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/F198278FX005

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/F198278BX007

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, II Journal of The Chemical Society, Perkin Transactions, I Journal of The Chemical Society, Perkin Transactions, II 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. Journal of The Chemical Society, Chemical Communications Journal of The Chemical Society, Dalton Transactions 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, Perkin Transactions, I and II These are, respectively, the organic chemistry and the physical organic chemistry sections of the Journal.P A R T I (organic and bio-organic chemistry) is designed to contain papers on all aspects of synthetic, and natural product organic and bio-organic chemistry and to deal with aliphatic, alicyclic, aromatic, carboncyclic and heterocyclic compounds. Papers on organometallic topics are considered for either the Dalton or the Perkin Transactions. 1PART I I (physical organic chemistry) is for papers on reaction kinetics and mechanistic studies of organic systems and the use of physico-chemical, spectroscopic, and crystallographic techniques in the solution of organic problems. Notice to Authors (1) Although authors need not be members of the Royal Society of Chemistry it is hoped that they will be.(2) Authors must indicate the Part of the Journal they wish their paper to appear in. This preference will be respected unless it is obviously erroneous in terms of the scientific content of the paper. (3) Since all papers will be subjected to refereeing, in parallel, by two independent referees, the original typescript (quarto or A4 size) and two good-quality copies should be provided. (4) All papers should be sent to the Director of Publications, The Royal Society of Chemistry, Burlington House, Piccadilly, London W 1V OBN. (5) For details of manuscript preparation, preferred usages, etc. the Instructions to Authors, previously available from the Faraday Society, and now obtainable from The Royal Society of Chemistry, should be consulted.(6) The Society will adopt the following abbreviations for the new journals in all its publications. J. Chem. SOC., Chem. Commun. J . Chem. SOC., Dalton Trans. J . Chem. SOC., Faraday Trans. I J. Chem. SOC., Faraday Trans. 2 J . Chem. SOC., Perkin Trans. I J . Chem. SOC., Perkin Trans. 2 * The author to whom correspondence should be addressed is indicated by an asterisk after his name in the heading of each paper. 11THE F A R A D A Y D I V I S I O N O F THE R O Y A L SOCIETY O F CHEMISTRY GENERAL D I S C U S S I O N N O 73 VAN D E R WAALS MOLECULES St Catherine's College, Oxford, 5-7 April 1982 Organising Committee Dr E B Smith (Chairman) Professor A D Buckingham Dr G Duxbury Dr B J Howard Dr G C Maitland There has been increasing interest in recent years in molecules whose components are bound only by relatively weak Van der Waals Interactions The study of their structure and properties, and of the binding forces involved, has been stimulated both by new developments in spectroscopy and molecular beams specifically designed to study such molecules, and by the concerted application of a range of more standard techniques to these molecules high resolution spectroscopy, inelastic scattering, theoretical calculation.measurement of thermophysical properties The aim of this discussion is to bring together workers in these diverse fields, to highlight the complementary nature of the information obtained using the various techniques and t o examine the direction that future work should take t o increase our understanding of these molecules The programme and application form may be obtained from: Mrs Y.A. Fish, The Royal Society of Chemistry Burlington House, London W1V OBN F A R A D A Y D I V I S I O N O F T H E R O Y A L S O C I E T Y O F C H E M I S T R Y A S S O C I A Z I O N E 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 I i T i D E C H l M l E P H Y S I Q U E F A R A D A Y D I S C U S S I O N N O . 7 4 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 Electron and Proton Transfer University of Southampton, 14-1 6 September 1982 This meeting will be concerned with fundamental aspects of the chemical kinetics of electron and proton transfer reactions in solution and with particular reference to well defined biological systems.Attention will be focused on (i) the theory of charge transfer, (ii) critical experiments designed to test those theories and (iii) their application to the understanding of charge transfer reactions in molecules of biological interest. The meeting will encompass well characterised reactions in solution, redox reactions and elementary biochemical reactions, particular attention will be paid to isotope effects, to electron and proton tunnelling, to intermolecular and intramolecular transfers and to related questions concerning the organisation of biological systems. Among those who have agreed to take part are R. A. Marcus, R. R. Dogonadze, H. Gerischer, J. Jortner, R. M. Kuznetsov, N.Sutin and R. J. P. Williams. The preliminary programme may be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry Burlington House, London W1V OBN ... 111FARADAY DIVISION O F THE ROYAL SOCIETY OF CHEMISTRY SYMPOSIUM NO. 1 7 The Hydrophobic Interaction University of Reading, 15-1 6 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 descri pt ion.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 Hebrew University, Jerusalem B. J. Berne Columbia University, New York S. 0. Christian Oklahoma University The preliminary programme may be obtained from: Mrs Y. A. Fish, The Royal Society of Chemistry Burlington House, London W1V OBN THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY GENERAL DISCUSSION NO. 75 Intramolecular Kinetics University of Warwick, 18-20 April 1983 Organising Committee Professor J.P. Simons (Chairman) Dr M. S. Child Professor R. J. Donovan Dr G. Hancock Dr D. M. Hirst Professor K. R. Jennings Dr R. Walsh Experimental and theoretical interest in the time-dependent behaviour of isolated molecules, radicals or ions is strong and increasing. The Discussion will be concerned with the kinetics of processes which occur in isolated species following their preparation in states with non-equilibrium energy distributions (e.g. by photon absorption or collisional activation). Topics covered will include: ( a ) theoretical and experimental studies of energy redistribution in isolated species; (6) observation and theoretical modelling of the competition between intramolecular energy redistribution and radiative decay or radiationless processes (e.g.internal conversion, fragmentation, isomerisation). Contributions for consideration by the Organising Committee are invited. Titles should be submitted as soon as possible and abstracts of 300 words by 31 May 1982. Full papers for publication in the Discussion Volume will be required by 15 December 1982. Titles and abstracts should be sent to: Professor J. P. Simons, Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD. I VFARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY 1982 ANNUAL CHEMICAL CONGRESS 30 March to 1 A p r i l University of Aston in Birmingham The Faraday Division is arranging a symposium jointly with the NMR Group at the Society's Annual Congress on: NMR IN PHYSICAL CHEMISTRY Further information and application forms may be obtained from Dr J.F. Gibson, The Royal Society of Chemistry, Burlington House, London W1V OBN THE ROYAL SOCIETY OF CHEMISTRY: MOLTEN SALTS GROUP IONIC TRANSPORT IN MELTS 29-30 March 1982 University of Aberdeen The meeting will review recent developments including computer simulation, fast ion conduction in glasses and electrochemical applications. For further information and offers of papers contact D. G. Winter, Warren Spring Laboratory, PO Box 20, Gunnels Wood Road, Stevenage, Herts SG1 2BX FARADAY DIVISION INFORMAL AND GROUP MEETINGS Division Half day symposium Including the Liversidge Lecture by D W Turner To be held at Imperial College, London on 16 March 1982 Further information from Mrs Y A Fish, The Royal Society of Chemistry, Burlington House, London W1 V OBN Polymer Physics Group with Macro Group UK and the Plastics and Rubber Institute Deformation, Yield and Fracture of Polymers To be held at Churchill College, Cambridge on 29 March to 1 April 1982 Further information from Mr J.N. Ratcliffe, Plastics and Rubber Institute, 11 Hobart Place, London SW1 W OHZ Division with the NMR Group Annual Congress NMR in Physical Chemistry To be held at the University of Aston in Birmingham on 30 March to 1 April 1982 Further information from Dr J. F. Gibson, The Royal Society of Chemistry, Burlington House, London W1 V OBN CoNoid and Interface Science Group with the SCI Stability and Control of Surfactant Concentrates To be held at the University of Nottingham on 5-6 April 1982 Further information from Dr M J Jaycock, Department of Chemistry, University of Technology, Loughborough LE11 3TU VFARADAY DIVISION INFORMAL AND GROUP MEETINGS Division: Half day symposium Laser Spectroscopy (including the Centenary Lecture by T.Oka) To be held at University College, London on 28 April 1982 Further information from Mr S. S. Langer, The Royal Society of Chemistry, Burlington House, London W1V OBN industrial Physical Chemistry Group with the British Corrosion Society Transport Mechanisms and Structure in Corrosion Films To be held at the University of Surrey, Guildford in July 1982 Further information from Professor T. Edmonds, BP Research Centre, Chertsey Road, Sunbury on Thames Tw16 7LN 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 8-1 0 September 1982 Further information from Dr W D Cooper, Department of Chemistry, University of Edinburgh, West Mains 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 Neutron Scattering Group and Polymer Phys~cs 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 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 Or P. Richmond, Unilever Research, Port Sunlight, Wirral, Merseyside L62 3JW viEvaluated Kinetic Data for High Tempemture Reactions, Volume 4: Homogeneous Gas Phase Reactions of Halogen- & Cyanide-Containing Species D.L.Mulch, J. Duxbury, SJ. Grant, D.C. Montague Department of Physical Cbemimy, University of Lee& U.K. The results of decades of research are now located in this one convenient publication! Published by the American Chemical Society and the American Institute of Physics for the National Bureau of Standards, this monograph presents in 721 pages, kinetic data for 300 homogeneous gas phase reactions involving halogens, the cyanide radical, and their compounds. Wherever possible, the data have been critically evaluated to give the best estimates of reaction rate parameters and their associated limits and temperature ranges.The supplement also offers relevant thermodynamic data, discusses it thoroughly, and lists recommended rate constants for each reaction in tabular form. A comprehensive bibliography lists all pertinent literature in the field. American Chemical Society, B&J Business Operations 1155 16th St., N.W., Washington, D.C. 20036 Please send m e ~ hardcover copies of Evaluated Kinetic Data for High Temperature Reactions, Volume 4: Homogeneous Gas Phase Reactions of Halogen- and Cyanide-Containing Species at $80.00 each. Please include payment with order.. Name City State Zip 'Foreign orders add $4 00 each for postage and handling. Foreign payment must be made In U.S. currency, by international money order, LNESCO coupons, or order through your subscription agencv California residents, add 6% state use taxNOTES 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 1500 words or word-equivalents. Nomenclature and Symbolism. For many years the Society has actively encouraged the use of standard IUPAC nomenclature and symbolism in its publications as an aid to the accurate and unambiguous communication of chemical information between authors and readers. In order to encourage authors to use IUPAC nomenclature rules when drafting papers, attention is drawn to the following publications in which both rules themselves and guidance on their use are given.Physicochemical Quantities and Units. Manual of Symbols and Terminology for Physicochemical Quantities and Units. (Pure and Appl. Chem., Vol. 51, No. 1, 1979, pp. 1-41. Also available as a soft-cover booklet from Pergamon Press, Oxford.) Surface Chemistry. ‘ Definitions, Terminology, and Symbols in Colloid and Surface Chemistry - I.‘ (Purr 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 Purr and Applied Chemistry in recent years: activities;* chromatography; electrochemistry; electron spectroscopy; equilibria, fluid flow; ion exchange; liquid-liquid distribution; molecular force constants; Mossbauer spectra; nuclear chemistry; pH; polymers; quantum chemistry; radiation;* Raman spectra; reference materials (recommended reference materials for the realization of physico- chemical properties : general introduction, enthalpy, optical rotation, surface tension, optical refraction, molecular weight, absorbance and wavelength, pressure-volume- temperature relationships, reflectance, potentiometric ion activities, testing distillation columns); solution chemistry; spectrochemical analysis ; surface chemistry; thermo- dynamics, and zeolites. Finally, the rules for the naming of organic and inorganic compounds are dealt with in the following publications from Pergamon Press: ‘Nomenclature of Organic Chemistry, Sections A, B, C, D, E, F, and H’, 1979. ‘Nomenclature of Inorganic Chemistry’, 1971. pp. 71L90.) ... Vlll

ISSN:0300-9599    

DOI:10.1039/F198278FP009

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraduj Trans. I , 1982, 78, 313-321 Effect of Temperature on the Limiting Excess Volumes of Amines in Aqueous Solution BY MURLIDHAR V. KAULGUD," VIJAY S. BHAGDE A N D ANJALI SHRIVASTAVA Department of Chemistry, Nagpur University, Nagpur 440 010, India Received 29th July, 1980 Limiting partial molal volumes (v:) from measurements of density in the concentration range 0.002-0.1 mol dm-, for aqueous solutions of NH,, MeNH,, EtNH,, n-PrNH,, N-BuNH,, t-BuNH,, (Me),NH and (Me,)N have been obtained at 5, 15 and 25 "C. The method of Cabani and coworkers for applying hydrolysis corrections in dilute solutions of amines had to be modified slightly in order to make it applicable to the lower concentration range used in this work. It is shown that the modified method applies equally well to both the low (0.002-0.1 mol drn-,) and slightly higher (0.1-0.3 mol dm-3) concentration ranges, yielding identical results.The present results show that limiting excess volumes become less negative with increasing temperature for lower amines, but higher amines show an opposite trend. Comparison of volumetric properties of amines with those of alcohols indicates that in monofunctional solutes the functional group plays an important role in solute-solvent interaction. It is found in particular that the -NH, group tends to offset the structure promoted by the hydrophobic group in amines. Volumetric and compressibility behaviour at 20 "C of some straight chain, mono- and di-alkylamines has been previously1* , reported from this laboratory and the results discussed from the point of view of solute-solvent interaction.The limiting values were obtained by extrapolation from higher concentrations (> 0.3 mol dm-3) where the hydrolysis of amines was negligible. The results obtained indicated similarity to the behaviour of alcohols and hence amines were also said to behave as hydrophobic structure formers. Cabani et ~ 1 . ~ have studied the volumetric behaviour of amines at 25 O C in a low concentration range (but > 0.002 mol dm-3). A special correction procedure was developed4 by them to ascertain the true value of the limiting partial volume of the amine molecule from that observed by subtracting the contribution of the ions produced due to hydrolysis. Though various physicochemical properties of solutions of alcohols and amines show similarity, there are still subtle differences.Thus alcohols produce a rise in the temperature of maximum density (t.m.d.) of water while amines lower it.576 The characteristic constant ' a ' in the expression for the structural contribution of the solute to the change in t.m.d. of water A8,,, = ax+5x2 (x = mole fraction of the solute), which is an index of its structure strengthening ability, is larger for alcohols than for amines (for n-BuOH a = 237 OC, and for n-BuNH, a = 173 "C). Similarly, ACF, AGP and ASP7q8 values for amines are much less than for the corresponding alcohols. Further, solid clathrates have been isolated for some amine~,~,lO but not yet for alcohols. The polar group thus seems to be playing a role in this case.The volumetric properties of alcohols at different temperatures are well documented. Similar studies in amines especially at low temperatures are lacking, probably on account of the difficulties associated with the correction for hydrolysis. It was hence thought worthwhile to evaluate the true limiting excess volumes of amines at 5, 15 and 25 OC and see how they compare with those of alcohols. 313314 EXCESS VOLUMES OF AMINES EXPERIMENTAL The differential float balance method described and used earlier" for measurements at 25 "C for solid solutes in aqueous solution was modified in the present work as follows: A single float method using a Sartorius model 2474 single pan semi-microbalance having facility to weight below the pan was adopted. The glass float weighted with mercury (volume z 170 cm3) was suspended by a thin nylon thread into distilled water kept in a stainless-steel cylinder.The steel cylinder in turn was kept in a well-insulated and covered constant-temperature bath (capacity 30 dm3) whose temperature was regulated by circulating liquid from a MK-70 cryostat through large copper coils kept in the bath. Replacement of the glass cylinders of the earlier work" by the stainless-steel cylinder in this work meant faster thermal equilibration between solution and bath and reduction in experimentation time, leading to better stability of the bath temperature around the nominal value. n-Propylamine (Riedel-De-Haen), n-butylamine (Merck, Germany) and t-butylamine (Fluka, pract.), were dried over KOH for ca.48 h, distilled twice and the miadle fraction was used to prepare stock solutions. Ammonia (local grade), methylamine (40%), ethylamine (70 %), dimethylamine (40 %) and trimethylamine (50 %) (all Riedel-De-Haen) solutions were slightly warmed and the amine vapours were led into doubly distilled water to obtain a stock solution. Dilutions were made in situ and the concentration of the solution was determined by volumetric titration with standard AnalaR HCl using a suitable indicator. To begin with water was degassed by passing nitrogen and to avoid carbonation a layer of nitrogen gas was maintained over the solution throughout the experiment. Amine solutions of increasing concentration were obtained by adding successively 10 cm3 of concentrated amine solution to the water in the steel jar and estimating the concentration of the resulting solution by titration with standard hydrochloric acid solution by withdrawing two aliquots of 5 cm3 each.Addition of 10 cm3 of solution and subsequent withdrawal of the same for estimation each time ensured that the float dipped into the water at the same level and the small errors arising due to additional nylon thread dipping in the event of successive addition could be avoided. The volume of the plunger was determined accurately first by weighing it in air and then in distilled water held at 5, 15 and 25 OC. Measurement of the density of a given amine solution were obtained in 2 or 3 different runs, each run comprising of 5 or 6 different concentrations. The difference in the weight of the plunger in water and solution at the lowest concentration was ca.10-15 mg yielding a difference in density of 1 unit in the fifth decimal place. This coupled with the readability of the balance to 0.01 mg results in an accuracy of ca. 0.1 ppm in the measured density value. The temperature fluctuation as read on the Beckman thermometer was kO.002 O C . From the measured density, apparent molal volumes were first calculated and these were subjected to the hydrolysis correction as described below. METHOD OF CORRECTION FOR HYDROLYSIS Cabani et al.4 adopted the following procedure for hydrolysis correction. They assume that the hydrolysis of the amines (B) in aqueous solution at a concentration c takes place through a hypothetical neutral aquo-amine species BHzO as: (1-a)c ac ac where a is the degree of hydrolysis $v = z-(;- A4 I)-.I000 The fictitious apparent molal volumes &!'s, which are calculated by introducing in eqn (2) the molecular weight of the hypothetical species BHzO, can then be considered to be made up of proportionate contributions of #V(BH,o), &(BH+) and &(OH-), that is: hbSd = (1 -a) 4v(BHzo) +a(&(BH+) +&(OH-))= (1 -a) dV(BH,o) +a&(BH+OH-) - (3)M. V. KAULGUD, V. S. BHAGDE A N D A. SHRIVASTAVA 315 Since BHZO is a neutral species, the value of &(RH,o) can be assumed to depend linearly on the (4) concentration as: where &(BHro) is the value of &(BE,zo) at infinite dilution and h is the slope factor. BH+ and OH-, the hydrolytic products, were assumed to form partners of a strong electrolyte.Hence the value of &,(BH+OH ) at any concentration was found by using the modified Debye-Hiickel theory as applied to the apparent volume of 1 : 1 electrolyte: 4V(BH,<)) = -a) $,,, - 1.8682/~ = 4; + h+c &(BH+OH-) = #;(BH+oH-) + I .868y'(ac) + h+ca. ( 5 ) (5') where h+ is called the deviation constant,12 or in the present case Introducing eqn (4) and (5') in eqn (3) one gets: &bsd = (1 - a)&(BHno) + hc( 1 - + 4;(BH+oH-) + 1.8682/(ac) + h+a2c. On rearranging and dividing both sides by (1 -a) one obtains: where K; is thermodynamic hydrolysis constant and 4; represents the numerator of the 1.h.s. of eqn (6). The a values at the actual ionic strength were calculated by an iterative procedure using the known values of the thermodynamic hydrolysis constant K z and molal activity coefficient values given by the Debye-Huckel limiting law.The term &BH+oH- at 25 O C was calculated using values of &(,,,+,,-, = 16.61 cm3 mol-l and 4",,,,+,,-, = 4.60 cm3 mol-l taken from literature, while the values of &(BH+CI-) were experimentally determined. For each amine considered and for each concentration the term K z h+ is constant and negligible. Hence &(BH,o) is determined as an intercept by extrapolation of the linear plots of d:/( 1 -a) against (1 -a) c to zero concentration. The limiting partial molal volume of free amine $"VcB, is then calculated by subtracting the molar volume of pure water from &(BH,o). But if the above method is applied to our data for n-propylamine at 25 "C and in the low-concentration range, the points in the plot of &/( 1 - a) against (1 - a) cdo not lie on a straight line as shown in fig. 1 (a).Such non-linear trends are also observed for curves at 5 and 15 O C , in the lower concentration range. Evaluation of limiting volumes by use of eqn (7) is thus difficult, as unequivocal extrapolation to zero concentration cannot be made. Similar difficulty was experienced in the previous work from this laboratory while evaluating true limiting partial molal compressibilities from measurements in dilute aqueous amine s01utions.l~ Kaulgud et al. while deriving their equations for hydrolysis correction for compressibilities on the same lines as Cabani et al. tor volume, slightly modified their equation so as to yield linear extrapolations.We felt that a similar procedure could be adopted to handle density data at lower concentrations to arrive at q5;(B) unambiguously. By multiplying both sides of eqn (7) by (1 -a)c and neglecting as before the term K z h+ one gets: 4;c = (1 - ~ ) C [ ~ ; ( B H ~ O ) + ( l - ~ ) C h l . (8) Eqn (8) shows that provided the solute-solute interaction term (1 -a) ch in the square bracket is negligible, one should get a straight line passing through origin if 4; c is plotted against (1 -a) c with &(,,, o) as slope. In order to see if this is so, we have plotted in fig. l(6) points showing 4: c as a function of (1 -a) c for n-propylamine at 25 OC (full circles). The a values were evaluated by the same iterative method adopted by Cabani mentioned above.The PKb values needed for this calculation were taken from the 1iterat~re.l~ It can be seen that all the points lie on a straight line passing through the origin. Moreover, the true limiting volume of the unhydrolysed n-propylamine, 74.18 cm3, obtained by subtracting the molar volume of pure water from the slope of this straight line, agrees very well with the value 74.12 obtained by316 EXCESS VOLUMES OF AMINES ii *> 8 16- 4 12 -3 8 - 2 I I I I I I I I , 0 LO 80 120 160 ( 1 -a) c~ 103 FIG. 1 .-(u) Plot of #;/( 1 -a) against (1 -a) c for n-propylamine at 25 O C . [The scale for abscissae is same as the inner scale in (b).] (6) Plots of 4; c against (1 -a) c for n-propylamine at 25 O C . Our data (@); inner scale for ordinate and abscissae).Cabani's data (0); (outer scale for ordinate and abscissae) (see text). Cabani et al.3 by applying eqn (7) to their density measurements. To further verify if (1 -a) ch remains negligible at higher concentrations, we have applied eqn (8) to Cabani's density data* for n-propylamine at 25 "C in the concentration range up to 0.3 mol dmP3. These points are plotted in fig. 1 (b) (open circles), with a four-fold contraction in both the abscissa and ordinate. These points are seen to lie on the same straight line as ours, indicating the same value of slope and consequently the same value for d",,,,. This method applied to Cabani's density data for MeNH,, EtNH, and n-BuNH, at 25 OC at higher concentration (up to 0.3 mol dm-3) also showed good straight lines and yielded an almost identical value for d",,,, as that found by Cabani.3 These facts show the validity of eqn (8) at higher concentrations.Similar straight lines passing through the origin were obtained for ammonia and other amines handled in this work at 25 OC and our values agree with Cabani's values in all cases (table 1). We have plotted 4;c against (1 - a ) c in fig. 2 for ethylamine at 5, 15 and 25 OC and in fig. 3 for ammonia, methyl-, n-propyl-, dimethyl- and n-butyl-amine at 5 OC, to indicate the applicability of eqn (8) at lower temperatures. Though the straight line plots at 15 "C are not shown, they are all found to be so. Values of limiting partial molar volume thus evaluated are given in table 1. The evaluation of the slope was carried out in all the cases by fitting the points into a straight line by the method of least squares.The excess molar volumes at infinite dilution v:E were calculated as = Vf- V: with F: = 4;. The molar volumes at 25 "C are taken from the literature3 and at 5 and 15 OC are evaluated from density data reported in the 1iterat~re.l~ * We are grateful to Prof. Cabani for making these data available.M. V. K A U L G U D , V. S. B H A G D E A N D A. SHRIVASTAVA 317 TABLE I.-vALUES OF THE LIMITING PARTIAL MOLAL VOLUME OF AMINES AT DIFFERENT TEMPERATURES amines 25 "C 5 "C 15 "C 25 OC Cabani, ref. (3) ammonia me thy lamine ethylamine n-propy lamine n-butylamine t-but ylamine dime t hy lamine trimet hylamine 22.01 f 0.2 36.75 f 0.08 54.95 f 0.03 7 1.64 & 0.024 87.23 f 0.04 89.66 & 0.05 58.62 f 0.12 77.85 f 0.1 22.86 & 0.1 39.79 f 0.04 56.94 f 0.02 72.56f0.012 89.43 f 0.02 90.1 f0.04 58.86 f 0.1 78.32 f 0.1 24.80 f 0.2 41.98 f 0.02 58.85 f 0.02 74.18f0.01 89.84+0.01 91.24 f 0.02 59.44 f 0.1 79.34 f 0.08 24.85 41.68 58.37 74.12 89.80 59.80 78.40 - FIG.2.-Plots of q5$c against (1 - a ) c for ethylamine at 25 (O), 15 (0) and 5 O C ( 0 ) (origin shifted towards right by 5 and 10 units for 15 and 5 O C , respectively). RESULTS AND DISCUSSION The straight lines passing through the origin obtained by application of eqn (8) for all amines indicate that the solute-solute interaction term (1 -a)ch in the square bracket is negligible up to quite high concentration (0.3 mol dm+) and even at low temperatures. The h values for amines at 25 OC tabulated by Cabani4 are of the order318 EXCESS VOLUMES OF AMINES (1 -a) x 1 0 3 FIG.3.-Plots of 4; c against (1 -a) c at 5 OC for NH, (O), MeNH, (a), Me,NH (A), n-PrNH, (0) and n-BuNH, (A). of unity with negative sign for amines handled in this work. Hence the maximum value of the interaction term (1 - a) ch z - 0.3 at a concentration of 0.3 mol dmP3 where a = 0.05. This value is very small as compared with #$(BH,o) which is of the order of 50-120 cm3 and the term (1 - a ) ch is, at the worst, of the order of magnitude of experimental errors. Hence, this modification of the original method is applicable for the determination of reliable 4: values even by using data at higher concentrations. It is also found that 4; values determined using eqn (8) are independent of slight uncertainty (ko.1 if any) in the measurement of p& values for amines.This fact is to be contrasted with the Cabani's procedure, which is shown to be sensitive to small uncertainties in p&, values. In this sense, the variant proposed by us in this work [eqn (S)] can be said to be better suited for reliable evaluation of #$(R). Table 1 shows that the limiting partial molal volumes of ammonia and t-BuNH, are practically constant between 5 and 25 OC, while for other amines these show a distinct increase with temperature. This is to be contrasted with the behaviour of alcohols for which Vz are practically constant in the temperature range 0-30 OC.16 Both amines and alcohols show negative limiting excess volumes ( V,""), the magnitude being larger for amines.This means that the extent of accommodation of amines in open cage structures existing in water is greater than for alcohols. Some insight into the further differences in behaviour of alcohols and amines might be gained by considering the temperature coefficient of the limiting excess volumes d "/dT = d( - V,O)/dT where V i is the limiting partial molar volume and V: is the molar volume of the pure liquid. Neal and Goring" studied the temperature coefficients of the apparent specificM. V. KAULGUD, V. S. BHAGDE A N D A. SHRIVASTAVA 3 19 TABLE 2.-TEMPERATURE COEFFICIENTS OF THE LIMITING EXCESS VOLUMES FOR AMINES AND ALCOHOLS amines 102 d V i E/dT alcohols lo2 dV'E/dT ammonia 2.2 met hylamine 16.6 methyl alcohol - 4.6" e t h y lamine 9.1 ethyl alcohol - 6.0" n-prop ylamine 1.2 n-propyl alcohol - 7.0" t-butylamine - 3.2 t-butyl alcohol - 10.3b dimethy lamine - 8.8 trimethylamine - 6.8 n-butylamine - 6.5 n-butyl alcohol - 4.3" (' Ref.(16). Evaluated from the vo against Tcurves of F. Franks and H. G. Smith, Trans. Furaduy SOC., 1968, 64, 2962. W 0 5 15 25 T/"C FIG. 4.-Vanation of vF with temperature for NH, (a), MeNH, (O), EtNH, (A), n-PrNH, (O), n-BuNH, (A), t-BuNH, (01, Me,NH (8) and Me,N (m). volumes dd,/dT of a large number of non-electrolytes in water at low concentration and concluded that the difference (dd,/dT- d V,/dT) (where d V,/dTis the temperature coefficient of the specific volume of the pure liquid) is positive for structure-breaking solutes and negative for structure-making solutes. The correspondence between d r: E/d T and (d@,/d T - d K/d T ) is quite obvious. Values of d V: E/d T for ammonia and all amines handled in this work (table 2) can hence be subjected to similar interpretation.Since alcohols are well known to be structure stabilizers, they show negative values for d E/dT as expected (table 2). Amines, which are, like alcohols,320 EXCESS VOLUMES OF AMINES monofunctional aliphatic organic solutes, would be expected to show similar behaviour. Surprisingly, the first three members, methyl-, ethyl- and n-propyl-amine and ammonia show positive d Fi "/d7', whereas the rest of the amines including di- and tri-methylamine show negative values. Since in both alcohols and amines the hydrophobic part is the same, the observed differences must be due to the different ways in which the -NH, and the -OH groups interact with solvent water.In fact, the observed results can be rationalized by assuming that the water-structure compatible OH group reinforces water structure and the NH, group disrupts it in its vicinity due to its inability to participate in cooperative hydrogen bonding with water. It can be seen that ammonia, which can easily be accommodated into the natural cavities in water, acts as a weak structure breaker. In the case of methylamine, the effect of NH, group outweighs that of a lone CH, group on water and the overall molecule turns out to be a structure breaker according to the sign of d "/dT. The magnitude of d r t E/dT decreases with increasing chain length as in ethyl- and n-propyl-amine showing the opposite influences on water structure of the hydrophobic group and the NH, groups.With a sufficiently long hydrocarbon chain and/or bulky groups on the nitrogen atom as in n-butylamine or di- and tri-methylamine, the effect of the hydrophobic group seems to outweigh that due to NH, and the molecule behaves as a structure stabilizer. These findings are in agreement with the results of a similar comparative study of limiting partial molal compressibilities (d",) of amines and alcohols. It was shown by Kaulgud18 that the relative structure-strengthening ability of a solute could be described quantitatively by the expression d(Q)KS-Pi V,O)/dT, where Pi and VF are the compressibility and molar volume of the pure liquid. It was foundlg that this expression gave values which are considerably less for amines than for the corresponding alcohols. Concluding, it can be said that the volumetric behaviour of a monofunctional solute in water is also influenced by the nature of the polar group.In particular, comparison between the hydroxyl group (-OH) and the amine group (-NH,) shows that the latter acts as a structure breaker. Two of us (V. S. B.) and (A. S.) thank the University Grant Commission, New Delhi for a Faculty Improvement Programme Fellowship and a Junior Research Fellowship, respectively. We also thank the referees for their valuable suggestions and comments. M. V. Kaulgud and K. J. Patil, J. Phys. Chem., 1974, 78, 714. M. V. Kaulgud and K. J. Patil, J. Phys. Chem., 1976, 80, 138. S. Cabani, G. Conti and L. Lepori, J. Phys. Chem., 1974, 78, 1030. S. Cabani, G. Conti and L. Lepori, J. Phys. Chem., 1972, 76, 1338. G. Wada and S. Umeda, Bull. Chem. Soc. Jpn, 1962, 35, 646. G. Wada and S. Umeda, Bull. Chem. Soc. Jpn, 1962, 35, 1797. ' D. M. Alexander and D. J. T. Hills, Aust. J . Chem., 1969, 72, 347. J. Konicek and I . Wasdo, Acta Chem. Scand., 1971, 25, 1541. R. K. McMullan, G. A. Jeffery and T. H. Jordan, J . Chem. Phys., 1967, 47, 1229. lo F. Franks, Water, A Comprehensive Treatise (Plenum Press, New York, 1973), vol. 2, chap. 1, p. 35. l 1 M. V. Kaulgud, H. G. Dole and K. S. M. Rao, Indian J. Chem., Sect. A , 1978, 16, 955. l 2 F. J. Millero, Chem. Rec. 1971, 71, 147. l 3 M. V. Kaulgud, M. R. Awode and A. Shrivastava, Indian J . Chem., Sect. A , 1980, 19, 144. l4 D. D. Perrin, Dissociation Constants of Organic Bases in Aqueous Solution (Butterworths, Washington, l5 K. Raznjeric, Handbook of Thermodynamic Tables and Charts (Hemisphere, Washington, 1976). 1975).M. V. KAULGUD, V. S. BHAGDE AND A. SHRIVASTAVA D. M. Alexander, J . Chem. Eng. Data, 1959, 4, 252. '' J . L. Neal and D. A. I. Goring, J . Phys. Chem., 1970, 74, 658. l R M. V. Kaulgud, J . Chem. SOC., Faraday Trans. I , 1979, 75, 2246. M. V. Kaulgud, A. Shrivastava and M. R. Awode, Indian J. Pure Appl. Phys., 1980. 18, 864. (PAPER O/ 1203) 32 1

ISSN:0300-9599    

DOI:10.1039/F19827800313

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 323-339 Phospholipid Monolayers at Non-polar Oil/Water Interfaces Part 3.-Effect of Chain Length on Phase Transitions in Saturated Di-acyl Lecithins at the n-Heptane/Aqueous Sodium Chloride Interface BY JAMES MINGINS,* J. A. GORDON TAYLOR A N D BRIAN A. PETHICAT Unilever Research, Port Sunlight Laboratory, Quarry Road East, Bebington, Wirral, Merseyside L63 3JW AND CRAIG M. JACKSON AND BEATRICE Y. T. Y U E ' Department of Biological Chemistry, Washington University, Division of Biology and Biomedical Research, St. Louis, Missouri 63 1 10, U.S.A. Received 30th October, 1980 Isotherms of surface pressure (n) against area per molecule ( A ) are reported for a homologous series of pure synthetic saturated 1,2-di-acyl glycerophosphocholines (lecithins) (C14 to C2,) spread at n- heptane/aqueous electrolyte interfaces.The lecithins show second-order phase transitions which for a given temperature move to higher 7c and lower A as the chain length is decreased until for di-C,, lecithin no phase transition can be distinguished. The calculations of Clapeyron heats in Part 1 are extended to the new data using a single value of the reference area for the solid state for all temperatures and chain lengths. These heats vary with chain length and decrease linearly with temperature. The high heat capacity calculated in Part 1 for di-C,, lecithin at low temperatures is shown to be an artefact. Entropies of compression have been calculated from the free energies obtained by integrating the full isotherms between a reference area in the expanded region above any phase transition and another area in the solid-condensed region.These entropies vary linearly with chain length, having a slope given by the configurational entropy term R In 3 per methylene group, where R is the gas constant. The entropy change for each phospholipid is approximately 2(n- 1) Rln 3, where n is the chain length. This finding, taken together with a comparison of the heats of the monolayer phase transitions with the calorimetric heats of melting of phospholipid chains, and with related published data on the heats of fusion and the heats of solution of the alkanes, suggests that the chains of the phospholipids are fully flexible at low monolayer densities and very restricted on the condensed side of the phase transitions.Classic studies on spread monolayers of fatty acids and related species at the air/water (A/W) interface (reviewed by Adam,l Harkins2 and Gaines3) have shown complex phase behaviour with intermediate phases such as the liquid-expanded and liquid-condensed regions appearing on occasion between the two-dimensional analogues of the solid, liquid and gaseous states. The form of the transitions depends markedly on the chain length of the monolayer molecules, with increase of chain length condensing the surface pressure (n) against area per molecule ( A ) isotherms and often precluding the formation of some monolayer states. Comparable behaviour was shown much later with phospholipid monolayers at the A/W interface. The thermo- dynamic status of these transitions has been under scrutiny recently; some are certainly metastable, with the isotherms being dependent on the history of the mono layer^.^^ Because of diminished cohesion among the chains of the monolayer molecules at t Present address: Clarkson College, Potsdam, New York 13676, U.S.A.3233 24 MONOLAYER PHASE TRANSITIONS an oil/water (O/W) interface the n against A isotherms are more expanded than at the A/W i n t e r f a ~ e . ~ . ~ In addition, the phase properties are different, with no evidence to date of the liquid-expanded-liquid-condensed transitions occurring. Second-order phase transitions have now been reported for both single-component and mixed monolayers spread at the O/W interfa~e.~? 8~ Although these transitions have been variously referred to by us as first order, near first order or degenerate first order, the isotherms show a discontinuity in slope and not in density and are treated here as second order in character.No systematic study of their chain-length dependence has been described hitherto. In this paper we examine the temperature dependence of the phase transitions in a homologous series of saturated lecithin monolayers and use the chain length dependence of the entropies of the transition to gain insight into the nature of the monolayer interactions. EXPERIMENTAL The results to be described are again from two collaborative research programmes, one at Port Sunlight (P/S) and a subsequent one at St. Louis (S/L). The purification of materials making up the n-heptane/aqueous electrolyte solution interfaces and the techniques of spreading on the Brooks trough* and measuring n with a recording electronic microbalance were described in earlier papers.l53 l7 Additional information for the experiments discussed here is as follows.Solutions of low pH were prepared from a constant boiling fraction of AnalaR hydrochloric acid distilled twice from Pyrex (P/S) or from untreated AnalaR grade (S/L). A homologous series of saturated 1,2-di-acyl glycerophosphocholines (lecithins) in steps of two methylene groups over the range C,, to C,, was obtained, using highly puri- fied synthetic samples from various sources. A general formula for the lecithins would be R,C0,CH,CH(0,CR,)CH20P0,(CH,),N(CH~)3 and in the present work the alkyl groups R, and R, are identical.As in our previous papers we refer to C, lecithin, where n is the length of the alkyl chain, i.e. n is the number of carbon atoms in R, plus 1. These lecithins were either synthesised from glycerol phosphoryl choline (GPC)l* and fatty acid anhydrideslg by the method of Cubero-Robles and van der Berg2* or from the cadmium chloride adduct of GPC and fatty acyl chloride2' by the method of Baer and Buchnea.22 The final products were purified by chromatography on silicic or alumina2, and for the S/L experiments the samples were - + TABLE 1 .-FATTY ACID ANALYSIS OF SYNTHETIC LECITHINS composition (wt 7:) lecithin source 12:O 14:O 16:O 18:O 18:l 20:O 22:O miscellaneous Utrec ht Welwyn Welwyn Welwyn S/L S/L S/L S/L S/L - -- - - - 100 - - 99.9 0.1 - - trace - 100 - 99.3 0.4 - 0.2 -- trace 99.9 < 0.1 - 0.1 - 0.1 98.3 - - - - - - - - - - - - - - - 100 - - - - - - - - - 1.6 ~- - - - ._ - 100 - - __ 0.2 - 99.8 - - - - Figures are rounded off to 0.1. Utrecht: L.L. M. van Deenen and colleagues, Organisch Chemisch Laboratorium, Utrecht. Welwyn : D. Chapman and colleagues, Unilever Research Laboratory, The Frythe, Welwyn. S/L : Mary E. Neubert, Department of Biological Chemistry, Washington University School of Medicine, St. Louis.MINGINS, TAYLOR, PETHICA, JACKSON AND YUE 325 subjected to preparative t.1.c. immediately prior to use.25 Results on the chain length distribution in the lecithins obtained from gas-liquid chromatography on the fatty acid methyl esters26 are summarised in table 1. The samples were stored in the cold over desiccant.Spreading solutions were made up in the concentration range 0.1-0.5 g dm-3 in 9/1 v/v heptane/ethanol (P/S) and 0.05-0.33 g dmP3 in various mixtures of heptane and ethanol or as .fresh solutions in pure ethanol (S/L). The equilibrium status of the measurements was checked by measuring the equilibrium spreading pressure (n,) of one of the lecithins, namely C,, lecithin. This was accomplished by placing dry crystals of the lecithin on the air/heptane surface where they were immediately wetted by the heptane and sank to the O/W interface. Here they spread spontaneously to give high surface pressures which were monitored by a dipping carbon plate attached to an electronic microbalance. RESULTS The 71-A isotherms for the five lecithins of chain lengths CI4, CI6, CIS, C,, and C,, are shown in fig.1-5, respectively. The results merge smoothly with the high A (P/S) data reported in Part 2. I I a 5 2 2.5 A /nm2 molecule-' FIG. I .-Surface pressure against area isotherms for 1,2-dimyristoyl lecithin at the n-heptane/aqueous NaCl solution interface. Inset : specimen isotherms depicting the phase transition of low temperatures. Numbers on curves give temperature in OC. In some experiments a lack of overlap of portions of the 71-A isotherms at high inonolayer density signifies a loss of monolayer on spreading so that recourse had again to be made to the scaling procedure justified in an earlier paper.I7 In view of the importance of this procedure in enabling access to the high-density isotherms it is worth briefly reiterating the routine.An accurate n-A isotherm is built up from high A by spreading successive amounts of lecithin on top of a resident stabilised monolayer326 MONOLAYER PHASE TRANSITIONS 0 5 I I 5 2 L A /nm2 molecule-' 5 FIG. 2.-Surface pressure against area isotherms for 1,2-dipalmitoyl lecithin at the n-heptane/aqueous NaCl solution interface. Numbers on curves give temperature in "C. and noting superposition of n at common values of A. As soon as stable 7c for a new spread fall below the established isotherm all the (apparent) values of A for the new monolayer are multiplied by a single factor which gives superposition of the isotherms in the region where the A overlap. In this way progress is made along the isotherm to lower A and the process is repeated for further spreads.For some (S/L) data where results at high A were not available, this meant scaling to high A data on C,, lecithin from (P/S), a manoeuvre justified by the chain-length independence of the n-A isotherms shown in Part 2 and confirmed here with the remaining (P/S) data. When temperatures for the isotherms from the two laboratories were not congruent, fitting the (S/L) results to high A data was accomplished by interpolation on sets of TL against temperature plots from (P/S) C,, lecithin data and scaling the deficient runs to the thus-obtained isotherms. No phase transition is found in the temperature range down to near 0 O C for C,, lecithin whereas a degenerate phase transition appears at high n at the lowest temperatures (10 O C and below) for C , , lecithin.Reducing the pH to 2.1 has no effect on the isotherm for C,, lecithin at 20 O C , confirming the findings in Part 1 on the C,, lecithin monolayers. As the chain length is increased to C,, the second-order phase transition is more clearly distinguished and persists to higher tempera- tures.? The phase transitions for C,, lecithin were reported in Part 1 and the t The (P/S) data on C , , lecithin were erroneously quoted in an earlier p a p e P where it was stated that there is no phase transition. The statement should read that there is no first-order phase transition.MINGINS, TAYLOR, PETHICA, JACKSON AND YUE 0 5 D 1.5 2.0 2 A /nm2 molecule-' FIG. 3.-Surface pressure against area isotherms for 1,2-distearoyl lecithin at the n-heptane/aqueous NaCl solution interface.Numbers on curves give temperature in O C . 40 327 1 1 1 1 I I I I 1 I I 0 5 09 I *3 I *7 2. I 2.5 r L A /nm2 molecule-' 3 FIG. 4.-Surface pressure against area isotherms for 1,2-di-arachidoyl lecithin at the n-heptane/aqueous NaCl solution interface. Numbers on curves give temperature in O C .328 MONOLAYER PHASE TRANSITIONS 0 5 I 1.5 2 2 A/nm* molecule-' FIG. 5.--Surface pressure against area isotherms for 1,2-dibehenoyl lecithin at the n-heptane/aqueous NaCl solution interface. Numbers on curves give temperature in OC. data are shown here again for completeness, together with results in the range 1,3 < A/nrn2 molecule-l < 2.0 to abut the results at the lower limit of A reported in Part 2. Isotherms on CIS, C,, and C,, lecithins continue the trend seen above with the lower chain length lecithins, namely that increase of chain length decreases IT at the transition (q) and increases the area at which the transition begins (At).In fig. 6 a composite plot of all the available isotherms at 20 O C shows that the phase transitions appear as breaks from a master curve for the lecithins at this temperature. Similar master curves can be drawn for the other temperatures. This behaviour follows from the chain-length independence of the high A isotherms shown in fig. 2 of Part 2. In fact, the small scatter in the isotherms on that diagram can be substantially reduced by application of the criteria for quantitative spreading described above. Plots of n, and A , against temperature for all the lecithins above C,, are given in fig.7 and 8,5c 4c 3c - I E z E -.. k= x IC C MINGINS, TAYLOR, PETHICA, JACKSON AND YUE \ cl4 I I 0.5 I .o 1.5 2.0 2. A /nm2 molecule-' 329 Fie;. 6.-Composite plot of surface pressure against area isotherms for various-chain-length lecithins typically at 20 "C. and the dependence on chain length is shown in fig. 9. The consistency of the data supports the validity of the methods used to judge monolayer spreading. With the absence of monolayer leakage, residual spreading solvents and drifts in temperature, the isotherms at A > A , are fully reversible except at high n where the lower-chain-length lecithins show a slight desorption. For A < A , the isotherms are again reversible provided monolayer collapse, plate-displacement by viscous mono- layers and desorption are also avoided.On traversing the isotherms through the transition point, pressures are often obtained which depend on the rate of compres- sion/expansion of the monolayer. Pressures at A just less than .4, are especially sensitive and rapid compression through the transition generates pressures which fall to the steady values given in the figures. Rapid compression to values of A much lower than A , give n readings only marginally higher than the steady values except at the lower chain lengths when desorption occurs. At lower A , in the solid-condensed region, the values of 71 can be held constant with time but they are often subject to the many time-dependent effects discussed in ref. (1 7), particularly at low temperatures and for long chains.The rapid expansion of a monolayer from A < A , to A > A , gives low pressures w3ich on standing increase slowly to the values on the given isotherms. Values of n, are proportional to temperature and range from ca. 25 mN m-l at 5 O C to ca. 38 mN rn-l at 30 O C .330 MONOLAYER PHASE TRANSITIONS I tem pera t ure/O C FIG. 7.-Temperature dependence of the surface pressure at the onset of the transition for various- chain-length lecithins at the n-heptane/aqueous NaCl solution interface. DISCUSSION THE NATURE OF THE PHASE TRANSITIONS For the longer-chain lecithins at the lower experimental temperatures, the phase transitions at the O/W interface appear at first sight to be first order. However, the pressure definitely rises as the area is reduced in the transition region and the high- density end of the transition is always in the form of a smoothly rising curve.As the chain length decreases and the temperature rises, the departure from the first-order behaviour becomes very pronounced and the phase transitions are perhaps best described as second order. and MarEelJa29 have argued that departure from first-order behaviour in phospholipid monolayers at the A/W interface is due to impurities. Several arguments against this conclusion have been made p r e v i o ~ s l y . ~ ~ We now add further arguments based on the analytical data of table 1. All the phospholipid samples show second-order behaviour and we could not distinguish differences in the isotherms of the various samples of a given homologue, either at high areas or in the transition regions.The purest samples (loo%, no detectable impurities) showed definite second-order behaviour at higher temperatures. Nor can it be maintained that impurities from the bulk aqueous or heptane phases cause the departure from first-order behaviour, since penetration of such impurities should be less at the higher pressures and temperatures for which the departure is most obvious. We therefore conclude that these transitions are not first order. At the onset of the transition, it is probable that aggregated clusters are formed, the size of these being larger at lower temperatures and for longer chains.31MINGINS, TAYLOR, PETHICA, JACKSON AND YUE 33 1 temperature/"C FIG 8.-Temperature dependence of the area at the onset of the transition for various-chain-length lecithins at the n-heptane/aqueous NaCl solution interface. The areas for C,, lecithin are approximate because of monolayer desorption at high 7 ~ .Throughout the whole temperature range studied the values of n, for C,, lecithin are substantially higher than nt at the corresponding temperatures and they exceed all the pressures depicted in fig. 3, except for the extreme values in the solid-condensed region. The isotherms depicted for C,, lecithin can therefore be considered equilibrium ones and we would expect the same to be true for the other lecithins. USE OF THE CLAPEYRON EQUATION In Part 1 of this series,15 the use of the Clapeyron equation for calculating the heats of the phase transition for the di-C,, lecithin monolayers was discussed at some length.The Clapeyron equation is strictly applicable to first-order transitions and the heats obtained were therefore approximate only. However, since the pressure (q) at the onset of the transition is a reproducible and definite feature of each isotherm, there is value in applying the Clapeyron equation for comparative purposes. The collected data on nt and At (the area at the onset of the transition) for the homologous series are presented in fig. 7-9. The results are self-consistent and follow smooth trends. The Clapeyron heats have been calculated from332 MONOLAYER PHASE TRANSITIONS taking A , as 0.47 nm2 molecule-' throughout because the solid-condensed isotherms are not well-defined for the lower homologues.This value is an experimentally accessible molecular area for the condensed regions of all the isotherms except for the di-C,, compound at higher T. The results are shown in fig. 10. The Clapeyron heats vary significantly with temperature and increasingly as the chain length diminishes and as the departure from first-order behaviour becomes more pronounced. The variation of AH with T for the di-C,, lecithin shown in fig. 10 is very much less at T c 20 OC than is shown in fig. 5 of Part 1 [ref. (1 5)]. The difference is entirely due to re-estimating the slopes of the nt-T curve at the graphical extremum at the lower temperatures as a result of adding extra experimental determinations of nt at 0.3 and 3.6OC. These recalculations for the di-C,, lecithin show that, contrary to the conclusion in Part 1, there is not a large heat capacity for the transition at low temperature.Indeed, the variation of AH with 7' for the shorter-chain compounds is now seen as reflecting the departure from first-order behaviour, assumed in the use of the Clapeyron equation. Since this departure from first-order behaviour is so pronounced for the shorter-chain phospholipids, it is not very useful to comment on the chain-length dependence of the Clapeyron heats, except to say that they follow the trends considered more closely in the next section, in which the entropy changes in the transition region are obtained by a rigorous thermodynamic argument not depending on treating the transitions as first order.25 2c I5 I - 0 E - m 24 5 ' IC E - 5 L I 0 10 20 30 L temperaturel'c 30 '5 - I - 0 E 25 FIG. 10.-Clapeyron heat of the phase transition plotted against temperature for various-chain-length lecithins at the n-heptanejaqueous NaCl solution interface. Molecular area of solid condensed monolayer for the Clapeyron equation taken as 0.47 nm* molecule-'.Values for C , , lecithin are approximate only because of uncertainties in A,. THERMODYNAMIC FUNCTIONS FOR MONOLAYER COMPRESSION ACROSS THE TRANSITION REGION In Part 2, it was shown that for areas > 2.0 nm2 molecule-l the Helmholtz free energies and the entropies of compression of the phospholipid monolayers are independent of chain length. The Helmholtz free energies of compression are obtained from the area under an isotherm between two chosen areas per molecule, and the entropies from the temperature coefficients of this area.For the data given here, the free energies and entropies of compression are also independent of chain length for compressions down to the area for the onset of the phase transition for a particular homologue. It is thus useful to choose a conveniently high area, outside the transition region, and to integrate the isotherms to a chosen area on the condensed side of the transition to obtain the Helmholtz free energy for the compression process including the transition. From the effect of temperature, the entropy of compression, including the phase transition, is obtained directly. The reference areas chosen were 2.0 and 0.47 nm2 molecule-' and the changes in Helmholtz free energy, entropy and energy for this range are denoted by AFc, ASc and AE,, respectively. The higher area lies outside the transition region for all the isotherms given, except for the di-C,, at 5 OC,334 MONOLAYER PHASE T R A N S I T I O N S tempera ture/O C FIG.11 .-Change in Helmholtz free energy on compressing the monolayer from 2 to 0.47 nm2 molecule-' plotted against temperature for various-chain-length lecithins at the n-heptane/aqueous NaCl solution interface. and as stated earlier the lower area is experimentally available for all homologues except the di-C,,. The transition region for the di-C,, lecithin was very degenerate and solution occurs as 0.47 nm2 molecule-l is approached, even at the lowest temperatures. (The values of n,, however, could be reasonably estimated and are included in fig.7.) Values of AF, are shown in fig. 11 as a function of temperature for the series of lecithins. The variation of AF, with Tis linear and the derived AS, for the compression is correspondingly independent of T within experimental error. The ASc values for the homologous series are given in fig. 12 and show a linear variation with chain length. The phase transitions for the di-C,, phospholipid are observed only below 10 OC. The data of fig. 1 are for reversible, stable monolayers, but at the higher pressures of the transitions there is some tendency for the C,, lecithin monolayer to dissolve. The value shown for di-C,, in fig. 12 is instead calculated by extrapolating ASc against log A plots for C,, lecithin (above A,) to 0.47 nm2 molecule-l and corresponds to the value for lower homologues, which do not show transitions above 0 "C.Given that above A, the entropy of compression is independent of chain length, the entropy variations in fig. 12 refer to the phase transition itself. To complete the description of monolayer compression through the transition region, the AEc values have been obtained from AF, and AS, (fig. 13). These values are, of course, independent of T. The linear variation with chain length of the entropy of compression across the phase transition region (fig. 12) is rather close to R In 3 per methylene group. This corresponds to the loss of configurational entropy to be expected if the phase transition were regarded as the restriction of the paraffin chains to one close-packed configuration on a tetrahedral lattice of methylene groups in the ' solid ' monolayer, as distinct from the multiple configurations available to the chains in the molten or solution states.This interpretation is obviously qualitative, if only because the packing335 FIG. 12.-Entropy of monolayer compression from 2.0 to 0.47 nm2 molecule-' plotted against chain length for lecithins at the n-heptane/aqueous NaCl solution interface. Flagged point for C,, lecithin obtained from extrapolation on ASc against log,, A graph for C,, lecithin because of monolayer desorption at the high n near 0.47 nm' molecule-' on C,, lecithin. Slope of line is Rln 3. of the phospholipid molecules at the selected reference area in the condensed monolayer phase (0.47 nm2 molecule-l) is much looser than for close-packed chains (0.38 nm2 per pair of chains).The similarity between the monolayer transition, the melting of paraffins and the dissolution of solid into liquid alkanes is strengthened by the fact that ASc extrapolates to zero near fo zero chain length. The extrapolation from C,, is long, but the total entropy of the phase transition for a given phospholipid homologue with n carbon atoms in the acyl groups is given adequately by 2(n - 1) R In 3. This suggests that all the 2(n - 1) methylene and methyl groups in each molecule take part in the monolayer transition, in a manner corresponding to the crystallisation of the chains from a melt, or more appropriately from a two-dimensional solution in heptane.If this interpretation is correct, the contribution of the phosphatidyl choline head group to AS, is small. This implies that the configurations of the head group are restricted, not necessarily in the same conformation, in both dilute and dense monolayers. Several ways in which the separated or close-packed zwitterions could be so restricted are conceivable on purely electrostatic grounds. The influence of the head group on the phase transition is revealed more directly by extrapolation of the linear plot of AE, with chain length (fig. 13). This shows that AEc would be close to zero for the di-C,, molecule and 45 J mol-l for a hypothetical di-C, molecule. The336 MONOLAYER PHASE TRANSITIONS chain length FIG. 13.-Energy of monolayer compression from 2.0 to 0.47 nm2 molecule-' plotted against chain length for lecithins at the n-heptane/aqueous NaCl solution interface.first carbon of the chain is a carbonyl group and could be regarded as part of the head group in the aqueous phase, as distinct from the methylene chain proper. The contribution of the head group to A& will be the subject of further c~nsideration.~~ Direct comparison of the ASc with the experimental data on the fusion of solid alkanes and on their dissolution in liquid alkanes is instructive. The even-numbered alkanes show no solid-solid phase transitions below C2,, and the entropies of fusion up to c18 are to be compared with the sum of the entropies of the solid-solid transition and of fusion for C,, and above. On this basis, the entropies between C,, and c36 vary by close to R In 3 per methylene group.In the range C,, to C,,, corresponding to our experiments, the variation is 1.14R In 3."9 34 Masden and Boi~telle~~ measured the solubilities of the c28, C,, and C,, alkanes in heptane and pentane as a function of temperature. The derived entropies of solution show a variation with chain length of 0.73Rln 3 for heptane and 0.83Rln 3 for pentane as solvents, with a standard deviation of 10- 15 %. These entropies refer to the respective saturated solutions at concentrations which vary twenty-fold from c28 to c36. Making a correction to equal mole fraction for the entropies of solution gives a variation with chain length of R In 3 within experimental error, in agreement with the monolayer transition case.Another set of useful comparisons can be made with the calorimetric heats of chain melting observed in phospholipids. These heats have been determined in St. Louis for the same phospholipids used there in the monolayer experiments, with the phospho- lipids dispersed in the presence of water and heptane. The results are to be described in detail el~ewhere.~~ The melting points observed in the calorimeter were, for each homologue from C,, to C2,), comparable with the two-dimensional critical points, as approximately estimated from the n-A isotherms. The chain melting in the bilayers is accompanied by changes in packing as the melting occurs. These packing changes are much smaller than the changes observed, for example, in the surface phase transitions for the longer-chain phospholipids at low temperature.The packing changes for the surface phase transitions with the di-C,, and di-C,, lecithins will,MINGINS, TAYLOR, PETHICA, JACKSON AND YUE 337 however, be fairly close to those occurring in the bilayers and a reasonable correspondence may therefore be expected between the Clapeyron heats and the calorimetric heats at these lower chain lengths. This is in fact found. The calorimetric heats are 24 and 33 kJ mol-l for the di-C,, and di-C,, lecithins, in the range of the results shown in fig. 10. For the higher homologues, the calorimetric heats fall increasingly below the Clapeyron heats. The general similarity of the phase changes in bilayers and monolayers of the phospholipids is reinforced by these comparisons. It does not seem necessary, in accounting for the chain-length dependence of AH and ASc, to invoke significant contributions from effects arising from re-ordering of the heptane molecules.This conclusion is supported by a substantial body of monolayer data (obtained at St. Louis, to be published) on the lecithin monolayers at the interfaces of aqueous salt solutions with iso-octane and cyclohexane. There are various points of difference with the data using heptane as the oil phase, but the values of ASc> and its chain-length dependence with these other oils are very similar to those given here for the heptane interface. RECONSIDERATION OF THE HIGH-AREA REGION The lecithin isotherms can be described naturally in two regimes where the molecular interactions differ.The first and at first sight simpler regime covers areas above those of the second-order phase transitions, where the intermolecular forces are overall repulsive and there are no inflections or discontinuities in the isotherms. The second regime, discussed in the above section, encompasses the second-order phase transitions down to the solid-condensed region of the isotherms, where co-operative cohesive interactions are at work. The first regime was dealt with at length in Part 2, partly in the expectation that the interpretation of the P A isotherms in the dilute region would facilitate understanding of the phase-transition region. In the event some reconsideration of certain of the ideas put forward in Part 2 is now required as a result of the analysis of the complete set of data on the phase transitions given above.It is useful to recall several features of the high-area data at this stage. The lecithin isotherms in the high-area region are not affected by substantial changes in pH and salt concentration. This finding holds also down to the condensed region, thus confirming that long-range forces of the ionic double-layer type are not involved in the monolayer behaviour of these zwitterionic molecules. In the high-area region, the isotherms are the same for all chain lengths. The data presented here show that this is true at all areas down to the onset of the phase transition for each homologue and temperature (fig. 6). From nA-n plots in the high-surface-area region we concluded that dimers and other aggregates are not formed.Again, this conclusion can be accepted to close up to the phase transitions. The experimental finding in the high-area region that for all chain lengths the apparent co-areas ( A , ) ) are much greater than the close-packed area (ca. 0.38 nm2 molecule-l) and that these A , values increase with temperature were the most striking results in Part 2 and were discussed in terms of repulsion arising from the zwitterion dipoles and from steric repulsions between a portion of the ends of the aliphatic chains. It was argued that the dipole contribution was insufficient to explain all the repulsion and that if A,, is ascribed to chain-chain interactions then at a given temperature a definite portion of the aliphatic chains, independent of chain length, must be involved to give A , invariant with chain length.This portion would, it was argued, increase with temperature to cover the A , variations. The data on the phase transitions contradict this account of the A , values because the transition entropy varies systematically with chain length. This objection to the explanation of the A,, values given in Part 2 is also supported by published data showing small heats and entropies of mixing of paraffins of differing chain lengths.37 40 12 F A R 1338 MONOLAYER PHASE TRANSITIONS We still conclude that dipole repulsions alone as calculated from the lattice energies given by ToppingA1 will not account for the high A, values and their dependence on temperature, but a good explanation of the high-area phospholipid isotherm is lacking.Better estimates of the dipole forces are in hand32 which with analysis of unpublished data on surface potentialsJ2 will enable us to resolve the above problem and to account in part for the contribution of the head groups of the phospholipids to the AEc of the transitions. CONCLUSIONS Provided the chain length is high enough, neutral glycerophospholipids spread to give stable monolayers at the n-heptanelwater interface. With careful attention to experimental technique and sample purity, reproducible second-order phase transitions can be established. With an increase of chain length or a decrease in temperature the onset of these transitions occurs at lower surface densities and pressures. The co-operativity of the chain interactions is upset by chain branching or unsaturation as at the A/W interface and the phase transitions disappear.The phase transitions depend on the nature of the head group but for the zwitterionic phospholipids there is no long-range double-layer interaction and even when the zwitterion is charged positively there is only a marginal effect. As with other lipid monolayers the phospholipids are more expanded at the O/W interface here than at the A/W interface because of reduced attractive dispersion interactions between the lipid chains in the n-heptane solvent. As a consequence there is no nmer formation at low and moderate surface densities. Instead the n-A data show that the phase transitions arise from a background of strong repulsion between the monomers in contrast to the A/W interface where co-operative interactions occur within an already attractive milieu.One of the most intriguing features of the data is that the repulsion is independent of chain length whereas the chain+hain attractions at the phase transition increase with chain length. No satisfactory explanation of the high repulsion is available and its large positive temperature coefficient runs counter to the expected behaviour of dipole and water-structuring forces. Calculated entropies of compression over areas up to the phase transition are approximately twice those for single-chain molecules at the O/W interface which in turn are approximately twice those for similar species in expanded monolayers at the A/W interface. There is therefore substantial freedom of rotation for both chains of the phospholipid molecule in the heptane.Restrictions on the chains at the A/W interface are in keeping with the model of rigidified chains lying on the water surface which was recently used to explain the chain-dependent two-dimensional second virial coefficient^^^ of A/W monolayers. Entropies of compression through the phase transition vary by R In 3 per methylene group on the phospholipid and extrapolate to zero entropy at the acyl carbons. The phase transitions can therefore be taken as a condensation of freely kinking chains into aggregates with virtually complete loss of entropy as the phospholipids de-mix from the heptane, The amount of entropy release per methylene group is similar to that calculated from bulk measurements on fusion of solid alkanes or solution of n-alkanes in n-alkane solvents.This work would not have been possible without a range of highly purified synthetic lecithins and we are very grateful to L. L. M. van Deenen, D. Chapman and their colleagues for furnishing us with samples and to M. E. Neubert for synthesising the pure S/L lecithins.MINGINS, TAYLOR, PETHICA, JACKSON A N D YUE 339 ' N. K. Adam, The Physics and Chemistry of Surfaces (Oxford University Press, 3rd edn, 1941). j G. L. Gaines, Insoluble Monolayers at Liquid-Gas Interfaces (Interscience, New York, 1966). W. D. Harkins, The Physical Chemistry of Surface Films (Reinhold, New York, 1952). N. L. Gershfeld, Annu. Rer. Phys. Chem., 1976, 27, 349. H. W. Horn and N. L. Gershfeld, Biophys. J., 1977, 18, 301. J. T. Davies, Proc.R. Soc. London, Ser. A , 1951, 208, 224. J. N. Phillips and E. K. Rideal, Proc. R. SOC. London, Ser. A , 1955, 232, 149. J. H. Brooks and B. A. Pethica, Trans. Faraday Soc., 1964, 60, 208. J. Mingins, N. F. Owens, J. A. G. Taylor, J. H. Brooks and B. A. Pethica, in Monolavers, ed. E. D. Goddard, Adti. Chem. Ser., 1975, 144, 14. l o J. H. Brooks and B. A. Pethica, Proc. 4th Int. Congr. Surface Actice Substances, Brussels, 1964, vol. 2, p. 191. I ' J. A. G. Taylor, J. Mingins, B. A. Pethica, B. Y. T. Tan and C. M. Jackson, Biochim. Biophys. Acta, 1973, 323, 157. l 2 C. M. Jackson and B. Y. T. Yue, in Monoiayers, ed. E. D. Goddard, Adi). Chem. Ser., 1975,144,202. J. A. G. Taylor, J. Mingins and B. A. Pethica, J . Colloid Interface Sci., 1976, 55, 2.l 4 E. Llerenas and J. Mingins, Biochim. Biophys. Acta, 1976, 419, 381. l 3 B. Y . Tan, C. M. Jackson, J. A. G . Taylor, J. Mingins and B. A. Pethica, J. Chem. Soc., Faraday Trans. I , 1976, 72, 2685. J. Mingins, E. Llerenas and B. A. Pethica, Ions in Macromolecular and Biological Interfaces, Colston Papers no. 29, ed. D. H. Everett and B. Vincent (Scientechnica, Bristol, 1978). l 7 J. A. G. Taylor and J. Mingins, J. Chem. SOC., Faraday Trans. 1, 1975, 71, 1161. In H. Brockerhoff and M. Yurkowski, Can. J . Biochem., 1965, 43, 1777. lY Z. Selinger, J. Lipid Res., 1966, 7, 179. 2n E. Cubero-Robles and D. van der Berg, Biochim. Biophys. Acta, 1969, 187, 520. 21 H. E. Fierz-David and W. Kuster, Heh. Chim. Acta, 1939, 22, 82. 22 E. Baer and D. Buchnea, Can. J . Biochem. Physiol., 1959, 37, 953. 23 C. C. Sweeley, Methods in Enzymology, Lipids, ed. J. M. Lowenstein (Academic Press, New York, 24 D. N. Rhodes and C. H. Lea, Biochem. J., 1957, 65, 526. 25 V. P. Skipski and M. Barclay, Methods in Enzymology, Lipids, ed. J. M. Lowenstein (Academic Press, 26 R. G. Ackman, Methods in Enzymology, Lipids, ed. J. M. Lowenstein (Academic Press, New York, 27 H. L. Scott Jr, J. Chem. Phys., 1975, 62, 1347. Zn J. Mingins and J. A. G. Taylor, Proc. R. SOC. Med., 1973, 66, 383. 29 S. MarEelja, Biochim. Biophys. Acta, 1974, 367, 165. 3o G. M. Bell, J. Mingins and J. A. G. Taylor, J. Chem. SOC., Faraday Trans. 2, 1978, 74, 223. 31 I. Langmuir, J . Chem. Phys., 1953, 1, 756. 32 M. L. Glasser, J. Mingins and B. A. Pethica, unpublished results. 33 A. A. Schaerer, C. J. Busso, A. E. Smith and L. B. Skinner, J. Am. Chem. Soc., 1955, 77, 2017. 34 J. F. Messerley, G. B. Guthrie, S. S. Todd and H. L. Finke, J. Chem. Eng. Data, 1967, 12, 338. 35 H. E. L. Masden and R. Boistelle, J. Chem. SOC., Faraday Trans. 1, 1976, 72, 1078. 36 C. M. Jackson and B. Y. T. Yue, unpublished results. 37 M. L. McGlashan and A. G. Williamson, Trans. Faraday SOC., 1961, 57, 588. 38 M. L. McGlashan and K. W. Morcom, Trans. Faraday SOC., 1961, 57, 581. 39 H. F. Stoeckli, J. G . Fernandez-Garcia and L. G. Boissonas, Trans. Faraday Soc., 1966, 62, 3044. 3o J. S. Rowlinson, Liquids and Liquid Mixtures (Butterworths, London, 1959). 41 J. Topping, Proc. R. SOC. London, Ser. A , 1927, 114, 67. 4 2 G. M. Bell, J. Mingins and J. A. G. Taylor, unpublished results. 43 B. A. Pethica, M. L. Glasser and J. Mingins, J. Colloid Interface Sci., 1981, 81, 41. 1969). vol. 14, p. 255. New York, 1969), vol. 14, p. 530. 1969), vol. 14, p. 329. (PAPER O / 15 15)

ISSN:0300-9599    

DOI:10.1039/F19827800323

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 341-346 Identification and Atomization Energies of Gaseous Lac, and Lac, by High-temperature Mass Spectrometry BY KARL A. GINGERICH,* REZA HAQUE AND MARIO PELINO Department of Chemistry, Texas A & M University, College Station, Texas 77843, U.S.A. Received 10th November, 1980 The atomization energies, Di, and standard heats of formation, AHf, 298,15, in kJ mol-l, of the new molecules Lac, and Lac, have been obtained from Knudsen effusion mass-spectrometric measurements as 4320 60 and 1095 f 50 for LaC,(g) and 4992 70 and 1 140 k 60 for LaC,(g). Since the first Knudsen effusion mass-spectrometric study of gaseous metal carbides by Chupka et a1.l a large number of stable gaseous carbides, mainly of transition metals including the rare earths and actinides, have been found and their bond energies reviewed and discussed in terms of empirical relation^.^.^ The first observation of gaseous metal with more than four carbon atoms, namely up to six for cerium, is quite recent5 and has since been confirmed for other metals such as uranium,6 yttrium7 and thorium.8 One of the objectives of the present investigation was the search and measurement of even higher carbides, e.g.with seven or eight carbon atoms. Lanthanum appeared to be a favourable candidate. It is one of the few metals for which the partial pressure of the dicarbide, the most abundant carbide molecule over the lanthanum-carbide- graphite system, is comparable with or even higher than the partial pressure of the corresponding metal.'T EXPERIMENTAL The system chosen for study was lanthanum-iridium-graphite.The presence of iridium permitted us to carry the Knudsen effusion studies to temperatures above 2800 K without the vapour pressure exceeding molecular flow conditions, since iridium had the effect of significantly reducing the lanthanum activity in the sample through formation of a very stable intermetallic compound. In addition, the high temperatures were favourable for the observation of high carbon-containing species. The graphite effusion cell, containing the sample, was enveloped by a tantalum Knudsen cell in order to avoid an excessive loss of carbon at the very high temperatures of measurement. The measurements reported here were taken in series two, subsequent to the study of the LaIr where additional specific details of the experimental conditions can be found.The instrument and the experimental procedure used have been described previously. l1 The parent ions of lanthanum and lanthanum carbide species were identified by their mass-to-charge ratio, shutter effect and, when possible, ionization efficiency and appearance potential. The new metal carbides, Lac, and Lac,, were observed between 2668-2840 and 2827-2840 K, respectively. Their ion currents were too small for a determination of their appearance potential. However, there were no higher carbides present which could have contributed measureably to the measured ion currents of LaCf and Lac:. The estimated contribution of Lac, by fragmentation to the measured value for LaCf is < 4%, in analogy to the values measured for fragmentation by 19.5 V electrons of UC, to UC+ and C, or C, into C+.'j In table 1 the ion currents of the species used are listed. 34 1342 ATOMIZATION ENERGIES OF Lac, AND Lac, RESULTS AND DISCUSSION The reactions used for evaluating the molecules Lac, and Lac, were La(@ + nC(graphite) = LaC,(g) LaC,(g) + Lac&) = La(& + Lac,(& 2LaC,(g) = La(& + Lac&).(1) (2) (3) Reactions (2) and (3) were included because during the set of measurements at 2840 K the ion currents started to decrease with time. In a subsequent set of measurements this decrease accelerated for La(& as well as for the lanthanum carbide species, indicating that the sample and possibly the graphite liner had almost completely evaporated.In addition, at these high temperatures the graphite liner had certainly reacted with the tantalum of the outer tantalum cell to form sub-stoichiometric tantalum carbide.', Since the elemental carbon species were not monitored in an effort to make more rapid measurements of the lanthanum-containing species, no direct evaluation of the carbon activity in the condensed phase could be performed for the moment of measurement of Lac: and Lac: at 2840 K, and reaction (1) was accordingly not used for the set measured at this temperature. The intensities of La+, Lac,+ and Lac,+ corresponding to the moment of measurement of Lac: and Lac: and shown in table 1 could, however, be evaluated by interpolation from measurements of these species prior to and following that of Lac: and Lac:.The limited data available for Lac, and Lac, permitted only third-law evaluations. The necessary thermal functions were taken from Hultgren et al.13 for La(g) and C(g), from Gingerich et al.I4 for Lac, (assumed structure La-C-C-C, linear) and Lac, (assumed to be C,-La-C,, linear). The thermal functions for LaC,(g) were essentially identical with those given by Kohl and st earn^.^ Those for Lac, and Lac, were calculated by standard thermodynamic expressions from estimated molecular parameters assuming linear C,-La-C, and C,-La-C, structures, respectively. Here the bond distance and force constants were taken to be the same as for the C2 molecule, 1.31 A and 9.25 mdyn k l , re~pective1y.l~ The bending force constants Kd (mdyn A rad-l) were estimated to be one tenth of the average of the two adjacent stretching force constants.The vibrational frequencies were calculated by the F-G matrix method of Wilson.16 The electronic contribution was approximated by that of the La2+ ion.', The calculated Gibbs-energy functions, (GO, - g ) / T , and heat- content functions, IIoT-G, for Lac&) and Lac&) in the temperature range 2400-3200 K are listed in table 2. In evaluating the equilibrium constants for reactions (1)-(3), which are all pressure- independent, the simplifying assumption was made that the relative ionization cross-sections and multiplier gains cancel; i.e. the experimental ion currents in table 1 were treated as if they were relative partial pressures. The log K values thus obtained for the experimental temperatures are listed in table 3, together with the Gibbs-energy function changes and resulting third-law reaction enthalpies.In table 4 the average values for the reaction enthalpies are given for 298.15 and 0 K reference temperatures as well as the derived atomization energies, Di, 298.15 and DZ, o, and heats of formation, AH; 298.15 and AH; o. Here the errors quoted correspond to the estimated total errors connected with the measurement of the ion currents and temperatures and from the assumptions made for instrument constants, as well as for the estimated molecular parameters used in the calculation of the thermal functions. For obtaining theK. A. GINGERICH, R. HAQUE A N D M. P E L I N O 343 TABLE 1 .-EXPERIMENTAL CURRENTS~ (MULTIPLIER ANODE CURRENT IN A) OF LANTHANUM AND LANTHANUM CARBIDE SPECIES usm IN THE EVALUATION OF LaC,(g) AND Lac&) temperature La+ Lac,+ Lac: Lac; Lac,+ /K ( x 10-8) ( x 10-10) ( x 10-9) ( x 10-13) ( x 10-13) 2668 1.44 0.770 0.580 0.30 - 2708 2.025 0.75 2767 2.8 1 2.40 1.95 1.7 - 2799 3 .OO 2.70 2.22 2.2 283 1 3.25 3.00 2.28 1.7 2827 3.14 2.85 2.15 1.8 1 .o 2840 2.80 2.20 1.60 1.5 0.6 - - - _- - a The values correspond to the most abundant isotopes. TABLE 2.-cALCULATED GIBBS FREE-ENERGY FUNCTIONS (GEF), - (GO,, - Hi)/ r, IN J K-' m01-l AND HEAT-CONTENT FUNCTIONS (HCF), WT-G, I N kJ mol-l FOR LINEAR LaC,(g) AND Lac&) ~~~ C,-La-C,, C,-La-c, ~ __ _ _ _ _ ~ _ _ temperature/ K GEF HCF GEF HCF 298 286.13 21.31 292.55 23.30 2400 530.56 401.71 563.84 427.29 2600 543.34 438.73 578.18 469.15 2800 555.30 475.82 59 I .67 51 1.1 1 3000 566.56 5 12.93 604.3 1 553.12 3200 577.17 550.13 61 6.26 595.2 I atomization energies from the enthalpies of reaction (1) the heat of sublimation of graphite, A P V , 298.15 = 716.7 2.1 kJ mol-l and AHO,, = 71 1.2 2.1 kJ mol-', was taken from Hultgren et al.13 Likewise the atomization energies (in kJ molP1)l4 of Lac, ( A x , 298.15 = 1832.9 7; A x , = 18 19.1 f 7) and Lac, ( A x , 298,15 = 2544.2 _+ 7; AH:, = 2527.1 & 7) were used to obtain the corresponding atomization energies of Lac, and Lac8 from the enthalpies of reactions (2) and (3).In tables 3 and 4 the error terms correspond to standard deviations for the reactions involving the molecule Lac, and to the deviation from the mean for reaction (3).The atomization energies in table 4 have been based on the average value obtained from reactions (1) and ( 2 ) for LaC,(g). For LaC,(g) each data set measured was given the same weight. The heats of formation, AHf, ,, and A H f , 298,15, have been based on the corresponding enthalpies of reaction (1). The errors given for the atomization energies and heats of formation represent estimated overall errors. As supporting evidence that equilibrium conditions prevailed in the present investigation, a comparison between the results for the previously known molecules Lac,, Lac, and Lac, from the present investigation with those obtained under similar conditions in an independent study', and with other literature dataly is given. Table 5 shows that the third-law enthalpies for reaction (1) measured to temperatures of 2626 K in the present investigation are in excellent agreement with the corresponding344 ATOMIZATION ENERGIES OF Lac, AND Lac, TABLE 3.-THERMOCHEMICAL EQUILIBRIA INVOLVING GASEOUS Lac7 AND Lac8 AND THEIR EVALUATION BY THE THIRD-LAW METHOD La(g) + 7C(graphite) = LaC,(g) 2668 5.6812 135.55 65 1.8 2708 5.4314 135.44 648.4 2767 5.2183 135.28 650.7 2799 5.1347 135.19 653.5 283 1 5.28 14 135.10 668.7 2827 5.241 7 135.1 1 665.6 average AG = 656.5k7.8 Lac&) +Lac,@ = W g ) + Lac,(& 2668 2.0144 - 28.21.3 27.6 2767 1.991 1 -27.810 28.5 2799 1.9582 - 27.680 27.5 283 1 2.0927 - 27.552 35.4 2827 2.035 1 - 27.568 32.2 2840 1.9233 - 27.516 26.4 average A% = 29.6 k 3.2 La(g) + 8C(graphite) = LaC,(g) 2827 5.4969 142.51 700.4 2LaC4(g) = + LaC8(g) 2827 3.1679 - 37.868 64.4 2840 3.1829 - 37.827 65.6 average A x = 65.0k0.6 TABLE 4.-sUMMARY OF THERMODYNAMIC PROPERTIES OF Lac7&) AND LaC8(g) ~ ~~~~~~ ~ ~ A x (R1) 656.5 k 7.8 ' e 9 8 .1 5 (R1) 664.1 f 7.8 AG (R2, 3) 29.6 f 3.2 ' e 9 8 . 1 5 (R2? 3, 29.5 f 3.2 D: 4320 k 60 Di98.15 4351 60 A H f , l l 1088 f 50 AHf, 298.15 1095 4 50 156.9 f 1.9 158.7f 1.9 7.1 0.8 7.1 & 0.8 1033 f 14 1040t 14 260f 12 262+ 12 700.4 708.9 65.0 & 0.6 66.3 f 0.6 4989 f. 70 5024 f 70 1132+60 1140k60 167.4 169.4 15.5 kO.1 15.8 f 0 . 1 1193k 17 1201 17 271 k 14 2725 12 results in the literature, indicating the attainment of equilibrium. It is concluded that equilibrium must have prevailed at the even higher temperatures (2668-2840 K) at which the new molecules reported here have been measured.The observation of the molecules Lac, and Lac, under equilibrium conditions and their measured high stability supports the previously suggested analogy6 betweenK. A. GINGERICH, R. HAQUE AND M. PELINO 345 TABLE 5.-cOMPARISON OF THIRD-LAW ENTHALPIES, AG, FOR THE REACTION La(g) + nC(graphite) = LaC,(g) FOR n = 2, 3 OR 4, IN kJ temp. no. of range data A* n in Lac, /K sets /kJ ref. 2444-2626 2248-2609 2267-2600 2240-2650 2444-2626 2331-2609 2444-2605 2273-2609 24 1 7-2600 11 18 20 8 8 11 5 17 18 159.3 & 0.5 159.5 0.6 162.9 k 0.4 159 315.9+ 1.1 3 16.5 k 0.9 318.0 k 0.9 316.4+ 1.0 318.4f 1.3 this work 14 9 i this work 14 this work 14 9 a Error terms correspond to standard deviations. higher complex carbides and corresponding carbon species with the same number of atoms.Thus the corresponding carbon molecules C, and C, have been observed by Milne et al.lS above 3100 K, employing high-pressure sampling techniques to the effusate from a graphite Knudsen cell. The molecules C, and C, have also been observed by Steele and Bourgelas at 2860 and 2890 K under conditions of free evaporation. The fact that the carbon molecules C, and C, have not yet been observed under equilibrium conditions at the highest temperatures of the present investigation (2840 K) also indicates that the presence of the transition metal in Lac, and Lac, strengthens the bonding of the carbon atoms. How this takes place depends in part on the assumed geometry of these carbide molecules. From previous studies6T carbon chains with the metal atom in the end position appear to be a likely alternative.Optical spectroscopic information from matrix-isolation studies would be of great value to answer such questions. The authors are grateful for support given to this work by the Robert A. Welch Foundation under Grant A-387. 1 2 .3 3 3 6 7 8 9 10 11 12 13 W. A. Chupka, J. Bcrkowitz, C. F. Giese and M. G. Inghram, J . Phys. Chem., 1958,62, 61 1. G. DeMaria and G. Balducci, International Revue of Science, Physical Chemistry Series One, 1972, 10, 209. K. A. Gingerich, in Natl. Bur. Stand. ( U S . ) , Spec. Puhl., 1979, 561, 289-300. K. A. Gingerich, in Current Topics of Material Science, ed. E. Kaldis (North-Holland, Amsterdam, 1980), vol. 6, pp. 345-462. K. A. Gingerich, D.L. Cocke and J. E. Kingcade, Inorg. Chim. Acta, 1976, 17, L1. S . K. Gupta and K. A. Gingerich, J . Chem. Phys., 1979, 71, 3072. K. A. Gingerich and R. Haque, J . Chem. Soc., Faraday Trans. 2, 1980, 76, 101. S. K. Gupta and K. A. Gingerich, J . Chem. Phys., 1980, 72, 2795. F. J. Kohl and C. A. Stearns, J . Chem. Phys., 1971, 54, 5180. R. Haque, M. Pelino and K . A. Gingerich, J. Chem. Phys., 1979, 71, 2929. K. A. Gingerich, J . Chem. Phq’s., 1968,49, 14; D. L. Cockeand K. A. Gingerich, J. Phys. Chem., 1971, 75, 3264. E. K. Storms, The Rejractorj Carbides (Academic Press, New York, 1967). R. Hultgren, P. D. Desai, D. T. Hawkins. M. Gleiser, K. K. Kelley and D. D. Wagman, Selected Values of the Thermodimmic Properties of the Elements (American Society for Metals, Metals Park, Ohio, 1973).346 ATOMIZATION ENERGIES OF Lac, AND Lac, l4 K. A. Gingerich, R. Haque and M. Pelino, High Temp. Sci., 1981, in press. l 5 G. Herzberg, Molecular Spectra and Molecular Structure, vol. 1, Infrared and Raman Spectra of Diatomic Molecules (Van Nostrand, New York, 2nd edn, 1950), Appendix. l 6 E. B. Wilson Jr, J. C. Decius and P. C. Cross, Molecular Vibrations(McGraw-Hill, New York, 1955). C. E. Moore, Natl. Bur. Stand. (U.S.), Circ., 1958, 3, 467. IX T. A. Milne, J. B. Beachey and F. T. Greene, Vaporization Kinetics and Thermodynamics of Graphite Using the High Prexsure Mass Spectrometer, AFML-TR-74-57, May 1974 (U.S. Govt. Report). l 9 W. C. Steele and F. N. Bourgelas, Studies of Graphite Vaporization Using a Moduluted Beam Mass Spectrometer, AFML-TR-72-222, Sept. 1972 (U.S. Govt. Report). (PAPER O/ 1734)

