摘要:

Journal of the Chemical Society, Faraday Transactions I ISSN 0300-9599Journal of the Chemical Society, Faraday Transactions I ISSN 0300-9599Journal of the Chemical Society, Faraday Transactions I SUBJECT INDEX-VOLUME 71, 1975 I, 1 Adsorption Active Sites in Zeolites. Part 4. n-Butene Isomerization over Deammoniated and partly Hydrolysed NHzY Zeolites. (Jacobs, Declerck, Vandamme & Uytterhoeven) . , Adsorption and Decomposition of Nitrous Oxide on y-Irradiated Magnesium Oxide. Adsorption of Carbon Dioxide on Goethite (x-FeOOH) Surfaces and its Implications for Adsorption of Non-swelling Vapours on the Surface of Cellulose. (Tremaine & Gray) . Adsorption of Propionitrile at the MercurylAqueous Solution Interface. (Abd-El-Mabey &Trasatti) . . . . . .. . . . . . Adsorption of Water Vapour by Calcium Fluoride, Barium Fluoride and Lead Fluoride. (Barraclough & Hall) . . . . . . . . . . . Calculation of the Adsorption Interaction of Argon and Krypton on the Relaxed (100-Faces) of Sodium Chloride and Potassium Chloride. (House & Jaycock) . . . . Carbon Monoxide Adsorption on Iron in the Temperature Range 85 to 350 K as revealed by X-Ray and Vacuum Ultraviolet [He(II)] Photoelectron Spectroscopy. (Kishi & Roberts) . . . . . . . . . . . . . Changes in the Sieving Action and Thermal Stability of Zeolite' A produced by Ion-exchange. (Takaishi, Yatsurugi, Yusa & Kuratomi) . . . . . . Dielectric Studies of Zeolite Systems. Part 1. Na+Fe3+-Y Zeolite with Polar and Non: dipolar Adsorbates. (Jones & Davies). .. . . . . Part 2. Ag-X Zeolite with Dipolar and Non-Dipolar Absorbates. (Jones) Electron-transfer at Alunlina Surfaces. Part 5. Oxidation of Aromatic Amines. (Flock- hart, Mollan & Pink) . . . . . . . . . Hydrogen Mordenite and Hydronium Mordenite. Ion Exchange and Thermal Stability. (Barrer & Klinowski). . . . . . . . . . . . Infrared Spectroscopic Study of the Adsorption of Nitriles on Aluminium Oxide. Fermi Resonance in Coordinated Acetonitrile. (Knoezinger & Krietenbrink) . . . Infrared Spectroscopic Study of the Adsorption of Water, Methanol and Acetone on Synthetic X and Y Zeolites. (Senkyr & Noller) . . . . . . Infrared Structural Study of Various Type L Zeolites. (Pichat, Franco-Parra & Bathomeur) . . . . . . . . . . . . . Infrared Study of the Adsorption of Oleic and Linolenic Acids onto the Surface of Silica immersed in Carbon Tetrachloride.(Marshall & Rochester) . . . . Infrared Studies of Reactions on Oxide Surfaces. Part 3. HCN and CzN2 on Silica. (Morrow & Cody) . . . . . . . . . . Infrared Study of the Interaction of Hydrogen Sulphide and Water with a Magnesium Oxide Surface. (Deane, Griffiths, Lewis, Winter & Tench) . . . . Infrared Study of the Interactions between the Surface of Silica and Methyl Fluorosulphate Vapour. (Eley, Kiwanuka & Rochester) . . . . . . Interlamellar Sorption of Ethanol and Montmorillonite Clays with Different Layer Charges. (Stul & Uytterhoeven) . . . Ion Exchange in Chabazite. Part 2. 'Comparison of Experimental and'Theoieticai Rates of Na+/K+ Exchange. (Duffy & Rees) .. . . Ion Exchange of Synthetic Zeolites X and Y with Co2+, NiZ+,Cu2+and Zu2+ Ions. (Mad & Cremers) . . . . . . . . . . . Ion-exchange Equilibria of Synthetic 13X Zeolite with Ni2+, Co*+, Zn2+ and Cd2+ Ions. (Gal &Radovanov) . . . . . . . . . Isotope Effect on Physical Adsorption on a Non-homogeneous Surface. (Moiseyev) . Mechanism of the Interaction of Hydrogen Sulphide with Adsorbed Oxygen on Lead Studied Molecular Theory of Adsorption in Pore Spaces. Part 1. Isotherms for Simple Lattice (Breakspere, Hassan & Roberts) . . . Anion Adsorption. (Russell, Paterson, Fraser & Farmer) . . , . . by X-Ray Induced Photoelectron Spectroscopy. (Kishi & Roberts) . Models. (Nicholson). . . . . . . . . . . PAGE 1545 2251 1623 2170 1230 2266 1597 1715 97 1791 2085 1192 690 242 1 997 991 1754 1021 1005 2340 1396 602 265 1671 1830 1721 2384 SUBJECT INDEX-VOLUME 71, 1975 Mossbauer Studies in the Colloid System p-FeOOH-@-Fez03 : Structures and Dehydration Mechanism. (Howe & Gallaker) .. . . . . . . . Optical Spectra of Co-Zeolites. (Hoser, Krzyzanowsoki & Trifiro) . . . . Paramagnetic Intermediate in the Oxidation of Ethylene over Magnesium Oxide. (Tarit & Naccache) . . . . . . . . . . . . . . Reflectance Spectra of Surface States in Magnesium Oxide and Calcium Oxide. (Zecchina, Lofthouse & Stone) . . . . . . . . . . . . Selective Adsorption of Methyl Esters of n-Fatty Acids at the SilicalBenzene and Silica/ Carbon Tetrachloride Interface. Part 1. Adsorption Isotherms.(Mills & Hockey) Part 3. Infrared Studies and the Structure of the Solid Surface. (Mills & Hockey) . Simultaneous Measurements of Infrared Spectra and Adsorption Isotherms for the Adsorp- tion of Phenol on Silica at the Solid/Liquid Interface. (Marshall & Rochester) . Spectroscopic Study of Surface Properties of Various NH4-exchanged X Zeolites. (Guilleux & Delafosse) . . . . . . . . . . . . Stability of Silver-Thiourea Complexes in Montmorillonite Clay. (Pleysier & Cremers) . Surface Potential of Hydrogen Adsorbed on Platinum Films. (Dus & Tompkins) . . Thermochemical Properties of Ammonium and Hydrogen Exchanged Mordenite. (Weeks Jr., Hillery & Bolton). . . . . . . . . . . . Thermodynamics of Adsorption from Solution.Adsorption by Graphon from Binary Mixtures of Benzene, Cyclohexane and n-Heptane. (Ash, Brown & Everett) . . Thermodynamics of Adsorption from Solution. The Systems (Benzene+ Ethanol)/Graphon Tin Oxide Surfaces. Part 1. Surface Hydroxyl Groups and the Chemisorption of Carbon Dioxide and Carbon Monoxide on Tin(w) Oxide. (Thornton & Harrison) . . Part 2. Infrared Study of the Adsorption of Pyridine and Ammonia on Tin(w) Oxide. Part 3. Infrared Study of the Adsorption of Some Small Organic Molecules on Y-Type Zeolites. Their Heterogeneous Nature as shown by Electron Spin Resonance Zeolites X, Y and A enriched with Trivalent Cations: Sorption of Carbon Dioxide and Part 2. Entropy and Physical State of Sorbed Phase. (Coughlan & Kilmartin) . Part 2. Heats of Adsorption.(Mills & Hockey). . . . . . and (n-Heptane+Ethanol)/Graphon. (Brown, Everett & Morgan). . . . (Harrison & Thornton) . . . . . . . . . . Tin(1v) Oxide. (Thornton & Harrison) . . . . . . . . following 7-Irradiation. Ammonia. Part 1. Isotherms and &ties. (Coughlan & Kilmartin) . . . (Abou-kais, Vedrine & Massardier) . . . . . PAGE 22 665 1402 1476 2384 2392 2398 2478 1777 256 930 205 1 123 883 461 101 3 2468 1697 1809 1818 I, 2 Biophysical Chemistry Adhesive Interaction with Lignin and Various Cellulose Monolayers. (Casilla & Eley) . 1469 Effect of Linear Energy Transfer on the Radiation-induced Inactivation of Dilute Aqueous Ribonuclease Solutions. (Bisby, Cundall &Burns) . . . . . . 1582 Free Radical Reactions in the Coenzyme B12 System.(Blackburn, Kyaw, Phillips & Nuclear Magnetic Resonance Relaxation in Aqueous Bovine Albumin Solutions. (Eley, Orientation of Poly-L-lysine Hydrobromide by Magnetic Fields. (Finer & Darke) . . 984 Self-diffusion of Zinc in Cross-linked Crystalline Carboxypeptidase A. (Dyer, Phillips & Swallow) . . . . . . . . . . . . . . 2277 Hey & Ward) . . . . . . . . . . . . . 1106 Townsend) . . . . . . . . . . . . . 803 I, 3 Catalysis (including heterogeneous and homogeneous catalysis and surface reactivity) Catalysis by Group Ib Metals. Part 1. Reaction of Buta-1,3-diene with Hydrogen and with Deuterium catalysed by Alumina-supported Gold. (Buchanan & Webb) . . 134 Catalytic Hydrogenation of Olefins over 1 : 1 and 1 : 2 Electron Donor-Acceptor Complexes of Polynuclear Aromatic Hydrocarbons with Alkali Metals.(Ichikawa & T a m ) . 2132 Decomposition of Propan-2-01 on Pure and Doped Rutile. (Gentry, Rudham & Wagstaff) 657 Electron Paramagnetic Resonance Studies of some Supported Organometallic Catalysts. Electron Spin Resonance and Catalytic Measurements of the Oxidation of Olefins on Exchange of Alkanes with Duterium over y-Alumina. A Bronsted Linear Free Energy Exchange Reactions of Propene on Oxide Catalysts Investigated by Microwave Spectroscopy. (Howe) . . . . . . . . . . . . . . 1689 Cadmium Molybdate Catalyst. (Burlamacchi, Martini, Trifiro & Caputo) . . 209 Relationship. (Robertson, Scurrell & Kemball) . . . . . . . 903 (Hughes & Kemball) . . . . . . . . .. . . 1285SUBJECT INDEX-VOLUME 71, 1975 5 PAQI Low Temperature Parahydrogen Enrichment Catalysed by Paramagnetic Ions Supported on Rutile. (Rudham & Tullett) . . . . . . . . . Nitric OxidelNitrogen Isotope Exchange on Tungsten. (Gasser and Jackson). . . Outer Sphere Catalysis of Cerium(1v)-Iron@) Electron Transfer Reactions by Pyridine Derivatives. (Ulstrup) . . . . . . . . . . . Reaction of Alkanes on Platinum-Tin and Platinum-Rhodium Alloy Films. (Karpinski and Clarke) . . . . . . . . . . Reactions of Alkanes on Iridium' and Iridium-Gold Catalysts. (Karpinski & Clarke) . Reactions of Alkanes on Rhenium and Rhenium-Gold Films. (Clarke & Taylor) . Solvent Effect on the Catalysis by Cetyltrimethylammonium Bromide Micelles of the Reaction between Hydroxide Ions and 2,4-Dinitrochlorobenzene.(Blandamer & Reid) Theoretical Studies of Polyelectrolyte " Catalysis " of Ionic Reactions. Part 1. Interionic Reactions between Oppositely Charged Species. (Mita, Kunugi, Okubo & Ise) . Water as a Base in the Acid-Base Catalysis of Hydrogen Exchange in Ketones. (Earls & Jones) . . . . . . . . . . . . . . I, 4 Colloid Science (including birefringence, electrophoresis, light scattering, sedimentation, thioxotropy, soluble and insoluble monolayers, micelles) Analytical Theories of the Steric Stabilization of Colloidal Dispersions. (Smitham, Evans & Napper) . . . . . . . . . . . . . Dynamic Surface Properties of Nonionic Surfactant Solutions. (Lucassen & Giles) . Electrokinetic and Surface Potentials at Liquid Interfaces.(Carroll and Haydon) . . Investigation of Equilibrium Wetting Films of n-Alkanes on a-Aluminia. (Blake) , . Menisci at a Free Liquid Surface; Surface Tension from the Maximum Pull on a Rod. Properties of the Non-polar OillWater Interface. Part 1. Procedures for the Accurate Measurement of the Interfacial Pressure of an Insoluble Monolayer. (Taylor & Mingins) . . . . . . . . . . . . . Proton Magnetic Resonance Investigations of the Interaction of Aerosol-OT with Imidazole, Methanol and Pyrazole in Carbon Tetrachloride. (El Seoud & Fendler) . . . Regulation of Surface Potential at Amphoteric Surfaces during Particle-Particle Interaction. (Chan, Perram, White & Healy) . . . . . . . . . Thermodynamics of Micellization of Homologous Series of n-Alkyl Methyl Sulphoxides and n-Alkyl(dimethy1)phosphine Oxides.(Clint & Walker) . . . . . (Padday, Pitt & Pashley) . . . . . . . . . . . 2361 23 3 435 893 2310 2063 21 56 936 21 86 285 21 7 361 1 92 1919 1161 452 1046 946 I, 5 Combustion and Flames (including explosions, shock waves; see also I, 8) Mass-spectrometric Tracer and Photometric Studies of Catalysed Radical Recombination Proton Afkity of Water and the Mechanism and Kinetics of Production of H30+ in Flames in Flames. (Jensen and Jones) . . . . . . . . . . 149 of HP, Oz and NZ. (Hayhurst & Telford) . . . . . . . . 1352 I, 6 Diffusion (including transport processes, thermal diffusion, viscosity, thermal conductivity; see also II, 8) Diffusion and Sorption of Simple Ions in Cellulose : Ion Exchange.(Bender, Moon, Stine, Fried, Klein &Bonjoukian) . . . . . . . . . . 491 Diffusion controlled Kinetics in the Reaction of Ferroprotoporphyrin IX with Carbon Monoxide. The Reaction Studied as a Function of Pressure, Temperature and Solvent by Means of Laser Flash-photolysis High-pressure Apparatus. (Caldin & Hasinoff). 515 Diffusion in Viscous Solvents. Part 2. Planar and Spherical Molecules in Propane1 ,Zdiol at 15,25 and 35OC. (Skipp & Tyrrell) . . . . . . . . . 1744 Difhsion of Monatomic and Diatomic Gases in 4A and 5A Zeolites. (Ruthven & Derrah) 2031 Donnan-membrane Effects in Hypefiltration of Ternary Systems. (Lonsdale, Pusch & Measurement of Diffusion Coefficients of Electrolytes by a Modified Open-ended Capillary Method.(Agar & Lord) . . . . . . . . . . . 1659 Measurements of Self-diffusion Coefficients of Water in Pure Water and in Aqueous Electrolyte Solutions. (Tanaka) . . . . . . . . . . 1127 Molecular Reorientation and Self-diffusion in the Plastic Solid Hexamethyldisilane studied using Nuclear Magnetic Resonance and Radiotracer Techniques. (Chadwick, Chezeau, Folland, Forrest & Strange) . . . . . . . . . . 1610 Walch) . . . . . . . . . . . . 5016 SUBJECT INDEX-VOLUME 71, 1975 Rate Theory of Irreversible Thermodynamics to Isotope Diffusion. Part 1. Isotope- Isotope Coupling Coefficients for Ions and Water in Concentrated Aqueous Solutions of Alkali Metal Chlorides at 298.16 K. (Anderson & Paterson) . . . . Solute Diffusion in Hydrated Polymer Networks. Part 1.Cellulose Gels. (Brown & Chitumbo) . . . . . . . . . . . . Part 2. Polyacrylarnide, Hydroxyethylcellulose and Cellulose Gels. (Brown & Chi tumbo) . . . . . . . . . . . . . Self-diffusion in Benzene under Pressure. (Collins & Woolf) . . . . . Thermal Diffusion and Convective Stability. Experimental study of the Carbon Tetra- chloride+Chlorobenzene System. (Sparasci and Tyrrell) . . . . . Test of the Curtiss and Muckenfuss Spherocylinder Model: Thermal Conductivities of Mixtures of the Noble Gases with Carbon Dioxide, Carbon Disulphide or Carbon Oxide Sulphide at 323 K. (Clifford, Colling &Gray) . . . . . . Testing Intermolecular Potential Functions using Transport Property Data. Part 1. Viscosity of Hydrogen from 273 to 1060 K. (Clifford, Dickinson, Gray & Scott) .Part 2. Thermal Conductivities of Mixtures of Helium with the Hydrogen Isotopes. (Clifford, Colling, Dickinson & Gray) . . . . . . . . . Tracer Diffusion of Tritiated Water (THO) in Ordinary Water (H20) under Pressure. Transient Diffusion through a Membrane Separating Finite and Semi-infinite Volumes. (Barrie, Spencer & Quio) . . . . . . . . . . . Viscosity Br) Coefficients in 100 % Sulphuric Acid as Solvent. (Spiro) . . . . Viscosities of Gaseous Argon, Oxygen and Carbon Monoxide Between 273 and 1300 K. (Clifford, Gray & Scott) . . . . . . . . . . . Water Transport in Aqueous Organic Solvents. Dimethyl Sulphoxide+ Water+ Hydro- chloric Acid at 25°C. s h o o & Chee-Yan) . . . . . . . . (Woolf) . . . . . . . . . . . . . .PAGE 1353 1 12 2296 42 1590 1953 1962 784 2459 988 875 446 I, 7 Electrochemistry (including electrolytes, activity coefficients, electrical conductivity, Cathodic Reduction of Acetophenone in acidic Methanol. Electrode Kinetic Study of a Charge Saturation in Splashing of Large Drops on Solid Spheres. (Maxwell & Iribarne). Determination of Thermoelectric Powers in Aqueous Hydrobromic Acid using Hydrogen Electrical Resistivity of Liquid Sodium+ Lithium Mixtures. Evidence for Incipient Electromotive Force Studies of Dilute Aqueous Solutions of Hydrochloric Acid+ Methyl- substituted Ammonium Chloride Mixtures at 25°C. (Macaskill & Pethybridge) . 1460 Osmotic Coefficients and Activity Coefficients of Methyl-substituted Ammonium Chlorides and Ammonium Methyl Sulphates in Water at 25°C.(Macaskill & Pethybridge) . 1465 Hydrogenation of Ethylene on Metal Electrodes. Part 3. Isotopic Exchange and Deutera- tion Reactions at a Platinum Electrode on Open Circuit. (Fujikawa, Kita, Miyahara Oxygen Electrode. Part 5. Enhancement of Charge Capacity of an Iridium Surface in the Anodic Region. (Buckley and Burke) . . . . . . . . . 1447 Properties of Molten Carboxylates. Part 1. Electrical Conductance and Molar Volumes of some Molten Lead and Zinc Carboxylates. (Ekwunife, Nwachukwu, Rinehart & X-Ray Photoeiectron Spectroscopic Studies of Oxide Fiims on Platinum and Gold electrode processes) Novel Vicinal Diether Synthesis. (Ud Din Bhatti & Brown) . . . . . Electrodes. (Payton, Angelos, Shuck & Zimmerman) . . . . .. 2111 Immiscibility? (Down, Hubberstey & Pulham) . . . . . . . 1387 106 1033 & Sato) . . . . . . . . . . . . . . 1573 Sime) . . . . . . . . . . . . 1432 Electrodes. (Dickinson, Povey & Sherwood) . . . . . . . 298 I, 8 Kinetics of Reaction (including photochemistry, reaction of gases, solids, liquids, two phase systems) Addition of n-C4Hs and C4Hs to Slowly Reacting Mixtures of Hydrogen and Oxygen at 480°C. Part 1. Formation of Hydrocarbon Products. (Baker, Baldwin, Fuller & Part 2. Formation of Oxygenated Products. (Baker, Baldwin & Walker) . . 756 Application of Time-dependent Rate Constant Theory to Reactions of Solvated Electrons. Reaction Distances, Rate Constants and Diffusion Coefficients in Concentrated Aqueous Solutions. (Buxton, Cattell & Dainton) .. . . . . 115 Arrhenius Parameters for Silene Insertion into Silicon-Hydrogen Bonds. (Cox & Purnell) 859 Walker) . . . . . . . . . . . . . . 736SUBJECT INDEX-VOLUME 71, 1975 7 PAGE Arrhenius Parameters for the Unimolecular Decompositions of Azomethane and n-Propyl and Isopropyl Radicals and for Methyl Radical Attack on Propane. (Camilleri, Marshall & Purnell) . 1491 Association of Di-t-butyl Nitroxide with Schardinger Dextrins. (Atherton and Strach) . 357 BEBO Calculations. Part 4. Arrhenius Parameters and Kinetic Isotope Effects for the Reactions of CH3 and CF3 Radicals with H2 and D2. (Arthur, Donchi & McDonell) 2431 Part 5. Arrhenius Parameters and Kinetic Effects for the Reactions of C2H5 and C2F5 Radicals with Hz and Dz.(Arthur, Donchi & McDonell) . . . . 2442 Effect of Substituents on the Geometry of Transition States for Slow Proton Transfer Electron Spin Resonance Studies of Photooxidation by Metal Ions in Rigid Media at Low Free Radical Addition to Olefins. Part 14. Addition of Trifluoromethyl Radicals to Formation and Dissociation of Cz from High Temperature Pyrolysis of Acetylene. (Beck Gas Phase Chain Reaction'of H;02+N02+C0. (Campbell, Handy & Kirby) . . 867 Gas-phase Thermal Unimolecular Decomposition of Oxetan. (Holbrook & Scott) . . 1849 Heavy Atom Kinetic Isotope Effects. Part 3. The Chlorine Kinetic Isotope Effect in the Gas-phase Unimolecular Decomposition of Ethyl Chloride. (Christie, Johnson, Hg(63P1) Photosensitization of 3,3-Dimethylbut-l-ene and the 2,Z-Dimethylbuta-l,3-diyl Homogeneous Gas Phase Pyrolysis of N-Isopropyiacetamide. (Maccoll '& Nagra) .Infrared and Raman Spectral Study of the Aqueous Nickel(rr)-Nitrite System. Evidence Interaction of Carbonyl Triplets with Aliphatic Amines. (Abuin, Encina, Lissi and Scaiano) . . . . . . . . . . . . . . 1221 Kinetic and Amplitude Measurements for the Process of Association of Acridine Orange studied by Temperature-jump Relaxation Spectroscopy. (Robinson, Seelig-Loffler & Schwarz) . . . . . . . . . . . . . 815 Kinetic Isotope Effects and Tunnelling in the Proton-transfer Reaction between 4-Nitro- phenylnitromethane and Tetramethylguanidine in various Aprotic Solvents. (Caldin & Mateo) . . . . . . - . . 1876 Kinetic Isotope Effects in Some Reactions of Tk(4-nitropheny1)methane with Alkoxide Bases in Various Alcoholic Media.(Caldin, Dawson, Hyde & Queen) . . . 528 Kinetic Studies for Diatomic Free Radicals using Mass Spectrometry. Part 3. Elementary Reactions involving BrO X211 Radicals. (Clyne & Watson) . . . . . 336 Kinetics and Mechanisms of Unimolecular Gas-phase Reactions of 2-Met hylpropane Radical-ions at Times as Short as 20 Picoseconds and at longer Times. (Derrick, Falick & Burlingame) . . . . . . . . . 1503 Kinetics of Reaction of Low-spin I&(rrj Complexes in Aqueous Gels. (Blandamer, Kinetics of the Reaction of Nickel@) Ion with Pyridine-2-axodimethylanilene: Solvent Kinetics of the Thermal Gas-phase Decomposition of 1,2-Epoxybutane. (Flowers & Penny) Kineticspf the Thermal Gas-phase Decomposition of 2,3-Epoxy-2-methylbutane.(Flowers Kinetics of the Thermolysis of Octamethylcyclotetrasiloxane in the Gas Phase. (Davidson Lauric Acid/Pentaerythrityl monolaurate : A Model Melt Esterification. Part 1. Kinetics. . . . . . . . . . . . Reactions. (McLennan) . . . . . . . . . . 1516 Temperatures. Part 5. Photo-oxidation by the Iron(@ Ion. (Cox & Kemp) . 2490 Fluoroethylenes. (Cape, Greig, Tedder & Walton) . . . . . . 592 & Mackie) . . . . . . 1363 Loudon, MacColl & Mruzek) . . . . . . . . . . 1937 Biradical. (Montague) . . . . . . . . . 398 for Photochemical Alteration of the Chemical Equilibrium. (Brooker) . . . 647 . 2450 Burgess and Membrey) . . . . . . . 145 Effects and Correlations. (Bennett0 & Imani) . . . . . . 1143 851 & Ozturk). .. . . . . . . 1509 &Thompson) . . * . . . . . . . . 2260 (Gordon & Leonis) . * . . . 161 Part 2. Statistical Distribution of Products. (Gordon and Leonis) . . . 178 with Ground State N 4S Atoms. (Clyne & McDermid). . . . . . 2189 . . . . . Mass Spectrometric Determinations of the Rates of Elementary Reactions of NO and NO2 Mechanism of Thermolysis of Hexamethyldisilane and the Silicon-Silicon Bond Dissociation Energy. (Davidson & Howard) . . . . . 69 Photochemistry of Anhydrides. Part 3. Primary 'Processes 'in the Photolyses of (CF3C0)z0 and (C2F5C0)20 and the use of (CF3CO)20 and (C2F5C0)20 as Gas Photolysis of Diazo-n-butane. Study of the Unimolecular Decomposition of Activated Phase Actinometers, (Chamberlain & Whittle) .. . . . . . 1978 Intermediates and Energy Partitioning. (Figuera, Perez & Wolf) . . . . 19058 SUBJECT INDEX-VOLUME 71, 1975 Photolysis of Perbromate in Aqueous Solution. (Klaning, Olsen & Appelman) . . Photoreactivity of Michler’s Keton in Solution. (Suppan) . . . . . . Pulsed Stirred-flow Technique for Gas Kinetics. (Baldwin, Davidson & Howard) . Radicals Generated by the Heterogeneous Decomposition of Peracetic Acid Studied by Electron Spin Resonance Spectroscopy. (Nalbandyan, Oganessyan, Vardanyan & Griffiths) . , . . . . . . . . . Reaction between Hydrogen and Nitrous Oxide. (Baldwin, Gethin, Plaistowe & Walker) Reaction of Hydroxyl Radicals with NO, NO2 and SOz. (Harris & Wayne) . . . Reaction of Oxygen Atoms with Methyl and Ethyl Nitrites.(Davidson & Thrush). . Reaction of Singlet Methylene with Methylenecyclopropane. Part 1, Evidence for Multistep Collisional Deactivation of Chemically Activated Spiropentane. (Frey, Reaction Kinetics of Ground State Fiuorine, FzP, Atoms. Part 2. Reactions Forming Inorganic Fluorides, Studied Mass Spectrometrically. (Appelman & Clyne) . . Reactions of Chlorine Oxide Radicals. Part 5. The Reaction 2CIO (X211) Products. (Clyne, McKenney & Watson) . . . . . . . . . Reactions of the Hydroperoxyl Radical (HO,) with Nitrogen Dioxide and Tetranitro- methane in Aqueous Solution. (Sutton) . . . . Recombination of Ethyl Radicals in the Range 693-803 K. iHugdes & Marshall) . . Sequence Studies in Liquid Phase Hydrocarbon Oxidation. Part 3.Temperature De- pendence of Alchohol Ketone Transition in the Oxidation of Ethylbenzene. (Danoczy, Nemes, Vidoczy & Gal) . . , . . . . . . . Single-pulse Shock Tube Studies of Hydrocarbon Pyrolysis. Part 4. Isomerization of Allene to Methylacetylene. (Bradley & West) . . . . . . . Solvent-Iodine Atom Charge Transfer Absorption in Liquid Alkanes and Cycloalkanes. (Logan, Bonneau, Joussot-Dubien & Fornier De Violet). . . . . . Sorption and Desorption Kinetics of the Cellulose and Water System. Part 3. The Rates of Sorption on the Lowest Sorption Limb of a Hysteresis Loop. (Newns) . . Studies of Reactions of Atoms in a Discharge Flow Stirred Reactor. Part 1. The Thermal Decomposition of 3-E t h yl - 3-me t hyloxet an and 3,3-Diet h yloxe tan. (Clement s, Frey & Frey) .. . . . . . . . . . . . Thermal Decomposition of Isopropanol. (Trenwith) . . . . . . . Thermal Determination of the Kinetics of the Iron(@-Tin(@ Redox Reaction in Chloride Solution. (Scott, Glasser & Nicol) . . . . . . . . . Jackson, Smith & Walsh) . . . . . . . . . O+H2+N0 System. (Campbell & Handy). . . . . . . . I, 9 Polymers and Polymerization (including physical properties of polymers and their solutions) Chemical Reactivity and Conformational Properties of Growing Chains. Part 2. Effect of the Reaction Medium on the Termination Kinetics in StyreneSMethyl Metha- crylate Solution Copolymerization. (Bonta, Gallo & Russo) . . . . Emulsion Polymerization Kinetics. General Solutions for Smith-Ewart Cases I & 11. (Hawkett, Napper & Gilbert) .. . . . . . . . . Kinetics of the Polymerization of Methyl Methacrylate initiated by Butylmagnesium Bromides and Dibutylmagnesium in Tetrahydrofuran+ Toluene. (Allen & Bateup) . Photo-initiation of Free-radical Polymerization by Dimethyl-(2,2’-bipyridyl)platinum(11)+ Tetrafluoroethylene. (Bamford, Mullik & Puddephatt) . . . . . . Photoinitiation of Free-radical Polymerization by Rhenium Carbonyl in the presence of Tetrafluoroethylene. (Bamford & Mullik) . . . . . . . . Photoinitiation of Free-radical Polymerization by Vanadium Chelates. Part 3. Methoxo- oxobis(8-quinolyloxo)vanadium(v). (Aliwi & Bamford). . . . . . Photoinitiation of Polymerization by Chloro-oxobis(2,4-pentanedionato)vanad~um(v) in the Presenceof ElectronDonors.(Aliwi &Bamford) . . . . . . Polymerisation of N-Vinylcarbazole. Part 1. End-group Studies on Polymers Prepared Using Azobisisobutyronitrile and Benzoyl Peroxide. (Bevington & Dyball) . . Primary and Chain Carrier Ions Participating in the Gas Phase Polymerization of Ethylene Oxide. (Giardini-Guiodon, Mele, Platania & Zocchi) . . . . . . Radical Ions Derived from Photosynthetic Polyenes. (Dawe & Land) . . . . Rate Theory of Irreversible Linear Random Polymerisation. Part 1. Basic Theory. (Stanford & Stepto) . * . . . . . . . . . . Part. 2 Application to Intramolecular Reaction in A-A+B-B Type Polymerisations. (Stanford, Stepto & Waywell) . . . . . . . . . . PAGE 473 539 972 1203 1265 61 0 241 3 1991 2072 322 2142 41 3 84 1 967 2148 278 2097 2485 2405 1413 1727 2288 2203 221 3 625 1733 52 2226 35 1 2162 1292 1308SUBJECT INDEX-VOLUME 71, 1975 9 PAGE Reaction of Polymer Radicals with Chromium(n) Acetate.(Lee & Minoura) . . . 1649 Thermodynamic Properties of Sodium Metaphosphate Polymers. (Jeffes & Warner) . 679 Chemical Consequences of Low Energy I+ or I: Implantation in Solid Hydrocarbons. Gamma Irradiation Induced Isotope Exchange in Nitrogen Sensitized by Helium, Neon and Low Temperature Pulse Radiolysis of Concentrated Aqueous Solutions. Dry Electron Capture Efficiencies and Tunnelling Processes of Trapped Electrons in 9.5 mol dm-3 LiCl and 10 mol dm-3 OH- Solutions. (Buxton & Kemsley). . . . . 568 OH Radical Induced Oxidation of Ethanol in Oxygenated Aqueous Solutions. Part 1.Formation of Acetic Acid. (Schultze & Schulte-Frohlinde) . . . . . 1099 Protonation Reaction of the Benzonitrile Radical Anion and Absorption of the Product. (Holcman & Sehested) . . . . . . . . . . . 121 1 Pulse Radiolysis of Methanol and Ethanol. Acid-Base Behaviour of Hydroxymethyl and Pulse Radiolysis Studies of Aromatic Liquids. Part 1. Electrons and Negative Ions. Pulse Radiolysis Study of Monovalent Cadmium, Cobalt, Nickel and Zinc in Aqueous Solution. Part 1. Formation and Decay of the Monovalent Ions. (Buxton & Pulse Radiolytic Investigation of the Reduction' of Cadmium(@ ions. (Kelm, Lilie & Radiation Mechanisms. Part 4. Electron Spin Resonance Studies of y-Irradiated Tri- chloronitrosomethane and Trichloronitromethane. (Mishra, Symons & Tattershall) .1772 Reaction of Charged Species in Irradiated Cyclohexane with Added Solutes as studied by Steady State Radiolysis. Relative Rate Constants for Reaction of Positive and Negative Species with Different Solutes. mavids, Warman & Hummel) . . . Reactions of Hydrogen Atoms in 6 mol dm-3 Sulphuric Acid. Part 1. Temperature Dependence of Activation- and of Diffusion-controlled Reactions. mainton, Solid Ammonia Radiolysis. Temperature Effect in the Radiolysis of Solid Ammonia. Unimolecular Decomposition of [80Br]Bromocyclopropane formed by Reaction of Recoil I, 10 Radiolysis (including nuclear transformation in solids, neutron capture, etc.) (Cailleret, Paulus &Abbe) . . . . . . . . . . 637 Argon, (Wood & Mascall) . . . . . . . 1678 Hydroxyethyl Radicals.(Johnson &Salmon) . . . . . . . 583 (Robinson &Rodgers) . . . . . . . . . . 378 Sellers) . . . . . . . . . , 5 5 8 Henglein) . . . . . . . . . . . . . 1132 1252 Phillipson & Pilling) . . . . . . . . . . . . 2377 (Blum) . . . . . . . . . . . . . . 2299 Bromine Atoms with Cyclopropane or with Bromocyclopropane. (Saeki &Tachikawa) 2121 I, 11 Solid-state Chemistry Defect Formation in Crystals Grown from Aqueous Solutions. Growth Rate of Barium Sulphate Crystals and its Influence on their Defect Content. (Melikhov & Vukovic) Effect of Hydrostatic Pressure on Self-diffusion and Plastic Deformation in Plastic Crystals. (McKay &Sherwood) . . . . . . . . . . . Effectsof MechanicalGrinding on the Texture and Structure of Calcium Carbonate.(Criado & Trillo) . . . . . . . . . . . . . . High Field Electrical Conduction in some Organic Charge-transfer Complexes. (Pethig &Soni) . . . . . . . . . . . . Infrared Study of the'Interaction between Caesium Chloride and Kaolinite. (Yariv) . Isotope-mass-effect and Self-diffusion in Crystalline Naphthalene. (Elampton & Sherwood) Role of Structural Imperfections in the Decomposition of Dibenzoyl Peroxide. (Morsi, Thomas &Williams) . . . . . . . . . Semiconductivity of Organic Substances. Part 17. Effects of Ultraviolet and Visible Light on the Conductivity of the Sodium Salt of Deoxyribonucleic acid. (Eley, Metcalfe & White) . . . . . . . . . Structure and Catalytic Activity of Iron Oxide and Magnesium Oxide Solid Solutions. Part 1.Structural and Magnetic Investigations. (Valigi, Pepe & Schiavello) . . Part 2. Catalytic Activity for N20 Decomposition. (Schiavello, Valigi & Pepe) . 201 7 2331 961 1534 674 1392 1857 955 1631 1642 I, 12 Thermodynamic and Equilibrium Properties (including multiphase systems) Absorption of Hydrogen by Palladium-Copper Alloys. Part 1, Experimental Measure- ments. (Burch & Buss) . . . . . . . . . . . 913 Part 2. Theoretical Analysis. (Burch & Buss) . . . . . . . 92210 SUBJECT INDEX-VOLUME 71, 1975 Acid Ionization Constants of 4-Substituted Phenols in Methanol+ Water Mixtures, (Parson &Rochester). . . . . . . . . . * . Apparent Molal Volumes of KCN, K2Zn(CN)4 and K2Cd(CN)4 in Aqueous Solutions at 15-45°C. Effect of Ionization and Hydrolysis.(Mathieson & Curthoys) . . Dissociation Constant of Benzoic Acid in H20+D20 Mixtures. (Lowe & Smith) . . Enthalpies and Entropies of Ionization of 4-Substituted Phenois in Methanol + Wate; Mixtures. (Parsons & Rochester) . . . Enthalpies of Formation of n-Alkan-1-01s. illlosselman & Dekker) . . Excess Enthalpies of Some Binary 1,4-Dioxan Mixtures at 293.15 K. (Singh, Phutela & Arora) . . . . . . . . . . . . . Excess Volumes of Mixtures containing o-Dichlorobenzene. (Dhillon) . . . . Excess Volumes of Mixtures of Tetrabutylin with Branched and Linear Alkanes. (Delmas, De Saint-Romain & Purves) . . . . . . Free Energies and Entropies of Transfer of Hydrobromic and Hydrdiodic Acids from Water to t-Butyl Alcohol+ Water Mixtures from Electromotive Force Measurements at Different Temperatures (5-35°C).(Bose, Das & Dundu) . . . . . Free Energies of Transfer of Alkali Metal Fluorides from Water to Hydrogen Peroxide+ Water and Methanol+ Water Mixtures using Ion-selective Electrodes. (Covington & Thain) . . . . . . . . . . . . . Heat Content and Thermal Electromotive Force of the Molten System, AgCI+ CuCl. (Sime, Sime &Johnson) . . . . . . . . . . . Heats of Dilution of Aqueous Solutions of Polyethylenimine Hydrochloride and Sodium Polyphosphate and their Low Molecular Weight Analogues. (Mita, Okubo & Ise) . Heats of Sublimation of some Cage Hydrocarbons by a Temperature Scanning Technique. (Clark, Knox, Mackle, McKervey & Rooney) Influence of Pressure on the Ionization of Substituted Anilinium Ions.(Hamann & Linton) Intramolecular Hydrogen Bonding and the Dissociation Constant of Salicylic Acid in H20+D20 Mixtures. (Lowe & Smith) . . . . . . Investigation into the Effect of Temperature and added t-Butyl Alcohol on the Dynamic Properties of the Belousov Reaction. (Blandamer & Morris). . . . . Ion Pairing in Protic and Aprotic Solvents of Intermediate Dielectric Constant at 25°C. (DAprano, Goffredi & Triolo) . . . . . . . . . . Ionic Solvations in Water+Co-solvent Mixtures. Part 3. Free Energies of Transfer of Single Ions from Water into Water+Ethylene Glycol Mixtures. (Wells) . . . Isothermal Joule-Thomson Coefficient of Nitrogen. (Pocock & Wormald) . . Microcalorimetric Studies. The Enthalpy of Formation of Hexadecacarbonylhexarhodium, Rh6(CO)16.(Brown, Connor & Skinner) . . . . . . . . Molecular Complexes of Heteroaromatic Five Membered Ring Compounds with Tetra- cyanoethylene. Charge Transfer Spectra, Equilibrium Constants and Ionization Molecular Complexes of Naphthalenes. Part 2. Electromotive Force and Thermochemical Multinuclear Nuclear Magnetic Resonance Studies of Aqueous Solutions of Tetrafluoro- Nuclear Magnetic Resonance Studies of Preferential Solvation. Part 5. Magnesium Perchlorate in Water+Acetone Mixtures at 185 K. (Covington & Covington) . . Orientational Order in Non-alkane Chain Molecules. Heats of Mixing of Tetralauryltin, Dioctyl Ether, trans and cis-Dec-5-ene in Linear and Branched Alkanes. (Delmans & Partial Molal Volumes and Volumes of Ionization of Hydroxycarboxylic Acids in Aqueous Properties of Aqueous Thorium Nitrate Solutions.Part 3. Partial Molal Heat Capacities Effect of Deuteration on Hydrogen Bonds. (Rao) . . . . . . . . . . Potentials of the Donors, (Aloisi, Santini & Savelli) . . . . . . Studies of Picrate Formation. (Abdel-Rehiem, Farrell & Westwood) . . . borate Salts. (Akitt) . . . . . . . . . . . Thanh) . . . . . . . . . . . . . Solution at 25, 30 and 35°C. (Hoiland & Vikingstad) . . . . . . and Heats of Dilution at 30°C. (Apelbiat & Sahar) . . . . . . Reassessment of HydrophobicBonding. (Howarth) . . . . . . . Salting-out of Alkanols by Inorganic Electrolytes. (Aveyard & Heselden) . . . Solubility of Argon in Water+Alcohol Systems. (Cargill & Morrison) .. . . Solvolysis of some Basic Solutes. (Nour, Hussein & Wasif) . . . . . (Dudeney & Irving) . . . . . . . . . . . . Second Virial Coefficient of Benzene and its Temperature Derivatives. (Wormald) . . Solutions in Selenic Acid. Part 8. Cryoscopic Studies in Selanic Acid: Ionisation and Spectrophotometric Investigations of Aqueous Solutions at Elevated Temperatures. PAGE 1058 1114 389 980 1069 41 7 1528 189 1181 1838 78 2366 1932 2107 485 1379 2319 1188 1868 705 699 2045 1762 1557 83 1 1172 2007 1667 2303 312 726 61 8 1041 1215SUBJECT INDEX-VOLUME 71, 1975 Standard Potentials of Ag-AgBr and Ag-AgI Electrodes in Urea+ Water Mixtures. Free Energies and Entropies of Transfer of the Hydrogen Halides. (Kundu & Mazumdar) Studies in Ion Solvation in Non-aqueous Solvents and their Aqueous Mixtures.Part 17. Free Energies of Transfer of Alkali-metal Chlorides from Water to 10-40 % (w/w) Dioxan+ Water Mixtures, and of Potassium Bromide and Iodide to the 20 % Mixture, Thermal Transpiration Correction of Hydrogen Equilibrium Pressure Measurements in Metal/Hydrogen Solution. (Wallbank and McQuillan) . . . . . . Thermodynamic Properties of 1 : 1 Adducts between Water and Various Bases in Carbon Tetrachloride. (McTigue & Renowden) . . . . . . . . Thermodynamic Properties of Fluorine compound. Part 15. Vapour Pressures of the Three Tetrafluorobenzenes and 1,3,5-Trichloro-2,4,6-trifluorobenzene. (Ambrose, Ellender, Sprake & Townsend) . . . . . . . . . Thermodynamic Properties of the Calcium+ Calcium Chloride System measured by an Electrochemical Technique.(Dosaj, Aksaranan & Morris) . . . . . Thermodynamic Study of Aqueous Dilute Solution of Organic Compounds. Part 3. Morpholines and Piperazines. (Cabani, Conti, Giannessi & Lepori) . - Part 4. Cyclic and Straight Chain Secondary Alcohols. (Cabani, Conti, Mollica & Thermodynamics of Dilute Solutions of Ethanol (1) in p-Xylene (2) from Freezing-point and Enthalpy of Dilution Measurements. (Stokes & Adamson) . . Thermodynamics of Ionization of Amino-Acids. Part 6. The Second Ionization Constants of Some Glycine Peptides. (King) . . . . . . . . Thermodynamics of Liquid Mixtures of Xenon and Hydrogen Chloride. (Calado, Kozdon, Morris, Da Ponte, Staveley & Woolf) . . . . . . . . Thermodynamics of Mixed Electrolyte Solutions Comparison of HCI + NH4CI + H20 and HC1+ KCI + HzO.(Downes) . . . . . . . . Thermomechanical Studies on Natural Rubber in' Torsion and Simple Extension. (Allen, Price & Yoshimura) . . . . . . . . Total Vapour Pressures, Thermodynamic Excess Functions and Complex Formation in Binary Liquid Mixtures of Some Organic Solvents and Sulphur Dioxide. (Lorimer, Smith &Smith). . . . . . . . . . . . . Vapour/Liquid Equilibrium in Ethylene+ Carbon Dioxide and Ethane+ Carbon Dioxide. (Mollerup) . . . . . . . . . . . Volumes of Ionization of Dicarboxylic Acids in Aqueous Solution from Density Measure- ments at 25°C. (Hoiland) . . . . , . . . . . at 25°C. (Feakins, Hickey, Lorimer & Voice) . . . . . . . Lefori) . . . . . . . . . . . . 11 PACE 1422 780 685 1784 35 1083 1154 1943 1707 88 1372 425 548 2232 2351 797AUTHOR INDEX-VOLUME 71.1975 Abbe,J.Ch. . . AbdeLRehiem. Ahmed G . Abd.El.Nabey. Beshier A . Abou.Kais. Antoine . Abuin. ElsaB . . . Adamson. Marion . Agar. J . N . . . . Akasaranan. Chokechai . Akitt. J . W . . . Alisi. S . M . . . . Aliwi. S . M . . . Allen. Geoffrey . . Allen. Peter E . M . . Aloisi. G . Gaetano . Ambr0se.D . . . Anderson. John . . Apelblat. Alexander . Appleman. Evan H . . Arora. P . S . . . Arthur. N . L . . . Ash. Stuart G . . . Atherton. Neil M . . Aveyard. Robert . . Baker. Richard R . . Baldwin. Alan C . . . Baldwin. Robert R . . Baldwin.RoyR . . . Bamford. Clement H . . Barraclough. Peter B . . Barrer. Richard M . . Barrie. James A . . . Barthomeuf. Denise . Bateup.Brett 0 . . . Beck,W.H. . . Bender. Max . . Bennetto. H . Peter . Bevington. John C . . Bisby. Roger H . . . Blackburn. R . . . Blake. Terence. D . . Blandamer. Michael J . . Blum. Aleksander . Bolton. Anthony P . . Bonjouk1ian.R . . . Bonneau. R . . . Bonta. Giorgio . . Bose. Kumardev . . Bown. Richard . . Bradley. John N . . . Breakspere. Robert J . . Brooker,M.H. . . Brown. Christopher E . . Brown. D . Lalage S . . Brown. Oliver R . . . PAGB . 637 . 1762 . 1230 . 1697 . 1221 . 1707 . 1659 . 1083 . 1557 . 52 . 1733 . 548 . 2203 . 2045 . 35 . 1335 . 1667 473. 2072 . 1528 2431 . 2442 . 123 . 357 . 312 736. 756 . 972 . 1265 736. 756 Brown. Wyn . Buchanan. Douglas Buckley. Denis N . Burch. R . . . Burgess. John . Burke. Laurence D . Burlamacchi. Leo Burlingame. A .L . Burns. William G . Buss. R . G . Bwton. George B . Buxton. George V . PAGE . . . . 1. 12 A . . . . . 134 . . . . 1447 . . . 913. 922 . . . . 145 . . . . . 1447 . . . . 209 . . . . 1503 . . . . 1582 . . . 913. 922 . . . . 568 . . . 115. 558 Cabani. Sergio 1 154. 1943 Caillbret. J . . . . . . . 637 Calado. Jorge C . G . . . . . 1372 Caldin. Edward F . . 515. 528. 1876 Camilleri. Patrick . . . . 1491 Campbell. Ian M . . . . . 867. 2097 Caputo. G . . . . . . . 209 Cargill. Robert W . . . . . 618 Carroll. Brendan J . . . . . 361 Casilla. Romulo C . . . . . 1469 Cattell. Frank C . R . . . . . 115 Chadwick. A . V . . . . . . 1610 Chamberlain. Geoffrey A . . . . 1978 Chan.D . . . . . . . 1046 . . . Cape. JohnN . . . . . . 592 52.625. 1733. 2213 CheeiYan Chan . . .. . 446 . . . 2266 Chezeau. J.M. . . . . . 1610 . . . 690 Chitumbo. Kaluba . . . . 1. 12 . . . 2459 Christie. J . R . . . . . . 1937 . . . 991 Clark. Timothy . . . . . 2107 . . . 1362 Clements. Allan D . . . . . 2485 . . . 491 Clifford. Anthony A . 875. 1590. 1953. 1962 . . . 1143 Clint. JohnH . . . . . 946. 1327 . . . 2226 Clyne. Michael A . A . . 322.336.2072. 2189 . . . 1582 Cody.1.A . . . . . . . 1021 . . . 2277 Colling. Lynne . . . 1590. 1962 . . . 192 Collings. A . F . . . . . . 2296 . 145.2156. 2319 Connor. Joseph A . . . . . 699 . . . 2299 Conti. G . . . . . . 1154. 1943 . . . 2051 Coughlan. Brendan . . . 1809. 1818 . . . 491 Covington. Anthony D . . . . . 831 . . . 2148 Covington. Arthur K . . . . 78. 831 . . . 1727 Cox. Alan . . . . . . 2490 . . . 1838 Cox.Bruce . . . . . . 859 . . . 123 Cremers. Adrian . . . . 256. 265 . . . 967 Criado. J.M. . . . . . 961 . . . 2251 Cundall. Robert B . . . . . 1582 . . . 647 Curthoys. GeoErey . . . . 1114 . . . 883 . . . 699 Dainton. Frederick S . . . . 115. 2377 . . . 106 Danoczy. Eva . . . . . 841 . . . 2203 Clarke. JohnK . A . . 893.2063. 2310 12Da Ponte. Manuel Nunes Darke. Arthur . Das. AsimK . . Davids. Erik L . . Davidson. Iain M . T . Davidson. J . A . . Davies. Manse1 . Dawe,E.A. . Dawson. Eleanor . Deane. Michael . Declerck. Luc . J . Dekker. Harm . Delafosse. D . Delmas. Genevieve Derrah. Russell I .. Derrick. Peter J . . Dhillon. M . S . . Dickiison. Eric . Dickinson. Thomas Donchi. K . F . . Dosaj. Vishu . Down,M.G. . Downes. Colin J . . Dudney. William Duffy.Stephen C .. DMS. Rikard . Dyball. Christopher J . . Dyer. Alan . . . Earls. Derek W . . . Ekwunife. M . E . . . Hey. Daniel D . . Ellender. J . H . . . El Seoud. Omar A . . Encina. Maria V . Evans. Robert . . Everett. Douglas H . . Falick.A . M . . Farmer. Victor C. Farrell. Patrick G . Feakins. David . Fendler. Janos H .. Figuera. Juan M . Finer. Elliot G . . Flochart. B . D . . Flowers. Michael C . F0Uand.R . . . Fornier De Violet. P . Forrest. J . W . . Fox. Malcolm F . . Franco.Parra. Carlos Fraser. A . R . . Frey. Henry M . . Frey. Jeremy G . . Fried. A . . . Fujikawa. Keikichi Fuller. Alan R . . Gal. Dezso . . . Gal. Ivan J . . . . AUTHOR INDEX-VOLUME 71. 1975 PAGE . . 1372 . . 984 . . 1838 . . 1252 69.972. 2260 . . 2413 . . 1791 . 2162 . 528 . 1005 .1545 . 417 . 1777 1172. 1181 . 2031 . 1503 . 189 1953. 1962 . 298 2431. 2442 . 1083 . 1387 . 425 . 1215 . 602 . 930 . 2226 . 803 . 2186 PAW Gallsgher. Kevin J . . . . . 22 Gallo. Bianca M . . . . . 1727 Gasser. Robert P . H . . . . . 233 Gentry. Stephen J . . . . . 657 Gethin. Allan . . . . . 1265 Giannessi. D . . . . . . 1154 GiardWGuidoni . Anna . . . 351 13 Gilbert. Robert G . Giles. Dennis . Glasser. David . Gordon. Manfred Gray. Derek G . . Gray. Peter . . Greig. Alan C . . Griffiths. David L . Griffiths. John F . . Guilleux. M . F . . Hall. Peter G . . Hamann. Sefton D . Hampton. Eric M . Handy. Brian J . . Harris. G . W . . Harrison. Philip G . Hasinoff. Brian B . Hassan. L . A . R . . Hawkett. Brian S . Haydon. Denis A .. Hayhurst. Allan N . Helv . T . W .. . . 1432 Henglein. A . . 955.1106. 1469. 2340 Heselden. Roy . . . 35 Hey. MichaelJ . . . 452 Hickey.BrianE . . . 1221 Hillery. Herbert F . . . 285 Hockey. John A . . . 123. 883 Hoiland. Harald . . . 1503 . . 1623 . . 1762 . . 780 . . 452 . . 1905 . . 984 . . 1192 . 851. 1509 . . 1610 . . 2148 . . 1610 . . 1407 . . 991 . 1623 . 1991. 2485 . . 2485 . . 491 . . 1573 . . 736 . . 841 . . 1671 H H H H H H H H H H H H H H olbrook. Kenneth A . olcman. Jerzy . oser. Hanna . owe. William A . oward. Anthony V . owarth. Oliver W . .owe. Arthur T . . Lowe. Russell F . . hbberstey. P . . iughes. Brian T . . [ughes. David G .. [ummel. Andries [ussein. M . A . . [yde. Richard M . . . . . 2288 . . . . 217 . . . . 1413 . . . 161. 178 . . . . 2170 . 875. 1590. 1953. 1962 . . .592 . . . 1005 . . . 1203 . . . 1777 . . . 2266 . . . 485 . . . 1392 . . 867. 2097 . . . 610 . 461. 1013. 2468 . . . 515 . . . 2251 . . . 2288 . . . 361 . . . 1352 . . . 1046 . . . 1132 . . . 312 . . . 1106 . . . 780 . . . 2051 . 2384.2392. 2398 Ichikawa. Masaru . Imani. Z . Sabet . . Iribarne. J . V . . . Irving. Roger J . . . Ise. Norio . . . Jacobs. Peter A . . . Jackson. George E . . Jackson. Graham. V . . 797. 2007 . 1849 . 1211 . 665 . 1597 69. 972 . 2303 . 22 . 1689 . 1387 . 1285 . 413 . 1252 . 1041 . 528 . 2132 . 1143 . 1033 . 1215 936. 1932 . 1545 . 1991 . 233Jaycock. Michael J . Jeffes. James H . E . Jensen. David E . Johnson. David W . Johnson. K . E . . Johnson. W . D . . Jones. George A . . Jones. Gwyn . Jones. John R . . Joussot.Dubien. J . Karpinski. Zbigniew Kelm.M .. . Kemball. Charles . Kemp. Terence J . Kemsley. Kenneth G . Khoo.KeanH . . Kilmartin. Sean . King. Edward J . . Kirby. R . M . . Kishi. Kosaku . Kita. Hideaki . Kiwanuka. Gerald M . KEning. Ulrik K . Klinowski. Jacek . Knoezinger. Helmut Knox. Trevor . Kozdon. Andrzej F . Krietenbrink. Hans Krzyzanowski. Stanislaw Kundu.KironK . . . Kunugi. Shigeru . . Kuratomi. T . . . Kyaw.M . . . . Land. E . J . . . . Lee. Munam . . Leonis. Constantine. G .. hpori. L . . . . Lewis. Ivor A . . . Lilie. J . . . . Linton. Max . . Lissi. Eduardo A . Lob0,V.M.M. . . Lofthouse. Michael G . . Logan. S . R . . . Lonsdale. Harold K . . Lorimer. John Phillip . Lorimer. John W . . Loudon. A . G . . . Lowe. Barrie M . . . Lucassen. Jacob . . Macaskill. J . B .. . Maccoll. Allan . . Mackie. J . C . . . Mackle. Henry . . Maes. Andre . . Marshall. Kenneth . Marshall. Roger M . . Martini. Giacomo . . AUTHOR INDEX-VOLUME 71. 1975 PA08 . . 1597 . . 670 . . 149 . . 583 . . 2366 . . 1937 . . 149 . 1791. 2085 . . 2186 . . 2148 . 893. 2310 . . 1132 . 903. 1285 . . 2490 . . 568 446 1 1809. 1818 . . 88 . . 867 . 1715. 1721 . . 1573 . . 2340 . . 473 . . 690 . . 2421 . . 2107 . 1372 . . 2421 665 1422. 1838 . 936 . . 97 . 2277 . . 2162 . 1649 . 161. 178 . . 1943 . . 1005 . . 1132 . . 485 . . 1221 . . 1659 . . 1476 . . 2148 . 501 . 780 . . 2232 * . 1937 . 389. 1370 . . 217 . 1460. 1465 . 1937. 2450 . . 1363 . . 2107 . . 265 . 1754. 2478 . 413. 1491 . . 209 M M M M M M M M M M M M M M n4 M M M M M M M M M M M M M M M M M M M M M M M M M ascal1,R.A. .. assardier. Jean . . ateo. Salvador . . athieson. John G . . axwell. J . B . . . azumdar. Kalyan . cDermid. I . Stuart . cDonel1. J . A . . . cKay. Peter . . cKenney. Donald J . . cKervey. M . Anthony . clennan. Duncan J . . cQuillan. Alan D . . eTigue. Peter . . ele. Aldo . . . elikhov. I . V . . . embrey. Jill R . . . etcalfe. Edwin . . ills. A . K . . . . ingins. J . . . . inoura. Yuji . . ishra. Shuddhodan P .. ita. Kazuei . . iyahara. Koshiro . loiseyev. N . . . [ollan. P . A . F . . . ollerup. Jdrgen . . ollica. V . . . . ontague. Derek C . . oon. J . K . . . .organ. Christopher J .. [orris. David R . . . 'orris. Peter J . . . [orris. Stephen H . . [orrison. Thomas J . . [orrow. B . A . . . rorsi. Salah E . . . [osselman. Cornelis . [ruzek.M . N . . . [ullik. s . u . . . Naccache. Claude . . Nagra. Surjit S . . . Nalbandyan. Azam B . . Napper. Donald H . . Nemes. 1stua.n . . Newns. Alan C . . . Nicholson. David . . Nicol. Michael J . Noller. H . . . . Nour,M.M. . . Nwachukwu. M . U . . Oganessyan. Elia A . . Okubo. Tsuneo . . Oztiirk. Turgut . . Olsen. Kjeld J . . . Padday. J . F . . . Parsons. Gerald H . . Pashley. R . M . . . ?AOI . . 1678 . . 1697 . 1876 . . 1114 . . 1033 . . 1422 . . 2189 . 2431. 2442 . 2331 . . 322 . . 2107 . 1516 . . 685 . 1784 . . 351 . . 2017 . . 145 . . 955 2384.2392. 2398 . 1161 * 1649 . 1772 936. 1932 . 1573 . 1830 . 1192 . 2351 . 1943 . 398 . 491 . 883 . 1083 . 1372 . 2319 . 618 . 1021 . 1857 . 417 . 1937 625. 2213 . 1402 . 2450 . 1203 285. 2288 . 841 . 278 . 238 . 1413 .997 . 1041 . 1432 . 1203 936. 1932 . 473 . 1509 . 1919 1059. 1069 . 1919AUTHOR INDEX-VOLUME 71. 1975 Paterson. E . . Paterson. Russell . Paulus. J . M . . Payton. Arthur D . Penny. David E . . Pepe.F . . . Perez. Juan M . . Perram. J . W . . Pethig. Ronald . Pethybridge. Alan D . Phillips. G . 0 . . Phillipson. Nigel A . Phutela. R . C . . Pichat. Pierre . Pilling. M . J . . Pink. R . C . Pitt. A . R . . . Plaistowe. John . Platania. Rosario Pleysier. Josef . Pocock. Geoffrey . Povey. Andrew F .. Price. Colin. . Puddephatt. R . J . Pulham.R . J . . Purnell. Howard . Purves. Patricia . Pusch. Wolfgang . Queen. Alan . Quig. Alexander . Radovanov. Petar Rao. C . N . Ramachandra Reid. Donald J . . . Renowden. Peter V . . Rinehart. F . P . . . R0berts.D . K .. . Roberts. M . Wyn . Robertson. Philip J . . Robinson. Alain J . . Robinson . B . H . . . Rochester. Colin H . 1058. 1069.1754.2340. 2478 ?AOB . 1623 . 1335 . . 637 . . 2111 . 851 . 1631. 1642 . . 1905 . . 1046 . . 1534 . 1460. 1465 . 803. 2277 . 2377 . 1528 . 991 . . 2377 . 1192 . 1919 . 1265 . 351 . 256 . . 705 . 298 . . 548 . . 2213 . . 1387 . 859. 1491 . . 1181 . . 501 . . 528 . . 2459 . . 1671 . . 980 . . 2156 . . 1784 . . 1432 . 2251 . 1715. 1721 . . 903 . . 378 . 815 Rodger. Michael A . J . . . Rooney. John J . . . . Rudham. Robert . . . Russell. J . D . . . . Russo. Saverio . . . Ruthven. Douglas M . . . Saeki. Masakatsu . . . Sahar. Ayah . . . Saint.Romain. Pierre . . Salmon. G.Arthur . . Santini. Sergio . . Sato. Shinri . . . Savelli. Gianfranco .. Schiavello. Mario . . Schulte.Frohlinde. Dietrich . Schultze. Hartmut . Schwarz. G . . . . . . 378 . . 2107 . 657. 2361 . . 1623 . . 1727 . . 2031 . . 2121 . 1667 . . 1181 . . 583 . 2045 . . 1573 . . 2045 . 1631. 1642 . . 1099 . . 1099 . . 815 Scaiano. Juan C . . . Scott. Alan C . . . Scott. Peter D . . Scott. Robert A . . Scurrell. Michael S . . Seelig.LOffler. A . . Sehested. Knud . . Sellers. Robin M .. Senkyr. Gerda . Sherwood. John N . . Sherwood. Peter M . A . . Shuck. Edward L . . Sime. S . J . . . . She. W . J . . . . Singh. Prem Paul . . Skinner. Henry H . . Skipp. Christopher J . . Smitharn. James B . . Smith. Barry C . . Smith. David G . . . Smith. Graham H . . Smith. Robert A . . Soni. Vimal . Sparasci. Antonio . . Spencer. H . Garth . Spiro. Michael . . Sprake. C . H . S . . . Stanford. J . L . . . Staveley. Lionel A . K . . Stepto. R . F . T . . . Stokes. R . H . Stone. Frank S . . . Strach. Steven J . . . Strange. J . H . Stul. J . S . . . . Suppan. Paul . . Sutt0n.H. C . . Swallow. A . J . . . Symons. Martyn C . R . . . . Taarit. Younes Ben . Tachikawa. Enzo . . Takaishi. T . . Tamaru. Kenzi . Tanaka. Kazuko . Tattershall. Bruce W . . Taylor. Joseph F . . Telford. Roland . . Tench. Anthony J . . Thain. Jennifer M . . Thanh. Ngyhen Thi . Thomas. John M . . . Thompson. John F . . Thornton. Edward W . . Thrush. B . A . . . Tornpkins. Frederick C .. Townsend. R . . . Townsend. Rodney P . . Trasatti. Sergio . . Taylor. J . A . G . . . Tedder. John M . . . 15 . 1221 875. 1953 . 1413 . 1849 . 903 . 815 . 1211 . 558 . 997 1392. 2331 . 298 . 2111 1432. 2366 . 2366 . 1528 . 699 . 1744 . 285 . 2232 389. 1379 . 2232 . 1991 . 1534 . 42 . 2459 . 988 . 35 1292. 1308 . 1372 1292. 1308 . 1707 . 1476 . 357 . 1610 . 1396 . 539 . 2142 . 2277 . 1772 . 1402 * 2121 . 97 . 2132 . 1127 . 1772 . 1161 . 2063 . 592 . 1352 . 1005 . 78 . 1172 . 1857 . 2260 PAQL 461. 1013. 2468 . . 2413 . . 930 . . 35 . . 803 . . . 123116 AUTHOR INDEX-VOLUME 71. 1975 Tremaine. Peter R . . Trenwith. Anthony B . . Trif~ro. Ferniccio . . TriIlo. J . M . . . Tullett. Arthur D . . . Tyrrell. H . J . Valentine . Ud Din Ghatti. Maraj . ulstrup. Jens . . Uytterhoeven. Jan . B . Valigi.M . . . Vandamme. Luc . J . Vardanyan. Irma A . Vedrine. Jacques C . Vidoczy. Tam’s . VWgstad. Einar Voice. Philip J . . Vukovic. 2 . . . Wagstaff. Kenneth P . Walch. Axel . . Walker. Raymond W . . Walker. T . . . . Wallbank. Andrew D . . Walsh. Robin . . Walton. John C . . . Ward. Anthony J . I . . Warman. John M . . PAGE . . 2170 . . 2405 . 209. 665 . . 961 . . 2361 . 42. 1744 . . 106 . . 435 . 1396. 1545 . 1631. 1642 . . 1545 . . 1203 . . 1697 . . 841 . . 2007 . . 780 . . 2017 . . 657 . . 501 736. 756. 1265 . . 946 . . 685 . . 1991 . . 592 . . 1106 . . 1252 Warner. Anthony E . M . Wasif. S a d . . Watson. Robert T . . Wayne,R.P. . . Waywell,D.R. . . Webb. Geoffrey . . Weeks. Thomas J . Jr . . Wells. CecilF . . . West. Kenneth 0 . . . Westwood. J . Vincent . White,L.R. . . White. Michael P . . Whittingham. K . P . . Whittle. Eric . . Williams. John0 . . Winter. John A . . . Wolf.AlfredP . . . Wood,C.J. . . W0olf.L . A . . . Wormald. Christopher J . Yariv. S . . . . Yatsurugi.Y . . . Yosbimura. Nobuya . Yusa. A . . . . Zecchina. Adriano . Zimmerman. Albert H . . Zocchi. Fernando . . PAOE . 670 . 1041 322. 336 . 610 . 1308 . 134 . 2051 . 1868 . 967 . 1762 . 1046 . 955 . 1407 . 1978 . 1857 . 1005 . 1905 . 1678 784. 1372. 2296 . 705. 726 . . 674 . . 97 . . 548 . . 97 . . 1476 . . 2111 . . 351

ISSN:0300-9599    

DOI:10.1039/F197571BA001

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Solute Diffusion in Hydrated Polymer Networks Part 2.-Polyacrylamide, Hydroxyethylcellulose and Cellulose Gels BY WYN BROWN* AND KALUBA CHITUMBO Institute of Physical Chemistry, Uppsala University, P.O. Box 532, 751 21 Uppsala 1, Sweden Received 6th March, 1974 Diffusion measurements on oligomeric solutes (n-alcohols, polyhydric alcohols, oligosaccharides) and polyethylene oxide polymers in sparsely crosslinked, water-swollen, polymer networks are described. Diffusion coefficients in the gel phase are much lower than those in the bulk liquid, which may be expressed as being mainly due to a change in the local viscosity. Diffusion coefficients ranged from 4.8 x (methanol) to 1.4 x lo-' cm2 s-l (polyethylene oxide 4000) in the hydroxy- ethylcellulose gel. The corresponding ratios to the free diffusionvalues are 0.31 and 0.106.Thereduc- tion of D for small solutes is shown to depend on the concentration and polar character of the matrix polymer but not on the degree of crosslinking of the network. Much of the stability of the postulated water structuring is attributed to the inertia of the polymeric component, with a minor contribution from the strength of the polymer-solvent interaction. Estimates of polymer-solvent interaction energies in the gels are obtained from diffusion data at different temperatures. In Part 1 were described measurements of the translational diffusion coefficients of the same solutes as discussed here in a densely crosslinked, water-swollen, cellulose gel C80.I It was found that diffusion coefficients in the gel phase are substantially lower than the free diffusion values.It was concluded that this effect may be expressed as a network-induced increase in the microscopic viscosity of the medium. This paper describes an extension of these investigations to solute diffusion in loosely crosslinked gels of cellulose, hydroxyethylcellulose and polyacrylamide. The purpose was to establish the generality of the findings of Part 1 and furthermore to examine to what extent the " tightness '' (average pore size) of the matrix determines the magnitude of the microscopic solvent viscosity. In contrast to highly crosslinked gels, steric exclusion of the solutes with the gels used here is negligible (that is, there is an insignificant variation in the effective internal volume of the gel with respect to both solute size and temperature), except in the case of the polymeric solutes.EXPERIMENTAL Experiments were made using the apparatus and technique described in Part 1.l The preparation of the cellulose and hydroxyethylcellulose gels, employing epichl oro- For clleu- hydrin as the crosslinking medium, followed the procedure described in Part 1. TABLE 1 .-STRUCTURAL PARAMETERS OF GEL SAMPLES (WATER-SWOLLEN DISCS ; DIAMETER 4 Cm) plolymerl water content effective internal x (wlw) cm3 volume a/cm3 hicknesslcm pol yacrylamide 6.8 3.03 2.05 0.26 hydroxyethylcellulose 5.4 5.22 3.60 0.43 cellulose (C5) 17.0 6.63 5.60 0.60 a volume of gel solvent diluting the bathing solution. 12W. BROWN AND K . CHITUMBO 13 lose ((3, 0.47 g epichlorohydrin per 100 g viscose was used.With hydroxyethylcellulose a large excess of epichlorohydrin was necessary in the reaction : 100 g epichlorohydrin per 5 g polymer was used. The polyacrylamide gel was a gift from Pharmacia Fine Chemicals, UppsaIa, Sweden. Further details concerning the gels are given in table 1. RESULTS AND DISCUSSION The results of the diffusion measurements are summarized in table 2. TABLE 2.-DIFFUSION COEFFICIENTS IN VARIOUS WATER-SWOLLEN GELS AT 25°C solute polyhydric alcohol ethylene glycol glycerol erythrit ol arabitol mannit 01 n-alcohol methanol ethanol propanol butanol oligosaccharide glucose sucrose raffinose stachyose polyethylene oxide PEG 1000 2000 3000 4000 hydroxyethyl- cellulose D x I06[cmz s-l 3.9, 2.76 2.56 2.32 3-06 4.78 2.50 4-02 2.62 1.98 1.47 1.22 DX 107fcrnZs-1 9.01 3.3 (j 1.41 4.05 cellulose (C5) D x 106/cmz s-l 3.46 2.82 2.57 2.19 1.88 3.88 - - - 2.05 1.51 1.15 0.94 - - - - polyacryl- amide D x 106/cm2 s- 2.12 1.99 1.84 1.72 1.61 - - - - 1.68 1.25 1.12 0.96 DX 107/crn2s-1 3.93 2.26 1.51 0.88 Fig.1 shows diffusion coefficients for some non-ionic solutes in the polyacrylamide gel. The diffusion coefficients are considerably smaller than those characterizing diffusion of these solutes in bulk water, as was observed in similar measurements in a densely crosslinked cellulose ge1.l These findings also apply to the diffusion data for these solutes obtained in sparsely crosslinked cellulose and hydroxyethylcellulose gels (fig. 2 and 3). With the smallest solutes, there is a reduced dependence of D on molecular size with each gel as was noted in Part 1.Diffusion measurements were made on tritiated water in the polyacrylamide, hydroxyethylcellulose and cellulose gels and the results are summarized in table 3(a). In agreement with our earlier findings, we conclude that the reduction in. D, relative to the diffusion coefficients in the bulk liquid, arises from an increased local viscosity of the solvent caused by its interactions with the polymer matrix, i.e. the factor limiting the diffusion coefficients of small solutes in such systems is the mobility of the solvent in the gel.14 DIFFUSION IN POLYMER NETWORKS An alternative viewpoint is provided by the two-phase model developed by Wang in which the diffusing entity is considered to encounter barriers in the form of macro- molecules with a hydration layer surrounding them.From the ratio of the gel and free diffusion coefficients for a given solute, one may estimate the volume of the hyd- rated polymer. The resulting hydrations of the hydroxyethylcellulose and cellulose gels are 6.7 g H20 and 1.7 g H20 per g gel, respectively (cf. the value of 3.4 g H20 per g for agar gels 3). Fig. 4 compares the results for the “ tight ” and “ open ” cellulose gels which have the same polymer concentration but which differ substantially in degree of crosslink- ing. The divergence observed for the larger solutes reflects the greater probability of solute-segment contacts in the more extensively crosslinked network when the solute 0.5 0- n U ?, - \ GEL DIFFUSION 1 I I 1 1.7 1.9 2.1 2.3 2.5 2.7 2.9 3.1 3.3 3.5 3.7 log vm FIG.1.-A comparison of transport in the polyacrylamide gel with free diffusion data. Double logarithmic plot of diffusion coefficients as a function of solute partial molar volume for polyhydric alcohols (0); oligosaccharides ( A) and polyethylene oxide polymers (0). - -A- - - I --.-- GEL DIFFUSION I 1 I I 1 I I I 13 1.5 I .7 1.9 2. I 2.3 2.5 2.7 log iim FIG. 2.-Double logarithmic plots of gel diffusion data (cellulose gel) and free diffusion data as a function of solute partial molar volume. Polyhydric alcohols (0); oligosaccharides (A). Points for methanol (m) and tritiated water (A) are included.W. BROWN AND K . CHITUMBO 15 n - 0 - L -2 2 ,.$ -0.5- % W tQ - - 1.0- size approaches the mean segment-segment dimension.The convergence at small molecular size shows that the increase in local viscosity is the same in both systems, that is, it is a function of the cellulose concentration only. In studies on the perme- ability of polyacrylamide gels, White also observed that the diffusion coefficient is independent of the degree of crosslinking. 1 7 1.0 cn 0.5 \ R E DIFFUSION \\ GEL DIFFUSION O t I I I I I I I I I 1.1 1.3 1.5 I .7 1.9 2 .I 23 2.5 2.7 log Em (4 1 1 I I I I t 2.7 2.9 3.1 33 3.5 3.7 log urn (6) FIG. 3.-(a) Gel diffusion data (hydroxyethylcellulose gel) and free diffusion data in a double logarithmic plot as a function of solute partial molar volume for n-alcohols (m), polyhydric alcohols (0), oligosaccharides (A) and tritiated water (b. (b) An analogous plot for polyethylene oxide polymers.Nishijima and Oster studied the translational diffusion of sucrose in concentrated aqueous solutions of polyvinylpyrolidone using an interference technique. They found that the diffusion coefficient decreased with increasing polymer concentration and reached a minimum asymptotic value of about 4 to 5 times lower than the diffusion16 DIFFUSION I N POLYMER NETWORKS value in water alone. These results were interpreted in terms of a corresponding increase in the local viscosity of the medium. The magnitude of the effect is similar to that found here [table 3(a)]. TABLE 3.-DIFFUSION COEFFICIENTS OF TRITIATED WATER IN GELS (a) sparsely crosslinked, water-swollen at 25°C gel gel diffusion self diffusion D g e i X 106/cmz ssi DgX 106/cm2 s-l l?pe1*/10-3 Pa s ~g:/lO-~ Pa s h y drox ye t hylcel lul ose 5.42 24.4 4.0 0.89 cellulose (C5) 4.03 24.4 5.39 0.89 polyacrylamide 2.9, 24.4 7.3 0.89 (b) densely crosslinked cellulose gel (C80) at different temperatures temperaturelac Dgel x 106/cm2 s-l vgel*/10-3 Pa s 10" 3.1 6 6.87 17" 3.48 6.24 25 " 3.87 5.61 35" 5.91 3.67 31" 5.04 4.30 ; see ref (15).D: 3: +?gel = p HTO L 2 !3 0 - M 0 - -05 - I I 13 1.5 1.7 1.9 2.1 2.3 2.5 log urn FIG. 4.-A comparison of gel diffusion data [cellulose gels: (a) = tightly crosslinked, (b) = sparsely crosslinked] in a double logarithmic plot against solute partial molar volume. Polyhydric alcohols (0); oligosaccharides (A). Points for methanol (M) and tritiated water (A) are included. Similarly, Biddle estimated microscopic viscosities in concentrated hydroxyethyl- cellulose solutions by measuring the rotational diffusion of dissolved fluorescein molecules using the fluorescence depolarization technique.' His estimated local viscosities are similar to those given in table 3(a) and he showed that this parameter increased linearly with polymer concentration.The results of Nishijima and Oster and Biddle for concentrated solutions are thus in agreement with the present findings for diffusion in the gel phase. One concludesW . BROWN A N D K . CHITUMBO 17 that the water mobility depends on the nature of the polymer and its concentration, independently of whether it is present as a tightly or loosely crosslinked network or is in true solution. Similar conclusions were reached by Pika1 and Boyd in a study of ion transport in solutions of sodium polystyrenesulphonate and in the crosslinked polyelectrolyte.Taken together, the results suggest that the increased local viscosity attains a limiting value, as found by Nishijima and Oster.6 THE INFLUENCE OF TEMPERATURE I N SPARSELY CROSSLINKED GELS Data for the temperature dependence of the diffusion coefficient of glucose in the gels are summarized in table 4(a) and shown in the Arrhenius plot, fig. 5. Apparent activation energies, calculated from the slopes in the intervals greater and less than 25"C, are given in table 4(b). The following interpretation is put forward. In the low temperature interval ( < 25"C), the apparent activation energy is smaller than that describing transport in the bulk liquid by the quantity of energy characterizing the interactions between solvent and the relatively immobile polymer segments.Above the observed transition temperature of 25"C, at which point the polymer segments acquire a greatly increased freedom, the apparent activation energy will exceed, by a corresponding amount, that characterizing transport in the bulk liquid. 103 KIT FIG. 5.-Arrhenius plots of glucose diffusion data in various gels. Hydroxyethyl cellulose (0): cellulose (0) ; polyacrylamide ( A). Equivalently, one may regard the apparent activation energy in the low temperature interval as reflecting the small scale segmental motion of the polymer and in the higher temperature interval the energy associated with the dynamic motion of large segments of polymer chains.The change in slope is apparently analogous to that found by Kumins et aL8 for a polymer-solvent system at the glass transition temperature of the polymer (i.e. the point of change from a quasi-crystalline state to a liquid-like one). These workers studied the diffusion of water through a vinyl chloride-vinyl acetate copolymer. Such behaviour in polymer-solvent systems is now well-established, see, for example, the discussions of Stannett and William~.~~ lo Values of the total energy involved in the transition, LIET, are included in table 4(b). Note that, in contrast to the18 DIFFUSION I N POLYMER NETWORKS densely crosslinked network (see next section), there is no change in the effective inter- nal volume with temperature change in any of the gels and glucose simply serves as a label reflecting the dependence of the average local viscosity of the medium on temp- erature.The slopes in the two temperature intervals are consistent in that they separately yield identical values of the " excess energy ", AEE, associated with the solvent mobility [table 4(b)]. The excess energy characterising the solvent in the swollen gel is defined here as the difference between the apparent activation energy in the relevant temperature interval and that characterizing the bulk solvent (average value estimated from the present data 20.5 kJ mol-l) and should provide a measure of the stability of the polymer-solvent interaction to thermal fluctuations (i.e. it is essentially a binding energy). It is then to be expected that D will decrease with increasing interaction, as is shown to be the case in fig.6. AEElkJ mol-' FIG. 6.-The dependence of the diffusion coefficient of glucose (25°C) on the excess energy character- izing the water-polymer interactions in the different gels : hydroxethylcellulose (O), cellulose (0) and polyacrylamide (A). It is also to be expected that the magnitude of AEE is related both to the concentra- tion of polymer segments and their chemical nature, the value of AEE increasing with polar character. One observes, however, that there is a large reduction in the D- value for glucose in comparison with free diffusion even when the interaction energy is small as in the case of the hydroxyethylcellulose gel (2.5 kJ mol-l). As pointed out by Fenichel and Horowitz," this must mean that the main contributing factor to the stability of the postulated water structuring is provided, not by the strength of this short-range interaction, but by long-range forces imparted by the great inertia of the macromolecular component. This structuring profoundly influences transport properties.TEMPERATURE DEPENDENCE-DENSELY CROSSLINKED GEL Fig. 7(a, b) show Arrhenius plots of the diffusion data for tritiated water [table 3(b)] and glucose, respectively, in the tightly crosslinked cellulose gel (C80). Appar- ent activation energies in the high and low temperature intervals are included in table Fig. 7(a) has the same symmetry as the curves in fig. 5 and the apparent activation energies and the derived excess energy are identical to the values characterising solvent mobility in the " open " cellulose gel.Since the polymer concentration in the two gels is the same, this is expected. Fig. 7(b) is markedly different. Since in the tight gel glucose is physically excluded to a degree dependent on the temperature,l the slope in the low temperature interval probably reflects the total energy associated with 4(b).W. BROWN AND K . CHITUMBO 0.8 0.7 h I cn - "E -2 0.6 E: X 9 W M 2 05 19 . 32 5- 3.3 3.4 3.5 3.3 3.4 35 103 K/T 103 KIT (4 (6) FIG. 7.-Arrhenius plots of diffusion data for : (a) tritiated water and (b) glucose in a tightly cross- linked cellulose gel. TABLE 4.-(a) DIFFUSION COEFFICIENTS FOR GLUCOSE IN VARIOUS GELS AS A FUNCTION OF TEMPERATURE hydroxyethylcellulose cellulose (C5) polyacrylamide temp./'C D X 106/cm2 s-l temp./"C DX 106/cm2 s-l ternp./OC DX 106/cm2 s-' 10 1.79 10 1.70 10 1.63 15 2.1 6 13 1-76 17 1.67 20 2.28 17 1 .go 21 1.68 25 2.64 25 2.05 25 1.68 30 3 -08 30 2.56 30 2.26 35 3.83 35 3.18 35 2.80 (6) THERMODYNAMIC PARAMETERS CHARACTERIZING POLYMER-WATER INTERACTIONS IN THE GEL PHASE gel apparent activation excess energy/kJ rno1-I energy*/ D25~ 1061 solute >25'C t25'C kJ mol-i kJ mol-I (glucose) - AET/ AEE/ crn2.s1 h ydroxyet hylcel lulose glucose 23.0 18.0 5.0 2.5 2.64 cellulose (C5) glucose 32.2 9.6 22.6 11.3 2.05 pol yacrylamide glucose 38.9 2.1 36.8 18.4 1.68 cellulose (CSO) tri tiated water 3 2.6 9.6 22.6 11.5 - glucose 21.3 32.2 - 10.9 1.55 * relative to the value (20.5 kJ mol-') characterizing bulk water.20 DIFFUSION IN POLYMER NETWORKS solvent mobilization accompanying activation of the polymeric component.Above 25°C the apparent activation energy has decreased by an amount equal to the polymer- solvent interaction energy (i.e. 10.9 kJ mol-l). Such temperature dependences are characteristic for diffusion in water-swollen networks and have been noted, for example, by Shuler et al. for the diffusion of sucrose through collodion membranes and by Boyd and Soldano l3 for the diffusion of sodium ions in cationic exchangers. INFLUENCE OF SOLUTE CHARACTER Diffusion coefficients for some alkali-metal chlorides in the polyacrylamide gel (fig. 8) are compared in table 5 with values in a cellulose gel. The diffusion coefficients in polyacrylamide are exceptionally low and, as shown by the figures for the effective volume fraction, these solutes are partially excluded from the polyacrylamide matrix.As pointed out by Franks, who compared ionic transport in water and ice,14 any factor which promotes water structure will reduce the mobility of alkali-metal ions. The lower D-values in polyacrylamide are thus to be expected owing to the stronger polymer-water interactions in this system. 0 - 0.5 GEL DIFFUSION O t I I I I I I I log vm FIG. 8.-T\ouble logarithmic plot of gel diffusion data (polyacrylamide gel) and free diffusion data for alkali metal chlorides as a function of solute partial molar vo1ume.l 0.7 0.9 1.1 1.3 1.5 1.7 TABLE S.-DIFFUSION COEFFICIENTS FOR ALKALI METAL CHLORIDES AND OTHER SOLUTES IN POLYACRYLAMIDE AND CELLULOSE GELS IN WATER AT 25°C polyacrylamide cellulose free diffusion solute D X 106/cmZ s-l vr* D x 106/cm2 s-l Vf* D x 106/cm2 s-l LiCl 1.22 0.428 3.66 1 .o 13.12 NaCl 1.88 0.373 4.25 1 .o 15.45 a KCI 1.95 0.35, 4.42 1 .o 19.17 butylamine 0.50 0.24, - - - butyric acid 1.77 0.546 - - 9.18 - 10.1 butanol 1.62 0.68 6 - * Vf = effective internal volume expressed as a fraction of the total solvent contained in the gel.a Handbook of Chemistry and Physics (Chemical Rubber Co., Ohio, 1971-72); b ref. (16).W. BROWN AND K . CHITUMBO 21 In addition, the diffusion coefficients of butylamine and butyric acid are compared with the value for butanol in table 5. As expected with the basic amide groups of the polymer, butylamine has a much lower D-value than butanol. Furthermore, butylamine is most excluded from the network.On the other hand, the diffusion coefficient of butyric acid is significantly higher than the value for butane€. Table 2 includes values of the diffusion coefficients for some polyethylene oxide polymers in the hydroxyethylcellulose and polyacrylamide gels. The values decrease rapidly with increasing molecular weight and, furthermore, the values for the effective volume fraction indicate a progressive increase in the steric exclusion as found for a tightly crosslinked cellulose ge1.l This work is part of a research programme financially supported by the Swedish Forest Products Research Laboratory. Financial support from the Swedish National Science Research Council is also gratefully acknowledged. W. Brown and K. Chitumbo, J.C.S. Farau'ay I, 1975,71, 1. J. H. Wang, J. Amer. Chem. SOC., 1954, 76, 4755. A. G. Langdon and H. C. Thomas, J. Phys. Chem., 1971,75,1821. M. L. White, J. Phys. Chem., 1960,64,1563. M . J. Pika1 and G. E. Boyd, J. Phys. Chem., 1973,77,2918. Y . Nishijima and G. Oster, J. Polymer Sci., 1956, 19, 337. ' D. Biddle, Arkiv Kemi, 1968, 29, 553. * C. A. K d s , C. J. Rolle and J. Roteman. J. Phys. Chern., 1957, 61, 1290. V. T. Stannett and J. L. Williams, J. PuZymer Sci. C, 1965, 10, 45. 1968,118,177. lo B. P. Tikhomirov, H. B. Hopfenberg, V. T. Stannett and J. L. Williams, Mukromol. Chem., l 1 R. Fenichel and S. B. Horowitz, Ann. I?. Y. Acad. Sci., 1965, 125, 290. l2 K. E. Shuler, C. A. Dames and K. J. Laidler, J. Chem. Phys., 1949, 17, 860. l3 G. E. Boyd and B. A. Soldano, J. Amer. Chem. SOC., 1953, 75,6091, 6099. l4 F. Franks, Chem. and I d , 1968, 560. l5 J. H. Wang, C. V. Robinson and I. S. Edelman, J. Amer. Chem. SOC., 1953, 75, 466. l6 Landolt-Bornstein, 5 a, Section 2522 (Springer-Verlag, Berlin/Heidelberg, 1969).

ISSN:0300-9599    

DOI:10.1039/F19757100012

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Mossbauer Studies in the Colloid System P-FeOOH-P-Fe,O, : Structures and Dehydration Mechanism BY ARTHUR T. HowE*-~ Inorganic Chemistry Laboratory, South Parks Road, Oxford OX1 3QR KEVIN J. GALLAGHER Chemistry Department, University College, Swansea SA2 8PP AND Received 6th February, 1974 The Mossbauer spectrum of the recently reported P-iron(ur) oxide has been measured and that of j3-iron(uI) oxide hydroxide re-investigated. The NCel temperature of P-iron(m) oxide lies between 300 and 380 K and the average magnetic field at 4.2 K is 49.5 tesla. Magnetic relaxation effects are observed. The spectrum above the NCel temperature is a broad asymmetric doublet, consisting of peaks having a continuous distribution of velocities, with the most probable quadrupole splitting occurring at 1.0 mm s-l, and the most probable isomer shift, extrapolated to 295 K, occurring at 0.31 mm s-'.While the distribution is within the known range of parameters of 6 or 5 co-ordinated Fe3+ in oxides, a proportion of tetrahedral Fe3+, with a lower isomer shift, could also exist within the broad spectral envelope, and a possible concentration range of from 0 to 40% was estimated. The spectrum of /3-iron(m) oxide hydroxide is consistent with octahedral Fe3+. The high pro- portion of ions on the internal surfaces of the tubular structure results in a non-stoichiometric surface anion excess, and the conventional formula 13-Fe00H has been reformulated as #?-FeOx(OH)3-,,, where xw 0.9. Four models of the anion vacancy distribution in /.?-iron(rrr) oxide are considered within the frame- work of the tubular structure retained during dehydration, and a defect structure is proposed which is consistent with the Mossbauer, X-ray, magnetic and B.E.T.data: all the internal Fe3+ ions are 5 co-ordinated whilst the surface Fe3+ ions are tetrahedrally co-ordinated. It has been reported that a new structural form of iron(r1r) oxide is obtained by vacuum dehydration of B-iron(Irr) oxide hydroxide. The new form, designated as the /3-form, retains all of the structural features of the parent compound. Colloidal particles of p-iron(rrr) oxide hydroxide are composed of tubular subcrystalline straws packed into cigar-shaped bundle^.^-^ The micro-structure is based on the hollandite unit cell (fig. l), which has a small central tunnel. The internal tube walls of the colloid are approximately 2 unit cells thick and define the large tunnels which are approximately 3 unit cells wide.2 A supercell, 5 x 5 unit cells, which describes such a structure is shown in fig.2. (The assignment of ions to the surface sites will be justi- fied later in the paper.) When the colloid is initially formed chloride ions occupy the small tunnels, as represented approximately by the formula FeO -%(OH) +&1, with water in the large tunnels. The structure shown in fig. 2 is obtained by thorough washing and subsequent drying, and is the material usually referred to as washed and dried /3-FeOOH. Upon dehydration of this compound to P-iron(rr1) oxide, of nominal formula P-Fe203, neither the morphology, as shown by electron microscopy, nor the total j- present address : Department of Inorganic and Structural Chemistry, The University of Leeds, Leeds LS2 9JT.22A . T. HOWE AND K . J . GALLAGHER 23 a = 1.048 nm 0 O/O H at z = 0 0 O/O H at z = &+ O F e a t z = O O F e a t z = &+ FIG. 1.-Projection along (001) of' the unit cell of the hollandite skeletal structure, upon which the structure of p-iron(m) oxide hydroxide is based. The small tunnels are shown unoccupied. The A site anions (uncrossed) are co-ordinated to 3 Fe3+ ions in a trigonal pyramid. The B site, bridging anions (crossed), are co-ordinated to 3 Fe3+ ions in a plane. This arrangement, together with the fact that the A sites are probably occupied by OH- ions while the B sites are occupied by 02- ions, inevitably leads to a distorted octahedral arrangement around each Fe3+ ion.FIG. 2.-One unit cell of the idealised super-structure of /3-iron(m) oxide hydroxide, FeI28O1 6(0H)15a, in which the numerous smaller hollandite cells can be recognised. The inset shows how replication of the supercell produces the regularly repeating large tunnels within the structure. B site anions are crossed, and the surface sites S1 and $2 are shaded.24 M o SSB A u E R ST u D I E s o F /3-FeOOH-p-Fe203 surface area, as obtained from B.E.T. measurements, alters significantly. Similarly no change can be observed in selected area electron diffraction, while the X-ray diffraction pattern reveals only a 2 % contraction along the c-axis of the tetragonal, hollandite-like cell, together with minor intensity changes, indicative of a topotactical reaction.It is surprising that removal of 25 % of the anions from such an open lattice does not result in more noticeable structural changes, and we have therefore investigated the Mossbauer spectrum of the new p-iron(m) oxide, together with that of the parent (washed and dried) oxide hydroxide, since the particle dimensions, important in determining the magnetic relaxation effects, were not known in a previous investiga- tion of the parent c~mpound.~ Mossbauer studies have also been reported *-lo for the unwashed oxide hydroxide (misleadingly also referred to as p-FeOOH), which, we point out, has some significantly different Mossbauer parameters from the washed and dried compound. EXPERIMENTAL The washed and dried p-iron(m) oxide hydroxide was prepared by the same method as for the previous B.E.T.and X-ray studies.ll The material thus obtained from slow hydrolysis of a hot ferric chloride solution was washed 12 times with distilled water and vacuum dried at room temperature to remove all water from the inner and outer surfaces. Dehydration in high vacuum at 443 K produced p-iron(n1) oxide which was stable in air at room temperature. This product, however, still retains 0.2 moles of H20 per mole of Fe203, as determined by the weight loss upon transformation at 873 K to the stable form, a-Fe203, and subsequent high- t emperature firing. Thin Mossbauer absorbers containing 3.7 mg cm-2 Fe were used to obtain the high quality spectra in the paramagnetic region so that saturation corrections would not be necessary. These samples were mixed with a tenfold excess of inert boron nitride powder to obviate the possibility of non-random particle orientation.The other spectra were ob- tained from undiluted samples having 15 mg cm-2 Fe. Mossbauer equipment of the type described by Cranshaw l2 was used with a 30 mCi J7Co/Pd source. Isomer shifts are quoted with respect to Fe metal at 295 K as zero. Several samples were heated in vacuum in the BN sample holder of a Ricor Mossbauer furnace regulated to within 1 K by a Eurotherm temperature controller. Allowance for a parabolic baseline caused by the geometric effect was included in the Harwell computer fitting programs. RESULTS AND DISCUSSION MOSSBAUER SPECTRA AND MAGNETIC PROPERTIES The spectrum of the oxide hydroxide is shown in fig.3. At 295 K only a broad paramagnetic doublet is observed. The parameters are given in table 1 and lie within the error range of those previously rep~rted.~ DCszi identified 295 K as the NCel ternperat~re,~ and found that, below this temperature, a magnetically-split 6 peak pattern developed while the paramagnetic doublet, although diminished in intensity, simultaneously persisted down to about 200 K. Our very broad spectrum at 77 K shows the residual effects of such magnetic relaxation behaviour. The most pro- nounced magnetic hyperfine field of 46.0 tesla (1 tesla = 10 kG), when extrapolated to 0 K, gives a limiting value of 47.0 tesla, as compared to 47.5 tesla found by DCszi.’ These values are significantly different from the value of 48.5 tesla at 80 K obtained from unwashed samples of the oxide hydroxide.1° The difference, which probably arises from the change in Fe3+ co-ordination due to the additional protons needed to balance the charge of the C1- in the small tunnels of the unwashed samples, has not been previously noticed, and is to be compared with the similarities in the otherA .T. HOWE AND K. J . GALLAGHER 25 Mossbauer parameters and in the magnetic behaviour. We found that the presence of water in the tunnels, as in washed but undried material, did not alter the spectrum compared to the washed and dried samples. FIG. -10 -5 0 5 10 velocity/mm s-' 3.-Mossbauer spectra of B-iron(nI) oxide hydroxide (more exactly ~-FeO~OH)3-2x where x = 0.91) at 77 and 295 K.The baseline is the same for both spectra. TABLE 1 .-MOSSBAUER PARAMETERS most Drobable isomer shift/ most probable mm r1 at 295 K w.r.t. Fe at 295 I< quadrupole sqlittingl most probable H / mm s- T at 4.2 K p-iron(m) oxide hydroxide 0.37 0.7 47.0 p-iron(m) oxide 0.31 1 .o 49.5 Fig. 4 shows the Mossbauer spectrum of p-iron(III) oxide over a range of tempera- tures from 4.2 to 383 K. The parameters are given in table 1. The most noticeable change from the spectrum of the oxide hydroxide is the increase in the Ndel temp- erature, which now lies above 295 K but less than 383 K. Because of slight variations in the shape of the spectrum from sample to sample a more accurate determination of the NCel point was not warranted. The superimposition of magnetically-split and paramagnetic components in the spectrum at 295 K resembles the behaviour of the oxide hydroxide, although the peaks in the spectrum at 77 K are now much sharper (cf.fig. 3 and 4). At 4.2 K relaxation effects should be absent, and the large half- widths of the peaks, which are up to 4 times the expected single peak value, providing a range of magnetic fields from 48 to 51 tesla about the mean of 49.5 tesla, indicate a range of Fe3+ environments in the lattice. The application of a 0.5 tesla magnetic field did not alter the shape of the spectrum at 295 K, indicating an antiferromagnetic ordering, in agreement with the absence of particle movement in a magnetic field. Our samples, of mean dimensions 60 x 60 x 350 m, were at no stage ground, and the two methods of preparing the Mossbauer samples gave the same Mossbauer spectra.26 MOSSBAUER STUDIES OF P-FeOOH-P-Fe,O, .. h -10 - 5 0 5 I0 velocity/mm s-l FIG. 4.-Mossbauer spectra of ,!3-iron(rn) oxide (more exactly fl-Fe0,(OH)3-2x where x = 1.45) over the temperature range 4.2 to 383 K Since the basic double tunnel structure and antiferromagnetic ordering are com- mon to both the unwashed and washed oxide hydroxide and to the oxide, it is likely that the origins of the magnetic relaxation effects found for the oxide are lsimilar to those proposed for crystals of the other two compounds of comparable size. DCszi ’ attributed the effects in the washed and dried oxide hydroxide to domain magnetisa- tion reversal in the small particles,l in a way resembling superparamagnetic behaviour in ferromagnetic mafe~ia1s.l~ However, Syzdalev l has proposed that similar data for the antiferromagnetic a-FeOOH and ferrimagnetic a-Fe,O, are best interpreted in terms of low-temperature paramagnetic behaviour in particles having less than a critical volume.Voznyuk and Dubinin l6 proposed, along the same lines, the pres- ence of paramagnetic surface layers on the otherwise antiferromagnetic particles of unwashed p-iron(II1) oxide hydroxide. These effects in small antiferromagnetic particles are, however, not yet well understood. DEHYDRATION OF P-IRON(III) OXIDE HYDROXIDE TO P-IRON(III) OXIDE A N D THE PARAMAGNETIC MOSSBAUER SPECTRA The asymmetry just evident in the paramagnetic spectrum of the oxide at 383 K (fig. 4) was further investigated using a low-velocity scan on a sample prepared in situA .T . HOWE AND K . J . GALLAGHER 27 in the Mossbauer furnace from a thin absorber of the oxide hydroxide. Fig. 5(a) shows the spectrum of the oxide hydroxide taken in vacuum at 295 K. The tempera- ture was then slowly raised to 443 K and after the initial vacuum of approximately Torr had been restored the spectrum, now of the oxide, shown in fig. 5(b), was collected. The spectrum was unaltered by continued heating at 443 K. A compar- ison of the two paramagnetic spectra could not be made at the same temperature since the oxide is below its Ntel temperature in the stability range of the oxide hydrox- ide. The oxide sample was heated further to 523 K. Between 383 and 523 K there was no observable change in either the value of the quadrupole splitting or the asymmetry of the peaks, providing evidence that the asymmetry did not originate from magnetic relaxation effects., I I I I I l I l I - 1 0 1 2 velocity/mm s-* FIG. 5.-The change in the paramagnetic Mossbauer spectrum in going from p-iron(rr1) oxide hyd- roxide (a), recorded at 295 K, to &iron(m) oxide (b), brought about by heating in vacuo at 443 K and recorded at this temperature. The horizontal lines show the previously observed ranges of the isomer shifts of tetrahedral (lower velocity) and octahedral (higher velocity) Fe3+ in oxides (a) at 295 K and (b) extrapolated to 443 K. One of the many possible computer fits for the oxide is shown, and this particular one is consistent with the structural model proposed, and has 15% tetrahedral Fe3+ (dashed lines), compared to 17 % in the model.The remaining doublets have isomer shifts at the low end of the octahedral range and are therefore consistent with the 83 % 5 co-ordinate Fe3+ proposed. The asymmetry of the oxide spectrum could not have arisen from preferred particle orientation since the material was diluted with boron nitride. Dilution did not noticeably alter any of the spectra, but was done as a precautionary measure in view of the cigar shape of the crystals. However, in such tubular structures the surface ions could possess an anisotropic recoil-free fraction, or f factor, leading to unequal areas of the quadrupole components. ' 9 * The effect, however, is not evident in the spectrum of the oxide hydroxide, and computer fits to both spectra, to be discussed later, were consistent with components of equal area. Neither the oxide hydroxide nor the oxide showed any evidence of having unusually low overallffactors due to the tunnel or defect structures.Only a small reduction (ca. 10 %) in the total area response at room temperature accompanied the conversion of the oxide hydroxide to the oxide, showing that theffactor of the oxide was only marginally lower than that of the oxide hydroxide. Available evidence indicates that the f factors of the surface ions would still be appreciable. Iron ions located in the large tunnels of zeolite A have anffactor at room temperature of 0.57, corresponding28 to a Debye temperature of 265 K. The conclusion is also supported by calculations for the surface layer of Fe The asymmetry of the oxide spectrum would appear to arise from a range of isomer shifts of the many unresolved components arising from the tubular structure, and in view of the high proportion of anion vacancies the isomer shifts may indicate the presence of 5 or 4 co-ordinate Fe3+.In oxides, the isomer shift decreases with increas- ing proportion of s character in the bonds, the more covalently bonded tetrahedral configuration producing a lower isomer shift than octahedral coordination. Isomer shifts can be validly compared between different oxides provided that the actual temperature of measurement was not appreciably lower than the Debye temperature of the solid. At low temperatures departures from the classical limit of the second order Doppler shift occur and render isomer shift comparisons less reliable.A limit of the second order Doppler shift of -7 x mm s-l K-l is indicated by studies on Fe3+ oxides,21 and applies to typical oxides above approximately 200 K,22 for which the Debye temperatures are 300-400 K.23 The possible low Debye temperature of our tubular materials indicates that our measurements have been made well within the applicability of the high temperature limit and can be validly compared to values from other oxides. For comparison, the velocities of components in the oxide spectrum measured at 443 K were calculated, by use of the above figure, to be 0.10 mm s-1 lower than they would have been at 295 IS. A survey of the numerous oxides studied 24p 2 5 shows that, at 295 K, the isomer shifts of Fe3+ in a state of 6 co-ordination range from 0.31 to 0.41 mm s-l, and those in a state of 4 coordination from 2 5 * 26 0.13 to 0.27 mm s-l.The few systems studied in which Fe3+ is 5 coordinated 27-31 cover a range from 0.22 to 0.33 mm s-l at 295 K. The established octahedral and tetrahedral ranges are drawn as horizontal lines for the oxide hydroxide (at 295 K) and the oxide (at 443 K) in fig. 5. It can be seen that both spectral doublets have a centroid within the range for octahedral Fe3+. That of the oxide is towards the lower end of the octahedral range and is therefore also consistent with 5 fold coordination. Furthermore, the asym- metry of the oxide envelope may indicate the presence of a small proportion of tetra- hedral Fe3+.It is clear, though, that the spectrum could not represent a predomin- ance of tetrahedral Fe3+ in the structure. Working on the basis of only the hollandite unit cell (fig. l), where each anion is coordinated to 3 Fe3+ ions, one might expect the removal of 25 % of the anions in going from FeOOH to Fe203 to result in 75 % tetrahedral Fe3+, with 25 % being 6 coordinated. Such a discrepancy with the observed spectrum may suggest the presence of 5 coordinated Fe3+ or that the features introduced by the supercell (fig. 2), primarily the internal surface, play a crucial role in determining the coordination of Fe3+ in the oxide. Since the coordination of the Fe3+ on the internal surfaces of the parent oxide hydroxide has not been considered previously, we shall, after the next section, establish this for use as a basis for considering the dehydration to the oxide.We shall first, however, estimate more precisely the maximum proportion of tetra- hedral Fe3+ in the oxide consistent with the observed Mossbauer envelope and the known isomer shift ranges for 4 and 6 coordinated Fe3+. M OSSB A u ER s T UD I ES OF P-FeOOH-jI-Fe,O, SPECTRAL ANALYSIS The surface induced distortions of the structure would result in a large number of Fe3+ environments and hence a large number of considerably superimposed Moss- bauer peaks, and a complete resolution of the envelope into these components would not be possible. The maximum proportion of the oxide envelope consistent withA . T. HOWE AND K. J . GALLAGHER 29 tetrahedral Fe3+ was therefore estimated by the following process of trial and error and successive approximation.Initially doublets (e.g. AA’), where the intensities and halfwidths of the lorentzian peaks A and A were constrained to be equal, were fitted to the envelope. A minimum of 4 doublets was required, and by choosing starting parameters covering a wide range of possibilities, several solutions were found with satisfactory l2 values in the range 190 to 230 for 205 degrees of freedom. The simplest solution was of the symmetrical form ABCDD’C’B’A‘. The 4 isomer shifts were all higher than the tetrahedral range and were consistent with 6 or 5 coordinate Fe3+. Tetrahedral components could only be introduced by choosing the doublet pairs in an unsymmetrical fashion. A satisfactory fit with the grouping ABCDC‘D’B’A was found with the doublet CC’ having an isomer shift at 295 K of 0.27 mm s-l, at the top of the tetrahedral range and accounting for 8 % of the total area, while the isomer shift of the doublet DD’ (0.40 mm s-l at 295 K) was at the top of the octahedral range.A satisfactory fit could also be obtained with the grouping ABCDD’C‘A’B’ provided that an additional small doublet with an isomer shift in the 6 or 5 coordinate range was included at the velocities of peaks B and A’. In this case the doublet AA’ had an isomer shift in the tetrahedral range and accounted for 33 % of the area, while the doublet BB’ was still consistent with 6 or 5 coordinate Fe3+. Satisfactory groupings of the type ABCDD’B’C’A’ could not be found.As some of the peaks in the above fits had halfwidths up to 0.50 mm s-l, which is considerably Jarger than the value expected from a single absorption (0.23 mm s-l in the present case), the possibility exists of further subdivisions of these peaks to yield component peaks with a new set of isomer shifts. This was achieved by further constraining halfwidths of groups of doublets to be equal, a procedure which maxi- mised the choice of final doublet assignment since any one peak could be proportioned between two peaks at the same velocity and reallocated to two new doublets. Succes- sive introduction of components into the fits led to a statistically satisfactory 14 peak fit with component halfwidths of 0.23, 0.38 or 0.45 mm s-l. The symmetrical doublet grouping gave 7 doublets with isomer shifts, 6, between 0.30 and 0.36 mm s-l when extrapolated to 295 K, in the 6 or 5 coordinate range, with quadrupole splittings, A, between 1.71 and 0.40 mm s-l.By reassignment into 8 doublets, one new doublet consistent with tetrahedral Fe3+ (6 = 0.26 mm s-l at 295 K, A = 1.16 mm s-l) and another doublet having 6 = 0.35 mm s-1 at 295 K and A = 1.15 mm s-l could be introduced. The area of the tetrahedral doublet could be varied from 0 to 15 %, and the 15 % grouping, which is consistent with the model proposed later (17 % tetrahedral and 83 % 5 coordinate Fe3+) is shown in fig. 5(b). The remaining isomer shifts of between 0.30 and 0.36 mm s-l at 295 K are at the lower end of the octahedral range where 5 coordinate Fe3+ would be expected.Additional tetrahedral components could not be introduced without having isomer shifts higher than the known limit for both tetrahedral (0.27 mm s-l) and octahedral Fe3+ (0.41 mm s-I). In order to assess the effect of a rather unlikely uncertainty in these limits, all possible doublets with isomer shifts of up to 0.30 mm s-1 were assigned to tetrahedral Fe3+, which raised the total tetrahedral proportions to 38 %, with the highest isomer shift in the envelope being 0.42 mm s-l at 295 K. We have therefore taken 40 % as the upper limit for the tetrahedral Fe3+ proportion. Further refine- ments, such as an allowance for possible non-lorentzian behaviour, saturation effects and cosine effects would not be expected to alter the basic conclusion of the analysis. Fits to the spectrum of the oxide hydroxide were consistent with octahedral Fe3+, in accord with the known structure. The range of possible fits to both spectra prevents a unique assignment of the isoiner shifts and quadrupole splittings, and for30 MOSSBAUER STUDIES OF P-FeOOH-P-Fe,O, this reason the parameters derived from the individual fits have not been quoted.The data in table 1, indicating the parameters of the general envelope, should be interpreted in this light. The limiting magnetic hyperfine fields for the two compounds are consistent with the above possible assignments. This parameter is not generally diagnostic of coordination number. THE SURFACE STRUCTURE OF P-IRON(III) OXIDE HYDROXIDE: SURFACE CONTROLLED NON-STOICHIOMETRY I N P-Fe0,(OH)3-2x WHERE XEO.9 In the basic hollandite unit cell (fig.1) all the cations sites are identical, and there is an equal proportion of A and B anion sites. Four double strings containing edge- sharing [FeX,] octahedra run the length of the crystal in the c direction, and contain the A site (non-bridging) anions. The double strings are linked together through corner sharing via the B site (bridging) anions to form the small central tunnel. A recent neutron diffraction study 32 of the unwashed material, containing C1- in the small tunnels, showed that the octahedra were distorted in a manner which retained the unique Fe site, but which allowed for the preferential proton positioning closer to the A site oxygens within the double strings, but on the side of these oxygens facing the two adjacent bridging (B site) oxygens.A similar site preference of the protons for the A site oxygens may exist in the washed and dried p-iron(Irr) oxide hydroxide. Because the crystal diameter is only of the order of 10 times the 5a x 5a supercell dimensions, diffraction techniques would not detect the supercell, and the neutron diffraction study was therefore unable to determine the internal surface structure, as evidenced by the high R factor of 11 %. Inspection of the supercell (fig. 2) shows that the internal surface of the large tunnels, as well as containing A and B anion sites, also has 12 sites coordinated to one Fe3+ (S1 sites) and 12 coordinated to two Fe3+ (S2 sites). These are both shaded in the figure. The absence of anions from the S1 sites would create twelve 5 coordinate Fe3+ ions per supercell, while the absence of site S2 anions would create twelve 4 coordinate Fe3+ ions.Such a significant breakdown of the stable [FeX,] octahedra upon which the structure is based, resulting in a change of coordination of 18 % of the Fe3+, would be unlikely to occur under the mild conditions of room temperature vacuum drying, suggesting that these sites are in fact occupied. However, if the surface anion sites were occupied, there would be an excess of anions in the crystal over and above that represented by the formula FeOOH. Each hollandite cell exposed to the internal surface will be of composition M8X1,, compared to the normal cell of M8X16. The total composition will be M128X268. The form- ula which satisfies the charge balance is Fe12s0116(OH)152, or FeOo.906(OH)1.188. provide definite evidence for the anion excess structure.After a 5 % weight loss during prolonged vacuum drying at room temperature, during which water is removed from the internal and external surfaces, the oxide hydroxide lost a further 12 % water, as a percentage of the initial weight, upon dehydration and final conversion at 873 K to a-Fe203. This value agrees well with the value of 11 -2 % water loss calculated for the OH- excess structure above, but is considerably higher than the value of 7.6 % calculated for the OH- deficient structure corresponding to unoccupied S1 and S2 sites. The general formula which expressed the variability of the 02- to OH- ratio depending on the extent of surface controlled non-stoichiometry is Fe0,(OH)3 - 2x.This formula is preferable to the more conceptual formula for the anion excess oxide hydroxide of Fe01-x(OH)1+2, since it covers the whole range from Fe(OH), to Fe203. Crystalline products with the composition Fe(OH)3 are unknown, so the The results of the previous thermogravimetric studiesA. T . HOWE AND K. J . GALLAGHER 31 oxide hydroxide, in which x = 0.906, is probably as OH- rich as can be obtained in a crystalline phase. Since all the investigations of p-iron(II1) oxide hydroxide indicate the presence of both large as well as small tunnels, it would appear that the stoichio- metric formula of FeOOH does not apply to the normally prepared material, which should be represented as FeO,(OH),-,, with x M 0.9, the exact value depending on the exact dimensions of the large tunnels.Inclusion of the effects of the external surface of the crystals in this study only altered x from 0.906 to 0.898. The anion excess structure, in which all the Fe3+ ions are 6 coordinated, is consist- ent with the small range of isomer shifts found in the Mossbauer spectrum of the oxide hydroxide, and the surface induced distortions would alone be sufficient to account for the range of quadrupole splittings. Possible proton disorder may contribute further to the spread of quadrupole splittings. [The structure would allow an ordered proton arrangement, with protons occupying all the A (non-bridging) sites, as suggested by Gallagher,, together with all the surface sites to give Fe12 8(OH)t2 80: 1 (j(OH);:(OH);;l. THE DEFECT STRUCTURE OF P-IRON(III) OXIDE : P-FeO,(OH),-,, WHERE XM 1.45 Fe00.9(OH)1.2-+FeOo.g+,(OH)l.2-2,+yH,0.Anion vacancies (0-) are created in the retained supercell according to the equation 20H-+H20+02-+ 0-. If complete dehydration to P-Fe203 (supercell formula FelzzOl 92) owurred, 76 anion vacancies must be distributed amongst the 128 A sites, the 116 B sites and the 24 surface anion sites in Fe12801,,(OH)l,2. This figure may be reduced to about 62 anion vacancies due to the small quantity of water equivalents (2 % of the initial weight) still present in the oxide after prolonged evacuation. The reported i.r. peak indicates that these residual protons, like those in the oxide hydroxide, are not hydrogen bonded, but the shift of the stretching mode to even higher frequencies compared to the latter indicates that the protons are more tightly bound to oxygen, which could be produced by OH- bound to Fe3+ in a state of low coordination in the highly defective structure of the oxide. If OH- was present in the structure it would imply that the dehydration process had ceased before the eventual expected product of P-Fe203 had been reached, and we have therefore assumed, in the following, the minimum extent of dehydration corresponding to the residual presence of OH-.The maximum value of x in the formula Fe0,(OH)3-2, is thus calculated to be 1.45, slightly less than the value of 1.5 for Fe203. Four model anion distributions have been considered for the removal of 62 anions. If the surface anions are retained, the average Fe3+ coordination drops from 6 to 4.55 (model I).If the surface anions are removed together with anions from the interior, the average Fe3+ coordination drops to 4.83 (model 11). For each of these two basic possibilities, we have calculated the proportions of Fe3+ in each type of coordination assuming (a) that only 6 and 4 coordination exists and (b) that only 5 and 4 coordina- tion exists. The values are given in table 2. The models show such a wide range in the proportions of the three coordinations that we can confidently discriminate between them on the basis of (1) the Mossbauer results, showing less than 40 % tetrahedral Fe3+, (2) the constancy, to within 2 %, of the unit cell dimensions upon dehydration, and (3) the increased magnetic coupling in the oxide. Removal of water from the oxide hydroxide can be expressed by32 MOSS B AUER s T UDI E s OF /?-FeOOH-P-Fe,O, The Mossbauer results are highly inconsistent with the presence of 73 % tetrahedral Fe3+ [model I(a)] and are also inconsistent with model II(a) with 59 % tetrahedral Fe3+.The results would, however, be in very good agreement with model II(b), with only 17 % tetrahedral Fe3+, the remaining Fe3+ being 5 coordinated. TABLE 2.-POSSIBLE PERCENTAGE DISTRIBUTIONS OF DIFFERENT Fe3+ COORDINATION NUMBERS IN B-IRON(III) OXIDE six five four model I 73 (b)* - 55 45 - no surface anions removed (a) 27 model I1 all surface anions removed (a) 41 - 59 (6)" - 83 17 * shows models consistent with Mossbauer results. The a unit cell dimension depends primarily on the axial distances of the coordina- tion polyhedra, while the c dimension depends primarily on the equatorial distances.Since Fe3+ in a square planar coordination is unknown in oxides, 4 coordinate Fe3+ would distort to the tetrahedral position with a consequent substantial reduction in both the axial and equatorial polyhedral dimensions, and this would result in decreased unit cell dimensions. On the other hand, the removal of an A site anion from an octahedron would, in the absence of site relaxation, leave a square pyramid without altering either the axial or equatorial distances. Relaxation of the former equatorial anions towards the more symmetrical trigonal bipyramidal configuration would only be expected to affect the equatorial distances slightly, predicting only small changes in the unit cell dimensions, as found.The X-ray evidence thus strongly favours model II(b), with a predominance (83 %) of 5 coordinated Fe3+, and only 17 % tetrahedral Fe3+. The magnetic evidence also favours this model since a major alteration of the Fe-0-Fe angles is likely to significantly alter the superexchange couplings. The weight of evidence therefore suggests model II(b) as the most likely structure. Consideration of the supercell shows that the site distribution is quite a stable one for such a structure. Of the 128 Fe3+ ions in the supercell, 24 are at surface sites and 104 at internal sites. Now least disruption to the structure and magnetic interactions will occur if the tetrahedral ions are those at the surface, and 17 % of 128 is 21.8, which corresponds closely to the 24 surface Fe3+.The internal sites are therefore all 5 coordinated (106.2 calculated, as compared to 104 in the supercell). Creation of the internal anion vacancies exlusively on the A sites would leave all the B (bridging) sites occupied to interlock the double strings of ions. Five coordination can then be achieved by the staggered removal of every third A site anion down each c-axis string. Further dehydration would necessitate the energetically unfavourable creation of internal tetrahedral Fe3+ ions, suggesting a reason why the dehydration ceased at this stage when OH- still remained in the lattice. In fig. 5(b) a particular doublet grouping consistent with 17 % tetrahedral Fe3+ is shown for a 16 peak fit.The isomer shift at 295 K is 0.26 mm s-l. The other doublets have isomer shifts at 295 K ranging from 0.30 to 0.36 mm s-l and are consistent with 5 fold coordination. Five coordinate Fe3+, in the form of trigonal bipyramidal coordination, while not common, is well established in a number of oxides.A . T . HOWE AND K. J . GALLAGHER 33 For instance, in FeV04 there are 2 octahedral Fe3+ sites and one trigonal bipyramidal Fe3+ site 29 (having an isomer shift 6 = 0.325 mm s-l and quadrupole splitting A = 1.11 mm s-l at room temperature). In YMn03 the Mn site is trigonal bipyra- midal, and Fe3+ doped into this site retains the configuration 30 [6 = 0.30 mm s-l, A = 2.13 mm s-l, H(0) = 45.5 TI. In the oxygen deficient phase Sr,-,La,Fe0,.5+x/2 a trigonal bipyramidal site has been proposed 28 (6 = 0.33 mm s-', A = 0.33 mm s-', H(0) = 48.8 T), together with tetrahedral and octahedral sites in the structure.It would therefore not seem unreasonable for most of the Fe3+ [as in model II(b)] to have this coordination. The magnitudes of the quadruyole splittings in the oxide, which are larger than those in the oxide hydroxide, arise from the electric field gradients created around the 4 and 5 coordinate Fe3+ by the unsymmetrical anion environment. The range of quadrupole splittings would result from both the surface-induced distortions and the distribution of tetrahedral Fe3+ and residual OH-, which are probably randomly distributed over the A sites. The presence of the initially occupied surface sites and the comparative freedom of movement of the internal anions due to the tubular structure has allowed the structure to sustain such a large degree of ionic removal and rearrangement without collapse.The analysis has provided a valuable insight into the structure, which, because of its tubular form and range of surface-induced distortions would not be readily amenable to an accurate positional analysis by either X-ray or neutron diffrac- tion. Financial support from the S. R.C. through the Physico-Chemical Measurements One of us (A. T. H.) is grateful to the Royal Unit, Harwell, is acknowledged. Commission for the Exhibition of 1851 for a Fellowship. H. Braun and K. J. Gallagher, Nature Phys. Sci., 1972, 240, 13. K. J. Gallagher, Nature, 1970, 226, 1225. W. Feitnecht in 4th Int.Symp. Reactivity oj*Solids, ed. J. H. de Boer (Elsevier, Amsterdam, 1960), p. 583. A. L. MacKay, Miner. Mag., 1960, 32, 545. A. L. MacKay, Miner. Mag., 1962, 33, 270. J. H. L. Watson, R. R. Cardell and W. Heller, J. Phys. Chem., 1962, 66, 1757. 22, 617. M. J. Rossiter and E. A .M. Hodgson, J. Inorg. Nuclear Chem., 1965, 27, 63. 9T. Takada, M. Kiyama, Y . Bando, T. Nakamura, M. Shiga, T. Shinjo, N. Yamamoto, Y. Endoh and H. Takaki, J. Phys. SOC. Japan, 1964, 19, 1744; N. Yamamoto, T. Shinjo, M. Kiyama, Y . Bando and T. Takada, J. Phys. SOC. Japan, 1968,25,1267. l o P. 0. Voznyuk and V. N. Dubinin, Fiz. Tverd. Tela, 1973, 15, 1897 (Sou. Phys. Solid State, 1973, 15, 1265). l 1 K. J. Gallagher and D. N. Phillips, Chimia, 1969, 23, 465. T. E. Cranshaw, Nuclear Znstr. Mehthods, 1964, 30, 101. l 3 A. A. van der Giessen, Philips Res. Rep. Suppl., 1968, 12,40. l4 D. W. Collins, J. T. Dehn and L. N. Mulay, in Mussbauer Methodology, ed. I. J. Gruverman l 5 I. P. Syzdalev, Fiz. Tverd. Tela, 1970, 12, 988 (Sou. Phys. Solid State, 1970, 12, 775). l6 P. 0. Voznyuk and V. N. Dubinin, Fiz. Tuerd. Tela, 1973, 15, 1897 (Sou. Phys. Solid State, l7 P. A. Flinn, S . L. Ruby and W. L. Kehl, Science, 1964, 143, 1434. l8 A. A. Maradudin in Solid State Physics, ed. F. Seitz and D. Turnbull (Academic Press, New l9 L. Kumer, H. Posch and J. Kaltseis, Phys. Letters A , 1972, 40, 59. 2o L. Dobrzynski and P. Masri, J. Phys. Chem. Solids, 1972, 33, 1603. 21 J. J. van Loef, Physica, 1966, 32, 2102. ' I. Dkzsi, L. Keszthelyi, D. Kulgawczuk, B. Molnar and N. A. Eissa, Phys. Stat. Solid., 1967, (Plenum, New York, 1967), p. 103. 1973, 15, 1265). York, 1966), vol. 19, p. 1. 1-234 M o SSB A UER STUD IE s OF ~-FeOOH-/3-FFe,O3 22 F. van der Woude, Phys. Stat. Solid., 1966, 17, 417. 23 G. A. Sawatzky, F. van der Woude and A. H. Morrish, Phys. Rev., 1969, 183, 383. 24 N. N. Greenwood and T. C. Gibb, Miissbauev Spectroscopy (Chapman and Hall, London, 2 5 J. Danon in Chemical Applications of Mossbauer Spectroscopy, ed. V. I. Goldanskii and R. H. 26 C. Le Corre, A. Malve, C. Gleitzer and J. Foct, Compt. rend. C, 1972, 274, 466. 27 J. G. Rensen and J. S. van Wieringen, Solid State Comm., 1969, 7, 1139. 28 H. Yamamura and R. Kiriyama, Bull. Chem. Soc. Japan, 1972,45,2702. 29 B. Robertson and E. Kostiner, J. Solid State Chern., 1972, 4, 29. 30 J. Chappert, J. Physique, 1967, 28, 81. 31 F. P. Glasser, F. W. D. Woodhams, R. E. Meads and W. G. Parker, J. Solid State Chem., 32 A. Szytula, M. Balanda and Z . Dimitrijevic, Phys. Stat. Solid. (a), 1970, 3, 1033. 1971), Ch. 5. Herber (Academic Press, New York, 1968). 1972, 5, 255.

ISSN:0300-9599    

DOI:10.1039/F19757100022

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Thermodynamic Properties of Fluorine Compounds Part 15.l-Vapour Pressures of the Three Tetrafluorobenzenes and 1,3,5-Trichlor0-2,4,6-trifluorobenzene BY D. AMBROSE,* J. H. ELLENDER, C. H. S. SPRAKE AND R. TOWNSEND Division of Chemical Standards, National Physical Laboratory, Teddington, Middlesex TW 1 1 OLW Received 11 th April, 1974 The vapour pressures were measured of 1,2,4,5-tetrafluorobenzene from 5 kPa to the critical pressure, of 1,2,3,5-tetrafluorobenzene from 5 kPa to 470 kPa, of 1,2,3,4-tetrafluorobenzene from 7 kPa to 202 kPa, and of 1,3,5-trichloro-2,4,6-trifluorobenzene from 3 kPa to 180 kPa. The results were fitted by Antoine equations (applicable at pressures up to 200 kPa) and by Chebyshev equations. For each compound one of the latter equations is applicable over the whole range up to the critical temp- erature, and it is shown that the variation with temperature of AHlAZcalculated from this equation is similar to that found for other compounds.The vapour pressures, critical temperatures and critical pressures of some highly fluorinated aromatic hydrocarbons have been reported in earlier papers in this ser- ies 2-4 ; extension of the work to the four compounds named in the title is now described. Findlay has already published values of vapour pressure in the range 3 to 30 kPa for two of them, 1,2,3,4-tetrafluorobenzene and 1,2,3,5-tetrafluoro- benzene; the values presented here are more precise, and we believe them to be more accurate. EXPERIMENTAL The samples were from the same batches of material as were used in other investigations in this laboratory; their purities (table 1) had been assessed by low-temperature calor- imetry,6* but as a precaution the samples were re-dried by treatment with calcium hydride or molecular sieve immediately before use.For all the compounds, ebulliometric measurements were made in the range from about 5 to 200 kPa,* and their critical temperatures Tc and pressures pc were determined.g In addition, ebulliometric measurements were made at higher pressures on 1,2,3,5tetrafluoro- benzene and 1,2,4,5-tetrafluoroben~ene,~~ and on the latter compound measurements were made in the range 330 kPa to the critical pressure by both dynamic and static method^.^^ l 1 Temperatures were measured by means of platinum resistance thermometers used with equipment sensitive to 0.001 K, an automatic a.c.bridge (Automatic Systems) linked through a selector and control unit to an automatic typewriter and paper-tape punch for ebullio- metry, and a Mueller bridge (Leeds and Northrup) for the other measurements. Experiments with benzene, including some made with different samples, have shown that by comparative ebulliometry at pressures between 5 and 200 kPa the boiling temperatures of a compound of well established purity that is unaffected by decomposition may be reproduced within 0.005 K. At pressures p above 200kPa measured points obtained by the other methods seldom deviate from the curve best fitting all the points by more than 0.001~ ; this corresponds to an uncertainty in temperature of about 0.03 K near atmospheric pressure, increasing to about 0.1 K near the critical temperature.For benzene, points obtained in two other investiga- tions believed to be reliable as well as those measured in this laboratory all lie within 0.001~ of acurvefitted to the three sets of data. The estimated range of uncertainty in the measure- ments of critical temperatures is kO.05 K. 35TABLE 1 .-PURITIES/MOLES PER CENT ; VAPOUR PRESSURES BELOW 205 kPa ; Ap = Pobs-PcaIc WHERE Pcalc HAS BEEN OBTAINED FROM EQN (I), (11) AND (111). EQN (I) WAS OBTAINED BY FITTING TO THE RESTRICTED RANGE OF OBSERVED VALUES INDICATED BY THE ENTRIES IN THE COLUMN FOR Ap ; EQN (11) AND (111) WERE FITTED TO ALL THE VALUES TABULATED AplPa APlPa TIK plkPa (1) (11) (111) TIK plkPa (1) (11) (111) 1,2,3,4-tetrafluoro benzene (99.85) 1 ,2,4,5 - t e trafluor o benzene (99.94) 300.800 304.567 308.223 3 12.01 8 315.890 3 19.61 6 324.471 328.626 333.357 338.009 342.899 348.275 352.947 356.490 363.273 367.378 368.232 373.020 377.346 382.509 387.449 391.651 7.437 8.945 10.646 12.688 15.095 17.757 21.794 25.823 31.137 37.177 44.535 53.942 63.331 71.305 88.788 100.917 103.593 119,685 135.850 157.351 180.294 201.822 2 -0 0 -0 -1 0 -2 1 -4 0 -4 -0 - 5 -1 -4 -0 5 7 - 5 -5 0 -3 12 7 -4 -1 -7 0 3 -3 1 3 -0 -0 1 1 - 5 -3 -3 -2 -2 4 -8 -10 2 4 -2 15 -14 -24 -8 4 -10 -4 1,2,3, 5-tetrafluorobenzene (99.99) 287.593 291.414 295.169 298.091 303.399 307.963 3 12.406 314.635 318.751 324.852 329.202 334.506 3 3 8.643 343.224 347.715 352.951 356.719 357.291 357.743 362.582 366.716 371.603 376.423 381.524 5.719 6.987 8.457 9.773 12.598 15.541 18.930 20.846 24.802 3 1.779 37.666 46.007 53.501 62.91 1 73.374 8 7.269 98.504 100.305 101.755 118.237 133.926 154.528 177.195 203.907 4 1 1 -1 -1 -1 -1 0 -3 -0 -4 0 -3 1 -3 0 - 3 -0 0 1 2 -0 3 -3 8 -0 1 13 3 4 14 3 0 8 -3 -1 4 - 5 -1 3 -5 7 11 3 3 3 -0 4 -1 2 4 -9 3 6 -16 5 -1 -34 - 5 293.198 296.489 300.519 303.459 308.374 3 13.302 315.839 320.056 324.073 330.275 334.497 339.931 344.142 348.439 353.3 84 3 58.929 362.593 363.049 363.591 368.237 372.427 377.572 382.398 387.568 5.855 6.944 8.505 9.822 12.399 15.524 17.370 20.841 24.655 31.652 37.286 45.699 53.227 61 395 73.237 87.873 98.750 100.1 77 101.896 1 17.566 133.286 154.790 177.330 204.254 3 -0 1 -0 -1 -0 -1 1 -3 1 -3 1 -4 -0 -2 1 -2 -0 -2 -3 4 0 5 -2 2 12 2 -1 10 -3 -2 8 - 5 2 11 0 -1 7 -1 1 8 1 4 11 5 - 4 0 0 0 0 8 1 -6 10 -9 -25 -4 3 -26 -10 trichlorotrifluorobenzene (99.8) 3 64.225 367.951 372.877 377.171 382.099 386.563 391.427 396.136 401.004 405.974 410.718 416.936 422.234 428.262 434.156 440.352 447.167 453.087 459.341 466.122 471.301 472.3 80 478.412 483.851 490.344 2.845 3.360 4.159 4.987 6.104 7.290 8.799 10.498 12.539 14.949 17.604 21.658 25.690 3 1.008 37.048 44.400 -1 53.802 1 63.223 6 74.562 -0 88.653 -1 103.470 - 5 119.561 -6 135.721 -0 100.790 2 157.227 13 4 0 3 0 -2 -2 -2 -0 -3 1 -4 1 -4 2 -7 -1 - 5 1 -7 -2 -2 1 3 2 5 0 9 -1 14 -2 22 2 23 -1 26 2 15 -5 8 -3 4 3 - 4 - 4 -15 1 -21 8 -26 16 496.543 180.149 -2 -61 -18D.AMBROSE, J . H. ELLENDER, C. H. S . SPRAKE, R. TOWNSEND 37 RESULTS Temperatures throughout are expressed as International Practical Kelvin Temp- eratures T6 s , which are treated as interchangeable with thermodynamic temperatures T.The experimental values at pressures below 205 kPa are listed in table 1, those at higher pressures in table 2, and the coefficients of the four equations fitted for each substance, two (I), (11) eqn (l), and two (111), (IV) eqn (2), in table 3. The critical temperatures and pressures have already been reported. T/K p/kPa 104Aloglo(plPa) lY2,3,5-tetrafluorobenzene ebulliometric measurements 385.724 228.28 2 390.810 260.59 1 396.260 299.04 1 402.166 345.46 1 408.935 405.26 1 415.599 471.52 0 1 , 2,4,5- tetra fluoro benzene ebulliometric measurements 392.840 234.88 0 395.913 254.29 0 402.207 297.91 0 408.641 348.26 0 414.978 404.00 0 421.514 468.06 -1 430.210 564.38 -8 437.130 651.79 - 10 dynamic measurements 406.79 333.39 2 413.51 390.65 1 424.26 497.64 2 433.19 602.23 1 447.16 799.04 5 455.93 944.97 5 TIK p/kPa 104Alog~~(p/Pa) 1,2,4,5-tetrafluorobenzene (cont’d) dynamic measurements 464.92 473.82 483.28 492.89 501.94 510.92 520.55 530.70 535.83 538.51 11 14.3 1304.6 1530.3 1792.2 2067.8 2372.1 2736.6 3168.8 3415.5 3544.0 5 6 2 4 3 0 -5 -9 -2 - 6 static measurements 492.10 1770.1 5 496.65 1903.8 4 504.61 2160.0 13 510.81 2369.6 2 516.08 2560.9 -5 526.75 2994.1 -8 53 1.32 3204.1 0 53 1.70 3212.4 - 12 535.07 3374.4 -7 535.22 3385.8 -2 536.29 343 1.7 - 10 539.11 3569.8 - 12 543.35 a 3800.6 a 1 a critical temperature and pressure The values in table 1 were fitted by Antoine equations, both over the full range (11) and also, more exactly, over a restricted range close to the normal boiling point (I).They were also fitted by Chebyshev equations (111),12 (2) (T/K) log,o(P/kPa) = 4 2 + k a,E,(x), s= 1 where Es(x) is the Chebyshev polynomial in x of degree s, and x is defined as [ZT- (T,,, + Tmin)]/(Tmax - Tmin). The values in table 1 and table 2 and the critical temper- atures and critical pressures were also fitted by eqn (2), the values in table 1 being given38 THERMODYNAMIC PROPERTIES OF FLUORINE COMPOUNDS tenfold the weight of those in table 2 (IV). Residuals from eqn (I), (11) and (111) are given in table 1 as Ap = Pobs-pcalc, and from (IV) in table 2 as A log,, p = log,, Pobs - log1, Pcalc, the values of pcalc having been obtained from the appropriate equa- tions.The residuals from (IV) of the values in table l are only marginally larger than those from (111). TABLE 3.-cOEFFICIENTS OF EQUATIONS. (I) SHORT-RANGE ANTOINE EQUATION (1) ; (11) FULL-RANGE ANTOINE EQUATION (l), AND (111) CHEBYSHEV EQUATION (2) FITTED TO ALL VALUES IN TABLE 1 ; (Iv) CHEBYSHEV EQUATION (2) APPLICABLE TO THE CRITICAL TEMPERATURE (1) (11) A - B - C A - B - C 1,2,3,4-tetrafluorobenzene (A) 6.158 54 1291.080 56.617 6.161 07 1292.550 56.453 1,2,3,5-tetrafluorobenzene (B) 6.154 14 1255.781 54.898 6.155 07 1255.981 54.919 1,2,4,5-tetrafluorobenzene (C) 6.173 40 1277.452 56.899 6.177 87 1279.904 56.642 trichlorotrifluorobenzene (D) 6.278 66 1721.246 68.694 6.280 75 1721.899 68.739 - - - - - hexafluoro benzene (El - fluorobenzene (F) - - - - - - A 3 300 392 B 4 287 382 C 3 293 388 D 4 364 497 A 4 300 551 B 5 287 536 C 5 293 544 D 5 364 685 E 5 278 517 F 5 312 561 a0 1167.856 1105.072 1 129.229 1298.238 2246.392 2149.572 2191.684 2599.105 2002.068 2467.161 a1 (111) 324.210 335.195 33 7.09 1 479.339 (IV) 854.391 848.157 858.521 1117.004 823.237 822.894 a2 - 3.256 - 3.662 - 3.676 - 5.71 1 - 8.634 - 8.976 - 9.400 - 13.415 - 10.618 - 5.186 4 0.242 0.283 0.282 0.439 3.721 4.082 4.332 5.236 4.408 3.220 a4 a5 - - 0.238 -0.159 0.025 - 0.341 -0.075 0.075 0.128 - 0.065 0.158 0.071 0.123 Measurements over the whole range up to the critical temperature were made only for 1,2,4,5-tetrafluorobenzene, and where observed values were lacking for the other compounds, estimated values calculated according to the following equation were included (with low weighting) in the sets to which the equations (IV) were fitted, (3) In eqn (3), lOglo(p/kPa)A,t is the value obtained from the full-range Antoine equation (11) for a given value of the reduced temperature T, = T/T, > 0.7 and D is obtained by insertion in eqn (3) of the critical pressure at T' = 1.It has been shown that for non-associated compounds, and in particular for halogenated aromatic hydrocar- b o n ~ , ~ eqn (3) reproduces observed values of vapour pressure in the range 0.7 -= T' < 1.0 within 0.5 %, and this was confirmed in this investigation in respect of 1,2,4,5-tetrafluorobenzene and also, from published 3-1 in respect of hexa- fluorobenzene and fluorobenzene. The equations (IV) detailed in table 3 for the other three compounds included in this study reproduce almost exactly the values given by the corresponding eqn (3).The values of I03D found are: 1,2,3,4- tetrafluorobenzene, 5.65 ; 1,2,3,5-tetrafluorobenzene, 5.87 ; 1,2,4,5-tetrafluoro- benzene, 5.57 ; trichlorotrifluorobenzene, 5.01 ; and for hexafluorobenzene, 5.86 ; log,,(p/kPa) = lOglo(p/kPa)A,t+ D(Tr-0.7) +200D(Tr -0.7)3.D . AMBROSE, J . H . ELLENDER, C. H . S . SPRAKE, R . TOWNSEND 39 and fluorobenzene, 6.09 : they correspond for all the compounds to the critical pres- sure being 7 to 8.5 % above the Antoine value, a usual amount. Properties calculated from the fitted equations are reported in tables 4 and 5. TABLE 4.-vAPOUR PRESSURES AND VALUES OF dp/d T CALCULATED FROM EQUATIONS DETAILED IN TABLE 3 276.91 298.15 367.510&0.001 385.58 429.69 464.75 506.17 533.73 550.83 a 2 6.509 101.325 171.33 500 1000 2000 3000 3791 0.121 0.331 3.116 4.707 10.75 18.21 30.86 42.14 50.55 (111) 274.33 2 (111) 298.15 7.556 (111) 363.413 kO.002 101.325 (111) 380.34 167.43 (IV) 424.50 500 (IV) 459.04 1000 (IV) 499.84 2000 (IV) 526.76 3000 (IV) 543.35a 3800 0.123 0.381 3.172 4.708 10.94 18.47 31.41 43.52 53.74 1,2,3,5-tetrafluorobenzene trichlorotrifluorobenzene 269.39 298.15 357.610+0.002 374.68 418.24 452.45 492.83 519.58 535.25 a 2 9.80 1 101.325 168.69 500 1000 2000 3000 3747 0.125 0.479 3.197 4.771 11.02 18.66 31.71 43.54 52.02 (111) 298.15 (111) 356.63 (111) 471.519-0.004 (111) 479.40 (IV) 549.37 (IV) 592.81 (IV) 643.65 (IV) 677.30 (IV) 684.85 a 0.0651 2 101.325 122.36 500 1000 2000 3000 3270 0.0047 (111) 0.095 (111) 2.474 (111) 2.875 (111) 8.642 (IV) 14.76 (IV) 25.24 (IV) 34.55 (Iv) 36.92 (IV) a Critical temperature ; b Tr = 0.7 ; C normal boiling point with maximum difference that may arise from use of eqn (I), (11) or (IV).TABLE 5.-ACENTRIC FACTOR cr) AND PROPERTIES AT THE NORMAL BOILING POINT: SECOND VIRIAL COEFFICIENT, MOLAR VOLUME OF LIQUID AND ENTHALPY OF EVAPORATION - B/dm3 mol- AH/kJ mol-I VLI 0 a b dm3mol-' a b 1,2,3,4-te trafluorobenzene 1,2,3,5-tetrafluorobenzene 1,2,4,5 - t et rafluoro benzene trichlorotrifluorobenzene hexafluorobenzene fluorobenzene pentafluorotoluene pentafluorophenol 0.345 0.347 0.356 0.427 0.396 0.243 0.416 0.50 1.21 1.19 1.20 1.72 1.29 1.50 1.04 1.12 1.44 1.89 1.32 0.117 33.0 0.117c 32.1 0.117 32.9 0.166 42.9 0.126 31.9 31.68 0.100 31.3 31.20 l3 0.143 35.3 34.74 l8 0.128 40.8 w = loglogc-loglop-l at 7'' = 0.7.16 -B: aestimated by use of w ; bcalculated from ob- served value of AHand eqn (4).A H : a calculated from estimated value of B and eqn (4) ; b observed value. VL : c assumed to be identical with that of 1,2,3,4-tetrafluorobenzene. The enthalpy of evaporation AH was obtained from the equation for which the second virial coefficient B was estimated by the method proposed by Pitzer and Cur1,l6 and the molar volume of the liquid was obtained from measured AH = T[RT/I> + B- VJ dI>/dT, (4)40 THERMODYNAMIC PROPERTIES OF FLUORINE COMPOUNDS densities.l4* 15* l7 Included in table 5 are entries for hexafluorobenzene, fluoro- benzene [for both of which equations (IV) are detailed in table 31, and for pentaflouro- toluene,s together with observed values 2* 3* of AH and values of B calculated from them to indicate the accuracy for this class of compound that may be attributed to the method of estimating these properties.Included also is an entry for pentafluorophenol which supersedes that given for this compound previously. DISCUSSION Equations (1)-(111) are offered as the most convenient (11) or the most accurate (I), (111) expressions for representation of the measured values, and (111) is recommended as the most reliable for extrapolation to lower temperatures. Equations (IV), however, while being convenient and accurate representations of the values, have been chosen so that in addition they conform in their behaviour to a known characteristic of the vapour-pressure line over a long range, as will now be discussed.0.4 0.5 0.6 0.7 0.8 0.9 1.0 TI. FIG. 1.-Plots of AH/AZ; for clarity some of the lines are displaced vertically from their true posi- tions by the amounts stated in parentheses ; A, 1,2,3,4-tetrafluorobenzene (- 5 kJ) ; B, 1,2,3,5- tetrafluorobenzene (- 5 kJ) ; C, lY2,4,5-tetrafluorobenzene ; D, trichlorotrifluorobenzene ; E, hexafluorobenzene (5 kJ) ; F, fluorobenzene ; G, water (12.5 kJ). Waring l9 drew attention to the fact that for water if AH/AZ (A2 = Zv-ZL, the difference in the compression factors of the co-existing vapour and liquid phases) is plotted against T,, a curve is obtained of distinctive shape with a minimum at Trw0.85, and similar curves have been shown for other compounds.20 These curves were plotted from thermal measurements, but the quantity AH/AZ may also be obtained from vapour-pressure measurements by means of the Clapeyron equation written as (5) The curve for AH/AZ may therefore be obtained by differentiation of a vapour- pressure equation, and any such equation that does not give a result qualitatively similar to the usual curve should be suspect.Fig. 1 shows curves for AH/AZ ob- tained from the equations (IV) in table 3 and from an eleventh order Chebyshev equation for water ; 22 they are all of satisfactory shape according to this criterion although the flatness near T, = 1 of curves A, B and D for 1,2,3,4tetrafluorobenzene, 1,2,3,5-tetrafluorobenzene, and trichlorotrifluorobenzene suggests that the estimated AH/AZ = RT2 d lnp/dT.D.AMBROSE, J . H . ELLENDER, c . H . s. SPRAKE, R . TOWNSEND 41 vapour pressures of these three substances may be slightly in error. Curves obtained in the same way from the equations already published for other halogenated aromatic compounds are of a shape similar to those in fig. 1. When the quantity A = Tlog@/atm)/(T-T,,), where Tb is the normal boiling point, is plotted against T, a curve of more pronounced parabolic shape than those in fig. 1 is obtained, again with a minimum at Trz0.85. This A is a variable in the Cox vapour-pressure equation, and Thomson called the characteristic shape of its curve the " Cox criterion " that must be satisfied by reliable data or a valid equation 21 ; the " Waring criterion '' used here seems to be preferable since it is related to the defined thermodynamic quantities AH and AZ. We acknowledge the assistance in the experimental work of Miss R.F. Anthony and Mr. S. W. David, and computational help given by Mr. E. B. Lees. Part 14, R. J. L. Andon and J. F. Martin, J.C.S. Faraday I, 1974,70, 605. J. F. Counsell, J. H. S. Green, J. L. Hales and J. F. Martin, Trans. Faraday SOC., 1965,61,212. D. Ambrose, J. Chem. SOC. A, 1968,1381. D. Ambrose and C. H. S. Sprake, J. Chem. SOC. A, 1971, 1263. T. V. Findlay, J. Chem. Eng. Data, 1969, 14,229. R. J. L. Andon and J. F. Martin, J.C.S. Faraday I, 1973,69, 761. R. J. L. Andon and J. F. Martin, J.C.S. Faraday I, 1973, 69, 871. D. Ambrose and C. H. S. Sprake, J. Chem. SOC. A, 1971, 1261. D. Ambrose, B. E. Broderick and R. Townsend, J. Appf. Chem. Biotechnol., 1974, 24, 359, lo D. Ambrose, C. H. S. Sprake and R. Townsend, J. Chem. Thermodynamics, 1969,1,499. l1 D. Ambrose, B. E. Broderick and R. Townsend, J. Chem. SOC. A, 1967, 633. l2 D. Ambrose, J. F. Counsell and A. J. Davenport, J. Chem. Thermodynamics, 1970,2,283. l 3 D. W. Scott, J. P. McCullough, W. D. Good, J. F. Messerly, R. E. Pennington, T. C. Kincheloe, I. A. Hossenlopp, D. R. Douslin and G. Waddington, J. Amer. Chem. SOC., 1956,784,5457. l4 D. R. Douslin, R. T. Moore, J. P. Dawson and G. Waddington, J. Amer. Chem. SOC., 1958,840, 2031. l5 D. R. Douslin, R. H. Harrison and R. T. Moore, J. Chem. Thermodynamics, 1969,1, 305. l 6 K. S. Pitzer and R. F. Curl, J. Amer. Chem. SOC., 1957, 79, 2369. l7 J. L. Hales and R. Townsend, J. Chem. Thermodynamics, 1974, 6, 11 1. l9 W. Waring, Ind. Eng. Chem., 1954, 46, 762. 2o R. C. Reid and T. K. Sherwood, The Properties of Gases and Liquids, Their Estimation and 21 G. W. Thomson, Physical Methods of Organic Chemistry, ed. A. Weissberger (Interscience, 22 D. Ambrose and I. J. Lawrenson, J. Chem. Thermodynamics, 1972, 4, 755. J. F. Counsell, J. L. Hales, E. B. Lees and J. F. Martin, J. Chem. SOC. A, 1968, 2994. Correlation (McGraw-Hill, New York, 2nd edn. 1966), p. 116. New York, 3rd edn., 1959), vol. 1, part 1, chap. VIII, p. 357.

ISSN:0300-9599    

DOI:10.1039/F19757100035

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Thermal Diffusion and Convective Stability Experimental study of the Carbon Tetrachloride + Chlorobenzene System BY ANTONIO SPARASCI AND H. J. VALENTINE TYRRELL* Department of Chemistry, Chelsea College, Manresa Road, Chelsea, London SW3 6LX Received 17th April, 1974 The thermal diffusion behaviour of carbon tetrachloride + chlorobenzene mixtures has been studied (at a mean temperature of 25°C) in horizontal liquid films confined between two rigid, thermally conducting, surfaces 0.923 mm apart. The heavier component (CCI,) migrated to the cold wall for all mixtures studied. Soret coefficients, thermal diffusion factors and heats of transport have been calculated from experiments where the upper plate was heated with respect to the lower one and also from experiments where the direction of the thermal gradient was reversed.Provided that the gradient in the second group of experiments was below a critical limit the results from the two sets of experiments were in good agreement. When the lower surface was heated with respect to the upper one, there was a critical limit to the applied temperature interval above which the apparent Soret coefficient decreased. The observed critical limits agree quite well with the predictions of a recent linear stability analysis. The type of convection which sets in above this limit has not hitherto been directly demonstrated because it does not contribute to the total heat flux across the liquid layer. If a horizontal fluid layer, thickness d, is heated from below, the fluid shows no convective motion as long as the temperature gradient across it (p ; taken as positive when the layer is heated from below) is less than a certain minimum va1ue.l This minimum depends upon the thickness of the layer, the experimental conditions employed, and the physical properties of the fluid.For a single component system, the critical limit is usually related to a dimensionless number R (Rayleigh or Rayleigh- BCnard number) defined as : R - gpapd4/Ky. (1) In eqn (1) g is the gravitational acceleration, p the density, a the volume coefficient of expansion, K the thermal diffusivity and y the dynamic viscosity. When /3 is large enough for R to exceed 1708 (assuming the layer to be conhed between two rigid, plane, horizontal surfaces) regular patterns of convective motion can be seen in the fluid (BCnard cellules).l The onset of this convective motion can either be detected optically or, more commonly, by studying the heat transfer rate across the fluid.When the critical value of R is exceeded this rate shows a sharp increase, that is, the apparent thermal conductivity increases suddenly at the critical value of p. of the components under the influence of the temperature gradient (Soret phenomenon). The density gradient associated with the temperature gradient will be modified by the establishment of concentration gradients, and the question of convective stability in such layers becomes more complicated. The problem has been considered by several Schechter, Prigogine and Hamm pointed out that the thermal diffusivity of a liquid is very much greater than the diffusion coefficient D(K/Dw lo2).Conse- quently, thermal equilibrium is attained much more rapidly than concentration 42 In liquid systems of more than one component there may be partial separationA. SPARASCI A N D H . J . V . TYRRELL 43 equilibrium, and as a result, convective flow may occur even under conditions where the overall density gradient would be such as to stabilise the layer in the gravitational field. Such flows were earlier suggested as a possible reason for certain discrepancies in results of experiments on thermal diffusion in aqueous electrolytes. A linear stability analysis for the case where KID is large (2 100) has led to some interesting conclusions. If the heavier component (subscript 1) migrates to the cold wall, the Soret coefficient (a) is defined to be positive : where Xi, Ni are respectively the molar and weight fractions of component i and the suffix " stat " indicates that dX,/dT, dNl/dTare measured in the stationary state.A dimensionless separation parameter S with the same sign as Q is defined as : where y = (a In p/dN,), and a is the volume coefficient of expansion, - (8 In p/aT),. Since : then, When the heavier component migrates to the cold wall ( S positive) the density gradient due to the temperature gradient is reinforced. The sign of this gradient is however changed when the heavier component migrates to the hot wall provided that the separation is large enough ( S < -1). The fluid layer then becomes gravitationally unstable. Experiments on unequally heated fluid layers can be classified into four groups as follows.(i) p negative, S positive. The density gradient always stabilises the layer against convective flow (if wall effects are neglected). (ii) p negative, S negative. From eqn (5) it is possible for the density gradient to become inverted, although even when this does not happen, bulk flow may occur because K % D. This seems to be the reason for some of the anomalous results reported for thermal diffusion experiments on systems where the heavier compon- ent migrates to the hot wall.9 (iii) /3 positive, S positive. In respect of the type of convective flow considered here the system is essentially the same as the preceding one provided that fl is relatively small. (iv) p positive, S negative.In this instance the temperature gradient gives rise to an inverted density gradient which is opposed by the migration of the heavier compon- ent towards the lower (hotter) surface. When KID is large, as in liquid systems, oscillatory instabilities (" overstable " states) have been predicted independently by Hurle and Jakeman and by Velarde and Schechter.6 Finite amplitude temperature oscillations have been shown 4* 21 to occur at the centre of a horizontal cell, heated from below, when filled with a water + methanol mixture for which S was negative.44 THERMAL DIFFUSION AND CONVECTIVE STABILITY The instabilities were observed for the experimental conditions predicted by the theory. At low temperatures (< 10°C), an 0.5 mol dm-3 aqueous sodium chloride solution has a negative Soret coefficient.Caldwell 22 has examined the heat flux across a horizontal layer of this solution as a function of the temperature interval (with #l positive). The apparent thermal conductivities showed a sudden increase at Rayleigh numbers above the theoretical value for a single component system (1708). From these observations, and the theoretical treatment of Hurle and Jakema~~,~ Soret coefficients were calculated 22 which agreed well with those obtained by more con- ventional methods. There was, additionally, some evidence of finite amplitude instabilities near the critical point. Experiments with positive values of p are extensions of the Rayleigh-BCnard type of experiment to two component systems. If S is positive, theory 6* predicts that convective instability should occur for smaller values of /3 than would be required to induce instability in a one-component fluid layer having the same dimensions and physical properties.This is not solely due to the fact that the migration of the heavier component to the upper (colder) surface of the cell reinforces the inverted density gradient associated with a positive value of B, because an identical limit is predicted for the conditions described under (ii) above. In these conditions, the stabilising density gradient associated with the temperature gradient is opposed by a de-stabilising gradient associated with the migration of the heavier component to the heated upper surface. Even if the overall density gradient is such as, apparently to stabilise the system in the gravitational field, convective instability can occur because of the fact that K $ D.Consider, for example, a small volume element in the two component liquid film for which both p and S are negative and the density is greatest at the bottom of the cell. If a fluctuation occurs and the volume element rises slightly, it should, at first sight, sink back to its original position in a short time. Though less rich in the heavier component, it is colder, and therefore heavier, than the fluid surrounding it in the new position. However, since K 9 D, the volume element may reach thermal equilibrium with these new surroundings before its composition changes substantially, and it will then be lighter than its surroundings. Consequently, it will tend to rise further, the original fluctuation being thereby reinforced rather than damped out.In this way, slow convection currents may be set up even when the overall density gradient seems adequate to prevent them. A similar phenomenon is found in oceanology as the '' salt fountain effect ", c.f. ref. (7). The Velarde-Schechter theory predicts that convective instability should occur when p and S have the same sign if a thermal Rayleigh number a reaches a critical value of 720. a is defined as (RSKID), which, from the definitions (1) and (3), is equivalent to : = R*aNlN2#l where R* is defined as the dimensionless quantity the temperature gradient (Pcrit) above which this served is therefore given by (7) (gpyd4/Dy). The critical value of type of convection should be ob- When this critical temperature gradient is reached the wavelength of the convective cell will be large.Theoretically it should be infinite but R varies little with wave- number at small wavenumbers,6 and there may be a number of long wavelength modesA. SPARASCI AND H. Y. V. TYRRELL 45 operative together. If so, definite tessellated patterns similar to those observed for one component systems at the critical point would not be seen. The velocity of such non-regular convective motion is very small.6 The onset of convection in binary systems with both p and S positive has hitherto been examined by the technique of measuring the heat flux across the fluid layer as a function of p. Suitable systems are carbon tetrachloride +benzene and carbon tetrachloride + chlorobenzene. In both cases the heavier component (carbon tetrachloride) migrated to the cold wall in a thermal gradient O.contrary to an earlier report.12 However, no change in the apparent thermal conductivity occurred in a cell heated from below until the normal Btnard limit was reached.3* ' 9 l3 Jt was originally suggested that this was because no thermal diffusion separation can occur in a two component liquid layer heated in this l 1 but such a separation was later observed for the toluene+ethanol system l4 in a pure Soret effect cell. The true explanation seems to be 6* '* l5 that the convective instability described by eqn (6), (7) and (8) involves a very slow bulk motion of the liquid. Because of the large value of KID for liquids, this motion does not contribute to the heat flux across the cell, and its onset cannot therefore be detected in thermal conduction experiments.Beltoa and Tyrell l4 noted in their experiments that the apparent Soret coefficient observed in an inverse temperature gradient (in the toluene +ethanol system with S positive) was reduced when p became large. A systematic study of the apparent Soret coefficient of a binary system with S positive, using a range of positive and negative values of p has therefore been carried out to search for the existence of a critical limit below the normal Bknard limit, and to test the theoretical conclusions of Velarde and Schechter summarised in eqn (6)-(8). In particular, eqn (8) shows that Pcrit should be dependent upon composition and, if both 6 and R* vary little with composition, it should be a minimum when N1 = N2 = 0.5.The system chosen for the study was carbon tetrachloride + chlorobenzene, partly because there had been some dispute about the sign of the separation in a thermal gradient, and partly because this system had been used in some of the heat flux experiments carried out by Legros, van Hook and Thomaes. EXPERIMENTAL The design of the Soret cell and the experimental techniques have been described ear- lier.lou9 l6 In the present work the plate separation in the reference cell was 0.918 mm and in the test cell 0.923 mm. The method of controlling the temperature interval across the cell used earlier was equally suitable for positive or negative values of p. Because the tempera- ture gradient was to be applied in both directions, and both components have relatively large refractive index increments with temperature, the optical system used limited the useful temperature interval which could be used across the cell to about +O.5O0C, for normal Soret effect studies ( p negative) temperature intervals of up to 0.8"C were used earlier ; in principle, the necessity of using a smaller temperature interval reduces the accuracy with which the Soret coefficient can be measured.Refractive index increments with temperature for the NaD line at constant composition were measured at a mean temperature of 25°C using a Rayleigh refractometer and the technique described earlier. Oa Unpublished data on the refractive index of these mixtures as a function of composition for the mercury green line (2 = 546.1 nm) had been found l7 to fit the curve (valid at 23.5"C) : n:z& = A+ BX2 + kXq(1- X,)"'.(9) In this equation, A was the refractive index of pure carbon tetrachloride, B the difference between the refractive indices of pure chlorobenzene and pure carbon tetrachloride, and Xz the mole fraction of chlorobenzene. The third term on the right contains three empirical46 THERMAL DIFFUSION AND CONVECTIVE STABILITY constants (k,p, rn) determined by a least squares fitting procedures as : k = 2.116~ rn = 0.7698, p = 0.5891. This third term represents the deviation of n&!l from linearity in mole fraction, and it is here assumed that the parameters k, rn, p can be used unchanged to describe the curvature of the plot of nk5 with composition provided that appropriate values of A and B were used.Using a precision Pulfrich refractometer on our purified samples the following results were obtained : carbon tetrachloride : nk5 = 1.457 56 c hlor o benzene : nk5 = 1.521 86, hence Bi5 = 0.064 30. Differentiation of eqn (9) gives : (a~~A~/ax~)~ = B ; ~ + kxq( 1 - x2)"(p/x2 - mix,). (10) This equation was tested by comparing the value of (dn&5/dX2)T predicted from (10) for X2 = 0.2862 with a directly measured value using the techniques described earlier.loa* 16, l8 The calculated value was 7.19 x The agreement is good, and no further direct measurements of (anis/dX2), were undertaken. Carbon tetrachloride (AnalaR grade) was purified by initially refluxing over mercury for two hours. This was followed by washing with concentrated sulphuric acid to remove traces of sulphides, 5 % sodium hydroxide solution and water. After drying over fused calcium and the experimental value 7.08 x 8.4- 8 . 2 - 8.0 7.8 7.6 7.4 7.2 7.0- 6.8 El B E l El - - - - - I 5.2 4.8 4.4 4.0- 3.6 3.2 2.8 2.4 2.0 Xi = 0.1982 - r;] El - I3 - - - - - !3 I I Xi = 0.51 6.A 6.2 5.6- 5.2 4.8 4.4 4.0 3.6 Xi = 0.8059 El u m w n % I3 - - - - , FIG.1 .-The onset of convection in thin (0.923 mm) liquid films of carbon tetrachloride+ chloro- benzene solutions heated from below as shown by the change in apparent Soret coefficient as the temperature interval across the cell is increased.A. SPARASCI AND H. J . V . TYRRELL 47 chloride the sample was fractionally distilled through a 50 cm column packed with alternate 3 cm bands of multiple Fenske and single turn helices.A reflux ratio of 4 : 1 was employed, the middle fraction (66 %) being retained. This fraction was distilled twice more, the middle fraction being retained at each stage. The final product was collected in a specially designed vessel which enabled storage and dispensing to be achieved without contamination by atmospheric water vapour (ni5 = 1.457 56 ; lit., ni5 = 1.457 59). Chlorobenzene was shaken with portions of concentrated sulphuric acid until the latter was substantially colourless. This was followed by washing with water, 5 % potassium bicarbonate solution and water. After drying over fused calcium chloride the chlorobenzene was distilled in the column mentioned above. The middle fraction (66 %) was dried over phosphorus pentoxide for twenty four hours and distilled twice more, the middle fraction being retained at each stage.The refractive index of the final product was measured on a Pulfrich refractometer (nio = 1.524 61 ; lit., ni0 = 1.524 59). RESULTS For both positive and negative values of p the heavier component, carbon tetra- chloride, was found to migrate to the cold wall as reported by Legros, van Hook, and Thomaes l1 from flow-cell data i.e. S is positive for this system. Hence, for positive values of fl, there should be a critical value at which the observed Soret coefficient TABLE 1 .-THERMAL DIFFUSION FACTORS (298.1 0) FOR CARBON TETRACHLORIDE+ CHLORO- BENZENE MIXTURES IN HORIZONTAL LIQUD FILMS HEATED (a) FROM ABOVE ( p NEGATIVE), (6) FROM BELOW (p POSITIVE) WITH AN APPLIED TEMPERATURE GRADIENT LESS THAN THE CRITICAL VALUE FOR CONVECTION (MEAN TEMPERATURE 25°C) molar fraction carbon tetrachloride X I O.oo00 0.0989 0.1000 0.1982 0.2000 0.3000 0.3180 0.3730 0.3917 0.4OoO 0.5000 0.5024 0.5150 0.5896 0.5957 0.6000 0.6050 0.6977 0.7000 0.7172 0.7220 0.8Ooo 0.8059 0.8173 0.8921 0.8955 0.9000 1 .m 104( - a n & 5 / a ~ ) / K-1 5.73 5.69 5.65 5.61 - - - - - 5.57 5.54 - - - - 5.50 - - 5.48 - - 5.46 - - - - 5.44 5.43 thermal diffusion factor (02") negative values of B - 2.88( 0.1 9) 2.25(+ 0.09) - - - 1.69( + 0.04) 1.50(+ 0.10) 1.51(+ 0.1 1) 1.46( + 0.06) 1.56(+0.09) 1.47( + 0.05) 1.55(+ 0.07) 1.60( + 0.06) - - - - - - 1.58( & 0.05) - - 1.64(+0.09) 2.14(+ 0.07) 2.15(&0.07) - - positive values of B - 2.39(+0.08) - - .69(+0.08) .47( + 0.05) .42( + 0.12) .61( & 0.1 1) .53( & 0.05) - - - - - - - - 1.70( & 0.05) - - 1.78(+0.11) -48 THERMAL DIFFUSION AND CONVECTIVE STABILITY should begin to fall because of the onset of convection.According to the Velarde- Schechter theory [eqn (6)-(8)] this critical point should be composition-dependent. Fig. (1) shows results obtained in this work at three compositions. In each case, there is a fairly clearly defined critical temperature interval above which the measured Soret coefficient decreases, and, as predicted, this is less for Xl = 0.515 than for X , = 0.198 or X , = 0.806. Soret coefficients (0) have been calculated both for positive values of /3 less than Pcrit, and for negative values of p where Q was essentially independent of p as expected.Results are shown in table 1, along with the refractive index increment values used in the calculations. There was no significant difference between thermal diffusion factors (or Soret coefficients) measured with negative values of /3 and those with positive values of p, provided Pcrit was not exceeded. The precision with which Q could be measured was essentially the same for the two groups of measurements and similar to that achieved in other work using a similar technique. DISCUSSION The Soret coefficients found for this system are comparable in magnitude with those K-I for an Since the Soret coefficient is related to the heat of transport found for the system carbon tetrachloride+benzene lo (Q = 6.4 x equimolar mixture).by the relationship : (1 1) Q,* RT2x,(1 +d In f2/d In XZ)T,P changes in 0 with concentration may be due to changes in Q*/X, or to the thermo- dynamic term, or both. Measurements of the total vapour pressure over chloro- benzene +carbon tetrachloride mixtures have been obtained by Bittrich and his co-workers 23 at 40°C but they have not tabulated activity coefficients (f) for the individual components. Although the total pressure data are given only at mole fraction intervals of 0.1, we have calculated partial pressures by the Boissonnas method,24 and estimated activity coefficients from these without correcting for vapour phase non-ideality. Values of (1 + a Inf/a In X)j- were obtained from these using a simple finite-difference method and combined with Soret coefficients inter- polated from a smoothed curve, constructed from data in table 1, to give values of Q$/Xl (table 2).Although the activity coefficient terms are inevitably not very accurate, and strictly apply only at 40°C, it seems certain that the observed changes in CT with concentration cannot be explained in terms of a variation of (1 +a lnf/ d In X)T,p with composition. Fig. 1 illustrates the sharpness of the transition at the critical temperature interval at which the apparent Soret coefficient began to fall off in the experiments where /? was positive (fluid layer heated from below). The three plots are representative and clearly there is little doubt about the value of the limiting temperature interval for Xl = 0.8059. Iii all cases, the horizontal line can be regarded as well-established being derived from a smoothed plot of Soret coefficient against composition based on all the data in table 1 .The individual experimental points were obtained from several independent experiments and show the kind of scatter expected from experimental errors of measurement, which are particularly severe for small values of AT and sol- utions dilute in either component. In most cases (cf. data illustrated for Xl = 0.1982, 0.8059), the fall in the apparent Soret coefficient with increasing AT was apparently linear, and the critical temperature interval could be determined to a few 1- Later confirmed by communication from Professor Bittrich of his own activity data. Q =A . SPARASCI A N D H. J . V .TYRRELL 44 tenths of a degree with some certainty. In the centre of the composition range (cf. plot shown for X , = 0.5150) the fall was not linear in AT, and extrapolation to the critical limit was more difficult. In these cases we took the critical limit as being the highest value of AT at which the Soret coefficient, as measured with p positive, coincided closely with the value based on data in table 1. TABLE 2.-HEATS OF TRANSPORT (@/xi), THERMODYNAMIC FACTORS, AND SORET COEFFICIENTS (0) DERIVED BY INTERPOLATION ON A SMOOTHED CURVE molar fraction carbon tetrachloride, XI 0.85 0.75 0.65 0.55 0.45 0.35 0.25 0.15 (l+8 InfP In WT.P 1.10 1.10 1.10 1.09 1.06 1.08 1.02 1.06 103x alK-1 tQt/Xl)/kJrnol-' 6.60 5.60 5.13 5.00 5.00 5.38 6.50 8.58 5.4 4.6 4 . 2 4.0 3.9 4.3 4.9 6.7 However, in order to calculate K, other data are needed.Eqn (6) can be re-written : Rcrit = (gd 3 / ~ ~ ) ( a ~ I a ~ 1) 1 X2ATc (12) The variation of density with composition for this system was studied by Das and Roy l9 at 10°C intervals between 10 and 60°C. In each case we have found that the density could be represented as a linear function of mole fraction of carbon tetra- chloride (X,), the correlation coefficients being greater than 0.9999 in each case. Thus ap/a XI was effectively independent of concentration, the following values (g ~ m - ~ ) being obtained from the least squares analysis at each temperature : 0.496 24 (10°C) ; 0.487 72 (20°C) ; 0.479 42 (30°C) ; 0.470 31 (40°C) ; 0.462 33 (50°C) ; 0.452 26 (60°C). These values vary linearly with temperature to a good approximation and a least squares analysis gave (correlation coefficient 0.9996) : At 25°C the best estimate of dp/i?X, was therefore 0.483 43 g ~ r n - ~ , and this was used in the subsequent calculations of 8.The diffusion coefficient presented more serious difficulties since no experimental diffusion work seems to have been undertaken on these mixtures. Self diffusion coefficients for the pure components Dr were estimated by assuming that : 2o (i?p/i?X1)T = 0.505 23-8.719 x (T-273.15) g CM-~. (13) D*q = kT/4nrW. (14) For eqn (14) to apply, a suitable method of calculating the van der Waals' radius Y, is required. Using the technique described by Edward,20 rw was found to be 0.283 nm for chlorobenzene and 0.275 nm for carbon tetrachloride.This made it possible to estimate DT for carbon tetrachloride and DZ for chlorobenzene with a reasonable degree of accuracy. The simplest possible relationship between the mutual diffusion coefficient D and the self diffusion coefficients is [ref. (2) p. 1341 : In the absence of more complete data the estimates of DT obtained from eqn (14) were D = X2D;+XlD;. (15)50 THERMAL DIFFUSION A N D CONVECTIVE STABILITY combined with eqn (15) to give the following estimate of Dy and its variation with composition : Hence Rcrit may be calculated at a series of compositions. The simple Rayleigh- Benard theory [eqn (l)] can be modified to take account of the increased de-stabilisa- tion due to the concentration gradient arising from the establishment of the Soret equilibrium, by replacing a by a( 1 + S ) in eqn (l), cf. eqn (5).The critical temperature interval in this case would then be given by : Dy = (1.1998 - 0.424XJ x lo-’ dyn. (16) 1708Ky pgd4a( 1 + s>. The definition of S in eqn (3) is equivalent to : Hence, in addition to the pbysical constants already discussed, it is necessary to know K, and (ap/dT),. To obtain K, thermal conductivity and heat capacity data are required. Frontas’ev and Gusakov have measured thermal conductivities for the pure components, and Mukhamedzyanov et aZ.26 have pointed out that, for nearly- ideal mixtures, the thermal conductivities of mixtures vary almost linearly with mass fraction, though not with mole fraction. Hence, values of the thermal conductivity of each mixture studied was estimated in this way.Excess heat capacities have been obtained for this system, together with heat capacities of the pure and reliable viscosity data are also available.27 Densities and values of (ap/aT), can be obtained from the data of Das and Roy l9 already referred to. s = - aX,X2(aP/aX,)*/(ap/aT)x. (1 8) TABLE 3 .-OBSERVED AND CALCULATED TEMPERATURE INTERVALS AND CRITICAL VALUES OF THE -k CHLOROBENZENE SOLUTIONS HEATED FROM BELOW (TWO RIGID BOUNDARIES 0.923 lTllll APART, THERMAL RAYLEIGH NUMBER (&.Tit) FOR HORIZONTAL LIQUID FILMS OF CARBON TETRACHLORIDE MEAN TEMPERATURE 25°C) critical temperature intervalslOC calculated molar fraction mpdified carbon Velarde-Schechter Rayleigh-B&nard tetrachloride, X I Rcrit. observed (eqn 12) (eqn 17) 0.8059 0.7172 0.5150 0.5000 0.3917 0.3730 0.3180 0.1982 1100 1 200 1000 900 1100 1100 1200 1200 0.35 0.24 0.34 0.21 0.26 0.18 0.24 0.18 0.27 0.18 0.28 0.19 0.32 0.19 0.32 0.19 6.8 6.8 6.9 7.0 7.3 7.3 7.5 8.8 Temperature intervals were obtained experimentally from the observed optical deflection produced by the reference cell subject to appropriate corrections, see ref.(1Oa). The results of these calculations are shown in table 3. Obviously the onset of convective re-mixing occurs at much smaller temperature intervals than would be predicted from the modified Rayleigh-BCnard eqn (17). The observed value of Rcrit (mean value 1100) is 53 % higher than the Velarde-Schechter prediction of 720, though it does remain, as predicted, substantially independent of concentra tion. The uncertainty in the experimental value of ATcrit is not large enough to explain this apparent discrepancy, but, in deriving acrit from the experimental data, a number ofA .SPARASCI A N D H. J . V. TYRRELL 51 assumptions have had to be made, the effects of which are impossible to judge. In addition, the theory assumes that K/D= lo2, though for the mixture considered here, this ratio appears to be about 50. In any event, this apparent discrepancy between theory and practice is relatively unimportant in comparison with the unequivocal evidence we have now obtained for the existence of a critical limit for the onset of very slow convective motion in a two- component system for which both p and S are positive. This limit is far below that associated with the normal Rayleigh-B6nard motion at which increased heat flow has been observed for this ~ystern.~ The convective velocity in our experiments must be low enough for no excess heat flow to be associated with it, as predicted by theory, and it can only be observed, as here, by studying changes in the concentration distri- bution across the fluid layer.This kind of convective motion is almost certainly responsible for many of the discrepancies in the published literature on Soret coeffi- cients since it may start at the vertical boundary walls of the Soret cell, as was sug- gested originally by Agar and Turner.* We are indebted to Dr. M. G. Velarde, Professor J. Thomaes and Dr. J-C. Legros One of us (A. S.) is indebted for useful discussions and access to unpublished data.to the S.R.C. for the award of a studentship. cf. S. Chandrasekhar, Hydrodynamic and Hydromagnetic Stability (Oxford, 1961). cf. H. J. V. Tyrrell, Diflusion and Heat Flow in Liquids (Butterworth, London, 1961). J-C. Legros, W. A. van Hook and G. Thomaes, Chem. Phys. Letters, 1968,1, 696. D. T. J. Hurle and E. Jakeman, J. Fluid Mech., 1971, 47, 667. R. S. Schechter, I. Prigogine and J. R. Hamm, Phys. Fluids, 1972, 15, 379. M. G. Velarde and R. S. Schechter, Phys. Fluids, 1972, 15, 1707. J-C. Legros, J. K. Platten and P. G. Poty, Phys. Fluids, 1972, 15, 1383. J. N. Agar and J. C. R. Turner, Proc. Roy. SOC. A, 1960,255, 307. M. G. Velarde and R. S. Schechter, Chem. Phys. Letters, 1971, 12, 312. Turner, Trans. Faraday SOC., 1969, 65, 1523. lo cf. (a) L. Guczi and H. J. V. Tyrrell, J. Chem. SOC., 1965, 6576 ; (b) M. J. Story and J. C. R. l 1 J-C. Legros, W. A. van Hook and G. Thomaes, Chem. Phys. Letters, 1968,2,251. l 2 G. Thomaes, J. Chim. phys., 1956, 50,407. l3 J-C. Legros, D. Rasse and G. Thomaes, Chem. Phys. Letters, 1970,4, 632. l4 P. S. Belton and H. J. V. Tyrrell, Chem. Phys. Letters, 1970, 4, 619. l 5 M. G. Velarde and R. S. Schechter, Chem. Phys. Letters, 1971, 12, 312. l6 G. Farsang and H. J. V. Tyrrell, J. Chem. SOC. A , 1969, 1839. J. Demichowicz-Pigonawa and H. J. V. Tyrrell, Roczniki Chem., 1969,43,433. l 9 L. M. Das and S. C. Roy, Indiaiz J. Phys., 1930,5,441. quoted in J. Timmermanns, The Physico- chemical Constants of Binary Mixtures in Concentrated Solutions (Interscience, New York, 2o J. T. Edward, J. Chem. Educ., 1970,47, 261. 21 J. K. Platten and G. Chavepeyer, J. Fluid Mech., 1973, 60, 305. 22 D. R. Caldwell, J. Phys. Chem., 1973, 77, 2004. ’ J-C. Legros, personal communication. 1959-60), V O ~ . 1, p. 321. 23 R. Kind, G. Kahnt, D. Schmidt, J. Schumann and H-J. Bittrich, Z. phys. Chem. (Leipzig), 1968, 230, 277. 24 C. G. Boissonnas, Helu. Chim. Acta, 1939, 22, 341. 25 V. P. Frontas’ev and M. Y. Gusakov, Uch. Zap. Saratovsk. Gos. Uniu., 1960, No. 69, 237 ; see 26 G. K. Mukhamedzyanov, A. G. Usmanov and A. A. Tarzanimov, Izvest. V. U. Z., Neft i Gaz, 27 R. J. Fort and W. R. Moore, Trans. Faraday Soc., 1966, 62, 1112. Chem. Abs., l963,58,1923b. 1964, 7(10), 70; see Chem. Abs., 1965, 62, 7142f.

ISSN:0300-9599    

DOI:10.1039/F19757100042

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Photoinitiation of Polymerization by Chloro-oxobis(2,4- pentanedionato)vanadiurn(v) in the Presence of Electron Donors BY S . M. ALIWI AND CLEMENT H. BAMFORD* Department of Inorganic, Physical and Industrial Chemistry, Donnan Laboratories, University of Liverpool L69 3BX Received 17th April, 1974 When a strong electron donor D, such as dimethyl sulphoxide (DMSO) or one of a wide range of amino-compounds, is added to VO(a~ac)~Cl a colour change from deep blue to pale green-yellow occurs. The change appears to be complete if [D] is slightly in excess of [VO(a~ac)~Cl] and all strong donors examined give effectively similar final absorption spectra. At the same time the solution acquires marked electrical conductivity. It is proposed that ion-pair complexes of the type (I) / o\+ (0 are formed between donor and chelate The complexes are effective photoinitiators of frwradical polymerization ; the reactions occurring at A = 365 nm with D = DMSO and 3) = pyridine (Py) have been studied with methyl methacrylate as rnoa~mer.The polymerization is not complicated by retardation and the rate of initiation at a volume fraction of additive ua = 0.1 is independent of monomer concentration (with benzene and ethyl acetate as diluents). The quantum yields of initiation are relatively high, 0.59 and 0.125 for D = DMSO and Py, respectively ; these values are 29 and 6 times those obtained with VO(acac),Cl in the absence of D. When D = DMSO both chlorine and DMSO residues are found in the poly- mers, but with D = Py only chlorine has be& detected.Negligible quantities of acetylacetone fragments are incorporated in either case. The products are VIV derivatives, VO(acac)zD with D = Py and VO(acac),D and VOC12D3 with D = DMSO. In the latter case acetylacetone is also formed ; both this product and VOC12DJ are considered to arise from interaction of the chelate with HCl produced in another secondary reaction. Spectrophotometric measurements show that the rate of decomposition of (I) for a given incident intensity is independent of 0,; on the other hand, the rate of initiation decreases markedly with increasing ua for Va > 0.1, approximately. Several reaction mechanisms are discussed and it is concluded that the observations are consistent with a primary act consisting of electron transfer from D to vanadium.Secondary processes in- volving the resulting radical cation D+ and C1- occur leading to C1 atoms and also, when D = DMSO, to cH2SOCH3 and HCI; both el and CH2SOCHJ initiate polymerization. With increasing ua, solvent separation of Df- and C1- retards reaction between these species and D t is consumed in a competitive processes ; the decrease in the rate of initiation with increasing Ua is thus understandable. The relevant kinetic parameters for this mechanism are evaluated. During the course of this work it was found that pyridine has a greater influence on the propaga- tion coefficient of methyl methacrylate than any other solvent so far examined. The presence of electron donors often has a marked influence on the thermal initiation of free-radical polymerization by transition-metal chelates, evidenced by a greatly enhanced rate of initiation and the appearance of monomer-selectivity.v 2 Information about the photoinitiation by chelates is relatively scanty, but unexpected 52S, M. ALlWI AND C. H. BAMFORD 53 solvent effects have been found in polymerizations photoinitiated by MnIII(fa~ac)~ (facac = CF3COCHCOCH3). In this paper we report an extension of our work on photoinitiation by VO(a~ac)~Cl [(acac) = CH3COCHCOCH3] to systems containing the electron-donors dimethyl sulphoxide (DMSO) or pyridine (Py). These donors, and many others, convert VO(acac),Cl into ion-pair complexes [D-+VO(aca~)~l+Cl- (D = electron donor), so that the nature of the photosensitizing species is fundamentally changed by the additive.The systems with which we are concerned may therefore be expected to show analogies with the ion-pair complexes Fe"'X- (X- = OH- or C1-) first studied by Evans, Santappa and UriS These workers demonstrated that on irradiation of FelILX- (A = 310, 365, 405 nm) free radicals or atoms are generated which can initiate the free-radical polymerization of vinyl monomers such as acrylonitrile, methyl methacrylate or methacrylic acid in aqueous solution at 25"C, e.g. Fe1ILCl- + hv+ Fe" + c l I i I I e l + CH2=C+C1-CH2-C. etc. (1) Initiation occurs with low quantum yields ( 5 x lo-' and 0.13 for X- = OH- and Cl-, respectively) and the rate of initiation was shown to be independent of monomer concentration and equal to the rate of consumption of ferric ion. This type of photo-chemical activity is quite general and many examples of it are now known.6 We shall see that photoinitiation by [D+VO(acac),]+Cl- is more complex, exhibiting features which are not implicit in (1).EXPERIMENTAL The techniques were as described in a previous paper with the addition of the following. Gas-liquid chromatography was carried out with a Pye-Unicam gas-liquid chromato- graph (model 24) employing dual flame-ionization detectors, a column (2m) containing 5 % silicone on Embacel (60-100 mesh) at 82°C was used. The participation of acetylacetonate radicals in initiation was examined by comparing the radioactivities of polymers prepared from VO(gcac)2C1 having ligands labelled with 14C with those of the labelled chelate itself. Radioactivities were measured by internal sample liquid scintillation counting with a Packard 3003 Tri Carb liquid scintillation spectro- meter. A standard toluene-based organic scintillator was employed ; in all cases the chemical compositions of the solutions, and hence the counting efficiencies were identicaL2 MATERIALS VO(aca~)~Cl was synthesized as described by Funk, Weiss and Zeising ; dichlorotri- (dimethylsulphoxide)oxovanadium(w) (VO(DMS0)3C12) was prepared as reported by Selbin and Holmes.* Azobisisobutyronitrile (Koch-Light) was recrystallized twice from AnalaR grade methanol and once from chloroform. It was then stored in the dark in vacuum.Dimethyl sulphoxide (B.D.H.) was dried for one week by molecular sieves, then frac- tionated under nitrogen (14mmHg) as recommended by Martin, Weise and Niclas.lo Scavenging of DMSO was effected by the technique of Atkinson, Bamford and Eastmond.2*11 Pyridine (B.D.H., AnalaR) was dried over fresh barium oxide for 24 h before being frac- tionated under an atmosphere of nitrogen.12 Benzene (B.D.H., AnalaR) was dried with sodium-wire for 24h, then distilled in vacuum and stored in the dark under nitrogen.Ethyl acetate (Hopkin and Williams, AnalaR) was dried for one week over anhydrous potassium carbonate and then fractionally distilled (b.p. 77°C at 76 mmHg). Methanol (AnalaR) was used without further purification. VO(acac)zC1 with labelled ligands was prepared from [l ,3-14C2]acetylacetone, synthesized by Claisen condensation l3 from [1,3-14Cz]acetone (Radiochemical Centre, Amersham) and ethyl acetate.Methyl methacrylate was purified by the method of Bamford and Lind.954 PHOTOINITIATION OF POLYMERIZATION RESULTS AND DISCUSSION KINETICS OF PHOTOINITIATED POLYMERIZATION The rate o of photoinitiated polymerization (A = 365 nm) of methyl methacrylate in the presence of dimethyl sulphoxide or pyridine is shown in fig. 1 to be proportional to ~ O ( a ~ a c ) ~ C l ] + over the range studied for constant incident intensity I. and reactant composition (volume fraction of additive v, = 0.1 in each case). Fig. 2 indicates that, for constant [VO(a~ac)~Cl], (I) is proportional to I$ for the system containing DMSO. Polymerization under these conditions is therefore an uncomplicated free-radical process. The overall rate observed cu, contains small contributions from thermal and uncatalyzed reactions (rates cot, (I),, respectively). The latter were 102[VO(acac)2CI]~/mol~ dm-4 FIG.1 .-Dependence of initial rate of photosensitized polymerization of methyl methacrylate at 25°C on [VO(acac),Cl]*. 0, Dimethyl sulphoxide as additive, Va = 0.1 ; 0, pyridine as additive, va = 0.1 ; h = 365 nm ; I. = 1.72 x einstein dm-3 s-l. 1041t/einstein* dm-3 s-* FIG. 2.-Dependence of initial rate of photosensitized polymerization of methyl methacrylate at 25°C on 18. [VO(a~ac)~Cl] = 1 . 0 ~ mol dm-3 ; Va(DMS0) = 0.1 ; A = 365 nm.S. M. A L I W I AND C. H. BAMFORD 55 measured directly under the appropriate conditions, cot without irradiation but with all components present and co, with irradiation but with monomer alone in the dilato- meter.Both cot and co, were less than 3 % of coo. Values of co in fig. 1 and 2 have been derived from coo by means of the relation 0 = (co:-w:-co:)+. Values of the kinetic parameter k,k, (kp, k, being the rate coefficients of propaga- tion and second-order termination, respectively) were determined as a function of u, from observations of rates and degrees of polymerization, with the aid of relation (1) of the previous The transfer constants to monomer and DMSO have values of 2 x and 7.1 x (60°C), r e ~ p e c t i v e l y . ~ ~ * ~ ~ Under our conditions, transfer has a negligible effect on molecular weights and calculations of k,k;+. In these experiments the conversion of monomer was kept below 3%. The results, presented in fig. 3 and 4, show that in both systems k,k;+ increases markedly with u,.- - - - - 1.8 1.6 1.4 o a 1.2 5. 2 1.0 -0.8 -0.6 0.4 0.41 I I I ' ' ' I ' I .o 0 0.2 0.4 0.6 0.8 v,(DMSO) FIG. 3.-Dependence of kpkt-* and kp/k; for methyl methacrylate at 25°C on volume fraction of dimethyl sulphoxide. 0, kpkF* ; photoinitiation by VO(acac)zC1 + DMSO ; [VO(acac),CI] = 1 . 0 ~ mol dm-3 ; h = 365 nm. A, k,kt-* ; photoinitiation by azobisisobutyronitrile (5 x mol dm-3). 0, kp/kg. 0.4 2 2 .o 1.6 1.2 0.8 0.4 0.0 A? 0 0 . 2 0.4 0.6 0.0 1 .o v,(p yridine) FIG. 4.-Dependence of kpkt-* and kp/kg for methyl methacrylate at 25°C on volume fraction of pyridine. mol dm-3 ; 0, kPkt-* ; photoinitiation by VO(acac)2C1 + Py ; [VO(acac),Cl] = 1 .O x h = 365 nm. 0, k,/k;.56 PHOTOINITIATION OF POLYMERIZATION If it is assumed that k, is affected only by the viscosity of the reaction medium, being inversely proportional to the latter as demonstrated by North and Reed,16 relative values of kp may be calculated from these observations and are shown in fig.3 and 4. Clearly both dimethyl sulphoxide and pyridine belong to the class of additives which significantly increase the observed propagation coefficient of methyl methacrylate. Our findings with DMSO agree within experimental error with those of Bamford and Ferrar;2 the behaviour of pyridine in this context does not appear to have been previously recorded, but comparison of fig. 3 and 4 reveals that its “activity” exceeds that of DMSO and is therefore greater than that reported for any other diluent. A recent discussion of phenomena of this type has been given by Bamford.17 Some values of k,k;* obtained with photoinitiation by azobisisobutyronitrile are shown in fig.3 and fall into line with those already mentioned. Table 1 shows that the apparent values of k,k,& for reactions photo-initiated by VO(acac)2C1 + DMSO (0, = 0.166) are not significantly dependent on [VO(acac),Cl] ; these data are therefore consistent with the effective absence of chain-transfer, retardation and primary termination. TABLE l.--k,k;* AS A FUNCTION OF [VO(a~ac)~Cl] WITH DMSO AS ADDITIVE; Ua = 0.166, A = 365 nm, 25°C 104[VO(acac)~Cl]/mol dm-3 0.666 1.00 2.00 4.00 5.00 8.00 10.00 20.00 30.00 k P t k-*/mol-* dm3 s-3 0.0649 0.0695 0.0713 0.0675 0.0639 0.0642 0.0637 0.0669 0.0641 mean: kPkt-* = 0.0662 mol-* drn’s-’ The order of the overall reaction in monomer for 21, (DMSO) = 0.1 was determined with benzene and ethyl acetate as diluents.A small correction was applied in the latter case to allow for changes in viscosity of the medium,16 but no correction was necessary when benzene was the diluent. Fig. 5 presents log w against log [MI plots and shows that under the conditions studied the order in [MI is close to unity; the slopes in fig. 5 are 0.99 and 1.08 for benzene and ethyl acetate, respectively. The rate of initiation is therefore effectively independent of the monomer concentration. 0 0.2 0.4 0.6 0.8 1.0 log[MMA]/mol dm-3 FIG. 5.-Initial rate of photosensitized polymerization of methyl methacrylate at 25°C as function of monomer concentration. [vO(a~ac)~Cl] = L O X mol dm-3 ; A = 365 nm; lo = 1.8 x einstein dm-3 s-l.0, diluent ethyl acetate ; 0, diluent benzene.S. M. ALIWI AND C . H. BAMFORD 57 o, 0.1 0.2 0.3 0.4 va(DMSO) FIG. 6.-Dependence of initial rate of photosensitized polymerization of methyl methacrylate at 25°C on volume fraction of dimethyl sulphoxide. [VO(aca~)~Cl] = 1.4 x mol dm-3 ; h = 365 nm ; I , = 1.8 x 10-6einstein dm-3 s-l. ..I kn 40.8 Ua(DMS0) FIG. 7.-Dependence of rate and quantum yield of initiation at 25°C on volume fraction of dimethyl sulphoxide. [VO(aca~)~Cl] = 1.4 x mol dm-3 ; h = 365 nm ; lo = 1.8 x 10-6einstein dm-3 s-l. -, $calculated from rates of polymerization(fig. 6) and values of kPkt-+ (fig. 3) with the aidof eqn(2) ; ---- , -d[I]/dt. va(pyridine) FIG. 8.-Dependence of rate and quantum yield of initiation of 25°C on volume fraction of pyridine.[VO(a~ac)~Cl] = 1 .O x 10-6einstein dm-3 s-l. -, 9 calculated from rates of polymerization and of kpkc8 (fig. 4) with the aid of eqn (2) ; - - -, - d[I]/dt.58 PHOTOINITIATION OF POLYMERIZATION The results so far described show that the rate of polymerization conforms to the conventional relation (2), in which [MIo is the bulk monomer concentration and 9 the rate of initiation. U) = kpkc3[M]f~ = kpk,'[M]o(l -~,)9'. (2) This equation has been used in deducing rates of initiation from observations of w over a range of Va with the aid of the results in fig. 3 and 4. The variation of cr) with va (DMSO) is illustrated in fig. 6 and values of Y calculated from w against va curves for the two additives are presented in fig.7 and 8. The dependence of 9 on v, is similar for DMSO and Py ; a maximum in 9 occurs near o, = 0.1, It follows from the results in fig. 1-3 that, for oa = 0.1, the rates of initiation are given by DMSO : 9 = 136.9 Io[VO(acac)2Cl]mol dm-3 s-l ( 3 4 Py : 9 = 23.8 Io[VO(acac)2Cl]mol dm-% s-l (3b) QUANTUM YIELDS These were determined by the techniques described in the previous paper and the results are displayed in tables 2 and 3. Quantum yields for initiation and the overall reaction are denoted by 4i and 40, respectively. In bulk methyl metha- crylate 4i = 2.06 x ; thus it is clear that addition of DMSO or Py (u, = 0.1) produces a large increase in the quantum yield of initiation (by factors of 29 and 6, respectively). TABLE 2.-QUANTUM YIELDS IN METHYL METHACRYLATE+ DIMETHYL SULPHOXIDE (Ua = 0.I). [VO(aca~)~Cl] = 10-3mol dm-3 ; 3, = 365 nm ; 25°C ; kpkL3 = 0.062 mol-3 dms s-*. lO6I0/ lo71abs/ 1 0 4 4 10~91 40 4i einstein dm-3 s-* einstein dm-3 s-1 mol dm-3 s-1 mol dm-3 s-1 1.91 3.97 2.55 2.35 642 0.59 1.91 3.97 2.50 2.26 630 0.57 1.70 3.53 2.26 1.84 640 0.52 1.70 3.53 2.58 2.40 730 0.68 mean : 4i = 0.59 TABLE 3.-QUANTUM YIELDS IN METHYL METHACRYLATE+ PYRIDINE (Va = 0.1). [VO(aca~)~Cl] = mol dm-3 ; 3, = 365 nm ; 25°C ; k,k,* = 0.063 mol-3 dm3 s-3. 1061,1 1071~bs 1 0 4 4 1 0 8 9 1 40 4i einstein dm-3 s-1 einstein dm-3 s-1 mol dm-3 s-1 mol dm-3 s-1 2.37 4.64 1.31 5.88 282 0.127 2.37 4.64 1.28 5.63 276 0.122 mean : $i = 0.125 IDENTIFICATION OF TERMINAL GROUPS I N POLYMERS Experiments with VO(a~ac*)~Cl prepared from acetylacetone containing [1 ,3-l4CZ] acetylacetone were carried out to examine the participation of acac radicals in initia- tion.Poly(methy1 methacrylate) specimens prepared by photoinitiation with VO(a~ac*)~Cl in the presence of dimethyl sulphoxide (u, = 0.166) were purified by several reprecipitations into methanol and submitted to scintillation counting. In a typical experiment it was found that the number of counts per mole of polymer was 9.3 x 106/100 s, the number per mole of initiator (chelate) being 2.6 x 108/100 s.S . M. ALIWI AND C . H . BAMFORD 59 The proportion of growing chains with acetylacetone terminations was therefore less than 3 %. Thus acat radicals are not significantly involved in initiation. Chlorine contents were determined by neutron-activation analysis, with results given in table 4.Initial determinations by this technique of the sulphur contents of polymers prepared in the presence of DMSO gave high values, corresponding to 4.77 S-atoms per polymer chain. Related findings with DMSO labelled with 35S have been recorded by other workers '9' using different initiators (including azobis- isobutyronitrile) and were attributed to the presence in DMSO of radioactive im- purities capable of copolymerizing with methyl methacrylate. Attempts were there- fore made to purify the DMSO by scavenging the impurities by radicals obtained from azobisisobutyronitrile, following the technique of Atkinson, Bamford and Eastmond." Two such treatments led to a considerable reduction in the sulphur content of the polymers, although it was clear that a significant amount of impurity remained in the DMSO.It was therefore decided to estimate the extent of incorpora- tion of sulphur in the polymer during photoinitiation by VO(a~ac)~Cl+ DMSO by an indirect method. Polymers were prepared at 25°C from methyl methacrylate + (scavenged) dimethyl sulphoxide (u, = 0.05) with photosensitization (A = 365 nm) either by the VO(acac),Cl + DMSO complex or by azobisisobutyronitrile (AZO) ; conditions were arranged so as to obtain effectively equal rates of polymerization in these experiments. After suitable purification by reprecipitation, the polymers were submitted to neutron-activation analysis and the difference in sulphur contents was attributed to incorporation of sulphur-containing fragments on photoinitiation by VO(acac),Cl + DMSO.This technique has evident disadvantages, but a satisfactory alternative could not be devised without the availability of pure DMSO. Results are summarized in table 4. TABLE 4.-NEUTRON-ACTIVATION ANALYSIS OF POLYMERS. initiating system 10-5Fn no. of C1 atoms S atom per polymer molecule VO(acac),Cl+DMSO 0.962 0.40rfr0.03 1.12kO.13 AZO+DMSO 1.12 0.0 0.41 0.08 VO(a~ac)~Cl+ Py 1.39 0.92rfr0.11 0.0 (Ua = 0.05) (u, = 0.05) (Va = 0.1) [VO(acac),Cl] = mol dm-3 It appears that photoinitiation by VO(acac),Cl + DMSO leads to polymer molecules containing on the average 0.40 Cl and 0.71 S atoms. Since the ratio of combination to disproportionation in the termination reaction at 25°C is 0.52,19 the total endgroup content arising from initiation should be 1.21 ; hence initiation by e l and by DMSO fragments accounts for nearly all the observed photoinitiation by VO(acac),Cl + DMSO.In the case of VO(a~ac)~Cl+Py, the bulk of the initiation arises from C1 atoms. We have previously shown that photoinitiation by VO(a~ac)~Cl in the absence of strong electron donors proceeds through formation of chlorine atoms. SPECTRAL A N D CONDUCTIVITY OBSERVATIONS ON COMPLEX FORMATION Addition of a strong electron donor such as dimethyl sulphoxide, pyridine or one of a wide range of amino compounds to a solution of VO(acac),Cl in methyl metha- crylate produces (in inactive light) a colour change from deep blue to pale greenish- yellow as illustrated by the absorption spectra in fig. 9. The change appears to be60 PHOTOINITIATION OF POLYMERIZATION complete for concentrations of electron donor slightly in excess of [VO(acac),Cl] and further addition produces no effect on the spectrum.All strong donors give final spectra of the same type, which is independent of the solvent initially present (methyl met hacry late, benzene, cyclo hexane , acetone, e t h y 1 acetate) . 325 350 400 450 500 550 600 650 700 750 800 wavelength /MI FIG. 9.-Absorption spectra. Solvent methyl methacrylate ; 25°C. (1) VO(a~ac)~Cl(l.O x rnol dm-3); (2) VO(acac)2Cl ( 1 . 0 ~ mol dm-3)+DMS0 (Va = 0.1); (3) VO(acac)2C1 ( 1 . 0 ~ loA3 mol dm-3)+DMSO(~, = 0.1) after irradiation ( A = 365 nm) for 20 min ; (4) as (3), but after irradiation for 45 min. Pathlength 10 mm. We believe these results indicate the formation of a new complex between VO(acac),Cl and the donor additive D and propose that it is an ionic complex of the type shown in (I).(I) [ (acac), ij' V C1- D The band in the absorption spectrum of VO(acac),Cl in methyl methacrylate solu- tion with a peak near 600 nm corresponds to charge transfer from p-orbitals of C1 to d-orbitals of V, so that its disappearance on complex formation is under- standable. The extinction coefficients of the complexes at A = 365 nm are 101.2 and 96.0 mo1-ldm3 cm-' for D = DMSO and Py, respectively. TABLE 5 .-ELECTRICAL CONDUCTIVITIES OF VO(aCaC)2 c1 SOLUTIONS AT 25°C solution tVO(acac)~C11/ ua 1 0 6 ~ specific mcl dm-3 conductancelf2-1 cm-1 MMA 0.0 MMA+ DMSO 0.0 MMA+VO(acac)zC1 2.36 x MMA+ VO(acac)zC1 + DMSO 2 .0 ~ 10-3 2 . 0 ~ 10-3 2.0 x 10-3 2.0 x 10-3 2.0 x 10-3 M MA + VO(acac)2 C1 + PY 2.37 x 2.37 x 2.37 x 1W2 2.37 x 0.0 0.1 0.0 0.1 0.2 0.4 0.6 0.8 0.1 0.4 0.6 0.8 0.1 0.1 0.12 8.2 13.0 17.3 20.0 24.0 0.63 1.5 4.2 9.8S. M. ALIWI AND C . H . BAMFORD 61 In agreement with the views expressed above, we find that addition of the donor is accompanied by a large increase in electrical conductivity. Some typical results are presented in table 5. On addition of pyridine to a concentrated solution of VO(acac),Cl in ether light green crystals separated. After washing with ether and drying these gave the following elemental analysis ( %) : C 48.4, H 4.8, N 4.01. Calculated values for (I) (D = Py) are C 47.15, H 5.04, N 3.71. Attempts to isolate the complex formed between VO(acac)2C1 and DMSO have not been successful.CHEMICAL CHANGES PRODUCED BY IRRADIATION Irradiation at 25°C of a solution of VO(acac),Cl in methyl methacrylate containing DMSO by light with A = 365 or 436 nm produces a colour change from greenish- yellow to pure green; the corresponding spectra are shown in fig. 9. Essentially similar results are obtained when DMSO is replaced by Py (fig. 10). The low- intensity peaks at 770, 650 and 400 nm for the DMSO complex which develop on irradiation are near those typical of d-d transitions in V" (e.g. vanadyl ion) deriva- tives,20p21 indicating that reduction of Vv (do system) to VIV (dl system) has occurred. 1.0 - 0.9 = 0.8 - ?? 0.7 - 8 0 . 6 - rn - 0.5 - *= 0.4 - a 0.3 - 0.2 0.1 .-. 8 - - 350 400 450 500 550 660 650 o* ' wavelength /nm FIG. 10.-Absorption spectra.Solvent methyl methacrylate ; 25°C. (1) [VO(aca~)~Cl] = 1.0 x mol dm-3 +Py (ua = 0.1) ; (2) as (l), after irradiation for 45 min ; (3) as (l), after irradia- tion for 90 min ; (4) as (l), after irradiation for 150 min. - - -, spectrum of VO(acac),(l.O x lo-' mol dm-3)+Py (ua = 0.1). Wavelength of irradiation 365 nm ; pathlength 10 mm. The precise location of the bands depends on the nature of the E.s.r. observations confirm this photo-reduction. Both VO(acac),Cl and the complex (I) are diamagnetic and do not give an e.s.r. spectrum. On irradiation of the DMSO complex an e.s.r. spectrum develops steadily as shown in fig. 11 ; this spectrum is typical of those of V02+ derivatives e.g. vanadyl porphyrin 23 and vanadyl acetyl- acet onate.24 After prolonged irradiation at 25°C of VO(acac)2C1+DMS0 in ethyl acetate or benzene solution (A = 365 nm) pale green-blue crystals separated out with a composi- tion corresponding to VO(DMS0)3C12 : found C 19.41, H 4.97, Cl 19.09, S 27.26% ; calculated C 19.35, H 4.88, C1 19.20, S 26.00%. The crystals had U.V. and i.r. spectra identical with those of a specimen of VO(DMS0)3C12 prepared from VOClz and DMSO.* Examination of the irradiated benzene solution by gas-liquid chroma- tography established the presence of acetylacetone, this being the only detectable volatile product.62 PHOTOINITIATION OF POLYMERIZATION Different results were obtained with VO(acac),Cl+ Py in benzene solution. Prolonged irradiation in this case does not yield any precipitate and the final absorp- tion spectrum is indistinguishable from that of the VO(acac),Py adduct (fig.10, dotted line). Further, no acetylacetone could be detected by gas-liquid chromato- graphy after irradiation. The significance of these observations is discussed below. -+ H FIG. 11.-E.s.r. observation at 77 K ; methyl methacrylate solution. mo 1 dm-3. (1) VO(acac),Cl ; (2) VO(acac)2C1+DMS0 (0, = 0.1) before irradiation ; (3) as (2) after [VO(a~ac)~Cl] = 1.0 x irradiation for 3 min ; (4) as (2) after irradiation for 8 min. h = 365 nm. RATES OF COMPLEX DECOMPOSITION A N D RADICAL YIELDS The spectral changes accompanying irradiation described in the previous section and illustrated in fig. 9 and 10 may be used to estimate rates of decomposition of the ion-pair complex.These, together with the rates of initiation of polymerization, (fig. 7 and 8) enable the radical yield n (i.e. the number of initiating radicals produced by photolysis of one molecule of complex) to be deduced. TABLE 6.-VALUES OF k d AND F2 AT 25°C. OPTICAL DENSITY MEASUREMENTS AT 775 nm. [VO(a~ac)~Cl] = 10-3mol dm-3 ; u,(DMSO) = 0.1 ; DILUENT BENZENE ; A(irradia- tion) = 365 nm; I,, = 10-6einstein dm-3 s-' - 107 d[o1 volume 104ki/s-' d t I 10~91 n fraction cf mol dm-3 s-1 mol dm-3 s-1 monomer 0.90 1.72 6.19 4.97 0.80 0.55 1.51 5.48 4.97 0.91 0.18 1.71 6.21 4.97 0.80 Optical densities at wavelengths of 775,400 and 340 nm were measured as functions of the time of irradiation. In all cases it was found that A , - A , decreases exponen- tially with time (A, being the optical density after irradiation for time t ) , indicating first-order decomposition of the complex.From the slopes of the log (A, --At) against t lines the first-order decomposition coefficient k, was calculated. Table 6 presents values of kd for v,(DMSO) = 0. I, for a range of monomer concentrationsS . M. ALIWl AND C. H . BAMFORD 63 with benzene as diluent, together with the corresponding values of the rates of complex decomposition, and n. It is clear that under these conditions kd is not much affected by changes in [MI and that n has a value approaching unity. Table 7 shows kd for a range of values of v, for DMSO and Py as additives; no diluent was present in these experiments. Apparently kd is sensibly independent of Va in the range examined.Although n is approximately unity for small v,, it decreases steadily with increasing v,. The differences in kd values in tables 6 and 7 are probably within experimental error. Rates of decomposition of (I) calculated from the values of kd in table 7 are indicated in fig. 7 and 8 to facilitate comparison with #. Since the extinction coefficients are independent of v, the data presented in tables 2, 3 and 7 allow estima- tion of 4i as a function of u, (fig. 7 and 8). REACTION MECHANISMS We believe the experimental observations are best interpreted in terms of a primary photolytic act of electron transfer to vanadium. Two processes may be distinguished, depending on whether transfer occurs from (i) Cl- or (ii) D. Transfer from a chelate ligand appears to be excluded since there is no evidence for the participation of acetylacetonate radicals.In both cases a VIV derivative is formed ; the other primary products are (i) a C1 atom and (ii) a cation radical D'. We now consider the nature of the secondary reactions which must be invoked to explain the experimental data. (i) When D = Py, direct initiation of polymerization by el accounts for the ob- served chlorine content of the polymer. The spectral indications (fig. 10) that the final vanadium-containing product is VO(a~ac)~Py are consistent with this mechanism. If D = DMSO the situation is more complicated since the polymer contains not only C1 but also fragments of DMSO. This suggests that only a fraction of the C1 atoms initiate polymerization, the remainder entering into hydrogen-abstraction reactions with DMSO [eqn (4)] to give CH3SOCH2 radicals which subsequently initiate.(4) Reaction (4) is energetically unfavourable compared to direct addition of C1 to mono- mer, but this may not be important if (4) involves the complexed DMSO molecule in the original ion-pair or takes place before the excitation energy of the complex has been lost. The vanadium derivative resulting from the primary electron transfer is VO(acac),(DMSO), but we shall see that this may react further. (ii) If Dr- is the primary product it is necessary to assume that reaction with Cl- takes place. When D = Py this is simply electron transfer (5) ; the latter is followed by initiation by el : When D = DMSO two processes must be postulated to account for the end-groups in the polymer : (CH3)ZSO + Cl+CH,SOCH, + HC1 Pyt+Cl--+Py+Cl. ( 5 ) (CH3)2SOtCl-+(CH3)2S0 + C1 (6) (CH3)2SOf+ Cl-+CH3SOCH2 +HCl.(7) For reasons given later, we do not believe that direct initiation by Dt is important. Eqn (5), (6) and (7) are no doubt an oversimplification of the mechanism and all the reactions probably involve formation of (DMSO - Cl) as an intermediate within the solvent cages (11) [eqn (S)] : D ~ C I - ~ D - - ci.64 PHOTOINITIATiON OF POLYMERIZATION Complexes of this type have been suggested by Cooper et a125 On the basis of pulse radiolysis studies of hexamethylphosphoric triamide (HMPT) in the presence of NaBr, Koulkes-Pujo and co-workers 26 postulated the formation of a bromine atom charge-transfer complex (HMPT * * Br) which was considered to react with Br- to form Br;.Related processes are likely with DMSO (Koulkes-Pujo et aL2’) ; in our systems formation of (D - Cl) and reaction with monomer would be equivalent to (5) or (6) followed by initiation by el. Decomposition of (D Cl) into the pro- ducts on the right of (7) appears to be feasible on energetic grounds. Koulkes-Pujo et aL2’ did not report observation of CB3SOCH2, but its formation from (D * * Br) may be less likely. These considerations do not materially affect the kinetic treatment, which we base on reactions (6) and (7). The reactions involved in initiation according to mechanism (ii) are summarized in (8). D (1) k - 1 I ( a ) 0 I Ainactive products (4 0 (11) (D = DMS0)- k3 II --+ C1+ (acac),~ :el t DMSO 0 I I LCH3SOCH2 + HC1+ (acac),VIV :f) DMSO i II (acac) VIv t k s DMSO -+inactive products.( 9 ) 1 I I I C1+ CH2=C+Cl-CH2-C* I I I 1 CH3SOCH2 + CH2=C-+CH,SOCH2-CH2-C- The deactivation step (8a) is included to accommodate the observation that the quantum yield for initiation is less than unity (tables 2 and 3). Reaction (8b) is the competing electron transfer leading to the solvent-caged species (11) and reactionsS . M. ALIWI AND C. H. BAMFORD 65 (8c), (8e) and (8f) result from (5), (6) and (7), respectively. Processes (84 g), repre- senting deactivation of (11), will be discussed later. (8h) is merely the formation of the VO(acac),+DMSO adduct by coordination of DMSO. Reactions (9a, b) are directly responsible for initiation and will be supposed to consume all the radicals with 100% efficiency.TABLE 7.-vALUES OF kd AND IE AT 25°C. OPTICAL DENSITY MEASUREMENTS AT 340 AND 350 nm FOR DMSO AND Py, RESPECTIVELY. @-radiation) = 365 nm v, 104kd/s-1 [(I)ll - o7 d[o1 , 1 0791 n d t mol dm-3 mol dm-3 s-1 mol din-3 s-1 DMSO. I. = 1.0 x 10Weinstein dm-3 s-l 0.1 1.40 10-3 1.40 1.37 0.98 0.2 1.43 7 . 5 ~ 10-4 1.07 0.70 0.65 0.4 1.33 7 . 5 ~ 1 .oo 0.13 0.13 0.8 1.34 7 . 5 ~ 1.01 - - mean 1.38 Py. lo = 1.76 x 10-6einstein dm-3 s-l 0.1 0.436 10-3 0.436 0.43 0.99 0.2 0.415 10-3 0.41 5 0.395 0.95 0.4 0.460 10-3 0.460 0.27 0.59 0.6 0.426 10-3 0.426 0.17 0.40 0.8 0.442 10-3 0.442 0.073 0.165 mean 0.436 Mechanisms (i) and (ii) are stoichiometrically similar; however, they may be distinguished kinetically and we shall see that (ii) is preferable from this point of view. The nature of the final products indicates the occurrence of further secondary reactions which we now describe.The products of photolysis in systems containing DMSO, viz. VO(DMS0)&12 and acetylacetone, probably originate from interaction of VO(acac),(DMSO), formed in (8e, h) and HC1, arising from (Sf), e.g. by the sequence DMSO V0(a~ac)~(DMS0) + HCl+VO(acac)(DMSO),Cl+ acacH (10a) DMSO VO(acac)(DMSO)2C1 + HCl+VO(DMSO),Cl, + acacH. (lob) Disproportionation of VO(acac)(DMSO),Cl formed in (1Oa) into VO(DMS0)3C12 and VO(acac),(DMSO) is also possible. In either case, the overall reaction is ZDMSO VO(acac)2(DMSO) + 2HC1+VO(DMS0)3C12 + 2acacH. (1 1) We have obtained direct experimental evidence for this process. A solution of VO(acac), in DMSO, on treatment with anhydrous HCl, was found to produce VO(DMS0),C12, apparently instantaneously.These reactions elucidate satisfactorily two aspects of the experimental data : (a) the appearance of acetylacetone without intermediate formation of acetylacetonate radicals (the latter being excluded by the absence of acetylacetonate residues from the polymer) and (b) the fact that no acetylacetone is formed in the presence of pyridine as additive (no HCl is generated in this system). 1-366 PHOTOINITIATION OF POLYMERIZATION By assuming stationary concentrations of [(V0(a~ac)~D)+Cl-]*, (II), Cl and CH3SOcH2 we obtain from (8), (9) the following relations. k 2 k3+k4 3 = kl~o~{VO(acac)2D]+C~-]- k-,+k2 - k3+k,+k,’ (124 k3 +k4 k3 + k4 + kg’ n = In (12c), y is the ratio of combination to disproportionation in the termination reaction and fc, fD are the average numbers of C1 atoms and D residues per polymer molecule.Eqn (12a) is consistent with the experimental observations on rates of initiation summarized by (3a, b) (cf. fig. 1,2,5). According to tables 6 and 7, n for both systems is close to unity for v, = 0.1, approximately, hence we believe that under these condi- tions k5 is negligible [eqn (12e)l. Comparison of equations (12a, b, c) with the data in equations (3a, b) and tables 2, 3, 4 then yields the parameters in table 8. TABLE 8.-REACTION PARAMETERS, Ua = 0.1, APPROXIMATELY D DMSO PY k1/mol-l dm3 232 190 kZ lk-1 1.4 0.14 kdk3 1.8 0 For very weak absorption the theoretical values of k,/mol-l dm3 are 233 (DMSO) and 221 (Py).The main kinetic differences between the DMSO and Py systems reside in the relatively high value of k2 compared to the deactivation coefficient k-l in the former and in the absence of a reaction corresponding to (8f) in the latter. Fig. 7 and 8 show that the rates and quantum yields of initiation are functions of v, and we now consider the implications of these results. The observed increases in 9 and 6i with v, for v, <0.1 are probably not mechanistically significant ; in this region complex formation may be incomplete and the complex is not completely soluble in the reaction mixture. However, neither of these features is present when v, > 0.1. We have already seen (cf. fig. 3 and 4) that there is no evidence for retarda- tion with increasing v, so that there are no reasons for questioning the validity of the values of 3.Inclusion of the deactivation step (13) k D [(V0(a~ac)~D)+Cl-]* + Djinactive products (1 3) could account for decreasing 9 and +i with increasing v, It would require replace- ment of the factor k2/(k1 +k2) in (12a, b) by k2/(kd1 +k2 +kD[D]) and the resulting mechanism (which would be of the Stern-Volmer type) would predict a linear relation between 3-l or 4;’ and kD[D]. This is not obeyed for either system since such plots are strongly convex to the [D] axis.S . M. ALIWI A N D C. H. BAMFORD 67 It seems to us more likely that the effects we are considering arise from increasing separation of the constituent ions of the complex (VO(acac)2D)+C1- with increasing [D] which is reflected in the conductivity changes presented in table 5.Solvent- separation would reduce the probability of the D'f-Cl- interactions leading to (8c, e,f), since these depend on reactions of species within the solvent cage (11) as already described. At large separation, Cl- may be effectively outside the cage. On this view, the coefficients k3 and k4 [scheme (S)] are functions of Ua and decrease relatively rapidly as v, increases. Reduction in the rates of reactions (8c, e,f) naturally increases the relative importance of (8d, g), so that while k5 is insignificant at v, = 0.1 it becomes increasingly important as v, increases. The nature of (8d, g) is not established by the present results. These reactions are written as first-order processes, but they may be more complex without invalidating the general argument, although, of course, the simple kinetics in (12) may then not apply for v, > 0.1.Thus interaction of species derived from (11) could give rise to inactive (non-radical) products. This might proceed through monomer insertion between V and Df- in (11), forming an adduct which then reacts with a second molecule of the same kind : 0 0 I1 II D+ ! I1 2D II 1 t M-D+ D (acac), V C1- +M + (acac), V C1- I I M-D+ 0 0 2 (acac), V C1- -+ 2 (acac), v+cI-D+M-MD+C~-. I I These reactions resemble those suggested in the previous paper [eqn (12)-(14)] to account for a similar phenomenon, viz. the existence in the photolysis of VO(acac),Cl and Mn(acac)3 of reactions involving monomer which lead to chelate decomposition by non-radical route^.^'^ Alternatively, (11) may first react with the additive D to give a relatively unreactive adduct.It is important to note that k, does not depend on 21, (table 7); this is consistent with (12d) since the latter does not contain k3 or k,. On the other hand, n, which involves both these coefficients (12e), decreases markedly with increasing va (table 7). Mechanism (i), discussed earlier, is based on electron transfer from C1-, and is not easily reconciled with these observations. Decrease in the probability of electron transfer accompanying increased ionic separation would be expected to produce an increase in the rate of the competing deactivation reaction (8a) and hence a decrease in k,. This could only be avoided if the excited complex were to enter into some other decomposition process at precisely the rate required to compensate for the reduction in the rate of electron transfer.Such a mechanism, holding over a range of v, and for both systems, seems highly improbable. This difficulty does not arise in mechanism (ii) since the electron transfer in (8b) is not sensitive to changes in ionic separation. According to mechanism (ii), which we are advocating, direct electron transfer from CI- to vanadium does not occur. This is probably a consequence of the stereochemistry of the complexes (I), in which close approach of C1- and V may be difficult. The mechanism and experimental findings imply, however, that electron68 PHOTOINITIATION OF POLYMERIZATION transfer from dimethyl sulphoxide and pyridine is more facile than that from acetyl- acetone ligands.A variant of this mechanism which we have not so far considered is that initiation occurs exclusively by Dt, leading to the incorporation in a single polymer chain of both D and C1, the latter being bound ionically in terminal groups Cl-D+-. The analytical results in table 4 do not support such a view since (when D = DMSO) the C1 and D contents of the polymer are not equivalent ; further, a large proportion of the chains would have end-groups containing neither species and this, in the absence of chain-transfer, would be difficult to understand. An additional shortcoming of the mechanism is that it does not explain the dependence of 3 on q, for either additive. D. J. Lind, Ph.D. Thesis (University of Liverpool, 1967). C. H. Bamford and A. N. Ferrar, Proc. Roy. Soc. A, 1971,321, 425 ; for discussion and re- ferences see C. H. Bamford in Reactivity, Mechanism and Structure in Polymer Chemistry, ed. A. D. Jenkins and A. Ledwith (Wiley, New York, 1974), chap. 3. C. H. Bamford and A. N. Ferrar, J.C.S. Faraday I, 1972, 68, 1243. S. M. Aliwi and C. H. Bamford, J.C.S. Faraday I, 1974,70,2092. M. G. Evans, M. Santappa and N. Uri, J. Polymer Sci., 1951, 7,243. see V. Balzani and V. Carassiti, Photochemistry of Coordination Compounds (Academic Press, New York, 1970). ’ H. Funk, W. Weiss and M. Zeising, Z. Anorg. Allgem. Chem., 1958, 296, 36. * J. Selbin and L. H. Holmes, J. Inorg. Nuclear Chem., 1962,24, 1111. C. H. Bamford and D. J. Lind, Proc. Roy. SOC. A, 1968,302,145. lo D. Martin, A. Weiss and H. J. Niclas, Angew. Chem. Int. Edn., 1967, 6, 318. l1 W. H. Atkinson, C. H. Bamford and G, C. Eastmond, Trans. Furaduy Soc., 1970,66, 1446. l2 N. Rabjohn, Org. Synth., Vol. IV, 480. l3 A. I. Vogel, Practical Organic Chemistry (Longman, Green & Co., London, 1959), 863. l4 C. H. Bamford, R. W. Dyson, G. C. Eastmond and D. Whittle, Polymer, 1969, 10,759. l5 S. N. Gupta and U. S. Nandi, J. Polymer Sci., 1970,8, 1493. l6 A. M. North and G. A. Reed, Trans. Faraday SOC., 1961,57, 859. l7 C. H. Bamford Molecular Behaviour and the Development of Polymeric Materials ed. A. Led- l8 M. U. Mahmud, Ph.D. Thesis (University of Liverpool, 1972). l9 C. H. Bamford, R. W. Dyson and G. C. Eastmond, Polymer, 1969, 10, 885. 2o T. R. Ortolano, J. Selbin and S. P. McGlynn, J. Chem. Phys., 1964, 41, 262. 21 J. Selbin, T. R. Ortolano and F. J. Smith, Inorg. Chem., 1963,2, 1315. 22 J. Selbin, Chem. Rev., 1965, 65,153. 23 D. Kivelson and S. K. Lee, J. Chem. Phys., 1964,41, 1896. 24 I. Bernal and P. H. Rieger, Inorg. Chem., 1963, 2,265. 25 T. K. Cooper, D. C. Walker, H. A. Gilles and N. V. Klassew, Canaci’. J. Chem., 1973,51,2195. 26 A. M. Koulkes-Pujo, L. Gilles, B. Lesigne, J. Sutton and J. Y. Gal, Chem. Comm., 1974,71. 27 A. M. Koulkes-Pujo, L. Gilles, B. Lesigne and J. Sutton, Chem. Comm., 1971, 749. with and A. M. North (Chapman and Hall, London, 1974), chap. 2.

ISSN:0300-9599    

DOI:10.1039/F19757100052

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Mechanism of Thermolysis of Hexamethyldisilane and the Silicon-Silicon Bond Dissociation Energy BY IAIN M. T. DAVIDSON AND ANTHONY V. HOWARD Department of Chemistry, The University, Leicester LE1 7RH Received 25th April, 1974 A detailed investigation of the thermolysis of hexamethyldisilane in a novel stirred-flow system at temperatures between 770 and 872 K is described. Evidence is presented for a mechanism which accounts for the different behaviour of hexamethyldisilane on thermolysis under different conditions, and a new value of 337 kJ mol-I for D(Me3Si-SiMe3) is deduced. The silicon-silicon bond dissociation energy in hexamethyldisilane (HMDS) is a key quantity on which the values of bond dissociation energies in other organosilicon compounds have been based,l but attempts to measure it by electron impact and by gas kinetics 3 9 have led to values ranging from 201 to 359 kJ mol-l.This spread of values is partly due to the complex behaviour of HMDS on thermolysis, as revealed by several gas kinetic s t ~ d i e s . ~ - ~ At high pressures of HMDS the reaction is relatively simple, the isomeric trimethylsilyl(dimethylsilyl)methane being formed as the sole main product in a radical chain reaction of well-established mechani~m.~'~ At low pressures, however, other products predominate, the product composition and kinetic behaviour being strongly dependent on experimental condition^.^. 4 9 In a static system with mass spectrometric analysis of reaction products, kinetic data were obtained which were interpreted as relating to a radical non-chain mechanism, so that the observed Arrhenius parameters could be identified with those for the initial rupture of the silicon-silicon bond.4 The bond dissociation energy thus obtained was plausible and, in conjunction with electron impact data,g led to values for silicon- methyl and silicon-hydrogen bond dissociation energies which agreed well with those deduced from the thermolysis of trimethylsilane if that thermolysis was also inter- preted as non-chain.' O Although the overall picture was thus reasonably consistent,l some worrying discrepancies became apparent, notably a suspiciously low A factor for silicon-silicon bond rupture (particularly in view of rotating sector experiments on the recombination of trimethylsilyl radicals which gave normal values 9 2, and evidence for a chain mechanism for the formation of methane in the thermolysis of tetramethyl- silane,l in contrast to the supposed non-chain mechanism for the formation of meth- ane in the very similar compound, trimethylsilane. O The thermolysis of HMDS at low pressure therefore merits thorough re-investiga- tion.This paper, enlarging on our preliminary account,14 describes the results of this re-investigation and puts forward a detailed mechanism which accounts for the apparently divergent results described above. EXPERIMENTAL APPARATUS For this re-investigation, we have developed a " pulsed stirred-flow " technique which The technique uses a offers several advantages over conventional methods in gas kinetics. 6970 THERMOLYSIS OF HEXAMETHYLDISILANE quartz stirred-flow reactor, identical in design to that of Mulcahy and William~,’~ but substantially smaller with a total volume of 54.58 cm3.The reactor was housed in a smoothed electric furnace of conventional design. A dried and regulated stream of nitrogen flowed through a sample valve,16 through the reactor, and then through a gas chromatograph. The sample valve was connected to a simple greaseless vacuum line with provision for the storage of reactants and samples of products. This vacuum line was fitted with a pressure trans- ducer, enabling the pressure of vapour in the sampling volume of the valve to be measured. A pulse of reactant vapour of known quantity could thus be injected into the carrier gas stream and carried into the reactor where complete mixing and partial reaction occurred ; the reaction mixture then entered the gas chromatograph for analysis.Thus, one stream of nitrogen served as the kinetic and analytical carrier gas, the flow rate determining the reaction time. This apparatus is very much simpler than a conventional flow system for gas kinetics and eliminates problems due to minute leaks, since the reactor is always at above atmospheric pressure. The technique is of course very economical in reactant compared to the conven- tional continuous flow techniques. As we have found in previous flow experiments with organosilicon compounds,” rigorous purification of the carrier gas is necessary to avoid the formation of siloxanes. The nitrogen carrier gas was dried by passage through a molecular sieve column, then passed through traps containing molten sodium.The latter were subsequently replaced by an ‘‘ Oxy-Trap ’’ column. To analyse the wide range of products formed, the technique of fractional co-distillation was used, with a fore-column which could be cooled to 113 K followed by the main column, 3 m long and 0.64 cm 0.d. packed with 10 % wlw squalane on 60-100 mesh Embacel. This column and the gas density balance detector were housed in an oven at 354 K. The kinetic situation in a pulsed reactor appears to be greatly different from that in a continuous flow reactor, but similar kinetic equations may be derived provided that the rate of reaction is low compared to the rate of expulsion of the pulse from the reactor, i.e. provided that the degree of decomposition is small.Under these conditions, with perfect mixing in the reactor, the basic equation is for the formation of product P from reactant A in a reaction of order n and rate constant k. In the equation, (P) and (A) are the number of moles of P and A measured by the gas chromato- graph, Vis the volume of the reactor andz = V/u, where u is the volumetric flow rate through the reactor. z was about 40 s in most experiments, but was varied between 29 and 112 s. Perfect mixing in the reactor was verified by placing a thermal conductivity detector at the outlet and observing the shape of the response curve when a chemically inert pulse was in- jected. Details of this test and of the derivation of the kinetic equation for the reactor will be published separately.’* For a first order reaction, eqn (i) with n = 1 is valid for any degree of decomposition.MATERIALS HMDS was prepared by treating trimethylchlorosilane (a gift from Dow Corning) with potassium-sodium alloy in ethylbenzene, followed by fractionation and purification by preparative g.l.c6 The isomeric trimethylsilyl(dimethylsi1yl)methane was prepared from HMDS by thermolysis.6 Commercial rn-xylene (B.D.H. L.R. Grade) was purified by pre- parative g.1.c. lY1,3,3-tetramethyl-1 ,3-disilacyclobutane was a gift from I.C.I. Ltd. Other compounds were obtained commercially. The purity of all compounds was determined by mass spectrometry and by infra-red spectroscopy, n.m.r. and refractive index where appro- pr ia te . RESULTS The thermolysis of HMDS was studied at temperatures between 770 and 872 K in the pulsed flow reactor with nitrogen carrier gas above atmospheric pressure.Low pressure pulses of HMDS were used, corresponding to initial concentrations in theI . M. T . DAVIDSON A N D A . V. HOWARD 71 range 1.4 x to 5.5 x mol ~ m - ~ . The products were trimethylsilane (3MS) ; tetramethylsilane (4MS) ; trimethylsilyl(dimethylsilyl)methane (ISO), the isomerisa- tion product ; and 1,1,3,3-tetramethyl-l,3-disilacyclobutane (TMDS). A little meth- ane was also observed at the highest temperature. Excess of rn-xylene was added in a second series of experiments covering the same range of temperature and composition, substantially altering the kinetic behaviour and product composition. Throughout the work, the main kinetic emphasis was on the measurement of the rate of formation of 3MS, but kinetic data over the whole range of conditions were also obtained for the formation of 4MS and over restricted ranges with less accuracy for the other products.Without added rn-xylene, the thermolysis of HMDS alone was found to be considerably more complex kinetically than had previously been supposed. The formation of 3MS and TMDS followed the same rate law, which was about first order at low temperature, but became of higher order at higher temperature particularly at low concentration of HMDS ; the average order was about 1.3. At higher concen- trations and lower temperatures than were used in this work it is known that the isomerisation of HMDS has an order of 1.5. Whilst the order for isomerisation was about 1.5 in the present investigation around 770 K, it increased with increase in temperature, being about 1.8 at the upper end of the temperature range.These results will be analysed in more detail in the Discussion. The formation of 4MS on the other hand was kinetically simple, first order kinetics being obeyed under all conditions. First order rate constants were calculated for the formation of 3MS (and TMDS) and 1.5 order rate constants for isomerisation, to facilitate comparison with earlier results 4 9 although of course no quantitative deductions can be made from such " rate constants ". Arrhenius parameters are collected in table 1, error limits only being appropriate for the formation of 4MS. TABLE 1 .-ARRHENIUS PARAMETERS FOR THE FORMATION OF PRODUCTS IN THE THERMOLYSIS OF HMDS WITHOUT ADDED m-XYLENE product log A* E/kJ rnol-' apparent order t true order TMDS 3MS } 15.3 295 1 N 1.3 IS0 13.4 207 1.5 -1.8 4MS 13.75 0.7 282+ 12 1 1 .o * first order rate constants and A factors in s-' and second order in cm3 mol-' s-l.t see text. When excess rn-xylene (up to 300-fold) was added, the formation of TMDS was completely suppressed and 3MS was formed at a reduced rate in a first-order reaction (under our experimental conditions m-xylene was thermally stable). The rate of the isomerisation was also reduced and its order was about 1.5. No change was observed, however, in the kinetics of formation of 4MS. First order rate constants for the formation of 3MS in excess of m-xylene were given by log,, k/s-l = (17.53k0.25)- (336.654.0) kJ mol-'/2.303 RT.Although a substantial excess of m-xylene was usually present, relatively small amounts were also found to produce the above effects. DISCUSSION The simplest explanation of the above results is provided by schemes 1 and 2 ; the approximate Arrhenius parameters in scheme 1 will be estimated below. The formation of 4MS will be discussed separately later.72 THERMOLYSIS OF HEXAMETHYLDISILANE SCHEME 1 .-THERMOLYSIS OF HMDS WITHOUT ADDED W~-XYLENE log A EIkJ rno1-I Me3SiSiMe3-+2Me3Si- (1) 17.25 337+ Me3Si-+ Me3SiSiMe3-+Me3SiH+ Me5Si2cH2 (2) 13.2 72 Me3SiCH2SiMe2-+ Me3Si-+ Me2Si=CH2[-+Me2Si SiMe2] (4) 14.6 212 Me3SiCH2SiMez+ Me6Si2-+ Me3SiCH2Si(H)Me2+ Me5Si2cH2 (5) 13.2 80 0.3 4 Me5Si2cH2 -+ Me3SiCH2Si Me, (3) /\ v 2Me3Si--+ Me3SiSiMe3 (6a) 13 0 Me3Si.+ Me3SiCH2SiMe2+ Me3SiCH2Si2Me5 (6b) 12.4 0 2Me3SiCH2SiMe2 -+(Me3SiCH2SiMe2)2 (6c) 11 0 Arrhenius parameters without error limits are approximate estimates.Scheme 1 offers an explanation for the different product composition observed at high and low pressures of HMDS. At high pressures, reaction ( 5 ) is much faster than (4) and the dominant chain sequence is (l), (2), (3) and (5), as has previously been established,6 giving IS0 as the main product with only small quantities of 3MS. At low pressures of HMDS, reaction (4) can compete with (5), pro- pagating a chain reaction producing TMDS and increased quantities of 3MS (TMDS is assumed to result from the dimerisation of the double-bonded intermediate >- Me2Si=CH, 1,10.13,19 SCHEME 2.-THERMOLYSIS OF HMDS WITH ADDED W2-XYLENE (RH) Me3 SiSiMe3 -+2Me3Sie (1) Me3Si*+ RH+Me3SiH+ Re (7) Re+ Me3SiSiMe3-+RH+ Me5Si2CH2 (8) RH+ Me5Si2cH2+R*+ Me3SiyiMe3 (- 8) Me5Si2cHz-+ Me3SiCH2SiMe2 (3) (9) 2R*-+R2 (10) Me3SiCH2SiMe2 + RH-, Me3SiCH2Si(H)Me2 + Re With added m-xylene (scheme 2), chain isomerisation of HMDS continues, but by a different sequence involving xylene and xylyl radicals in reactions (7), (8), (9) and (lo), giving different kinetic behaviour.Formation of 3MS, however, is now non-chain, with the rate equalling 2kl[HMDS]. The experimental rate constant for the formation of 3MS in rn-xylene thus equals 2kl, and hence the Arrhenius parameters for reaction (1) in scheme 1 are obtained, the activation energy of 337 kJ mol-1 being identified with D(Me,Si-SiMe,). This value of E l , with those for E2 and E5 estimated previously,6 enabled the activation energy for the isomerisation of HMDS at higher concentration to be calculated as 249 kJ mol-I in good agreement with the experimental value of 251 & 8 kJ mol-I.Estimation of A factors is rather more difficult. The rate of combination of trimethylsilyl radicals has been measured in the gas phase and in solution. No activation energy was observed and log A6a was 14.26 in the gas phase and 12.74 in Whilst the former figure seems to be rather high, there is little doubt that trimethylsilyl radicals have a comparable recombination rate to methyl radicals. Log A l -log ACia has been put at 4.2 from estimates of the entropy of HMDS and the trimethylsilyl radical.ll Hence log A,, = 13.0 from our value of 17.2 for log A , .I .M. T. DAVIDSON AND A . V. HOWARD 73 There were indications from the kinetic study of the isomerisation of HMDS at higher concentration that the A factor for reaction (6c) should be as low as possible.6 A reasonable estimate in the light of current ideas on recombination would be log A 6c = 11, whence log A66 = 12.4 by the geometric mean rule. To reconcile the experi- mental results obtained in this work and in the isomerisation study at higher concen- tration with scheme 1 it is necessary to postulate that reactions (2) and (5) have rela- tively high A factors. There is no reliable experimental evidence on these,6 and we simply estimate log A2 = log A5 = 13.2. These estimates may well require revision as more information becomes available on abstraction and combination reactions of silicon-containing radicals.The relative rates of reactions (4) and (5) may be deduced from the relative yields of TMDS and I S 0 in our experiments. Since reaction (4) is first order and (5) is second order, it follows from eqn (i) that V(IS0) (ii) 5- - 2V( ISO) - k4 (Me,Si=CH,)(HMDS) - (TMDS)(HMDS) where the bracketed quantities are number of moles and not concentrations. The most consistent product analyses over a range of (HMDS) were found at 795 and 820 K, giving average values for k5/k4 of 1.894 x lo7 and 1.033 x lo7 respectively. These figures, with the Arrhenius parameters for reaction (5) estimated above, gave the values of E4 and log A4 shown in scheme 1.When m-xylene was added in excess, the rate of formation of 3MS was reduced by a factor varying from 4.5 at 770 K to 2.0 at 872 K, indicating that the formation of 3MS in the absence of m-xylene is a chain reaction of short chain length. For this chain reaction the approximate Arrhenius parameters in scheme 1 may be used to assess the relative importance of the three termination reactions from th? relative concentrations of the chain-carrying radicals, given by [Me,Si-]/[Me,SiCH, SiMe,] = k4/k2[HMDS], and hence to predict the order and Arrhenius parameters under different conditions. Over most of the range of temperature and concentration of HMDS, termination by (66) would be the most important, whence formation of 3MS would be first order, with rate constant (klk,k4/k6b)0-5, log A-16.3, and E-311 kJ mol-l. However, termination by (6a) would be more important at high temperature and low concentration, giving a 1.5 order reaction with rate constant (k,kg/k6a)0-5, log A z 15.3, and Ez241 kJ mol-l.The chain length of this 1.5 order reaction would drop below unity at the lowest concentrations of HMDS ; for example, the cut-off at 870 K would be below 3 x mol ~ m - ~ . Termination by (6c) is only important at low temperature and high concentration; this termination would give an order of 0.5, but the chain would be slower than reaction (1) under these conditions and this process can therefore be ignored. Formation of 3MS in the absence of m-xylene would thus be expected to have an order of either 1 or 1.5 in HMDS, depending on the experimental conditions.The extent to which the experimental results were consistent with this prediction was examined by comparing experimental points with lines of slope 1.5 or 1.0 and intercepts calculated from the appropriate Arrhenius parameters on a plot of log(3MS) against log(HMDS). For the regions of temperature and concentration where 1.5 order behaviour would be expected, experimental points were corrected by subtracting the amount of 3MS formed in reaction (I) from the total observed, to take account in a simple way of the cut-off discussed above. Results at three typical temperatures are given in fig. 1 , showing quite satisfactory agreement, and increasing confidence in the validity of the estimates of Arrhenius parameters for reactions (2), (4), (5) and (6).It is known that formation of I S 0 in the absence of m-xylene is a clean 1.5 order74 THERMOLYSIS OF HEXAMETHYLDISILANE reaction at lower temperature and higher concentration than were used in this work, with log A = 16.65 and E = 251 kJ mol-l. These kinetic characteristics would be observed in the present work also, but at high temperature and low concentration there would be a significant contribution from a sequence terminated by (6b), which would be second order in HMDS, with rate constant (klk2k-~/k4k6b)o*s, log A = 14.9, and E= 179 kJ mol-'. This accounts for the difference in order and Arrhenius parameters for isomerisation between this work (table 1) and the earlier.6 log (HMDS) FIG. 1 .-Thermolysis of HMDS. Comparison of experimental points with calculated lines.A, 861.3 K, n = 1.5; B, 820K, n = 1; C, 778.7K, n = 1.0. At 870 K with added rn-xylene some methane was observed, the molar ratio of methane to 3MS being 0.2, corresponding to a first order rate constant of 3.7 x s-l for the formation of methane at 870 K. Methane most probably results from the minor dissociation Me,SiSiMe3 -,Me3SiSiMe2 + Me- (la). Combining our estimate of k6, with the well established rate constant for the combination of methyl radicals we calculate by the geometric mean rule a value of 1013-6 cm3 mol-l s-1 for the A factor for the combination of methyl and trimethylsilyl radicals to form 4MS. The A factor for the reverse reaction, dissociation of 4MS, Me4Si+Me3Sib + Me., may then be estimated as 10 17*6 cm3 mol-l s-l from a recent value for the entropy change.ll If Al, is assumed to have about the same value then E l , = 350 kJ mol-l, which may be identified with the silicon-methyl bond dissociation energy in HMDS.The bond dissociation energies deduced from this work, and other recent values consistent with them are in table 2. The new value for D(Me,Si-SiMe,), with aI . M. T . DAVIDSON A N D A . V . HOWARD 75 recent literature value 2 o for AH"fMe,Si,),, gives AHi(Me,Si-), = - 11 kJ mol-'. These thermochemical quantities may be used to speculate upon the mode of forma- tion of 4MS in the thermolysis of HMDS. TABLE 2.-BOND DISSOCIATION ENERGIES bond D/kJ mol-' Me3Si-SiMe3 337 la Me5Si2CH2-H 400f Me3Si-H 368 b9c-d Me3SiCH2-H 406 Me,SiCH,Si(Me,)-H 360 f Me5Si2-Me 350 a a this work ; b ref.(14) ; C I. M. T. Davidson, M. Jones, and H. F. Tibbals, unpublished work ; dR. Walsh and J. M. Wells, Chem. Cumm., 1973, 513 ; e J. A. Kerr, A. Stephens and J. C. Young, Int. J. Chem. Kinetics, 1969, 1, 339 ;f ref. (6). 4MS was the only product formed in a first order reaction with Arrhenius para- meters unaffected by added m-xylene, which suggests that it may result from a uni- molecular elimination : Me,SiSiMe, +Me,Si + Me,Si:. (1 1) Dimethylsilylene would then react with HMDS, probably yielding Me,SiSiMe,CH,- SiMe,H by insertion into one of the eighteen carbon-hydrogen bonds. (Our g.1.c. apparatus would not have produced a well-defined peak for a product of such high molecular weight.) This explanation for the formation of 4MS could reconcile the results of our present and earlier investigations of the low-pressure thermolysis of HMDS.In the earlier work,4 concentrations of HMDS below lo-* mol ~ m - ~ were thermolysed between 796 and 828 K. Kinetic data were obtained by measuring the rate of increase of a distinctive peak at m/e 203 in the mass spectrum of the reaction mixture. This peak, due to Me5Si,CH,SiMe2+, was believed to come from Me5Si2CH,SiMe3, a radical combination product formed in the final step of a radical non-chain sequence. First order rate constants were given by log k/s-l = (1 3.5 &- 1 .O)- (281.6 k9.2) kJ mol-'/2.303 RT, the activation energy being identified with D(Me,Si-SiMe,) in accordance with the above interpretation. Over the entire temperature range, the above expression gives rate constants which are smaller than k , (scheme 1) by a factor of at least two; hence neither the rate of dissociation nor a chain reaction was being observed.However, the above Arrhenius parameters are essentially the same as those obtained in this work (table 1) for the formation of 4MS, and Me,Si,CH,SiMez would be a prominent peak in the mass spectrum of the in- sertion product Me5Si2CH2SiMezH. Hence, in the earlier in~estigation,~ kinetic measurements were probably made on a product resulting from reaction (11) and not reaction (1). The chain sequence shown in scheme 1 would have proceeded concurrently with the sequence initiated by reaction (1 1) ; 3MS was indeed observed but could not be measured mass spectro- metrically in the presence of so many other similar compounds, all with prominent peaks due to Me,Si+.At the very low pressures of these experiments no isomerisation would have been expected, but Me,Si=CH, should have been formed in reaction (4). The mass spectrum of TMDS is distinctive and could easily be detected ; TMDS and other cyclic compounds were in fact formed in preliminary experiments at higher pressure but were definitely not present in the main series of experiment^.^ Probably Me,Si=CH, diffused to the wall instead of dimerising in the gas phase. since the76 THERMOLYSIS OF HEXAMETHYLDISILANE pressure of HMDS was very low and there were no added gases, unlike the present study in a stream of nitrogen. If the earlier kinetic results do relate to reaction (1 I), then Ell = 282 kJ mol-l.The activation energy for the reverse reaction, (- 11) is not known but is unlikely to be negligible.,'. 22 From quantitative data 21 on the insertion reactions of silylene, could be as low as 21 kJmol-l, giving AH,, = 261 kJmol-l. Then, from published enthalpies of formation 2o AHf"(Me,Si), = 138 kJ mol-l. This gives the second bond dissociation energy in 4MS, D(Me-SiMe,) as 288 kJ mol-l, sub- stantially less than the first dissociation energy. However, the divalent state is relatively more stable for silicon than for carbon, and there is evidence that in silane the first and second bond dissociation energies are 398 and 249 kJ mol-1 respectively.21 The quantitative model developed here thus offers a coherent explanation for our two low-pressure investigations of the thermolysis of HMDS 4* I4 despite their widely differing results.The same model has been successfully applied to the kinetic results at high pressure of HMDS, where isomerisation occurs almost exclusively.6 Finally, it permits some comment on another mode of decomposition of organodi- silanes. Sommer and his co-workers 23 while studying the photolysis of phenyl- methyldisilanes obtained convincing evidence for an intramolecular hydrogen transfer which they suggested might be an important pathway in the thermolysis of organo- disilanes. The relevant reaction for HMDS would be : + Me,SiH + Me,Si=CH,. ( W The complete suppression of the formation of TMDS by added xylene is clear evidence against the occurrence of reaction (lb) in our thermolysis of HMDS.Whilst Elb cannot be estimated accurately at present, it would almost certainly be rather less than El. However, Alb would be several powers of ten lower than A , and hence reaction (lb) would not be expected to compete with (1) in thermolysis experiments. We thank the S.R.C. for financial support, Dow Corning (Barry) Ltd. and I.C.I. Ltd. for the gift of chemicals, and Dr. D. R. Deans of I.C.I. Heavy Organic Chemicals Division for his interest and encouragement in the development of the pulsed flow technique. I. M. T. Davidson Quart. Rev., 1971, 25, 111. G. G. Hess, F. W. Lampe and L. H. Sommer, J. Amer. Chem. Soc., 1965, 87, 5327 ; J. A. Connor, B. Finney, G. J. Leigh, R. N. Haszeldine, P. J. Robinson, R. D. Sedgwick and R. F. Simmons, Chem. Comm., 1966,178. J. A. Connor, R. N. Haszeldine, G. J. Leigh and R. D. Sedgwick, J. Chem. SOC. A , 1967, 768. I. M. T. Davidson and I. L. Stephenson, J. Chem. SOC. A, 1968,282. C. Eaborn and J. M. Simmie, Chem. Comm., 1968, 1426. I. M. T. Davisdon, C. Eaborn and J. M. Simmie, J.C.S. Faraday I, 1974, 70, 249. P. J. Robinson, personal communication. S. J. Band, I. M. T. Davidson and C. A. Lambert, J. Chem. SOC. A , 1968,2068 and references therein. ' N. Sakurai, A. Hosomi and M. Kumada, Chem. Comm., 1968, 930. lo I. M. T. Davidson and C. A. Lambert, J. Chem. SOC. A, 1971, 882. l1 P. Cadman, G. M. Tilsley and A. F. Trotman-Dickenson, J.C.S. Furaday I, 1972, 68, 1849. l2 G. B. Watts and K. U. Ingold, J. Amer. Chem. SOC., 1972,94,491. l3 R. P. Clifford, B. G. Gowenlock, C. A. F. Johnson and J. Stevenson, J. Organometallic Chem., l4 I. M. T. Davidson and A. V. Howard, Chem. Comm., 1973, 323. l S M. F. R. Mulcahy and D. J. Williams, Austral. J. Chem.. 1961, 14, 534. 1972, 34, 53.I . M. T . DAVIDSON AND A . V . HOWARD 77 l6 G. L. Pratt and J. H. Purnell, Anal. Chem., 1960,32,1213. l7 G. H. Cady and D. P. Seigwarth, Anal. Chem., 1959,31,618. l 8 A. C. Baldwin, I. M. T. Davidson and A. V. Howard, to be published. l9 M. C. Flowers and L. E. Gusel’nikov, J. Chem. SOC. By 1968,419. 2o J. B. Pedley and B. S. Iseard, Catch Tablesfor Silicon Compounds (University of Sussex, 1972). 21 P. John and J. H. Purnell, J.C.S. Faraday I, 1973,69, 1455. 22 I. M. T. Davidson, J. Organometallic Chem., 1970, 24,97. 23 P. Boudjouk, J. R. Roberts, C. M. Golino and L. H. Sommer, J. Amer. Chem. SOC., 1972, 94, 7926; J. R. Roberts, Ph.D. Thesis (University of California, 1970).

ISSN:0300-9599    

DOI:10.1039/F19757100069

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Free Energies of Transfer of Alkali Metal Fluorides from Water to Hydrogen Peroxide+Water and Methanol+Water Mixtures using Ion-selective Electrodes BY ARTHUR K. COVINGTON* AND JENNIFER M. THAIN Department of Physical Chemistry, University of Newcastle, Newcastle upon Tyne NEl 7RU Received 8th April, 1974 Alkali metal ion-responsive glass and fluoride- or chloride-, responsive ion-selective electrodes have been used to obtain free energies of transfer of alkali metal fluorides and chlorides from water to methanol + water mixtures. The precautions and safeguards necessary to obtain reliable results are discussed. The results obtained for the chlorides are in satisfactory agreement with existing data obtained using amalgam double cells. The method was extended to the determination of free energies of transfer for the alkali metal fluorides from water to hydrogen peroxide + water mixtures and the results obtained were compared with spectroscopically-derived values.The Gibbs free energy of transfer, that is the difference between the standard free energy per mole of electrolyte in a pure solvent, usually water, and that in another solvent or mixed solvent, is an important measure of the differences in interaction between the ions of the electrolyte and the solvent molecules in the two media.l Evaluation of this quantity may be made either from e.m.f. measurements or by solubility methods, but the former is the more satisfactory method. Most attention has been directed towards acids using hydrogen gas and silver-silver halide electrodes but in a number of ways values for salts are easier to interpret.For the latter, amalgam electrodes have been employed together with silver-silver halide electrodes. 39 However, Lowe and Smith reported the use of sodium responsive glass electrodes to obtain free energies of transfer of sodium chloride between H20 and H,0+D20 mixtures. No values have been reported for fluorides because of the lack, until recently,6 of a suitable electrode reversible to fluoride ions. The lanthanum fluoride, single crystal, ion-selective electrode developed by Ross and Frant has been used for precise thermodynamic studies in aqueous solutions 7-9 and is suitable for use in solvents other than purely aqueous ones.1o* l 1 In a preliminary study l 2 on the work described in this paper, we used sodium-responsive glass and fluoride ion-selective electrodes to obtain free energies of transfer of sodium fluoride from water to water + hydrogen peroxide mixtures.Neither amalgam electrodes nor silver-silver halide electrodes are usable in solutions containing hydrogen peroxide and this accounts for the paucity of electrochemical work in this solvent medium apart from the study of Mitchell and Wynne-Jones We now describe the precautions necessary for using two ion-selective electrodes to obtain free energies of transfer, and the test of the method for alkali metal chloride solutions in methanol + water mixtures, for which data are available for comparison purposes from amalgam electrode cells.4 Results will be given for three alkali metal fluorides in this solvent mixture which it has not been possible to determine previously, and for four alkali metal fluorides in hydrogen peroxide + water mixtures.Interest in these results arises from an important correlation l4 between the free energy of preferential solvation, a quantity determinable from n.m.r. chemical shift measure- ments and the free energy of transfer. using hydrogen-responsive glass electrodes. 78A . K . COVINGTON AND J . M. THAIN 79 THEORY The cells studied were without liquid junction, namely anion (X) responsive (1) I MX,H20yS I ion selective electrode. cation (M) responsive glass electrode The difference in e.m.f. (AE,) between two cells of type (I) containing solutions of composition MX(mS), S(x), H,O(l-x) with e.m.f. Ex and MX(mw), S(x = 0), H,0(1 -x = 1) with e.m.f.E,, is given by where k = RT(1n lO)/F; m and y k are respectively a suitable concentration variable, and the mean ionic activity coefficient relative to a hypothetical solution of unit concentration in the solvent mixture of mole fraction x of component S . The problem is to choose the most appropriate concentration scale ; for reasons related to our previous work l4 we have chosen the aquamolality scale, defined as number of moles of solute per 55.51 moles of solvent. If we make ms = mw, then A E g , is obtainable from the measured AE,,,, by making an activity coefficient term correction, and where y;, yl can be calculated from a Debye-Huckel-type expression. It can be shown using the method of Robinson and Stokes,15 that the aquamolality activity coefficient and the molality activity coefficient are the same for a given solution irrespective of the choice of aquamolality or molality scales.Thus the aquamolality activity coefficient can be calculated from expression (3) AE, = E, -Ex = AE& + 2k log(msy;/mwy;) (1) A q X ) = AEt(x)-2k h(Y%/Y3 (2) logy* = - -ln(l f0.03603 m) 1 + BaJc where c is the molarity given by 18.015 md WA+0.01S015 mW,' c = (3) (4) The symbols in eqn (3) and (4) have the following significance : A = molar Debye- Huckel limiting law constant ; a = ion-size parameter taken as 4.56 A ; B = molar Debye-Huckel constant in the denominator of eqn (3) ; d = solution density equated to the solvent density ; W, = molecular weight of the solvent calculated from W, = 18.015(1-x)+xWs where W, is the molecular weight of the co-solvent; W, = molecular weight of the solute.The last term in eqn (3) arises from the need to adjust the Debye-Huckel expression for the rational (mole fraction) activity coefficient to give the molal activity coefficient. The quantities A and B, which are functions of the relative permittivity (dielectric constant) of the solvent mixture (E,) and the temperature (T), were calculated from B/cm-l mol-* dm* K* = 50.29 x 108/(~,T)* where values for the dielectric constant were obtained from a recent compilation.16 The activity coefficient correction in eqn (2) amounts at a maximum to 3-4 % in the free energy of transfer when m = 0.01 aquamolal. Some measurements were made at m = 0.005 and 0.02 aquamolal, which confirmed that this method for making activity coefficient corrections obviated the need for an elaborate procedure for extrapolation to infinite dilution.A/mol-* dmt Kt = 1.8246 x 106/(~,T)*80 TRANSFER OF ALKALI METAL FLUORIDES alkali metal ion (M) responsive glass electrode EXPERIMENTAL The potentials of cells (I) were measured using an integrated circuit (MF1-3 electrometer amplifier : Computing Techniques Ltd., Billinghurst, Sussex) interfacing device to a digita 1 voltmeter (DVM) (Advance Electronics Ltd., Bishop’s Stortford, Herts.) with f 0.1 mV discrimination. The cell potentials could also be opposed by a vernier potentiometer, and the DVM used as a null-reading device. The presence of hum pick-up was detectable by a difference between the directly read DVM value and the backed-off value.With strict attention to shielding and grounding, the affects of hum and stray capacitance were rendered negligible. The output from the interfacing device could also be fed to a potentiometric recorder to give a trace of the variation of cell potential with time. The electrodes used and their sources are shown in table 1. As far as possible, several electrodes of the different types available for a given ion were used in the studies. Glass electrodes were conditioned by soaking in 1 rnol dm-3 solutions of the requisite alkali metal ion in 0.1 rnol dm-3 Tris+HCl solution. Conditioning of the halide-responsive ion- selective electrodes is not necessary. Cell vessels were simple H-type with a tap between the compartments.Sets of six cell vessels were placed in an air thermostat (25.0+O.l0C) and pairs of electrodes were transferred between solutions of steadily increased methanol or peroxide content. The Orion lanthanum fluoride electrode was found to be affected by hydrogen peroxide solutions. The trouble was traced to hydrogen peroxide attacking the epoxy cement used to fix the single crystal into the epoxy body. This attack could be prevented by coating the sensitive annular region with paraffin wax. Alternatively a lanthanum fluoride electrode was constructed from a 5 mm diameter, 2 mm thick, undoped, lanthanum fluoride crystal (Metals Research Ltd., Histon, Cambridge) sealed with Araldite into glass fibre-reinforced, epoxy, drilled rod (Bushing Co. Ltd., Hebburn, Co. Durham).The inner filling system was 0.1 mol dm-3 NaFSO.1 rnol dm-3 NaCl I AgCl I Ag. This home-made electrode was quite resistant to attack by hydrogen peroxide and substantially the same results were obtained from it. MCI, HzO, MeOH C1--responsive AgCl (11) (1-x) x 1 ion-selective electrode. TABLE 1 .-ELECTRODE TYPES AND SOURCES anion responsive cation responsive Orion 94-09 fluoride-responsive* Orion 94-17 chloride-responsive* Thermal electrolytic Ag/AgCl Orion Na+-responsive glass, 94-1 1 A* E.I.L. Na+-responsive glass, GEA 331 3 or 33 1048 loo? Corning monovalent cation-responsive glass 476220: E.I.L. K+-responsive glass 33 1057 2007 * Orion Research Inc., Cambridge, Mass., U.S.A. ; f Electronic Instruments Ltd., Chertsey, Surrey ; Corning-E.E.L., Halstead, Essex. Methanol+ water mixtures were prepared from AnalaR grade methanol and doubly distilled water.Hydrogen peroxide 85 % w/w (Laporte Chemicals Ltd., Warrington, Lancs.), free from stabilising additives was diluted with doubly distilled water, and stock solutions were analysed by titration with permanganate. Salts were the highest grade material readily available and were dried before use. All solutions were made 0.01 aquamolal in salt and rnol dm-3 in Tris to buffer the solutions to a pH where the alkali metal ion responsive glass electrodes would not show a mixed response to hydrogen ions.A . K. COVINGTON AND J . M. THAIN 81 The technique employed consisted of allowing the pair of electrodes to reach a steady potential in the purely aqueous solutions and then to transfer them successively to solutions of increasing methanol content.In this way the glass electrode is subjected to only small changes in solvent environment and large changes in solvent medium, which often provoke drifts in potential, avoided. Return transfers down the series to the purely aqueous solution showed that, although the actual steady potential reached in each cell was not necessarily the same as that reached in the up-series transfers, the dzflerences in potential between successive cells i and .j, AElj, differed from AE,, in the downward series by only 1 %. This is illustrated in fig. 1 which refers to an Orion Na+ responsive electrode conditioned and used in caesium chloride solutions. The transfer e.m.f. A,??,,,, in eqn (1) is obtained from AEtj and AEji by X X that is, the transfer e.m.f.from water into a methanol mixture of mole fraction x is obtained by summing the potential differences between successive cells. