published an isotherm for the Na-NH,/Y equilibrium in 1966, and Theng et al. obtained isotherms for this exchange in both X and Y two years later.3 In his review of ion exchange in zeolites Sherry4 included a Na-NH,/X isotherm which was not found in his earlier paper,2 and since then, as part of a wider study of Szilard-Chalmers recoil studies in X and Y, Lai and Rees5 published a further isotherm for the NH, Na equilibrium in Y.Finally, Herman and Bulko include some data on the effect of temperature on maximal levels of exchange of ammonium ion in Y in a recent study on copper zeolites.6 Apart from the work of Theng et al., none of the above studies was intended to be concerned primarily with the sodium-ammonium exchange equilibrium in X and Y. Most of the studies were, however, carried out under similar conditions, and it is therefore surprising that the results are not in good agreement, especially in the case of X.As part of an ongoing study of transition metal exchange in ammonium and sodium zeolite^,^ l 1 it became necessary to re-examine the sodium-ammonium equilibrium in the faujasite-type zeolites, and thence to compare and contrast these data with previously reported information. The results of this examination are reported in this paper. EXPERIMENTAL MATERIALS Synthetic sodium zeolites were supplied by Union Carbide. Chemicals used, whether for analysis or exchange purposes, were of AnalaR grade. 17411742 ION EXCHANGE IN SYNTHETIC FAUJASITES ANALYSES (i) Zeolite phase: samples were analysed for Na,O, SiO,, Al,03, Fe,O,, (NH,),O and H,O content by methods outlined previously in the literature.’.(ii) Solution phase: sodium was determined in solution using either atomic absorption spectroscopy or flame photometry. Ammonia was determined by a modified Kjeldahl method described previo~sly.~. * PREPARATION OF EXCHANGED FORMS OF ZEOLITES Initially all zeolite samples were treated with solutions containing 0.5 mol dm-3 sodium chloride in order to ensure that starting materials were in the homoionic sodium forms. Extensive washing of these samples thereafter was avoided in order to minimise hydrolysis of the zeolites, especially X. For the isotherms (which were constructed at a total solution normality of 0.1 g-equiv . dmP3) maximum levels of exchange of ammonium for sodium in X and Y were obtained by exhaust- ively exchanging 0.2 g samples of the sodium zeolite with 50 cm3 aliquots of 0.1 mol dm-3 ammonium nitrate solution at 298 K.The resulting maximum exchange levels agreed with those indicated by the general isotherm shapes (see Results). Other experiments on the maximum exchange level of ammonium in X and Y were also conducted usingsolutions ofammonium chloride ofdifferent concentrations at room temperature or with 0.5 mol dm-3 ammonium chloride solution over a range of temperatures. In agreement with Herman and Bulko,6 the maximum exchange level in Y was found to be sensitive to temperature. The temperature sensitivity of X was much less marked. These results are recorded in this paper but also discussed el~ewhere.~. l2 EQUILIBRIUM STUDIES All equilibria were constructed at a constant total normality of 0.1 g-equiv .dm-3 using ammonium and sodium nitrate solutions containing different ratios of the two cations, and at a temperature of 298 K. Forward isotherm points were constructed by conventional methods described previ~usly.~ Reverse isotherm points were constructed by a modified method recently outlined9-l1 in order to avoid problems arising from ion-site redistributions on drying the zeolite. THERMODYNAMIC TREATMENT OF DATA Values of the standard free energy of exchange were calculated by the usual methods, described in detail re~ent1y.I~ Important aspects, summarised here, are the exchange equation and thermodynamic equilibrium constant K,, uiz. NH:(s)+Na+(c)sNH,+(c)+Na+(s) (1) and K a = ( Q N H ~ ( ~ ) ~ ~ a ( s ) / a ~ ~ ~ ( s ) aNa(c)) (2) where a is the activity, and (c) and (s) refer to crystal and solution phases, respectively.Values of K, were determined using the procedure of Gaines and Thomas’, from (3) where (4) and ion r is (NH,): and (Na): are the n~rmalised~~’ l5 equivalent fractions of ammonium and sodium in the zeolites. mNa and mNHl are the concentrations (mol dm-3) of the ions in solution. the ratio of corresponding single-ion activity coefficients in solution, which may be evaluated using Glueckauf’s method.16 In practice, for the Na e NH, system, the ratio is close to unity, and it is also (for uni-univalent systems) independent of crystal phase composition,1° so this correction is near negligible.P. FLETCHER AND R. P. TOWNSEND 1743 Finally, application of the Gibbs-Duhem equation enables the calculation of the crystal phase activity coefficients fNH4 and fNa using the relations14 In fNH4 = - In KF + (NH4); In KF + { In KF d(NH,): ( 5 ) ( N H ~ ) ? RESULTS Analytical data for samples of X and Y before and after exhaustive exchange of the sodium forms with 0.5 mol dm-3 ammonium chloride solution at 298 K are shown in table 1.It is apparent that 100% exchange of sodium by ammonium ion was not TABLE 1 .-CHEMICAL ANALYSES OF ZEOLITES sample unit-cell composition Na-X Na84.6(A102)84. 7(si02)107.3 .25 .7H20 Na-Y Na6~.3(A102)61.4(Si02)130.6242 * 5H20 NH,/Na-X (NH4)61.7Na22.9(A102)84.8(si02)107.2 ' 236H20 NH,/Na-Y (NH4)43Nai8.~(A102)6i.4(si02)*~~.6 * 219.4H2O sample oxide formula Na-X Na-Y NH4/Na-X NH,/Na-Y Na20. A1203 - 2.53SiO2. 5.94H20 Na20. A120, * 4.25SiO2 0.002Fe203 * 7.89H20 ((NH,)20)o,73(Na20)o~27 * Al,O,. 2.53Si0, * 5.57H20 ((NH,)20)o~7,(Na20)o~30 * A12034 25Si0, * 0.002Fe203 - 7. 14H20 achieved under these conditions, the maximum levels of exchange being 73 and 70%, respectively.The effects of both temperature and concentration of ammonium salt on the maximal level of exchange in X and Y are shown in table 2. Ion-exchange isotherms for the Na c NH, equilibria in X and Y are shown in fig. 1. Both systems were reversible. Plots of In K , against (NH,), are shown in fig. 2 and 3, and in fig. 4 are the corresponding crystal-phase activity coefficients, which were calculated using eqn ( 5 ) and (6). Values for the standard free energies of exchange and the thermodynamic equilibrium constant Ka are given in table 3.DISCUSSION GENERAL COMMENTS The exchange isotherm for the Na + NH, exchange in Y obtained in this work may be compared with others in the 3 7 The isotherm shape in fig. 1 conforms to type d of Breck's classification,17 and agrees closely with those found by Sherry2 and Theng et aZ.3 The AG* values are similar (table 3); indeed, the small differences observed are probably the result of the different silica to alumina ratios found in the1744 ION EXCHANGE I N SYNTHETIC FAUJASITES TABLE 2.-EFTECT OF TEMPERATURE AND SALT SOLUTION CONCENTRATION ON DEGREE OF AMMONIUM EXCHANGE percentage sodium removeda /K /mol dmP3 Na-X Na-Y temperature “H4ClI 298 298 313 333 353 298 298 298 298 298 298 298 0.1 0.5 0.5 0.5 0.5 1 .o 1.5 2.0 3.0 4.0 5.0 saturated solution 70.0 72.7 73.4 73.4 73.3 73.1 93.0 93.1 92.9 93.2 93.0 93.0 70.1 70.1 76.3 85.4 93.1 70.1 70.3 70.3 69.9 69.8 70.0 70.1 a For all experiments 0.2 g of zeolite were exchanged five times for 24 h with fresh 50 cm3 aliquots of the appropriate ammonium chloride solution.(NH,), FIG. 1 .-Ion-exchange isotherms (not normalised) for the Na $ NH, exchange in (a) X (b) Y: 0, forward points; x , reverse points; 0, direct analysis of exhaustively exchanged samples.P. FLETCHER AND R. P. TOWNSEND FIG. 2.-Normalised Kielland plot of N a e N H , exchange in X. 1745 I 0.0 0.2 0.4 0.6 0.8 1.0 (NH.9): FIG. 3.-Normalised Kielland plot for Na e NH, exchange in Y. Y samples used by the different workers. (This point is discussed further below.) The observed maximum levels of exchange are also similar, uiz.68,2 703 and 70% (this work). In contrast, the isotherm given by Lai and Rees5 is different, appearing sigmoidal in shape (i.e. type b in clas~ification),~~ and the maximum exchange level appears higher. However, their Y zeolite5 had a substantially higher aluminium content (68 aluminium atoms per unit cell) than those used here or in the other 57 FAR I1746 ION EXCHANGE IN SYNTHETIC FAUJASITES 1 C C f 0 0 0 0.2 0.4 0.6 0.8 (NH4),N FIG. 4.-Plots of the phenomenenological activity coefficients f for the ions NH: and Na+ in X and Y. Data are plotted in terms of the normalised equivalent fraction of ammonium ion in the crystal phase (NH,)?. TABLE 3.-cOMPARISON OF THERMODYNAMIC DATA FOR THE NH,,Na EXCHANGE IN DIFFERENT ZEOLITES SiO, : A1,0, zeolite ratio AGe/kJ rno1-I ref.temp./K ref. X X Y Y Y mordenite chabazite clinoptilolite clinoptilolite 2.52 2.53 5.33 4.98 4.25 10.53 5.13 10.00 10.00 - 2.801 - 0.534 - 2.759 - 2.592 - 2.005 - 3.760 -4.138 - 5.392 - 5.727 293 298 298 293 298 298 298 303 333P. FLETCHER A N D R. P. TOWNSEND 1747 studies,2? which may account for the different isotherm shape (see comments below on X). Regarding the maximum exchange level, Lai and Rees extrapolated their isotherm from < 80 to loo%, and in general used different conditions for their experiments to the other studies* since, as the authors themselves state, ‘thermo- dynamic analysis. . .was not the prime purpose of the exchange measurements, no attempt was made to cover the complete range of A , with the thoroughness such an analysis would warrant’.5 Finally, further data on the maximum exchange level for ammonium in Y are found in a recent paper by Herman and Bulko.6 Using a Y sample containing 56 aluminium atoms per unit cell, they found a maximum level of exchange of 75 % after three equilibrations at 298 K, and an exchange level between 68 and 7 1 % after one equilibration.These data are in good agreement with this work and other In contrast to Y, agreement over the sodium-ammonium exchange in X between different workers is very poor. There are less data on X in the literature; the only studies with which this work can be compared are those of Sherry4 and Theng et aL3 Sherry’s isotherm4 is type b, in common with fig. 1. However, in this present study a maximum exchange level of 70% was obtained at a solution concentration of 0.1 mol dme3, yet Sherry extrapolates his isotherm from ca.85 to 100%. In the absence of comment in the paper4 it is not possible to ascertain if any markedly different experimental conditions were used which may explain this discrepancy. Sherry does not publish a value for the free energy of exchange of ammonium into sodium X. Theng et aL3 calculated AG* for this exchange and obtained a value very much higher than that found here (see table 3). In addition, their isotherm is type d, and differs very markedly from both Sherry’s4 and this work, with a maximum exchange level of only 63%. Although Theng et aL3 determined their isotherm at a total solution molarity of 0.05 mol dmP3, whereas the solution concentration in both Sherry’s study4 and this present work was double this, this difference cannot be the explanation for the observed3 different isotherm shape, as Barrer and Klinowski have shown that for uni-univalent exchanges the isotherm shape is independent of the external solution concentration.1 9 9 2o CONSIDERATIONS OF ION SIZES A N D DISTRIBUTIONS The potassium ion (Pauling radius 0.133 nm) can completely replace sodium in both X4 and Y.2 In contrast,2v4 neither rubidium nor caesium (Pauling radii 0.148 and 0.169 nm, respectively) exchange to 100% in either zeolite. The effective diameter of the six-oxygen windows (which lead into the sodalite units) is quoted3 as being between 0.266 and 0.288 nm. It is apparent on comparing the relative sizes of the potassium, rubidium and caesium ions with the six-oxygen window diameters in X and Y that the above observations are readily explained.The ionic radius of the ammonium ionz1 is 0.143 nm, which lies between those of potassium and rubidium; its diameter is in fact almost identical to the window size. It is therefore probable that this accounts at least partially for the variability in maximum exchange level that is observed when replacing sodium by ammonium ions in X. It is not obvious however why this variability should be much greater in X than in Y if ion sieving is the only important consideration. The structure of X is only marginally more open than Y,22 and in fact there is now much evidence that in the case of Y an exchange level which is > 70% does not necessarily imply that the ingoing ions have even removed all the original ions from the super cage^.