ISSN:0300-9599    

DOI:10.1039/F19827800341

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I , 1982, 78, 347-356 Association between Polar Molecules Part 2.-Equilibrium and Thermodynamic Studies on the Dipole Association of Benzoni trile Derivatives with Hexamet hylp hosp horam ide, Di -n- bu ty 1 Sulphone, Di-n-butyl Sulphoxide and Tetramethylurea in Non-polar Solvents BY HIDEAKI FUJIWARA,* TATSUYA TAKAGI, YUTAKA YAMAZAKI A N D YOSHIO SASAKI Faculty of Pharmaceutical Sciences, Osaka University, 133- 1 Yamadakami Suita, Osaka, 565 Japan Received 2 1 st November, 1980 Dipole association between p-substituted benzonitriles and polar substances such as hexamethylphos- phoramide(HMPA), di-n-butyl sulphone(DBSN). di-n-butyl sulphoxide(DBSX) and tetramethylurea (TMU) is investigated in non-polar solvent by means of n.m.r. spectroscopy assuming the formation of a 1 : 1 associate between the polar molecules.Association parameters (association constant and association shift) are derived and discussed in relation to the dipole moment and the substituent constant. The association constant is, as expected, dependent upon the dipole moment of HMPA, DBSN, DBSX and TMLJ, and also upon the substituent constants of the p-substituted benzonitriles. However, the correlation between this constant and the dipole moment of the benzonitriles is the reverse of that expected. Polarization in the Ph-X fragment in benzonitriles which opposes that in the C E N bond is thought to cause this phenomenon. The association shift, which is the change in shift on association, cannot be explained simply by the dipole moment of HMPA, DBSN, DBSX and TMU, and so the magnetic anisotropy of these compounds is considered to be responsible.However, when a comparison is made within each benzonitrile + polar substance system, the anisotropy effect seems to be cancelled out and a correlation is estahlished with the substituent constants. Thermodynamic parameters are also estimated for the association of HMPA with substituted benzonitriles and for association with substituted nitrobenzenes and several aliphatic nitro and nitrile compounds. These are in the range expected for the dipoles under consideration, and the model adopted to interpret the concentration and temperature dependences of the shift is proved to be sound. In a previous report' we have investigated intermolecular interactions between p-substituted nitrobenzenes and hexamethylphosphoramide in non-polar solvents by analysing the n.1n.r.concentration dependence. These interactions are thought to be dipolar and the results were discussed with regard to the dipole moment and the substituent constant. These approaches using n.m.r. as a diagnostic tool seem suitable for the elucidation of weak intermolecular interactions in solution. Although such approaches have been widely utilized for the study of hydrogen bonding and acidic-proton-aromatic-system interactions,2 the reported results are very limited regarding dipole-dipole interactions. In the present study the interaction ofp-substituted benzonitriles with several polar substances which contain well-known polar groups such as those shown below was investigated to show the effect of several types of polar group on dipole association. \ "\ \ / c=o 347348 ASSOCIATION BETWEEN POLAR MOLECULES For compounds of types (I)-( IV), hexamet hylphosp horamide( H M PA), di-n- butyl sulphone(DBSN), di-n-butyl sulphoxide(DBSX) and tetramethylurea(TMU), respec- tively, were used.Compounds having n-butyl groups were selected for the sake of solubility. RESULTS We have analysed the n.m.r. concentration dependence by assuming the formation of a 1 : 1 associate between the interacting molecules. Because there exists a definite orientation force of an electrostatic nature between the two polar molecules and the concentration of these is very dilute in a non-polar solvent, the above assumption is considered justified.Job plots are consistent with this idea (fig. 1). This will also be confirmed if similar results (for the association constants) are reached when the equilibrium is observed from both sides for the two interacting molecules. However, such attempts were often disrupted by the poor solubility of the nitrobenzenes and benzonitriles in methylcyclohexane and by the very small changes in the lH shifts (< 1 Hz) in HMPA and TMU on addition of the polar molecules. 2.01 0 0.5 1 .o CBN/(%N + CHMPA) FIG. 1 .--Job plots for the p-tohitrile+ HMPA system in methylcyclohexane. Similar results are found for other systems. A = doobsd --aA; total concentration = 0.2 dm3 mol-l. 0, Experimental values for the ortho protons; A, experimental values for the meta protons. Although only meta protons with regard to the NO, group gave rise to considerable concentration dependences in the previous nitrobenzene + HMPA system, both ortho and meta protons do so in the present case of benzonitrile + polar substance systems.Therefore, the association constant(K) and association shift(A,,) are derivable for both protons. Sometimes a better agreement is reached between the K values from the ortho and meta protons if the infinite-dilution shift(6;) is used as the shift of A in the free state. An example is given in table 1 for the p-tolunitrile+HMPA and p-tolunitrile + DBSN systems. The table, which depicts a critical example, clearly shows that slight changes (only of the order of ppm) modify the K values favourably. Calculations are given in table 2 for all systems studied here.Averages of the K values from the ortho and meta protons are used for further discussion.H. F U J I W A R A , T. T A K A G I , Y. Y A M A Z A K I AND Y . S A S A K I 349 TABLE COMPARISON OF K VALUES DETERMINED FROM DIFFERENT VALUES OF THE SHIFT IN A FREE STATE^ calculation 1 calculation 2 ___- systemb site S i c A A Bd K S i c AABd K I 0 584.2 24.9 2.0 584.3 25.3 1.9 I m 560.8 33.9 1.4 560.1 32.4 I .7 I1 0 584.2 12.4 2.6 584.3 12.6 2.2 I1 m 560.8 16.2 1.3 560.1 13.6 2.2 a The shift and association constant ( K ) are expressed in units of Hz and dm3 mol-l, respectively; S;, shift of A from TMS at its initial concentration (0.03 mol dm-3, calculation 1); di, shift of A from TMS at its infinite dilution (calculation 2); I, p-tolunitrile+ HMPA, 11, p-tolunitrile+ DBSN system; A A B = S A B - 6, (association shift).TABLE 2.-ASSOCIATION PARAMETERS FOR THE p-SUBSTITUTED BENZONITRILE + POLAR SUBSTANCE SYSTEMS IN METHYLCYCLOHEXANE AT 34.1 OCa polar substances benzonitriles HMPA DBSN DBSX TMU p-substituent pBNh site A A B K AAB K AAn K AAB K H 3.97 CH, 4.40 C(CH3)3 4.64 OCH, 4.97 N(CH3), 6.60 CI 2.50 a AAB and K are 0 P m 0 0 m 0 m 0 m 0 m 0.277 0.288 0.297 0.363 0.310 0.310 0.221 0.426 0.127 0.352 0.506 0.297 2.2 0.132 2.7 2.5 0.143 2.5 1.9 0.140 2.3 1.7 0.152 2.2 1.7 0.140 2.3 1.7 0.126 1.8 2.7 0.150 2.7 3.0 0.172 2.8 1.7 0.112 1.9 2.2 0.159 1.9 4.9 0.234 3.7 4.6 0.199 3.5 0.176 0.183 0.173 0.199 0.206 0.204 0.181 0.27 I 0.093 0.203 0.300 0.260 2.2 1.6 1.9 1.7 1.3 1.5 2.1 2.1 1.6 1.7 3.0 2.9 0.209 0.160 0.203 0.269 0.22 1 0.221 0.167 0.297 0.077 0.258 0.350 0.330 1.1 1.4 0.9 0.9 0.8 0.8 1.2 1.3 1 .o 0.8 1.7 1.6 n units of ppm and dm3 mol-l, respectively.Errors are within & 1 % for pB,, dipole moments of benzonitriles from ref. (3), the value in benzene AAU and is adopted consistently. 5% for K . TABLE 3.-TWO-PARAMETER TREATMENT OF In K BY THE EQUATION In K = p gi + 4 0,‘ -k r polar substance P 4 r f” HMPA DBSN DBSX TMU 2.47 0.38 0.73 0.03 1.31 0.59 0.90 0.04 1.52 0.30 0.58 0.11 1.72 0.36 0.00 0.23b Fitting parameter f = (root mean square deviation)/(root mean square of the data). Relatively largefvalue is due to smaller value of K in this case.3 50 ASSOCIATION BETWEEN POLAR MOLECULES DISCUSSION ASSOCIATION CONSTANT Plots of the association constants for the systems containing DBSN, DBSX and TMU against those for the systems containing HMPA show a linear relationship between them.A least-squares fit of the data according to the relation y = ax results in K(DBSN) = 0.90 K(HMPA) r.m.s.d. = 0.52 dm3 mol-I (1) K(DBSX) = 0.73 K(HMPA) r.m.s.d. = 0.33 dm3 mol-1 (2) K(TMU) = 0.41 K(HMPA) r.m.s.d. = 0.17 dm3 mol-l. (3) The proportionality constants in these equations vary with the ratio of dipole moments @) of the polar substances: p T , R s N / p H M p A = 0.81, , u D B S X / / L H M ~ A = 0.71, and p T M x T / p H M p A = 0.63 from the values for HMPA(5.54 D),3 DBSN(4.46 D),4 DBSX(3.93 D)4 and TMU(3.47 D),4 where the value of di-iso-butyl sulphoxide is cited for DBSX. This fact confirms that the dipolar term is dominant in the intermolecular interaction and also that the experimental determination of K is accurate enough to disclose simple relationships.TABLE 4.-vARIATlON OF 31P CHEMICAL SHIFT OF HMPA ON ADDITION OF CYAN0 COMPOUNDS~ compound addedb AC/Hz benzene - 1 benzoni trile - 1 1 p-tolunitrile -6 p-anisonitrile -8 p-chlorobenzoni trile - 13 propionitrile -3 a Data from INDOR experiments on a Hitachi R-20B n.m.r. spectrometer (60 MHz for lH). 0.3 mol dmP3 of the compound is added to a solution of 0.03 mol dm-3 HMPA in A, Variation of the shift on addition of the cyano compound. The minus methylcyclohexane. sign denotes a down-field shift on addition. In contrast, when association constants are plotted against pBN in each benzo- nitrile + polar substance system, a deceptively simple relation appears : K decreases with pRN. In the benzonitriles polarization in the CN group is delocalized to the rela- tively large phenyl ring, and the p-substituent modifies the charge distribution in the CN group as well as that in the phenyl group.This may make the simple dipolar representation of the molecule invalid for a microscopic interaction in this case. Plots of the K values against the substituent constants5 oi and o,t show a correlation between the K and oi. The two-parameter treatment for In K gives sufficiently low values off(tab1e 3), especially for higher K values, which indicates a dominance by the oi term. This oi dependence and the above inverse proportionality of ,YBN could be accounted for by considering polarization in the Ph-X fragment, where Ph is the phenyl ring.This polarization would cause a local dipole moment which opposes that in the CN group or a quadrupole moment over the whole molecule, which contributes to the increased value of K. The oi and pBN dependence does not necessarily mean that the main site of interaction in benzonitriles is the phenyl ring. In fact, the 31P chemical shift of HMPA is seen to move down-field on addition of benzonitrilesH. F U J I W A R A , T. T A K A G I , Y. Y A M A Z A K I A N D Y. S A S A K I 35 1 (table 4), supporting the absence of the P atom above the phenyl ring in an associate. Association around the CN group also seems likely because of the predominant excess charges on this group. ASSOCIATION SHIFT The association shift A A R is determined from the shift of both ortho and metu protons with regard to the CN group (table 2).However, the substituent dependence is so small at the meta position that no discernible correlation is found for this, and so the discussion below concerns the ortho position. Plots of A A B in the systems containing DBSN, DBSX and TMU against those in the system containing HMPA reveal the following relation, which is similar to eqn (1)-(3): AA,(TMU) = 0.71 AA,(HMPA) r.m.s.d. = 0.01 ppm (4) A,,(DBSX) = 0.63 AAB(HMPA) r.m.s.d. = 0.02 ppm ( 5 ) AA,(DBSX) = 0.49 AAB(HMPA) r.m.s.d. = 0.03 ppm. (6) The proportionality constant is in the reverse order to the ratio of dipole moments of the polar substances in contrast to the case of K . If the association shifts arise exclusively from the increased polarization of CN group which is induced by the electric-field effect of the partner molecule and which will lead to an electron-density decrease in the phenyl ring,' they must clearly be correlated with the dipole moment of the partner molecule.Since this is not the case here, another factor must be taken into account. Thus the magnetic anisotropy effect of the partner molecule may be regarded as having masked the electric-field effect because of the observation of the shift at the ortho position. TABLE 5.-TWO-PARAMETER TREATMENT OF AA13 BY THE EQUATION AAH = a Oi + b 0,' + c' polar substance a b C .f" HMPA DBSN DBSX TMU 0.46 0.52 0.34 0.1 1 0.21 0.13 0.15 0.04 0.18 0.26 0.22 0.08 0.10 0.29 0.23 0.03 a ,f= fitting parameter [see footnote (a) in table 31.In contrast, if the AAR are compared in each system of benzonitrile+polar substance, some characteristics are disclosed regarding the pBN, ai and a,+. The two-parameter treatment for ai and a,+ is given in table 5, and the resulting coefficients as well as the fitting parameter are shown to be comparable with the previous case.' In the present case, as mentioned above, the magnetic anisotropy effect of the partner molecule is considered to contribute to AAB. Therefore, the satisfactory results of the two-parameter treatment support the assumption that the anisotropy effect is cancelled out when one considers the substituent effect on A A R with the same partner molecule. Both the lH and 31P low-field shifts in the A,, values (table 3 and 4) are consistent with the increase in intramolecular polarization caused by the association : increased electrical polarization in the polar molecule is enhanced by the neighbouring dipole of the partner molecule.352 2 .o k E: 1 .2 - 0 . 4 ASSOCIATION BETWEEN POLAR MOLECULES I I 1 2.8 3.2 3.6 1 0 3 ~ 1 ~ FIG. 2.-van't Hoff plots for the dipole association between p-methoxynitrobenzene and HMPA. Association constants are determined from the shift data of the meta protons with regard to the nitro group. 0, Run I ; A, run 2; 0, run 3. J: Lo 6. 0 0 . 4 0.8 C HMPA FIG. 3.-Concentration and temperature dependences of the 'H shift of the metu protons in the p-methoxynitrobenzene + HMPA system in methylcyclohexane. Concentrations are ca. 0.03 mol dm-3 (constant) for nitrobenzene and ca.0-0.7 mol dmP3 for HMPA. Temperature: 1, 262.4; 2, 287.4; 3, 307.2; 4, 337.3; 5, 360.8 K.H. F U J I W A R A , T. T A K A G I , Y. Y A M A Z A K I A N D Y. S A S A K I 353 THERMODYNAMIC PARAMETERS Enthalpies (AH') and entropies (ASO) of association can be determined from van't Hoff plots after measuring the temperature dependence of the association constants (fig. 2). These plots are obtained from the dependence of the shift on concentration as well as on temperature (fig. 3). It is seen in fig. 3 that the addition of HMPA drastically enhances the temperature dependence of the shift. This phenomenon is readily interpreted as being due to the enhancement or suppression of the association by temperature. The AH0 values thus derived amount to 7-12 kJ mol-l (table 6), and an enthalpy-entropy compensation relation holds approximately (fig.4). The dipole-dipole interaction energy is expressed as :6 For the pair HMPAb = 5.54 D) and nitrobenzeneb = 3.93 D) at 300 K, this energy amounts to -8 kJ mol-1 when r = 5.5 A, and to - 16 kJ mol-1 when r = 5 A; these values are comparable with the enthalpies observed in the present study. Furthermore, the enthalpies may be regarded as comparable with those reported for the dimers of acetone or acetonitrile, i.e. - (1 6-22) kJ mol-l in the gas phase;7 the rather large value for the latter case might be ascribable to a difference in the dielectric constant of the media. TABLE 6.-ENTHALPIES (ALP) AND ENTROPIES ( A F ) FOR THE ASSOCIATION OF HMPA WITH VARIOUS NITRO AND NITKILE COMPOUNDS IN METHYLCYCLOHEXANE compounds - AHO/kJ mol-l - A F / J K-l mo1-1 (I) nitrobenzenes (p-substituent) OCH, CH, D OCOCH, c1 N(CH,), (11) benzonitriles (p-substituent) N(CH,), CH3 C(CH,), OCH, c1 (111) others acetonitrile nitromethane ni troe t hane 7.2 11.9 9.9 8.3 10.3 9.3 8.3 11.8 9.5 9.6 8.1 11.4 12.8 12.4 16.0 28.7 29.2 23.1 26.0 22.8 22.8 27.1 26.6 28.7 20.4 26.7 26.2 29.1 Although a definite correlation would be expected from eqn (6) between the AH0 and p values, no relations were observable between them.In contrast, a regular relation holds for plots of AH' against a,+. These results are similar to the plots of In K against p or 0; in our previous report. Therefore, entropy terms are thought to take part in the association in a regular manner, causing the AHO-As0 relationship and the similar behaviour for both In Kand AHO.For a full understanding of the AH03 54 30 5 2 5 - E - L4 c, --. 3 2 0 - 1 5 - ASSOCIATION BETWEEN POLAR MOLECULES - / 5 / 1 B /'O 7 6 9 12 -AH"/kJ mol-' FIG. 4.--Enthalpy+ntropy relationship for association between polar molecules. Solvent : methylcyclo- hexane. Polar substances: 0, p-substituted nitrobenzenes; A, p-substituted benzonitriles; 0 , others (8, acetonitrile: 9, nitromethane; 10, nitroethane). p-Substituents; 1, N(CH,),; 2, OCH,; 3, CH,; 4, C(CH,),; 5, D; 6, OCOCH,; 7, C1. 2 . 1 h % u, a v h z 5 2.0 W z (.o 1 . 9 0 0.4 0.0 HMPA FIG. 5.-Concentration and temperature dependences of the 'H shift of the acetonitrile protons in the acetonitrile+HMPA system in CCl,.Temperature: 1, 255.3; 2, 280.3; 3, 306.4; 4, 328.3; 5, 344.3 K.H. F U J I W A R A , T. T A K A G I , Y . Y A M A Z A K T A N D Y. S A S A K I 355 value, a quantitative estimation of the association energy would be necessary which takes into account the multipole interaction as well as that for dispersion. These studies are now underway using molecular-orbital calculations. In this report a saturated hydrocarbon is used as solvent. We have found in a thermodynamic study that CCl, is not suited as an inert solvent for the present study. That is, the shift of the nitrile compounds does show a significant dependence on temperature even if the interacting partner molecule is excluded from the solution (fig. 5). Such a phenomenon is observable in other systems although reduced to ca.f . These facts suggest that the intermolecular interactions between the polar substances and CCl, is considerable. A charge-transfer interaction might be responsible for this." Therefore, for a full interpretation of the data for CCI,, some investigation of solvation effects caused by the solvent is needed. This will be the subject of a future study. EXPERIMENTAL p-tert-butyl- and p-NN-dimethylamino-benzonitriles were synthesised from p-tert-butyl- benzoic acidg and p-chlorobenzonitrile,lO respectively. Other benzonitriles were of commercial grade. Solid benzonitriles were recrystallized several times and stored over silica gel after drying under vacuum. Liquid ones were distilled over CaH, under reduced pressure and stored over molecular sieve 4A.HMPA, DBSN, DBSX and TMU were distilled similarly using BaO as a desiccant and stored over molecular sieve. N.m.r. spectral measurements were made as in a previous report' using methylcyclohexane as solvent. Although the shifts were measured relative to the internal TMS standard, the same results were obtained in the present work when one of the solvent methyl peaks was adopted as the standard. 31P n.m.r. shifts were measured by the lH-("'P} INDOR method on a Hitachi R-20 B n.m.r. spectrometer operating at 60 MHz, using a synthesizer (Anritsu MGSI4C) and an r.f. power NH, CN CL CI CL CN CN [ref.( 11 c 11 SCHEME 1 .-Synthesis of the deuterated benzonitriles.356 ASSOCIATION BETWEEN POLAR MOLECULES amplifier (Hitachi R-209 PA); the shifts were determined relative to the internal TMS standard.The temperature used was 34 "C. Variable-temperature experiments were carried out according to a conventional procedure ; the probe temperature was measured with a thermocouple(CC) immersed in methylcyclohexane. The temperature range used was typically ca. 0-90 "C. Temperature fluctuations during the shift measurement for over 10 samples were within kO.5 "C (in < 2 h). The temperature was varied ca. 5 times for each series of the sample preparation, and such measurements were repeated more than twice to test the reproducibility of the experiment. Errors in the measurement of AH0 and AS" were estimated to be k0.4 kJ mol-l (5%) and k 1.2 J K-' mo1-l ( 5 7 3 , respectively. Since ortho and meta protons in p-substituted benzonitriles usually give AA'BB' spectra, the accurate determination of the absolute shift of both protons is not straightforward.However, the spectral pattern does not vary so much on the addition of the polar substance that the relative variation of the shift (and hence K and A A B ) can be determined accurately by the measurement of two intense peaks in both the AA' and BB' regions. Ex'ceptionally p- chlorobenzonitrile showed spectra of J/Ad 5 1 besides the large change in spectral pattern on addition of the polar substance. Accordingly, [o,u'-~HJ- and [m,m'-2H,]-p-chlorobenzonitriles were prepared, and K and A A B were determined separately. [m,m'-2H2]- and [o,o', p2H,]- benzonitriles were also synthesised to obtain the association parameters for the benzonitriles. The synthesis of these deuterated compounds is shown in scheme 1.A typical run of the acid catalysed H-D exchange reaction of anilines is as follows: 5 g aniline, 20 cm, D,O and 3 cm3 37% DCI were sealed in an ampoule at ca. lo-" Torr using a vacuum line. After heating at 140-150 OC for 40 h check was made of the extent of the deuterium exchange by n.m.r. (ca. 90 atom % usually); the aniline was then freed from its deuterochloride by 504 NaOH and extracted 3 times by 50 cm3 ether. A crude product of the deuterated aniline was obtained (above 95% yield) which was used for the next diazotization reaction, followed by an appropriate treatment to give final products. The acid-catalysed H-D exchange reaction in the evacuated state was found to be very successful because of the lack of any subsidiary oxidation reactions. The determination of K and A A I 3 was made in the same way as in the previous report using a NEAC S-900 computer at the Computation Centre, Osaka University. ' H. Fujiwara, T. Takaba, Y. Yamazaki and Y. Sasaki, J . Chem. Soc., Faraday Trans. I , 1979,75, 79. i? ( a ) P. Laszlo, Prog. Nucl. M a p . Reson. Spectrosc., 1967, 3, 231. (b) A. S. N. Murthy and C. N. R. Rao, App. Spectrosc. Reo., 1968, 2, 69. (c) J. C. Davis and K. K. Deb, Adt.. Magn. Reson., 1970, 4, 201. A. L. McClellan, Tables of E.xperimental Dipole Moments (W. H. Freeman, San Francisco, 1963). J. A. Riddick and W. B. Bunger. Organic Solvents (John Wiley, New York, 1970). Y. Yukawa and Y. Tsuno, Nippon Kagaku Zasshi, 1965, 86, 873. T. M. Reed and K. E. Gubbins, Applied Statistical Mechanics (McGraw-Hill Kogakusha, Tokyo, 1973), p. 161. ' (a) A. 1). Buckingham and R. E. Raab, J . Chem. Soc., 1961, 551 1. (b) T. A. Renner and M. Blander, J . Phys. Chem., 1977, 81, 857. (c) D. J. Frurip, L. A. Curtiss, and M. Blander, J . Phys. Chem., 1978, 82, 2555. R. Anderson and J. M. Prausnitz, J . Chem. Phys., 1963, 39, 1225. C. Ruechardt and S. Eichler, Berichte, 1962, 95, 1921. lo E. B. Pederson, J. Perregird and S - 0 . Lawesson, Tetrahedron, 1973, 29, 421 1. l 1 These reactions were achieved by reference to similar ones: (a) H. T. Clarke and R. R. Read, Org. Synth., 1,1932,514. (b) W. W. Hartman and M. R. Brethen, Org. Syrzth., 1.1932,162. (c) N. Kornblum, Org. Synth., 3, 1955, 295. (PAPER O / 1804)

ISSN:0300-9599    

DOI:10.1039/F19827800347

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I , 1982, 78, 357-375 Ionic Contributions to Partial Molal Volumes in Aqueous Solutions BY JEAN V. LEYENDEKKERS School of Biological Sciences, Building A 12, University of Sydney, 2006, New South Wales, Australia Received 26th November, 1980 The ionic components of the partial molal volumes of aqueous electrolyte solutions at 25 O C have been analysed on the basis of the TTG model with reference to molecular data. Twenty cations and seventeen anions were considered. The limiting value, DO, for each ion is summarised by oo(ion) = V(edi) + A V(ed H,O) + A V(HB) + V(Born), where V(edi) is the effective volume of the ion in solution, which includes the electrical deformation effects. V(edi) = Mtys = $nNx 10-24(r+Ai\Keom+A,,)3, where r is the crystal radius (Pauling), Ageom is the geometric packing factor and Ael the electrical deformation effect.For small cations Ae, is proportional to the number of water molecules that have lost their orientation polarization per mole per unit charge. AV(ed H,O) is the volume change due to the electrical deformation of the water molecules close to the ion and equals kzA,,, where z is the charge on the ion and k is a function of the internal pressures and dielectric properties of pure water. A V(HB) is the volume change associated with rupture or formation of hydrogen bonds and V (Born) is the Born volume. The Masson (or TTG) slope S , was split into the ionic components via an equation given previously. The cations all have positive slopes (in cm3 kgb mol-4 for I$, where I is the molal ionic strength) ranging from 1.03 for H+ to 5.67 for Ce3+.The monovalent anions have small slopes (usually < 0.5 and negative). The F- ion ( S ; = l.l8), HCO; ion ( S ; = 0.83) and OH- ion ( S ; = 2.22) are exceptions. The divalent oxyanions have large (> 3) positive slopes. The deviations of the ionic slopes from the Debye-Huckel slope are the result of an anion trersus cation factor, plus the specific ionic effect on the reorientation time of the water molecules (resulting either from rotational or translational diffusion). All the slopes can be summarised by the relationships and where Q represents either the self-diffusion coefficient of water or the reciprocal of the proton magnetic relaxation rate. For the halide ions, Q can also be represented by the relative molal shift of the dielectric relaxation time.This result and those for 8' are consistent with the concept of hydration water exchange of dielectrically mismatched water molecules near small cations. 1. INTRODUCTION Friedman and Krishnanl have discussed the problem of obtaining absolute single-ion properties and cite reviews on the subject. There is still some controversy' about whether the absolute single-ion values for the limiting partial molal volumes, ijo, are known. However, the values from measurements of ionic vibration potentials are in good agreement with those derived on the basis of various assumptions that are all mutually consistent.' Hepler and Woolley2 have selected i j o for H+(aq) at 25 O C as - 5 f 1 cm3 mol-l as the 'best' value on the basis of a large number of investigations which they cite.Most of the recent tabulations of uo(ion)1*3 are based on a value of 357358 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES - 5.4 cm3 mol-l for H+(aq) which agrees with the 'best' value range. These values are used in the present paper. The most recent compilation of v"(ion) values at 25 OC is that of Akitt,4 based on - 5.8 cm3 mol-l for H+(aq), which is also in the 'best' value range. Akitt's compilation derives from older tabulations but a large body of more recent data was also used in preparing the tables. The newer data, covering a large variety of salts (but few new ions), confirm the assumption of additivity. An important consequence of knowing the values of single-ion properties is that the role of the ionic charge in the solvation process can be more easily elucidated.In addition, the use of model concepts is often initiated since the relative values of the ionic properties indicate trends that can be associated with molecular and other properties of the solute and solvent. Attempts can thus be made to sort out the various contributions to the single-ion value. In this way, the interpretation of solution processes can be made more realistic. One of the simplest models for ionic solvation is that due to Born1 whereby the ions are represented as charged hard spheres and the solvent as a fluid with a dielectric constant D which is uniform even in the presence of ionic fields. The solvation free energy, AGo(ion), for this model, applied to water, (1) is given by AGo(ion) = -(Nz2e2/2r')[1 -(l/D,)] where N = 6.023 x e is the electronic charge, r' is the radius of the ion and the subscript w represents water.Throughout this paper z represents the charge (positive) of the ion. The corresponding ionic hydration volume, V(Born), is obtained by differentiating eqn ( I ) with respect to pressure to give V(Born) = - Az2/r' (2) where with co a constant and p the pressure. A = 4.175 cm3 A mol-1 at 25 OC.l to the volume occupied by the ion in solution. That is, A = c,(a In D,/ap)/D,, Eqn (2) has been widely usedl3 with the residual volume of v"(ion) being ascribed (3) where r is the crystal radius, C is a constant and A depends on the ion. Both C and A vary according to the model The Born model and related Debye-Pauling model together with the various refinements have been discussed in some detail by Friedman and Krishnan.' Alternative equations to eqn (3) usually have the same form, i.e.an intrinsic-volume term and a hydration (electrostriction) term. For example, Couture and Laidler5t developed the empirical equation vo(ion) = C(r+ A)3 + V(Born) uo(ion) = h, + b, r3 - 262 (4) where b, and b, have values that differ according to whether the ion is monatomic or is an oxyanion. Akitt4 has recently developed an equation analogous to eqn (4), based on the assumption that the electrostriction is a linear function of z. Conway and coworkers7~* had shown earlier that, for regions close to the ion, the electric field is such as to give an electrostriction volume that varies approximately linearly with z.For the intrinsic-volume interpretation Akitt4 assumed the dimensions of the solute to be secondary to the structural considerations (or molecular configurations). On the other hand, the Chemical Modell interprets the volume as being made up of four terms, viz. where V(int) is the actual volume of the ion, V, is the change in volume of the water vo(ion) = V(int)+ V, + hI + V(Born) ( 5 )J. V. LEYENDEKKERS 3 59 that enters the type I cosphere state, Qr is the change in volume of water that enters the type I1 cosphere state and V(Born) is the volume change associated with the polarization of the water by the field of the ion (according to the Born model). The type I cosphere state is one in which the water is oriented by ionic fields or other directional solute-solvent forces.The type I1 cosphere state is one in which the water is perturbed by the proximity of a solute particle but the effect cannot be ascribed to directional solute-solvent forces. In the present analysis an equation analogous to eqn (5) is developed on the basis of the TTGt model and using molecular data (from infrared, Raman, dielectric and nuclear magnetic resonance spectroscopy). The TTG molal volume, My,,9 is equated to V(int), & is taken to represent the volume change due to the electrical deformation of water molecules near the ion and is taken to represent the volume changes associated with hydrogen-bond rupture or formation. The last term is given by eqn (2). In addition, the effect of concentration changes of the solute is considered.The differences between the TTG and Debye-Hiickel (DH) slopes for the ions are equated to the changes in the molecular reorientation time of water using nuclear magnetic resonance (n.m.r.) and dielectric relaxation data. 2. DISSECTION O F T H E LIMITING PARTIAL MOLAL VOLUME O F AN ION From the TTG model the intrinsic volume of the ion is given by M y , ( i ~ n ) . ~ The corresponding volume for the salt is related to the compressibility and density of the solution and of pure water and values of M y , for a large number of salts have been calculated from such data.g In addition, an equation has recently been derivedg which enables M y , to be split into the ionic contributions [eqn (7)]. With the intrinsic volume thus determined, the residual volume of uo(ion) is ascribed to the hydration effects (i.e.the effects on the water that cause volume changes). The hydration volume is subdivided into components along the lines of the Chemical Model [eqn ( 5 ) ] . First, the volume V(Born) is calculated from eqn (2) using the Pauling crystal radius; V(Born) is then subtracted from the hydration volume to give a volume, V(ex). The use of r in eqn (2) is considered more appropriate than a modified radius. This is because in terms of a refined Born or Debye-Pauling model r' [eqn (2)] should equal r when adjustment of the dielectric constant is used (such an adjustment varies with the ionic species). Whilst adjustments to the dielectric constant are not used here in the same sense, the form of refinement is analogous.That is, the neglect of the molecular nature of the solvent in deriving V(Born) is compensated for by the contribution of V(ex). V(ex) is in fact a consequence of dielectric changes. In summary, the TTG equation for ijo(ion) will have the form ijo(ion) = My,(ion) + [ V(ex), + V(ex),] + V (Born). (6) The V(ex) volume has been split into two components to conform with the concepts of eqn (5). V(ex),, V(ex), represent the volume changes in type I and type I1 cosphere states, respectively. Because of the links with other models eqn (6) might be considered a hybrid equation.' However, this is only in relation to the interpretation of the components of the hydration volume, since the intrinsic volume derives directly from the TTG concepts. Each of the components of eqn (6) will now be analysed in more detail.Molecular data will be used as a guide to ensure the interpretation is realistic. t Tammann-Tait-Gibson.360 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES 2.1. THE IONIC TT(; MOLAL VOLUME 2.1.1. CALCULATION OF My,(ion) volume, M y , , has recently been d e r i ~ e d , ~ viz. where and v represents the number of moles of the subscripted ion. &., is a volume contribution attributed to the specific shape-structure characteristics of the ions (including microdynamic effects). Eqn (7) can be split into the ionic contributions, e.g. An equation relating the standard partial molal entropy, SO, and the TTG molal so = 1.6 1 3(Mys - Kh) - 22.258 w at 25 OC (7) w = t( v+ 2: + v- 22) v+sy = 1.6l3[v+(Mys+- Kh+)]- 11.129 V+Z: for the cation, and similarly for the anion.Since values for SO+, s? and <h+, <h- are available,1~3~9 the ionic TTG molal volumes can be estimated from eqn (8). The comments in the Introduction concerning the uncertainties inherent in values of single-ion properties of course apply here. The reported values for the ‘absolute’ partial molal entropy of H+(aq) (in the usual hypothetical 1 mol kg-l standard state at 25 “C) range from - 1.1 to -6.3 cal K-l mol-l and Hepler and Woolley2 chose -4+ 2 cal K-l mol-1 as the ‘best’ value. A value of 0.0 cal K-l mol-1 is used here. This is close to the ‘best’ value range and has been found to give the most consistent results, both previously and for the relationships derived below. Since K h is negligible for Na+ and C1- the ionic values for this quantity could be obtained in a straightforward way.9 Table 1 lists the values of the various ionic quantities in eqn (8).The latter was used to calculate My,(ion) values wherever possible in order to average out inconsistencies in the data. Otherwise the principle of additivity was used. The My,(ion) values, when combined, give values for the 47 salts listed previouslyg to within kO.1 cm3 mol-1 on average, with a maximum deviation of 0.8 cm3 mol-1 for less reliable data. 2.1.2. CALCULATION OF RADIUS OF ‘EQUIVALENT’ SPHERE The volume My,(ion) can be analysed in terms of an ‘equivalent’ sphere as in eqn (3). The sphere has a radius of (r+A) where r is the crystal radius (Pauling’s in A or 100 pm) and A is a term usually regarded as defining the positional correlation of the ion and adjacent water molecules, or the ‘packing effect’.lT3 Thus, (9) Values of A were calculated from eqn (9) and are listed in table 1.The values for the monovalent cations are generally around 0.8 A with an average of 0.85 A whilst values for the halide ions are very small. These values agree with those found to apply for free energies, entropies and enthalpies of cations and halide ions calculated via the Debye-Pauling mode1.l Such consistency indicates that the values of My,(ion) are reasonable. The values of A for the tetra-alkylammonium ions are also small, whiIe the values for the oxyanions are similar to those of the monatomic cations. Mly,(ion) = 4,nN x 10-24(r+A)3 = 2.52(r+A)3 cm” mol-l.2.1.3. INTERPRETATION OF A Friedman and Krishnan’ have discussed the significance of A in some detail. They suggest a need to include in A an allowance for ‘the interplay of repulsive forces between the ion and the solvent and between solvent molecules’ as well as the geometric packing effect. It seems clear that A is made up of at least two terms, andJ. V. LEYENDEKKERS 36 1 TABLE COMPONENTS OF THE TTG MOLAL VOLUME M y , (25 OC, 1 bar) ion so &h MY, r + A r A A,, H+ Li+ Na+ K+ Rb+ c s + NH,+ Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Zn2+ Cd2+ La3+ Ce3+ Me,N+ Et,N+ nPr,N+ nBu,N+ F- c1- Br- I- OH- NO3- ClO,- SCN- HC0,- C032- CrOd2- so,2- 0.0 3.4 14.4 24.5 29.7 28.7 31.8 26.9 26.4 - - 55 - 28.2 - 13.2 - 9.4 3.0 - 25.5 - 14.6 - - 36.5 - - - - - - - 2.3 13.2 19.3 26.1 - 2.5 - 35.0 43.6 36.0 22.5 9.2 4.1 - 12.9 - 1.7 - 4.2 0.0 0.8 3.0 7.2 2.5 9.7 4.1 0.3 - - - 0.3 - 1.5 - 4.2 - 3.8 - 18.0 - - - - - - - - 2.3 -0.3 0.3 4.5 - 0.6 - 4.4 - 0.8 - 0.8 -4.8 - 1.6 - 3.3 - - 5.2 13.21 15.83 22.89 28.3 1 27.70 33.81 26.08 25.77 3.0 14.2 19.71 2 1.47 27.96 7.59 12.08 22.9a 2 1.47 96.1 - - 156 222.5 286.2 - 3.17 14.78 19.16 27.58 4.75 24.20 38.2 28.42 20.05 14.80 31.7 26.84 - 1 .273 1.737 1 .845 2.087 2.240 2.220 2.376 2.179 2.171 1.060 1.763 1.985 2.042 2.230 1.444 1.686 2.087 2.042 3.366 3.956 4.453 4.843 1.080 1.803 1.966 2.220 1.235 2.126 2.475 2.243 1.995 1.804 2.326 2.200 - - - - 0.36(?) 0.0 1 0.68 0.97 1.33 1.47 1.67 1 .03b 0.35 0.66 0.99 1.12 1.34 0.74 0.97 1.14 1.07 3.05' 3.10 3.69 3.92 4.2 1 4.68 1.36 1.81 1.96 2.16 1.03 0.96 1.24 1.52 1 .4d 1.15 1.26 1.6 1.5 - - 0.913 1.263 1.057 0.875 0.757 0.77 0.75 0.706 1.149 0.71 1.103 0.995 0.922 0.890 0.704 0.7 16 0.947 0.972 0.316 0.266 0.266 0.243 0.163 - - - 0.28 - 0.007 0.006 0.060 0.205 0.275 0.886 0.955 0.84 0.845 0.544 0.726 0.70 0.363 0.7 13 0.507 0.325 0.207 0.22 0.20 0.156 0.599 0.16 0.55 0.445 0.372 0.340 0.154 0.166 0.397 0.422 - - 1.06 -1.11 - 1.11 - 1.137 - 1.217 -0.83 -0.557 -0.544 - 0.490 - 0.345 - 0.275 - 0.494 - 0.425 -0.54 -0.535 -0.836 -0.654 - 0.68 - so/cal K-l mol-l [ref. (l), (3) and (9)]; Ch/cm3 mol-l [ref.(9)]; M~,u,/cm~ mol-l [eqn (8 from eqn (9), r / 8 , is the Pauling crystal radius or interatomic distance (oxyanions) [ref. (l)]; Ael = A-Ageom with Ageom = 0.55 8, for monatomic cations, NH,+, OH- and halides and 1.38 8, for other ions; interatomic distance; R,N+ radii consistent with V(ex) results and ref.(1); mean for S-C and C-N interatomic distances, r(H+) is uncertain, 0.36 8, is consistent with overall ionic trends (see text). unless indicated]. a From 4; and K O data for salt, using principle of additivity, (r+A) in R362 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES it should be useful to analyse the residuals (A-Ageom), where Ageom is the geometric effect. G1ueckauf,lov l 1 using purely geometrical considerations (i.e. before deformation of the arrangement by the electrical charge takes effect), assumed that the packing space of the ion can be taken to correspond to a hollow sphere of constant thickness, with the additional condition of the isomorphic replacement of a water molecule by an ion of radius 1.38 A (radius of the water molecule) which, for spherical ions, leads to Ageom = ( VH2,/@V)i - rw = 0.55 A with VH,, = 18 cm3 (molal volume of water).For small tetrahedral oxyanions, (e.g. ClO;, SO:--, AsOi-) the ‘packing’ sphere ‘encloses already a fair amount of dead space when the oxygen atoms form the starting points of the water structure, utilizing the lone pairs of the oxygen atoms. ’lo Thus Ageom = rox ( I .38 A) and r of eqn (9) becomes MO (the radial distance of oxygen centre from the anion centre). In the present paper, the interatomic distance plus rox has been taken for the oxyanions. Assuming that these values of Ageom are correct, then the deviations of A from them (table I ) must be the result of the electrical deformation, A,, (table 1).The deviations are positive for the cations (the electrical effect increases their apparent size) whilst for the anions and tetra-alkylammonium ions the deviations are negative (the electrical effect offsets the geometrical increase in size). 2.1.4. RELATIONSHIP OF A A N D THE LOCAL DIELECTRIC PROPERTIES At low temperatures the static dielectric constant is a parameter highly sensitive to the amount of structural correlation between the dipoles. Therefore, as a check on the above results dielectric relaxation data should be useful. Pottel12 gives background information on the dielectric relaxation in aqueous ionic solutions, so that only a few points needed to explain the equations will be given here. From a simple modeP2 of the dielectric relaxation in solutions of non-dipolar solutes in polar solvents the static permittivity E, of the solution is reduced relative to E,,, of pure water according to &, = &o, w ( l -f +if) ( 1 1) wherefdenotes ‘that fraction of the total volume of the solution which is occupied by the solute particles, regarded as dielectric spheres with permittivity E , = 2.A solute particle may consist not only of an ion but also of those adjacent water molecules which have lost their orientational polarizability’,12 by the influence of the ion. Eqn (1 1 ) should be valid at least up to 1 mol kg-l. On the basis of the TTG model where hop is the volume of the water molecules that have lost their orientational polarizability (per mole of salt), and V is the volume of the solution, m is the concentration, mol (kg H,O)-l.Using the TTG equation for V,339 this gives m V,op = ( l 03uw + m d J f - mMy/, = ( I O3 + mM)f/dsol - mMy/, wheref= ( E ~ , , - E ~ ) / ( ~ E , + E ~ , ~ ) from eqn ( 1 I), dsol is the density of the solution, u, the specific volume of pure water, dV the apparent molal volume at concentration m. At zero concentration all terms are zero (since f is then zero). Eqn (1 3 ) may also be expressed where x is the concentration of the salt per g of solution. ( 1 3 4 h o p = .fM/xdso, - My/,J . V. LEYENDEKKERS 363 TABLE 2.-sTATIC PERMITTIVITY CHARACTERISTICS OF ELECTROLYTE SOLUTIONS AT 1 f?'l [mOl - salt HCl LiCl NaCl KCl RbCl CSCl M a BaCl, LaCl, Na,SO, NaOH Me,NBr Et,NBr H+ Li+ Na+ K+ Rb+ Cs+ M g2+ Ba2+ La3+ F-, C1 -, Br-, I-, R,N+ OH- ions so,,- NH; Ca2+ Sr2+ Ce3' NO3- c10, co32- CrO,*- (kg H,O)-'] (25 OC, 1 bar) 58.33 64.912 66.7 68.1 69.1 70.6 47.43 50.5 34.5 56.3 57.5 66.612 61 .O 19.01 18.47 18.45 28.79 34.09 41.32 20.25 30.43 27.9 23.63 - 2.52 115 169 170 95.4 75.6 56 41.8 22.2 266 219.5 404 153.7 173.4 0 0 9.3 5.3 4.2 3.1 2.3 1.3 14.8 12.2 22.4 8.5 9.6 0 0 9.3 5.3 4.2 3.1 2.3 1.3 14.8 12.2 22.4 0 5.4 0.1 ( 1 3.8) ( 1 2.8) (5.7) (22) (2) (2) ( 1 ) (0) 4, data from ref. (3), E ~ , ~ = 78.3 [ref.(12)], bracketed values were estimated from other nlop values, assuming linearity as in fig. I , t', = 1.002961 cm3 g-' [ref. (3)]. The limiting value of the change in E, with concentration is obtained as follows, (14) from eqn (1 I), with from eqn (12). The last term drops out if cop does not change appreciable with &,/am = - E,,, 1.5@f/am)/( 1 + v/am = ( M W s + VopV v- m(Mvs + Vop) (a v/am)/ v2 + (4 v) a v;,p/am,364 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES concentration (which seems valid at moderate concentrations).12 As the concentration approaches zero (denoted by L), (15) Eqn (14) gives the relationship between the partial molal volume (aV/2m) and the partial molal permittivity.When data are available this should be a useful relationship because, as discussed below, Fop only depends on the cation for many salts. Further, the limiting equation, eqn (15), indicates why M y , and its derivative dMys/dt are important components of limiting s10pes.l~ Values of Fop, and hence n, the number of water molecules occupying the volume [i.e. ( Vpp/ 18) moles] were calculated from eqn (1 3 ) with rn taken as 1.The values are shown in table 2 together with the values of Pottel.12 The agreement between the two sets of values is good. Assuming zAel is proportional to the work done on the ion in changing its volume ( i e . causing the electrical deformation of the ion, table l), then this quantity should be proportional to nlOp, the number of molecules that have lost their orientational polarization in doing the work (at least for small cations).12 This is illustrated in fig. 1, where the line for the small cations passes through ZA = 0.55 when nlop is * zero. This is apparently the limiting value for the series and is consistent with Ageom from eqn (10). There are few data for the anions but assuming a linear relationship as for the cations, the values of nlop for some oxyanions can be estimated (table 2).The predicted value of ca. 2 moles of water for the ClO; ion is compatible with Raman data.14 The latter data indicate that each mole of C10; is associated with 2 moles of (a&,,/am), = - 1.5 & o , w ( ~ y s + vOP)/ 10"~. 0 4 8 12 16 20 24 *lop FIG. 1 .-Ionic dielectric characteristics. nlop is the number of water molecules that have lost their orientation polarization due to the effect of the ion of charge z. A is from the TTG molal volume, M y s , uiz. 2.52(r+A)3 with r the crystal radius (Pauling).J . V. LEYENDEKKERS 365 disordered water (non-hydrogen-bonded), and that the effect is practically independent of concentration.From the above, it can be concluded that (Y+A) represents the effective radius of the ion in solution. The halide ions have zero effect on the orientation polarization as do the tetra-alklylammonium ions.12 This is verified also uiu eqn (13) (table 2). Apparently these ions bond in some special way with water, in contrast to the oxyanions and small cations (Section 2.2). The radii of the tetra-alkyl ions are not known with certainty but are estimated from mode1s.l The values used by Davies et aE.I5 are much lower than those of Friedman and Krishnan.l However, the latter are very close to the radii consistent with the present results (table 1, Section 2.2). 2.2. THE IONIC TTG HYDRATION VOLUME 2.2.1. CALCULATION AND INTERPRETATION OF V(ex) The terms suitable for representing the hydration volume, V(ex) and V (Born), have been discussed above [eqn (6)].These quantities will now be calculated and an attempt made to interpret the significance of V(ex) (that is, interpret the molecular processes giving rise to this volume change). From eqn (6) V(ex) = v"(ion) - Mly,(ion) - V (Born) (16) where V(ex) = V(ex),+ V(ex),. Values of V (Born) at 25 "C (table 3) were calculated from eqn (2) using the crystal radius (Pauling) for Y' (table 1). Values of V(ex) (table 3) could then be estimated from eqn ( I 6) using the data in table 1. All the cations (with the exception of Be2+ and H+) have negative values of V(ex) whereas the anions have positive values. The results in Section 2.1 suggest that the volume change V(ex) for the small cations and possibly the oxyanions is a consequence of the fact that polarized water molecules entering the inner hydration sphere (cosphere I) lose their preferred orientation.Since water molecules leaving cosphere I (i.e. entering cosphere 11) have no preferred orientation there arises what may be called an electrodynamic mismatch boundary. As a result of this, structural tensions are set up in cosphere 11, probably leading to the breaking of hydrogen bonds. Non-hydrogen-bonded groups of water exhibit strong cohesive interactions which are not highly directional, so that the concept of hydrogen-bond breaking in cosphere I1 would be compatible with the Chemical Model. For this model it is assumed that the water in the type I1 cosphere state is perturbed but the effect is not associated with directional forces.Since zAel is directly proportional to the volume of the water molecules that have lost their orientational polarization, a correlation between V(ex) and zAel could be expected for small cations and the oxyanions. Such a correlation could also apply to the other ions, for even though they do not cause any loss of orientational polarization to the adjacent water molecules they do induce electrical distortions in these water molecules and affect the local hydrogen-bond configuration. For example, n.m.r. spin-relaxation measurements indicate partial immobilization of water by alklyl groups,16 whilst halide ions appear to form hydrates with adjacent water molecules, and the fluoride ion is believed to substitute for H,O on the water 1atti~e.l~ When zAel and V(ex) were compared for all the ions it was found that the following relationships closely apply: ( 1 7 4 V(ex) = kz(A,, + 0.83) for anions and tetra-alkylammonium ions V(ex) = kz(Ael - 0.83) for small cations (17b)366 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES TABLE 3.-cOMPONENTS OF V(eX) (25 OC, 1 bar) H+ - 5.4 Li+ - 6.3 Na+ - 6.6 K' 3.6 Rb+ 8.7 Cs+ 15.9 NH,+ 13.0 Be2+ - 22.8 - Mg2+ - 32.0 - Ca2+ - 28.6 Ba2+ - 23.3 Zn2+ - 32.4 La3+ -55.3 Ce3+ - 56.3 Me,".84.2 Et,N+ 143.7 nPr,N+ 209.3 nBu,N+ 270.3 F- 4.2 CI- - 23.2 Br- 30.1 I- 41.6 OH- 1.4 Sr2+ - 29.0 Cd2+ - 30.8 - - - NO3- 34.4 cm- 49.5 SCN- 41.1 HC0,- 28.8 C032- 6.5 CrO,,- 30.5 so,,- 24.8 - 5.2 13.21 15.83 22.89 28.31 27.7 33.81 26.08 3.0 14.2 13.74 19.71 2 1.47 27.96 7.59 12.08 22.9 21.47 22.1 1 96.1 - 156 222.5 286.2 3.17 14.78 19.16 27.58 4.75 - 24.2 38.2 28.42 20.05 14.80 31.70 26.84 - 1 1.6h 6.14 4.3 3.14 2.84 2.5 4.05 47.6 25.32 16.8 14.92 12.4 22.5 17.2 33.3 35.1 - - .- 1.37 1.35 1.13 0.99 0.89 3.0 2.3 2.13 1.93 4.05 4.35 3.36 2.75 3.0 3.63 13.44 10.4 11.2 - 1 .o - 13.36 - 18.13 - 16.15 - 16.77 - 16.16 - 15.41 - 9.03 21.8 - 20.9 - 20.4 -31.51 - 35.55 - 38.86 - 17.5 - 25.7 - 44.9 - 42.67 -43.31 - 10.53 - 10.55 -11.17 - 12.21 - 15 4.03 10.72 13.07 15.93 0.7 1 .o 13.56 14.05 15.68 12.38 5.14 9.2 9.16 - 15.72 22.0 14.105 8.96 9.55 8.68 6.77 26.04 13.89 47.80 49.48 38.63 32.32 29.51 13.37 14.41 51.7 54.94 57.68 - 46.0 -48.17 -48.17 -49.35 - 52.82 - 36 - 24.3 - 23.44 -21.3 - 14.97 - 11.94 -21.44 - 18.23 - 23.44 - 23.22 - 72.56 - - 56.77 - 59.02 - 14.7 - 35.4 - 32.2 -25.1 - 25.6' - 22.2 - 35.1 7.9 - - 69 - 70 - 70.1 - 67.9 - 68.4 - 30.9 - 40.1 - 97 - 97.6 - 101 35.5 37.6 37.0 37.1 37.8 40 35.0 36.5 37.2 15.7 12.9 35.0 32.5 39 35.6 77.7 65 68.2 __ -2 - 5 -4.5 - 3.5 - 3.5 - 3.0 - 5 + 1 - 9.5 - 10 - 10 - -9.5 - 9.5 -4 - 5.5 - 13.5 - 13.5 - 14 + 5 + 5 + 5 +5 + 5.5 +5 + 5 + 5 +2 + 5 +4.5 + 5.5 +5 + 10.5 +11 + 9 + 9.5 - - (14.4) (36) (32.4) (25.2) (25.2) (21.6) (36) (7.2) (68.4) (72) (72) (68.4) (68.4) (3 1.6) (39.6) (97.2) (97.2) (1 00.8) (36) (36) (36) (36) (36) (36) (36) (36) (36) (79) - - (39.6) (14.4) - (32.4) (39.6) (75.6) (64.8) (63.4) " In cm3 mol-l; with r = 0.36 A; mean; 8' [ref.(l)], Mtys (see table I), V(Born) eqn (2), V(ex) eqn (16), k = 43.4 cm3 mol-1 k1 (see text), Ael from table 1.where k = 43.4 cm3 mol-l A-l; an interpretation of this constant is given below (Section 2.2.2). It can therefore be assumed that V(ex), = kzAe, ( 1 8 4 V(ex), = kzA,, (18W andJ. V. LEYENDEKKERS 367 where AHB z k0.83. Since the hydrogen-bond effects of the ions can be expected to vary with the type of ion, the small variations in AHH are probably a consequence of this. For the small cations the quantity kzAe, is directly proportional to nlop. From fig. 1 , kzAe1 = A,[ V(H2O) nlopl (19) where A, = 0.26 (dimensionless). Thus, the volume change V(ex), associated with these ions is about one-quarter the volume of the water molecules that have lost their orientational polarization. Since M y s is independent of concentration, A , will change according to the changes in eOp with concentration (negligible at moderate concentrations).12 The partial molal volume change for hydrogen-bond formation in water, 8,-fiD, is ca.7.2 cm3 m ~ l - ’ . ’ ~ Since 0.83 x 43.42 = 362, this represents 5 moles of hydrogen-bonds for unit charge for a mole of ions. The number of hydrogen-bonds affected by a particular ion has been estimated by subtracting kzA,, from V(ex) and dividing by 7.2 (table 3). These results predict that the small cations are associated with bond-breaking (except Be2+) whilst the other ions form hydrogen bonds. The hydrogen-bond effects tend to cancel out for the salts because of the opposing signs. The ions of the tetra-alkylammonium salts, however, reinforce each other so that the structure-making effect is greatly enhanced.2.2.2. INTERPRETATION OF k I N EQN (17) The complexity of the interplay of microdynamic and other forces in water is very great1* and the extent of the volume change due to electrical deformation of the water molecules near the ion and bonding changes (cospheres I and 11) will be reduced because of opposing forces in the bulk. Since B,, the Tait parameter for water,3, l9 represents the difference between the expansive pressure (due to thermal energy) and the cohesive pressure (due to molecular interactions that are independent of temperature), the constant, k , in eqn (17) could be expected to have the form which is dimensionally correct also (table 4). It is found thatf(B,) given by [BT+2P(st)] fits the results at 25 OC very well and gives realistic temperature trends for k.P(st) is the ‘structural’ internal pressure3 and (20) is given by where c’ = 0.4343 x 0.315, T is the Kelvin temperature and Pw the compressibility of water. A, is a constant related to the refractive index of water and is independent of temperature and pressure but slightly dependent on the 20* 21 A, is zero for liquids such as methanol and benzene and is a measure of the deviation of water from ‘normal’ b e h a v i ~ u r . ~ ~ 21y 22 Eisenberg21 has outlined the molecular significance of A, in terms of a model involving types of oscillators with given polarizabilities and suggests that An might reflect a change with temperature of the average polarizability of water molecules. The molecular significance of B, has been analysed by G i ~ ~ e l l .~ , 23 (-AnT/Pw), Or [-AnT[BT+ l)lC’1 (21 a) Using k = (e2N/Dw eJ/f(BT) where f(BT) = Br,(1-2A, T/c’), since B, 9 1, the values of AHR for Na+ at 0 and 40 O C were estimated and are compared with the value at 25 O C (table 4). The temperature trends of nlOp are consistent with those of nHB which indicates that the form off(&) is reasonable. Further evidence thatf(B,) is appropriate comes from the relationship between B,368 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES TABLE 4.-EFFECT OF TEMPERATURE ON k IN EQN (21 a) AND nHB FOR THE Naf ION (VOLUMES IN cm3 mol-l) A H B / A eqn ( 16) temp BT3 k V(Born) V(ex) AeJA and / O C /bar eqn (21a) eqn (2) eqn (16) eqn (9) (18a, b) UHB nlOpl2 0 2672 87.9 42.1 -3.7 -23.3 0.40 -0.95 -5.6 5.7 25 3005 78.3 43.4 -4.3 -18.13 0.325 -0.74 -4.5 4.6 40 3080 73.2 46.2 -4.8 -16 0.33 -0.68 -4.5 4.6 A , = 6.6 x deg-l [ref.(3)]; (e2N/D,ri) = 93 x lo3 cm3 mol-l A-l bar at 25 O C . and the dielectric constant. Recently, Bradley and P i t ~ e r ~ ~ fitted D, data to a form of Tait equation (covering the ranges 0-350 "C and 1-5000 bar). These results indicate that for 1 bar and at low temperatures (ca. 0-50 "C) alnD,/@ = C/[B,(l -fl T)] wheref, is a positive constant and C is slightly dependent on temperature. Thus, the coefficient A of V(Born) [eqn (2)] can be written A = C'/B,(l -fl T) which is similar in form to eqn (21b). For an organic liquid the pressure differential of the dielectric constant is inversely proportional to the corresponding Tait parameterlg? 25 (i.e.fl does not appear). This indicates thatf, is an index of the special structure of water (relative to normal liquids) just as A , is in eqn (21 b), A, being zero for normal liquids as noted above. 2.3. SUMMARY FOR 8" The results for the small cations are clear-cut in that the contributions of Ageom and A,, seem unambiguous. This also applies for the oxyanions such as SO:- and C10;. However, the division of A is not so certain for other ions. The splitting of V(ex) into components also has the same defect. Nevertheless, the results under the scheme adopted appear consistent with the present concepts of hydration (e.g. the Chemical Model') and molecular data (Sections 2.1.4 and 2.2). The interpretation of 0" in terms of the TTG model can be summarised as DO = V(edi) + A V(ed H20) + A V(HB) + V(Born) (22) where V(edi) is the volume of the ion, including the electrical deformation effect, AV(ed H20) is the change in volume due to the electrical deformation of the water molecules near the ion, AV(HB) is the volume change due to the breaking or making of weak hydrogen-bonds and V(Born) is the usual Born effect [eqn (2)]. In quantitative terms these contributions to 0" become V(edi) = Mv,(ion) = iniV x 1 O-24(r")3 = 2.52(r + Ageom + Ae1)3 with Y the crystal radius (Pauling); Ageom is from eqn (10) for spherical ions and is 1.38 A for oxyanions and (presumably) tetra-alkylammonium ions (Section 2.1.3).Ael is given by eqn (1 7 a or b) and eqn (19) or eqn (9). A V(ed H20) = [(e2N/D, r&)/flBT)] zAe, = kzA,,J. V.LEYENDEKKERS where AllT) is given by eqn (21 b) and BT = 2671.8+ 19.454t-0.27028t2+9.798 x (bar) for 0-45 "C 369 AV(HB) = V(ex)-AV(ed H,O) (ca. f0.83 kz at 25 "C) with V(ex) from eqn (16). V(Born) is given by eqn (2). 3. CHANGES O F THE IONIC PARTIAL MOLAL VOLUMES WITH CONCENTRATION 3.1. CALCULATION OF SLOPES The apparent molal volume of an aqueous electrolyte, q$", and the corresponding partial molal volume 0 are given by:9 4, = q$:+s,I~+s,I (23) v = V0+ l.5SVI+f2S,I (24) where showng that, for 25 "C, = 8" and I is the ionic strength [mol (kgH,O)-l]. It has recently been (25) S, = wsDH +0.20+0.0444(~+~- 1.4[v-f+nAP(H20)]) where M? = iz+z-v, sDH = 1.865 cm3 kgi mol-8, with AP(H,O) taken as the entropy of ionization of water and n the number of moles of water that associate or ionize due to one mole of salt.For strong electrolytes n = 0. The salts of most oxyanions and salts that form complexes have n = + 1 (or +2, etc. for 2 moles of anion). Salts such as NH4Cl and LaCl, have n = - I , as the effect is due to the cation. From the analysis of tio (Section 2) the H,O entropy term of eqn (25) might arise from hydrogen-bond effects. If 2 moles of bonds are involved the entropy change is 17 cal K-l rnol-l, which is close to the value for the association of water, viz. 19.3 cal K-l mol-l. The difference in the two entropy terms only represents a difference of 0.1 cm3 kgt mol-g in Sv. As shown later (Section 3.2) the hydrogen-bond effect in this case would be a consequence of events close to the ionic surface. The deviations of S, from the DH slope were interpreted as arising from structural effect^.^ Since the ionic contributions to eqn (25) are additive these can be considered separately and compared with molecular data so that the type of structural effects involved can be more clearly understood.Splitting eqn (25) into the ionic components (26) gives St = O.~Z$Y,,, + 0.1 +0.0444.~0, - 0.0444( 1.4)n+ As"(H,O) (27) Usually the water entropy term is related to the anion so that n, = 0 (except NH;, La3+ and Ce3+). The value of n- is zero for the halides. The values of the ionic slopes calculated from eqn (26) and (27) are listed in table 5 . Zeldes26 calculated the individual slopes for the ions of the alkali halides via an iterative least-squares method, using the data for fifteen salts. These values are in good agreement with those calculated from eqn (26) and (27) (table 5).Since the data used by Zeldes extended up to very high concentrations (e.g. 13.8 mol kg-l for LiC1) the individual ionic volume effects are additive over a wide concentration range; this is implied by eqn (26) and (27). The values of S, in eqn (23) and (24) are usually small. These results indicate that eqn (26) and (27) should give good estimates of the effects of concentration [I$ mol(kg H,O)-l] for the alkali halide ions. As a check on the overall reliability of the slopes it should be profitable to relate them to ionic spectral data.l2, 27 S; = 0 . 5 ~ 2 ~ ~ ~ + 0.1 - 0.0444( 1.4) [s? + n-As"(H,O)]. FAR 1 13TABLE 5.--I: SLOPES FOR THE IONIC PARTIAL MOLAL VOLUME (25 O C , 1 bar) cation nu B+ a' bd a-b anion nu B- a' H+ 0 Li+ 0 Na+ 0 K+ 0 Rb' 0 c s + 0 NH: - 1 Mg2+ 0 Ca2+ 0 Be2+ 0 - 1 Sr2+ 0 Baz+ 0 Zn'+ 0 Cd2+ 0 Pb2+ 0 Ce3+ - 1 0.06 0.14 -0.12 0.06 - 0.08 -0.01 0.0 I - 0.04 0.0 1 - 0.05 0.03 0.45 - 0.28 0.27 -0.21 0.23 -0.18 0.18 -0.14 0.20 1.76 1.26 1.17e 1.76 1.49e 2.19 2.21e 2.38 2.21e 2.44 2.37e 2.14 2.69e 3.26 3.25e 3.50 3.49e 3.81 3.81e 3.69 1.03 1.18 (1.25) 1.18 1.67 (1.78) 1.67 2.12 (1.93) 2.12 2.35 2.35 2.44 (1.88) 2.44 1 .o 1.39 0.19 2.58 2.58 3.24 3.24 3.41 3.41 3.96 3.96 2.70 3.18 3.82 5.67 (1.94) 0.73 0.08 -0.01 0.09 -0.18 0.07 0.09 0.03 -0.14 0.00 - 0.07 - 0.44 0.1 1 0.02 0.0 1 0.09 0.08 -0.15 -0.15 -0.13 F- c1- Br- I- OH- CN- NO,- NO,- Br0,- c10,- c10,- SCN- HC0,- co,2- Cr0,'- so,2- so42- 0 0 0 0 + 1 0 + 1 + 1 + 1 + 1 + 1 0 + 1 + 1 + 1 - 1 + l 0 + 1 0 + 1 0 0.14 -0.12 5 .-0.01 0.0 1 - 1 - 0.04 0.04 -3 - 0.08 0.05 0.18 -7 - 0.04 - 0.05 - 0.05 - 0.06 - 0.08 - 0.08 - 0.07 0.04 0.25 0.18 0.22 0.12 1.8 1 .77e 1.08f 0.23 0.21e 0.18f - 0.09 -0.15e -0.12f -0.51 - 0.27e - 0.72f 2.22 -0.1 - 0.2 - 0.2 - 0.3 -0.5 -0.5 -0.41 0.75 5.76 3.38h 5.0 3.3 1 5.44 3.35h 4.39 3.25h hd a-b 1.18 1.18 1.18 0.21 (0.28) 0.2 1 0.21 -0.17 (0.06) -0.17 -0.17 -0.59 (-0.34) -0.59 -0.59 2.13 - 0.7 + 0.4 0.06 -0.15 -0.2 - 0.48 - 1.139 0.07 0.83 5.78 3.43 4.5 3.26 5.14 3.94 4.76 3.56 0.6 0.6 -0.1 0.02 0.00 - 0.03 0.08 0.02 0.05 0.08 0.3 0.09 0.6 -0.13 -0.6 -0.26 -0.15 -0.3 - 0.03 0.72 - 0.48 - 0.08 - 0.02 - 0.05 -0.5 0.05 0.3 - 0.6 - 0.4 -0.3 a n number of moles of water [eqn (26) and (27)]; B coefficients [mol (kg H,OI-'] from ref.(27), unless otherwise indicated; eqn (28) and (29) unless otherwise eqn (32) and (33); J eqn (34) with Bd [IO-Z indicated; mol (kgH,O)-'l from ref. (12): eqn (26) and (27), derived from density data and table 1, bracketed values from ref. (26) (see text); entrotw data uncertain. ref. (1): ean (30).J. V. LEYENDEKKERS 37 1 3.2. COMPARISON WITH NUCLEAR MAGNETIC RESONANCE A N D DIELECTRIC If &(ion) values are compared with the corresponding values for the limitingeffect of the ion on the proton magnetic relaxation rate q, i.e. [( 1 / TJ'ntra]-l [d( 1 /T)intra/drn],,, (with L = rn -+ 0) or B,.,.,., the following set of equations are derived for 25 O C , (28) (29) where sAH = 0.5z:(1.865) and similarly for the anion. The values of Bn.m.r.are effectively zero for K+ and C1- (within experimental error),27 so that [$AH+ 1.2 = S,(K+)] and [s,, -0.6 = S,(Cl-)]. Values of the slopes calculated from eqn (28) and (29) are listed in table 5. The values of S ; for the divalent oxyanions are different for symmetric charge systems (e.g. MgSO,) and non-symmetric systems (e.g. Na,SO,). For the latter n = 0 and for the former n = + 1. This was shown previo~sly,~ but it was found that if so for the anion MSO; (e.g. NaSO;) was used, the discrepancy disappeared. Since S O and Bn.m.r. values for these types of ions are not readily available it is more convenient to use a modified equation. For example, for eqn (25) n is taken as zero for the non-symmetric systems and s" for the divalent anion is used.For the n.rn.r. data the following equation DATA St = s ~ I H + 1.2-6.2Bi.m,r. S , = sfi, -0.6+ 10.5B;.m.r. was found to apply for unsymmetric charge systems involving the divalent oxyanions. Obviously, the formation of the ion M(0xy)- reduces the disturbance of the water. The deviations (a-b) in table 5, for eqn (28)-(30), are partly due to differences in the ionic effects on the bulk water (water without an ion as nearest n e i g h b o ~ r ) . ~ ~ This can be better understood by dividing Bn.m.r. into its components, uiz. for the cation, z:, 5; and TO, are the rotational correlation times of water in the cationic and anionic hydration spheres (on the immediate surface of the and in bulk water, respectively, An additional * superscript represents the limiting values as rn approaches zero, 0.018nh (or nh/55.5) is the ratio of water near the ion to the total number of moles in the solution.The first term of eqn (31) is the limiting effect of the ion on the bulk water and is usually small for most ions. This quantity is proportional to (i31,/c?rn),l~ (pe is the TTG effective pressure) and can vary appreciably for some ions (e.g. the divalent oxyanions, Mg2+, H+ and F-) giving rise to the larger deviations from eqn (28)-(30) for these ions. These effects are generally small so that the above equations give a good summary of the data, and can be used to predict the slope or, if the slope is known, to predict the other properties. On the other hand, if slope, Bn.m.r. and S O are all known, the deviations give information on the competing effects.Even at infinite dilution the correlation time of water is changed because of the presence of the ions. For example, at 25 OC for pure water, the correlation time zEo = 2.5 x s whilst for the water close to, say Na+, zC+O = 5 x lo-" s and for Mg2+, zC+O = 1.3 x 10-l1 s.,' A change in the molecular reorientation time of water means a change in the structure16 and it is in this sense that deviations from the DH slope occur in the case of the partial molal volume. 13-23 72 IONIC CONTRIBUTIONS TO PARTIAL MOLAL VOLUMES Since 1/T, cc 1/D, where D is the self-diffusion coefficient of water,2i equations similar to eqn (28) and (29) for D can be expected. Although the data are less reliable than for it is found that &(ion) is well-represented (table 5 ) by = sAH + 1.2 + 8( 1 /Do) (?D+/c?m), (32) S , = s~H-0.6- 12(l/D,)(?D-/?~t?)L (33) where D+ and D- are the water self-diffusion coefficients in the cationic and anionic hydration spheres, respectively, and D, is the corresponding coefficient in pure water.The concentration range for which the mean water correlation time is shorter than that of pure water extends up to ca. 6 mol kg-1.27 This is because z;O and z;O describe the reorientation motion for water close to the ionic surface, and explains why S , applies from very low to very high concentrations. The dielectric and proton magnetic relaxation data contain the same information about the influence of the halide ions on their nearest-neighbour water molecules.12 This is illustrated via the equation (34) S;(halides) = S G ~ - 0.6 + 0.15Bi where B; [in of the dielectric relaxation time Y.uiz. mol (kg H,O)-l] is the anionic component of the relative molal shift (1 /Yw) (dY/dm), B, = B$ + B,. ( 3 5 ) Values of S; for the halides from eqn (34) are listed in table 5. On the other hand, the dielectric and proton magnetic relaxation data contain different information about the influence of small cations.12 The dielectric effect has been discussed in relation to the static permittivity (Section 2.1.4), i.e. the hydration water exchange of molecules that have either lost or retained their orientation polarization. The n.m.r. data relate to the reorientation of the vector connecting two protons within the H,O of the first hydration sphere (i.e. the rotational correlation time of the intramolecular proton-proton vector in the hydration complex).27 Therefore there is no correlation with Bd+ in line with eqn (34).The cations all have positive values of St which means they increase the volume, a change usually attributed to structure promotion. The sign of B:e,,r. does not affect the result. The anions have both positive and negative values of S ; . However, within the experimental errors, the anions with negative slopes (structure-breaking effect) all have negative values of B L . ~ . ~ . . Thus, the n.m.r. classification of structurez7 agrees with the volume classification for anions but not for cations. This is because of the constant factor in S$., (uiz. -0.6 for anions and + 1.2 for cations). All the eqn (28)-(30) and (32)-(34) can be represented by ( S , - SDH)ion = A , + AH,O(config) (36) where AH,O(config) is proportional to the change in the dynamic configuration of the water molecules near the ion (which is specific to the ion).A,, however, only varies according to whether the ion is a cation or anion. A , is probably related to the configuration of the ion relative to the nearest-neighbour water molecule. There seems general agreement that the configurations are specific for anion versus cation.l*! Alternatively, since (St - SF') and ( S ; - SVc1-) are proportional to the orientation effects it is the deviations from sDH for these two ions that are of interest. There might be a contribution due to a change in some non-electrostatic contribution when the state of the formal charge of the ion changes1 For the present, the consistency of the set of equations derived for SFn [eqn (26)-(30) and (32)-(34)] indicate that the values in table 5 are probably close to the true values.J.V. LEYENDEKKERS 373 4. DISCUSSION In common with other studies concerning u"(ion), the first assumption is that the single-ion values are equal to or at least close to the absolute values. In the present analysis, the next assumption is that Mv/,(ion), as calculated, represents the intrinsic volume of the ion, with reasonable accuracy. The consistency of the results (in terms of the values of r + A ) with experimental results for free energies, entropies and enthalpies,l together with the correlation with qOp support this second assumption. The recent study by Akitt4 generally gives quite different values for the intrinsic volumes. Amongst the exceptions are the values for Na+, Ca", Sr2++, Ba2+ and La3+ which are in good agreement with the values in table 3.These ions are characterised by complete or nearly complete dielectric saturation.12 The agreement of the intrigsic volumes for these ions is taken as evidence that the TTG values are reliable (the values being additive). The reason is that these ions are more likely to satisfy the assumptions of Akitt's m0de1.~ Hence, intrinsic volumes calculated via Akitt's model for these ions will be more reliable than the corresponding volumes for other ions. The ions Sr2+ and La3+, in particular, feature with high weight in the derivation of the empirical equation used by Akitt.4 In addition, the theoretical justification for the latter equation does not take into account the large variations in the degree of dielectric saturation of the ions.The value of the hydration volume is given by subtracting My,(ion) from co(ion), and this is assumed to be made up of V(ex) and V(Born). In the analysis of the components of this volume reliance has been placed on the concepts of the Born (or Debye-Pauling) model and the Chemical Model,' together with evidence from molec- ular data. The assumption here is that the approximation of the hydration effect that V(Born) represents is simply additive with respect to the additional hydration effects, due to the molecular nature and special structure of the solvent water. The resulting equation is a function of z and Y, which seems reasonable, and the coefficients are compatible (having similar functional dependencies, involving for example the dielectric constant and internal pressure functions).V (Born) is proportional to z 2 / r (constant T and p ) , whereas, under the same conditions, the component V(ex), is proportional to z , or approximately so. V(ex),, however, depends not only on z but also on the electrical deformation of the water molecules. The latter affects the volume according to the type of deformation and also the number of water molecules involved. The interpretation of the components of V(ex) gives consistent results for all the ions and is compatible with the details of the Chemical Model states of cosphere water,' and with results from dielectric relaxation measurements and the conclusions from Raman and i.r.spectra (Section 2.2). V(ex) has been assumed to result from the electrical distortion of the water molecules near the ion [causing the volume change V(ex),] and subsequent disruption in the hydrogen-bond population in cosphere 11, or the boundary region with cosphere I [causing the volume change V(ex),]. These assumptions are supported by the correlation of nlop with V(ex), and the fact that the number of hydrogen-bonds estimated to be affected are rea~onable.~~? Since non-hydrogen-bonded groups exhibit strong cohesive interactions that are not highly directionalli?l8 the change in concentration of such groups would account for the hydration effect of the second kind (state I1 cosphere water of the Chemical Model') which is classified as 'states of cosphere water in which the water is perturbed by the proximity of a solute particle but the effect cannot be ascribed to directional solute-solvent forces ' .l The results for the small cations are compatible also with the work of Vaslow.28 By using the general series expansion for the potential due to an arbitrary charge distribution and the quadrupole moment of a water molecule deduced by3 74 IONIC CONTRIBUTIONS T O PARTIAL MOLAL VOLUMES Buckingham31 he was able to show that the minimum potential energy of small positive ions (e.g. Li+, Na+, K+ and Cs+) is not on the dipole axis of a water molecule but at a substantial angle to that axis. He found that the variation in energy of an ion (over large regions of the hemisphere of a water molecule, away from the hydrogen atoms) is smaller than the energy of breaking a hydrogen bond.The conclusion from this is that hydration for small ions is consistent with the normal structural groupings in liquid water. Vaslow visualised the hydrogen-bonds near the ion as stretching or bending. This agrees with the conclusion that the hydrogen-bonds break in the 'mismatch' boundary away from the ion. Vaslow's values for the minimum electro- static energies of the ions in the field of a water molecule (cylindrical symmetry assumed) at angles 51'42, 30 and 12' (for Li+, Na+, K+ and Cs+, respectively) to the axis along the dipole (two-fold axis), are directly proportional to zA and hence to the number of water molecules that have lost their preferred orientation. In considering the magnitude of the effect of the ion concentration on the solution volume only the value of vo(ion) and the slope S p are needed.In this regard, the dissection of z;'(ion) (important in the physical interpretation of the slope) is of secondary interest. The values of S P derived here appear reasonable in view of the iterative results based on the principle of additivitylg9 26 and the correlations with molecular data. The latter are significant since n.m.r. spectral data provide fairly unambiguous information on the ionic effects.27 Even for the correlations where the ionic contributions are not so clear-cut (e.g. BFn values from dielectric relaxation data12) the consistencies between the thermodynamic and spectral data analyses are noteworthy.As shown here, the multiple components of the total structural effect give opposing volume changes so that the conflicting structural classification for electrolytes and ions in the literature is understandable. The reasons why M y s and dMys/dt are important structure indicatorsgl 1 3 7 32 also are brought out. I H. L. Friedman and C. V. Krishnan, in Water: a Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1973), vol. 3, chap. 1. L. G. Hepler and E. M. Woolley, in Water: a Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1973), vol. 3, chap. 3. :3 J. V. Leyendekkers, Thermodynamics of Seawater as a Multicomponent Electrolyte Solution (Marcel Dekker, New York, 1976), part I. J. W. Akitt, J . Chem. Soc., Faraday Trans. I , 1980, 76, 2259. A. M. Couture and K . J. Laidler, Can. J . Chem., 1956, 34, 1209. A. M. Couture and K . J. Laidler, Can. J. Chem., 1957, 35, 207. B. E. Conway, R. E. Verrall and J. E. Desnoyers, 2. Phys. Chem., 1965, 230, 157. J. V. Leyendekkers, J . Chem. Soc., Faraday Trans. I , 1981,77, 1529. * J. E. Desnoyers, R. E. Verrall and B. E. Conway, J. Chem. Phys., 1965, 43, 243. lo E. Glueckauf, Trans. Faraday Soc., 1964, 61, 914. l 1 E. Glueckauf, Trans. Faraday Soc., 1968, 64, 2423. R. Pottel, in Water: a Comprehensice Treatise, ed. F. Franks (Plenum Press, New York, 1973), vol. 3, chap. 8. l 3 J. V. Leyendekkers, J. Chem. Soc., Faraduy Trans. I , 1980, 76, 1206. l 4 T. H. Lilley, in Water: a Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1973), l5 J. Davies, S. Ormonroyd and M. C. R. Symons, J . Chem. Soc., Faraday Trans. 2, 1972, 68, 686. l6 F. Franks, in Hydrogen-bonded Solvent Systems. ed. A. K. Covington and P. Jones (Taylor & Francis, London. 1968), p. 3 1 . l 7 G. E. Walrafen, in Hydrogen-bonded Solcent Systems, ed. A. K. Covington and P. Jones (Taylor & Francis, London, 1968), p. 9. I R Water: L( Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1972), vol. 1. I y H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolyte Solutions (Reinhold, New York, 1958). vol. 3, chap. 6.J. V. LEYENDEKKERS 375 2o J. V. Leyendekkers and R. J. Hunter, J . Phys. Chem., 1977, 81, 1657. 21 H. Eisenberg, J . Chem. Phys., 1965, 43, 3887. 22 E. Reisler and H. Eisenberg, J . Chem. Phys., 1965, 43, 3875. 23 R. Ginell, J. Chern. Phys., 1961, 34, 1249. 24 D. J. Bradley and K. S. Pitzer, J . Phys. Chem., 1979, 83, 1599. 25 B. B. Owen, R. C. Miller, C. E. Milner and H. L. Cogan, J. Phys. Chem., 1961, 65, 2065. 26 H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolyte Solutions (Reinhold, New York, 27 H. G. Hertz, in Water; a Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1973), 28 F. Vaslow, J. Phys. Chem., 1963, 67, 2773. 29 C. N. R. Rao, in Water; a Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1972), 30 S. Levine and J. W. Perram, in Hydrogen-bonded Solvent Systems, ed. A. K. Covington and P. Jones 31 A. D. Buckingham, Discuss. Faraday Soc., 1957, 24, 15 1. 32 J. V. Leyendekkers, J . Solution Chem., 1979, 8, 853. 1958), p. 397. vol. 3, chap. 7. vol. 1, chap. 3. (Taylor & Francis, London, 1968), p. 115. (PAPER O/ 183 1)

ISSN:0300-9599    

DOI:10.1039/F19827800357

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 371-381 Equimolar Mixtures of Divalent Transition Metal Perchlorates as Constant Ionic Media in Studies on Complex Formation in Non-aqueous Donor Solvents Chloro Complex Formation in Dimethyl Sulphoxide BY WLODZIMIERZ LIBUS,* ROMAN PASTEWSKI A N D TERESA SADOWSKA Department of Physical Chemistry of the Institute of Inorganic Chemistry and Technology, Technical University of Gdansk, 80-952 Gdansk, Poland Received 28th November, 1980 Equimolar mixtures of divalent transition metal perchlorates in dimethyl sulphoxide (DMSO) have been used as constant ionic media for studying the formation of complexes of metal cations with weakly coordinating anions. Isopiestic experiments indicate the same solvent activities and accordingly the same activity coefficients of the solutes in such mixtures.In the present study Co(ClO,), is used as one component of the mixture and optical absorptions due to CoC1,DMSO-, which is formed on addition of Et,NCl, are used to study chloride complexing with several transition metals. Determination of stability constants of ionic complexes is either performed under conditions of constant activity coefficients or variations in the coefficients are taken into account in the calculations. The latter approach, in which dilute ionic solutions are used,'? has the advantage of providing thermodynamic characteristics of complex formation in the absence of a supporting electrolyte and is better for discussing the solvent effect on complex formation.3 However, determination of stability constants in solutions of variable low ionic strength in non-aqueous solvents suffers from our present inadequate knowledge of the concentration dependence of the activity coefficients.On the other hand, the use of a supporting electrolyte such as alkali metal, ammonium or tetra-alkylammonium perchlorates to maintain a constant ionic strength cannot ensure that the activity coefficients remain constant when using weakly complexing ligands, since a wide concentration of the latter is needed to provide sufficient data.4 An alternative possibility for maintaining constant activity coefficients in studies of metal complexes with anions is to use equimolar mixtures of the respective perchlorates or some other metal salts involving non-coordinating anions as the reaction media.In a series of papers Z. LibuS et al. have shown that aqueous equimolar divalent transition metal perchlorates and magnesium perchlorate behave as effective constant ionic This was ascribed to the fact that the respective metal cations all occur as the hexa-aquo complexes, M(OH,)i+, probably preserving constant second layers of hydrogen-bonded water molecules. This suggested that equimolar solutions of the divalent transition metal perchlorates or tetrafluoroborates could be used as constant ionic media in non-aqueous donor solvents like acetonitrile or dimethyl sulphoxide (DMSO). Previously, we have found that the respective solutions contain the metal cations exclusively in the form of octahedral hexasolvo complexes whose mobilities and associating abilities with the non-coordinating anions 377378 COMPLEX FORMATION I N DMSO vary within relatively narrow limits as the nature of the central metal atom changes.12-15 Some results indicated constant activity coefficients of the metal perchlorates in their equimolar mixtures as we11.16 In the present work we have examined the chloro complex formation of divalent transition metal cations with the chloride anion in DMSO solution using equimolar mixtures of cobalt(r1) perchlorate with the other metal perchlorates as effectively constant ionic media. Light absorption due to the CoC1,DMSO- complex anion17 has served as a measured solution property depending on the free chloride anion concentration.EXPERIMENTAL The DMSO-solvated metal perchlorates, Mn(C10,), .6DMSO, Co(ClO,), * 6DMS0, Ni(CIO,), .6DMSO, Cu(ClO,), -4DMSO and Zn(ClO,), * 6DMS0, were obtained from the respective hydrated metal perchlorates by dissolving the latter in anhydrous DMSO, followed by boiling the solutions for several minutes and evaporating the solvent under reduced pressure.18 The crystals were filtered off and further purified by repeated crystallizations from DMSO.Reagent grade tetraethylammonium chloride was recrystallized twice from anhydrous acetonitrile and dried in uclcuo at 60 "C. Dimethyl sulphoxide, reagent grade, was dried using calcium hydride and then distilled under reduced pressure. The product was further purified by repeated fractional crystallization carried out under anhydrous conditions. The specific conductivity of the final material was 3 x The stock solutions of the metal perchlorates in DMSO were analysed for the respective metals by standard EDTA titrations.Prior to titration, weighed portions of the stock solutions were dissolved in water. In addition, the stock solutions of Co(ClO,), and Cu(ClO,), were analysed for the respective metals by electrodeposition and the stock solution of Ni(CIO,), was analysed for nickel gravimetrically using dimethylglyoxime. At least five determinations were carried out by each method and the results obtained by different methods agreed to within 0.25 % for cobalt and 0.15% for nickel. Solutions used in the final measurements were obtained by weighing from the respective stock solutions and the solvent. Their concentrations were calculated using the densities determined independently. Preparation of the solutions and other manipulations were carried out in a dry-box.Absorption spectra were measured by means of a VSU 2-P Zeiss spectrophotometer equipped with a thermostated cell compartment. The isopiestic experiments were carried out using the apparatus and techniques described in ref. (19). W1 cm-l at 25 OC. RESULTS ISOPIESTIC MOLALITIES Taking into account the similarity of the solution forms of the divalent metal perchloratesin DMS0,13j l5 as already mentioned, their osmotic and activity coefficients might be expected to display essentially the same concentration dependences. In order to check this expectation, we have attempted t o determine the isopiestic molalities of the metal perchlorates in DMSO solution.It has been found, however, that the isopiestic equilibrium in these systems at 25 "C is only attained after very long times, 1 to 2 months, or even longer for the more dilute solutions. As a result, only a few experiments have been completed and the results are listed in table 1 . Inspection of the data shows that the isopiestic molalities are the same, to a good approximation, for Co(ClO,),, Ni(ClO,), and Zn(C10,),, while they are slightly higher for Cu(CIO,),. This pattern is very similar to that observed for the aqueous perchlorates,lg indicating that in DMSO solution the pattern of the osmotic and activity coefficients is also the same, viz. the coefficients for Cu(ClO,), are slightly lower than those for the otherw. LIBUS, R.PASTEWSKI A N D T. SADOWSKA TABLE 1 .-ISOPIESTIC MOLALITIES (mol kg-l) OF THE DIVALENT TRANSITION METAL PERCHLORATES IN DMSO, AT 25 "c 3 79 Co(ClO,), Ni(C10,), Cu(ClO,), Zn(C10,), 0.1970 0.1964 0.2022 0.1976 0.2017 0.2007 0.2034 0.2003 0.2865 0.2824 -~ ~ _ _ _ _ _ _ _ ~ _ ~ _ _ _ _ _ _ _ - - ~~~~ ~ ~ ~- . ~ _ _ _ divalent metal perchlorates belonging to this group. However, apart from these second-order differences, the present results support the expectation that equimolar mixtures of the divalent transition metal perchlorates in DMSO solution should behave as effectively constant ionic media. SPECTRAL EFFECTS U N D E R L Y I N G THE APPLIED METHOD Fig. 1 shows the visible absorption spectra of cobalt(I1) in a series of solutions containing Co(CIO,), at a constant concentration of 0.175 mol dm-, and Et,NCl at concentrations varying from 0 (curve 1) to 0.0253 mol dm-, (curve 7).The absorption band observed at 535 nm for the pure Co(ClO,), solution (curve 1) is due to the 4TT,,(F) 7 lT,JP) transition within the Co(DMSO):+ octahedral complex.lR As is seen, addition of Et,NCl to the solution containing Co(ClO,), in a large excess results in the appearance of a new band with maximum at ca. 682 nm. As the concentration of Et,NCl increases further, the more subtle features of the new band become clearer and the band shape characteristic of the CoC1,DMSO- tetrahedral complex finally develops. For comparison, the spectrum of the latter complex, known from the earlier study,17 has also been indicated in fig. 1 (curve 8, right-hand scale).It may be seen that the intermediate spectra are linear combinations of those of curves 1 and 8. The conclusion might be drawn from these observations that CoC1,DMSO- is the only chloro complex formed. However, formation of an octahedral complex, e.g. CoCl(DMSO)', at a concentration comparable with that of CoC1,DMSO-, having relatively low absorptivity may well remain without a noticeable effect on the measured spectra under conditions of excess cobalt(r1) perchlorate. An indirect demonstration that an octahedral chloro complex of cobalt(r1) in fact is formed is provided by the observation that substituting Co(ClO,), for Ni(ClO,), in the equimolar mixtures containing Et,NCl at a constant concentration results in a marked increase in intensity of the 'blue' band due to the CoC1,DMSO- complex, as illustrated by curve 2 in fig.1. The effect indicates increasing concentration of the 'free' chloride anion which, in turn, may only result from the fact that a part of the total chloride content in the solution was initially bound to cobalt(I1) in the form of octahedral complexes, while nickel(r1) forms weaker chloro complexes. It appears that substitution of Co(ClO,), for Mn(ClO,), has an opposite effect on the intensity of the band due to the CoC1,DMSO- complex (curve 2' in fig. 1). This we take as an indication that the MnCl+ complex has a higher stability than CoCl+. Still more effective in depressing the intensity of the band under consideration appear to be Cu(ClO,), and Zn(C10,),, indicating the high stability of the chloro complexes of the respective cations in DMSO solution.3 80 COMPLEX FORMATION I N DMSO 500 400 309 200 100 0 400 500 600 700 A / nm FIG.1 .--Visible absorption spectrum of cobalt(i1) in the DMSO solutions of: 1, (Co(CIO,), (0.175 rnol drnP3); 2-7, Co(CIO,), (1.175 mol dm-")+ Et,NCl (concentration from 0.00596. curve 2. to 0.0253 mol drnP3, curve 7); 2', Co(CIO,), (0.101 rnol dm-:l) + Mn(CIO,), (0.077 rnol dm-:$) + Et,NCl (0.0062 mol drn-"); 2", Co(CIO,), (0. I46 mol dm-:3)+Ni(CI0,), (0.029 rnol dm-:')+ Et,NCI (0.0062 rnol dm-3); at 25 "C. The broken line 8 represents the spectrum of the CoCl; complex, right-hand scale. O U T L I N E O F T H E METHOD The applied method consists of using a series of solutions of two metal perchlorates, M(ClO,), and M'(ClO,),, of a constant total concentration ct to which a small amount of the complexing anion X is added, so that c, << ct.The equilibrium concentration [XI of the free anion is then determined as a function of the composition of the mixture either directly or from the measured equilibrium concentration of one of the complexes formed. Evaluation of the results makes use of the relation n n C nPTL [MI [XIn + C n& [M'] [XI" +[XI - c, = 0 (1) n=o 71 -0 arising from the material balance for the anion, where are the 'medium' stability constants of the complexes MX, and M'X,, respectively, and the bracketed symbols denote equilibrium concentrations of the respective species. These formulae, as well as those used to denote the chloro complexes studied, neglect possible coordination of the solvent molecules.This is in accord with the known factw. L I B U S , R . P A S T E W S K I A N D T. SADOWSKA 38 1 that the stability constants determined under conditions of a constant ionic medium may, in principle, be sums of the 'true' stability constants of a number of different solution species, all corresponding to the non-committal formula MX,."> 23 Further argument is needed to ascribe the derived stability constants to particular solution species, such as MXL: or MX,L- (L denotes the solvent molecule), or to resolve them into the stability constants of the single species. and pi', the stability constants at zero ionic strength, are related to the above stability constants by where Y , and Y/, are quotients of the activity coefficients for the respective complex-forming equilibria.The assumption is made that under the conditions stated all the activity coefficients remain constant while the relative contents of the two metal perchlorates are varied. Provided that the free anion concentration has been determined for at least 2n solutions of the above type, the best solution of a set of eqn (1) in the stability constants may be found using a computer-based program. In the zeroth approximation initially estimated values of the stability constants must be used and the free cation concentrations [MI and [M'] may be approximated by the total concentrations, provided that association with the perchlorate anion is negligible. The latter seems to be a sufficiently good approximation for the divalent transition metal perchlorates in DMSO solution, as may be seen from recently determined association ~0nstants.l~ EVALUATION OF T H E SPECTROPHOTOMETRIC RESULTS In the present study use was made of the readily determinable equilibrium concentration of the CoCI,DMSO- complex, which served as an indicator of the free chloride anion concentration.Accordingly, equimolar mixtures of Co(ClO,), with the other divalent transition metal perchlorates, containing a small amount of Et,NCl, were studied, the concentration of CoCl; (apart from its solvation) being found as [CoCl,] = c(F- E,,)/E,, where Fdenotes the measured mean molar absorption coefficient of cobalt(I1) at 680 nm and E , = 0.45 and c4 = 524 are the known molar absorption coefficients of Co2+ and CoCl;, respectively, at this wavelength.Taking into account the possible formation of three consecutive chloro complexes of either metal cation, the computer-based iterative procedure of finding the best values of the stability constants from the spectrophotometric data consisted in minimizing the function while the equilibrium concentrations in the ith solution at each step were calculated as and [COCI,], 5 ci - [CoCl,], 1 +pp[Cl], +p:"[cl]; iC11i = (-1 [CO], = ci = 1 +pp[cl], +pp[cl]; +p,"[Cl]; * The necessary values of pfo, the stability constant of the CoC1; complex in the given medium, were found in separate experiments, as described below. The program written for an ODRA 1305 computer permitted elimination of any of the five382 COMPLEX FORMATION I N DMSO complex-forming reactions represented in eqn (1) and (4), apart from that for the CoCl; complex, as well as admission at thejth iteration of any of these complex formations not so far considered, thus providing an insight into the convergence process.Using the above approach, convergent values of the stability constants of the MC1+-type complexes were obtained in few iterations for the equimolar mixtures involving either Mn(C10,), or Ni(C10,), in addition to Co(ClO,), on the assumption that pFo = Dz* = pp = 0. On the other hand, negative values of the constants, lacking physical meaning, were obtained for these systems if the respective terms were admitted into the iterative calculations. These results we take as a confirmation of the expected exclusive formation of the monochloro complexes of the respective metal cations under conditions of excess metal perchlorates.On the other hand, for the systems involving copper(I1) and zinc(II), convergence could only be obtained either by using very well estimated initial values of the stability constants or by performing the calculations in steps. In the first step, formation of only the predominant complex of the metal other than cobalt(r1) was taken into account, viz. of CuCl+ for copper(I1) and of ZnC1; for zinc(II), in addition to the two cobalt(I1) complexes, CoC1+ and CoCI;. After a number of iterations providing a preliminary convergence of the stability constants, formation of the other chloro complexes was admitted. For copper(II), admission ofeither CuClt or CuCl; resulted in equally good approximations.Of the respective results those corresponding to the formation of CuCli in addition to CuCl+ have been assumed as real, in accordance with the earlier work in which formation of these latter complexes in the dilute solutions of CuC1, in DMSO was inferred from the conductometric data and the visible absorption spectra.,O This does not mean excluding formation of the higher chloro complexes of copper(I1) under conditions of excess chloride anions or in the more concentrated solutions of CuCl, in DMSO. For the system involving Zn(C10,),, self-consistent results have been obtained on the assumption that the ZnC1; neutral complex is formed in addition to ZnCl;, while no significant improvement of the approximations was obtained upon admission of the formation of ZnC1+.Values of /?Fo in the particular media used were needed for the above calculations. One possible way was to calculate them from the previously determined stability constant at zero ionic strength, = (8.4+ 1.5) x 108.17 Magnell and Reynolds have reported (4.21 k0.58) x los assuming the absence of complexes other than CoCl;.l However, in view of the obvious uncertainty involved in selecting a proper equation for the activity coefficients, an attempt has been made to derive together with /?Fo, from the spectrophotometric data for the DMSO solutions of Co(ClO,), containing Et,NCl at smaller and variable concentrations. However, the concentrations of Et,NCl had to cover a fairly wide range in order to obtained significant changes in the mean molar absorption coefficient of cobalt(1r) within the CoCl; band.Accordingly, the Debye-Hiickelequation, log y = - z+z-A %/I/( 1 + Ba'dI), involving the ion-size parameter estimated as 8.0 A,17 was used to make allowance for the possible small variations in the activity coefficients. Values of 1.11 1 and 0.4262 were assumed for the constants A and B, respectively, valid for DMSO at 25 OC. On the assumption that only the CoCl+ and CoCl; complexes are formed under conditions of excess Co(ClO,),, the stability constants pFo and p f o were then derived from the material balances for both ColI and C1- using a separate convergence program written for an ODRA 1305 computer. In this case, the sums of the deviations squared were minimized with respect to the parameters [Co], [CoCl], [Cl], p: and pf, the latter two being constants for all the solutions in which the equilibrium concentration of CoCl; was determined experimentally.The derived values of the constants are listed inw. LIBUS, R . PASTEWSKI A N D T. SADOWSKA 383 TABLE 2.---STABILITY CONSTANTS OF THE MClf-, MC1;- AND MCI, -TYPE COMPLEXES IN EQUIMOLAR MIXTURES OF Co(C10,)2 WITH THE OTHER METAL PERCHLORATES OF TOTAL CONCENTRATION c,(mol dmP3) in DMSO, AT 25 OC Values at ct = 0 are the calculated 'thermodynamic' stability constants (see text). Co(C104)2 + Et4NClu 0.0 620f40 0.086 74.0 k 4.0 0.134 62.0 k 4.0 0.176 56.0 f 4.0 0.0 0.027 1 13.1 k 7.7 0.079 68.9 f 1.4 0.127 57.3k 1.5 0.175 53.7 f 0.9 0.225 48.0 k 0.8 0.285 45.5 & 1.5 - 0.0 0.075 65.8 k 1.5 0.125 56.7 k 3.2 0.175 53.1 2.4 - 0.0 - __ - - - 1720 348.7 f 12.4 201.1 k2.3 175.5k4.5 159.0 2 1.9 146.3 f 1.4 130.4 & 1.8 256 30.4 k 0.7 24.5 f 1 .O 23.5k 1.1 1.02 x lo5 0.1 10 65.0k0.3 (1.25k0.05)~ lo4 0.152 58.0k0.2 (0.81 k0.02) x lo4 0.195 48.0k0.1 (0.89f0.02) x lo4 0.0 0.101 66.0k0.5 - 0.151 58.0k0.5 - 0.185 50.0 k 1 .O - ~ - - (6.3 f0.8) x lo8 - (2.63 k0.35) x lo7 - ( 1 .9 2 k 0 . 3 ) ~ lo7 - (1.68 f 0.3) x lo7 - - (7.82 f 2.2) x 105 - (1.43f0.37)~ lo6 - (2.29k0.29) x lo6 - - __ (1.1+0.14)x lo7 ( 2 . 1 f 0 . 0 6 ) ~ 10" (2.9 k 0.2) x lo7 (2.1 & 0.07) x lo1' (1.8k0.8) x lo7 (3.8k0.3) x 10" a Concentration up to 15 that of CO(CIO,)~. table 2 along with the 'medium' stability constants determined in the equimolar mixtures of Co(CIO,), with the other divalent metal perchlorates.We note that /l: = (6.3 0.8) x lox is intermediate between the two values cited above, at the same time being lower than the value derived from the study of the dilute solutions of CoCl,. The agreement is not particularly important for the present purposes, since the value of a," derived here should only be considered as an empirical constant providing the best approximation of the variation in the 'medium' stability constant /lFo within the range of moderate concentrations of the ionic medium, as arising from the Debye- Hiickel equation. DISCUSSION The stability constant of the CoCl+ complex has in the present work been determined in a number of mixtures of Co(ClO,), with the other divalent transition metal perchlorates playing the role of constant ionic media.Fig. 2 shows a plot of384 COMPLEX FORMATION I N DMSO the respective values against total concentration of the mixtures. As is seen, there is a good agreement between the values obtained from independent experiments involving the different metal perchlorates. This we take as evidence of the real formation of the CoCP complex in DMSO solution, contrary to the earlier claims by some other At the same time the agreement confirms the expected constancy of the activity coefficients of the species involved in the respective complex-forming equilibrium in equimolar mixtures of the divalent transition metal perchlorates, apart from the nature of the metal cation. It also appears that the variation of the activity coefficients with the total concentration of the metal perchlorates is the same for both CoC1+ and MnCI+, as shown by the fact that the ratio of their derived stability constants remains constant when the total concentration of the metal perchlorates forming the ionic medium is varied (see table 2).( 3.0 I ( 2.5 L n - % 2.0 1.5 I I I 1 0.0 0.1 0.2 0.3 0.4 clmol dm-3 FIG. 2.-Dependence of the stability constant /I:'() of the CoCI' complex in DMSO solution on the totai concentration of the metal perchlorates, Co(ClO,),+M(CIO,),, forming the ionic medium: 0, M = Mn; 0, M = Ni; A, M = Cu; +, M = Zn; at 25 "C. The upper curve represents the Debye-Huckel plot on the assumption that ii = 8.0 A. These two observations confirm an earlier suggestion that, for some systems, the activity coefficients of solution species may be uniquely determined by the solution coordination state, being the same function of the latter for analogous complexes of different metals,24> 25 There is little doubt that the MCl+-type complexes are inner-sphere octahedral, MCl(DMSO),f, at least for Co" and Nil1, as indicated by the visible spectral effects accompanying their formation.17.2o Outer-sphere association producing the M(DMSO)g+ - Cl- ion-pairs, which are the other possible forms of occurrence of the MCI+ 'empirical' complexes, seems to be of minor importance in these systems, as also found for the weakest NiCl+ complex.2o It follows that the complex-forming reactions should be formulated as (8) and the above inferences concerning the activity coefficients relate to the quotient of the activity coefficients for these type of equilibria.Using equimolar mixtures of the divalent transition metal perchlorates as effectively constant ionic media in studies on complex formation of the respective metal cations M(DMSO);+ +C1- e MCl(DMS0); + DMSOw. LIBUS, R. PASTEWSKI AND T. SADOWSKA 385 provides a method of devising conditions where the activity coefficients remain constant. This approach to controlling activity coefficients seems to be superior to the more traditional one in which constant formal ionic strength is adjusted while the composition and the coordination state of the ionic medium are varied considerably. However, changing the concentration of the reactants, necessary in any equilibrium study, only in rare cases may be accomplished without changing the solution coordination state, as was achieved in the present work.Another example is provided by the isoconducting mixtures of ZnC1, and CoC1, in acetonitrile, containing complex electrolytes of the type M(AN)i+. 2MC1,AN- at constant concentrations, in addition to the neutral coordination forms MCl,(AN), believed to be of minor importance in determining the activity coefficients (AN denotes the acetonitrile There are reasons to believe that equimolar mixtures of some trivalent metal perchlorates, or some other trivalent metal salts involving non-coordinating anions, in aprotic donor solvents may also behave as effectively constant ionic media owing to the correspondence of their coordination statesz7 The expected agreement of the quotient of the activity coefficients for all the complex-forming reactions analogous to reaction (8) may be made use of for estimation of the respective stability constants at zero ionic strength from the ‘medium’ ones.In order to obtain a standard of the variation of the quotient of the activity coefficients for the MC1+-type complexes, the data for CoCl+ are used. The upper part of fig. 2 shows that the stability constant of the CoCl+ complex shows a variation with the concentration of the metal perchlorates forming the ionic medium which is approximately consistent with that arising from the Debye-Huckel equation with the ion-size parameter of 8.0 A (corresponding to Bao = 3.41). The latter value has been estimated from the molecular model of the complex cations involved in reaction (8) and the ionic radius of the unsolvated anion.The indicated values of the stability constant at zero ionic strength of this complex are those found previously from spectrophotometric and conductometric studies, respectively, of very dilute solutions of CoCl, in DMS0.I7 The straight line is that of the theoretical Debye-Huckel slope providing the best fit to the presently obtained ‘medium’ stability constants. The observed reasonable agreement of its intercept with the independently determined values of the stability constant at zero ionic strength indicates the reasonable validity of the assumed equation for the activity coefficients. Assuming its validity for the other complex-forming reactions similar to reaction (8), the respective stability constants at zero ionic strength have been calculated from the ‘ medium ’ ones determined in the ionic media of the metal perchlorates.The stability constants derived in the present work relate mainly to the MCP-type complexes (apart from their solvation) preferentially formed under conditions of excess metal perchlorates. Exceptional in this respect are the systems involving zinc(I1). The ZnC1: and ZnCl; complexes formed in them, most probably both tetrahedral, are so much stronger that they depress formation of the ZnC1+ complex, making derivation of its stability constant impossible. Qualitatively, this is in accord with the earlier findings of Ahrland and Bjork, who studied chloro complex formation of zinc(rr) using the NH,ClO, + NH,Cl solutions as reaction media of constant formal ionic strength.lqj For the other systems, the MCl+ complexes have been well- characterized and their real formation seems to be beyond doubt.It seems most probable that they all are of the MCl(DMS0): inner-sphere type. For M = Mn, confirmation of this solution form of the complex is provided by the agreement, discussed above, of the quotients of the activity coefficients of the complex-forming reactions for MnCl’ and CoCl+ Note that it is formation of the MC1+-type complexes that has the largest effect on the properties of the most dilute solutions of the respective386 COMPLEX FORMATION I N DMSO metal chlorides in DMSO, as well as in the other donor solvents. As a result, the stability order of these complexes is important for an understanding of the observed differentiation in behaviour of the divalent transition metal chlorides in their dilute solutions in these solvents.Fig. 3 shows plots of log PI for the MCl+-type complexes in DMSO solution against position of the metal within the Mn-Zn series. The plotted values relate to the same total concentration of 0.1 mol dmP3 of the metal perchlorates forming the ionic media and have been found by interpolation from the data listed in table 2. Also indicated are the thermodynamic stability constants calculated from the 'medium' ones as discussed above. The indicated values for ZnC1+ have been calculated from the data of Ahrland and Bjork4 using the Debye-Hiickel aquation for the activity coefficients. For obvious reasons they are not strictly comparable with our data determined in different media. 5 4 - c2 w 2 3 2 1 0- 0- 0 I \ I I I 1 I 1 I Mn Fe Co N i Cu Zn FIG.3.-Variation in the stability of the MCF-type complexes in DMSO solution within the Mn to Zn transition metal series: lower curve, in 0.1 mol dm-3 metal perchlorates; upper curve, in pure DMSO; points for Zn2+ recalculated from the data of Ahrland and Bjork.4 Inspection of fig. 3 shows that the stabilities of the MC1+-type complexes in DMSO solution do not follow the regularity known as the Irving-Williams series,28 but rather an 'inverted' Irving-Williams series. This is in accord with our earlier observations concerning some others systemsg+ 29+ 3o and seems to be typical of complexes formed by metal cations with weakly coordinating anions in strongly coordinating donor solvents like DMSO.To our knowledge, regularities of the present type are not generally known and the Irving-Williams series is believed to be obeyed, as it is in aqueous systems.:51 We postpone a more detailed discussion of this question to a later paper in which the thermodynamic characteristics of the complexes will be completed by the enthalpy data. ' K. R. Magnell and W. L. Reynolds, Znorg. Chim. Acta, 1972, 6, 571. W. LibuS, L. Frqczyk and B. Chachulski, Pol. J . Chem., 1978, 52, 493. W. LibuS, Muter. Sci., 1979, 5, 85. S. Ahrland and N. 0. Bjork, Acta Chem. Scand., Ser. A , 1976, 30, 257; 265. L. Sestilli and C. Furlani, Inorg. Nucl. Chem., 1970, 32, 1997. Z. LibuS, J . Phj,s. Chem., 1970, 74, 947.Z. LibuS, Znorg. Chem., 1973, 12, 2972. Z. LibuS and H . Tialowska, J . Solution Chem., 1975, 4, 101 1. j S. Ahrland, N. 0. Bjork and R. Portanova, Acta Chem. Scand., Ser. A . 1976, 30, 270. l o Z. LibuS and G. Kowalewska. Pol. J. Chem., 1978, 52. 709. ' I Z. LibuS and W. Maciejewski, Roc;. Chem., 1976, 50, 166. '' W. LibuS and H. Strzelecki, Electrochim. Actu, 1970, 15, 703; 1971, 16, 1749.w. LIBUS, R . PASTEWSKI A N D T. SADOWSKA 387 ' : I W. LibuS and M. Pilarczyk, Bull. Acad. Pol. Sci., Skr, Sci. Chim., 1972, 20, 539; 1973, 21, 773. I J W. LibuS and B. Chachulski, J . Solution Chem., 1980, 5, 355. W. LibuS, W. Grzybkowski and R. Pastewski, J . Chem. Soc., Furaday Trans. I , 1981, 77, 147. Ifi W. LibuS and H. Strzelecki, Electrochim. Actu, 1972, 17, 577. W. LibuS, M. Pilarczyk, R. Pastewski and T. Szuchnicka, Electrochim. Acta, in press. H. L. Schlifer and W. Schaffernicht, Angew. Chem., 1960, 72, 618. '!I Z. LibuS and T. Sadowska, J . Phys. Chem., 1969, 73, 3229. ?'I W. LibuS, M. Pilarczyk and T. Szuchnicka, Electrochim. Actu, 1975, 20, 831; 1980, 25, 1033. .'I V. Gutmann, Coord. Chem. Rer., 1967, 2, 239. ?? F. J. Rossotti and H. Rossotti, The Determinution of Stability Constants (McGraw-Hill, New York, 1961), p. 3. W. LibuS, Rocz. Chem., 1976, 50, 1813. ?-I W. LibuS, XIIIrh Int. Conf: Coordination Chemistry (Zakopane, Krakow, 1970), Section Lectures, p. 241. 2,i W. Libus, T. Sadowska and Z. LibuS, J . Solution Chem., 1980, 9, 341. ''Ii W. LibuS. D. Puchalska and T. Szuchnicka, J . Phys. Chem., 1968, 72, 2075. " i W. LibuS and M. Pilarczyk, Bull. Acud. Pol. Sci., Ser. Sci. Chim., 1974, 22, 717. :'" H. Irving and R. J. P. Williams, J. Chem. Soc., 1953, 3192. 2B W. Libus, B. Chachulski, L. Frgczyk and H. Strzelecki, Rocz. Chem., 1975, 49, 19. .I1' W. LibuS, Mater. Sci., 1979, 5, 135. .'I S. Ahrland, The Chemistry of'Non-Aqueous Solvents, ed. J. J. Lagowski (Academic Press, New York, 1978), vol. VA, p. 45. (PAPER O/ 1836)

ISSN:0300-9599    

DOI:10.1039/F19827800377

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chcrvz. Sor., Farday Truns. I , 1982, 78, 389-401 Ultrasonic Relaxation Studies of Concentrated Surfactant Solutions and Liquid Crystals B Y GORDON J. T. rrIDDY* Unilever Research, Port Sunlight Laboratory, Quarry Road East, Bebington, Wirral, Merseyside L63 3JW A N D MICHAEL F. WALSH A N D EVAN WYN-JONES Department of Chemistry and Applied Chemistry, University of Salford, The Crescent, Salford M5 4WT Receiced 5th December, 1980 Ultrasonic relaxation studies have been carried out on non-ionic, anionic and cationic surfactants in the micellar phase and in hexagonal and lamellar liquid crystal phases, over the frequency range 1 - 105 MHL. For all the surfactants, a relaxation process with similar absorption characteristics occurs i n both the concentratedmicellarphaseand theadjacent liquidcrystal phase.For most surfactantsstudied quantitatively, the relaxation frequency decreases with surfactant concentration and the relaxation process is attributed to the perturbation of a dynamic equilibrium between surfactant-' bound' and ' free' water molecules, occurring on a time scale Qf lo-' to lopR s. A model to describe this 'bound'/'free' water-exchange process has been developed which allows the association and dissociation rate constants for the process to be determined. I t also gives an estimate of the hydration numbers of the surfactants. For caesium oleate solution the relaxation rate increases with surfactant concentration. This is tentatively attributed to exchange of free and bound counter-ions; values of the rate constants for the process are derived.The structures of most lyotropic liquid crystalline (1.c.) phases formed by anionic, cationic and non-ionic surfactants in water are now well characterised.l Aqueous mixtures of surfactants can exist as one of a number of different forms, these being the isotropic micellar solution and the lamellar, hexagonal, reversed hexagonal and various cubic 1.c. phases. The transition between isotropic micellar solution and 1.c. phase, or between different 1.c. phases, is usually first order and can arise as a result of concentration and/or temperature changes. However, the two-phase regions are often narrow for two-component systems. The structure of the surfactant aggregates in the different phases is very different. In dilute solutions, micelles are thought to be roughly spherical ; the lamellar liquid crystalline phase comprises a bilayer structure where the surfactant lamellae are separated by water layers; the hexagonal phase consists of long cylindrical micelles packed in an hexagonal array in a water continuum; the reverse hexagonal phase contains hexagonally packed water cylinders.There are at least three and possibly four or more types of cubic phase but their structures are not well established. A particular surfactant may form several of these liquid crystalline phases and, indeed, surfactants which do not form micelles can give liquid crystalline phases at sufficiently high concentrations. Lyotropic liquid crystals often exhibit high viscosities, and properties such as diffusion are anisotropic.Until now kinetic studies in these systems have been confined mostly to dilute micellar solutions. As a result of extensive chemical relaxation measurements, the kinetics of micelle formation is well understood.5-8 Of the various chemical 389390 ULTRASONIC RELAXATION OF SURFACTANTS relaxation methods used for these studies the ultrasonic method covering the frequency range 1-100 MHz has been used extensively to study the monomer/micelle exchange process in surfactants with C,-C,, hydrocarbon chains. The relaxation associated with this exchange (process 1) can be described by a single time constant and is observed only at surfactant concentrations in excess of the critical micelle concentration (c.m.c). The amplitude of this relaxation initially increases with increasing surfactant concentration, reaching a maximum at a concentration which is CQ.2-4 times the c.m.c. Further increases in surfactant concentration result in a steady decrease of the relaxation amplitude until it becomes almost undetectable. At this stage a new relaxation (process 2), with a lower relaxation frequency, is observed in many ~urfactants.~ This new relaxation has an amplitude which increases with increasing surfactant concentration and it is also observed in the liquid crystalline phases of all the surfactants studied. In a preliminary communication9 we assigned process 2 to the perturbation of the equilibrium between free water molecules in the solution and those 'bound' to the surfactant headgroups. In this paper we report further details concerning the ultrasonic relaxation process of concentrated surfactant systems.A kinetic model to describe this relaxation process is proposed and the relaxation data are considered in relation to the model. EXPERIMENTAL The ultrasonic absorption and velocity measurements were taken with the Eggers resonance methodlo over the frequency range 1-20 MHz and a conventional pulse method" over the range 15-105 MHz. MATERIALS Caesium perfluoro-octanoate (CsPFO) was prepared by neutralising perfluoro-octanoic acid (Koch Light, 99% min) with caesium carbonate (Koch-Light, 99% min) and dried in a warm oven (ca. 40 "C). The dry product was then recrystallised from a butanol+petroleum ether mixture and stored over phosphorous pentoxide under vacuum. Caesium oleate (CsOl) was prepared by neutralising oleic acid (B.D.H.AnalaR) with caesium carbonate and also dried in a warm oven. The CsOl was recrystallised from ethyl acetate and stored over P,O,. Triton X-100 (iC, 0 E9.,) and Triton X-114 (C, 0 E,) were obtained from B.D.H. and used without further purification. Tetraethylene glycol dodecyl ether (C,,E,) and pentaethylene glycol dodecyl ether (C,,E,) were obtained from Nikko Chemicals Ltd, Tokyo and used as supplied. Sodium octyl sulphate (SOS), sodium decyl sulphate (SDeS), sodium dodecyl sulphate (SDS), decyltrimethylammonium bromide (DTAB) and hexadecyltrimethylammonium bromide (CTAB) were obtained from Cambrian Chemicals Ltd and recrystallised from absolute ethanol. Aerosol OT was obtained from the Fisher Corporation, U.S.A.Dodecyldimethylammoniopropane sulphonate (C,,DAPS) was obtained from Unilever Research Port Sunlight Laboratory and was recrystallised from a hexanol + 60-80 petroleum ether mixture. Mono-caprylin was also obtained from Unilever Research and was prepared by the method described by Peel.', Mono-olein was obtained from Eastman Organic and used without further purification. ,H,O was obtained from Fluorochem Ltd (> 99% min). PREPARATION OF SAMPLES Micellar solutions were prepared by dissolving the surfactant in water. For liquid crystal samples the required amounts of surfactant and water were mixed by heating above the l.c./isotropic solution transition temperature. With very concentrated surfactant systems andG. J . T. TIDDY, M. F. WALSH AND E. WYN-JONES 39 1 where the I.c.melting temperature was high (> 60 "C), the surfactant and water mixture was heated in a pressure cooker for ca. 20 min at 1 15 "C in order to produce homogeneous samples. Samples were then transferred by pouring directly into the ultrasonic cell. The presence of the required 1.c. phase was verified by optical microscopy. RESULTS AND DISCUSSION The measurements were carried out on samples in the micellar phase, lamellar phase and/or hexagonal phase; a list of systems studied is given in table 1. For some systems we were unable to make measurements over the full frequency range because the Eggers resonance method (1-20 MHz) could not be used for samples with high viscosity. (These are indicated by footnote in table 1) monomer-micelle exchange (process I ) occurred at a similar frequency to process 2.In both these cases we were unable to make an accurate estimation of the amplitude and frequency for the second relaxation process. Previously9 we showed that for samples close to a phase boundary an additional relaxation occurred (process 3) which could be separated from process 2. For quantitative analysis we have chosen samples where the relaxation spectrum could be measured over the full range 1-105 MHz and where the contributions from processes 1 and 3 were negligible. In all these cases the relaxation observed could be described in table 1.) In other systems, (footnote TABLE 1 .-SURFACTANTS STUDIED surfactant structure results phases studieda (frequency /amplitude) caesium perfluoro- (CsPFO), CiF,,C02Cs octanoate Triton X- 100 Tetraethylene glycol C,2E,d pentaethylene glycol C,,ESd sodium octyl sulphate (SOS), (C, S0,Na) sodium decyl sulphate (SDeS), (Cl0 SO,$Na) sodium dodecyl sulphate(SDS), (C12 S0,Na) caesium oleate Triton X-114 (X- 1 14), CH 0 ETd (X-100).iCH 125E9.2d dodecyl ether dodecyl ether (CSO), c,i = C0,CS decyltrimethyl- (DTAB), c,, N(CH,),Br ammonium bromide ammonium bromide ammoniopropane sulphonate hexadecyltrimethyl- (CTAB), C,,N(CH,),Br dodecyldimethyl- (DAPS), C12N(CH3)2(CH2)3S03 Aerosol OT di-(iC, C02)CH2CHS03Na mono-octanoin C , monoglyceride mono-olein C , monoglyceride L,, LAM fig. 1 high L,, LAM fig. 1 HEX fig. 1 LAM b L,, HEX, LAM fig. 1 C C C fig. 1 Ll C Ll not obs. Ll fig. I LAM LAM b LAM b b high high highb high lowC lowC lowC low l0wC not obs.medium mediumb lowb lowb a L, = isotropic micellar solution, LAM = lamellar liquid crystal phase, HEX = normal hexagonal Not En represents an ethylene liquid crystal phase. analysed quantitatively because two relaxation processes overlap (see text). oxide group (OCH,CH,),. Not analysed quantitatively because of limited frequency range (see text).392 ULTRASONIC RELAXATION O F SURFACTANTS either by a single time constant or with a very narrow distribution of relaxation times centred around a mean value. In the earlier communicationsY we reported that for a given surfactant relaxation process 2 varied continuously throughout all the phases regardless of phase structure, with an amplitude which increased with increasing surfactant concentration and decreased with increasing temperature.The variation of relaxation frequency ( f c ) with concentration for a given surfactant was as large as the change caused by altering surfactant chemical structure. A schematic diagram summarising the relaxation frequency data for the compounds studied is shown in fig. 1. For a variety of different [ C& E,’((X- 11 4 ) 1 CSPFO I H 20 I pmqq types of surfactant in different phases the relaxation frequencies lie in the range 2-1 7 MHz. For compounds where a quantitative analysis of the data was not possible, the relaxation frequencies were of a similar magnitude. Ultrasonic relaxation is observed when a chemical exchange process occurs on a time scale comparable to the inverse frequency of the sound waves. In the present systems such a process could arise from several different sources, these being: (i) alkyl chain conformational exchange; (ii) a process associated with micelle shapes or inter-micelle interactions, such as an equilibrium between spherical and rod micelles or the formation of ordered micelle domains; (iii) free/bound counter-ion exchange; (iv) exchange between different headgroup conformations; (v) free/bound water exchange.Mechanism (i) can be dismissed because the process is independent of chain length and is the same for hydrocarbon or fluorocarbon chains. This process might be expected to have a relaxation frequency of > 100 MHz, as for paraffins and p01ymers.~~ Since the relaxation is independent of surfactant phase structure and is similar for ionic and non-ionic surfactants with very different headgroups, mechanisms (ii) and (iv) can be eliminated also.As we concluded previou~ly,~ the only molecular process that is common to all these surfactants in all the different phases is mechanism (v). Where the water is bound to counter-ions these could be involved also and so mechanism (iii) could be involved indirectly.G. J. T. TIDDY, M. F. WALSH AND E. WYN-JONES I I I 1 1 393 FIG. 2.-(a) Relaxation frequency for water exchange in normal ( I ) and heavy ( 8 ) water. (A) CsPFO, (B) X-114, (C) C,,E,. ( b ) Relaxation data and theoretical wave for 40% CsPFO in zHH,O. Data fitted to: alf2 = A / [ 1 + (f/L.)'] + B, where GC is the sound absorption coefficient,fis the measurement frequency,f, is the relaxation frequency and A and B are constants.A comparison of relaxation frequencies in normal and heavy water for samples having the same surfactant/water molar ratio shows that the relaxation frequency is lower by a factor of ca. 25% in heavy water [fig. 2(a)]. A typical experimental curve and fit is shown in fig. 2(b). Similar results were obtained for SOS and DTAB samples, but for the reasons given above quantitative analysis of these data was not carried out. This again is good evidence that the relaxation process involves water. The change is of the same order of magnitude as the difference in viscosities and self-diffusion coefficients between normal and heavy water. l 4394 80 60 40 20 ULTRASONIC RELAXATION OF SURFACTANTS - - - - 180 r 160 120 I 6 8 100 ' 5 IZ [no.of (OCH,CH,) groups] FIG. 3.-Variation in a / f for a series of C,2H2,(OCH2CH,),0H/H20 mixtures (0.5 mol dm-3). 0, 15 and 0, 45 MHz. m Y) 53 1 n b --. i I I I I 1 I ] 0.2 0.4 0.6 0.8 1.0 1.2 1.3 concentration/mol dm-3 FIG. 4.-Variation of reciprocal relaxation times ( l / ~ ) with concentration. A, CsOl; 0, C,,E,; x , X-114; 0, X-100; ., CSPFO. Further evidence in favour of mechanism (v) is obtained from a consideration of sound absorption as a function of surfactant structure. If we consider the series of surfactants C,,H,,(OCH,CH,),OH, we expect the sound absorption parameter ( a l p ) to increase with the number of binding sites on the headgroup and consequently with n for a given concentration of surfactant. Fig. 3 shows that this is the case. A number of authors have studied the kinetics of solute/solvent interactions in surfactant solutions and biological systems. 15-,0 These investigations have been mainly confined to n.m.r.15-lS or dielectric techniqueslg and relaxation times of s have been proposed.A review of these and other results has recently been given by Packer.20 The relaxation process observed here is obviously well within the frequency toG. J. T. TIDDY, M. F. WALSH AND E. WYN-JONES 395 FIG. 9 000 8000 7000 6000 5000 4000 3000 200 0 1000 0 0 . 2 0.L 0.6 0.8 1.0 1.2 i 4' / X ' - X-+ 0.2 0 4 0.6 0.8 1.0 1.2 concentration/rnol dm-3 5.-Variation of relaxation amplitude ( A ) with concentration. A, CsOl; 0, CI2E5; x , X-114; 0, X-100; m, CSPFO. i 12000 10000 i- 8 0 0 0 1 E 2 5 0 0 0 1 ii:@ 200 0 0.2 0.4 0.6 0.8 1.0 1 2000 I I 0 0.2 0.4 0.5 0.8 1.0 1.2 1.3 concentration/rnol dm-3 FIG.6.-Variation of relaxation amplitude ( A ) with concentration for samples in water and heavy water. A, x-i i 4 / ~ , 0 ; A, x-i M/H,O; 0 , CSPFO/D,O; 0, C~PFO/H,O; 0, c,,E,/D,o; m, c,,E,/H,o. range generally attributed to water exchange and is in good agreement with the work of others. The variation of reciprocal relaxation time (1 /z) with surfactant concentration for the systems studied quantitatively is shown in fig. 4 and the relaxation amplitudes ( A ) are given in fig. 5. The relaxation amplitude increases monotonically with surfactant concentration, while 1 /z can increase or decrease. Note that the relaxation amplitudes are larger in heavy water (fig. 6).396 ULTRASONIC RELAXATION OF SURFACTANTS QUANTITATIVE DESCRIPTION OF RELAXATION PROCESS In order to develop a quantitative model for the relaxation process we have to consider what is meant by ‘bound’ water and how the exchange between free and ‘bound ’ water is perturbed by the ultrasonic wave.The binding of water to surfactants has been studied by many workers3*,21-26 who have attempted to measure hydration numbers, i.e. the number of water molecules bound to the surfactant. The number appears to vary with the measurement technique, but generally ionic surfactants appear to bind less water than polyoxyethylene non-ionic surfactants. For example, SDS and C,&, have hydration numbers of ra. 8 and ca. 20, r e s p e c t i ~ e l y . ~ * ~ ~ This unremarkable observation no doubt has to do with the relative sizes of the headgroups.‘ Bound’ water really refers to water that exists in a different state from normal water due to the presence of the surfactant/water interface. It may not always involve a favourable enthalpy change. We can envisage three types of bound water: (i) water attached by specific hydrogen bonds to surfactant headgroups; (ii) water in the hydration sphere of counter-ions; (iii) water adjacent to the surface, which is not able to participate fully in normal hydrogen bonding to neighbouring water molecules because of conformation restraints imposed by the surface. An example of this type is water adjacent to hydrocarbon chains. For the first two types of water there is an obvious enthalpic contribution to binding.The third group can hardly be called ‘bound’ and the entropic change is likely to be energetically unfavourable. A better term would be ‘disturbed’ or ‘disrupted’ water. Since different techniques will sense the different types of water in different ways it is hardly surprising that the hydration number depends on the measurement technique. There is no reason to suppose that the three groups of water will all exchange with free water at the same rate. Also we have to consider the possibility that the observed relaxation could arise from exchange between the different varieties of bound water. In order to obtain an ultrasonic relaxation the sound wave must perturb an equilibrium between states that involves a molar volume and/or enthalpy change. The reciprocal relaxation time (1 /z) is related to the kinetics of the process with the exact form of the expression depending on the number of different species involved.The absorption is related to the volume change (AV) and enthalpy change (AH) by the where p,,, is the maximum absorption per wavelength, p is the density of the solution, u is the sound velocity in the solution, T is the absolute temperature, R is the gas constant, C , is the specific heat at constant pressure, E is the expansion coefficient of the solution and r expresses the concentrations of the species in equilibrium and is given by r = ( p L;l:)-l (2) where c j are the stoichiometric coefficients and i;i are the concentrations of the species in equilibrium. To apply these expressions in the present system we require a description of the equilibria concerned.It is usual to Lmploy equations derived for fast reaction studies of species in dynamic equilibrium under the assumption of a small p e r t ~ r b a t i o n . ~ ~ The theory allows easy correlation of relaxation rates (1 /T) with changes in concentration of the reactingG. J. T. TIDDY, M. F. W A L S H AND E. WYN-JONES 397 species, but estimations of the absorption are fraught with difficulty unless volume and heat changes for the process are known. In our case it is appropriate to first try the simple ‘two state’ process. We have a series of equilibria represented as follows; k, sw+ w e sw, . k-2 . (3) . k, . sw,-,+w e sw, k-n where S and W refer to surfactant and free water and SW, is the surfactant/water complex with n molecules of bound water.Each equilibrium has associated with it individual forward and backward rate constants. Obviously we need to simplify this multi-equilibria description. We assume that all bound sites are equivalent and that the rate constants are independent of the number of bound water molecules. In this case it is easy to show that:2i 1 1 = k , [ S ] + k - , . z (4) The results for caesium oleate could be described by this type of equation with appropriate values of k , and k-n. However, this is the only case where eqn (4) could be applied, since in all the other systems shown in fig. 4, l / t decreases as surfactant concentration increases. An alternative approach is to consider a model where we have a fixed number of binding sites per surfactant molecule ( p ) and only a fraction are occupied at a given concentration. Then the equilibrium is given by: k f kb w + s, e s, where S, and S, are unoccupied and occupied binding sites and W is free water.The concentration of bound water is equal to [S,]. Using the same derivation as for eqn (4), this leads to where Wc is the total water concentration and q is the number of bound water molecules per surfactant. Obviously if q > p / 2 , then 1 / T will decrease as [S] increases. This expression successfully accounts for the behaviour of 1 /z with surfactant concentration for all the surfactant systems except caesium oleate. For given trial values of p and k,/k,, values of q can be calculated from the equilibrium condition: The first bracket on the r.h.s.of eqn (6) can be evaluated and a plot against 1 /z should yield values of k f and k , from the slope and intercept. These values should be consistent with assumed ratio k,/kb and k , must be positive. A fit of data with reasonable values of the parameters was possible and the values obtained from an iterative procedure (using a computer programme) are given in table 2. Obviously the actual number of binding sites occupied (q) depends on concentration. Table 2 gives398 ULTRASONIC RELAXATION OF SURFACTANTS TABLE 2.-KINETIC AND EQUILIBRIA PARAMETERS ESTIMATED FROM THE MODEL k,/ 1 0-6 range of dm3 (kflkb) K/dm3 p for p value range of q value mol-l k,/1OP6 /dm3 surfactant solvent mol-l best fits quoted q quoted s-l s-l mol-l x-1 loa H 2 0 0.9 10-25 15 7.8-15 14.0 2.2 9.0 0.3 X-114 H 2 0 0.8 10-20 15 10-17 13.8 1.8 2.6 0.7 C I A H 2 0 0.1 8-15 15 6-10.3 10.3 1.3 14.6 0.09 D20 0.1 8-15 15 6-10.3 10.3 0.8 17 0.05 D20 3 13-18 16 13.4-18 15.4 2.7 1 2.7 C,,DAPS H 2 0 0.8 8-20 12 7.3-17.1 11.0 3.9 5.7 0.7 CsPFO H 2 0 3 13-18 16 13.4-18 15.4 3.3 1 3.3 a Only 2 concentrations were measured in the hexagonal phase due to the difficulties of handling and unreliable data.the calculated number at a surfactant concentration of 1 mol dmP3 and the range of p values over which a reasonable fit could be obtained. Because a range of values for p and q could be used to fit the data, we were not able to obtain very accurate estimates of headgroup hydration. The values quoted are in reasonable agreement with those measured at low concentrations by other techniques,26 although the latter are expected to be higher.Thus non-ionic surfactants have rather large hydration numbers, which increase with increasing ethylene oxide content. The zwitterionic surfactant would be expected to have about the sum of the hydration water of sulphonate and alkyl ammonium groups. This does appear to be the case since small anionic and cationic groups bind ca. 3-5 molecules of water each (excluding the counter-ion).26 The value for caesium perfluoro-octanoate is surprisingly large in view of the hydration number of sodium octanoate (8.5-8.9, including counter-ion hydration).28 It seems highly improbable that the additional ' bound ' water is associated with the different counter-ion, since caesium is thought to be less strongly hydrated than sodium.Also, the additional water is unlikely to be bound to the headgroup because the fluorocarbon chain is electron-withdrawing rather than (weakly) electron-donating as is the hydrocarbon group. Thus we tentatively conclude that the larger amount of ' bound' water with caesium perfluoro-octanoate arises from a contribution due to water adjacent to fluorocarbon chains. This water is not hydrogen-bonded to the chains, but has its number of available conformations restricted by the presence of the chains. Further measurements by more traditional techniques (such as self-diffusion coefficients) are required to verify this conclusion. With caesium oleate we can only describe the observed increase of l/t with concentration by using p z 46 and q z 3.This is physically unreasonable since the origin of the 43 empty 'sites' for water molecules is obscure and such a requirement does not exist with the other compounds in table 2. The model described above is also shown to be consistent with the amplitude parameter. From eqn (1) a plot of pmax2RT/pu2 against r should be linear and pass through the origin. The slope of the line will be given byG. J. T. TIDDY, M. F. WALSH AND E. WYN-JONES 399 8000 -i 6000 E 30 2 N m v. - . 2 4000 h N 3 Q . X 2 3 2000 0 0 . 5 1 .o 1.5 2.0 r/mol dm-j FIG. 7.-Variation of pmax/pu2 with r. A, CsPFO/H,O; 0, X-l14/H,O. As shown in fig. 7, this behaviour is observed and the value of n(AV-&/p C, AM2 is much greater than that predicted from A V changes alone.Therefore it is possible that both AV and AH are contributing to the relaxation process. Consequently the signs of these two terms must be different in such a way that they are additive in the whole bracket. However, the individual AV and AH cannot be estimated. The difference between caesium oleate and most of the other surfactants is the presence of the counter-ion. A similar qualitative increase in 1 /z with concentration could be deduced from the measurements on sodium octyl sulphate, its decyl and dodecyl homologues (where monomer-micelle exchange is also seen) and for Aerosol OT (where the frequency range used was limited). Thus we are entitled to enquire if the difference is associated with the counter-ion. At present this appears to be the best explanation for the result.The counter-ion will have a different hydration state when it is close to the interface from that in bulk solution. Thus we have an exchange process of the type: k€ X(Hzo)d + sd X(H20)a (8) kb where a and d refer to surfactant (S) and counter-ions (X) in the ‘bound’ and ‘free’ states. It is generally accepted that the fraction of counter-ions dissociated from surfactant aggregates (sphere, rods or bilayers) is ca. 30% or less and is fairly invariant with surfactant concentration and added electrolyte (the ion condensation hypothe~is).~~-~l In addition, while the process is accompanied by a change in hydration, the equilibrium monitored is mainly determined by the electrostatic interactions. The relaxation equation for the process will be of a similar form to eqn (4) above, since the concentration of bound ions will vary with surfactant concent- ration.If a is the degree ofcounter-ion dissociation from the micelle then the relaxation rate is given by:27 (9) 1 /Z = kf([Xd] + [Sd]) + k b = 2a kf [S] + k b .400 ULTRASONIC RELAXATION OF SURFACTANTS From the data in fig. 1 we estimate values for ak, and k , of 37 x lo6 dm3 mol-l s-’ and 4-7 x lo6 s-l. Using the equilibrium relationship derived assuming that the degree of micellar charge dissociation (a) is independent of overall surfactant concentration (ion condensation hypothesis)29 31 then we have: Substituting the values of ak, and k , above we calculate a = 0.12-0.23 and hence k , e 185 x lo6 dm3 mol-I s-l. The value for a is in reasonable agreement with other estimates26 after allowing for the crude nature of the approximations involved.In particular, k f is highly unlikely to be constant. It is very improbable that any barrier exists to prevent ions entering the ‘bound’ state since there will be an attraction due to the electrostatic interactions. Thus the forward rate will be diffusion controlled and hence the smaller the average separation between ions and the surfactant micelles (i.e. at higher surfactant concentrations) the higher the value of k,. Obviously, in order to keep a approximately constant k , must increase also. The exact form will depend on the nature of the ion distribution around the micelle and how close the ion has to approach the micelle surface to count as a ‘bound’ ion. This may be within 1-2 A of the micelle surface, the same order of magnitude as is required to distort the hydration sheath for the observation of n.m.r.quadrupole ~ p l i t t i n g s ~ l , ~ ~ . Assuming that k , is diffusion controlled, we can calculate its order of magnitude from the average separation of the micelles and the ion self-diffusion coefficient. (This is ca. lo2 larger than the micelle self-diffusion coefficient.) Assuming the micelles are spherical with a radius of 25 A then for a solution of ca. 0.5 mol dm-3 the average micelle separation is ca. 80° A. With a counter-ion diffusion coefficient of ca. 7 x m2 s-l the ions will collide with micelles on a time scale of ca. 1 x lop9 s. This is certainly of the correct magnitude to give the observed relaxation times.Thus for caesium oleate the relaxation process is probably due to the bound/free counter-ion exchange. Experiments to verify this proposal by making measurements as a function of added electrolyte and with a series of surfactants and counter-ions are in progress. We have attempted to measure the effect of added salt on the counter-ion exchange relaxation time for micellar solution of CTAB plus NaBr. CTAB shows a high water solubility and in addition micellar solutions of CTAB + water have no ultrasonic relaxation. Whilst a relaxation is observed in solutions containing added NaBr the effect is small, presumably due to the very low hydration number of the Br- ion (ca. 1). A much larger effect was observed for the anionic surfactant sodium dodecyl benzene sulphonate in the presence of NaCl or MgC1,.Here we found that the relaxation amplitudes for the surfactant in equimolar NaCl or MgCl, were approxi- mately in the ratio of the hydration numbers of the two cationic species, i.e. Na(6), Mg( 15). Thus we have qualitative confirmation that counter-ion exchange is involved in the ultrasonic relaxation process. Work is in progress to obtain further data to check the quantitative relationships. For a quantitative interpretation of trends in relaxation amplitudes with surfactant concentration, the volume and enthalpy charges accompanying the exchange process must be known. However, even without this particular knowledge a qualitative account is still possible, by considering the relative amounts of the species in equilibrium [eqn (2)].As is expected for an exchange process of this type the relaxation amplitude increases with surfactant concentration (fig. 4). Similarly it is anticipated that those surfactants with large hydration values will reflect large relaxation amplitudes. Our results show this to be the case.G. J. T. TIDDY, M. F. WALSH AND E. WYN-JONES 40 1 These results show that the models proposed here and their interpretation are consistent with processes involving the dynamic exchange of water molecules. V. Luzzati, in Biological Membranes, ed. D. Chapman (Academic Press, London and New York, 1968), chap. 3. P. A. Winsor, Chem. Rev., 1968, 68, 1. P. Ekwall, in Admnces in Liquid Crystals, ed. G. H. Brown (Academic Press, San Francisco and London, 1971), vol. 1, chap.1. G. J. T. Tiddy, Phys. Rep., 1980, 57, 1. P. J. Sam, J. Rassing and E. Wyn-Jones, Chem. Phys. Lett., 1972, 13, 233. E. A. G. Anniansson and S. Wall, J. Phys. Chem., 1974, 78, 1024; 1975, 79, 857. Holland, 1975). G. J. T. Tiddy, M. F. Walsh, and E. Wyn-Jones, J . Chem. Soc., Chem. Commun.. 1979, 6, 252. J. H. Andrae, R. Bass, E. L. Haesell and J. Lamb, Proc. Phys. Soc. London. 1958. 8, 131. fi R. J. Graber, J. Lang and R. Zara, Kolloid Z.Z. Polym., 1970, 238, 470. ti Chemical and Biological Application of Relaxation Spectrometry, ed. E. Wyn-Jones (D. Reidel, lo F. Eggers, Acoustica, 1968, 19, 233. l 2 W. E. Peel, Ph.D. Thesis (Sheffield Polytechnic, 1974). l 3 R. A. Pethrick, A. M. North and M. A, Cochran, J . Chem. Soc., Faraday Trans 2, 1972, 68, 1719 l4 J. S. Murday and R. M. Cotts, J . Chem. Phys., 1970, 53, 4724. l 5 E. G. Finer and A. Darke, Chern. Phjis. Lipids, 1974, 12, 1 . l 6 A. M. Gottlieb and Y. Lang, Biochim. Acta, 1973, 307, 444. l 7 G. J. T. Tiddy, J . Chem. Soc., Faraday Trans. I , 1972, 68, 369. l a J. Charvolin and P. Rigny, Chem. Phys. Lett., 1973, 18, 515. l 9 H. J. Berendsen and J. Migchelson, J . Chem. Phj?s., 1973,59,296; U. Kaatze, R. Henze and R. Pottel, Lo K. J. Packer, Philos. Truns. R. Soc. London, Ser. B, 1977, 278, 59. *' P. H. Elworthy and C. B. Macfarlane, J . Chem. Soc., 1963, 908; 1964, 31 1 . 22 V. Volkov, Kolloidn. Zh., 1971, 33, 545. " V. Volkov, Kolloidn. Zh., 1972, 33, 802. L4 F. Tokiwa and K. Ohki, J . Phys. Chem., 1967, 71, 1343. 25 A. Johansson and B. Lindman, in Liquid Crystals and Plastic Crystals (Ellis Harwood, Chichester, L6 B. Lindman and H. Wennerstrom, Phys. Rep., 1979,52,2; Topicsin Current Chemistrj) (Springer-Verlag, 27 C. F. Bernasconi, Relaxation Kinetics (Academic Press, London and New York, 1976). L8 P. Ekwall and P. Holmberg, Acta Chem. Scand., 1965. 19, 455. 29 G. S. Manning, Annu. Rev. Phys. Chem., 1972, 23, 117. 30 S. Engstrom and H. Wennerstrom, J . Phys. Chem., 1978, 82, 271 1. 3 1 B. Jonsson and H. Wennerstrom, Chem. Scr.. 1980, 15, 40; H. Wennerstrom, B. Lindman, G . Lindblom and G. J. T. Tiddy, J . Chem. Soc., Faruduy Trans. I , 1979, 75, 663. 32 G. Lindblom, B. Lindman and G. J. T. Tiddy, J . Am. Chem. Soc., 1978, 100, 2299; G. J. T. Tiddy, G. Lindblom and B. Lindman, J . Chem. Soc., Faraday Trans. 1, 1978, 74, 1290. and 1974, 70, 1274. Chem. Phys. Lipids, 1979, 25, 149. 1974), vol. 2, p. 192. Berlin and Heidelberg, 1980), vol. 87, p. 3. (PAPER O/ 1883)

ISSN:0300-9599    

DOI:10.1039/F19827800389

出版商:RSC    

年代:1982

数据来源: RSC