摘要:

Contents 4259 4269 4277 4287 4295 431 1 4321 4335 Protonation Constant of Caffeine in Aqueous Solution M. Spiro, D. M. Grandoso and W. E. Price Ionic Equilibria in Acetonitrile Solutions of 2-, 3- and 4-Picoline N-oxide Perchlorates, studied by Potentiometry and Conductometry L. Chmurzynski, A. Wawrzyn6w and Z. Pawlak Liquid-phase Adsorption of Binary Ethanol-Water Mixtures on NaZSM-5 Zeolites with Different Silicon/Aluminium Ratios W-D. Einicke, M. Heuchel, M. v.Szombathely, P. Brauer, R. Schollner and 0. Rademacher Influence of Oxidation/Reduction Pretreatment on Hydrogen Adsorption on Rh/TiO, Catalysts. An lH Nuclear Magnetic Resonance Study J. P. Belzunegui, J. M. Rojo and J. Sanz Volumetric Properties of Mixtures of Simple Molecular Fluids A. C. Colin, E. G. Lezcano, A.Compostizo, R. G. Rubio and M. D. Peiia Study of Ultramicroporous Carbons by High-pressure Sorption. Part 4.-Iso- thems and Kinetic Transport in Activated Carbons J. E. Koresh, T. H. Kim, D. R. B. Walker and W. J. Koros Kinetic and Equilibrium Studies associated with the Solubilisation of n- Pentanol in Micellar Surfactants G. Kelly, N. Takisawa, D. M. Bloor, D. G. Hall and E. Wyn-Jones The effect of Carboxylic Acids on the Dissolution of Calcite in Aqueous Solution. Part 1 .-Maleic and Fumaric Acids R. G. Compton, K. L. Pritchard, P. R. Unwin, G. Grigg, P. Silvester, M. Lees and W. A. House 130-2Contents 4259 4269 4277 4287 4295 431 1 4321 4335 Protonation Constant of Caffeine in Aqueous Solution M. Spiro, D. M. Grandoso and W. E. Price Ionic Equilibria in Acetonitrile Solutions of 2-, 3- and 4-Picoline N-oxide Perchlorates, studied by Potentiometry and Conductometry L.Chmurzynski, A. Wawrzyn6w and Z. Pawlak Liquid-phase Adsorption of Binary Ethanol-Water Mixtures on NaZSM-5 Zeolites with Different Silicon/Aluminium Ratios W-D. Einicke, M. Heuchel, M. v.Szombathely, P. Brauer, R. Schollner and 0. Rademacher Influence of Oxidation/Reduction Pretreatment on Hydrogen Adsorption on Rh/TiO, Catalysts. An lH Nuclear Magnetic Resonance Study J. P. Belzunegui, J. M. Rojo and J. Sanz Volumetric Properties of Mixtures of Simple Molecular Fluids A. C. Colin, E. G. Lezcano, A. Compostizo, R. G. Rubio and M. D. Peiia Study of Ultramicroporous Carbons by High-pressure Sorption. Part 4.-Iso- thems and Kinetic Transport in Activated Carbons J. E. Koresh, T. H. Kim, D. R. B. Walker and W. J. Koros Kinetic and Equilibrium Studies associated with the Solubilisation of n- Pentanol in Micellar Surfactants G. Kelly, N. Takisawa, D. M. Bloor, D. G. Hall and E. Wyn-Jones The effect of Carboxylic Acids on the Dissolution of Calcite in Aqueous Solution. Part 1 .-Maleic and Fumaric Acids R. G. Compton, K. L. Pritchard, P. R. Unwin, G. Grigg, P. Silvester, M. Lees and W. A. House 130-2

ISSN:0300-9599    

DOI:10.1039/F198985FX033

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE ITALIANA DI CHIMICA FlSlCA DEUTSCHE BUNSEN-GESELLSCHAFT FUR PHYSIKALISCHE CHEMIE KONINKLIJKE NEDERLANDS CHEMISCHE VERElNlGlNG SOCIETE FRANGAISE DE CHIMIE, DIVISION DE CHlMlE PHYSIQUE FARADAY DIVISION GENERAL DISCUSSION No. 90 Colloidal Dispersions University of Bristol, 10-12 September 1990 Orga nising Com mitte e Professor R. H. Ottewill (Chairman) Professor P. Botherol Professor E. Ferroni Or J. W. Goodwin Professor H. Hoff mann Professor A.L. Smith Professor P. Stenius Dr Th. F. Tadros Professor A. Vrij Dr D. A. Young The joint meeting of the Societies will be directed towards examining current understanding of the behaviour of colloidal dispersions. In particular, stability and instability, short range interactions, dynamic effects, non-equilibrium interaction, concentrated dispersions and order-disorder phenomena will form topics for discussion.The preliminary programme is now availablemay 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 SYMPOSIUM No. 26 Molecular Transport in Confined Regions and Membranes Oxford, 17-18 December 1990 Experimental, theoretical and simulation studies which address fundamental aspects of molecular transport will be discussed in the following main areas: a) Transport of atoms and molecules in pores, zeolite networks and other cavities; time-dependent statistical mechanics of small systems in confined geometries b) Molecular transport through synthetic membranes, biological membranes, smectic liquid crystalline phases and Langmuir Blodgett films; the dynamics of the molecules forming the membrane c) Diffusion, reorientation, conformational dynamics, viscosity and conductivity of polymer melts, to include papers dealing with bulk systems since the segments of the polymer will move in the anisotropic environment of the complete chain d) Applications of Brownian dynamics to the study of diffusion in porous media and across membranes including the transport of flexible aggregates such as microemulsions e ) The growth of crystals, colloidal aggregates and droplets on irregular surfaces and in pores Contributions for consideration by the Organising Committee are invited and abstracts of about 300 words should be sent by 31 December 1989 to: Dr D.J. Tildesley, Department of Chemistry, The University, Southampton SO9 SNH. Full papers for publication in the Symposium Volume will be required by August 1990.THE FARADAY DIVISION OF THE ROYAL SOCIETY OF CHEMISTRY ASSOCIAZIONE ITALIANA DI CHIMICA FlSlCA DEUTSCHE BUNSEN-GESELLSCHAFT FUR PHYSIKALISCHE CHEMIE KONINKLIJKE NEDERLANDS CHEMISCHE VERElNlGlNG SOCIETE FRANGAISE DE CHIMIE, DIVISION DE CHlMlE PHYSIQUE FARADAY DIVISION GENERAL DISCUSSION No. 90 Colloidal Dispersions University of Bristol, 10-12 September 1990 Orga nising Com mitte e Professor R. H. Ottewill (Chairman) Professor P. Botherol Professor E. Ferroni Or J. W. Goodwin Professor H. Hoff mann Professor A.L. Smith Professor P. Stenius Dr Th.F. Tadros Professor A. Vrij Dr D. A. Young The joint meeting of the Societies will be directed towards examining current understanding of the behaviour of colloidal dispersions. In particular, stability and instability, short range interactions, dynamic effects, non-equilibrium interaction, concentrated dispersions and order-disorder phenomena will form topics for discussion. The preliminary programme is now availablemay 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 SYMPOSIUM No. 26 Molecular Transport in Confined Regions and Membranes Oxford, 17-18 December 1990 Experimental, theoretical and simulation studies which address fundamental aspects of molecular transport will be discussed in the following main areas: a) Transport of atoms and molecules in pores, zeolite networks and other cavities; time-dependent statistical mechanics of small systems in confined geometries b) Molecular transport through synthetic membranes, biological membranes, smectic liquid crystalline phases and Langmuir Blodgett films; the dynamics of the molecules forming the membrane c) Diffusion, reorientation, conformational dynamics, viscosity and conductivity of polymer melts, to include papers dealing with bulk systems since the segments of the polymer will move in the anisotropic environment of the complete chain d) Applications of Brownian dynamics to the study of diffusion in porous media and across membranes including the transport of flexible aggregates such as microemulsions e ) The growth of crystals, colloidal aggregates and droplets on irregular surfaces and in pores Contributions for consideration by the Organising Committee are invited and abstracts of about 300 words should be sent by 31 December 1989 to: Dr D.J. Tildesley, Department of Chemistry, The University, Southampton SO9 SNH. Full papers for publication in the Symposium Volume will be required by August 1990.

ISSN:0300-9599    

DOI:10.1039/F198985BX035

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

The Future of Faraday Transactions The Faraday Editorial Board has recently given careful consideration to the future of Faraihy Transactions. As a result of this review a number of very important changes will be made h m January 1990 to the appearance, the organisation, and the publishing of the Transactions. We believe that these will further enhance the reputation of the Transactions and increase their attractiveness both to authors and to readers. One of the most obvious changes is that the rather arbitrary separation of the Transactions into two parts will cease. From January 1990, there will be a single journal entitled Journal of the Chemical Society, Far& Transactions, which will appear 24 times a year. The appearance of the journal will be improved markedly by the introduction of new covers, higher quality paper and a more modem, two-column A4 format.The rewLification of the two parts of the Transactions with issues published every two weeks will increase the flexibility of our operations, and the publication of Special Issues, such as Faruday Symposia, will cause virtually no delay to regular papers. Achieving short ‘times to publication’ is viewed as our highest priority. Since the move to the new Cambridge office very considerable improvements have been made in this respect. I am confident that with the journal in its new form we shall achieve publication times of four months after acceptance for normal papers. In addition, we are to introduce a letters section to be called Faraday Communications. Faradby Communications should describe important new advances in no more than 1600 words and we shall guarantee a publication time of 12 weeks, as long as there are no delays due to referees’ comments.It will be possible to submit Faraday Communications via a member of the new International Advisory Editorial Board. The composition of this board will be announced in a forthcoming issue. Faraday Editorial Board is very happy to announce the appointment of Dr Peter S a m of the University of Nottingham as Scientific Editor from 1st October 1989. Without in any way weakening our commitment to those areas of physical chemistry in which the Tranractions have been particularly strong, Dr Sam will be especially committed to attracting high-quality papers in experimental and theoretical chemical physics.Finally, the Editorial Board would like to place on record its appreciation of the huge contribution that Dr David Young made to the development and success of Farday Transactions and Faraa‘uy Discussions during the 19 years that he was Scientific Editor. Ian W. M. Smith Chairman, Faraday Editorial Board F A R IThe Future of Faraday Transactions The Faraday Editorial Board has recently given careful consideration to the future of Faraihy Transactions. As a result of this review a number of very important changes will be made h m January 1990 to the appearance, the organisation, and the publishing of the Transactions. We believe that these will further enhance the reputation of the Transactions and increase their attractiveness both to authors and to readers. One of the most obvious changes is that the rather arbitrary separation of the Transactions into two parts will cease.From January 1990, there will be a single journal entitled Journal of the Chemical Society, Far& Transactions, which will appear 24 times a year. The appearance of the journal will be improved markedly by the introduction of new covers, higher quality paper and a more modem, two-column A4 format. The rewLification of the two parts of the Transactions with issues published every two weeks will increase the flexibility of our operations, and the publication of Special Issues, such as Faruday Symposia, will cause virtually no delay to regular papers. Achieving short ‘times to publication’ is viewed as our highest priority.Since the move to the new Cambridge office very considerable improvements have been made in this respect. I am confident that with the journal in its new form we shall achieve publication times of four months after acceptance for normal papers. In addition, we are to introduce a letters section to be called Faraday Communications. Faradby Communications should describe important new advances in no more than 1600 words and we shall guarantee a publication time of 12 weeks, as long as there are no delays due to referees’ comments. It will be possible to submit Faraday Communications via a member of the new International Advisory Editorial Board. The composition of this board will be announced in a forthcoming issue. Faraday Editorial Board is very happy to announce the appointment of Dr Peter S a m of the University of Nottingham as Scientific Editor from 1st October 1989. Without in any way weakening our commitment to those areas of physical chemistry in which the Tranractions have been particularly strong, Dr Sam will be especially committed to attracting high-quality papers in experimental and theoretical chemical physics.Finally, the Editorial Board would like to place on record its appreciation of the huge contribution that Dr David Young made to the development and success of Farday Transactions and Faraa‘uy Discussions during the 19 years that he was Scientific Editor. Ian W. M. Smith Chairman, Faraday Editorial Board F A R IISSN 0300-9599 JCFTAR 8 5 ( 9 ) 2649-3078 (1 989) JOURNAL OF THE CHEMICAL SOCIETY Faraday Transactions I Physical Chemistry in Condensed Phases 2649 2665 2669 2675 2683 2695 2705 2713 2723 2737 2749 277 1 2785 2797 2809 2819 CONTENTS Experimental Activity Coefficients in Aqueous Mixed Solutions of KCl and KF at 25 "C compared to Monte Carlo Simulations and Mean Spherical Approximation Calculations A Phase Separation caused by the Solubility of Butane in 2-Methylpropan-2- 01-Water Mixtures Solubilization of some Tetramethylammonium Salts and of Ethyltrimethyl- ammonium Bromide by their Homologues in Chloroform Solubility of H,S, CO, and CH, in N-Formyl Morpholine F-Y.Jou, R. D. Deshmukh, F. D. Otto and A. E. Mather Solubilities, Solubility Products and Solution Chemistry of Lanthanon Trifluoride-Water Systems Solubility of Hex- I -ene, Hexane and Cyclohexane in Liquid Nitrogen E.Szczepaniec-Cieciak, M. Kurdziel and L. Ulman Transfer and Partition Free Energies of 1 : 1 Electrolytes in the Water- Dichloromethane Solvent System at 298.15 K A. F. Danil de Namor, R. Traboulssi, F. F. Salazar, V. Dianderas de Acosta, Y. Fernandez de Vizcardo and J. M. Portugal Evidence for Adduct Formation in the Solubilisation of Hydrophobic Compounds by Aqueous Solutions of Urea M. P. Byfield, V. L. Frost, J. L. J. Pemberton and J. M. Pratt Excess Thermodynamic Properties of some 2-Alkoxyethanol-Water Systems G. DouhCret, A. Pal and M. I. Davis Zeolites H-[GaIZSM-5 and H-ZSM-5. A Comparative study of the Introduction of Transition-metal Cations by a Solid-state Reaction A. V. Kucherov, A. A. Slinkin, G. K. Beyer and G.Borbely Characterization of Crystalline and Amorphous Phases during the Synthesis of (TPA, M)-ZSM-5 Zeolites (M = Li, Na, K) J. B.Nagy, P. Bodart, H. Collette, C. Fernandez, 2. Gabelica, A. Nastro and R. Aiello Crystal Structure of Different Dealuminated Y-type Zeolites Determination of Framework Vacancies and Non-framework Species J. Jeanjean, L. Aouali, D. Delafosse and A. Dereigne Effects of Electric Fields on the Rheology of Non-aqueous Concentrated Suspensions L. Marshall, C. F. Zukoski IV and J. W. Goodwin Effect of Solvent on the Reactions of Coordination Complexes. Part 9.-Kinetics of Solvolysis of cis-(Chloro)(cyclohexylamine)bis(ethylene- diamine)cobalt(m) in Methanol-Water, propan-2-ol-Water, Ethylene Glycol- Water and t-Butyl Alcohol-Water media A. C.Dash and J. Pradhan Fluorescence of Cyclotetrasilanes in Rigid Glass at 77 K, a Remarkably Large Stokes Shift H. Shizuka, K. Murata, Y. Arai, K. Tonokura, H. Hiratsuka, H. Matsumoto and Y. Nagai High-resolution 'H Nuclear Magnetic Resonance Study of Hydrogen Sulphide Adsorption on Heterogeneous Catalysts V. M. Mastikhin, I. L. Mudrakovsky, A. V. Nosov and A. V. Mashkina T. S. Ssrensen, J. B. Jensen and P. Sloth R. W. Cargill and D. E. MacPhee J. Czapkiewicz M. P. Menon and J. James 88-2Con tents 2827 2835 2847 2857 2867 2875 289 1 290 1 2909 2917 293 1 294 1 2945 2953 2963 2973 2983 299 1 2999 Nuclear Magnetic Resonance Studies of Preferential Solvation. Part 6.- Application of Blander's Coordinated Cluster Theory to the Methanol-Water Solvent System A. K.Covington and M. Dunn Nuclear Magnetic Resonance Studies of Preferential Solvation. Part 7.-Sodium Iodide in Ethylene Glycol-Acetonitrile and in Propylene Glycol-Acetonitrile Mixtures Infrared Study of the Thermal Decomposition of Heteropolyacids of the Series H3+z PMo,,-, V,O,, A. Bielanski, A. Malecka and L. Kubelkova Dielectric Properties of Water in the Coexisting Phases of Aqueous Polymeric Two-phase Systems B. Yu. Zaslavsky, L. M. Miheeva, M. N. Rodnikova, G. V. Spivak, V. S. Harkin and A. U. Mahmudov Growth, Characterisation and Ultrasonic Studies of a Rotator-phase Solid, Carbon Tetrabromide CO Hydrogenation using Cobalt/Manganese Oxide Catalysts. Comments on the Mechanism of Carbon-Carbon Bond Formation G. J. Hutchings, M. van der Riet and R. Hunter Application of the Competitive Preferential Solvation Theory to Coordinative Solute-Solvent Interactions M.Szpakowska and 0. B.Nagy Enthalpies of Interaction of Sodium Chloride and Potassium Chloride with some Amides in Water at 25 "C K. G. Davis, M. A. Gallardo-JimCnez and T. H. Lilley Enthalpies of Interaction of some Alkali-metal Halides with N-Methyl- acetamide and with N,N-Dimethylformamide in Water at 25 "C M. A. Gallardo-JimCnez and T. H. Lilley Cationic Lead(r1) Halide Complexes in Molten Alkali-metal Nitrate. Part 3.-The Structure of Pb,X3+ and the Solvated Pb" Ion, determined by Liquid X-Ray Scattering and Raman Spectroscopy Thermodynamic Classification - of Anions through Constituent Analysis of Transfer Enthalpies in Acetonitrile-Methanol Mixtures Y. Kondo, T.Fuji- wara, A. Hayashi and S. Kusabayashi Study of the Ammonia-Zeolite Interaction in Modified ZSM-5 by Temperature- programmed Desorption of Ammonia W. Reschetilowski, B. Unger and K-P. Wendlandt Glass/Rubber Transitions and Heat Capacities of Binary Sugar Blends L. Finegold, F. Franks and R. H. M. Hatley Shape-selective Adsorption of Aromatic Molecules from Water by Tetra- methylammonium-Smectite J-F. Lee, M. M. Mortland, s. A. Boyd and C. T. Chiou Separation of Small-particle Dispersions by the Preferential Accumulation in One of Two Liquid Phases, or by Static Flotation at their Interface E. A. Boucher High-resolution Solid-state 13C Nuclear Magnetic Resonance Study of the Dynamic Behaviour of Tetramethylammonium Ions trapped in Zeolites S. Hayashi, K.Suzuki and K. Hayamizu Infrared Study of the Effects of Oxidation/Reduction Treatments on Pt Dispersion in Pt/Al,O, Catalysts J. A. Anderson, M. G. V. Mordente and C. H. Rochester Effects of Oxidation-Reduction Treatments of Pt/A1,0, on Catalytic Activity and Selectivity for Hexane Reforming J. A. Anderson, M. G. V. Mordente and C. H. Rochester Kinetics of the Solvolysis of trans-Dichlorotetra(4-t-butylpyridine)cobalt(111) Ions in Water and in Water-Propan-2-01 Mixtures K. H. M. Halawani and C. F. Wells A. K. Covington and M. Dunn C. S. Yoon, J. N. Sherwood and R. A. Pethrick L. Bengtsson and B. HolmbergContents 301 1 Thermodynamic Properties of Amphiphilic Drugs in Aqueous Solution D. Attwood, V. Mosquera and V. P. Villar 3019 Preferential Solvation of Ions in Mixed Solvents. Part 4.-Comparison of the Kirkwood-Buff and Quasi-lattice Quasi-chemical Approaches 3033 Crystallochemical Characterization of Magnetic Spinels prepared from Aqueous Solution S.Mann, N. H. C. Sparks, S. B. Couling, M. C. Larcombe and R. B. Frankel Dielectric Studies of the Switch-over Mechanism in the Principal Relaxation Process of Alkan- 1-01s H. Mandal, D. G. Frood, M. Habibullah, L. Humeniuk and S. Walker Numerical Interpretation of Oscillatory Glow and Ignition during Carbon Monoxide Oxidation in a well-stirred Flow Reactor J. F. Griffiths and A. F. Sykes 3071 Reviews of Books A. Hamnett; M. Springford; D. R. Rosseinsky; M. B. Goatly; M. Streat; P. D. I. Fletcher Y. Marcus 3045 3059Contents 301 1 Thermodynamic Properties of Amphiphilic Drugs in Aqueous Solution D. Attwood, V. Mosquera and V. P. Villar 3019 Preferential Solvation of Ions in Mixed Solvents. Part 4.-Comparison of the Kirkwood-Buff and Quasi-lattice Quasi-chemical Approaches 3033 Crystallochemical Characterization of Magnetic Spinels prepared from Aqueous Solution S. Mann, N. H. C. Sparks, S. B. Couling, M. C. Larcombe and R. B. Frankel Dielectric Studies of the Switch-over Mechanism in the Principal Relaxation Process of Alkan- 1-01s H. Mandal, D. G. Frood, M. Habibullah, L. Humeniuk and S. Walker Numerical Interpretation of Oscillatory Glow and Ignition during Carbon Monoxide Oxidation in a well-stirred Flow Reactor J. F. Griffiths and A. F. Sykes 3071 Reviews of Books A. Hamnett; M. Springford; D. R. Rosseinsky; M. B. Goatly; M. Streat; P. D. I. Fletcher Y. Marcus 3045 3059