^^^^^ Thus, for example, a potassium Y sample exchanged with ammonium ions to a level of 72% * Samples were pretreated with saturated ammonium chloride solution in order to remove sodium.1s 57-21748 ION EXCHANGE IN SYNTHETIC FAUJASITES still contained seven potassium ions per unit cell in the type I1 More recently, CremersZ5 has given collected evidence for ion redistributions taking place during exchange, and Vansant and Uytterh~even~~ agree with Theng et aL3 in emphasising that ‘the maximum limit to exchange was determined by the same factors as the exchange ~electivity’.~~ They also note, in common with experimental observations made here, that the maximum exchange level obtained varied more with temperature in the faujasites with the higher silica: alumina ratios.Thus, in this present work (table 2) it was found that the maximum level of exchange of ammonium for sodium in Y was ca. 90% at 353 K using similar solution concentrations to Herman and Bulko.6 In contrast, raising the temperature of exchange with X had no effect on the observed maximum exchange level of 73% (table 2). It is apparent therefore that differences in framework charge density between X and Y zeolites are a primary factor affecting the Na and NH, exchange levels and selectivity in these zeolites. Framework charge densities can be considered using simple dielectric theory, or from a statistical thermodynamic viewpoint. APPLICATION OF DIELECTRIC THEORY Application of simple dielectric theory leads to the expressionl19 2 6 t 27 where rNHl and rNa are the ionic radii of ions NH: and Na+, E,, E, are the permittivities of the crystal and solution phases, respectively, e is the charge on the electron and N is Avogadro’s constant. Since26 E , < E,, and rNH4 > rNa, eqn (8) leads to the prediction that the standard free energy of exchange for the exchange reaction given in eqn (1) should be negative for all zeolites.This prediction is borne out by experimental data in the literature for a range of zeolites (table 3). In addition, it has been shownll that simple dielectric theory leads to the conclusion that if the same exchange is observed in two zeolites of differing framework charge where the subscripts ‘ hc’ and ‘lc’ refer to ‘high charge’ and ‘low charge’, respectively.Thus, if (as is the case here) dielectric theory predicts that AG* should be negative for the exchange of ammonium in sodium zeolites, then the theory also implies that AG* should be less negative for the zeolite of higher framework charge. This prediction is also confirmed from a plot of AG* values for the Na + NH, exchange against (Al), (fig. 5), where NA1 = Nsi + NA1 and NA1 and Nsi are the number of aluminium and silicon atoms, respectively, per unit cell of zeolite. The trend observed in fig. 5 involves data from zeolites of very different structures, and exemplifies the comments of Barrer and Davies2’ on simple dielectric theory, viz. ‘this model is simple in principle although possibly only semi-quantitative in application.It is however the most useful because no information is required about the geometric arrangement of cations. ’ The broad agreement between the predictions of simple dielectric theory and the plot in fig. 5 thus emphasises the significant contribution that the framework charge makes to the magnitude of AG*, irrespective of the different framework structures. Nevertheless, the overall significance of the trend observed in fig. 5 should be assessed in the light of the following points.P. FLETCHER AND R. P. TOWNSEND 1749 X X \ \V\ \ \ \ \ '.A 0 ? \ \ \ \ 0 \ FIG. 5.- Values of AG* for the Na S NH, exchange reaction in different zeolites plotted as a function of the equivalent fraction of aluminium in the framework (Al)c: x , clinoptilolite [ref.(30) and (31)]; V, synthetic mordenite [ref. (28)]; a, chabazite [ref. (29)]; A, zeolite Y [ref. (2), (3) and this work]; 0, zeolite X [ref. (3)]; 8, zeolite X (this work); (-------) best fit of results for all the synthetic zeolites excluding the X data. (i) Except for the value of A G e obtained for X by Theng et al., all the data for synthetic species follow a very clear trend. AG* decreases almost linearly with increasing aluminium content in the framework from synthetic mordenite (-AG* = 3.76 kJ mol-l) to X(-AG* = 0.53 kJ mol-l). (ii) Although all the measured AG* values for the natural zeolites are negative in agreement with prediction [see inequality (9)], the data lie well off the line followed by the synthetic species. Allowance can be made for the fact that the A G e values for c l i n ~ p t i l o l i t e ~ ~ ~ ~ ~ were measured at 303 and 333 K using the AH* value given by Howery and but the correction does not alter significantly the magnitude of AGe.The problems involved in purifying natural zeolite^,^ the care necessary in order to obtain homoionic forms of (for example) ~linoptilolite~~ and especially the ammonium form7133 make data for natural zeolites difficult to compare with those of the synthetic materials. In the context of these reservations, and considering therefore only the free energy values obtained for synthetic species in 5g. 5, it does appear that the value obtained for the N a e N H , exchange in X by Theng et aIq3 is rather high. STATISTICAL-THERMODYNAMIC CONSIDERATIONS Using statistical thermodynamics, Barrer and Klinowski3, have shown that the activity coefficients of the exchanging ions in the crystal phase are related to the framework charge density :1750 ION EXCHANGE IN SYNTHETIC FAUJASITES COAA is an excess interaction energy which arises when pairing of the entering A': ions takes place.q is related to the framework charge density by where N is the number of ion sites in the crystal and No the number of charges (i.e. equivalent to the number of framework aluminium atoms). Eqn (1 1) and (12) were derived assuming a random siting of cations.34 For uni-univalent exchange, such as the ammonium-sodium exchanges considered here, eqn (1 1) and (1 2) reduce to and Barrer and Klinow~ki~~ demonstrated that for this very simple case, eqn (14) give a theoretical basis to the original semi-empirical approach of Kielland.35 If eqn (14) are applicable to the Na + NH, exchanges in X and Y, inspection of eqn (14) shows that fNa and fNH, [which were phenomenologically determined using eqn (5) and (6)] should vary with (NH,), and (Na), in a symmetrical manner, a criterion which is met well with Y (fig.4). Applying eqn (14) to the Na$NH4 exchange in (for example) X gives P(NH,),N - 11 (15) ~ A A In (fNH,(X)/fNa(X)) = ___ VXkT when the subscript X refers to zeolite X. A similar expression holds for Y. Thus plots of ln(fNH,/fNa) against (NH,), should yield plots with intercepts equal to -(wAA./qkT), and gradients twice these values and of opposite sign. Using the data shown in fig. 4, plots of eqn (15) for both X and Y were constructed, and are shown in fig.6. The data conform well with the simple statistical-thermodynamic model3, in the case of Y, but the fit for X is much poorer. Correlation coefficients on the linear best fits to the data are found to be = 0.997 and R(x) .= 0.945, respectively, where (as usual) R = o,m/oy, with ox and oy the standard deviations ( N weighting) of the dependent and independent variables, respectively, and rn is the gradient of best fit. Comparing either the best fit gradients or the intercepts for X and Y in fig. 6 yields (16) a very similar result, uiz. Since X and Y are isostructural, it follows from eqn (13) and the analytical data in qXwAA(X) = o*4qY0AA(Y). table 1 that qy/rx = Nx/Ny = 1.435 so that wAA(X) = o*279wAA(Y) (18) where COAA(X) and wAA(y) are the interaction energies involved when the entering ammonium ions pair in X and Y, respectively.Examination of fig. 6 shows that both COAA(X) and wAA(y) are positive. To evaluate these, it is necessary to know qx and qy, which requires a knowledge of N [eqn (1 3)]. N is difficult to estimate, but since many of the ions in the supercages are unsited2*24 it seems likely that N c No for both zeolites. Values of wAA(X) and w A A ( ~ ) determined on this assumption are shown in table 4.P. FLETCHER AND R. P. TOWNSEND 1751 0.0 0.2 0.4 0.6 0.8 (NH,): FIG. 6.-Plots of the crystal-phase activity correction factor as a function of (NH,),N for zeolites X and Y. TABLE 4.-vALUES OF WAA, THE INTERACTION ENERGY BETWEEN AMMONIUM IONS, CALCULATED FOR DIFFERENT VALUES 0.25 0.100 248 0.358 887 0.50 0.200 496 0.717 1777 0.75 0.300 743 1.075 2664 1 .oo 0.400 99 1 1.434 35531752 ION EXCHANGE IN SYNTHETIC FAUJASITES Several comments are made on these data.(i) Barrer and Klinowski calculated34 theoretical isotherm plots for values of K, > 1 when o,,/kT was small and positive. These compare well with the experimental data found here {compare normalised plots of X and Y in fig. 1 (this work) with fig. 3(c) [l-41 for X and fig. 3(a) [2] for Y in ref. (34)). (ii) When coAA/kT > 0, the entering cations (ammonium in this case) avoid each The higher the value of oAA/kT, the stronger is this repulsion. Thus in Y, the ammonium ions avoid each other more strongly than in X, [Note that wAA for a given ion pair is not an intrinsic property of that ion pair, and therefore may vary from zeolite to zeolite. COAA is an excess energy function and refers to prescribed reference (in this case Na-X and Na-Y, respectively).It is therefore quite consistent that COAA for the same ion pair can be markedly different in different zeolites.] (iii) From the comments under point (ii) it is clear that the COAA values are consistent with the view that in X the framework charge is a dominant factor determining ammonium ion selectivity and sitings, whereas in Y the interactions between the ammonium ions themselves are far more important. This is consistent with Sherry’s view.2 (iv) It follows from point (iii) that ions in the supercages of Y are more randomly distributed than in X, since the important factor in Y is not ion binding to the framework2 but rather, as shown above, the tendency of the ammonium ions to avoid one another.Since eqn (1 1) and (12) were derived assuming a random distribution of cations, it follgws that the data for Y should fit eqn (15) much better than those for X. This is indeed the case (fig. 6 ) . CONCLUSIONS Experimental data and theoretical considerations combine to show that the ammonium ion is reluctant to enter the sodalite units in X and Y, and that the ammonium ions are more mobile and less strongly bound in Y than in X. It is probable therefore that the observed increase in maximum exchange level of ammonium ion with temperature in Y occurs by an initial migration of sodium ions out of both sites I and the sodalite units into the supercages, which is then followed by exchange, the ammonium ions then remaining in the supercages.In X, the sodium ions are less mobile and are bound much more strongly to the zeolite framework. This decrease in mobility, together with the reluctance of the ammonium ions to enter the sodalite units, makes the maximum exchange level for ammonium ion in X less sensitive to temperature changes. Dielectric theory provides a rationale for the observation that Y is more selective than X for the ammonium ion, and statistical-thermodynamic theory confirms that the ammonium ions bind less strongly to the framework in Y than in X. P. F. gratefully acknowledges a Scholarship from the British Gas Corporation, and subsequently a Research Fellowship from the City University.R. M. Barrer, W. Buser and W. F. Grutter, Helv. Chim. Acta, 1956, 29, 518. H. S. Sherry, J . Phys. Chem., 1966, 70, 1158. B. K. G. Theng, E. Vansant and J. B. Uytterhoeven, Trans. Faraday Soc., 1968, 64, 3370. H. S. Sherry, in Zon Exchange-A Series of Advances, ed. J. Marinsky (Marcel Dekker, New York, 1969), vol. 2, p. 89. P. P. Lai and L. V. C. Rees, J . Chem. Soc., Faraday Trans. I , 1976, 72, 1809.P. FLETCHER AND R. P. TOWNSEND I753 fi R. G . Herman and J. B. Bulko, in Adsorption and Ion Exchange with Synthetic Zeolites, ed. W. H. Flank, (ACS, Washington, D.C., 1980), vol. 135, p. 177. R. M. Barrer and R. P. Townsend, J. Chem. Soc., Faraday Trans. 1, 1976, 72, 661, 2650. R. M. Barrer and R. P. Townsend, J.Chem. Soc., Faraday Trans. 1, 1978, 74, 745. P. Fletcher and R. P. Townsend, Proc. 5th Int. Con$ Molecular Sieves, Naples, 1980, ed. L. V. C. Rees, (Heyden, London, 1980), p. 3 1 1. P. Fletcher and R. P. Townsend, J . Chromatogr.. 1980, 201, 93. * l P. Fletcher and R. P. Townsend, J. Chem. Soc., Faraday Trans. 1. 1981, 77, 497. l 2 P. Fletcher and R. P. Townsend, ‘Transition Metal Ion Exchange in Zeolites. Part 5’, paper in preparation. A. Dyer, H. Enamy and R. P. Townsend, Sep. Sci. Technol., 1981, 16, 173. R. M. Barrer, J. Klinowski and H. S. Sherry, J . Chem. Soc., Faraday Trans. 2, 1973, 69. 1669. D. W. Breck, Zeolite Molecular Sieres (Wiley-Interscience, London, 1974), p. 532. l 4 G. L. Gaines and H. C. Thomas, J . Chem. Phys., 1953, 21, 714. Ifi E. Glueckauf, Nature (London), 1949, 163, 414. IH L. V. C. Rees, personal communication. IY R. M. Barrer and J. Klinowski, J . Chem. Soc., Faraday Trans. I , 1974, 70, 2080. J. Klinowski, in The Properties and Applications of Zeolites, ed. R. P. Townsend (Special Publication, The Chemical Society, London, 1980), vol. 33, p. 288. 2 1 F. A. Cotton and G. Wilkinson, Adcanced Inorganic Chemistry (Interscience, New York, 1966), Lz D. W. Breck, Zeolite Molecular Sieces (Wiley-Interscience, London, 1974), pp. 176 and 177. 2:1 E. F. Vansant and J. B. Uytterhoeven, Adv. Chem. Ser., 1971, 101, 726. 24 W. J. Mortier, M. L. Costenoble and J. B. Uytterhoeven, J. Phys. Chem., 1973, 77, 2880. 25 A. Cremers ACS Symp. Ser., 1977, 40, 185. 2fi R. M. Barrer, L. V. C. Rees and M. Shamsuzzoha, J . Inorg. Nucl. Chem., 1966, 28, 629. 2 i R. M. Barrer and J. A. Davies, J . Phys. Chem. Solids, 1969, 30, 1921. 2x R. M. Barrer and J. Klinowski, J. Chem. Soc., Faraday Trans. I , 1974, 70, 2362. ey R. M. Barrer, J. A. Davies and L. V. C. Rees, J . Inorg. Nucl. Chem., 1969, 31, 219. :H’ D. G. Howery and H. C. Thomas, J . Phys. Chem., 1965, 69, 531. 31 R. M. Barrer, R. Papadopoulos and L. V. C. Rees, J . Inorg. Nucl. Chem., 1967, 29, 2047. 32 A. Ardya and A. Dyer, J . Inorg. Nucl. Chem., 1981, 43, 589. 3 3 M. Loizidou and R. P. Townsend, paper in preparation. 3 4 R. M. Barrer and J. Klinowski, Philos. Trans. R . SOC. London, 1977, 285, 637. p. 334. J. Kielland, J. Soc. Chem. Ind.. 1935, 54, 232T. (PAPER 1 /807)

ISSN:0300-9599    

DOI:10.1039/F19827801741

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1755-1765 Calorimetric Investigations on Association in Ternary Systems Part 2.-Thermodynamic Functions of Complex Formation of Pentahalogenobenzoic Acids and 1,l -Dinitroethane with Proton Acceptors BY STEFANIA TANIEWSKA-OSINSKA* AND PAWEL GORALSKI Department of Physical Chemistry, University of Lodi, 91 -416 Lodz, Nowotki 18, Poland Received 12th June, 1981 Enthalpies of complex formation (AH) by hydrogen bonding of pentachlorobenzoic (PClHBz) and pentafluorobenzoic (PFHBz) acids with a number of proton acceptors in CCl, have been determined. A calorimetric method of determining equilibrium constants and enthalpies of formation of hydrogen- bonded complexes which are in equilibrium with proton transfer complexes is described. The K and AH values of complex formation of a C-H acid (1,l-dinitroethane) with a number of proton acceptors are determined.The results obtained both by the calorimetric method and by infrared spectroscopy are compared with those obtained by the use of the i.r. method only. The AH values of complex formation are discussed as a function of their acidity, dipole moments and the Gutmann’s donor numbers of the investigated substances. The study of thermodynamic functions of the formation of 1: 1 complexes containing hydrogen bonds in systems where one of the components undergoes dimerization or where proton-transfer complexes form has not yet received much attention. This is probably owing to the difficulties caused by the competitive processes taking place in a non-polar medium which make it very hard to determine both the equilibrium constant and the enthalpy of complex formation.Using the calorimetric method described in an earlier article’ we have determined the enthalpies of complex formation in systems of pentachlorobenzoic acid (PClHBz) and pentafluorobenzoic acid (PFHBz) with such acceptors as dimethyl sulphoxide (DMSO), tetrahydrofuran (THF), methyl ethyl ketone (MEK) and acetonitrile (MeCN) in CCl,. The equilibrium constants and dimerization enthalpy values necessary for the calculations were taken from the The second part of this work consisted of a calorimetric study of complexes arising with proton donors of another type, i.e. those which derive their acidity from strongly polarized C-H bonds. Spectroscopic studies of C-H acids have thus far usually been confined to interactions with halogen derivatives of methane and ethane as proton donors., This leads to the conclusion that the hydrogen bonds which thus arise are very weak.The work of Denisov et aL5 on 1,l-dinitroethane (DNE) solutions in a number of donor-acceptor systems showed that DNE is almost as good a proton donor as phenol. This conclusion was reached by comparing the respective enthalpies of formation of donor-acceptor complexes. When a strong base is used, e.g. amine, DNE forms 1 : 1 proton-transfer complexes in CCl, which are in equilibrium with molecular complexes (hydrogen-b~nded).~ To date, acid-base complexes with DNE have not been investigated by the calorimetric method. I7551756 ASSOCIATION IN TERNARY SYSTEMS In the study reported here we chose to employ the calorimetric method as it is the only method which permits direct determination of thermal effects accompanying the processes under investigation.A comprehensive thermodynamic description of DNE-amine systems in CCl, using spectroscopic methods is not easy to obtain. Thus, i.r. spectroscopy (in the range 1000-1700 cm-l) enables one to determine the equilibrium constant of proton transfer in the 1 : 1 complex, but not that of molecular complex formation. Using the n.m.r. method one can determine the value K for complex formation via a hydrogen bond, but the procedure has only been employed with systems where no ion associates are formed. One exception is DNE-dibutylamine in CCl,, for which Denisov et ~ l ., ~ - ' using i.r. spectroscopy and n.m.r., obtained values of K and AH for both molecular complex formation and proton transfer. Equilibrium constants and enthalpies of 1 : 1 complex formation with a hydrogen bond can be determined in the presence of ion pairs by the calorimetric method, whose use requires a knowledge of the equilibrium constant and enthalpy of proton transfer determined in some other way, e.g. by i.r. spectroscopy. We have determined the values of K and AH of molecular complex formation using the following compounds as proton donors : NN-dimethylformamide (DMF), NN- dimethylacetamide (DMA), dimethyl sulphoxide (DMSO), 1,1,3,3,-tetramethylurea (TMU), hexamethylphosphoric triamide (HMPT), triethylamine (TEA), and di- isobutylamine (DIBA).In each case the proton donor was DNE, and the solvent CCl,. EXPERIMENTAL Calorimetric measurements of the heat of dissolution were made using the calorimeter employed in an earlier study.l The benzoic acid derivatives and 1,l-dinitroethane were synthesized according to literature methods.*-1° Following purification and vacuum sublimation (performed twice) the purity of the carboxylic acids was tested using the melting point: PFHBz, 103.8-104.0 OC(1it. 103.5-104.0 OC"); PClHBz, 207.6-207.8 OC (lit. 199-208 'Clz). Thecolourless DNE fraction obtained boiled at 24.2 "C/ 13.5 mmHg (ni0 = 1.4342, lit. nio = 1 .43461°3 13). The amines, DIBA and TEA (Loba Chemie, p.a.) and DMSO (Reachim SU 'p'), were dehydrated by heating with CaH,. The amides, DMF, DMA (Fluka AG fur UV-spektr.) and TMU (Fluka AG Buchs S.G. pract.), were dried over molecular sieve 4A.All liquids were distilled twice, in each case separating the middle fraction whose boiling point coincided with literature data. Hexamethylphosphoric triamide (Fluka A. G. Buchs S.C. pract.) was distilled and stored under an argon atmosphere to prevent the formation of peroxides. The dioxan used for di- electric measurements was purified according to the procedure described by Bates and Hobbs.'* The refractive index was determined using a Pulfrich refractometer (Carl Zeiss Jena), and dielectric measurements were performed using type GK-68 DK-meter (GDR). RESULTS AND DISCUSSION DETERMINATION OF THERMODYNAMIC FUNCTIONS FOR COMPLEXES CONTAINING HYDROGEN BONDS The procedure for determining the enthalpies of 1 : 1 complex formation with hydrogen bonds in carboxylic-acid-proton-acceptor systems was described in an earlier paper.l In the case of systems in which complexes involving hydrogen bonds and proton transfer may arise at the same time, the procedure is entirely different.t 1 mmHg z 133.322387 Pa.s. TANIEWSKA-OSINSKA AND P. GORALSKI 1757 The state of equilibrium in such a system can be described in the following way: AH [AH - * B] [AHl.[Bl ' A H + B e A H * - - B ; K = [A- HB+] AH. * * B e A - H B + ; Kp, =-- [AH. * *B]' AHPT* The measured value of the heat of dissolution of one of the components in the other dissolved in CCl, is a resultant of three effects: AHm = AHs+AHeomp~ex+AHse~f (3) where AHs is the dissolution heat of the given component in pure CCl,, AHcomplex is the enthalpy of formation of type (1) and (2) donor-acceptor complexes and AH,,,, is the enthalpy of decomposition of self-associates of the components in CCl,.The value of AHm was measured for each donor-acceptor system 30-40 times for various concentrations of base and acid in CCl,. The proton-donor concentrations ( a ) used ranged from 0.001 to 0.02 mol dm-3 and proton-acceptor concentrations (b) used ranged from 0.005 to 0.25 mol dmP3. Higher proton-donor concentrations were not used, to eliminate the possibility of formation of self-associates and more complex aggregates. The AH, value was determined using independent calorimetric measurements. In the calculations account was taken of the concentration dependence of the heat of dissolution of a given component in the solvent, AH, =f(c).We could find no published study that took this dependence into account. The enthalpy related to interactions between molecules of the component (AH,,,,) in the solvent used was determined in the following way: (4) where AH? is the standard heat of dissolution of the component in CCl, and AH; is the enthalpy of dissolution in CCl, corresponding to its concentration in a three-component system [eqn (3)]. In determining AH,,,, use was made of the enthalpy of dilution of the associating component. Such a procedure takes into account the self-associations present, but does not require a knowledge of their stoichiometry and of their equilibrium constants and formation en thalpies. The thermal effect (AHcomplex) of the formation of mixtures of molecular and ionic complexes in solution is expressed in the following way: AH,,,, = (AH?- AH;) ncomplex AHcomplex = AHc.m.+AHi.p. = [AH. * 'BI V [AH(l +KPT)+AHpTKpTI ( 5 ) since AHc.m. = [AH. * *B] VAH (6) is the thermal effect of molecular complex formation and = [A-HB+] I/ (AH+AH,,) (7) is the thermal effect of ion-pair formation. related to the value [AH * * B] by the following dependence The equilibrium constant of process (I), i.e. the value we want to determine, is1758 ASSOCIATION I N TERNARY SYSTEMS In the case of systems where no proton transfer is observed, the calculations were carried out putting AHpT = 0 and KpT = 0. Experimental determination of the effect AHcomplex [eqn (3)] and the fact that the measured values are related to the K and AH values sought [eqn ( 5 ) and (S)] makes it possible to determine both of these values from a series of experiments.In the calculations use was made of the minimization condition of the AHcomplpx function by the least-squares method: where i is the ordinal number of the experimental measurement. ENTHALPY OF DISSOLUTION OF THE COMPONENTS IN eel, AT 25 OC HEATS OF DISSOLUTION OF CARBOXYLIC ACIDS IN cc1, The calorimetric measurements of the heats of dissolution of PClHBz and PFHBz in CCl, were performed in the concentration range 1 x 10-,-3 x mol dm-3, and the results are collected in table 1. No measurements could be performed at higher TABLE HE HEATS OF DISSOLUTION OF PClHBz AND PFHBz IN CC1, C,Cl,COOH C,F,COOH c / ~ O - ~ mol dm+ AHg/kJ mol-l c / ~ O - ~ mol dm-3 AH:/kJ mol-l 0.07 1 0.137 0.285 0.294 0.725 0.729 1.433 1.447 35.80 33.16 30.10 29.64 26.88 26.38 24.58 23.9 1 0.129 0.225 0.455 0.475 0.930 1.466 2.054 2.719 ~~ 38.18 37.39 36.05 35.71 32.45 32.03 30.3 1 30.02 concentrations owing to the low solubility of the acids in CCI,.The standard dissolution heats of PClHBz and PFHBz were obtained by combining the Kl and AHl values (of dimerization) determined by spectroscopy with the dissolution enthalpies measured calorimetrically where c is the total acid concentration in the solution (in mol dmP3). If the AH? values were determined exclusively on the basis of calorimetric measurements, they would be much less precise owing to the small range of concentrations available for measurements. Using dependence (3) we obtained AHF(PC1HBz) = 39.4 1.4 kJ mol-1 and AH?(PFHBz) = 41.0+ 1.3 kJ mol-l.The error in the above values is due mainly to the errors in Kl (ca. 15 %) and AHl (ca. lo%)?s. TANIEWSKA-OSINSKA AND P. GORALSKI 1759 HEATS OF DISSOLUTION OF DNE I N CCl, The results are presented in table 2. The observed concentration dependence of dissolution enthalpy in CCl, points to intermolecular interactions of DNE molecules in solution. One can suggest the existence of DNE-DNE associates owing to the fact that DNE contains both a distinctly acidic proton and electron-donating NO, groups. TABLE 2.-ENTHALPIES OF DNE DISSOLUTION IN eel, AT 25 O C c / ~ O - ~ mol dmP3 AH,/kJ mol-l 0.0 0.265 0.690 0.8 15 1.333 1.763 2.333 3.41 5 3.855 5.843 7.922 10.173 60.800 17.18 15.72 14.85 14.58 14.06 13.49 13.34 13.10 13.15 13.07 13.00 12.95 12.44 Czerska et aZ.15 hold that the nitro groups of a number of compounds, including DNE, can form weak hydrogen bonds, e.g.with phenol. Gramstad and Simonsen16 determined the enthalpy of interaction between nitromethane and 2,6-di-t-butylphenol in C&1, to be - 3.5 kcal molt1. Therefore the interactions between DNE molecules in CCl, may be provoked by weak hydrogen bonds. In the pertinent literature there is no mention of the observed effect of self-association of DNE in non-polar solvents (if it does occur). HEATS OF DISSOLUTION OF PROTON ACCEPTORS The experimental results of dissolution heats of DMA, HMPT, TMU, TEA and DIBA in CCl, at 25 OC are presented in table 3.The dissolution enthalpies of the other proton acceptors in CCl, can be found e1sewhere.l The isotherms of the heats of dissolution of the proton acceptors point to weak interactions between the molecules of these bases in CCl,. The AH? values were determined by graphic extrapolation. In the relevant literature one can find the dissolution heats of some of the acceptors employed in the present i n v e ~ t i g a t i o n , ~ ~ ~ ~ ~ but the authors of those studies do not quote the concentrations to which their values correspond. In view of the fact that the observed concentration dependence of the heat of dissolution suggests considerable intermolecular interaction only in the case of DNE the value of AH,,,, [eqn (4)] was determined only for that component.DIPOLE MOMENTS OF THE PROTON DONORS EMPLOYED In order to arrive at a more complete comparative description of the acids under study, we also measured their dipole moments using the classical procedure of determining the dielectric permittivity, refractive indexes and solution densities.19 In1760 ASSOCIATION I N TERNARY SYSTEMS TABLE 3.-ENTHALPIES OF PROTON ACCEPTOR DISSOLUTION IN cc1, AT 25 OC (AHJJ mol-l, c/mol dm-3) 0.0000 0.0010 0.00 16 0.0027 0.0063 0.0104 0.0162 0.0233 0.0520 0.07 17 - 5238 - 5376 - 5501 - 5573 - 5686 - 5736 - 5803 -5816 - 5950 - 6008 0.0000 0.0024 0.0046 0.0087 0.0131 0.0173 0.0237 0.052 1 DIBA - 2472 - 2548 - 2623 - 2686 -2719 - 2740 - 2778 - 2887 DMA 0.0000 0.005 1 0.0109 0.0170 0.0232 0.0289 0.0426 0.0560 0.0690 0.090 1 - 1445 - 1555 - 1644 - 1688 - 1723 - 1758 - 1826 - 1885 - 1903 - 1947 C AHS 0.0000 0.0043 0.008 1 0.01 52 0.0249 0.0325 0.0473 0.052 1 - -251 - 398 - 507 - 541 - 557 - 566 - 591 -61 1 0.0000 0.002 1 0.0043 0.0088 0.0 148 0.0224 0.0287 0.0379 0.0454 0.062 1 2556 2480 2359 2212 21 54 2015 1940 1902 1852 1793 view of the low solubility of pentahalogeno substituted acids in CC1, we decided to use 1,4-dioxan as the solvent in this part of the study.Another reason for this choice was the proton-acceptor properties of dioxan, which emphasise the differences between the individual acids as proton donors; one can assume that in a dilute dioxan solution each acid molecule is hydrogen-bonded to a solvent molecule. Furthermore, dioxan ensures complete dissociation of carboxy acid dirners.,O The measurements were performed at concentrations ranging from x, = 5 x lo-, to x, = 1.5 x mole fraction.Apart from these measurements, the dipole moments of the acids under investigation were also determined by quantum-mechanical calculations using the CND0/2 method using the CNIND program.21 The interatomic distances and bond angles required for the calculations were adopted from X-ray data.22-25 Since in the literature there are no crystallographic data on PClHBz we assumed a ring structure with substituents, analogous to that of PClPh. The parameters describing the structure of the carboxy group were adopted from the data on PFHBz. Considering the fact that the angle between the ring plane and the carboxy group in pentafluorobenzoic acid is 29.8O and that chlorine atoms are larger than fluorine atoms, p(PC1HBz) was calculated assuming various values of the angle # for 30 6 #/O 6 90.The values obtained ranged from 2.23 to 2.28 D. The experimental results and theoretical calculations are collected in table 4.s. TANIEWSKA-OSINSKA AND P. GORALSKI 1761 TABLE 4.-EXPERIMENTAL AND THEORETICAL DIPOLE MOMENTS (D) OF THE BENZOIC ACIDS UNDER STUDY p (expt.1 p (calc.) acid dioxan gas phase p (lit.) HBz 2.19 k 0.10 1.93 1.64-2. 1226-32 PFHBz 2.96 k 0.07 2.09 - PClHBz 3.17k0.10 2.23-2.28 - The fact that the experimental results are higher than the calculated ones for isolated gas molecules and the high growth rate of the experimental values obtained in the HBz, PFHBz and PClHBz series could probably be attributed to the strong ‘dioxan effect ’.However, this effect should not disrupt the sequence of dipole-moment values in this series. In order to find out whether the observed direction of changes of p in acids is analogous to that in phenols, we performed a parallel calcuiation for phenols. Their experimental y values are available in the literature. The calculated and literature data are listed in table 5. Analysing the data in tables 4 and 5, one is struck by dissimilar growth sequences in the phenol and acid series that are hard to account for. The most polar compoimd in the acid series is PClHBz; in the phenol series it is PFPh. This follows both from the experimental findings and from theoretical calculations. TABLE 5.-DIPOLE MOMENTS (D) OF PHENOLS (Ph, PFPh, PClPh) p (calc.) phenol gas phase p (lit.) Ph 1.75-1.83a 1.45 (gas)-1.75 (C6H12)33 PFPh 2.38 2.14 (C6H12),34 2.54 (C6H6)35 PClPh 2.24 2.05 (C6H,2)36 a Differences are due to the various crystallographic structures of phenol. HEATS OF FORMATION OF 1 ; 1 DONOR-ACCEPTOR COMPLEXES W I T H HYDROGEN BONDS A N D ITS DISCUSSION CARBOXYLIC A C I D COMPLEXES Calorimetric measurements of the dissolution heats for each substance under study dissolved in the other component were made under conditions of both excess acid and excess base.In systems with PFHBz the measurements in both cases, i.e. a > 6 and b > a, yielded the same AH values. In the case of systems with PClHBz the enthalpies of complex formation in solutions with excess acid were always higher than those obtained in base excess.For example for the PClHBz-THF system we obtained for a > 6 -AH = 86.2 kJ rnol-l, while for 6 > a -AH = 36.0 kJ mol-l. The fact that the values obtained were so much higher than the known enthalpies of association with hydrogen bonds could be related to the formation of complexes other than AB, e.g. A,B. The formation of such complexes could be due to an excess of proton donor uis-a-vis acceptor. The use of excess base almost excludes the possibility of formation1762 ASSOCIATION I N TERNARY SYSTEMS TABLE 6.-ENTHALPIES OF 1 : 1 COMPLEX FORMATION WITH HYDROGEN BONDS IN SYSTEMS OF BENZOIC ACID AND ITS PENTAFLUORO- AND PENTACHLORO-DERIVATIVES WITH VARIOUS ACCEPTORS - AH/kJ mol-l (HBz) - AH/kJ mol-1 (PFHBz) - AH/kJ mol-' (PClHBz) base IU IIb IC I1 IC I1 MeCN 17.1 & 11.3 14.2 22.2 & 1.6 22.6 27.2 & 4.4 18.0 MEK 20.1 k 8.3 18.0 26.0 k 1.3 26.0 33.1 & 2.6 21.8 THF 23.4 & 3.3 21.8 30.1 f 1.1 30.1 36.0 f 1.9 25.1 DMSO 32.2k2.5 32.6 38.5+ 1 .1 40.6 41.0k 1.4 36.8 Ref. (1); ref. (3); this work. of associates containing more than one molecule of proton donor. Thus the AH values presented here correspond exclusively to the formation of AB complexes. All the results are collected in table 6 (column 1). This table also contains the data obtained earlier for benzoic acid.' The results are compared with those obtained by i.r. spectroscopy (column 2). The two sets of data are in good agreement in the case of complexes with HBz and PFHBz, while in the case of complexes with PClHBz there are considerable differences.The calorimetrically determined enthalpies of complex formation in systems of PClHBz with acceptors are the highest of all the benzoic acids studied. Besides benzoic acids one finds literature values of equilibrium constants and enthalpies of formation of hydrogen-bond complexes in systems of pentahalogeno substituted phenols with various acceptors. All systems mentioned have been studied by i.r. s p e ~ t r o s c o p y , ~ ~ - ~ ~ and it has been suggested that, regardless of the type of proton acceptor, complexes with pentafluoro substituted proton donors (PFPh, PFHBz) exhibit larger values of Av(OH), K and AH related to complex formation then those for unsubstituted donors or for donors with five chlorine atoms substituted in the ring.A h , pentahalogeno substituted donors exhibit deviations from the linear correlation of Av(OH), K and AH for hydroxylic proton donors with acidity expressed in terms of pK, values in water. Due to the effect of the substituents on the ring, the acidity ofpentachloro substituted donors is higher than that of their pentafluorine analogue^.^^ By contrast PClHBz (as well as PClPh) forms weaker hydrogen-bonded complexes than does PFHBz (and PFPh), sometimes also weaker than its unsubstituted Different authors account for this behaviour in different ways. Thus, Gramstad4* attributes the formation of weaker complexes by PClPh than by phenol to dipole association of pentachlorophenol with the proton acceptor. The apparent decrease in proton-acceptor concentration or the decrease in proton acidity in PClHBz in the systems under study should not be very marked.It seems to us that in the case of acids this phenomenon does not explain the higher K and AH values related to complex formation in pentachloro substituted donors vis-a-vis pentafluoro substituted donors. The difference in p values for pentachloro- and pentafluoro-substituted proton donors is not large enough (Ap = 0.21 _+O. 17 D) to explain why dipole association should occur in PClHBz but not in PFHBz. In the case of phenols, the p values (table 5) suggest that the polarity of the fluorine analogue is even higher than that of the chlorine one. Additional dipole association in PClHBz might provide its exothermic contribution to the AH value determined by calorimetry, but it would not be as high as the observed 10 kJ mol-l.According to Denisov, formation of a hydrogen bond by a PClPh molecule destroyss. TANIEWSKA-OSINSKA AND P. GORALSKI 1763 the planar structure of pentachlor~phenol.~~ The disturbances in the n-electron system lowers the acidity of PClPh and weakens the hydrogen bond. However, in the case of acids this phenomenon does not play a significant According to K ~ o p i o , ~ ~ the main factor accounting for the lack of correlation between the K , AH and Av(0H) values for pentahalogeno substituted proton donors and pK, is the formation of an intramolecular hydrogen bond. A hydrogen bond with a chlorine atom in the ortho position is in competition with the formation of a hydrogen bond with an acceptor molecule.In a sense, such a bond neutralizes and stabilizes the proton-donor molecule. This rather convincingly explains why the K and AH values of complex association with PClPh are lower than in the case of PFPh. However, such an intramolecular bond should also lower the acidity of PClPh, i.e. its ability to dissociate the hydroxylic proton hydrogen-bonded tochlorine. Recognizing that the intramolecular bond between the proton and the chlorine atom is stronger than that with fluorine, PClHBz should also exhibit the smallest susceptibility to dimerization among the acids used. This is because an intramolecular bond should stabilize the monomeric form of the acid in CCl,. Yet the results obtained by Zhukova et aL2 indicate that the dimerization constant for PClHBz is ca.3.5 times larger than that of its fluorine analogue. Futhermore, it seems that the intramolecular bond idea cannot exhaustively account for the anomalous behaviour of pentahalogeno substi- tuted proton-donors observed in the course of association. Our results are in good agreement with the directions of changes in acidity, but diverge from the spectroscopic data cited. To date, analogous systems have not been studied by the calorimetric method. The observed divergence might be a result of the marked interaction between the proton of the carboxy group and the chlorine atom in the ortho position. This interaction would have to be large enough to weaken the intensity of the v(0H) band of the free acid, which in itself serves as the basis for determining the dimerization constant.This would explain the higher Kl values obtained by Zhukova et al. as well as the high enthalpy values obtained by us for complexes with PClHBz. As can be seen from eqn (3) the high K , value increases the absolute AH? value, which in turn augments the calculated enthalpy of donor-acceptor complex formation. It seems that some of the doubts that arise in the study of pentahalogeno substituted proton donors might be dispelled and the problems involved interpreted more adequately if the equilibrium constants determined by spectroscopy were not related to the OH bands and if investigations used alternative experimental methods. 1,l -DIN I T ROE THANE COMPLEXES The calorimetric measurements reported earlier in this study made it possible to determine equilibrium constants and enthalpies of formation of complexes with a hydrogen bond for a number of DNE-proton-acceptor systems in CCl,. Also TABLE 7.-THERMODYNAMIC FUNCTIONS OF THE FORMATION OF HYDROGEN-BONDED DNE COMPLEXES WITH VARIOUS ACCEPTORS IN CC1, AT 25 *C base DN(SbC15)4v K/dm3 mol-l -AH/kJ mol-l -AG/kJ mol-I -AS/J mol-l K-I Ka/dm3 mol-' -AHa/kJ mol-I DMF 26.6 14.6f 1.5 18.69f0.62 6.64k0.24 40.4 f 3.0 10.0 18.02+ 1.26 - DMA 27.8 10.9f0.4 19.57k0.34 5.92f0.09 45.8 f 1.4 - DMSO 29.8 14.2f0.7 19.69f0.34 6.57k0.12 44.0 f 1.5 4.7 9.34 TMU 31.0 8.7f0.3 20.24k0.34 5.36+0.09 49.9 f I .5 - - TEA 61.0 16.4k2.5 27.40k3.14 6.93k0.35 68.7 f 1 1.7 - - DIBA - 28.6k3.8 29.20k2.01 8.31 f0.31 71.3f7.9 - HMPT 38.8 31.8f2.9 21.65+0.65 8.57k0.22 43.9k2.9 255 29.33f4.68 - a Denisov's data5.43 related to 25 OC.1764 ASSOCIATION I N TERNARY SYSTEMS determined were the AG and AS values of molecular complex formation. The results obtained are collected in table 7.In the range of systems studied proton transfer takes place when TEA and DIBA are used as proton donors. The values of KPT and AH,, required for the calculations were adopted from the study of Denisov et al.5 In the case of some other systems with amines in CCl,, e.g. trioctyl-and dibutyl-amines, we found it impossible to carry out sensitive calorimetric measurements owing to the formation of amine hydrochloride^^^ under the experimental conditions used. The results obtained in this study can be compared with those obtained by the n.m.r. m e t h ~ d ~ ? , ~ for systems containing DMF, DMSO and HMPT.In the case of DMF-DNE the equilibrium constants and enthalpies of complex formation obtained by the two methods are in good agreement, while in the case of the DMSO system the n.m.r. yields results which are lower than expected, considering the basicity of the compound. Earlier studiesl79 44-48 of various acceptors with phenols and alcohols in CCl, suggest that the enthalpy of formation of complexes with DMSO is as a rule higher than that for DMF or DMA. The large difference in the value of the equilibrium constant for HMPT-DNE complex formation, with AH being of the same order of magnitude (table 71, is difficult to explain. The direction of changes in the enthalpies of formation of complexes with DNE for the proton acceptors under study is qualitatively the same as in the case of phenol or p-fluorophenol.In fig. 1 we present the AH values obtained as a function of Gutmann’s donor number49 [DN(SbCI,)]; there appears to be a linear correlation between the two values. This in turn suggests that the ‘donicity’ of the bases used in this work is the main factor determining the changes in the enthalpies of various proton acceptors with the same C-H acid. 20 30 40 50 60 DN(SbC1,) FIG. 1.-Enthalpy of hydrogen bonded DNE complex formation with various acceptors as a function of Gutmann’s donor number. S. Taniewska-Osinska and P. Goralski, J. Chem. SOC., Faraday Trans. 1, 1981, 77, 969. W. A. Zhukova, L. J. Tarasova and M. J. Sheikh-Zabe, Teor. Eksp. Khim., 1978, 14, 396. G. S. Denisov and M. J. Sheikh-Zabe, Teor.Eksp. Khim., 1978, 14, 398. A. S. N. Murty, C. N. R. Rao and C. N. Ramachandra, Appl. Spectrosc. Rev., 1968, 2, 69. N. S . Golubiev and G. S . Denisov, Mol. Spektrosk., 1977, 4, 62. P. M. Borodin, N. S. Golubev, G. S. Denisov and Yu. A. Ignatiev, Org. Magn. Reson., 1975, 7 , 185. ’ G. S. Denisov, N. S. Golubev and V. M. Schreiber, Stud. Biophys., 1976, 57, 25. * D. E. Pearson and D. Cowan, Org. Synth., 1964, 44, 78.s. TANIEWSKA-OSINSKA AND P. GORALSKI 1765 D. E. Pearson, V. Mitchell, T. J. Sullivan and J. T. Watson, Synthesis, 1978, 2, 127. lo R. B. Kaplan and H. Shechter, J. Am. Chem. SOC., 1961, 83, 3535. l 1 E. Nield, R. Stephens and J. C. Tatlow, J . Chem. Soc., 1959, 166. l 2 E. H. Huntress, Organic Chlorine Compounds ( J . Wiley, New York, 1948), p. 464.l 3 H. Shechter and L. Zeldin, J . Am. Chem. Soc., 1951, 73, 1276. l 4 W. W. Bates and M. E . Hobbs, J . Am. Chem. SOC., 1951, 73, 2151. l 5 N. 0. Czerskaja, W. J. Szilenko, W. A. Szlapocznikov and S . S . Novikov, Izu. Akad. Nauk SSSR, l6 T. Gramstad and 0. R. Simonsen, Spectrochim. Acta, Part A , 1976, 32, 723. lH E. M. Arnett, L. Joris, E. J. Mitchell, T. S. S. R. Murty, T. M. Gorrie and P. v. R. Schleyer, J . Am. l 9 N. E. Hill, W. E. Vaughan, A. H. Price and M. Davies Dielectric Properties and Molecular Behaoiour 2o G. W. Gusakova, G. S. Denisov and A. L. Smolyansky, Zh. Prikl. Spectrosk., 1971, 14, 860. 21 Quantum Chemistry Program Exchange, no. 141. 2 2 T. Sakurai, Acta Crystallogr., 1962, 15, 1164. 28 V. Benghiat and L. Leiserowitz, J. Chem.SOC., Perkin Trans. 2, 1972, 1778. L5 G. A. Sim, J. M. Robertson and T. H. Goodwin, Acta Crystallogr., 1955, 8, 157. 26 H. A. Pohl and M. E. Hobbs, Ann. N.Y. Acad. Sci., 1940, 40, 389. 2 i C. S. Brooks and M. E. Hobbs, J. Am. Chem. SOC., 1940, 62, 2851. 28 G. J. Wilson and H. H. Wenzke, J . Am. Chem. SOC., 1935, 57, 1265. 29 G. K. Estok and S. P. Sood, J . Phys. Chem., 1957, 61, 1445. 3o K. Palm and H. Dunken. Z. Phys. Chem., 1961, 217, 248. :H D. Anbanathan, Ind. J . Chem., 1975, 13, 512. 32 M. Misrd and R. Gupta, Ind. J . Pure Appl. Phys., 1972, 10, 5367. 33 A. L. McClellan, Tables of Experimental Dipole Moments (Freeman, San Francisco and London, 34 M. J. Aroney, M. G. Cleaver, R. K. Pierens and R. J. W. Le Fevre, J . Chem. SOC., Perkin Trans. 2, 35 H. Hua Huang, J . Chem. SOC., Perkin Trans. 2, 1975, 903. 36 J. Jadzyn and J. Malecki, Acta Phys. Pol. A , 1972, 41, 599. 37 T. Gramstad, Spectrochim. Acta, 1966, 22, 1681. 38 G. S. Denisov, A. L. Smolyansky, A. A. Trusov, M. J. Shiekh-Zabe and W. M. Szraiber, Zh. Obsch. 39 R. Kuopio, Acta Chem. Scand., Ser. A , 1977, 31, 369. 4(1 T. Gramstad, Spectrochim. Acta, 1963, 19, 1363. 4 1 W. Vlasov and G. G. Yacobson, Usp. Khim., 1974, 43, 1642. 4 2 G. Drefahl and G. Heublein, J . Prakt. Chem., 1963, 20, 323. 43 P. M. Borodin, N. S. Golubiev and G. S . Denisov, Yad. Magn. Rezon., 1974, 5, 83. 44 R. S. Drago, G. C. Vogel and T. E. Needham, J . Am. Chem. SOC., 1971, 83, 6014. R. S. Drago, N. O’Bryan and G. C. Vogel, J. Am. Chem. SOC., 1970, 92, 3924. 46 K. F. Purcell, J. A. Stikeleather and S. D. Brunk, J . Am. Chern. SOC., 1969, 91, 4019. 47 R. S. Drago, J. A. Nusz and R. C. Courtright, J . Am. Chem. SOC., 1974, 96, 2082. 49 V. Gutmann, The Donor-Acceptor Approach to Molecular Interactions (Plenum Press, New York, Ser. Khim., 1972, 3, 620. E. M. Arnett, E. J. Mitchell and T. S. S. R. Murty, J . Am. Chem. SOC., 1974, 96, 3875. Chem. SOC., 1970, 92, 2365. (Van Nostrand Reinhold, London, 1969). H. Giller-Pandrand, Bull. SOC. Chim. Fr., 1967, 1988. 1963). 1976, 1854. Khim., 1975, 45, 2253. A. D. Sherry and K. F. Purcell, J . Am. Chem. SOC., 1972,94, 1848. 1978). (PAPER 1 /947)

ISSN:0300-9599    

DOI:10.1039/F19827801755

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1767-1784 Novel Application of the Radiotracer Technique in the Study of Molecular Complexes BY SHUDDHODAN P. MISHRA* AND RAM ADHAR SINGH Department of Chemistry, Faculty of Science, Banaras Hindu University, Varanasi 221005, India Received 15th June, 1981 A novel application of the radiotracer technique for inferring separately the thermodynamic constants of some model molecular complexes of iodine with n-donors (aromatics) and n-donors (alcohols) in solution has been studied. Labelled tetramethylammonium pentaiodide (TMAI;) was synthesized using molecular iodine (I3*I) as a tracer. The solubility equilibria of TMAI,* with and without donor were established, and iodine concentrations were analysed using a thin-wall end-window Geiger-Muller counter.These data yield equilibrium constants and enthalpies of a series of molecular complexes which are comparable with the literature values, lending support to the applicability of the tracer technique to the study of molecular complexes. Various advantages of this technique over others have been outlined. A new field of research touching the boundaries of physics, chemistry and biology, that of molecular complexes, has been gathering momentum in the past few decades.l-* Its applications vary from theoretical physics to material science, as seen from the vast literature concerning these c o m p l e x e ~ . ~ - ~ ~ Ever since Benesi and Hildebrand reported a spectrophotometric study of the interaction of iodine with aromatic^,^^ information about the molecular complexes of iodine has played a key role in developing and testing theories of electron-donor-electron-acceptor interactions, since iodine is an acceptor molecule which forms stable molecular complexes over the complete range of interactions, i.e.from the weakest to the stronge~t.~-ll However, several reports have appeared in the literature in which the experimental results do not conform to the theoretical predictions.l? 9 9 1 5 7 l6 On the one hand, attempts have been made to explain such results on the basis of deviations from Beer’s law of optical absorption which are characteristic of cornple~es,~~ the effect of solvent on molecular complex equilibria and activity coefficient^,'^^ the formation of termolecular complexes,6* l5 etc. ; on the other hand, doubts have been raised as to the reliability of thermodynamic and spectral parameters deduced by optical methods a10ne.15-18 Although one can obtain a sufficiently reliable value of the product K,E from the slope of Benesi-Hildebrand (BH) and related plots, its separation into Kc, the equilibrium constant (ca. 1 dm3 mol-l), and E , the molar extinction coefficient (ca.lo4 dm3 mol-1 cm-l), which is obtained from the intercept, is not an easy task because a slight error in the intercept results in a large error in E an4 K,. In addition there remains considerable disagreement in these values when different physical methods are employed owing to their inherent l9 Thus before considering theoretical explanations one must acquire accurate values of K , and E for weak complexes, for which conventional spectral methods cannot yield reliable resu1ts.l’ It is now felt that in order to supplement the spectral data, various non-spectral methods (such as solubility, vapour pressure, refractometry, positron-annihilation lifetime measurements, etc.) can be employed in conjunction with the spectral technique.Thus if a reliable value of Kc is obtained by a non-spectral method, a more authentic value of E can be obtained from the composite spectral data. 17671768 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES Any physical property which is related in a simple way to the concentration of one of the species, free or complexed, and is temperature-dependent allows the determination of K,. A recent report highlights the use of non-spectral methods in the study of molecular complexes.2o Various other attempts have also been made in this direction.21 Valuable information has been obtained from nuclear magnetic resonance,21 infrared ~pectroscopy,~ resonance Raman spectroscopy,22 refractometry,23 vapour ~olubility,~~? 25 etc.More recent uses have been made of positron- annihilation lifetime measurements,26g 27 Mossbauer spectroscopy28 and photoelectron ~ p e c t r o ~ c o p y . ~ ~ A single report has appeared recently on the use of the radioisotope labelling technique for vapour-densi ty measurements of volatile inorganic species, e.g. TbCl, + AlCl, mixtures using 16"Tb as tracer.30 In this paper we report for the first time the application of radiotracer labelling to the study of some model molecular complexes of iodine with some n-donors (aromatics, e.g.benzene, p-xylene, naphthalene, hexamethylbenzene, anthracene) and n-donors (alcohols, e.g. methanol, ethanol, n-propyl alcohol, n-butyl alcohol, iso-propyl alcohol, iso-butyl alcohol, t-butyl alcohol). The constant activity method coupled with the radioisotope labelling technique has been used to determine the equilibrium constants, K,, and enthalpies, -AH, of the above complexes in n-heptane solutions. EXPERIMENTAL LABELLING A N D PROCEDURES The experimental method for the determination of K, of iodine complexes in solution is based on a recently reported constant-activity methodl73 25 with suitable modifications allowing the determination of iodine concentrations by the tracer technique.Labelled tetramethylammonium pentaiodide, TMAI:, was synthesised by using molecular iodine labelled with carrier-free radioiodine (1311) having an activity of ca. 5 mCi and a specific activity of ca. 1.1 Ci mmol-l. The carrier-free radioiodine in molecular form was obtained from the Isotope Group, Bhabha Atomic Research Centre, Trombay, Bombay, India. Details of the procedure of synthesis and analysis are outlined below: (i) The amount of iodine (unlabelled) required for the synthesis of TMAI,* (5.0 g of I, for 2.0 g of TMAI)31 was dissolved in methanol and its concentration was determined by iodimetric titration or otherwise after suitable dilution. This gives the reference (blank) concentration of iodine. (ii) To the above solution, the required amount of radioiodine (l3II) was added so as to obtain a measurable activity in the final solution.The solution was mixed thoroughly and a small amount withdrawn and treated with sodium thiosulphate to convert the iodine into iodide. This was necessary because molecular iodine was found to escape from the stainless-steel planchet even at room temperature and thus affect the reproducibility of the results. A few (fixed) mm3 of solution was placed on a planchet and the radioactivities of the sample and reference were counted with the help of a thin-wall end-window Muller counter under identical conditions and geometries. A thin-wall end-window D,y Geiger-Muller counter coupled with an automatic timer (Electronics Corporation of India Ltd, type no. GCS13A) was used for the measurement of activities.The background count rate was very low (ca 10 counts min-l). The activities were ca. 1000-2000 counts min-l. At least ten countings were made for reference and sample to minimize random errors inherent in the radiochemical measurements. General safety precautions for radiochemical work were rigorously followed. Synthesis of TMAI,* was carried out in a fume cupboard. (iii) The labelled iodine solution was then added to an appropriate amount of TMAI and the mixtures was allowed to stand overnight. TMAI: was separated out, filtered, washed with methanol and dried. This was used as a constant activity source. The mixtures of TMAI: + n-heptane and TMAI: +(donor + n-heptane), prepared by placing ca. 50 mg of TMAI,* in 5 cm3 of solution in stoppered volumetric flasks, were kept in aS.P. MISHRA AND R. A. SINGH 1769 thermostatted container at different temperatures (25, 30, 35 and 4OoC, with variations of kO.1 "C) for three days to attain the solubility equilibrium t oc TMAI,*(~) + TMAI,*(~)+I,* solvent solution t oc TMAI:(s) $ TMAI,*(s) + I, + DI,. donor solution solution A small portion of the above solution was taken from the equilibration vessel and treated with an equal volume of aqueous sodium thiosulphate so as to convert the iodine into iodide. It was shaken thoroughly so as to extract all the iodide in the aqueous layer and a few mm3 of the solution were measured for their radioactivity. Parallel experiments were also made using unlabelled TMAI, and u.v.-visible spectroscopy, to check that the concentration of iodine determined by the radiotracer technique agrees with that determined by spectrophotometry, and also to check whether I; species are produced in greater excess (in the case of alcohol complexes) during our radiotracer experiment^.^^ A Cary- 14 spectrophotometer fitted with a thermospacer (in which water from a thermostat was continuously circulated) was used for recording the u.v.-visible spectra.No detectable I; species were produced in the case of complexes with aromatics; however, some production of I; species in the case of alcohol complexes could not be eliminated, and suitable corrections were made to account for the excess of iodine concentration estimated by the radiotracer technique. CALCULATION OF EQUILIBRIUM CONSTANT For a reversible reaction between two molecules, donor D and acceptor A, to form a complex DA, the equilibrium constant K , may be written as D + A e D A ( 3 ) C(DA) K, = [ C( D)O - C( DA)] [ C( A)' - C( D A)] (4) where C(DA) is the equilibrium concentration of the complex and C(D)O and C(A)O are the initial concentrations of D and A, respectively.In the constant-activity method for studying molecular complexes of iodine, a constant concentration is maintained between two polyiodides [eqn (l)] if we keep the temperature and solvent constant.25 However, the addition of a donor to this system [eqn (2)] establishes a second equilibrium D + I, e DI, ( 5 ) which consumes some of the iodine in complex formation, thus displacing the equilibrium in eqn (2) to the right so as to keep the amount of free iodine the same as in eqn (1).Thus the increase in iodine concentration in eqn (2) over eqn (1) yields the concentration of the complex and the concentration of iodine in eqn (1) yields the equilibrium concentration of free iodine. Hence the equilibrium constant, K,, for the interaction of a donor, D, with the acceptor, I,, forming a 1 : 1 complex can be given as where C(D1,) is the equilibrium concentration of complex and C(12)t that of iodine at the temperature c OC and C(D)O is the initial concentration of the donor. Therefore, by measuring the concentration of iodine in different solutions of donor (polyiodide) and solvent (polyiodide) one can derive the equilibrium constant of complex formation; and by studying the temperature dependence of the equilibrium constant, one can deduce other thermodynamic parameters.Note that we need not know, while using the radiotracer technique, the position of the isosbestic point : this is a requirement for a spectrophotometric determination of the concentration of1770 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES iodine.l7yZ5 Further, note that once the solution has been taken out of the equilibrium vessel, the analysis of the total iodine concentration can be made at different temperatures (and even in a different medium, as in the present case). The good agreement between the values obtained by the radiotracer technique and others (uiz. spectrophotometry) leads us to believe that the partition functions between n-heptane and water do not deter the extraction of iodine from the n-heptane layer to the aqueous layer in the iodide form.The data were analysed on an ICL 1900s computer at the Computer Centre, Banaras Hindu University, Varanasi. MATERIALS n-Heptane (Koch-Light) was shaken three times with concentrated H,SO, followed each time by washing with distilled water. This was dried over KOH pellets, distilled over sodium wire and stored over molecular sieve (4A). Benzene (Sarabhai M. Chemicals, India, special for spectroscopy) was dried over molecular sieve (4A) and used without further purification. p-Xylene (B.D.H., AnalaR grade) was purified by fractional crystallization (m.p. 13.3 "C, lit. 13.4 "C). Naphthalene (B.D.H., microanalytical reagent grade) was purified by sublimation and subsequent recrystallization from cyclohexane (m.p. 80.2 "C in sealed tube).Hexamethylbenzene (Ega Chemie) was recrystallized from an ethanol + cyclohexane mixture (m.p. 164.5 "C). Anthracene (B.D.H., AnalaR grade) was recrystallized twice from ethanol (m.p. 216.3 "C). Iodine (B.D.H., AnalaR grade) was purified by sublimation from a mixture of potassium iodide and iodine (in the ratio 1 : 2.5) (m.p. 113.5 "C) and was stored in dark, stoppered bottles. Methanol (B.D.H., spectral grade) was stored over molecular seive [n.m.r. data: 60 MHz, 6(CH3) = 1.42, lit. 1.43, 6(OH) = 3.46, lit. 3.471. Ethanol (LR grade) was refluxed with fused lime for 10 h and distilled. The middle fraction was refluxed with magnesium turnings and iodine for 2 h and fractionally distilled. The middle fraction (b.p.78.3 "C) was used for preparing stock solutions [6(OH) = 3.69, lit. 3.701. n-Propyl alcohol (LR grade) was distilled and dried over K2C03, followed by refluxing with magnesium ribbon and iodine, and was fractionally distilled [(b.p. 97.1 "C, lit. 97.2 "C), 6(OH) = 3.58, lit. 3.581. iso-Propyl alcohol (LR grade) was distilled and dried over CaC1, followed by BaO and then fractionally distilled [b.p. 82.2 "C, &OH) = 4.02, lit. 4.001. t-Butyl alcohol (LR grade) was distilled from lime and fractionally crystallized three times (m.p. 25.4 OC, lit. 25.3 "C). Tetramethylammonium iodide was prepared by the action of methyliodide on an aqueous solution of trimethylamine at low temperatures (using ice + salt mixture) and was purified by crystallization from aqueous alcoholic The composition of TMAI, was established by iodimetry separately and was found to be correct.Special care was taken to exclude moisture from the solutions containing alcohol donors as this leads to the production of I; species.32 Dry nitrogen was flushed through the solvent and solutions to exclude dissolved oxygen in the solvent.' RESULTS AND DISCUSSION The equilibration of TMAI, in n-heptane at different temperatures yields a constant concentration of iodine at the respective temperatures. The dissociation constant of the equilibrium may be given as Kd TMAI,(s)~TMAI,(s)+I, Since TMAI, and TMAI, are both solids, the apparent dissociation constant, K;, may be given as The concentration of iodine in n-heptane at different temperatures yields the heat of dissociation of TMAI, in n-heptane, which can be obtained from Van't Hoff's plot of log K; against 1/T (fig.1). The heat of dissociation of TMAI, in n-heptane KA = [I,]. (9)S. P. MISHRA A N D R. A. S I N G H 1771 3’3.25 3.35 lo3 KIT FIG. 1 .-Van? Hoff plot for the heat of dissociation of TMAI, g TMAI, + I, in n-heptane. 0.5 0 . 4 E U - z E 1 N I W n i, 0 . 3 0.2 0 0.5 1 .o C(D)/mol dm-3 1.5 FIG. 2.-Plot of total iodine concentration [CD(I,)] determined by the radiotracer technique against the concentration of donor [C(D)] in n-heptane at 25 OC: (1) benzene, (2) p-xylene, (3) naphthalene, (4) hexamethylbenzene.1772 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES TABLE 1 .-EXPERIMENTAL CONCENTRATION DATA FOR AROMATIC-IODINE SYSTEMS IN n-HEPTANE AT 25 "C - - 0.394 0.406 0.406 0.404 0.402 0.401 activitya W)O /counts C(I,)b C(D1,) C(D) K, average K , /mol dmP3 min-' /mmol dm-3/mmol dmP3 /mol dmP3 /dm3 mol-' /dm3 mol-' - - 0.528 0.528 ref.0.000 00 0.020 00 0.050 00 0.100 00 0.200 00 0.500 00 1.000 00 2.000 00 ref. 0.000 00 0.020 00 0.050 00 0.100 00 0.200 00 0.500 00 1.000 00 2.000 00 ref. 0.000 00 0.010 00 0.020 00 0.050 00 0.100 00 0.250 00 0.500 00 1.000 00 ref. 0.000 00 0.005 00 0.010 00 0.020 00 0.050 00 0.100 00 0.250 00 0.500 00 ref. 0.000 00 0.005 00 0.007 50 0.010 00 0.020 00 0.030 00 0.040 00 0.050 00 0.528 0.530 0.532 0.531 1350 1350 1360 1377 1404 1458 1620 1890 2430 1225 1225 1238 1257 1290 1354 1550 1877 2525 1336 1336 1343 1353 1377 1419 1543 1751 2165 1288 1288 1307 1329 1370 1488 1692 2294 3305 1300 1300 1338 1361 1380 1456 1539 1615 1697 - - 0.6 19 0.616 0.619 0.620 0.619 0.621 0.197 0.197 0.198 0.201 0.205 0.213 0.236 0.276 0.355 0.197 0.197 0.199 0.202 0.207 0.218 0.249 0.302 0.406 0.197 0.197 0.198 0.199 0.203 0.209 0.227 0.258 0.319 ) benzene - - 0.00 1 0.004 0.008 0.0 16 0.039 0.079 0.168 p-xylene - - 0.002 0.005 0.010 0.02 1 0.052 0.105 0.209 naphthalene - - 0.001 0.002 0.006 0.01 1 0.030 0.08 1 0.122 - - 3.05 3.20 3.11 3.14 3.12 3.13 3.18 - - 0.020 00 0.050 00 0.099 99 0.199 98 0.499 96 0.999 92 1.999 83 - - 0.020 00 0.499 99 0.099 99 0.199 98 0.499 95 0.999 90 1.999 79 - - 0.010 00 0.020 00 0.049 99 0.099 99 0.249 97 0.499 92 0.999 88 > hexamethyl benzene 0.197 - 0.197 - - 0.200 0.003 0.005 00 0.203 0.006 0.009 99 0.209 0.012 0.019 99 0.228 0.03 1 0.049 97 0.259 0.062 0.099 94 0.351 0.154 0.249 85 0.505 0.308 0.499 61 0.197 - 0.197 - 0.203 0.006 0.000 499 0.206 0.009 0.007 49 0.209 0.012 0.009 99 0.221 0.024 0.019 98 0.233 0.036 0.029 96 0.245 0.048 0.039 95 0.257 0.060 0.049 94 - anthracene - - - - 6.00 6.24 6.15 6.00 6.14 6.06 6.11 > 0.402 0.529 0.619 3.13 6.10 a Error limit is 5 counts min-'; error limit is L-0.001 mmol dm-3.S.P. MISHRA A N D R. A. SINGH 1773 TABLE 2.-EXPERIMENTAL CONCENTRATION DATA FOR AROMATIC-IODINE SYSTEMS IN n-HEPTANE AT 30 "C activitya C(D)O /counts C(I.$ C(D1,) C(D) K, average K, /mol dm-3 min-l /mmoldm-3/mmoldm-3 /mol dmP3 /dm3 mo1-1 /dm3 mol-l ref. 0.000 00 0.198 76 0.496 90 0.993 80 1.987 60 ref. 0.000 00 0.099 42 0.198 85 0.497 12 0.994 24 1.988 48 ref. 0.000 00 0.099 38 0.248 56 0.496 92 0.993 84 ref.0.000 00 0.049 70 0.099 39 0.248 49 0.496 97 ref. 0.000 00 0.009 94 0.029 83 0.049 72 800 1003 1080 1195 1387 1771 1336 1707 1793 1879 2139 2572 3436 786 986 1043 1130 1274 1562 1335 1731 1972 2212 2935 4125 1360 1737 1830 2015 220 1 0.197 0.247 0.266 0.294 0.341 0.437 0.193 0.247 0.259 0.272 0.310 0.372 0.497 0.197 0.247 0.261 0.283 0.319 0.391 benzene - - 0.019 0.047 0.094 0.189 p-xylene - __ 0.012 0.025 0.063 0.125 0.250 naphthalene - - 0.014 0.036 0.072 0.144 hexamethyl benzene - - 0.193 0.247 - 0.281 0.034 0.049 67 0.316 0.069 0.099 32 0.419 0.172 0.248 32 0.589 0.342 0.496 63 0.193 0.247 - 0.260 0.013 0.009 93 0.286 0.040 0.029 79 0.313 0.066 0.049 65 - ant hracene - - - - - 0.198 74 0.496 85 0.993 71 1.987 41 - - 0.099 41 0.198 82 0.497 06 0.994 12 1.988 23 - - 0.099 37 0.248 52 0.496 85 0.993 70 0.385 0.385 0.385 0.385 0.505 0.507 0.509 0.510 0.509 0.583 0.588 0.588 0.588 2.806 2.801 2.802 2.786 1 5.384 5.370 5.376 0.385 0.508 0.586 2.800 5.377 a Error limit is & 5 counts min-l; error limit is 0.001 mmol dmP3.determined by the radiotracer technique is found to be 8.23 0.1 kcal mol-l, which is in good agreement with the most recent value of Tse and Tamres (8.18 kcal m ~ l - l ) . ~ ~ A R 0 MAT I C-I 0 DINE SYSTEMS The increase in radioactivity vis-a-vis the total concentration of iodine in solution with the concentration of donor for systems of iodine with various aromatics in n-heptane at 25, 30 and 35 *C are reported in tables 1-3 and shown graphically for1774 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES TABLE 3 .-EXPERIMENTAL CONCENTRATION DATA FOR AROMATIC-IODINE SYSTEMS IN n-HEPTANE AT 35 "c activitya C(D)O /counts C(12)b C(D1,) C(D) K, average K, /mol dm-3 min-l /mmoldm-3/mmoldm-3 /mmol dmP3 /dm3 mol-l /dm3 mol-l ref.0.000 00 0.197 52 0.493 80 0.987 60 1.975 20 ref. 0.000 00 0.098 80 0.197 61 0.494 02 0.988 04 1.976 07 ref. 0.000 00 0.098 76 0.256 91 0.493 82 0.987 64 ref. 0.000 00 0.049 39 0.098 77 0.246 94 0.493 87 ref. 0.000 00 0.009 88 0.029 65 0.049 41 1035 1618 1736 1914 2210 2801 1030 1641 1720 1799 2037 243 3 3225 1020 1595 1684 1816 2037 2480 1044 1662 1868 2075 2693 3723 1052 1656 1754 1912 2070 0.197 0.308 0.330 0.364 0.420 0.533 0.193 0.308 0.323 0.338 0.382 0.457 0.605 0.197 0.308 0.325 0.351 0.393 0.479 benzene - - 0.022 0.056 0.112 0.225 p-x ylene - - 0.01 5 0.030 0.074 0.149 0.297 naphthalene - - 0.017 0.043 0.085 0.171 - - 0.197 50 0.493 74 0.987 49 1.974 97 - - 0.098 78 0.197 58 0.493 95 0.987 89 1.975 77 - - 0.098 74 0.246 87 0.493 73 0.987 47 hexame t h y 1 benzene 0.193 - - 0.308 0.346 0.038 0.049 35 0.384 0.076 0.098 69 0.499 0.191 0.246 75 0.689 0.381 0.493 49 an thracene 0.193 0.308 __ 0.322 0.014 0.009 87 0.35 1 0.043 0.029 61 0.380 0.073 0.049 34 - - - - - 0.370 0.371 0.371 0.370 0.486 0.486 0.488 0.488 0.488 0.566 0.562 0.56 1 0.562 2.509 2.51 1 2.510 2.510 4.740 4.762 4.773 0.3705 0.487 0.563 2.510 4.755 a Error limit is 5 counts min-l; error limit is fO.OO1 mmol dmP3.25 O C in fig. 2. It varies linearly with the donor concentration up to ca. 2 mol dm-3.The concentration of the complex, determined from the increase in iodine concentration in donor solution over that in the pure solvent, is found to be linear when plotted against the equilibrium concentration of the donor. This substantiates the argument that predominantly 1 : 1 complexes are formed within the experimental range of donor concentrati~ns.~~ Values of the equilibrium constants and enthalpies for systems of iodine with various aromatics have been calculated using eqn (6) point by point, andS. P. MISHRA AND R. A. SINGH 1775 TABLE 4.-EQUILIBRIUM CONSTANTS AND ENTHALPIES OF COMPLEXES OF IODINE WITH AROMATICS IN n-HEPTANE SOLUTION, DETERMINED BY THE RADIOTRACER SOLUBILITY METHOD donor K, (25 OC)/dm3 mol-1 - AH/kcal mol-I benzene p-xy lene naphthalene hexamethylbenzene anthracene 0.40 1 f 0.006 (0.356,a 0.246b) 0.532 & 0.008 (0.642,a 0.41 la) 0.620 f 0.008 (0.544,d 0.2Y) 3.13 f 0 .10 (2.99d) 6.08 _+ 0.15 (3. 70e) 1.51 k0.05 (1.3OC) 1.68 & 0.08 1.70 f 0.10 (1.91,d 1.80a) 4.04 f 0.20 4.50k0.20 (2.209 (4.339 Values in parentheses refer to the literature data: a ref. (4); M. Tamres and S . Searles, J. Am. Chem. SOC., 1960, 82, 2129; W. B. Person, J. Am. Chem. SOC., 1964, 86, 10. ref. (33); S. M. Brandon, ref. (23b); J. Peters and the K, values thus obtained are plotted against C(D). A linear plot results with an intercept at C(D) = 0 equal to K,. Such plots are almost parallel to the abscissa for all the aromatic-iodine systems. The K, values thus obtained are reported in table 4 along with the literature data (wherever available) and enthalpies of complex formation.It is evident from table 4 that the equilibrium constants obtained by the radiotracer technique are higher than those obtained by conventional spectral methods. However, they are comparable to the values obtained by the spectral-solubility method.l79 25 The values of the equilibrium constants determined by the present technique, the spectral-solubility method and the conventional Benesi-Hildebrand procedure at 25 'C in n-heptane are (respectively) as follows (in dm3 mol-l): (i) benzene-iodine, 0.401 f 0.006 (0.3564 and 0.24633), (ii) p-xylene-iodine, 0.532 f 0.008 (0.642 and 0.41 14), (iii) naphthalene-iodine, 0.620 & 0.008 (0.54425b and 0.2537b), (iv) hexamethyl- benzene-iodine, 3.13 f 0.10 (2.9923b and 1 .734), (v) anthracene-iodine, 6.08 k 0.15 (3.7037b).The solubility method yields a higher value of K, because it accounts for the total effect of donor-acceptor interaction (specific plus non-specific), whereas K, determined by the spectral method reflects only the effects of donor-acceptor contacts (specific) occurring in excess of random collision^.^^ The specific solvent effects include hydrogen bonding, complexation, isomerization, etc., while non-specific solvent effects include solvent cage strains, dispersion and polarization effects.34 In a later com- munication Christian et a/.35 showed that there is marked variation in the K, value, if activity-coefficient corrections are made based on the simple non-electrolyte theory. It is surprising to note that if the above corrections are made, the corrected values of K, are in good agreement with the values obtained by the Benesi-Hildebrand method.Both Tamres et ~ 1 . ~ ~ and Foster et pointed out that activity-coefficient effects in themselves appear to be too small to provide an alternative explanation to the observed values. The uncorrected results are in better agreement with the trend predicted by Mulliken's the0ry.l Since Kc& remains almost constant in the two methods, the higher values of K, would result in a lower value of E , which may follow an increasing order from benzene to hexamethylbenzene; this is what would be1776 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES TABLE 5.-EXPERIMENTAL CONCENTRATION DATA FOR ALCOHOL-IODINE SYSTEMS IN n-HEPTANE AT 25 O C W)O activity a C(I2Ib C(DI2) C(D) Kc /mol dm-3 /counts min-' /mmol dmP3 /mmol dmP3 /mol dm-3 /dm3 mol-' ref.0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 950 950 984 1018 1090 1235 1389 1553 1726 876 876 9.12 947 1023 1178 1352 1530 1730 840 840 876 913 989 1147 1322 1505 1701 726 726 762 79 1 857 1002 1157 1316 1491 856 856 895 934 1015 1 I86 1373 1573 1790 methanol 0.197 - 0.197 - 0.204 0.007 0.21 1 0.014 0.226 0.029 0.256 0.059 0.287 0.090 0.322 0.125 0.358 0.161 ethanol 0.197 - 0.197 - 0.205 0.008 0.2 13 0.0 16 0.230 0.033 0.265 0.068 0.303 0.106 0.344 0.147 0.389 0.192 iso-propyl alcohol 0.197 - 0.197 - 0.205 0.008 0.214 0.017 0.232 0.035 0.269 0.072 0.310 0.113 0.354 0.157 0.399 0.202 n-propyl alcohol 0.197 - 0.197 - 0.206 0.009 0.215 0.018 0.233 0.036 0.272 0.075 0.314 0.117 0.357 0.160 0.405 0.208 t-butyl alcohol 0.197 - 0.197 0.206 0.009 0.215 0.018 0.234 0.037 0.273 0.076 0.318 0.121 0.362 0.165 0.413 0.216 - - - 0.024 99 0.499 9 0.099 97 0.199 94 0.299 91 0.399 87 0.499 84 - - 0.024 99 0.049 98 0.099 97 0.199 93 0.299 89 0.399 85 0.499 81 - - 0.024 99 0.049 98 0.099 96 0.199 93 0.299 89 0.399 84 0.499 80 - - 0.024 99 0.049 98 0.099 96 0.199 92 0.299 88 0.399 84 0.499 79 - - 0.024 99 0.049 98 0.099 96 0.199 92 0.299 88 0.399 83 0.499 78 - - 1.42 1.46 1.48 1 S O 1.53 1.59 1.64 - - 1.61 1.63 1.67 1.73 1.80 1.87 1.94 - - 1.72 1.75 1.77 1.83 1.91 1.99 2.05 - - 1.78 1.80 1.82 1.90 1.98 2.03 2.1 1 - - 1.80 1.82 1.86 1.92 2.02 2.10 2.19S.P. MISHRA A N D R. A. SINGH 1777 TABLE 5. (cont.) activitya /counts min-l Kc /dm3 mol-l ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 ref. 0.000 00 0.025 00 0.050 00 0.100 00 0.200 00 0.300 00 0.400 00 0.500 00 720 720 756 787 857 1005 1170 1338 1530 556 556 582 609 663 782 912 1056 1094 iso-butyl alcohol 0.197 - 0.197 0.207 0.010 0.215 0.018 0.234 0.037 0.275 0.078 0.320 0.123 0.366 0.169 0.419 0.222 n-butyl alcohol 0.197 0.197 0.206 0.009 0.216 0.019 0.235 0.038 0.276 0.079 0.323 0.126 0.374 0.177 0.423 0.226 - - - - 0.024 99 0.049 98 0.099 96 0.199 92 0.299 88 0.399 83 0.499 78 - 1.84 1.87 1.90 1.99 2.09 2.15 2.25 - - 0.024 99 1.87 0.049 98 1.90 0.099 96 1.93 0.199 92 2.03 0.299 87 2.14 0.394 82 2.25 0.499 77 2.30 a Error limit is & 5 counts min-l; error limit is +0.001 mmol dm-3.expected from Mulliken’s theory. The higher values of K , obtained by the radiotracer technique than those from the spectral method using solubility principles can be attributed to the fact that there may be an enhancement of iodine solubility owing to other effects, e.g. orientation isomers in a 1 : 1 complex and contact charge transfer, which are not normally accessible by spectral methods in the normal range of study. The more accurate determination of concentrations by the present technique than via optical measurements may also account for the slightly higher values of K,.Both the equilibrium constants and enthalpies obtained by the radiotracer technique (cf. table 4) vary in the order of the electron-donating capacity of the donor (e.g. ionization potential), viz. benzene < p-xylene < naphthalene < hexamethylbenzene < anth~acene.~~ The anthracene-iodine system is of particular interest. Because of the lower solubility of anthracene in n-heptane and various problems involved in its study by spectrophotometry, no reliable thermodynamic data are available for this Such problems as a weak charge-transfer band in the visible region, interference by the donor, iodine and complex bands and the difficulty in determining the isosbestic point are encountered when studying this complex by spectrophotometry20 but not when using the radiotracer technique; thus the anthracene-iodine system was used to demonstrate the advantages of the latter technique over the former.Although the measurements were restricted to the low concentration range, values of K , (6.08 dm3 mol-l) and -AH (4.50 kcal mol-l) could be obtained. The K, value, although considerably lower than that reported by Bhattacharya and Basu (52.35 dm3 mol-l), is only slightly higher than that reported by Person et al. (3.7 dm3 m01-9.~~ 58 FAR 11778 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES TABLE 6.