ISSN:0300-9599    

DOI:10.1039/F198985FP111

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

JOURNAL OF THE CHEMICAL SOCIETY Faraday Transactions 11, Issue 9 Molecular and Chemical Physics For the benefit of readers of Faraday Transactions I , the contents list of Faraday Transactions II, issue 9, is reproduced below. 1377 1391 1401 1413 1425 1439 1447 1465 1417 1487 1505 1519 1531 1539 1553 1575 1585 The Microscopic Origin of Cooperativity and its Effect on Long-lifetime Kinetic Modes for Template-free RNA Synthesis A. Fernhdez Electronic and Geometric Structure of the Titanium Hydrides, T i p and TiH2+ A. Mavridis and J. R. Harrison Quenching of Singlet Molecular Oxygen, 0 2 (a 'Ag) and 0 2 (b '&+), by Ha, D2, HCl and HBr P. Borrell and D. S. Richards A Molecular Beam Mass-spectrometric Study of Isopropyl Nitrate Pyrolysis Reactions at Short Residence Times and Temperatures up to 700 K T.Hansson, J. B. C. Pettersson and L. Holmlid Kinetics and Mechanism of the Thermal Gas-phase Oxidation of 1 ,I-Dichlorofluoroethane, CF2CC12 by Molecular Oxygen in the Presence of Trifluoromethyl Hypofluorite, CF30F J. Czarnowski Molecular-mechanical Study of Cyclodextrin and its Inclusion Complexes T-X. Lii, D-B. Zhang and S-J. Dong Reaction of H Atoms with some Silanes and Disilanes. Rate Constants and Arrhenius Parameters N. L. Arthur, P. Potzinger, B. Reimann and H. P. Steenbergen Tautomerism and Intramolecular Hydrogen Bonding in P-Thioxoketones. An Ab Initio Study on Monothiomalondialdehyde S. Millefiori and A. Millefiori Polarized Absorption Spectra of some Aromatic Radical Anions in Stretched Polyethylene Films Y. Kubozono, M. Ata, M.Aoyagi, T. Fujita, N. Fukami and Y. Gondo The Probability of Capture and its Impact on Floc Structure R. D. Cohen Effective Spherical Potentials for Liquid Propane R. J. Bowles, D. J. Tildesley and N. Quirke Momentum Space Studies of Large Polyenes D. L. Cooper, N. L. Allan and P. J. Grout Effect of Diffuse Functions on Potential-energy Surfaces of the Alkylation Reactions X- + CH3F -+ XCH3 + F (X = OH, CH3, H2CCHO) E. Sanchez Marcos and J. Bertran A Time-resolved Study of the Photoinduced Tautomerization of 2-Aminopyridine and Derivatives within Specific 1:1 Complexes with Carboxylic Acids J. Konijnenberg, A. H. Huizer and C. A. G. 0. Varma Thermal Shape Fluctuations of Spherical Swfactant Micelles S. Ljunggren and J. C. Eriksson Double Ionization Energies of the Fluorethane Molecules C~HSF, CH3CHF2, CH2FCHF2, CH3CF3, CHF2CF3 and C&6 measured by Double-charge-transfer Spectroscopy W.J. Griffiths and F. M. Harris An Ion-Dipole Mixture against a Charged Hard Wall with Specific Adsorption C. W. Outhwaite and M. Molero1601 Reactions of Methyl Radicals with Cyclohexane and Methyl-substituted Cyclohexanes K. Al-Niami, K. A. Holbrook and G. A. Oldershaw 1609 Corrigendum to: A Photophysical and Theoretical Study of Styrylanthracenes G. Bartocci, F. Masetti, U. Mazzucato, A. Spalletti, G. Orlandi and G. Poggi Corrigendum to: Polarizability Anisotropies of C-H, C-C, C-Cl and C-Br Bonds G. W. Allen, R. S. Armstrong, M. J. Aroney, R. K. Pierens and A. J. Williams 1611 Reviews of Books J. N. L. Connor; J. N.Sherwood; D. M. Hirst.The following papers were accepted for publication in Faraduy Transactions I during June 1989. 9/00281B 9/00931K 9/00938H 9/00939F 9/00943D 9/00947G 9/01035A 9/01181A 9101351B 9/01402K 9/0 150 1 I 9/01546I 9/01556F 9/01576K 9/01587F 9/0 159 1 D The Influence of Ionic Association on the B Coefficient of the Jones-Dole Equation for NaI in Water-t-Butyl Alcohol Mixtures at 26 "C Kaoperska, A., Bald., A, Szejgis, A. and Taniewska-Osinska, S. Intercalation Mechanism of Nitmgenated Bases into V205-Xerogel Casal, B., Crespin, M,, Galvan, J. C., Tinet, D. and Ruiz-Hitzky, E. An EPR Study of the Effect of Temperature upon Copper-impregnated Titanium Dioxide Powders Amorelli, A., Evans, J. C. and Rowlands, C. C. Effect of Carboxylic Acids on the Dissolution of Calcite in Aqueous Solution.Part 1.Maleic and Fumaric Acids Compton, R. G., Pritchard, K. L. and Unwin, P. R. Molecular Self-diffusion of Methane in Zeolite ZSM-5 by Quasi-elastic Neutron Scattering and NMR Pulsed Field Gradient Techniques Jovic, H., Bee, M., Caro, J., Bulow, M. and Karger, J. The General Isotherm Equation for Adsorption of Surfactants at SolidLiquid Interfaces. Part 1.-Theoretical Zhu, B-Y. and Gu, T. The General Isotherm Equation for Adsorption of Surfactants at SolidLiquid Interfaces. Part 2.4xperimenta.l Zhu, B-Y. and Gu, T. Interlayer Water Molecules of Vanadium Pentoxide Hydrate. Part l.-Phase Equilibrium with Water Vapour Kittaka, S., Ayatsuka, Y., Ohtani, K. and Uchida, N. Solubilization of Pentanol in Sodium Dodecyl Sulphate Micelles.Interpretation of Calorimetric Results using a Theoretical Model Johnson, I., and Olofsson, G. Viscosity B Coefficients of Alkali-metal Chlorides and Bromides in 2-Methoxyethanol at 25 and 35 "C Hazra, D. K. and Nandi, D. The Protonation Constant of Caffeine in Aqueous Solution Spiro, M., Grandoso, D. M. and Price, W. E. Travelling Waves in the Acidic NitrateFerroin Reaction Pota, G., Lengyel, I. and Bazsa, G. Liquid-phase Adsorption of Binary Ethanol-Water Mixtures on NaZSM-5 Zeolites with Different Silicon-to-aluminium Ratios Einicke, W-D., Heuchel, M., von Szombathely, M., Brauer, P., Rademacher, 0. and Schollner, R. Catalytic Activity of SAP05 for the Cracking of Butane and Hexane Lercher, J. A., Halik, C. and Chaudhuri, S. Kinetics of Bridging Flocculation: The Role of Relaxations in the Polymer Layer Pelssers, E.G. M., Cohen Stuart, M. A. and Fleer, G. J. Influence of Oxidation-Reduction Pretreatments on the Hydrogen Adsorption on Rh/Ti@ Catalysts. A 'H Nuclear Magnetic Resonance Study Rojo, J. M., Sanz, J. and Belzunegui, J. P. The Chemisorption of Linear and Cyclic Polymethylsiloxanes on Alumina studied by Fourier-transform Infrared Spectroscopy Vincent, B., Cosgrove, T. and Prestidge, C. A. An EPR Study of Monoclinic Zirconium Dioxide Polycrystalline Powders doped with Paramagnetic Transition-metal Ions Evans, J, C., Owen, C. R. and Rowlands, C. C. An EPR Study of Transition-metal Ion Impregnated Brookite Titanium Dioxide Powders Evans, J. C., Amorelli, A. and Rowlands, C. C. (iii)9/01611B 9/01616B 9/01783F 9/018711 The Reactions of Scandium Ions in Hydrocarbon Matrices Howard, J.A., Mile, B., Hampson, C. A. and Morris, H. The Use of EPR Techniques in the Molecular Approach to Heterogeneous Catalytic Processes on Oxides Che, M., Louis, C. and Sojka, Z. An EXAFS Investigation of Structural Changes induced during the Pretreatment of a Titania-supported Iron-Ruthenium Catalyst Berry, F. J. Kinetic and Equilibrium Studies associated with the Solubilization of n-Pentanol in Micellar Surfactants Takisawa, G., Kelly, N., Bloor, D. M., Hall, D. G. and Wyn- Jones, E. 9/01903K Photocurrent Kinetics of Group I11 Metal-(Phthalocyaninato)Halogen Couves, J. W., Tamizi, M. and Wright, J. D. 9m132I Calorimetric and Spectrophotometric Studies of Complexation of Manganese(II), Cobalt(II) and Nickel(1I) with Bromide Ions in N,N-Dimethylformamide Ishiguro, S-I.and Ozutsumi, K. 9/02181G Effect of Pressure on the Electrical Conductivity of the Molten Dichlorides, Bromides and Iodides of Cadmium, Tin and Lead Cleaver, B. and Kumar, S. P.Cumulative Author Index 1989 Abe, M., 1493 Adachi, K., 1065, 1075, 1083 Agathonos, P., 1357 Aguilella, V. M., 223 Aiello, R., 2749 Akitt, J. W., 121, 2035 Albery, W. J., 1181, 1189 Al-Bizreh, N., 1303 Albuquerque, L. M. P. C., 207 Allen, G. C., 55 Almeida, B. S., 1217 Amodeo, P., 621 Anderson, J. A., 11 17, 1129, 2983, 2991 Anderson, M. W., 1945 Anpo, M., 609 Anzai, S., 2499 Aouali, L., 2771 Apelblat, A., 373 Arai, T., 929, 1451 Arai, Y., 2369, 2809 Archer, M. D., 1027 Aruga, T., 2597 Asakura, K., 441, 2021 Attwood, D., 3011 Austin, J.C., 1159 Azenha, M. E. D. G., 2625 Bahra, G. S., 1979 Baiker, A., 999 Bakshi, M. S., 2285, 2297 Bald, A., 479 Barlow, M. T., 1945 Barone, G., 621, 2087 Barone, V., 621 Bartle, K. D., 2347 Beckett, M. A., 727 Bie, M., 2525 Bellotto, M., 895 Bengtsson, L., 305, 317, 2917 Berry, F. J., 467 Bertoldi, M., 237 Bertran, J., 1207 Beyer, G. K., 2737 Beyer, H. K., 2127 Bicelli, L. P., 1685 Bielanski, A., 2847 Bird, R., 2173 Black, S. N., 1795 Blandamer, M. J., 735, 1809 Bloor, D., 2099 Bodart, P., 2749 Bolis, V., 855, 1383 Bolton, J. R., 1027 Bond, G. C., 168 Borbdy, G., 2127 Borbely, G., 2737 Borowko, M., 343 Bosch, H., 1425 Boss, R. D., 11 Boucher, E. A., 2963 Bowker, M., 165, 2635 Boyd, S. A., 2953 Brimblecome, P., 157 Brookes, B.I., 2173 Brown, P., 2099 Brown, R., 2159 Bruce, J. M., 2647 Biilow, M., 1501 Burgess, J., 735, 1809 Burkhardt, I., 21 13 Burrows, H. D., 2625 Busca, G., 137, 237 Byfield, M. P., 2713 Cai, F. X., 1991 Calado, J. C. G., 1217 Campbell, J. A., 843 Campelo, J. M., 2535 Carbonara, M., 1257 Cargill, R. W., 2665 Carlstrom, G., 1049 Caro, J., 1501 Cassol, A,, 2445 Castronuovo, G., 2087 Cattania, M. G., 801 Chadwick, A. V., 166, 1979 Chandra, H., 1801 Che, M., 609 Chen, J., 829 Chen, L-f., 33 Chien, S-H., 2199 Chiou, C. T., 2953 Claesson, P. M., 1933 Cleaver, B., 2453 Clegg, S. L., 157 Clifford, A. A,, 2347 Cohen, H., 1169 Collette, H., 2749 Colling, C . N., 1303 Coluccia, S., 609, 1655 Comninos, H., 633 Compton, R. G., 761, 773, 977, 1821, 2255, 2273 Condlyffe, D.H., 2453 Conway, B. E., 2355 Conway, S. J., 71, 79, 1841 Cooper, J., 1365 Copperthwaite, R. G., 633 Costas, M., 221 1 Cottrell, M. R., 1809 Coudurier, G., 1607, 2615 Couling, S. B., 3033 Couves, J. W., 1979 Covington, A. K., 2827, 2835 Cox, B.G., 187 Cristiani, C., 895 Cristinziano, P., 621 Cruz, M. J., 2071, 2079 Cucinotta, V., 2445 Czapkiewicz, J., 2669 da Silva Pereira, M. I., 2473 da Costa, F. M. A., 2473 da Costa, M. A., 907 Dalas, E., 2465 Danil de Namor, A. F., 2705 Das, I., 201 1 Das, P. K., 2405 Das, S., 1531 Dash, A. C., 2405, 2797 Datka, J., 47, 837 Davey, R. J., 1795 Davis, K. G., 2901 Davis, M. I., 2723 Dawber, J. G., 727 de Vizcardo, Y. F., 2705 de Acosta, V. D., 2705 De Giglio, A., 23 Delafosse, D., 2771 Dell’Atti, A,, 23 Del Vecchio, P., 2087 Dereigne, A., 2771 Deschaux, M., 2605 Deshmukh, R.D., 2675 Di Bernardo, P., 2445 Dianoux, A. J., 2525 Dickel, G., 1463, 1671 Ding, J., 1599 Domen, K., 929, 1451 Dong, S., 1575, 1585, 1599 Donini, J. C., 91 Douhiret, G., 2723 Downs, G. W., 1841 Drummond, C. J., 521, 537, 551, Dunn, M., 2827, 2835 Easteal, A. J., 1091 Eden, J., 991 Egawa, C., 2597 Elders, J. M., 2035 Elisei, F., 1469 El Jamal, M., 2615 el Torki, F. M., 349 Endoh, A., 1327 Ernst, S., 2127 Espinos, J. P., 1279 Esteso, M. A., 2575 56 1Fahim, R. B., 1723 Falconer, J. W., 71, 79, 1841 Fernandez, A., 1279 Fernandez, C., 2749 Fernandez-Merida, L., 2575 Fernandez-Pineda, C., 1019 Ferranti, F., 2241 Finch, J. A., 91 Finegold, L., 2945 Flanagan, T. B., 1787 Fletcher, P. D. I., 147, 3075 Foerch, R., 1139 Foo, C.H., 65 Forissier, M., 1607, 2615 Formosinho, S. J., 2625 Forster, H., 1149 Forzatti, P., 895 Franchini, G., 1697 Frankel, R. B., 3033 Franks, F., 2417, 2945 Frey, H. M., 167 Frood, D. G., 3045 Frost, V. L., 2713 Fubini, B., 237, 855, 1383 Fujiwara, T., 2931 Fulop, V., 2127 Gabelica, Z., 2749 Gabriel, C. J., 11 Gabrys, B., 168 Gadzekpo, V. P. Y., 1027 Gallardo-Jiminez, M. A., 290 1, Gans, P., 1835 Garcia, A., 2535 Garrone, E., 585, 1373, 1383 Garst, J. F., 1245 Gasser, D., 999 Gavish, B., 1199 Gelsthorpe, M. R., 2641 Gervasini, A., 801 Geus, J. W., 269, 279, 293, 1267 Giamello, E., 237, 855, 1373 Giancola, C., 2087 Gilbert, P. J., 147 Gill, D. S., 2285, 2297 Gill, J. B., 1835 Gillette, G., 2369 Girault, H. H., 843 Goatly, M. B., 3074 Golding, P. D., 2229 Gonzalez-Diaz, 0. M., 2575 Gonzalez-Elipe, A.R., 1279 Gonzalez-Lafont, A., 1207 Goodwin, J. W., 2785 Gorner, H., 1469 Gottschalk, F., 363 Grieser, F., 521, 537, 551, 561 Griffiths, J. F., 3059 Guardado, P., 735 Guest, A., 1897 Guil, J. M., 1775 Gutierrez, C., 907 Gutschick, D., 21 13 Guy, P. D., 1795 Guyan, P. M., 2647 2909 AUTHOR INDEX Habibullah, M., 3045 Hagele, G., 1409 Hakin, A. W., 1809 Halawani, K. H., 2185 Halawani, K. H. M., 2999 Hall, D. G., 1881, 2099 Halle, B., 1049 Hamnett, A., 3071 Hampton, S., 773 Han, S., 829 Handreck, G. P., 645 Harkin, V. S., 2857 Harland, R. G., 761, 2273 Harris, R. K., 1409, 1853 Harrison, P. G., 1897, 1907, Hasselaar, M., 1267 Hasted, J. B., 99 Hatano, M., 199 Hatley, R. H. M., 2945 Hayamizu, K., 2973 Hayashi, A., 2931 Hayashi, S., 2973 Hazra, D.K., 1531 Healy, T. W., 521, 537, 551, 561 Heatley, F., 917 Hegarty, B. F., 1861 Herder, C. E., 1933 Herder, P. C., 1933 Hernandez-Luis, F. F., 2575 Hesselink, W. H., 389 Hester, R. E., 171, 1159 Hey, M. J., 1743 Hibino, T., 2327 Higgins, J. S., 170 Higuchi, A., 127 Hill, W., 691 Hirai, T., 969 Hiratsuka, H., 2809 Hisada, O., 2555 Holmberg, B., 305, 317, 2917 Holz, M., 1257 Hong, C. T., 65 Hori, Y., 2309 Horn, I. M., 1809 Howard, J., 1233 Howarth, 0. W., 121, 2035 Hubbard, C. D., 735 Humeniuk, L., 3045 Hummel, A., 991 Hunger, M., 1501 Hunter, R., 363, 633, 2875 Hussein, G. A. M., 1723 Hutchings, G. J., 363, 633, 2507, Hwang, L-P., 2335 Ichikawa, K., 175 Ikeda, R., 11 1 Ikeda, S., 1619 Ikeda, Y., 1099 Imamura, H., 1647 Imanishi, Y., 1065, 1075, 1083 Indelli, A., 2241 Inoue, Y., 1765 Ishida, H., 111 1921 2875 (vi) Ishiguro, S., 2587 Itaya, K., 1351 Ito, K., 2555 Ito, T., 2381 Itoh, N., 493 Iwasawa, Y., 441, 2021, 2597 James, J., 2683 Jeanjean, J., 2771 Jensen, J.B., 2649 Jiang, R., 1575, 1585 Jin, T., 175 Jobic, H., 2525 Johnson, G. R. A., 677 Johnston, C., 11 11 Jonkers, G., 389 Jorgensen, N., 11 11 Joiwiak, M., 2141 Juillard, J., 1337, 1709 Jutson, J. A., 55 Kaneko, K., 869 Kanno, T., 579 Karaiskakis, G., 1357 Karger, J., I501 Kato, S., 1619 Kato, T., 2499 Katoh, T., 127 Keeler, J. H., 163 Kelebek, $., 91 Kemball, C., 2159, 2173 Kermarec, M., 1991 Khan, S. U. M., 2001 Kim, T. H., 1537, 1545, 1557 Kinjo, K., 2555 Kishi, R., 655 Kishimoto, S., 1787 Kitajima, K., 1647 Kiwi, J., 1043 Klinowski, J., 1945 Knijff, L.M., 269, 293 Kobayashi, M., 579 Koda, S., 957 Kondo, Y., 2931 Koresh, J. E., 1537, 1545, 1557 Koros, W. J., 1537, 1545, 1557 Kosugi, N., 869 Kotaka, T., 1065, 1075, 1083 Koutsoukos, P. G., 2465 Kozjowski, Z., 479 Kubelkova, L., 2847 Kucherov, A. V., 2737 Kurdziel, M., 2695 Kuriacose, J. C., 2249 Kuroda, H., 869 Kusabayashi, S., 2931 Kuwabata, S., 969 Lamotte, J., 2397 Lancaster, N. M., 1303, 1315 Land, E. J., 2647 Larcombe, M. C., 3033 Larramona, G., 907 Laschi, F., 601 Lavalley, J. C., 2397 Lawrence, D. G., 1365 Lawrence, K. G., 23 JOU, F-Y., 2675Leach, H. F., 2173 Lee, J-F., 2953 Lelj, F., 621 Lepetit, C., 1991 Lewis, T. J., 1009 Leyendekkers, J. V., 663 Lhermet, C., 1709 Li, C., 929, 1451 Lilley, T. H., 2901, 2909 Lillford, P. J., 2417 Liu, J.-Y., 1027 Liu, T., 1607 Lluch, J.M., 1207 Longdon, P. J., 1835 Lorenzelli, V., 137 Loudon, R., 169 Louis, C., 1655 Lowe, B. M., 945 Lubetkin, S. D., 1753 Luna, D., 2535 Lund, A., 421 Machin, W. D., 2229 MacPhee, D. E., 2665 Maeda, M., 2555 Mafk, S., 223 Maffi, S., 1685 Mahmudov, A. U., 2857 Maignan, A., 783 Makcka, A., 2847 Malet, P., 1279 Malitesta, C., 1685 Manchado, M. C., 1775 Mandal, H., 3045 Mann, S., 3033 Manzurola, E., 373 Marcantonatos, M. D., 248 1, Marchese, L., 1655 Marchetti, A., 1697 Marchettini, N., 2149 Marcus, Y., 381, 3019 Marinas, J. M., 2535 Markovits, G., 373 Marshall, L., 2785 Maruya, K., 929, 1451 Mashkina, A. V., 2819 Masiakowski, J. T., 421 Mastikhin, V. M., 2819 Mather, A. E., 2675 Matsuhashi, N., 11 1 Matsui, H., 957 Matsumoto, H., 2369, 2809 Matsumoto, T., 175 Matthews, R.W., 1291 Mayagoitia, V., 2071, 2079 Mazzucato, U., 1469 McAleer, J. F., 783 McLure, I. A., 1217 McMurray, N., 2047, 2055 Meima, G. R., 269, 279, 293, Melo, M-J. B. V., 2473 Menon, M. P., 2683 Merwin, L. H., 1409 Levy, o., 373 2605 I267 AUTHOR INDEX Meyerstein, D., 1169 Miessner, H., 691, 2113 Miguel, M. de G. M., 2625 Miheeva, L. M., 2857 Mills, A., 503, 2047, 2055 Mills, D., 2347 Mitani, T., 1485 Mitchell, B., 1795 Miyoshi, H., 1873 Mizoe, K., 1327 Mizuno, K., 1099 Morazzoni, F., 801, 2581 Mordente, M. G. V., 2983, 2991 Morel, J-P., 1709 Morel-Desrosiers, N., 1709 Moreno, M. S., 2535 Morrison, C., 1043 Morterra, C., 1383, 21 13 Mortland, M. M., 2953 Morton, J. R., 1963 Moseley, P. T., 783 Mosier-Boss, P. A., 11 Mosquera, V., 3011 Moulder, R., 2347 Mount, A.R., 1181, 1189 Mousset, G., 1337 Mudrakovsky, I. L., 2819 Mukherjee, T., 2647 Munuera, G., 1279 Murakami, Y., 2327 Murata, A., 2309 Murata, K., 2369, 2809 Nagai, Y., 2369, 2809 Nagy, J. B., 2749 Nagy, 0. B., 2891 Nakagawa, T., 127 Nakamura, D., 111 Nakamura, T., 493 Nandi, D., 1531 Nastro, A., 2749 Natarajan, P., 813 Nazhat, N. B., 677 Neagle, W., 429, 719 Neilson, G. W., 1365 Newman, K. E., 485 Nicholas, A., 773 Nicol, J. M., 1233 Niwa, M., 2327 Nomura, H., 957, 1619 Northing, R. J., 2273 Nosov, A. V., 2819 Nowak, R. J., 11 Nowicka, B., 479 Nunes, M. R., 907, 2473 Ohlmann, G., 691 Ohtaki, H., 2587 Ohyama, Y., 749 Okubo, T., 455, 749 Olier, R., 2615 Oliva, A., 1207 Oliveira Jr, 0. N., 1009 Olivier, D., 1991 Onishi, H., 2597 Onishi, T., 929, 1451 Orchard, S.W., 363 Osada, Y., 655 Otsuka, K., 199 Otto, F. D., 2675 Pacynko, W. F., 1397 Padmaja, S., 2249 Pal, A., 2723 Pandey, J. D., 331 Paniego, A. R., 1775 Patterson, D., 221 1 Pawlowska, M. M., 2481 Pellicer, J., 223 Pemberton, J. L. J., 2713 Pereira, I., 907 Peter, L. M., 2473 Petersen, R. L., 2435 Pethrick, R. A., 2867 Pilkington, M. B. G., 2255 Piwowarska, Z., 47, 837 Pope, C. G., 945 Portanova, R., 2445 Portugal, J. M., 2705 Portwood, L., 711, 1801 Pradhan, J., 2797 Pratt, J. M., 2713 Preston, K. F., 1963 Preti, C., 1697 Price, W. E., 415, 1091 Pritchard, T. N., 1853 Pudney, P, 2635 Rai, R. D., 331 Rajaram, J., 2249 Ramakrishnan, V., 2249 Ramaraj, R., 813 Ramis, G., 137 Rao, K. J., 251 Rastogi, R. P., 201 1 Raven, C. I., 1743 Reed, W.F., 349 Rees, L. V. C., 33, 1501 Reis, J. C. R., 207 Reller, A., 855 Reschetilowski, W., 2941 Rhodes, C. J., 711 Roberts, G., 2635 Robinson, G., 2417 Rochester, C. H., 71, 79, 429, 719, 1111, 1117, 1129, 1841, 2983, 2991 Rodnikova, M. N., 2857 Rodriguez-Mellado, J. M ., 1567 Rojas, F., 2071, 2079 Rooney, J. J., 1861 Rosen, D., 99 Rosseinsky, D. R., 3073 Rossi, C., 601, 2149 Rowlinson, J. S., 171, 172 Ruddick, A. J., 1795 Ruiz, J. J., 1567 Saadalla-Nazhat, R. A., 677 Sabbatini, L., 1685 Sacco, A., 23, 1257 Said, M., 99 Sakata, Y., 929, 1451 Salazar, F. F., 2705 Salvagno, S., 1009 (vii)Sanchez, F., 1809 Sarkany, A., 1511, 1523 Sato, Kazunori, 1765 Sato, Kiyoshi, 1765 Saussey, J., 2397 Sayari, A., 1963 Schmehl, R. H., 349 Schmidt, J. A., 1027 Schneider, H., 187 Schneider, I., 187 Schumann, M., 1149 Sciotto, D., 2445 Scotti, R., 801, 2581 Scurrell, M.S., 2507 Seimiya, T., 2499 Selvaraj, U., 251 Sharma, A., 2011 Sheppard, N., 1723 Sherwood, J. N., 2867 Shido, T., 441 Shindo, Y., 1099 Shizuka, H., 2369, 2809 Shukla, R. K., 331 Sijpkes, A. H., 2563 Sinot, P. J., 1425 Slinkin, A. A., 2737 Sloth, P., 2649 Smith, E. G., 1853 Smith, G. W., 91 Smith, J. J., 11 Smith, M. R., 467 Smith, T. D., 645 Soares, V. A. M., 1217 Somsen, G., 2563 Song, S., 1575 Sorek, Y., 1169 S~rrensen, T. S., 2649 Sozzani, P., 2581 Sparks, N. H. C., 3033 Spivak, G. V., 2857 Spoto, G., 21 13 Stevens, A. D., 1439 Stewart, A. A., 843 Stirling, C. J. M., 1009 Streat, M., 3075 Stroka, J., 187 Strumolo, D., 801 Sugawara, S., 1351 Sundar, H. G. K., 251 Sutton, H. C., 883 Suzuki, H., 2587 Suzuki, K., 2973 Swallow, A.J., 2647 Sykes, A. F., 3059 AUTHOR INDEX Symons, M. C. R., 711, 1439, Szczepaniec-Cieciak, E., 2695 Szejgis, A., 479 Szpak, S., 11 Szpakowska, M., 2891 Taiwo, F. A., 2427, 2435 Takagi, T., 1099 Takagi, Y., 493 Takahashi, R., 2309 Takaishi, T., 1327 Takano, S., 2499 Takisawa, N., 2099 Tamura, K., 1493 Tanabe, T., 1787 Tanaka, H., 2369 Taniewska-Osinska, S., 479, Tashiro, T., 2381 Tassi, L., 1697 Taylor, D. M., 1009 Terai, M., 1493 Thamm, H., 1 Themistocleous, T., 633 Theocharis, C. R., 2641 Thomas, J. M., 1945 Tiddy, G. J. T., 1397 Timmennann, E. O., 163 1 Tissier, M., 1337 Toi, K., 2381 Tolazzi, M., 2445 Tomat, G., 2445 Tonokura, K., 2369, 2809 Tosi, G., 1697 Traboulssi, R., 2705 Tsuchiya, S., 1647 Tsutsumi, K., 1327 Tsyganenko, A.A., 2397 Ugliengo, P., 585, 1373 Ulgiati, S., 2149 Ulman, L., 2695 Unger, B., 2941 Unwin, P. R., 1821 Urch, D. S., 1139 Vaccari, A., 237 van Buren, F. R., 269, 279, 293, Van-Den-Begin, N., 1501 van der Riet, M., 2875 van Dillen, A. J., 269, 279, 293, van Leur, M. G. J., 279 van Lith, D., 991 van Rensburg, L. J., 633 1801, 2427, 2435 2141 1267 1267 van Veen, J. A. R., 389 Vazquez-Gonzalez, M. I., 1019 Vedrine, J. C., 1607 Vedrine, J. C., 2615 Villar, V. P., 301 1 Vink, H., 699 Vis, R. J., 269, 279 Vuilleumier, J-J., 2605 Wacker, T., 33 Walker, D. R. B., 1545, 1557 Walker, P. A. M., 1365 Walker, S., 3045 Waller, A. M., 773, 977 Wang, P-L., 2335 Wang, Y-P., 2199 Warman, J. M., 991 Watanabe, T., 2381 Waugh, K. C., 163 Weale, K. E., 165 Weitkamp, J., 2127 Wells, C.F., 2185, 2999 Wendlandt, K-P., 2941 West, R., 2369 Wilkinson, D. P., 2355 Willett, M. J., 1907, 1921 Williams, D. E., 783 Williams, G., 503 Woodhouse, J. R., 2507 Woolf, L. A., 1091 Wormald, C. J., 1303, 1315 Woznicka, J., 1709 Wright, J. D., 1979 Wyn-Jones, E., 2099 Yamada, Y., 609 Yao, Z., 2211 Yarwood, J., 1397 Yeates, D., 2641 Yeh, C., 2199 Yeh, C-t., 65 Yoneyama, H., 969, 1873 Yoon, C. S., 2867 Yoshida, N., 1787 Yoshioka, H., 1485 Yoshitake, H., 2021 You, X., 829 Young, D. A., 173 Zaki, M. I., 1723 Zambonin, P. G., 1685 Zanonato, P., 2445 Zaslavsky, B. Yu., 2857 Zecchina, A., 609, 1655, 2113 Zhan, R., 1599 Zielinski, R., 1619 Zukoski IV, C. F., 2785 (viii)~~ FARADAY DIVISION INFORMAL AND GROUP MEETINGS Gas Kinetics Group Developments in Gas Kinetics: New Techniques, Results and their Interpretation To be held at the University of Yorkon 3 4 July 1989 Further infwnation from Professor R.J. Donovan, Depariment of Chemistry, Unimsity of Edinburgh, West Mains Road, Edinburgh EH9 3JJ Industd Physical Chemistry Group with the Thin Films and Surfaces Group of the IOP Materials for Non-linear and Electrooptics To be held at GirtMl College, Cambridge on 6 7 July 1989 Further information from The Meetings officer, Institute of Physics, 47 Belgravle Squam, London SW1 X 8QX Electrochemistry Group with Electroanalytical Group Graduate Students' Meeting To be held at Imperial College, London on 12 July 1989 Further information from Dr G. H. Kelsali, Depatiment of M i d Resources Engineering, Imperial College, London SW7 2BP Pofymer Physics Group Biologically Engineered Polymers 89 To be held at Churchill Cdlege, Cambridge, on 31 July to 2 August 1989 Further infwmation from Dr M.J. Miles, AFRC Institube of Food Research, Colney Lane, Notwich NR4 7UA Carbon Group with surface Reactiity and Catalysis Group Carbons and Catalysis To be held at laughborough Universii of Techndogy on 11-13 September 1989 Further information from Dr J. W. Pafrick, Dimcbr, Carbon Research Group, Loughborough Consultants Lid., UnMtyof Technology, Loughborough LE113TF Polymer Physics Group Biennial Meeting: Physical Aspects of Polymer Science 25th Anniversary To be held at the University of Readng on 13-15 September 1989 Further information from Dr G. R. Mitchell, Poiymer Physics Labomtwy, Un'krsity of Reading, WhiBeknights, Reading RG6 2AF.Colloid and lnterface Science Group Inorganic Particulates TobeheldatChestevCollegeon 1921 Sepaember1989 Further infomation from Dr R. Buscall, ICI plc, Corporats Colbidscience Group, PO Box 11, The Heath, Rw~xrm, Cheshire WA7 4QE Polar Mi& Group with Low Temperature Group of the IOP and the Institute of Ceramics High Temperature Semiconductors To be held at ?he University of Birmingham on 19-21 September 1989 Further infomation from Dr C. Greaves, Department of Chemistry, Universii of Birmingham, P.0. Box 363, Birmingham B15 211 Division with the lnstitute of Physics Sensors and their Applicatlons To be hekl at the Univtmity of Kent at Canterbury on 19-22 September 1989 Further incormation from The Meetings Officer, Institute of Physics, 47 Belgm Squam, London SW1X 8QXDivision with the Deutsche Bunsen Gesellschaft, Division de Chimie Physique of the Socidte Franwise de Chimie and Assodazione Italiana di Chimica F i s h Transport Processes in Fluids and Mobile Phases To be held at the Physacalisdw InstiM, m, West Gwmany on 2528 Seprwnber 1989 Further information from Professor G.LuddKIrst, Department of Chemistry, Universiity of Southampton, Southampton SO9 5NH Division Autumn Meeting: Chemistry at Interfaces To be held at Loughborough University of Techndogy on 26-28 September 1989 Further informalion from Professor F. Wilkfnson, Depaftment of M k t t y , Loughbomgh Uni- of Tdndogy, Loughborough Ell 3TU _ _ _ _ _ ~ Polymer Physics Group Polymers in Motion To be held at the Universiof Cambridgeon 14 December 1989 Further information from Dr M.J. Richadm , National Physical bbombry, Queen's Road, Teddngton, Miesex lw11 OLW Colloid and Interface Science Group Wetting and Spreading To be held atthesCientificsodelies LechrreThealm, London on 15 December 1989 Further information from Dr R. B u d l , ICI plc, Corporate Colloid Science Group, PO Box 11, The Heath, Runcorn, Cheshire WA7 4Q€ Theoretical Chemisw Group Electronic Structure Calculations on Large Molecules: Novel Methods and Applications To be held in Cambndge on 1517 Deoember 1989 Further information from Dr P. Fowler, Department of Chemistry, Uniwxsity of Ex-, Sodw Road, Exeter EX4 4QD H&h Resolution Specboscopv Group with the Mdewlar Beams Group Spectroscopy and Molecular Beams TO be held a! the University of Nottingham on 1819 December 1989 Further information from Dr P.G. Sam, Department of Chemistry, Univlersity of Nottingham, Unkrsity Park, Nottingham NG7 2RD Neutron Scattering Group Applied Neutron Scattering To be heldatthe U n W i o f Warwidcon 21 December 1989 Further information from Dr C. Windsor, B 521.1, ERE, Harwell, Didcot, Oxfordshire Colloid and Interface Science Group with Protein and Peptick Group, Biotechnology Group and Gesellschafr f i r Biologische Chemie Stability of Proteins: Theory and Practice To be heldat Gkaon cdlege, Cambidgem B29 March 1990 Further information from Professor F. Franks, PAFRA Lad., 150 Cambndge Science Park, Cambridge CB4 4GG Division Annual Congress: The Solid State: Reactivity and Electrical Properties To be held at Queen's UniverSi, Belfast on 9-12 Aptj 1990 Further information from Dr J.F. Gibson, The Royal Socjety of Chemistry, Burtingbn House, h d o n W1V OBN Polymer Physics Group New Materials and their Applications To be held at the University of Warwick on 10-1 2 April 1990 Further information from Dr M. Richardson, National Physical Laboratory, Cueen's Road, Teddingbon, M i e s e x lw11 OLWNeutton Scattering Group Neutron Scattering Data Analysis To be M a t t h e Rutherford Appleton LaboraQory, Chilton on 1 May 1990 Further infwmadion from Dr M. W. Johnson, Neutnxl Sdenoe Division, Rutherford Appleton htmratoty, Chilton, OidootOxfordshire Ektmchemistry Group with the SCI Electrosyn thesis To be held at the Uniwsity d southampton on 911 July 1990 Furtherinfcmabm .from Dr D. pledcher, bparimmt of chemistry, University d southampBon, Southampton so9 5NH. Gas Kinetics Group Xlth International Symposium on Gas Kinetics Furtherinfwmabon * from professor R J. Donovan, Department of Chemistry, Uniwdy d Edinburgh, West Mains Road, Edinburgh EH9W TO be held h Assisi, ltaly 2-7 September 1990 ~ ~~~~~~~~~~~~~~~~~~~~~ ~ Colloid and Interface Science Group Conference of the European Colloid and Interface Science Society (EICS) To be held at the UnivF#si of Mstol on 1514 September 1990 Further intormation from Professor R. H. OWwiIl, School of Chemistry, U M i of Bristol. Canbxk's Close, Bristd BS8 1TSNOMENCLATURE 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. Nomenclature. The following publications provide the IUPAC nomenclature rules and guidance on their use: Nomenclature of Organic Chemistry, Sections A, B, C, 0, E, F, and H (Pergamon, Oxford, 1979 edn.) Nomenclature of lnorganic chemistry (Butterworths, London, 1971 , now published by Pergamon). Biochemical Nomenclature and Related Documents (The Biochemical Society, London, 1978). Where there are no IUPAC rules for the naming of particular compounds or authors find difficulty in applying the existing rules, they should seek the advice of the Society’s editorial staff. Units and Symbols. A detailed treatment of units and symbols with specific application to chemistry, based on the Systeme Internationale d’llnites (SI), is given in Quantities, Units and Symbols in Physical Chemistry, published for IUPAC by Blackwell Scientific Publications, Oxford (1 988 edn.). A comprehensive list of IUFAC publications on nomenclature and symbolism appears in the January issue of J. Chem. SOC., faraday Transactions. (xii)

ISSN:0300-9599    

DOI:10.1039/F198985BP117

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1989, 85(9), 2649-2664 Experimental Activity Coefficients in Aqueous Mixed Solutions of KCl and KF at 25 "C compared to Monte Carlo Simulations and Mean Spherical Approximation Calculations? Torben S. Serensen,* Jergen B. Jensen and Peter Sloth Fysisk- Kemisk Institut and Center for Modelering, Ikke-lineare systemers Dynamik og Irreversibel Termodynamik (MIDIT), Technical University of Denmark, Bldg. 206, DK-2800, Lyngby, Denmark A system of three different ions in solution has been studied experimentally and theoretically. Mean molar ionic activity coefficients have been measured in pure and mixed, aqueous solutions of KF and KCI at 25 "C and 1 atm by means of valinomycin, LaF, and Ag/AgCl electrodes. More than 200 independent electrometric measurements were considered.The ionic strength varied from 0.0005 to 4 rnol dm-3. The activity coefficients were closer to unity for pure KF than for KF in equimolar mixture with KCl at the same ionic strength for ionic strengths higher than 1 mol dm-3. The activity coefficients for KCI in pure and mixed solutions could not be statistically separated up to 4 rnol dm-3. The Harned coefficients are estimated to be Ok0.025 dm3 mol-' for KCI and 0.055+0.025 dm3 mol-' for KF. The Debye-Huckel limiting law is obeyed within 1.5% in the region from 0.0005 to 0.01 rnol drnp3, indicating that the ions involved are small. Comparison with calculations for the primitive electrolyte model using the Kirkwood-Buff equations and the generalized DHX theory has shown, that the expcrimental data are approximately fitted using diameters of 2.9, 2.9 and 3.4A for the K+, C1-, and F- ions, respectively.The same values fit approximately the data in the mean spherical approximation (MSA). The MSA calculations demonstrate the validity of Harned's rule. From the latter theory we also obtain single-ion activity coefficients. Monte Carlo (MC) simulations of single-ion activity coefficients have been performed for the primitive, electrolyte model for the above-mentione$ ionic diameters and in addition for a diameter of the F- ion equal to 3.7 A. Widom's test-particle method is used in conjunction with the extrapolation procedure suggested by Sloth and Smensen. The MSA as well as the MC calculations support earlier suggestions, that the single-ion activity coefficient of F- in KF-KCI mixed solutions is almost independent of the salt ratio.The same is approximately torue for the C1- ion. Increasing the diameter of the F- ion from 3.4 to 3.7 A does not alter this conclusion, but the separation between single-ion activity coefficients becomes larger. Bagg and Rechnitzl found the interesting experimental result, that the single-ion activity coefficient of the fluoride ion for mixtures of trace concentrations of NaF in NaCl up to 1 rnol kg-' or of trace concentrations of KF in KX solutions up to 4 mol kg-' (X = C1, Br, and I) were the same as in pure NaF of KF solutions at the same ionic strength. Similar results were also obtained by Leyendekkers,* who studied trace activities of F- in NaF-NaCI and NaF-KCl mixed solutions. Although such electrochemical measurements of single-ion activities based on the use of salt bridges are always t Paper presented at the Third International IUPAC Symposium on Solubility Phenomena.held at the University of Surrey, 23-26 August, 1988. 2649valinomycin membrane aq. KCl-KF solution with electrode (K+) KCl fraction = X and ionic The valinomycin and LaF, electrodes are both close approximations to ideal ion- selective membranes for K+ and F- ions, respectively.1*2*19.20 We have used the electrodes F23 12K and F1052F manufactured by Radiometer A/S. The Ag/AgCl electrodes were high-precision electrodes ( & 10 pV) of the thermal-electrolytic kindly lent to us from the pH-calibration department of Radiometer A/S. A detailed description of the performance of these electrodes have been given elsewhere.'.21 The cells were built up as a water-jacketed Pyrex glass cell with lid. Through holes in the lid, the electrodes were immersed in ca. 100 cm3 of the solution. The temperature was kept at 25.00+0.05 "C by means of circulating water from a thermostat. The e.m.f. values were registered on a pH-meter (PHM64, Radiometer A/S) with a precision of t 1 atm = 101 325 Pa. Ag/AgClT. S. Sorensen, J. B. Jensen and P. Sloth 265 1 k 0.1 mV [old data from ref. (8)] and on a Radiometer PHM84 with the same measuring precision (new data). The potentials were followed on an REC 80 recorder from Radiometer, until constant potentials were obtained (within 0.1 mV). Normally, that took, 2-4 h depending on the concentration.The old data' were grouped in 15 different series with KC1 fractions X = 0, 0.1,0.2, 0.5, 0.8, 0.9 and 1. In each series, the ionic strength was varied from 0.001 to 0.2 mol dm-,. Additional measurements were performed in more concentrated solutions up to 4 mol dm-3 for X = 0 and 1 (pure salts) and up to 2 mol dmP3 for X = 0.5. Each series was prepared and selected at random, and measurements were made independently. In ref. (8), standard potential differences and Debye-Huckel limiting slopes were determined for each series from polynomial fits of the e.m.f. values (corrected for the simple Nernst concentration terms) against h. There seemed to be no systematic variation of the standard potentials or the limiting slopes with X , which is a test of the lack of interference of the foreign ions on each of the three ion-selective electrodes.However, we found a slow variation of the normal potentials of the valinomycin and LaF, electrodes over a timescale of several months.' Thus, it is necessary to find new normal potentials after a few experimental series. The activity coefficients have to be calculated with the 'local ' normal potential difference for the series. Afterwards, however, activity coefficients may be pooled together in order to improve the precision of the final mean data. The procedure followed in order to find the normal potential of a given experimental series was the same in the new experiments as described earlier. The e.m.f. values of the cells were corrected for the Nernst concentration term and plotted against the square root of the ionic strength.The most significant polynomials in h at the 95% level of significance were found, and the normal potentials were determined by extrapolation to In the earlier experiments,' the limiting slope for all the series pooled together was statistically indistinguishable from the Debye-Huckel value for the KC1 measurements (cell 1). However, the limiting slope for the KF data (cell 2) was ca. 10% too low. Nevertheless, with the new data added we shall see, that lny+(KF) is also compatible with the limiting law. The old experiments consisted of 56 measurements of y+(KCl) on cell 1, and 46 measurements of y+(KF) on cell 2. The new experiments-added 73 independent measurements on cell (f) and 82 independent measurements on cell (2).The analysis presented here is therefore based upon 129 measurements on cell (1) and 128 on cell (2). The chemicals used were Merck (pro analysis) KF and KCl. They were dried at 110 "C for at least 24 h and then used without further purification. They were cooled in desiccators with silica gel at least 30 min before weighing with a precision *0.05 mg. KF especially is hygroscopic. Solutions with I 2 0.1 mol dm-3 were made independently by weight (k 0.05 mg). The rest was made by careful dilution of stock solutions. The water used was distilled and deionized water with a conductivity < 1.0 pS at 25 "C. The Monte Carlo calculations were exploiting Metropolis sampling on the primitive electrolyte model (charged hard spheres in a dielectric continuum) as described e a ~ l i e r .~ ~ - ~ ~ The new features are that there are one cation and two anions which may all be of different sizes. Furthermore, 64 test particles of each type are inserted in the central box in a regular lattice. The interactions of the test particles with the 'real ions' are averaged in order to calculate the excess chemical potential of each ion according to the method of Wid~m."*'~ The real ions pass freely through the test ions, however. Thus, the presence of the test particles does not disturb the Metropolis Markov chain. In order to extrapolate to an infinite number of particles ( N = a), we have proved that the obtained values for the excess chemical potentials for each ion should be plotted against the cube root of 1 / N and extrapolated linearly to zero.This was proven in ref. (16) (fig. 1 and table 3). fi = 0.2652 Aqueous Mixed Solutions of KCl and KF Statistical Treatment of Experimental Data and MC Data There was a considerable scatter in the lny+ values at very low concentrations. However, if we collect all measuring points for Iny+(KCl) and 0.0005 < I / m ~ l d m - ~ < 0.01 without regard to X (66 data points), only a fir%-9egree1 polynomial in $ is significant at the 95 O/O level with a slope of - 1.22f0.10 dm? mol-?. Similarly, the 70 independent measurements for lny+(KF) in the same concentration range without regard to <gave ? straight line at the 95 % level of significance with a slope equal to - 1.10 0.07 dmz mol-2. Both values are cFmpatjble with the limiting-law value from the Debye-Hiickel theory - 1.1779 dm.mo1-p.26 When all the 136 data points were pooled, without regard to X and to the type of electrolyte, we still obtain:d a stlraight line at the 95 YO level of significance, and the slope was - 1.160 f 0.025 dm? mol-5, which is only ca. 1.5 YO from the limiting-law value. The reason why the limiting law is followed in the whole region from 0.005 to 0.01 mol dmP3 is that the diameters of the hydrated ions for this specific system are extraordinarily small (see later). Thus, there seems to be no reason to believe that the limiting law should not be fulfilled for the present pure and mixed solution:, andlwe proceed by enforcing the initial slope of the lny+ us. fi plots to be - 1.1779 dmz mol-5 for lny+(KCl, X = l), lny+(KCl, X = O S ) , lny,(KF, X = 0) and lny+(KF, X = 0.5) us. fi.- From the data points, we then determine the highest polynomials which are significant at the 95 % level. We use orthogonal Forsythe polynomial^^^ for the ‘ free polynomials ’ for which the coefficient of each new orthogonal polynomial of one higher degree can be tested without any covariance by the Student’s After having determined the significant degree and the error belts of the polynomial, we determine the least-square polynomial of the same degree subject to the restrictions, that it should go through (0, 0), which the free polynomials do within the uncertainty, and that they should have the correct limiting slope. We observed only a very weak statistical tendency at intermediary concentrations for the lny+(KCl) values to move towards the lny+(KF) values with admixture of KF.Therefore, we have preferred to pool together dl data for Iny+(KCl) - for X = 1 and X = 0.5. - We obtained the following results : p = $ ( I in mol dmP3) (3) (0 < p d 2.0) (4) lny+(KCl, - pure and X = 0.5) = - l.l779p+ 1.0958~~-0.532823p3+0.10659p4 Number of experiments: 13 (previous, X = 1) + 21 (new, X = 1) + 17 (previous, X = 0.5) + 52 (new, X = 0.5). lny,(pure KF) = - l.l779p+ 1.5956~~- 1.5179p3+0.78008p4-0.142732~~ (0 < p < 2.0) ( 5 ) Number of experiments : 10 (previous) + 30 (new). lny+(KF, - X = 0.5) = - l.l779p+ 1.0935~~-0.47354p3+0.083062p4(0 < p < 1.9) (6) Number of experiments: 12 (previous) + 52 (new). The results are shown in fig.1-3 together with the best least-square polynomials (full curves) and some Mean Spherical Approximation calculations (dashed lines). Experimental points with I < 0.01 mol dm-3 have been omitted on these figures, since the great amount of low-concentration data cannot be properly represented. However, they have not been dropped from the least-squares calculations. Fig. 4 exhibits the error belts inside which the true (smoothed) values of -In y+ - (KF)T. S. Sorensen. J. B. Jensen and P. Sloth 0.6 0.7k A 2653 I " " " " " " " ' * " ' ~ . p 0.1 0.5 1.0 1.5 2.0 I /mol* dm- 2 Fig. 1. The negative natural logarithm of the mean ionic, molar activity coefficient of KCI, -Iny+ ure and mixed aqueous solutions of KCI us. the square root of the molar ionic strength, (KC1), at 25 "C and atmospheric pressure.a, fraction of KCI = X = 1. A, X = 0.5. ., coincident data points (X = 1 and 0.5). Solid curye: polynomial [eqn (4)]. Dashed curve: MSA calculation for X = 1 with diameters 2.9 and 2.9 A for K' and C1-. All measurements below the vertical dashed line have been omitted from the figure, but not in calculating the least-square polynomial [eqn (4)]. 1 1 1 Fig. 2. Aqueous solutions of KF at 25 "C and atmospheric pressure. Salt fraction of K F = X,, = 1 . Plot of -Iny+(KF) us. fi. Filled circles: experiments. All experimental points below vertical dashed line have been omitted from the plot, but not in calculating the least-square polynomial [eqn (5)]. Solid curve: polynomial [eqn (5)].oDashed curve : MSA calculation with diameters 3.4 and 2.9A for F- and K'. lie within a probability of 66 % for pure KF as yell as for equimolar KCl-KF mixtures.The two curves clearly separate above = 1 molTdmi, and the separation at higher ionic strengths is much more pronounced and unequivocal than for KCl. The KF curves move in the direction of the combined KCl curve with admixture of KC1. The Monte Carlo calculations were also subjected to statistical analysis. The values of the negative logarithms of the single-ion activity coefficients (- lny,) were plotted against the inverse cube root of the number of particles. In all cases investigated, straight lines only were significant at the 95% significance level.2654 Aqueous Mixed Solutions of KCl and KF Fig. 3. Aqueous solutions of equimolar mixtures of KF and KCI (X = 0.5) at 25 "C and atmospheric pressure.Plot of -In y,(KF) us. k. Filled circles : experiments. All experimental values below the vertical dashed line have been omitted from the plot, but not in calculating the least-square polynomial [eqn (ti)]. Solid curve : polynomial [eqn (6)]. 0.7 i 1 1 3 1 2 /mol! dm- 2 Fig. 4. Plots of the smoothed polynomial values of -Iny+(KF) us. b for (a) pure (XKp = 1 ) and (b) mixed (XKV = 0.5) aqueous KF solutions at 25 "C and atmospheric pressure. The error belts are shown, within which the 'true' values of -Iny+(KF) lie with 66% probability. A clear separation of the two curves is seen- above I = 1 mol dm-3. MSA and MC Calculations From fig. 1 and 2 we learn, that the experimental data for pure KCl and pure KF solutioons are fitted quite well by MSA calculations assuming a contact distance equal to 2.90 A between K+ and C1- and 3.15 A between K' and F-.The fit is valid up to ca. 2 mol dm-3. Furthermore, an equal division of the 'primitive model effective ionic radii' has been assumed for K+ and C1-, since the limiting conductivities of those two ions areT. S. Sorensen, J. B. Jensen and P . Sloth 2655 I ' I I I I I I I I I Fig. 5. Plots of -Iny+(KCI) [(a), (b) and (e)] and of -Iny+(KF) [(c), ( d ) and (f)] us. A',, and XKc,, respectively, at constant total ionic strength (4, 1 rnol dm-3; B, 0.5 rnol dmP3) calculated by the MSA (solid lines). Bjerrum length &, = 7.135 A (water, 25 "C), diameters 2.9, 2.9 and 3.4 A for K', C1- and F-, respectively. Dashed lines are similar M C test particle calculations for a total ionic strength equal to 1 mol dm-3.The MSA calculations exhibit clearly the Harned linearity of the 'artificial' KCl-KF system. The MSA calculations, as well as the MC calculations, verify the empirical rule of thumb, that the two trace activity coefficients are nearly, but not completely, identical. approximately equal. Thus, the effective ionic diameters are 2.9, 2.9 and 3 . 4 i for K+, C1- and F-, respectively. Next, we consider mixtures. We want to test Harned's rule' which can be written in the form lny+(KCl, - fixed I ) = lny,(pure KCl, Z)-aKc.(l - X ) (7) lny+(KF, - fixed I ) = lny,(pure KF, I)-a,,X. (8) The (non-dimensionalized) Harned coefficients are aKCl and aKF. As first derived by Glueckauf er al.,29 the Harned coefficients with dimension dm3mol-l (or kg solvent mol-') may vary with ionic strength, but their sum must be independent of ionic strength in order for the system of equations (7) and (8) to satisfy the Maxwell equations of cross- differentiation of free energy.This means, that the sum of the non-dimensional Harned coefficients must be proportional to the total ionic strength: aKCl + aKF = constant x I. (9) Fig. 5 and 6 demonstrate the Harned linearity of MSA calculations of Iny+ for mixtures of KCl and K F at constant ionic strength. The strict linearity is observed-at total ionic strengths equal to 0.5, 1 .O, 1.5 and 2.0 mol dmP3. In all cases, the logarithms of the trace activity coefficients of one salt in great excess of the other tend almost to meet each other, approximately in the middle between the values for the two pure salts.The tendency for trace activity coefficients to be almost, but not exactly, equal has long been known as an empirical rule of thumb for many electrolyte mixture^.^2656 Aqueous Mixed Solutions of KCl and KF 0.50 0.40 I I I 1 I 1 I I I I 0 0.2 0.4 0.6 0.8 1 .o XKCI ~ ~ X K F Fig. 6. Plots of -lny+(KCl) [(a) and (b)] and of -Iny+(KF) [(c) and ( d ) ] us. A', and XKC,, respectively, at constant total ionic strength (A, solid linei: 1.5 mol dmP3; B, dashed lines: 2 mol dm-:) calculated by the MSA. Bjerrum length AH = 7.135 A (water, 25 "C), diameters 2.9,2.9 and 3.4 A for K+, C1- and F-, respectively. The Harned linearity is clearly exhibited, and the two trace activity coefficients are nearly, but not completely, identical.Table 1. MSA values for the dimensionless Harned coefficients for KF-KCl" total I/mol dmP3 &KCl ~ K F ~ K C I +a,, 0.5 - 0.0294 0.304 0.00 10 1 .o - 0.0268 0.0280 0.00 12 1.0 (MCb) -0.050 & 0.0 1 1 0.008 f 0.01 7 - 0.042 & 0.020 1.5 - 0.0256 0.0270 0.0014 2.0 -0.0255 0.0260 0.0005 " Diametets: 2.9,2.9 and 3.4 A for K', C1- and F-, respectively; Bjerrum length: AR = e:/(4nekT) = 7.1355 A. MC values calculated from table 2. Table 1 shows calculated dimensionless Harned coefficitpts from the MSA theory. All calculations are performed for di9meters 2.9,2.9 and 3.4 A for K+, C1- and F- and with a Bjerrum length equal to 7.135 A (ca. 25 "C in water). The Bjerrum length is given by: AB = ei/4nek, T (10) where e, is the elementary charge, k, is Boltzmann's constant, T the absolute temperature and E the dielectric permittivity of the solvent.From table 1, it is obvious, that the dimensionless Harned coefficients vary slowly with the total ionic strength. The sum (last column) does not seem to be proportional to the total ionic strength. However, the sum is only 2-5 O/O of the value of the individual Harned coefficients. Thus, it might be argued that the sum should really be zero if not for some inevitable lack of consistency in the MSA approximation. Then, the data are compatible with eqn (9) with the constant = 0. In order to check the MSA calculations, we have also made a direct determination of the single-ion activities in a 1 mol dmW3 mixture of KC1 and KF by Monte Carlo simulations using Widom's test-particle formalism.W? have made two sets of calculations, one with the diameter of F- set equal to 3.7 A and one with the diameterT. S. Smensen, J. B. Jensen and P. Sloth 2657 Table 2. Monte Carlo results for lnyi for KF-KCl mixtures with dF = 3.4Au Nb N(K+) N(CI-) N(F-) -Iny, -InY,, - In Y , no. config./ lo6 32 16 64 32 80 40 100 50 150 75 216 108 350 175 512 256 1000 500 c o - 100 50 140 70 160 80 216 108 360 180 512 256 c o - 64 32 80 40 100 50 130 65 150 75 216 108 350 175 512 256 c o - 0 X = 0, pure KF, d, = 2.9 A 0 16 0.2096 - 0. I950 0 32 0.2530 - 0.2435 0 40 0.2709 0.2566 0 50 0.2763 - 0.2686 0.291 1 0 75 0.2986 - 0 108 0.3 141 0.3 127 0 175 0.3421 0.3161 0 256 0.3682 - 0.3513 0 500 0.3679 - 0.3544 0.45 13 - 0.4345 - + 0.0079 0.0072 - - - - - X = 0.5, KF : KCl = 1.1 , dK = 2.9 A, d,, = 2.9 A 25 25 0.308 1 0.3404 0.2742 35 35 0331 1 0.3597 0.295 1 40 40 0.3340 0.3657 0.2978 54 54 0.347 1 0.3743 0.3109 90 90 0.3565 0.403 1 0.3333 128 128 0.3827 0.4047 0.3345 - - 0.469 1 0.501 3 0.4255 f 0.0 134 f 0.0095 f 0.0089 0 0 X = 1, pure KCl, d, = 2.9 A, d,, = 2.9 A 32 40 50 65 75 108 175 256 - 0 0 0 0 0 0 0 0 - 0.3082 03093 0.3237 0.3259 0.3379 0.3367 0.3495 0.3509 0.3562 0.3576 0.3723 0.3764 0.3977 0.3965 0.4 1 73 0.4032 0.5101 k0.0037 1 .o 0.5 0.5 0.5 0.5 0.5 0.5 0.67 0.5 - 0.5 0.5 0.5 0.5 0.5 0.5 - - 0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.5 - a Total ionic strength, I = 1 mol dm-3; Bjerrum length, N = co by linear regression us.cube root of I / N . set to 3.4 A. In both cases, the diameters of K+ and C1- 0 EL, = 7.1355 A.Extrapolations to were set to 2.9 A and the Bjerrum length to 7.135 A. The former set was made to have a better separation between the activity coefficients and to test the sensitivity of the conclusions drawn to changes in the diameter of the larger ion. Tables 2 and 3 show the result of the MC simulations for differential total numbers of ions ( N ) in the central box. As proven in two earlier ~ a p e r s , ~ ~ . ~ ~ the extrapolation to the thermodynamic limit should be made by means of straight-line regressions of lny, us. the cube root of (l/N). Examples are shown in fig. 7. Fig. 8 and 9 show the 'spectra' of lny, and lny, obtained from the MC simulations for the two sets at 1 mol dm-3 and ca. 25 "C. For comparison, the MSA values for the single-ion activity coefficients are also shown.The correspondence between the two methods is quite good, and the most important feature is, that not only is lny, approximately independent of composition (even to the limit of trace activities), but the same feature is exhibited by lny,,. Furthermore, lny, seems to follow the dominant anion in the mixture, but the K+ ion has to strike a compromise in the case of an equal2658 Aqueous Mixed Solutions of KCI and K F Table 3. Monte Carlo results for Iny, for KF-KCI mixtures with d , = 3.7 A - 80 100 150 216 3 50 512 1000 co 130 260 390 520 650 so 64 80 100 I40 160 216 360 512 co I04 I30 260 390 520 650 M 40 50 75 108 175 256 500 - 65 I30 195 260 325 - 32 40 50 70 80 108 180 256 - 52 65 I30 195 260 325 - X = 0, pure KF, d , = 2.9 A - - 40 0.23 10 0.21 13 50 0.2544 0.2240 - 75 0.2637 0.2477 - I08 0.2865 0.2647 I75 0.2988 0.2896 - 256 0.3095 0.3064 - 500 0.3456 0.3289 - .- - 0.4 149 0.4 174 +0.0104 & 0.0028 - - - - - - - __ - X = 1/13, KF:KCI = 12: 1, d , = 2.9A, d(,l = 2.9A 5 60 0.2678 0.3521 0.238 I 10 120 0.3007 0.3846 0.2670 15 180 0.3 148 0.4020 0.2930 20 240 0.3206 04223 0.2952 25 300 0.3438 0.4327 0.3208 - 0.4327 0.5409 0.4203 k0.0152 +0.0119 Ifr0.0186 - 0 0 X=0.5,KF:KCI= l : l , d K = 2 .9 A , d ( . , = 2 . 9 A 16 16 0.2690 0.308 1 0.20 18 20 20 0.2786 0.3248 0.2171 25 25 0.2940 0.3406 0.2299 35 35 0.3 156 0.35 1 I 0.2434 0.2546 40 40 0.3 192 0.3671 54 54 0.3274 0.3748 0.2749 90 90 0.36 I6 0.4034 0.2970 I28 128 0.3661 0.4247 0.303 1 0.4692 0.5299 0.41 12 k 0.007 I k 0.0082 k 0.0060 - - 0 0 X = 12/13, KF:KCI = 1 12, d , = 2.9 A, d,., = 2.9 A 48 4 0.3292 0.3398 0.2345 60 5 0.3455 0.349 1 0.241 5 I20 10 0.3759 0.3956 0.2879 180 15 0.3874 0.4024 0.2972 240 20 0.4042 0.4194 0.3072 0.3 160 300 25 0.41 75 0.4 I22 .- 0.5097 0.5 168 0.4161 k 0.0082 0.01 30 f 0.0077 - no.config./ 10" 0.5 0.5 0.5 0.5 0.5 0.5 0.5 __ 0.78 0.5 0.5 0.5 0.5 - 0.5 0.64 0.6 0.5 0.55 0.5 0.5 0.5 - 1.535 0.5 0.5 0.5 0.5 0.5 - a Total ionic strength, I = 1 mol dm-3; Bjerrum length, & = 7.1355 A. by linear regression us. cube root of l / N . Extrapolation to N = M mixture of the two salts. These conclusions seem to be independent of the precise magnitude of the larger ion. Table 4 exhibits a comparison between mean ionic activity coefficients and Harned coefficients for the ionic system with d , = 3.7 A calculated by MC and MSA.It is observed, that although the correspondence between MSA and MC is quite good for the mean activity coefficients, the values are not exactly the same, and the Harned coefficients are very sensitive to the method of calculation.T. S. Sorensen, J . B. Jensen and P. Sloth 0 0 0 0 0 2659 0.05 0.10 0.15 0.20 0.25 Fig. 7. MonJe Carlo test particle simulations. Bjerrum length AB = 7.135 A; ionic diameters 2.9, 2.9 and 3.7 A for K', C1- and F-. Total ionic strength = 1 mol dmP3. Linear plots of the negative natural logarithm of the single ion activity coefficients of K+, C1- and F- us. the inverse cube root of the total number of particles in the simulation.The extrapolated values are close to the MSA values. At least one half million of Metropolis configurations per simulation. Periodic boundary conditions and minimum image cut-off of the configurational energies. 1 N - 3 Table 4. Mean activity and Harned coefficients from Monte Carlo and MSA" - - - 0 - 1 -0.487f0.011 -0.4734 -0.025+0.012 -0.0444 13 $ - 0.4995 & 0.0054 - 0.4923 -0.021 f0.013 -0.0442 - 13 l2 -0.5133f0.0056 -0.51 10 0.042 f 0.087 - 0.0442 1 -0.5101 f0.0037 -0.5144 - - - 0 -0.4161 f0.0040 -0.4152 - - 1 0.13 0.21 0.048 1 - 13 -0.426f 0.016 - 0.41 89 2 - 0.4402 0.0060 - 0.439 1 0.048f0.014 0.0478 0.0507 & 0.0069 0.0475 - l 2 -0.4629 f 0.0049 - 0.4590 1.3 1 a d, = d,, = 2.9 A, d, = 3.7 A.0.38 0.40 0.42 0.44 0.16 0.46 0.50 0.52 0.54 0.56 0.58 Fig.8. 'Spectrum' of negative natural logarithms of activity coefficients obtained by MC test particle simulations (66 % error belts OF extrapolated values) and by the MSA method (dashed lines) with a 'large' F- ion (diameter 3.7 A) at 1 mol dmP3 total ionic strength and 25 "C. The single- ion activity coefficients of F- and C1- are almost independent of the composition of the mixtures, even close to trace conditions. In solutions with two ions (perhaps with a trace of a third), the activity coefficients of the two ions are almost identical also for unequally sized ions. For an 0.40 0.45 0.50 1 I 1 - 1 1 ' 1 I 1 I 1 I 1 - I < ICn I I I - - 0 1 - 1 -!- ln * I1 - G; s; I ' V U l I U 31 I 1 - , 2 2 r I I I b zi mI I m=l-l h l N 0 q a. 2' K t r; + I Yl II 1; II + I GYg a" I I I I 1 ,KCI I I 1.T F- K + mi ., K l , . ,. 2!, , t I I I KCI 'a -1: $ $2 E $ 3 L; 'x 11' pzzi7 I I - 11- - 1 I I I I I I I I I 1 , $0 I 0.40 0.45 0.50 Fig. 9. Similar 'spectrum' as in fig. 8, but with a smaller F- ion (3.4 A). The qualitative features are the same as in fig. 8, but the separations of the values are less pronounced.T. S. Sorensen, J. B. Jensen and P. Sloth 266 1 Discussion It is far from evident, that the primitive model is a good model for electrolytes having concentrations of the order of 1 mol drnp3. One might anticipate, that the structure of the solvent, specific ionic interactions, dielectric saturation and lowering of the dielectric constant with increased salt concentration all contribute to modify the results given here.Furthermore, as a McMillan-Mayer theory, the primitive model yields activity coefficients at the given osmotic pressure rather than at atmospheric pressure, which leads to a significant correction at concentrations higher than ca. 1 mol dmP3. However, a great amount of ignorance as to the precise correction to carry out would be manifest, if one tried to perform all these corrections. As an alternative, we can follow the ‘Golden Rule of Physical Chemistry’ stating that, in complex systems, it is better in comparison with experiments to ignore the sophisticated features in models rather than to correct for a part of these features in a more or less rigorous manner. This is so, since the ignored corrections very often tend to counteract each other, so that they partly cancel.In this spirit, Ebeling and Scherwinski14 demonstrated recently, that the primitive model by means of simple MSA calculations, was very efficient in mimicking the real behaviour of 1 : 1 electrolytes (alkali-metal halides) in water at 25 “C up to at least 1 mol dm-3. The effective ionic radii seemed to be quite realistic when compared to crystallographic radii and to the expected amount of hydration of the different ions. Similar results were obtained much earlier by Triolo et al.30 However, the contact distances quoted in the two above-mentioned papers differ considerably, and we think that the distances from the latter paper are less reliable, since a fit was made up to 2 mol dmP3 without any corrections, and since the less exact Percus-Yevick equation was used for the hard-sphere contribution.The inaccuracies also showed up in a bad fit at moderate concentrations. In contrast, Ebeling and Scherwinski fitted only up to 1 mol dm-3 and used a generalized Carnahan-Starling approximation for unequal hard spheres. Ebeling and Scherwinski attempted a least-squares determination of the individual ionic radii by fitting to mean ionic activities of all combinations of Li+, Na+, K+, Rb+ and Cs+ with C1-, Br- and I-. This ought to be an impossible task, and so it proved, since the authors had to fix one of the 8 radii ‘to obtain a stable fit’. Finally, the extra- thermodynamic assumption made was that the iodide ion was not hydrated, and that the effective radius of this ion was equal to the Pauling radius.A table of 8 radii was then constructed, which yields approximately correct activity coefficient curves for 15 alkali- metal halides. The radii for I$+ and for C1- were found to be 1.36 * 0.04and 1.78 * 0.04 A, where we have used 1.45 A for both of these ions. We think it is more natural to ascribe similar radii to K+ and C1- in the framework of the primitive model, since the limiting conductivity in water is almost equal for these two ions. In a primitive model, only the size of the ion, the absolute value of the ionic charge and the viscosity of the surrounding medium should be of importance for the conductivity. In addition, one might envisage a considerable ‘dielectric drag’ on such small ions due to the finite relaxation time of the reorientation of the water m01ecules.~~-~~ However, this drag is also equal for equally sized and charged cations and anions. Th: difference in the contact distance between K’ and C1- (3.14 A compared to our 2.90 A) may be ascribed to some difference in the underlying experimental data.Ebeling and Scherwinski have used the compilations of Hamer and WU.~’ We discussed in ref. (8) the discrepancy between our data for KCl and the data compiled by Hamer and Wu (transformed from molal to molar values) and concluded, that our data for KC1 should be considered to be more reliable, even with the reduced number of data points in ref. (8) compared to the present study. This was so, because the majority of data in the Hamer and Wu compilation was made with cells with liquid junctions, and because the extended Debye-Hiickel equation was used in the ‘dilute end’ of the cell to calculate2662 Aqueous Mixed Solutions of KCI and KF Iny+(KCl) in the ‘concentrated end’.The great increase in the number of experiments made in this paper has not changed the mean values of Iny+(KCl) - very much, but the uncertainty belts have shrinked considerably. Other primitive model values for the contact distance in KCI solutions have been stated by Triolo et al. (osmotic coefficients at 25 0C),3(’ by Moller (osmotic coefficients at 0 “C), ref. (38) fig. I 15, p. 41 7, and by Onsager and Fuoss (concentration dependence gf conductance data), ref. (38), p. 451. The values obtained were 2.80, ca. 2.0 and 3.07 A, rFspectively. If we disregard the very low value from the theory of Moller, the value 2.9 A obtained in this study is quite compatible with the other values.It is interesting to compare the contact distance obtained for KCI in aqueous solutions to the ion-ion distance calculated from crystals. Adams has given a survey over ion hard- sphere dimensions estimated in various ways from solid-state structural ~hemistry.~’ The best evaluations are probably derived from electron density maps for ionic crystaJs obtained by X-ray diffraction. The radius of the K+ and the C1- ions are i.49 and 1.64 A, respectively. Thus, the ‘contact distance’ without hydration is 3. I3 Ad This is very close to the value in solution given by Ebeling and Scherwinski (3.14 A). Thus, KC1 seems from such studies to be little or very weakly hydrated, but this might be the result of cancellation of different effects neglected in the primitive model.One should note, however, that it is quite possible to have a contact distance between two unhydrated ions in solution which is somewhat less than the distance in crystals, since the ion pairs in crystals are pulled apart by their other neighbours. Fo! KF, we fit a contact distance equal to ca. 3.15 A. This is not far from the value 3.08 A obtained by Triolo et al.30 from osmotic coefficients. Thus, the effective size of the F- ion in aqueous solution seems to be greater than the size of the C1- ion. The contact dista5ce without hydration in a KF crystal estimated from electron density maps is only 2.65 A.39 Thus, the only plausible explanation for the difference seems to be, that the F- ion is ?trongly hydrated.This is not surprising, when one takes the small size of this ion (1.16 A from electron density maps) and its hydrogen-bond forming ability into account. We have not found any significant deviation between MSA calculations and MC test particle simulations at I mol dm-3 using the same radii. Also, the qualitative behaviour of mixtures of KF and KCI at 1 mol dm-3 total ionic strength is the same (fig. 5). The quantitative relations (Harned coefficients) are quite different, however. The conclusion is that MSA is not quite good enough to calculate activity coefficients in mixtures, at least in cases with very small differences between the lny+ us. h curves of the two salts, see table 1 . The rule of thumb, mentioned in the monograph of Robinson and Stokes,9 that the trace activity coefficient of one salt in excess of another almost has the same value as the trace activity coefficient of the second salt in excess of the first, seems to be verified by the MSA calculations as well as the MC calculations, see fig.5, 6, 8 and 9. However, we also obtain a deeper look into the variations of the single-ion activity coefficients with the degree of mixture between two salts with a common ion. The single-ion activity coefficients of the two counterions to the common-ion (Cl- and F-) seem to be very nearly independent of the salt fraction at constant ionic strength even to the limit of trace conditions. The activity coefficient for K+ at I mol dm-3 is very near to the activity coefficient of the counter-ion in purc solutions evenjor differing ionic radii (KF with diameter of the F- ion equal to 3.4 A as well as 3.7 A).However, in a 1 : I mixture the activity coefficient of K’ has to strike a compromise between the single-ion activity coefficients of the two counterions. Passing to the trace limit, the activity coefficient of the K+ ion moves towards the activity of the counterion in excess. The two trace activities will be identical, if the single-ion activities of the two counterions are unaffected by the mixing, and if the single-ion activities of the two ions in a pure (or nearly pure) solution are the same. The MC simulations with diameters 2.9, 2.9 and 3.4A for K+, C1- and F- (fig. 5,T. S. Surensen, J . B. Jensen and P. Sloth 2663 Table 5.Mean activity and Harned coefficients from experimental KF data ~ - Iny+(pure KF) -lny+(KF, X = 0.5) I/mol dm-3 [polynomial, eqn (5)] [polynomial, eqn (6)] ~ K F - 1 .o 0.463&0.010 0.475 t0.010 0.024 & 0.028 1.5 0.476 0.01 0 0.485f0.010 0.0 19 & 0.028 2.0 0.455 &- 0.01 5 0.486 & 0.01 5 0.062 f 0.042 3 .O 0.345 f 0.025 0.473 k 0.020 0.256 & 0.064 4.0 0.203 & 0.030 0.441 f 0.020 0.477 k 0.072 _ _ _ _ _ _ ~ dashed curves) show some variation in the mean ionic activity coefficients with the salt fraction. The values of lny+(KCl) should vary more than the values of lny+(KF). Contrarily, the experiments-seem to indicate, that In y+(KCl) is almost unaffected, whereas there is a clear variation in lny+(KF) (fig. 1 4 ) . Thus, either we have not yet hit the best effective ionic radii, or the physkal properties of real mixtures are too subtle to be completely explained by primitive model calculations. One might do a lot of fitting with the MSA model, but in view of the different Harned coefficients obtained from MSA and from MC, this does not seem to be worth while.We cannot test the Harned linearity in the experimental data, since we have only measured pure salts and 1 : 1 mixtures. However, the MSA calculations as well as the MC simulations exhibit clear-cut Harned linearity. Thus, assuming Harned's rule to hold, we may calculate the dimensionless Harned coefficient czKF at various total ionic strengths (table 5). The Harned coefficient aKC, we have assumed to be zero at all ionic strengths by pooling together the data for lny,(pure KCl) and lny+(KCl, X = 0.5).The Harned coefficients aKF found have a considerable uncertainty. They cannot really be statistically distinguished from zero before above 2 mol dmV3. In order for the Maxwell conditions to be fulfilled, a,, should be strictly proportional to I when aKCl is zero, see eqn (9). The experimental data are not in conflict with such a proportionality, since there is a considerable and significant increase in aKF in passing from I = 2 mol dm-3 to 4 mol dmP3. The proportionality constant is 0.055 0.025 dm3 mol-1 when estimated as the average of the values at the five concentrations stated in table 5 . Thus, from the present study we extract the following estimates for the Harned coefficients: a K c , / I = 0 0.025 dm3 mol-1 (1 1) aKF/I = 0.055 0.025 dm3 mol-'.(12) The uncertainties on the two Harned coefficients have been taken as equal. With the above very large uncertainties, the experimental Harned coefficients are not completely incompatible with the MSA Harned coefficients (table 1). (The MSA Harned coefficients should be preferred to MC Harned coefficients, since experimental data have been fitted to MSA curves and not to MC curves.) To measure Harned coefficients in the KF-KCl system with a precision significantly better than obtained in this paper would require at least 1000 independent data points, since the precision of each single measurement cannot be improved with the present electrodes. We thank Teknologistyrelsen (The Danish National Agency of Technology) for economic support for the present study.We are grateful to Dr Hans Bjarne Kristensen and Dr Hans Boye Nielsen, the pH Calibration Department, Radiometer A/S, for the delivery of high-precision Ag-AgC1 electrodes.2664 Aqueous Mixed Solutions of KC1 and KF References 1 J. Bagg and G. A. Rechnitz, Anal. Chem., 1973, 45, 1069. 2 J. V. Leyendekkers, Anal. Chem., 1971, 43, 1835. 3 T. S. Serrensen and K. F. Jensen, J. Chem. Soc., Faraday Trans. 2, 1975, 71, 1805. 4 N. A. 0sterberg, J. B. Jensen and T. S. Serrensen, Acta Chem. Scand., Part A , 1978, 32, 721. 5 N. 0. Osterberg, J. B. Jensen, T. S. Serrensen and L. D. Caspersen, Actu Chem. Scund., P u t A , 1980, 6 N. 0. Osterberg, T. S. Serrensen and J. B. Jensen, J . Electroanal. Chem., 198 I , 119, 93.7 A. K . Covington, Ion Selective Electrodes, 1983, 5, 93. 8 J. B. Jensen, M. Jaskula and T . S . Serrensen, Acta Chem. Scand., Part A, 1987, 41, 461. 9 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 2nd edn, 1965), chapt. 32, 721. 15. 10 J. G. Kirkwood and F. P. Buff, J . Chem. Phys., 1951, 19, 774. 11 T. S. Serrensen and J. B. Jensen, Acta Chem. Scand., in press. 12 E. Waisman and J. L. Lebowitz, J . Chem. Phys., 1972, 56, 3086. 13 L. Blum, in Theoretical Chemistry; Advances and Perspectives (Academic Press, New York, 1980), vol. 14 W. Ebeling and K . Scherwinski, 2. Phys. Chem. (Leipzig), 1983, 264, 1. 15 P. Sloth and T. S . Srarensen, Chem. Phys. Lett., 1988, 143, 140. 16 P. Sloth and T . S . Ssrensen, Chem. Phys. Lett., 1988, 146, 452. 17 B. Widom, J . Chem. Phys., 1963, 39, 2808. 18 B. Widom, J . Phys. Chem., 1982, 86, 869. 19 R. D. Armstrong, J. C. Lockhart and M. Todd, Electrochim. Acta, 1986, 31, 591. 20 R. D. Armstrong and M. Todd, Electrochim. Acta, 1986, 31, 1413. 21 T. S. Serrensen and J. B. Jensen, J. Nonequilibr. Thermodyn., 1984, 9, 1. 22 N. Metropolis, A. W. Rosenbluth, M. N. Rosenbluth, A. H. Teller and E. Teller, J . Chem. Phys., 1953, 23 D. N. Card and J-P. Valleau, J . Chem. Phys., 1970, 52, 6232. 24 P. Sloth, T. S. Ssrensen and J. B. Jensen, J . Chem. Soc., Faraday Trans. 2, 1987, 83, 881. 25 T. S. Ssrensen, P. Sloth, H. B. Nielsen and J. B. Jensen, Acta Chem. Scand., Part A , 1988, 42, 237. 26 J. S. Newman, Electrochemical Systems (Prentice-Hall, Englewood Cliffs, N.J., 1973) table 28-1, p. 84. 27 G. E. Forsythe, J. Soc. Ind. Appl. Math., 1957, 5, 75. 28 T. S. Ssrensen and P. Schack, in Analysis and Simulation of Biochemical Systems, ed. H. C. Hemker 29 E. Glueckauf, H. A. C. McKay and A. R. Mathieson, J . Chem. Soc., 1949, 299. 30 R. Triolo, J. R. Grigera and L. Blum, J . Phys. Chem., 1976, 80, 1858. 31 M. Born, 2. Phys., 1920, 1, 221. 32 R. M. Fuoss, Proc. Natl Acad. Sci. U.S.A., 1959, 45, 807. 33 R. H. Boyd, J. Chem. Phys., 1961, 35, 1281. 34 R. Zwanzig, J . Chem. Phys., 1963, 38, 1603. 35 T. S. Serrensen, Acta Chem. Scand., Part A , 1978, 32, 571. 36 T. S. Serrensen, Acta Chem. Scand., Part A , 1979, 33, 583. 37 J. W. Hamer and Y-C. Wu, J . Phys. Chem. ReJ Data, 1972, 1, 1047. 38 H. Falkenhagen and W. Ebeling, Theorie der Elektrolyte ( S . Hirzel Verlag, Leipzig, 1971). 39 D. M. Adams, Inorganic Solids. An Introduction to Concepts in Solid-State Structural Chemistry. (John 5 , pp. 1 4 6 . 21, 1087. and B. Hess (Elsevier, Amsterdam, 1972). Wiley & Sons, London, New York, Sidney, Toronto, 1974), chapt. 2, table 6. Paper 8/04446E; Received 4th November, 1988

ISSN:0300-9599    

DOI:10.1039/F19898502649

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1989, 85(9), 2665-2668 A Phase Separation caused by the Solubility of Butane in 2-Me thylpropan- 2- ol-Wa ter Mix t urest Robert W. Cargill* and Donald E. MacPhee Department of Molecular and Life Sciences, Dundee Institute of Technology, Bell Street, Dundee DDl 1HG The solubility of butane and propane at a partial pressure of 101.3 kPa has been measured at 278-335 K in a solvent consisting of water and t-butyl alcohol (2-methylpropan-2-01), of composition up to 0.08 mole fraction t-butyl alcohol. In the solvent, a phase separation occurs when butane gas dissolves into it at low temperatures, and over a wide composition range. Conditions for this separation have been identified and related to the high solubility of butane in the solvents at atmospheric pressure.The phenomenon is compared to conventional salting-out of non-electrolytes from water. Liquid mixtures consisting of water and an alcohol possess special features which have been recognised for some time. The variation in the excess enthalpy of mixing, and in the absorption of sound, with mole fraction of the mixture, are well known examples.' The solubility of gases in these liquid mixtures also reveals some of their concentration- dependent anomalies.2* When solubility is studied as a function of the mole fraction of alcohol in the liquid phase, maxima and minima occur on the solubility isotherms at low enough temperatures (generally below 25 "C). These maxima and minima, and the points of inflection to which they change at higher temperatures, have been interpreted by the influence exerted by the alcohol molecules and the gas molecules on the hydrogen- bonded structure of liquid water.The nature of the gas molecule, and the size of the alkyl group of the alcohol molecule are the important factors which govern the position and extent of the extrema on the solubility i s ~ t h e r m s . ~ ' ~ In the water-ethanol system, the greatest anomalies have been generated by some gaseous hydrocarbon solutes and the hydrophobic interaction has been invoked for an explanation.6 This paper describes the solubility of propane and butane in mixtures of water with t-butyl alcohol (2-methylpropan-2-01) ; this alcohol exaggerates anomalies. First, solubility measurements are given, then attention is drawn to a phase separation which occurs in this solvent system when sufficient butane dissolves into it.Experimental Propane and butane were supplied by Air Products Ltd, and stated to be 3 99.5 YO pure. Water was distilled and deionised. 2-Methylpropan-2-01 was better than 98 YO pure. The solubilities were measured to an accuracy of 2 YO, by the flowing-film te~hnique,~ over the temperature range 5-60 "C, and at a total pressure of gas plus solvent vapour of ca. 1 bar. Ideal behaviour of the gases in the gas phase and in the solution phase was assumed. It has been shown in previous work that this approximation is justified for the conditions of our experiments.6 t Paper presented at the Third International IUPAC Symposium on Solubility Phenomena, held at the University of Surrey, 23-26 August, 1988.26652666 Solubility of Butane in 2-Methylpropan-2-01-water Table 1. Solubility data for propane and butane at a partial pressure of 101.3 kPa in mixtures of water and 2-methylpropan-2-ol" log (S,,/cm3 kg-') log (S,/cm3 kg-I) ~~ T / K propane butane T / K propane butane 277.0 277.1 279.3 282.3 284.3 293.1 303.5 313.0 322.2 323.0 327.8 331.1 282.2 286.3 290.5 296.9 306.4 316.1 317.7 326.2 334.8 280.5 283.0 287.6 295.3 305.7 317.4 x = 0.000 I .873 1.810 I .767 1.720 I .566 1.443 1.351 1.262 1.240 - - x = 0.005 1.789 1.690 I .633 1.544 1.416 1.315 1.256 1.218 x = 0.026 1.837 1.761 1.691 1.585 1.473 1.390 - - I .828 I .752 1.710 1.663 1 SO7 1.340 1.233 1.159 1.123 1.129 1.751 1.560 1.463 1.301 1,232 1.188 1.137 1.103 1.796 I .70l 1.521 1.378 I .303 - 326.8 1.349 I .253 333.6 1.310 1.218 278.8 282.4 290.5 299.8 309.I 318.9 334.2 280.6 285.2 295.7 306.1 326.8 333.6 278.2 283.2 296.9 306.9 331.2 33 1.6 290.3 298.4 308.8 319.1 330.1 331.1 x = 0.038 1.790 1.736 1.653 1.572 1 SO6 I .493 1.736 x = 0.050 I .687 1.692 1.716 1.736 1.772 1.788 x = 0.066 1.950 I .946 2.022 2.086 2.309 2.344 x = 0.080 2.392 2.472 2.472 2.460 1.743 1.674 1.520 I .452 1.459 1.429 - 1.735 1.736 I .794 1.832 1.904 I .923 - 2.551 - - 2.594 2.623 3.044 2.877 2.832 2.776 2.699 2.685 = mole fraction of t-butyl alcohol. Results and Discussion Table 1 gives the experimental solubility data: So is the volume of gas in cm3 at 273.1 K and 101.3 kPa, dissolved in 1 kg of solvent under a partial pressure of 101.3 kPa. Fig. 1 shows the large increase in the solubility of butane at various temperatures when the mole fraction of t-butyl alcohol increases from 0 to 0.08.A similar but smaller increase occurs for propane. Accurate measurements of the solubility of butane in this mixed solvent could not be made at low temperatures when the mole fraction of t-butyl alcohol was much above 0.08, because phase separation took place in the liquid phase. The saturation with butane changed the miscibility characteristics of the pair of liquids making up the solvent. The curve on fig. 2 shows the temperature-composition limits of theR. W. Cargill und D. E. MacPhee 1 I 1 f 2667 0.02 0.04 0.06 0.08 mole fraction of t-butyl alcohol Fig. 1. Solubility isotherms for butane in water-t-butyl alcohol mixtures.e, 4.7 "C; x , 21.0 "C; 0, 39.4 "C; 0, 60.2 "C. h Fig. 2. 1 I I 1 1 I 1 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 mole fraction of t-butyl alcohol Conditions for phase separation in water-t-butyl alcohol mixtures.2668 Solubility of Butane in 2-Methylpropan-2-01-water Table 2. Water and t-butyl alcohol mixtures in which phase separation occurs volume of butane gas /cm3, dissolved in mole fraction of alcohol before separation upper layer lower layer 1 cm3 solvent 0.092 0.4 0.084 0.19 0.4 0.074 0.39 0.5 0.1 15 0.51 0.6 0.15 13 54 200 > 1000 phenomenon. At temperatures down to within a few degrees of the boiling-point of butane (272.6 K), no phase separation occurred in the solvent if its t-butyl alcohol content was less than 0.09 or greater than 0.7 mol fraction.The area under the curve identifies temperatures and compositions where phase separation was observed. The maximum on the curve occurs at ca. 9.5 "C and 0.085 mole fraction. Some representative data are given in table 2 for the compositions of the two phases (deduced from density measurements) and for the approximate amounts of butane gas which dissolved in the mixture to achieve the phase separation. The volume of butane dissolved and the composition of the upper phase were measured to ca. 10% accuracy. Of the monohydric alcohols, methanol, ethanol, both propanols and t-butyl alcohol are each miscible with water in all proportions, whereas the other butyl alcohols are partially miscible with water. When a certain amount of butane gas dissolved into the t-butyl alcohol-water system, it changed from total to partial miscibility. This resembles the conventional 'salting-out ' effect, where an added salt creates a phase separation or a precipitation from a liquid mixture, due to the affinity of the ions for the water molecules in the mixture.In the system described here, it is the affinity between the hydrocarbon gas molecules and the alcohol molecules with their bulky alkyl radical which causes the separation in the liquid phase. Phase separation caused by the dissolution of another gas, ethene, has been reported7 for this and other aqueous organic mixtures, but at pressures around 40 bar. With butane, the phase separation occurs at 1 bar because the solubility of this gas in this mixture is great enough at the temperatures indicated. This ' salting-out ' by a gas has been studied as a method of separating t-butyl alcohol from water. However, an analysis revealed that there was no cost advantage in competition with extractive and azeotropic distillation procedures. References I F. Franks and D. J. G. Ives, Q. Rev. Chem. SOC., 1966, 20, 1. 2 R. W. Cargill, J. Chem. SOC., Faraday Trans. I , 1978, 74, 1444. 3 R. W. Cargill, J. Chem. Res. 1982, ( S ) 230; (M) 2313. 4 R. W. Cargill and T. J. Morrison, J. Chem. SOC., Faraday Trans. I , 1975, 71, 618. 5 A. Ben Naim and S. Baer, Trans. Faraday SOC., 1964, 60, 1736. 6 R. W. Cargill and D. E. MacPhee, J. Chem. Res., 1986, (S) 276; (M) 2301. 7 J. C. Elgin and J. J. Weinstock, J. Chem. Eng. Datu, 1959, 4, 3. Paper 8/04448A ; Received 4th Nouember, I988

ISSN:0300-9599    

DOI:10.1039/F19898502665

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chrrii. Soc.. Fizr-nd~j~ Trans. I . 1989, 85(9). 2669-2674 Solubilization of some Tetramethylammonium Salts and of Ethyltrimethylammonium Bromide by their Homologues in Ch1oroform-t Jan Czapkiewicz Institute of Chemistry, Jagiellonian University, 30460 Krakow, Poland The solubilities of tetramethylammonium chloride, bromide, thiocyanate and perchlorate and of ethyltrimethylammonium bromide in chloroform at 25 f 0.2 "C have been determined. The solubilities of these salts increase markedly in the presence of a variety of higher homologues of common co- ions. It is suggested that this phenomenon involves the cooperative formation of reversed micelles. The low solubility of tetramethylammonium halides and of ethyltrimethylammonium bromide in chloroform had been noted already in 1907 by Wagner,' who also found that at an elevated temperature the latter salt separates out from the solution and forms a solute-rich liquid phase. Walden et a1.2 reported the extremely low solubility of tetramethylammonium perchlorate in this solvent.These findings and more recent solubility data published by Abraham and Danil de Namor3 for an extensive series of quaternary onium salts in 1,2- and 1,l-dichloroethanes as well as the comments and some data on the solubilities of such salts in dichl~romethane~ indicate that halogenated hydrocarbons are, as a rule, extremely poor solvents for tetramethyl and other low- molecular-weight onium salts. The solubility of homologous salts in these solvents increases markedly with growth of the hydrocarbon chain.This effect, however, should not be freely extrapolated to long-chain salts because a reversed situation may occur at some chain length as was shown for di-n-alkyldimethylammonium chlorides, the solubility of which in chloroform and carbon tetrachloride decreases in the order C,,-C,, > C,,-C,, > C,,-C,, over a broad range of temperat~res.~ A similar behaviour was found by Kertes' for normal dialkylammonium chlorides in carbon tetrachloride and benzene. The break in solubility occurred in both solvents upon passing from dioctyl to the didodecyl homologue. The present work was inspired by results of vapour-pressure osmometry studies carried out first for a series of long-chain (C16-Clo) alkyltrimethylammonium bromides in chloroform' and extended later to the short-chain (C,-C,) homologues.8 It was found that the mean aggregation number of these salts increases with the decrease in number of the carbon atoms in the alkyl chain.The concentration range covered for ethyltrimethylammonium bromide was, however, limited by its low solubility. It thus seemed interesting to explore the possibility of enhancing the solubility of this salt and of the tetramethylammonium salts, presumably possessing strongest aggregational properties, through the formation of mixed reversed micelles with freely soluble higher homologues. The results here presented should be treated as model studies on compartmentalization of sparingly soluble salts in micellar assemblies formed in chloroform. 1. Paper presented at the Third International IUPAC Symposium on Solubility Phenomena, held at the University of Surrey.23-26 August, 1988. 26692670 Solubilization of Tetramethylammonium Salts Experimental Tetramethylammonium bromide and chloride were commercial samples which were additionally recrystallized from isopropyl alcohol. The perchlorate salt was prepared by mixing aqueous solutions of the chloride and of sodium perchlorate which was an analytical-grade reagent. The precipitate was recrystallized several times from water. The tetramethylammonium thiocyanate was prepared according to the procedure described by Bekkevoll et al.'. Butyl-, propyl- and ethyl-trimethylammonium bromides were prepared by mixing a cooled (ca. - 15 "C) solution of trimethylamine in acetone (100 cm3 of amine per 1 dm3 of solvent) with a cooled acetonic solution of the appropriate alkyl bromide used in a ca.10 Yo molal excess. The mixtures were tightly stoppered and kept in the fridge for two days and then left for a few more days under the hood at room temperature. The products were than filtered off and thoroughly washed with acetone. Recrystallizations were best carried out from nitromethane, chloroform, ethyl acetate and from 1,2- dichloroethane. These and further operations were carried out in a dry box because the salts, especially the ethyl homologue, are hygroscopic. Tetrapropylammonium bromide was prepared by quaternization of tripropylamine with propyl bromide in butan-2-one in a sealed glass ampoule heated for 4 days at ca. 80 "C. Tetrabutyl-, decyltrimethyl- and decyl-tripropylammonium bromides as well as dodecyltrimethylammonium chloride and tetradecyltrimethylammonium thiocyanate were prepared in this laboratory ear lie^.^ All the samples were thoroughly dried in a heated vacuum pistol containing P,O,.Analytically pure chloroform was washed with concentrated sulphuric acid, then five times with a double volume of water. It was initially dried with CaCI,, then with type A-4 zeolites and finally distilled, the middle fraction yielding 50 %. The purified solvent was used to test for solvate formation by applying a modified procedure of de Ligny et al."' The salts were stored in evacuated closed vessels over chloroform and P,O,. Removal of air secured the bromides from turning yellow-brown in the presence of chloroform vapour.The test was negative for all tetramethyl- ammonium salts studied. The ethyltrimethylammonium bromide appears, however, to form solvates. At room temperature it absorbs from the vapour phase c w . 1.5 mol chloroform per mole of salt. For the solubility measurements in chloroform an excess amount of the salt was added to the solvent and the flasks containing the mixture were tightly stoppered and shaken for 50 h in a thermostat at 25 & 0.1 "C. In the case of tetramethylammonium bromide the crystals assumed a slight yellow-green colour after several hours of contact with chloroform and a yellow-brown colour of the solution could be detected at the end of the experiment. This side-effect was eliminated by purging chloroform with argon prior to solubility measurements.It was also observed that this bromide remained colourless in the presence of higher alkyltrimethylammonium salts in solution. Aliquots of the saturated solutions were removed, weighed and evaporated to dryness by applying slightly reduced pressure. A similar procedure was adopted for solubility measurements of the Me,NX salts in solutions of their homologues of known molal concentration in chloroform. The weight of the solution used for the solubilization was noted and checked at the end of the experiment. The residue left after the removal of solvent was dissolved in water. The halides and the thiocyanate were determined potentiometrically by titration with a standard AgNO, solution. Appropriate anion-selective electrodes and an Ag-AgCI reference electrode with an NH,NO,-agar bridge were used for this purpose. Blank experiments without added Me,NX salts were carried out, with a precision of 1 %.The perchlorate anion was determined spectrophotometrically in the form of an ion pair with methylene blue extracted quantitatively to 1,2-dichloroethane."J. Czapkiewicz 267 1 Table 1. Solubilities of Me,NX salts and of EtNMe,Br in chloroform, at 298 K" _ _ _ _ ~ . ~ - ______ _ _ _ _ ~ - salt solubility/mol kg-' Me,NCI Me,NCNS Me,NClO, ca. I x (ca. 1.5 x EtNMe,Br 2.85k0.05 x (4.7 x lo-*) -. ~- ~ _ _ _ _ _ ~ 9.7kO.l x lo-' (1.45 x 8.2kO.l x 10-5 ( 1 . 2 ~ lo-'; 8.3 x lo-, in DCM') 1.14 x lo-' in I,2DCEb) Me,NBr 3.7+0.1 x 10-4 (5.5 x 10-4) 3.58 x lo-' in I,2DCEb and 8.08 x 10-5 in I,IDCEb In parentheses are given values expressed on a molar scale and included are available data for 1,2-dichIoroethane ( I ,2DCE), 1 , l -dichloroethane ( I , 1 DCE) and dichloromethane (DCM).Taken from ref. (3). 'Taken from ref. (4). The solubility of ethylthrimethylammonium bromide was also measured by preparing a warm solution of the salt and leaving it in a thermostat for 50 h at 25 "C. Both methods gave comparable results. Since the solubilization of this salt in the presence of its homologues occurs almost instantaneously, the following procedure was used. A small flask tightly stoppered with a polyethene plug was weighed accurately and reweighed after addition of the salt. Samples weighing 10&500 mg were used. The flask was then placed in a thermostatted oil bath contained in a large glass beaker and the solubilizer of known molar concentration was injected dropwise using a glass syringe with a steel needle which was pierced through the plug.The flask was illuminated, shaken by hand and observed with the aid of a lens. After complete dissolution of the solid material or of a second liquid phase the flask was thoroughly wiped and reweighed. Such a titration procedure lasted 1-2 h and gave reasonably reproducible results. Results and Discussion The solubilities of the Me,NX salts and of the solvated EtNMe,Br in chloroform are shown in table 1. Included are the relevant data reported for other halogenated solvents. The values confirm the early on the low solubilities of Me,NX salts in chloroform. There seems, however, to be no correlation between the data for this solvent and for the two isomeric dichloroethanes as well as for dichloromethane.The chloride appears to be by an order of magnitude more soluble in chloroform than in 1,2-dichloroethane whereas a reversed situation occurs for the perchlorate. The thiocyanate is by an order of magnitude less soluble in chloroform than in dichloromethane. Since solvate formation by Me,NX salts in both dichloroethane solvents was also unobserved3 it seems that the apparent lack of consistency of the solubility data may account for the fact that the dichloroethane and dichloromethane solvents have dielectric constants which are ca. two times higher than that for chloroform (e = 4.7) and that, in consequence, ion-ion pair equilibria are most important for the latter solvents, whereas the ion pair-n-mers equilibria are the predominant processes in chloroform.In table 2 values are given for the solubility of Me,NCl in the presence of dodecyltrimethylammonium chloride. Least-squares analysis of the data for the solubility of Me,NCl in the presence of the long-chain solubilizer indicates a linear relation with a correlation coefficient of 0.999, a slope of 0.140 0.002 and the intercept of 0.1 1 1 kO.01 I x lo-' mol kg-l. It appears that this extrapolated value is slightly higher than that determined for Me,NCl in pure chloroform. This shift may well account for the defects of the surface of the solubilized crystals caused by adsorption of the higher homologue.2672 Solubilization of Tetramethylammonium Salts Table 2.Solubilization of Me,NCl by C,,H,,NMe3Cl in chloroform at 298 K concentration of solubility of C12H25NMe3C1/ Me,NCl/ lo-, mol kg-' lo-, mol kg-' 0.42 0.57 0.68 1 S O 2.03 2.60 3.64 6.0 1 8.98 9.40 0.097 0.19 0.20 0.22 0.34 0.36 0.44 0.62 0.96 1.36 1.45 The data for solubilization of Me,NBr by PrNMe,Br and by C,,H,,NMe,Br as well as for solubilization of Me,NCNS by C,,H,,NMe,CNS confirm the results obtained for the chloride. An increase in solubility of the bromide and the thiocyanate was observed systematically. However, owing to their low solubilities they contributed an amount of the anions which never exceeded 5% of the total concentration in the systems studied. Thus, the results are not so convincing and are not here included. Much more interesting are the results obtained for the solubilization of EtNMe,Br by its various homologues as shown in fig.1. A manifold increase in solubility is also observed here. The situation, however, is reversed in the sense that the solubilizate plays the role of the host in the mixed micelles. The results for the Et-Pr ternary system exhibit some pecularities. A second liquid phase separates out upon addition of the solution of PrNMe,Br and exist over a broad concentration range. Below the dotted curve on fig. 1 a single liquid phase exists. As mentioned earlier, such a phenomenon was already observed by Wagner' for warm solutions of EtNMe,Br in chloroform. This phase separation resembles the so-called cloud point observed in the case of aqueous solutions of polyoxyethylene non-ionic detergents.It is caused by dehydration of the solute at elevated temperature. Indeed, the two-phase liquid systems studied presently could be reversibly cooled down below 25 "C to form a homogeneous solution. The temperature of such transitions depends on the concentration of PrNMe,Br. The separation of a second phase is caused most probably by desolvation of the ion pairs and n-meric species. Knowledge of the composition of such solute-rich phases might perhaps give an insight into the structure of reversed micelles. Their cores may still be substantially solvated. Higher homologues, which are surface active, do not form liquid two-phase systems with Et-Me,Br as the solubilizate. Their behaviour is similar to that observed for the C,,H,,NMe,Cl-Me,NCl system as exemplified by the Pr,NBr-EtNMe,Br and C,,H,,NMe,Br-EtNMe,Br systems.In both cases the extent of solubilization grows linearly with increase in concentration of the additive and the relation satisfactorily extrapolates to the solubility values determined for EtNMe,Br in pure chloroform. There are two additional data points on fig. 1 corresponding t o Bu,NBr and C,,H,,NPr,Br used as solubilizing agents, both 0.05 mol kg-'. These were chosen to illustrate the relation between the ability of the salts to solubilize and their tendency to form micellar aggregates. In table 3 the mean aggregation number, 6, determined by vapour pressureJ. Czapkie w icz 2673 0.20 ym 0.15 24 - 8 6 3 g 0.10 1 h v F 0.05 Pr4NBr f DecNMe3 Br A Bu4NBr A DecNPr3Br J / ,A' -0- PrNMesBr 0.01 0.03 0.05 concentration of solubilizing saltlmol kg- ' Fig.1. Solubilization of EtNMe,Br by homologous salts as a function of their concentration. Table 3. Mean aggregation number, n', of quaternary ammonium salts (0.05 molal) in chloroform and equilibrium concentration, C,,,, of solubilized EtNMe,Br salt ii C,,/mol kg-' C,,H,,NMe,Br 2.50 0.175 C,oH,,NPr,Br 1.86 0. I45 Pr,NBr 2.19 0.205 Bu,NBr 1.87 0.160 __ ~~~ ~- ______________ .- _- osmometry for the four 0.05 mol kg-' salts are compared with the equilibrium concentration of solubilized EtNMe,Br. The present results support the general finding' that in the group of RNRiBr salts, where R 3 R', the mean aggregation number grows with the decrease in R and/or R'. The pairwise comparison of the solubilization abilities of the ammonium salts, Pr,NBr > Bu,NBr, C,,H,,NMe,Br > C,,H,,NPrBr and Pr,NB > C,,H,,NPr,Br, correlates well with their aggregational tendency.This is not true for the Pr,NBr > C,,H,,NMe,Br pair. These two salts, however, do not have common structural features. It may be expected that the relation PrNMe,Br > C,,H,,NMe,Br would hold if measurements were carried out below the critical temperature at which phase separation occurs. It is noteworthy that for the system which is 0.05 molal with respect to Pr,NBr and2674 So 1 u h ilizu t ion o j Te t r urne t h y lam rn on iu rn Salts 0.205 mol kg-I with respect to EtNMe,,Br, the vapour pressure osmometry yields a value of the mean aggregation number of 10.1. The results discussed here represent model studies. They suggest, however, that solubilization of scarcely soluble electrolytes through the formation of mixed assemblies in chloroform may be a potential tool in extraction and separation procedures. This work supported by grant no. R.P.1-08. References I L. Wagner, Z. Kryslullogr., 1907. 43, 148. 2 P. Walden, H. Ulich and B. Busch, 2. Phq”. Chcm., 1926, 123, 443. 3 M. H. Abraham and A. F. Danil de Namor, J. Chem. Soc., Furuduy Truns. I , 1976, 72. 955. 4 S. Bekkevoll, I. Svorstol. H. Hoiland and J . Songstad, Ac/u C ’ l i m / . Scrrntl.. See/. B, 19x3. 37. 935. 5 H. J. Harwood and P. L. Du Brow, J. Org. Chcw., 194X, 13, 186. 6 A. S. Kertes, J. Inorg. Nucl. Chem., 1965, 27, 209. 7 J. Czapkiewicz, J. Colloid Inlcr-ucc Sci., in press. 8 J. Czapkiewicz, to be published. 9 J. Czapkiweicz and B. Czapkiewicz-Tutaj, J. Chcm. Soc., Furuduy Trunr. I , 1980, 76, 1663. 10 C. L. de Ligny, D. Bax, M. Alfenaar and M. G. L. Elferink, Rrcl. Truti. Clihi. Puj-.s-Bu.s, Bdg., 1969, 88, 1183.