-EXPERIMENTAL CONCENTRATION DATA FOR ALCOHOL-IODINE SYSTEMS IN n-HEPTANE AT 30 OC C(WO activity" Wdb W I , ) W) Kc /mol dmd3 /counts min-' /mmol dm-3 /mmql dm-3 /mol dm-3 /dm3 mol-l ref.0.000 00 0.024 98 0.049 97 0.099 95 0.199 41 0.299 80 0.399 63 0.499 20 ref. 0.000 00 0.024 88 0.049 83 0.099 25 0.198 94 0.298 48 0.398 00 0.497 63 ref. 0.000 00 0.024 85 0.049 62 0.099 18 0.198 90 0.298 22 0.397 68 0.496 93 ref. 0.000 00 0.024 80 0.049 55 0.099 23 0.198 88 0.298 60 0.397 70 0.496 83 ref. 0.000 00 0.024 70 0.049 39 0.099 15 0.198 82 0.298 45 0.397 98 0.497 49 840 1053 1087 1122 1192 1339 1486 1637 1795 985 1248 1292 1340 1434 1628 1833 2052 229 1 968 1226 1274 1322 1420 1624 1830 2050 228 1 925 1172 1219 1267 1364 1564 1776 1997 2230 1000 1254 1304 1356 1462 1678 1905 2148 2402 methanol 0.197 - 0.247 - 0.255 0.008 0.263 0.016 0.280 0.023 0.314 0.067 0.348 0.101 0.384 0.137 ethanol 0.195 - 0.247 - 0.256 0.009 0.265 0.018 0.284 0.037 0.322 0.075 0.363 0.116 0.406 0.159 0.454 0.207 iso-propyl alcohol 0.195 0.247 - 0.257 0.0 10 0.266 0.019 0.286 0.039 0.327 0 1080 0.369 0.122 0.41 3 0.166 0.459 0.212 n-propyl alcohol 0.197 - 0.247 - 0.257 0.010 0.267 0.020 0.288 0.041 0.330 0.083 0.375 0.128 0.421 0.174 0.470 0.223 t-butyl alcohol 0.197 - 0.247 - 0.257 0.010 0.267 0.020 0.288 0.041 0.331 0.084 0.375 0.128 0.423 0.176 0.473 0.226 0.421 0.r74 - - - 0.024 97 0.049 95 0.099 93 0.199 84 0.299 70 0.399 47 0.499 03 - - 0.024 87 0.049 61 0.099 21 0.198 90 0.298 36 0.397 84 0.497 52 - - 0.024 84 0.049 60 0.099 14 0.198 82 0.298 10 0.397 51 0.496 71 - - 0.024 79 0.049 53 0.099 19 0.198 80 0.298 47 0.397 53 0.496 61 - - 0.024 69 0.049 37 0.099 11 0.198 74 0.298 52 0.397 80 0.497 26 - - 1.29 1.30 1.32 1.35 1.37 1.39 1.41 - - 1.45 1.48 1 S O 1.53 1.57 1.62 1.68 - - 1.54 1.57 1.59 1.63 1.65 1.69 1.73 - - 1.61 1.63 1.65 1.68 1.73 1.77 1.82 - - 1.62 1.64 1.67 1.70 1.74 1.79 1.84S. P.MISHRA A N D R. A. SINGH 1779 TABLE 6. (cont.) W)O activity a C(12)b W I 2 ) W) Kc /mol dm-3 /counts min-l /mmol dm-3 /mmol dm-3 /mol dm-3 /dm3 mol-l ref. 0.000 00 0.024 85 0.049 35 0.098 90 0.198 55 0.298 25 0.397 86 0.497 17 ref. 0.000 00 0.024 75 0.049 28 0.099 06 0.198 68 0.298 15 0.397 60 0.496 82 892 1118 1165 121 1 1306 1503 1706 1920 2148 818 1026 1068 1112 1202 1385 1577 1780 I994 iso-butyl alcohol 0.197 - 0.247 - 0.257 0.010 0.267 0.020 0.288 0.041 0.332 0.085 0.377 0.130 0.424 0.177 0.474 0.227 n-butyl alcohol 0.197 - 0.247 - 0.257 0.010 0.268 0.02 1 0.289 0.042 0.335 0.088 0.380 0.133 0.429 0.182 0.480 0.233 - 0.024 84 0.049 33 0.098 86 0.198 46 0.298 12 0.397 68 0.496 94 - 1.65 1.67 1.69 1.73 1.76 1.80 1.85 - 0.024 74 0.049 26 0.099 02 0.196 59 0.298 02 0.397 42 0.496 59 - 1.67 1.69 1.73 1.76 1.81 1.25 1.90 a Error limit is & 5 counts min-l; Error limit is f 0.001 mmol dm-3.A L C 0 HO L-I0 D I N E SYSTEMS Alcohols, particularly methanol and ethanol, are classical examples of n-type donors which form molecular complexes with iodine.l*s-ll In two recent papers Tse and Tamres have presented a detailed report on these two systems using the spectral-solubility m e t h ~ d . ~ ” ~ ~ They have discussed the data in the light of the self- association of alcohols and specific and non-specific effects included in the solubility experiments.However, detailed data on a series of alcohols is lacking. Note that self-association in an alcohol increases with an increase in itsconcentration in n-heptane solution; i.e. the association number increases from 1.2 to 4.1 when the alcohol concentration is increased from 0.04 to 0.60 mol dm-3.38 However, the degree of self-association of alcohols in a homologous series is expected to decrease owing to steric effects. Tamres et aZ.38 point out that the donor strength of the alcohol may change after association in the order trimer > dimer > monomer. Thus, the iodine molecule in a concentrated alcohol solution interacts with a stronger donor than in dilute solution.This should in principle result in the dependence of K, on concentration. However, they have found that no such concentration dependence is observed [shown by the linearity of A (isosbestic) against C(D) plots]34 and have proposed that the K, value obtained by this method is a weighted average of all the equilibrium constants for all different aggregates complexing with iodine. The hydrogen-bond energies of aliphatic alcohols are in the order methanol > ethanol > iso-propyl alcohol > t-butyl alcohol.3s It reflects steric and/or electronegativity effects on the strength of hydrogen bond. Note that the hydrogen-bond energies of alcohols are comparable with the enthalpies of alcohol-iodine complexes. ls? 3s Thus a competition between alcohol- 58-21780 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES TABLE 7.-EXPERIMENTAL CONCENTRATION DATA FOR ALCOHOL-IODINE SYSTEMS IN n-HEPTANE AT 35 "c C(D)O activitya C(Mb C(DI2) C(D) Kc /mol dmP5 /counts min-l /mmol dmP3 /mmol dm-3 /mol dm-3 /dm3 mol-1 ref.0.000 00 0.024 40 0.049 00 0.098 50 0.197 95 0.296 80 0.395 90 0.494 76 ref. 0.000 00 0.024 44 0.049 05 0.098 90 0.198 10 0.297 64 0.396 25 0.494 20 ref. 0.000 00 0.024 66 0.049 14 0.098 95 0.198 36 0.297 75 0.396 96 0.495 45 ref. 0.000 00 0.024 52 0.049 00 0.098 65 0.198 65 0.296 92 0.396 16 0.495 35 ref. 0.000 00 0.024 46 0.048 92 0.098 65 0.198 10 0.297 78 0.397 06 0.496 86 840 1312 1351 1391 1474 1644 1825 2005 2198 628 992 1024 1058 1127 1269 1420 1590 1772 600 948 98 1 1015 1085 1226 1376 1539 1709 546 862 894 925 990 1124 1265 1416 1576 636 994 1031 1068 1145 1302 1469 1638 1830 methanol 0.197 - 0.308 - 0.317 0.009 0.326 0.018 0.346 0.036 0.386 0.078 0.428 0.120 0.470 0.162 0.515 0.207 ethanol 0.195 - 0.308 - 0.318 0.0 10 0.328 0.020 0.350 0.042 0.394 0.086 0.441 0.133 0.494 0.186 0.550 0.242 iso-propyl alcohol 0.195 - 0.308 - 0.319 0.01 1 0.330 0.022 0.352 0.044 0.399 0.091 0.447 0.139 0.500 0.192 0.555 0.247 n-propyl alcohol 0.195 - 0.308 - 0.319 0.01 1 0.330 0.022 0.354 0.046 0.40 1 0.093 0.452 0.144 0.506 0.198 0.563 0.255 t-butyl alcohol 0.197 - 0.308 - 0.319 0.01 1 0.331 0.023 0.355 0.047 0.403 0.095 0.455 0.147 0.508 0.200 0.568 0.260 - - 0.024 39 0.048 98 0.098 46 0.197 87 0.296 78 0.395 74 0.494 55 - - 0.024 43 0.049 03 0.098 86 0.198 01 0.297 51 0.396 06 0.493 96 - - 0.024 65 0.049 12 0.098 91 0.198 27 0.297 61 0.396 77 0.495 30 - - 0.024 51 0.048 98 0.098 60 0.198 56 0.296 78 0.395 96 0.495 09 - - 0.024 45 0.048 90 0.098 60 0.198 00 0.297 63 0.396 86 0.496 60 - - 1.19 1.21 1.24 1.27 1.31 1.33 1.36 - - 1.32 1.34 1.37 1.41 1.45 1.52 1.59 - - 1.41 1.43 1.45 1.48 1.52 1.57 1.62 - - 1.46 1.47 1 S O 1.53 1.57 1.62 1.67 - - 1.48 1 S O 1.53 1.56 1.60 1.63 1.69S.P. MISHRA A N D R. A. SINGH 1781 TABLE 7. (cont.) C(D)O activity" C(IAb C(DI2) C(D) Kc /mol dm-3 /counts min-l /mmol dmP3 fmmol dm-3 /mol dmP3 /dm3 mol-l ref. 0.000 00 0.024 36 0.048 88 0.098 49 0.197 89 0.297 67 0.397 10 0.496 34 ref. 0.000 00 0.024 58 0.049 00 0.098 48 0.197 97 0.297 20 0.396 66 0.495 95 580 907 940 975 1046 1191 1345 1509 1681 525 82 1 852 884 948 1083 1221 1368 1525 iso-butyl alcohol - 0.197 0.308 0.319 0.01 1 0.33 1 0.023 0.355 0.047 0.405 0.097 0.457 0.149 0.512 0.204 0.57 1 0.263 n-butyl alcohol 0.197 - 0.308 0.320 0.012 0.332 0.024 0.356 0.048 0.406 0.096 0.458 0.150 0.513 0.205 0.572 0.264 - - - 0.024 35 0.048 86 0.098 44 0.197 79 0.297 52 0.396 90 0.496 08 - 1 S O 1.52 1.55 1.58 1.62 1.67 1.72 - 0.024 57 0.048 98 0.098 43 0.197 87 0.297 05 0.396 45 0.495 69 - 1.52 1.54 1.57 1.61 1.64 1.68 1.73 " Error limit is + 5 counts min-l; ,!I error limit is +0.001 mmol dmW3 iodine complex formation and alcohol-alcohol hydrogen-bonding equilibria is ex- pected to occur in the solubility experiments.The increase in radioactivity, i.e. the total iodine concentration in solution, with the donor concentrations of various alcohol-iodine systems in n-heptane is reported in tables 5-7 and shown in fig.3 (25 "C). An interesting feature of these curves is the slight departure from linearity at higher alcohol concentrations. Such behaviour was not observed in the case of aromatic-iodine complexes (vide infra). Such a departure might be due to a small but significant production of I; species at higher alcohol concentrations. To check this possibility parallel experiments were made with unlabelled TMAI, and I; was analysed by spectr~photometry.~~? 3 4 3 38 The I; species was produced at higher alcohol concentrations but in an amount less than the experimental error in determining the concentration ( mol dmP3) of iodine. Thus, such a curvature is real. This conforms to the proposition by Tamres et al.34 that the complexes in concentrated alcohol solutions are stronger than those in dilute solutions owing to increased hydrogen bonding.However, it clearly opposes their concept of weighted average for K,. In fitting the data for such systems, the total concentration of iodine determined by the radiotracer method was analysed by a numerical least-squares technique on an ICL 1900s computer as a quadratic function of C(D), and a limiting slope at C(D) = 0 was obtained.'' These results were further checked by calculating K , point-by-point using eqn (6), and the values thus obtained were plotted against C(D). A linear plot was obtained, and the intercept at C(D) = 0 gives the limiting value of K,, supposedly for the monomer-iodine complex (fig.4). The results for various alcohol-iodine systems are reported in table 8. Equilibrium1782 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES 0.4 m E -0 - 2 E 6 0.3 . n 3 v 0.2 I I I 0 0.2 0.4 C (D)/mol dm-3 FIG. 3.-Plot of total iodine concentration [CD(I,)] determined by the radiotracer technique against concentration of donor [C(D)] in n-heptane at 25 OC: (1) methanol, (2) ethanol, (3) iso-propyl alcohol, (4) n-propyl alcohol, ( 5 ) t-butyl alcohol, (6) iso-butyl alcohol, (7) n-butyl alcohol. I I I 1 I C(D)/mol dm-3 0 0.2 0.4 FIG. 4.-Plot of e P P against donor concentration [C(D)J of various alcohol-iodine complexes in n-heptane at 25 OC: (1) methanol, ( 2 ) ethanol, (3) iso-butyl alcohol, (4) n-propyl alcohol, ( 5 ) t-butyl alcohol, ( 6 ) iso-butyl alcohol, (7) n-butyl alcohol.S. P.MISHRA A N D R. A. SINGH 1783 TABLE 8.-EQUILIBRIUM CONSTANT AND ENTHALPIES OF COMPLEXES OF IODINE WITH ALCOHOLS IN n-HEPTANE SOLUTION, DETERMINED BY THE RADIOTRACER SOLUBILITY METHOD ~~ donor K, (25 OC)/dm3 mol-l - AH/kcal mol-l methanol ethanol 1.41 f 0.05 (1 .29a) 1.59 f 0.06 (1 S O a ) 3.10 f 0.2 (3.0a) (3.2a) iso-propyl alcohol 1 . 7 0 f 0 . 0 8 3 . 5 f 0.2 n-propyl alcohol 1.76 f 0.08 (1 .03b) 3.5 k 0.2 t-butyl alcohol 1.77 f 0.08 3.6 & 0.2 iso-butyl alcohol n-butyl alcohol 1.80 f 0.06 (1 .37b) 1.84 & 0.06 3.6 & 0.2 (3.47 3 . 7 & 0.2 Values in parentheses refer to the literature data: a ref. (34); R. M. Keefer and L. J. R. M. Keefer and L. J. Andrews, J. Am. Chern. Andrews, J. Am. Chem. SOC., 1952,75,3561; SOC., 1955, 77, 2164.constants (K,) and enthalpy values (- AH) for various alcohol-iodine systems are also reported in table 8, along with the literature data whenever available. The values of K, and -AH obtained by different methods, viz. the radiotracer technique, the spectral-solubility method and the Benesi-Hildebrand procedure, can be outlined as follows ( K , in dm3 mol-l, -AH in kcal mol-l): (i) methanol-iodine, K, 1.41 kO.05 (1 .29,34 0.6F4), -AH 3.1 f 0.2 (3.034); (ii) ethanol-iodine, K, 1.5 k 0.06 ( 1.50,34 0.94 k0.0634), -AH 3.40 f0.2 (3.2,34 3.534); (iii) iso-propyl alcohol, K, 1.70k0.08, -AH 3.5 & 0.2; (iv) n-propyl alcohol, K, 1.76 f 0.08, 1 .03,9 - AH 3.5 f 0.2, etc. It is evident that the values obtained by the radiotracer technique are comparable with the literature data.The enthalpy values are also in good agreement with the literature data. Note that the enthalpies vary in the same order as the ionization potential of the We expect similar behaviour based on Mulliken's theory of molecular complexes. CONCLUSIONS The findings presented in this paper clearly show that the radiotracer technique deserves consideration as an alternative to the spectral method for evaluating thermodynamic parameters of molecular complexes in solution. The present method, which-combines the advantages of the constant-activity method and the sensitivity of the radiotracer technique yields the various thermodynamic parameters separately. Also, prior knowledge of the isosbestic point is not required for the determination of the iodine concentration, as is the case in the spectral method.Furthermore, this method would be of great help in deducing the thermodynamics of molecular complexes, which is otherwise rather difficult by spectral studies owing to the overlap of optical bands of the donor, acceptor and/or complex. In situations where both techniques can be used, complementary data will be available to aid in the better characterization of the systems. In addition to these advantages, the radiotracer technique allows the estimation of very low concentrations not normally tackled by1784 RADIOTRACER STUDIES OF MOLECULAR COMPLEXES other methods. Thus, the present paper is a novel application of the radiotracer technique to the study of molecular complexes. We thank Prof.B. M. Shukla, Head of the Department of Chemistry, B.H.U., for providing facilities. This work was supported by the research fellowships from the D.A.E. and C.S.I.R., which are gratefully acknowledged by R.A.S. R. S. Mulliken and W. B. Person, Molecular Complexes (John Wiley, New York, 1969). Advances in Chemical Physics, vol. 12, Intermolecular Forces, ed. J. 0. Hirshfelder (Interscience Publishers, New York, 1967). K. Morokuma, Ace. Chem. Res., 1977, 10, 294. Spectroscopy and Structure of Molecular Complexes ed. J. Yarwood (Plenum Press, London, 1973). G. Briegleb, Elektronen-Donator- Acceptor Komplexe (Springer Verlag, Berlin, 1961). R. 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ISSN:0300-9599    

DOI:10.1039/F19827801767

出版商:RSC    

年代:1982

数据来源: RSC