ISSN:0300-9599    

DOI:10.1039/F19898502669

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Cireni. Soc'.. Fui*aduj. Trans. I , 1989, 85(9), 2675-2682 Solubility of H,S, CO, and CH, in N-Formyl Morpholinet F-Y. Jou, R. D. Deshmukh, F. D. Otto and A. E. Mather" University of Alberta, Edmonton, Alberta, Canada T6G 2G6 The solubility of H,S, CO, and CH, in N-formyl morpholine has been measured at five temperatures between 298.15 and 403. I5 K at pressures up to ca. 7 MPa (1 3.6 MPa in the case of methane). The experimental data were correlated by the Peng-Robinson equation of state. Parameters for the N-formyl morpholine were estimated from liquid molar volumes and vapour pressures. Values of the binary interaction parameter were obtained from the experimental data. Using the expressions connecting the values of the binary interaction parameter with the parameters of the Krichevsky- Ilinskaya equation, Henry's coefficients for the three solutes were obtained.Most physical solvents that are used for the removal of the acid gases (H,S and CO,) from gas streams are polar organic compounds which have a strong affinity for hydrogen sulphide and/or carbon dioxide. At low pressures, the solubility of gases in physical solvents is usually proportional to the partial pressure of the gas at a given temperature. The enthalpy of solution is generally small (< 5 kJ mol-') and it is often possible to regenerate the solvent by flashing to a lower pressure. Hence processes using physical solvents require smaller amounts of energy than the traditional processes using chemical solvents such as alkanolamines. N-Formyl morpholine is a physical solvent which has been proposed for the separation of the acid gases, H,S and CO,, from natural and synthesis gas streams.A knowledge of the solubility of these gases is necessary to decide if this solvent has any advantage over those solvents now in use. The solubility of methane is also important, since its magnitude is a measure of the loss of hydrocarbons in the solvent. To answer these questions, the solubility of these gases in N-formyl morpholine was measured at five temperatures between 298.15 and 403.15 K at pressures up to ca. 7 MPa. Experimental The experimental apparatus and procedure are similar to that outlined by Jou et al.' The liquid and vapour phases were brought to equilibrium in a windowed Jerguson cell. A 250 cm3 cylindrical reservoir was attached to the top of the cell to increase the volume of the vapour phase.The vapour from the reservoir was recirculated through the solvent by a magnetically driven piston pump. The cell and pump were enclosed in a 0.4 m3 air bath maintained at f 0.5 "C of the set-point temperature. The heater consisted of a series of electrically heated fins which allowed for measurements to 130 "C. The air within the bath was mixed with a blade fan. With the exception of the pump piston, all metallic materials in contact with the fluids were type 316 stainless steel. The pump piston was made from Carpenter 450 steel as it must be ferromagnetic. The pressure in the cell was measured with calibrated Heise gauges, which have an accuracy of f O . 1 YO of full-scale span.The temperature of the fluid in the cell was monitored by an iron-constantan thermocouple. A gas sample line extended from the reservoir to the sample loop of the gas chromatograph. The liquid sample line led from the base of the cell to a needle valve located outside of the air bath. University of Surrey, 23-26 August, 1988. f Paper presented at the Third International IUPAC Symposium on Solubility Phenomena, held at the 89 2675 F A R I2676 Solubilities in N-Formyl Morpholine Table 1. Solubility of CO, in NFM total pressure mole fraction / kPa" of co, S/kmol kg-' T = 298.15 K 6410 0.971 0.287 6378 0.958 0.183 6359 0.941 0.137 6298 0.910 0.087 5 6267 0.899 0.076 9 6216 0.854 0.050 8 6126 0.803 0.035 3 5906 0.720 0.022 3 2950 0.369 0.005 07 1120 0.145 0.001 48 379 0.052 4 0.000 48 1 550.8 (45.5) 0.007 05 0.000 06 1 7 540.4 (4.56) 0.000 663 0.000 005 76 487.4 (0.452) 0.000 0704 0.000 000 612 T = 313.15 K 7090 0.620 0.0142 5530 0.504 0.008 8 1 3240 0.294 0.003 63 1320 0.136 0.001 37 334 0.039 3 0,000 352 553.0 (38.2) 0.00422 0.000 036 8 503.2 (26.4) 0.002 99 0.000 026 1 506.8 (4.24) 0.000 464 0.000 004 04 45 1.2 (0.788) 0.000 094 3 0.000000819 T = 343.15 K (NFM vapour pressure = 0.13 kPa) 6590 0.389 0.005 52 5200 0.304 0.003 79 4660 0.285 0.003 46 3810 0.232 0.002 62 1930 0.1 18 0.001 16 560.4 (72.3) 0.004 86 0.000 042 5 506.4 (6.86) 0.000 4 1 9 0.000 003 64 502.3 (2.79) 0.000201 0.000001 75 489.3 (2.22) 0.000 136 0.000001 19 T = 373.15 K (NFM vapour pressure = 0.81 kPa) 6750 0.286 0.003 48 5440 0.232 0.002 62 3730 0.162 0.001 67 2210 0.094 8 0.000 9 10 87 1 0.038 3 0.000 346 264 0.012 3 0.000 108 554.3 (22.3) 0.001 03 0.000 008 98 T = 403.15 K (NFM vapour pressure = 3.25 kPa) 491.7 (5.31) 0.000 244 0.000 002 12 6270 0.20 1 0.002 19 4950 0.164 0.001 71 31 10 0.099 3 0.000 958 808 0.029 1 0.000 261 299 0.0106 0.000093 1 468.2 (22.3) 0.000 83 1 0.000 007 22 53 1.4 (4.74) 0.000 188 0.000 00 1 63 "The partial pressure of the solute is given in parentheses when nitrogen was added in order to raise the total pressure.F-Y. Jou, R.D. Deshmukh, F. D. Otto and A . E. Mather Table 2. Solubility of H,S in NFM total pressure mole fraction /kPa' of H,S S/kmol kg T = 298.15 K 1900 0.968 0.264 1500 0.799 0.0344 942 0.563 0.01 12 464.2 (1 32.0) 0.124 0.001 23 471.4 (93.4) 0.095 3 0.00091 5 449.8 (1 8.4) 0.022 2 0.000 198 481.3 (4.85) 0.006 31 0.000055 1 464.6 (0.670) 0.001 17 0.000010 1 468.7 (0.0797) 0.000 236 0.000 002 06 T = 313.15 K 278 1 0.984 0.529 1760 0.684 0.0184 1370 0.583 0.012 1 1030 0.480 0.008 01 753 0.366 0.005 03 555 0.304 0.003 79 458.7 (197.0) 0.121 0.001 20 496.4 (44.5) 0.032 1 0.000 288 445.3 (18.1) 0.0144 0.000 127 339.4 (5.89) 0.005 34 0.000 046 7 455.3 (0.657) 0.000 852 0.000 007 4 1 393.2 (0.169) 0.000 258 0.000 002 54 T = 343.15 K (NFM vapour pressure = 0.13 kPa) 5230 0.970 0.280 3140 0.670 0.0177 1720 0.434 0.006 65 655 0.199 0.002 16 473.4 (1 65.0) 0.057 9 0.000 534 468.4 (29.8) 0.0120 0.000 106 500.6 (5.27) 0.002 92 0.000 025 4 469.0 (1.13) 0.000 744 0.000 006 47 T = 373.15 K (NFM vapour pressure = 0.8 1 kPa) 472.3 (0.0905) 0.000 116 0.000001 01 6500 0.812 0.0374 4840 0.668 0.0175 3650 0.540 0.0102 2040 0.347 0.004 6 1 933 0.180 0.001 91 254 0.059 1 0.000 545 493.8 (53.1) 0.0148 0.000 13 1 464.3 (6.23) 0.002 33 0.000 020 3 414.2 (0.796) 0.000457 0.000 003 97 462.5 (0.237) 0.000 168 0.000 00 1 46 T = 403.15 K (NFM vapour pressure = 3.25 kPa) 6750 0.636 0.0152 4950 0.506 0.008 90 3060 0.350 0.00468 1230 0.161 0.001 67 364 0.054 1 0.000497 495.6 (34.1) 0.007 26 0.000 063 6 440.3 (3.62) 0.001 11 0.000009 65 444.3 (1.23) 0.000430 0.000 003 74 470.4 (0.100) 0.000 055 0 0.000 000 478 2677 "The partial pressure of the solute is given in parentheses when nitrogen was added in order to raise the total pressure.89-22678 Solubilities in N-Formyl Morpholine Table 3. Solubility of CH, in NFM total total /kPa 103X /kPa 103X pressure 1 05S/kmol kg-l pressure 1 05S/ kmol kg-' I3 450 10 030 7 360 4 840 2410 1230 452 93 13 620 1 1 650 7 200 3 800 2 020 988 295 13 470 10 800 6610 4 420 2 230 847 198 T = 298.15 K 55.3 50.9 43.8 39.8 35.2 31.7 24.7 22.0 13.1 11.5 7.06 6.17 2.60 2.26 0.555 0.482 T = 313.15 K 56.9 52.4 50.5 46.2 34.9 31.4 20.0 17.7 11.2 9.82 5.48 4.78 I .75 1.52 T = 343.15 K 61.4 56.8 51.0 46.7 34.3 30.8 23.8 21.2 12.5 11.0 4.79 4.18 1.20 I .05 13 080 9 720 6 660 3 550 1200 576 26 1 1 3 090 9 190 6 490 4110 2 760 1300 629 185 T = 373.15 K 64.2 59.6 50.6 46.3 36.1 32.5 20.2 17.9 6.96 6.09 3.45 3 .OO 1.60 0.92 1 T = 403.15 K 67.7 63.1 50.6 46.3 37.6 33.9 24.9 22.2 16.7 14.7 7.92 6.93 3.90 3.40 1.22 1.06 Table 4.Equation of state constants for NFM for the Peng-Robinson equation T/K u , , Pa m'; molP h, cm:' mol-' ~ ~~ ~- 298.15 6.7909 93.203 313.15 6.57 19 93 -404 343.15 6.2 157 93.718 373.15 5.9336 93.909 403.15 5.6995 93.965 ~~ - ~~ Table 5. Correlation parameters (a,,) for d,, ~~ ~- ~~ ~ co, H,S CH, ~~ -0.0174 - 0.06 1 7 0.1 182 ~~~ - Prior to introduction of the fluids, the apparatus was brought to the desired temperature and purged with nitrogen to remove traces of oxygen. Ca.50cm3 of the solvent was fed by gravity to the equilibrium cell. Hydrogen sulphide or carbon dioxide was added to the cell in an amount indicated by the pressure. The pump was started and the vapour recirculated through the solvent. Additional amounts of acid gas were addedF-Y. Jou, R . D. Deshmukh, F. D.Otto and A . E. Mather 2679 8 7 6 5 --. $ 4 !?i. v1 n 3 2 1 0 I I I I P 1 0 0.2 0.4 0.6 0.8 1 mole fraction CO, Fig. 1. Solubility of CO, in N-formyl morpholine. T/K: @, 298.15; 0, 323.15; V, 348.15; 0, 373.15; ., 398.15. until the desired partial pressure had been approximately obtained. When necessary, nitrogen was added to maintain the system pressure above 200 kPa. At equilibrium, as established by a constant cell pressure, the pump was stopped and the phases analysed. A portion of the vapour was released to the sample loop of a gas chromatograph. A sampling valve was used to inject the gas into a 3 m long, 6.35 mm OD column packed with Chromosorb 104. The liquid sample was withdrawn from the equilibrium cell into a vessel containing 1 mol dm-3 NaOH, thus converting free dissolved acid gas into the involatile ionic species.A 40 cm3 high-pressure sample bomb containing the sodium hydroxide was used for sampling at loadings of CO, where the reaction rate was not great enough to prevent pressure build-up over the basic solution. A 50 cm3 Erlenmeyer flask fitted with a rubber septum served as a collection vessel for the sampling of solutions of CO, at partial pressures < 1000 kPa and for the sampling of solutions containing H,S. The procedure was somewhat different when no acid gases were present. The liquid sample was passed into a 50 cm3 weighed sample bomb while constant pressure was maintained in the equilibrium cell by gas addition. The bomb was reweighed to determine the amount of sample and then attached to a mercury-filled burette and the pressure brought to atmospheric.The gas which evolved from the liquid was collected in the calibrated 50 cm3 burette. The amount (in mol) collected was calculated from the P-V-T data after subtracting the vapour pressure of the liquid. In addition, it was necessary to account for the small amount of gas, equivalent to the solubility at atmospheric pressure, which remained in the liquid sample. Gas chromatography was used to measure these residual solubilities.2680 7 6 5 2 4 2 : Solubilities in N-Formyl Morpholine I I I I 0 0.2 0.4 0.6 0.8 1 mole fraction H2S Fig. 2. Solubility of H,S in N-formyl morpholine. Symbols as for fig. 1. The CO, content in an aliquot of the liquid sample was determined by adding excess 0.05 mol dm-, BaCl, to precipitate the carbonate as BaCO,.The precipate was washed and titrated with standardized 0.1 mol dm-, HCl using methyl orange-xylene cyanol indicator to a grey-green endpoint. The H,S content in an aliquot of the sample was determined by reacting the liquid with a solution of acidified 0.05 mol dmP3 I,. The unreacted I, was back-titrated with 0.05 mol dm-, Na,S,O, using starch indicator. The experimental error in the liquid-phase concentration is estimated to be +2-3 % in the range studied. Results and Discussion The solubility of H,S and CO, in N-formyl morpholine was determined at 298.15, 313.15, 343.15, 373.15 and 403.15 K at partial pressures of the acid gases up to ca. 7 MPa. The data are presented in tables 1 and 2 for CO, and H,S, respectively.The solubility of CH, in N-formyl morpholine was measured at the same five temperatures listed above. Partial pressures ranged up to 13.6 MPa. The data for methane are given in table 3. At a given partial pressure there is only a small effect of temperature on the solubility of methane. Both the mole fraction of solute in the liquid phase and the solubility S (kmol of solute per kg of solvent) are reported. For the acid gases, the partial pressure, the product of the mole fraction solute in the vapour phase and the total pressure, is reported when nitrogen was used to maintain the total pressure above the atmospheric pressure. The experimental data were correlated using the Peng-Robinson2 equation of state. The procedure followed was analogous to that described by Jou et aL3 The parameters for the pure solvent (NFM) were obtained from the liquid density and vapour pressure obtained from the fragmentary points of Cinelli et al.* and Vetere et al.5F-Y.Jou, R. D. Deshmukh, F. D. Otto and A . E. Mather 15 10 2 2 [ 5 I I I I I i / 0.00 0.01 0.02 0.03 0.04 0.05 0.06 0.07 mole fraction CH4 Fig. 3. Solubility of CH, in N-formyl morpholine. Symbols as for fig. 1. Table 6. Parameters for the Krichevsky-Ilinskaya equation H2,/MPa z?;/cm3 mol-' A / R T 298.15 313.15 343.15 373.15 403.15 298.15 313.15 343.15 373.15 403.15 298.15 313.15 343.15 373.15 403.15 6.89 8.92 13.8 19.6 25.9 1.15 1.62 2.93 4.70 6.92 187 178 167 160 155 co, 35.55 36.60 38.93 41.60 44.69 35.07 35.97 37.97 40.23 42.83 H2S 37.32 CH, 38.56 41.29 44.39 47.94 0.099 0.105 0.121 0.142 0.168 -0.151 -0.137 -0.1 10 - 0.084 -0.059 0.736 0.702 0.656 0.632 0.623 268 12682 Solubilities in N-Formyl Morpholine I 275 325 375 42 5 TIK Fig.4. Effect of temperature on the Henry’s constant: (----) CH,, (--- -) co,, (..-. .) H,S. The parameters are given in table 4. The experimental data were used to obtain the binary interaction parameter which appears in the mixing rule of the equation of state. The values of aij were found to be independent of temperature. The values are presented in table 5 for the three solutes. The results of the correlation are compared with the experimental data in fig. 1-3. In general, the correlated values are in good agreement with the experimental data over the wide ranges of temperature and pressure involved.Bender et aL6 have shown the connection between the binary interaction parameter and the three parameters in the Krichevsky-Ilinskaya7 equation : the Henry’s constant, the partial molar volume of the solute at infinite dilution, and the Margules parameter. Their equations (corrected) have been used to obtain these parameters, which are presented in table 6. The results for the Henry’s constants are shown in fig. 4. No comparison with other data are possible as only Zawacki et al.8 have measured data in these systems. The authors are grateful to the Alberta/Canada Energy Resources Research Fund for financial support of this research. References 1 F-Y. Jou, A. E. Mather and F. D. Otto, Ind. Eng. Chem. Process Des. Dev., 1982, 21, 539. 2 D-Y. Peng and D. B. Robinson, Ind. Eng. Chem. Fundum., 1976, 15, 59. 3 F-Y. Jou, R. D. Deshmukh, A. E. Mather and F. D. Otto, Fluid Phase Equilibria, 1987, 36, 121. 4 E. Cinelli, S. Noe and G. Paret, Hydrocurbon Process., 1972, 51(4), 141. 5 A. Vetere, R. De Simone and A. Ginnasi, Ind. Eng. Chem. Process Des. Dev., 1975, 14, 141. 6 E. Bender, U. Klein, W. P. Schmitt and J. M. Prausnitz, Fluid Phase Equilibria, 1984, 15, 241. 7 I. Kritchevsky and A. Iliinskaya, Acta Physicochim. U.R.S.S., 1945, 20, 327. 8 T. S. Zawacki, D. A. Duncan and R. A. Macriss, Hyclrocurhon Process., 1981, 60(4), 143. Paper 8104463E ; Received 4th Nocember, 1988

ISSN:0300-9599    

DOI:10.1039/F19898502675

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. I, 1989, 85(9), 2683-2694 Solubilities, Solubility Products and Solution Chemistry of Lanthanon Trifluoride-Water Systems7 Manchery P. Menon and Jeffrey James Department of Chemistry, Savannah State College, Savannah, Georgia, U.S.A. Solubilities and solubility products of 14 lanthanon trifluorides (LnF, . 0.5H20) have been determined using conductometric, potentiometric and radiometric techniques. The solubilities and the averages of selected pK,, values from literature and from our own work, that do not differ by more than one unit, show three maxima when plotted against the atomic number of lanthanons. Standard enthalpy change measured from the temperature-dependence plots of K,, and the free-energy change calculated from Ksp of the trifluorides fall within the range 3 M O kJ mol-' and 78- 103 kJ mol-', respectively, in most cases.Standard entropy change calculated from the enthalpy and free energy changes ranges from -47 to - 294 J K-' mol-'. The stability constants determined for the monofluoride complexes fall within the range of reported values, but those for the difluoride complexes are much higher than the fewer values that are found in literature. Lanthanon trifluorides were found to be least soluble in water at pH 3-4. The Commission for the Solubility Data Project organized by the International Union of Pure and Applied Chemistry almost a decade ago seeks to establish a data base for the solubility of gases, liquids and solids in liquids and solids. Siekierski and Miqduski have been deputed to compile the solubility data on lanthanon halides.' In contIast to other lanthanon trihalides the solubility information concerning the sparingly s( duble lanthanon trifluorides is generally presented in the form of the solubility product.KSB2 It is not sometimes clear whether the solubility product is the true thermodyiiamic solubility product, K,,, based on activities or those computed in terms of stoichiometric concentration of ions in aqueous solution. Burgess and Kijowski' have listed several sets of pK,, values for all lanthanon trifluorides (except PmF,) reported by several authors using various techniques. However, agreement among values for a given fluoride is less than satisfactory, the difference being as high as three units (three orders of magnitude in Ksp) in some cases.Although two sets of older data show a steady decrease in pK,, from LaF, to LuF,~** more recent values show an increase from LaF, to EuF, ancI then a decrease to ErF,.5 Furthermore, it is not very clear whether allowance for hydrolysis and/or complexation of Ln3+ and F- ions has been made in the calculation of earlier value^.^*^ Since thermodynamic K,, values of lanthanon trifluorides are very useful to solve several analytical problems including complexation of lanthanide ions addi .ional data and better techniques for the measurement of Kso are needed. Although the enthalpies of solution of other lanthanon trihalides in water at 25 "C are well known,2 those for the dissolution of lanthanon trifluorides in water were non- existent before the commencement of our work.This is probably due to the impracticability of direct measurement of enthalpies of solution in view of their poor solubilities in water. However, enthalpies of solution of the trifluoride can be estimated from the temperature dependence of Kso, using the standard van't Hoff equation. Since University of Surrey, 23--26 August, 1988. t Paper presented at the Third International IUPAC Symposium on Solubility Phenomena, held at the 26832684 Solution Chemistry of Lanthanon Trifluoride- Water the standard free energy change can also be estimated from the thermodynamic Kso, all the three functions, AH&., AGZiss. and AS&, can be computed using standard thermodynamic relations. Lanthanide ions are known to form complexes with fluoride ions in aqueous solutions, but there is considerable disagreement among values reported for the monofluoride complexes.'* ' There is also a paucity of information on the overall stability constants of lanthanon difluoride complexes.'.This paper presents the final results of our measurement of the solubilities, pK,, values, thermodynamic functions and the stability constants of the 14 lanthanon fluoride-water systems carried out over several years. Experimental Materials and Equipment Lanthanon chlorides (99.9% pure), mostly in the hydrated form were obtained from Thikol/Ventron Division of Alpha Products and the radioisotopes used in the study were supplied by Amersham Corporation. Only reagent chemicals were used for making solutions needed for the study.Doubly distilled water was employed for making solutions. Polyethylene ware was chosen for handling fluoride solutions and their mixtures with lanthanons. The equipment used for the study are the following: Harshaw model 3 in? x 3 in NaI(T1) well crystal pray spectrometer, Giger-Muller counter, Orion ion analyser meter model 407, Orion expandable and digital ion analyser model E 920 (for more recent work on stability constants) and their fluoride ion electrode model 96.09, Yellow Springs Instrument Co. Inc. (YSI) model 32 conductance meter and their model 3403 conductivity cell, a Burnell model AB wrist-action shaker and a Fisher Scientific model 137 Versa Bath. Solubility Measurements The detailed procedures for the measurement of solubilities of the lanthanon trifluorides (LnF, .0.5H20) by the conductometric, potentiometric and radiometric methods have been reported el~ewhere.~-l, Lanthanon trifluorides were precipitated by mixing hot solutions of the respective chlorides and sodium fluoride. Precipitated fluorides were washed several times with doubly distilled water to remove the absorbed ions and then dried in the oven at 110°C. Small portions of the purified and dried precipitate (fluoride semi-hydrates) were agitated with doubly distilled water in four set of graduate cylinders made of polyethylene, intermittantly with the shaker for more than 48 h to prepare saturated solutions. Attainment of saturation condition at 25 "C was tested with repeated measurement of the conductance of solutions. The conductance of saturated solutions of the fluoride as a function of temperature, ranging from 20 to 55 "C, was measured by agitating the mixtures in constant-temperature bath for a minimum of 7 h.The conductance of doubly distilled water and the pH of the saturated solutions were also measured at each temperature. The total solubility of LnF, was measured using equations described in previous publications. '-13 The free fluoride ion concentration in saturated solutions of LnF, was also measured potentiometrically using the Orion ion analyser meter. An Orion standard (0.100 mol dm-, NaF) was diluted with doubly distilled water to prepare the working standards ( 10-3-10-4 mol drn-,) for calibration of the instrument. The pH and ionic strength of the standards were very similar to those of the samples.The pH of the fluoride solutions, in general, ranged from 5.0 to 6.0. No other electrolytes were added to the system. Radiometric procedure was employed only to such cases where a radioisotope of moderate half-life was available. In this method the solutions of lanthanon chlorides of t I in = 2.54 x lo-' m.M . P. Menon and J. James 2685 known concentration was mixed with appropriate amount (ca. 500 pCi) of the respective radioisotope and the radioactive fluoride was precipitated as discussed before. After the precipitates had been washed and dried, small portions were agitated with water in four polyethylene bottles until saturation was attained. The specific radioactivity (cpm mmol-l) of the original stock solution and of a solid standard, after digestion with concentrated H,SO, and dilution with doubly distilled water, was determined by counting either y or B radiation.The specific radioactivity of the stock solution and of the solid standard agreed well within experimental errors. After filtration and centrifugation radioactivity of 2 cm3 of the clear filtrate from each sample was measured under identical geometry and then compared with the specific activity of the labelled standard to determine the total concentration of lanthanon. l 3 Measurement of the Stability Constants of Lanthanon Fluoride Complexes Stability constants for the monofluoride complexes were measured, in the past, using an Fe"/Fe'" ele~trode,'~ solvent extraction using radioisotopes15* l6 and potentiometric titration using an LaF, ele~trode.~ Hefter et al.17 used a simple potentiometric method for the measurement of the stability constants of fluoride complexes of a few monovalent and divalent metals.Since there is no evidence for the formation of trifluoride complexes of l a n t h a n ~ n s ' ~ we attempted to measure the stability constants PI and Bz for the mono- and di-fluoride complexes of lanthanons using a simplified potentiometric technique. Unsaturated solution mixtures in 0.5 mol dmP3 NH,NO, ( I = 0.5 mol drn-,) were prepared by mixing various amounts of lanthanon and a constant amount of F- to obtain a final concentration of 5 x rnol dmP3 F- and ( 1 .&I .5) x lo-, mol dm-, Ln3+. Sodium fluoride standards of appropriate concentration were also prepared in the same medium.The concentration of free fluoride ion and the pH of each mixture were measured at room temperature, 25 0.5 "C to compute the stability constant, PI of the monofluoride complex. In early experiments the Orion ion analyser meter was used for measurement of the concentration of the free fluoride. One of the problems we experienced in using the meter is to know when the specific fluoride-ion electrode reached equilibrium with fluoride ion in solution. It takes quite some time to reach this equilibrium and record a constant reading on the meter. The other problem was the inability to measure from the instrument the linearity or slope of the calibration curve. We used only one standard for our measurements with the meter. We recently acquired an expandable and digital Orion ion analyser E 920 which uses two standards for calibration and lights a ready signal when the electrode-F- ion equlibrium is reached. Furthermore, it also provides the slope of the calibration curve.We have therefore, used the new instrument also to measure the stability constants, p1 and /Iz of the mono- and di-fluoride complexes, respectively, of lanthanons. For the measurement of p2 for LnFk, unsaturated solution mixtures containing 5 x mol dm-, Ln3+ and varying amounts of F- (< 5 x lo-, mol drn-,) were prepared in 0.5 mol dmP3 NH,NO, (I = 0.5 mol drn-,). After mixing the solutions thoroughly, the concentration of the free fluoride ion and pH of each solution were measured. Measurement of the Solubilities as a Function of pH Buffer solutions of the same ionic strength but with pH ranging from 1 to 8 were prepared. Each of these solutions was saturated with the desired LnF, and the free fluoride ion concentration was measured by the ion analyser meter.Stoichiometrically one-third of this concentration represents the solubility of the lanthanon trifluoride in each buffer solution.2686 Solution Chemistry of Lanthanon Trijluoride- Water Results and Discussion Solubilities and Solubility Products The total solubility of lanthanon trifluoride without any allowance being made for the hydrolysis and/or complexation of lanthanide and fluoride ions are presented in table 1. The major processes that involve complex formation of Ln3+ and F- and the hydrolysis of the former are according to Vasile'v and Kojlovski5 the following: Ln3+ + F- + LnF2+; PI = (LnF2+)/(Ln3+) (F-) Ln3+ + H 2 0 g Ln(OH)2+ + H'; K , = (Ln(OH)2+) (H+)/(Ln3+) (1) F - + H + e H F ; K,, = (HF)/(H+)(F-) (2) (3) Mean activity coefficients are omitted in these equations as their values in dilute solutions (< Menon et al.', have shown that the concentration of the free metal ions and of free fluoride ions can be computed from the measurements of either the total concentration of the metal (radiometric and conductometric) or the concentration of free fluoride ion (potentiometric) using the reported values of P1, K,, and K,, and the measured pH of the saturated solutions.In this work we have used our measured values of P1 and the reported values of KH'* and KHF12 to compute the values for the concentration of free lanthanide ion or the fluoride at 25 "C from the appropriate solubility data.The thermodynamic solubility product was then computed using the relation : mol drn-,) are close to unity. where y + is the mean activity coefficient of the electrolyte. The negative logarithm of the solubility products (pK,,) so obtained from three different types of measurements are listed in table 1, and the literature values of pKsp are also included for comparison. It appears that most of the literature values of pK,, are based on the stoichiometric concentrations without any correction for complexation and hydrolysis. Although the contribution from the hydrolysis of Ln+ to its total concentration is minimal (ca. 1 %), that from complexation is significant (> 10 %).Although there is significant difference between stoichiometric and thermodynamic solubility products for a particular LnF,, there is only slight difference in their pKsp values. In most cases the literature cites only the pK,, values, but these values differ in some cases by as much as three units.2,4 The disagreement among pK,, values may be attributed to the difference in the aging of the precipitate2 or the method of preparation of the f l u ~ r i d e . ' ~ We have studied the effect of aging of the precipitate in the mother liquor, pH and nature of the precipitant on the solubility of DyF,. The results are shown in table 2. It is obvious from this table that all of these variables do affect the pKsp values. We have used 10 min aging time and NaF as precipitant in most cases. The pH of the saturated solutions varied from 5 to 6.Selected values of pK,, from literature and from our own work that do not differ by more than one unit (one order of magnitude in Ksp) were averaged to obtain reasonably accurate values for this parameter. The selected values for pKs,, their average values, the respective values for Ksp, the concentration of free lanthanide ion and the estimated activity coefficients used for computation are shown in table 3. The general trend in the variation of solubilities of LnF, in aqueous solution at 25 "C as a function of atomic number is depicted in fig. 1. Our solubility values and those computed from the average pKsp values show three maxima when plotted against the atomic number of lanthanons.No such data exist in literature for the whole lanthanide series for comparison. However, the data reported by Nikolaev et aL20 for the solubilities of the fluorides from EuF, to LuF, also shows a maximum at Er. In fig. 2, the variation of pKsp values for LnF, are plotted as a function of the atomicM . P. Menon and J . James 2687 Table 1. Solubilities and solubility products (pK,,) of lanthanon trifluorides in aqueous solutions at 25 "C lanthanon trifluoride LaF, CeF, PrF, NdF, SmF, EuF, GdF, TbF, DYF, HoF, ErF, TmF, YbF, LuF, total solubility" / l 0-5 mol dm-, pK,, a pK,, ' 7.37 f 0.24 (1 3)c' 7.61 f0.28 (13)' 7.05 f 0.68 (13)" 4.58k0.61 (12)c 4.57 f 0.55 (I 2)' 3.14k0.53 (12)R 3.01 k0.31' 2.09 _+ 0.17' 8.14 f 0.68 (1 3)c 11.13f0.72 (13)p 1.59 f 0.09 (1 3)R 4.82 f 0.72 (24)c 4.76 f 0.60 (24)' 1.69 f 0.32 (24)" 26.90 f 2.4 (8)' 34.00 f 2.90 (8)p 7.39 f 0.85 (8)R 2.96 & 0.45 (24)c 2.14f0.31 (24)' 1.92 f 0.40 (24)R 5.87 f 0.15 (9)c 4.17f0.17 (9)p 3.59 k0.52 (9)R 9.03 f 0.46" 8.66 f 0.88' 20.8 f 1.4" 15.90 f 1 .6OP 13.1 f 1.4 (1 1)' 11.10+2.90 (11)' 1.98 f 0.20 (1 l)R 4.23 f 0.66 ( 10)' 4.53f0.41 (10)' 1.49 f 0.04 ( 8.55+0.62c 6.85 f 0.54' 8.95 f 0.62R 27.7 f 1 .7r 37.5 & 3.2' 15.3 15.3 15.4 16.1 16.1 16.7 16.7 17.4 15.2 14.9 17.9 16.1 16.0 17.8 13.7 13.1 15.6 17.2 17.3 17.6 15.9 16.3 16.6 15.3 15.1 13.9 14.1 14.7 14.6 17.6 16.4 16.1 18.0 15.3 15.4 15.2 13.6 12.9 14.9 (4)C 24.0 (26)' 18.7 (5)R 19.2 (3)' 17.1 (4)' 19.1 (5)R 18.9 (3)' 17.0 (4)' 17.08 (22)' 18.6 (3)' 20.3 (5)R 17.9 (3)p 19.3 (3)T 16.0 (4)' 17.2 (3)p 21.9 (5)R 15.4 (4)' 16.8 (3)' 18.1 (3)T 15.3 (4)p 16.7 (3)p 18.0 (3)T 14.9 (4)' 16.3 (3)' 14.6 (4)' 17.6 (3)T 15.8 (3)' 14.6 (4)' 17.2 (3)T 15.5 (3)' 14.5 (4)' 15.8 (3)' 14.6 (4)' 17.1 (5)T 15.0 (3)' 14.6 (4)' 16.3 (3)T 15.0 (3)' 14.6 (4)' 16.4 (3)T 20.2 (3)T 18.0 (5)R C, Conductometric ; P, potentiometric ; R, radiometric ; T, thermochemical.a This work. Previous work. number of lanthanons. Two of these curves also show three minima corresponding to lanthanons that exhibit the maximum solubilities. The pK,, values reported by Vasile'v and Kozlovski5 also show an increase in pK,, values from La to Eu and a drop to Er. Fig. 3 depicts the variation of pK,, values of Ln(OH), and of oxidation potentials for Ln/Ln3+ systems [taken from ref. (18)] as a function of the atomic number of2688 Solution Chemistry of Lanthanon Trijluoride- Water Table 2.Effect of aging, pH and nature of precipitant on the solubility and pK,, of dysprosium trifluoride" aging time solubility solubility /min pH / mol dmP3 pKsp pH / lop5 mol dmP3 pK,,, * NaF as precipitant 10 4.8 8.66 C 30 - 10.90 60 - 20.0 90 - 37.4 120 - 28.4 150 - 24.3 180 - 29.8 HF as precipitant 14.96 4.18 2.78 16.79 5.30 4.13 16.10 14.57 3.97 2.98 16.75 5.27 5.13 15.73 13.42 4.18 1.57 17.78 5.26 2.20 17.27 12.55 3.92 1.65 17.76 5.25 4.10 16.2 1 13.00 -d 2.49 17.06 17.28 13.25 - 2.18 12.92 - 2.4 1 17.11 a Concentration of fluoride ion was measured by potentiometric method. * Stoichiometric value for Ksp was computed. "H of these samples is ca. 5. pH of these samples is ca.4. Table 3. Averages of selected values of pK,, from literature and from this work" lanthanon selected values average value activity [L~,+I, fluorides of PK,, of pKsp coefficient 1 O1'KSp / 1 0-5 mol dmW3 LaF, 15.3, 15.3, 15.4, 14.9 CeF, 16.1, 16.1, 16.7, 17.1 PrF, 16.7, 17.4, 17.0 NdF, 17.8, 17.9 SmF, 16.1, 16.0, 16.0 EuF, 15.6, 15.4 GdF, 17.2, 17.3, 17.6, 16.8 TbF, 15.9, 16.3, 16.6, 16.7, DYF, 15.3, 15.1, 16.3, 14.6 HoF, 13.9, 14.1, 14.6 ErF, 14.7, 14.6, 15.5, 14.5 TmF, 16.4, 16.1, 15.8 YbF, 15.3, 15.4, 15.2, 15.0, 14.6, 16.3 LuF, 13.6, 15.0, 14.6 14.9 15.2 16.5 17.0 17.9 16.0 15.5 17.2 16.1 15.3 14.2 14.8 16.1 15.3 14.4 0.93 1 0.953 0.958 0.939 0.944 0.917 0.958 0.945 0.9 18 0.9 14 0.928 0.948 0.935 0.919 63.10 3.16 1 .oo 0.13 10.00 3 1.60 0.63 7.94 50.1 63 1 .O 158.5 50.1 7.94 398.0 7.47 3.45 2.58 1.56 4.65 6.38 2.30 4.38 7.15 13.53 9.43 4.37 7.02 11.99 "The pK,, values from the literature and from our work that do not differ by more than one unit were averaged.The selected values are shown in the table. lanthanons. These graphs also do not follow a pattern of gradual increase or decrease in the respective quantities with increase in atomic number. Thermodynamic Functions for the Dissolution of LnF, in Water at 25 "C The only enthalpies of dissolution of lanthanon trifluorides available in literature seem to be those reported by Afanas'ev et aL21. 22 Owing to the low solubility of LnF, in water they used a solvent with composition, HC1- 5.05H20 - 0.02H3BO, to measure enthalpies of solution of eight lanthanon trifluorides at 50 "C with the aid of a microcalorimeter.M.P. Menon and J. James 2689 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 La Fe Pr Nd Pm Sm Eu Gd Tb D Ho Er Tm Yb Lu 57 58 59 60 61 62 63 64 65 6g 67 68 69 70 71 lanthanides and their atomic number Fig. 1. General trend in the solubilities of LnF, in aqueous solution at 25 "C as a function of their atomic number. 0, Our average values from conductometric and potentiometric measurements ; 0 , average values from selected literature and our data; A, reported by Nikolaev et 11 10 ;;; La Fe Pr Nd Prn Sm Eu Gd Tb Dy Ho Er Trn Yb Lu 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 lanthanides and their atomic number Fig. 2. General trend in the variation of pK,, values of LnF, in aqueous solution at 25 "C as a function of atomic number of lanthanons.0, Our values from conductometric and potentiometric measurements; 0, average values from selected literature and our data; A, reported by Va~il'ev.~ We have used the temperature dependence of the K,, to measure the standard enthalpy change for the dissolution process. Such temperature-dependent graphs of K,, have been reported in previous publications.s-lO* 12, l3 . The enthalpy changes were measured from the slopes of these straight-line graphs using the van't Hoff equation. The K,, values at 25 "C measured from these graphs were used t o compute the standard free energy2690 Solution Chemistry of Lanthanon TriJluoride- Water " 2.4 3 ! 2.3 *Y 2.2 4 Li' 2.0 c +E: 2.1 * c 1.9 0 4 1.8 La Fe Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 lanthanides and their atomic numbers Fig.3. General trend in the variation of pK,, values of Ln(OH), and of oxidation potentials for Ln/Ln3+ system in aqueous solution as a function of atomic number (0) oxidation potentials, (0) pK,, of Ln(OH), (data taken from Morss18). Table 4. Values for the thermodynamic functions for the dissolution of lanthanon trifluorides in water at 25 "C lanthanon fluorides AH"/kJ mol-la AG"/kJ mol-la ASo/J K-' rnol-la LaF, CeF, PrF, NdF, SmF, EuF, GdF, TbF, HoF, ErF, TmF, YbF, LuF, DYF, 59.7 k 2.9 53.0 & 17.4 54.2 f 2.1 43.4 f 3.1 51.0f 1.8 39.8 k4.0 50.8 k 8.0 27.4 & 0.5 56.1 k 9.9 36.4 f 2.8 34.2& 3.5 14.7f0.7 135.7 f 2.3 59.8 f 4.0 89.9 k4.7 91.7 f. 0.1 97.2 k 1.7 90.7 & 5.3 91.3f 1.0 86.9 f 8.0 97.6 f 2.3 93.8 k4.7 85.1 f 0.3 78.2 f 1.4 99.6 k 18.2 102.3f 10.1 86.5 k 1.0 73.0f 1.4 - (92.4 f 7.3) - ( 129.7 f 58.3) - (145.0 f 2.0) - (141.5 f.8.8) - ( 1 3 5.3 f 7.0) -(157.0f 16.0) -(157.0f 19.0) -(222.8 f 16.0) - (97.8 f 34.1) -(140.1 k 14.2) - (21 7.3 k 62.0) - (294.2 f 34.1) -(44.4 & 18.0) +(165.2+ 11.1) These are averages of the values obtained from conductometric and potentiometric measurements. change and these values, in turn, were used to calculate the standard entropy change for the dissolution of the fluoride. The average values obtained from the conductometric and potentiometric data for the above three functions are listed in table 4. The enthalpy and free-energy changes are both positive for all the systems, but the entropy changes, except for YbF,, are all negative.The enthalpy change for YbF, system is also very high when compared with the rest. Krestove and EgorvaZ3 have reported negative entropyM. P. Menon and J . James 269 I changes for the dissolution of all lanthanon trifluorides and attribute the negative entropy change to the structural changes of water in the region of long-range hydration of the ions. It appears that the hydration of the lanthanide ions play an important role in the solution properties of lanthanon fluoride-water systems. Stability Constants for the Mono- and Di-fluoride Complexes of Lanthanons The stability constant, pl, for the monofluoride complex was computed using the expression :24 ( 5 ) [LnF2+] ([F-I,,,. - [F-] - 6,[H+] [F-] - 26,[H+] [F-],) [Ln3+] [F-] - ” = {Ln3+] [F-] - where [Ln3+] and [F-] are the concentration of free lanthanide and free fluoride ions, respectively, and 6, and 6, are the association constants (6, = 1160, 6, = 4.8 x lO4)lZ for the formation of H F and HF;.The concentration of Ln3+ was obtained by subtracting [F-]comp,ex. from [Ln3+Itot.. The method is based on the fact that when the lanthanide ion is taken in large excess over that of fluoride ion, the only possible fluoride complex is LnF2+ in the mixture. The value for p1 reported in this work is the average of six measurements using an ion analyser meter and/or a digital ion analyser. On the assumption that Ln3+ forms only the mono- and di-fluoride complexes we have derived the following equation for the computation of p2, from the analytical data:25 a- (1 -A)Pl[F-] p2 = (1 -H)[F-], where p1 is the stability constant for the monofluoride complex, [F-] the measured concentration of free fluoride, A the average number of fluoride ions that complex with one lanthanide ion.The value of n is given by the following equation:25 Experimentally measured values of p1 and other quantities were used to compute p2 analytically with the use of eqn (6) and (7). Eqn (6) may be converted into a linear form, to determine p1 and p2 simultaneously, as follows: (8) n (2 - a) [F-] = & + I 3 2 ( 1 - q - [F-](1 - H ) The intercept and the slope of the plot of the term on the left against ( 2 - ~ ) [F-]/(l -n) were used to determine the values of p1 and p2 simultaneously by a linear-regression method.The stability constants for the mono- and di-fluoride complexes of lanthanons measured in this work at an ionic strength of 0.5 mol dm-3 along with the literature values are presented in table 5. Becker and Bila17 and also Gmelin Handbuch list the stability constants for the monofluoride complex at different ionic strengths reported by various authors. In general, the value for Dl tends to increase with decrease in ionic strength. We feel the values we obtained with our digital ion analyser are more precise than the ones we measured with ion analyser meter, In some cases there are differences between the previously reported values and ours. Our values for p2 are, in general, not as precise as those for P1, as indicated by the error limits. Errors associated with the results are standard deviations from the mean of 8-10 individual measurements.Our values are much higher than those reported in literature for a few lanthanons at an ionic strength of 0.5 mol dm-3.’. 16, 24 We have also included in parentheses the values obtained for the complexes of a few lanthanides by the linear regression method using eqn (8). Corresponding r values are also given. The results with r values < 0.7 were omitted. Previous authors have used a pH of 3.6 and a higher concentration of Ln3+ and/or F- ions in their experiments. It is important that the measurement of free fluoride ion2692 Solution Chemistry of Lanthanon Tr fluoride- Water Table 5. Stability constants for the mono- and di-fluoride complexes of lanthanons in aqueous solutions of I = 0.5 mol dm-3 at 25 "C lanthanon methoda La Ce Pr Nd Sm Eu Gd Tb DY Ho Er Tm Yb Lu A B C D A B C E A B A B D A B A D C A B C D E A D A E A B A D E A A B E A D E P* this work previous work P, this work previous work ( x 10-6) ( x 10-6) 453 f 29 - - 427f 17 (13) - 485f 10 (16) - 600f 100 (7) - 997 +98 (12) - 1288 (6) 790 k 44 - - - 912+ 12 - 637 67 - 618k51 - - 724 (1 3) - 670f 100 (7) 1416+41 - - 1076+ 14 (24) 1172k54 (1025, r = 0.91)' - 2455 (6) - 2512 (6) 1638 61 - - 1215f70 (24) - 2692 (6) - 2344 (6) - 2344 (6) - 2320 f 50 (7) - 3802 (6) 1720 f 50 (2378, r = 0.72) 2390 k 79 (2004, r = 0.85) 2028f74 (1911, r = 0.97) 1269 f 41 - 2452 f 42 (2406, r = 0.92) 1770 f 50 (7) - - 3467 (6) 1784 k 35 (1 766, r = 0.99) 203 1 f 53 (2752, r = 0.97) 1900 f 72 - 2122k58 (2022, r = 0.92) - 4074 (6) 1730 + 50 (7) 2951 (6) - - 13.5f 1.4' - - 15.40k3.1 (13) - 0.15 ( 6 ) 3.7 + 1.6 - - 0.10f0.04 (12) - 0.91 (6) 7.5 k 2.1 - 3.7k0.8 - 4.1 f0.8 - - - - - - 25.10f2.9 (13) - - 9.0 & 3.7 - 1.9k0.6 (2.38, r = 0.91)' 2.95 (6) 3.6+ 1.1 - 1.10f 1.1 (24) - - - - - 1.4f 1.2 (24) - - - - - - 4.6f 1.7 (3.38, r = 0.72) 0.13 f 0.04 (7) - 2.7f 1.1 (3.75-tO.85) - - 6.4f 1.0 (6.6, r = 0.97) 4.3f 1.5 - 10.8f2.7 (1 1.0, r = 0.92) - - - - 8.7f0.9 (8.7, r = 0.99) 8.9+ 1.2 (7.8, r = 0.97) 5.7 + 3.0 - 4.5f 1.5 (4.6, r = 0.92) - - - - - - - a Method (A) potentiometric F- ion concentration measurement using an expandable and digital Orion ion analyser and F- ion-selective electrode, (B) potentiometric F- ion concentration measurement using an Orion ion analyser meter, (C) solvent extraction using radioisotopes, (D) potentiometric titration using LnF, electrode, (E) potentiometric titration using FeT1/FelIT electrode.bThese are analytical results obtained using eqn (6). cThese are results obtained by a linear-regression method using eqn (8).M . P. Menon and J . James 2693 3.0- 0 'E '0 2.5- - --- E I 5 2.0- E .& 1.5- .- 8 1.0- - .6 3 3 0.5- z .- - 0.0 - ip 1 1 1 1 1 1 1 1 1 1 1 2 3 4 5 6 7 8 9 10 PH Fig. 4. pH dependence of the solubility of LnF, (potentiometric measurement). 0, LaF,; 0, TmF,; A, TbF,. concentration be made in unsaturated solution. The measured pH of each of our mixtures was in the range 4.5-5.0. pH Dependence of the Solubility of Lanthanon Trifluorides Fig. 4 shows the solubility versus pH of a few lanthanon trifluorides in various buffer solutions of ionic strength equal to 0.5 mol dmP3.The solubilities of all lanthanon trifluorides in buffered solutions were measured by the potentiometric method. In all cases the solubility was found to reach a minimum value at a pH of ca. 4. Menon' has shown that the effect of complexation of F- with H' and Ln+ on solubility will be minimum at this pH so that the fluoride will be least soluble at this pH. Conclusions The solubilities and pK,, values of lanthanon trifluorides do not follow a pattern of gradual increase or decrease with the atomic number of lanthanons. The solubilities and the averages of selected pK,, values from literature and from our own work, that do not differ by more than unit, show three maxima when plotted against the atomic number of lanthanons.The wide difference in the reported pK,, values may be attributed to the difference in the aging of precipitate, techniques used for the measurement and also in the evaluation methods. The solubilities of all lanthanon trifluorides in buffered solutions of constant ionic strength (0.5 mol dmP3) show a minimum at a pH between 3 and 4. Standard enthalpy and free energy changes for the dissolution of the fluorides in water are both positive while the entropy change, except for YbF,, are negative. Although the stability constant for the monofluoride complexes can be measured with great precision, those for difluoride complexes have considerable error in measurement. The financial support provided by the U.S. Department of Energy through their grant no.DE-FG09-84SR 140 14 for this work is gratefully acknowledged. Acknowledgement2694 Solution Chemistry of Lanthanon Trijluoride- Water is also made to the donors of the Petroleum Research Fund, administered by the American Chemical Society for support of initial phases of this research through their grant no. 11218-83. The authors are also thankful to numerous undergraduate chemistry majors who contributed to this work at different times. References 1 IUPAC Solubility Data Project Combined Newsletter (Pergamon Press Oxford, February 1980). 2 J. Burgess and J. Kiljowski, Adu. Znorg. Chem., 1981, 24, 57. 3 J. J. R. Frausto da Dilva and M. M. Queimado, Rev. Port. Quim., 1973, 15, 29. 4 G. N. Koroleva, S.A. Gava, N. S. Poluektov, A. I. Kirillov and M. E. Kornelli, Dokl. Akud. Nuuk 5 V . P. Vasil'ev and E. V. Kozlovskii, Russ. J. Znorg. Chem. (Engl. Transl.), 1977, 22, 472. 6 Gmelin Handbuch der Anorganischen Chemie (Springer-Verlag, Berlin, 8th edn, 1976), no. 39, part C3, 7 P. Becker and B. A. Bilal, J. Solution Chem., 1985, 14, 4067. 8 M. P. Menon, J. Radioanal. Chem., 1981, 63, 283. 9 M. P. Menon, J. Chem. Eng. Data, 1982, 27, 81. 10 M. P. Menon, Int. J. Appl. Radiat. Isotop., 1982, 33, 1375. 1 1 M. P. Menon, J. Radioanal. Nucl. Chem. Lett., 1985, 96, 31 1. 12 M. P. Menon, J. James and J. D. Jackson, J. Radioanal. Nucl. Chem., 1986, 102, 419. 13 M. P. Menon, J. James and J. D. Jackson, Lanthanide and Actinide Res., 1987, 2, 49. 14 J. B. Walker and G. R. Choppin, Thermodynamic Parameters of Fluoride Complexes of the Lanthanides in the LanthanidelActinide Chemistry, Adv. Chem. Ser. 71 (ACS, Washington D.C., 1967), p. 127. 15 B. A. Bilal and J. Kob, J. Inorg. Nucl. Chem., 1980, 42, 629. 16 A. Aziz and S. J. Lyle, J. Inorg. Nucl. Chem., 1970, 32, 1925. 17 G. T. Hefter, C. B. Chan and N. H. Tioh, Anal. Chem., 1984, 56, 749. 18 L. R. Morss, Yttrium, Lanthanum and Lanthanide Elements in Standard Potentials in Aqueous Solution, 19 A. I. Popov and G. E. Krundson, J. Am. Chem. SOC., 1954, 76, 3921. 20 N. S. Nikolaev, Sh. A. Abdurakhmanov and Kh. Sh. Dzhuraev, Russ. J. Inorg. Chem., 1974, 19, 618. 21 Yu. A. Afans'ev, E. I. Khanaev and M. G. Kotov, Radiokhimiya (Engl. Transl.), 1975, 17, 203. 22 T. P. Storozhenko, E. I. Khanaev and Yu. A. Afans'ev, Russ. J. Phys. Chem., 1975, 49, 1241. 23 G. A. Krestov and I. V. Egorova, Radiokhimiya (Engl. Transl.), 1970, 12, 903. 24 B. A. Bilal and P. Becker, J. Znorg. Nucl. Chem., 1979, 41, 1607. 25 M. P. Menon, J. James and T. L. Hill, Int. J. Appl. Radiat. Isotop., 1988, 39, 949. 26 J. L. Weaver and W. G. Purdy, Anal. Chim. Actu, 1959, 20, 376. SSSR., 1976, 228, 1384. p. 123. ed. A. J. Bard, R. Parsons and J. Jordon (Marcel Dekker, New York, 1985), p. 587. Paper 8104465A ; Received 4th November, 1988

ISSN:0300-9599    

DOI:10.1039/F19898502683

出版商:RSC    

年代:1989

数据来源: RSC

摘要:

J . Chem. SOC., Furuduy Trans. I , 1989, 85(9j, 2695-2703 Solubility of Hex- 1 -ene, Hexane and Cyclohexane in Liquid Nitrogen7 Elzbieta Szczepaniec-Cieciak" and Magdalena Kurdziel Cryogenic Department, Faculty of Chemistry, Jagiellonian University, u1.M. Karasia 3, 30460 Cracorv, Poland Lucyna Ulman Regional Laboratory of Physicochemical Analysis and Structure Research, Jagiellonian University, u1.M. Karasia 3, 30460 Cracow, Poland The solubilities of hex- I-ene, hexane and cyclohexane have been measured in liquid nitrogen at a temperature of 77.4 K by the filtration method. The solubilities of the C, hydrocarbons in LN, at 77.4 K vary from 0.5 x mole fraction for cyclohexane, to 0.7 x lop8 mole fraction for hexane and 0.16 x lo-, mole fraction for hex-1-ene. The results are compared with the solubilities calculated on the basis of the Scatchard- Hildebrand equation of regular solution theory.Correlation between the solubilities of alkanes, alkenes and cyclic hydrocarbons in liquid nitrogen, and some properties of solutes {normal boiling point (q), critical point (q), heat of vaporization at normal boiling point (AHhj, the mean of the heat of vaporization and the enthalpy of melting [(AHh+AH,)/2] and Lennard-Jones force constant (&/kH) are presented. In recent years, there has been a considerable interest in the volatile organic compounds and, among them, the non-methane hydrocarbons (NMHC) present in the atmosphere. The amounts of NMHC found in air are normally in the low-ppb range or even lower. The abundance of these substances may vary considerably between polluted urban and industrial areas and remote rural and maritime areas, as well as at different altitudes of the tr0posphere.l The principal sources of hydrocarbons in atmospheric air are internal combustion engines, the petroleum refining industry, metallurgical and coke-making industries, power generation, as well as solvent evaporation.Separation of air by cryogenic methods requires the removal of those impurities in air which may undergo solidification in the low-temperature part of the installation and in consequence would cause blockages and soon make the process inoperative, leading to the blanketing of heat-transfer surfaces, corrosion and even an explosion hazard. The aim of the present work was to investigate the solubility of chosen C , hydrocarbons (hexane, hex- 1-ene and cyclohexane) in liquid nitrogen in order to supplement the data on the solubility of hydrocarbons in liquid nitrogen.In table 1, a review of experimental data concerning the solubility of C, hydrocarbons in cryogenic liquids is presented. The studies on the solubility of C, hydrocarbons in LN, were planned on the basis of selected properties of pure components of the solutions (table 2) and on the calculations of the solubility from the Scatchard-Hildebrand equation (by the Preston-Prausnitz met hod). 3 3 l7 Since it was expected that the solubility of the C, hydrocarbons in LN, would be very small ( 10-6-10-8 mole fraction) a filtration method was used. t Paper presented at the Third International IUPAC Symposium on Solubility Phenomena, held at the University of Surrey, 23--26 August, 1988.26952696 Solubility of Hex- 1 -ene, Hexane and Cyclohexane Table 1. Summary of literature data of the solubility of C, hydrocarbons in cryogenic liquids solute solvent temperature range/K ref. hexane hex- 1 -ene nitrogen 77.4 7 5-92 { 7;: oxygen argon 90-100 124.3-176.8 103.4-1 50.0 { 93.8-163.7 methane 161.3-176.2 this paper 2 3 4 5 6 7 8 9 nitrogen 77.4 this paper oxygen 77.9-92.5 3 argon 9&100 5 cyclohexane nitrogen 77.4 this paper argon 90- 100 5 methane 99.7-1 50.0 154.0-1 90.0 7 8 2,3-dimethylbutane argon 92.3-109.3 18 methane 100-123 18 2-methylpent- 1-ene argon 90- 100 5 4-methylpent-2-ene argon 90-1 00 5 hex-2-ene argon 90- 1 00 5 methylcyclopentene argon 90-100 5 benzene oxygen 77.3 2 argon 10 1-1 45 10 103.8-1 53.1 7 162.8- 185.4 I 1 99.4- 199.8 8 122-160 10 methane krypton Table 2.Selected properties of pure components of so1utions12-16 property hexane cyclohexane hex- 1 -ene nitrogen molecular weight boiling point under pressure of 1013.25 hPa, T,/K melting point, T,/K transition temperature, q,,,,/K critical temperature, TJK critical pressure, P,/ lo5 Pa critical volume, K/cm3 mol-' Pitzer acentric factor, ~ r ) dipole moment, p/D heat of vaporization at normal boiling point, AH,,/kJ mol-' heat of melting at normal melting point, AH,/kJ mol-' Lennard-Jones force constant, ( c / k , ) / K ASJR Astrans/' 86.178 341.9 84.162 353.9 84. I62 336.6 28.013 77.4 177.8 8.8 1 - 507.4 29.7 370 0.296 0.0 28.85 13.03 399.3 279.7 186.1 553.4 40.7 308 1.15 4.36 0.213 0.3 30.08 2.929 297.1 133.3 8.43 504.0 31.7 350 0.285 0.4 28.03 9.347 63.5 35.5 126.2 33.9 89.5 0.04 0.0 3.577 1.37 0.78 0.721 71.4E. Szczepaniec- Cieciak, M.Kurdziel and L. Ulman 2697 The purpose of the present work was also to develop a new version of apparatus to investigate the solubilities of solidified substances in liquid nitrogen by means of the filtration method. Experimental Method Determining the solubilities of the heavier hydrocarbons in liquid nitrogen by means of a modified filtration method is carried out in the following stages: (1) saturating liquid nitrogen with hydrocarbon, together with separating the saturated solution of the hydrocarbon in liquid nitrogen from the excess of solidified hydrocarbon by filtration ; (2) evaporating liquid nitrogen from the solution ; (3) preparing a solution of the hydrocarbon in an organic solvent (1,4-dioxane), with an addition of an internal standard (benzene) ; (4) determining the hydrocarbon content in the 1,4-dioxane solution by means of gas chromatography.Apparatus The apparatus used to determine the solubility of solidified hydrocarbons in liquid nitrogen with a modified filtration method consists of the following parts : (A) solidified hydrocarbon generating unit (fig. 1) ; (B) unit producing a saturated solution of the hydrocarbon in liquid nitrogen (fig. 2); (C) saturated solution evaporation unit (fig. 3). Determination of the Hydrocarbon Contents Solidified hydrocarbon was obtained in fine globules by pouring it slowly into a perforated vessel (fig. 1) immersed in liquid nitrogen.The solidified hydrocarbon was placed in unit B (fig. 2) by screwing the vessel to the nut attached to the tip of the electrical stirring rod and inserting both in a Dewar flask filled with liquid nitrogen. The solution was stirred vigorously for 5-6 h and then the vessel with solidified hydrocarbon was withdrawn and the Dewar flask was stoppered with a plug, fitted with a heating element to speed up liquid-nitrogen evaporation and with a connection for a gas meter (fig. 3). A coil immersed in a warm-water bath was placed between the Dewar and the gas meter thus preventing the water in the gas meter from cooling excessively. When the liquid nitrogen had completely evaporated, the final reading of the gas meter was taken and the volume of gas was converted to STP.After disconnecting the gas meter, the inside of the Dewar flask was rinsed with 10 cm3 of 1,4-dioxane with an appropriate amount of benzene as an internal standard in order to wash away the hydrocarbon. The solution was quantitatively transferred to a measuring flask, in which the hydrocarbon content was determined using gas chromatography. [Perkin-Elmer F-21 gas chromatograph with a flame-ionization detector (FID)]. The conditions of gas chromatography analysis were : stainless-steel column, 3 m long and with an internal diameter of 4 mm, filled with 4% OV-17 and 6% QF-1 on a Gas Chrom P 60/80 mesh (Applied Science Laboratories Inc.); carrier gas: nitrogen, rate of flow 50 cm3 min-'; column temperature: 80 "C (353 K); injection volume of sample: 5 x Another procedure, with different apparatus and analytical technique were also employed to determine the solubility of solid hexane in liquid nitrogen. The apparatus and experimental details are described elsewhere.l9 In order to separate the precipitates of solid hexane, the solution was filtered through cm3.2698 Solubility of Hex- 1 -ene, Hexane and Cyclohexane Fig. 1. Diagram of unit A for obtaining solidified hydrocarbon : 1, Dewar flask ( I .5 dm3); 2, liquid nitrogen; 3, perforated polyethylene vessel ( 1 mm hole size); 4, cotton cloth pouch; 5, ceramic beads; 6 , solidified hydrocarbon ; 7, wooden handle. 1 Fig. 2. Diagram of unit B for obtaining a saturated solution of a hydrocarbon in liquid nitrogen: 1 , Dewar flask (2 dm3); 2, liquid nitrogen; 3, perforated polyethylene vessel ( 1 mm hole size); 4, cotton cloth pouch; 5, ceramic beads; 6, solidified hydrocarbon; 7, nut; 8, Teflon blades; 9, electrical stirring rod.E.Szczepaniec-Cieciak, M. Kurdziel and L. Ulman 2699 6- 8 7 7- Fig. 3. Diagram of unit C for evaporating the saturated solution: 1, Dewar flask (2 dm3); 2, saturated solution of the hydrocarbon in liquid nitrogen; 3, heating element; 4, plug; 5, plastic filler sealing; 6, power supply; 7, thermostat; 8, coil; 9, gas meter. a tube, equipped with a sintered glass filter of average pore diameter 9 pm. The filtrate was collected in a glass flask placed in a vessel with liquid nitrogen. After the liquid nitrogen had completely evaporated, the residue of hexane from the flask walls was washed out with a known volume of 1,4-dioxane containing a known amount of a different internal standard, i.e.cyclohexane. The determination of the hexane concentration was carried out using a gas chromatograph coupled with a LKB- 9000s mass spectrometer. The conditions for separation of the substances investigated were: column 282 cm long and with an internal diameter of 4 mm, filled with 3 YO Benton, 3 % SE-30 on a Chromosorb G 60/80 mesh; carrier gas: helium, rate of flow 10 cm3 min-' ; column temperature : 80 "C (353 K) ; molecular separator temperature : 160 "C (433 K); flash-heater temperature : 180 "C (435 K) ; injection volume of sample : 50 x lop3 cm3. Reagents The following reagents were used without further purification : hexane and cyclohexane (spectroscopy grade, Merck, Darmstadt), hex- I-ene (purris, Koch-Light), 1,4-dioxane (spectroscopy grade, Serva, Feinblochemie, Heidelberg and Fluka AG, Buchs), benzene (spectroscopy grade, BDH), liquid nitrogen [quality grade (99.8 NJ, PPH Polgaz, Siewierz, Poland].Results and Discussion The results of the solubility determinations are given in table 3. A comparison of the solubility values of the six-carbon hydrocarbons under investigation (representatives of three homologous series, namely alkanes, alkenes and cycloalkanes) yields the following order of increasing solubility in liquid nitrogen at 77.4 K : cyclohexane (0.5 x lops mole fraction), hexane (0.7 x lop8 mole fraction) and hex-1-ene (0.16 x lop6 mole fraction).The solubilities of hexane and cyclohexane are2700 Solubility of Hex- 1 -ene, Hexane and Cyclohexane Table 3. Results of solubility determinations for solid hexane, hex- I -ene and cyclohexane in liquid nitrogen at 77.4 K property hexane hex- 1 -ene cyclohexane experimental values of solubility x, (mole fraction) mean value of solubility X, (mole fraction) amount of hydrocarbon pg m-3 N, results of solubility calculation^^^ (mole fraction) 1.5 x 0.7 x lop8 0.2 x 0.3 x 1.3 x 0.5 x 0.7 x lop8 0.5 x lop8 26.6 26.9 1.4 x lop8 0.16 x 0.10 x 0.21 x 0.15 x 0.18 x 0.16 x lov6 0.04 x 10.0 601.2 1.97 x 0.8 x 0.3 x 0.2 x 0.6 x lo-* 0.5 x lo-' 0.5 x 0.2 x 19.6 18.8 0.97 x lop8 The results were obtained using different apparatus and different analytical technique (g.c.-m.s.). I 1 I 2 3 4 5 6 -91 ' 1 number of carbon atoms Fig.4. Relationship between the solubilities of (0) alkanes and (A) alkenes in liquid nitrogen (77.4 K) and the number of carbon atoms in a carbon chain. The data for the solubilities of the alkanes and alkenes in liquid nitrogen are given in table 4. very similar, while the solubility of solidified hex-1-ene in liquid nitrogen at the same temperature is more than an order of magnitude higher than that of the remaining hydrocarbons. In table 3, the values of the solubility of solid hexane, hex-1-ene and cyclohexane calculated from the Scatchard-Hildebrand equation are also given.17E. Szczepaniec- Cieciak, M . Kurdziel and L. Ulman 270 I Table 4. Correlations between the solubilities of solid hydrocarbons in liquid nitrogen at 77.4 K and the properties of solutes - solubility .GS" (mole AHtl no.solute fraction) ref. 7JK T,/K /kJ mol-' /kJ mol-' (&/k,,)/K 1 2 3 4 5 6 7 8 9 10 1 1 methane 6.27 x 10-' (20), (21) ethane I .46 x lo-' (22) propane I .08 x 10-3 (22) butane 1.02 x 10-5 (22) isobutane 1.22 x (22) hexane 0.7 x this paper ethene 4.41 x (23), (24) propene 7.78 x lo-' (24) but-1-ene 1.17 x (24) hex-1-ene 0.16 x this paper cyclohexane 0.5 x lop8 this paper 111.7 184.6 231.1 272.7 261.3 341.9 169.4 225.4 266.9 336.6 353.9 190.6 305.4 369.8 425.2 408.1 507.4 282.4 365.0 419.6 504.0 553.4 8.18 14.70 18.77 22.39 20.30 28.85 13.54 18.42 21.92 28.03 30.08 4.559 8.778 10.647 13.525 12.92 20.94 8.445 10.71 12.884 18.688 1 6.504 148.6 215.7 237.1 531.4 330.1 399.3 224.7 298.9 297.1 The experimental results for hexane and cyclohexane are somewhat lower (approxi- mately half as high) than the solubilities calculated from the Scatchard-Hildebrand equation, while for the hex- 1-ene-nitrogen system the experimental result is more than an order of magnitude lower than the calculated result.The experimental solubilities of hexane and hex-1-ene in liquid nitrogen at 77.4 K were compared with the respective values for other hydrocarbons belonging to the alkane and alkene homologous series. Fig. 4 shows the relationships between the experimental solubilities of alkanes and alkenes in liquid nitrogen and the number of carbon atoms in a carbon chain. In the homologous series of alkanes and alkenes solubility decreases as the molecular weight increases, corresponding to the increase of boiling point and heat of vaporization of solute.For the first four hydrocarbons of the linear molecule alkane homologous series, the drop in solubility with the growth of the number of carbon atoms in a molecule is approximately an order of magnitude per atom. This relationship is true only for linear molecules. The solubility of isobutane is ca. an order of magnitude higher than that of butane. For hexane, the drop in solubility with the number of carbon atoms in a molecule is greater than that for the first four hydrocarbons of the alkane homologous series. For alkenes the decrease of solubility with increasing number of carbon atoms in a carbon chain is smaller than that of alkanes having the same number of carbon atoms.Table 4 lists the experimental values of solubility of solid hydrocarbons in liquid nitrogen at the temperature of 77.4 K as well as the following parameters characteristic of the solute: normal boiling point (Q, critical temperature (q), heat of vaporization at normal boiling point (AH,), half the sum of the heat of vaporization and the heat of melting at normal melting point [(AH, + AHm)/2] and the Lennard-Jones force constant Fig. 5 shows the sample dependence between the logarithm of solubility (logx,) and the boiling points (TJ of solutes. Linear dependences of solubility and also critical points or heats of vaporization of solutes are a good approximation of experimental solubilities of solid alkanes, alkenes and cycloalkanes in liquid nitrogen.Grouped along the line are non-polar or slightly polar substances (permanent dipole moment of molecules not exceeding 0.4 D).? (Elks). t 1 D z 3.33564 x C m.2702 Solubility of Hex- 1 -ene, Hexane and Cyclohexane Fig. 5. The logarithm of the solubility in liquid nitrogen of hydrocarbons at 77.4 K, as a function of their normal boiling points. (0) Alkanes; (V) alkenes; (0) cycloalkanes. The data for the solubilities of the hydrocarbons, the boiling points and the numbers of solutes are given in table 4. The solubilities of solid hydrocarbons in liquid oxygen and liquid argon give similar correlation plots as the solutions in liquid nitrogen. A review of literature data concerning solid solubility in liquid oxygen and liquid argon is given in ref.(25) and (26). The authors thank Dr K. Nagraba for his help in developing the method of determining hydrocarbon content and to R. Dudzik for assisting in some of the measurements. The work has been financed by the Institute for Low Temperature and Structure Research, Polish Academy of Sciences, from the funds of central research and development project Applying Cryogenics in the National Economy. References 1 F. J. Reineke and K. Bachmann, J . Chromatogr., 1985, 323, 323. 2 C. McKinley and F. Himmelberger, Chem. Eng. Progr., 1957, 53, 112M. 3 C. McKinley and E. S. Wang, Advanced Cryogenic Engineering (Plenum Press, New York, 1960), 4 W. L. Ball, Safety Air Ammonia Plants, 1966, 8, 12. 5 R. G. Amamchyan, Zh. Fiz. Khim., 1976, 50, 513. 6 F. Kurata, Research Report RR-14 (Center for Research, Inc., Lawrence, Kansas, 1975) p.28. 7 A. Neumann and R. Mann, Kaltetechnik-Kfimatisierung, 1970, 22, 182. 8 G. P. Kuebler and C. McKinley, Aduanced Cryogenic Engineering (Plenum Press. New York, 1974). 9 E. Dickinson, C. M. Knobler and R. L. Scott, J . Chem. Soc., Faraday Trans. I , 1973, 69, 2179. vol. 4, p. 11. vol. 19, p. 320. 10 G. Ya. Zelikina, T. G. Meister and T. B. Mamchenko, Zh. Prikl. Spektrosk., 1980, 32, 629. 1 1 A. Neumann, R. Mann and W. D. Szalghary, Kaltetechnik-Kliniuti.c.ierung, 1972, 24, 145. 12 R. C. Reid, J. M. Prausnitz and T. K. Sherwood, The Properties qfGases and Liquids (McGraw-Hill, New York, 1977). 13 G. T. Preston and J. M. Prausnitz, Ind. Eng. Chem., Process Des. Dev., 1970, 9, 264.E. Szczepaniec- Cieciak, M. Kurdziel and L. Ulman 2703 14 CRC Handbook of Chemistry and Physics, ed. R. C. Weast (CRC Press, Boca Raton, Florida, 1986). 15 Landolt- Bornstein, Zahlenwerte und Funktionen aus Physic, Chemie, Astronomie, Geophysik und Technik (Springer-Verlag, Berlin, 1951), vol. 1, part 3; 1961, vol. 2, part 4; 1967, vol. 4, part 4a. 16 Poradnik, Fizyko-Chemiczny (Wydawnictwo Naukowo-Techniczne, Warszawa, 1974) (Handbook of Chemistry and Physics). 17 E. Szczepaniec-Cieciak, B. Dabrowska and J. M. Lagan, Cryogenics, 1977, 17, 621. 18 G. T. Preston, E. W. Funk and J. M. Prausnitz, J. Phys. Chem., 1971, 75, 2345. 19 E. Szczepaniec-Cieciak and Z. Wojtaszek, Proc. Sixth Intern. Cryog. Eng. Conf. (Grenoble, 1976), 20 V. G. Fastowskii and Yu. A. Krestinskii, Zh. Fiz. Khim., 1941, 15, 525. 21 M. H. Omar, Z. Dokoupil and H. G. M. Schroten, Physica, 1962, 28, 309. 22 E. Szczepaniec-Cieciak, V. A. Kondaurov and S. M. Melikowa, Cryogenics, 1980, 20, 48. 23 N. M. Tsin, Zh. Fiz. Khim., 1940, 14, 418. 24 E. Szczepaniec-Cieciak, V. A. Kondaurov and S. M. Melikova, Cryogenics, 1979, 19, 649. 25 E. Szczepaniec-Cieciak, Ciecze Kriogeniczne Jako Rozpuszczalniki Zestalonych Substancji (Drukarnia Uniwersytetu Jagiellonskiego, Krakow, 1984), pp. 1-207. (Cryogenic Liquids as Solvents of Solidified Substances. Jagiellonian University Press). p. 278. 26 E. Szczepaniec-Cieciak and B. Ciejek, Cryogenics, 1979, 19, 639. Paper 8/04469D ; Received 4th November, 1988

ISSN:0300-9599    

DOI:10.1039/F19898502695

出版商:RSC    

年代:1989

数据来源: RSC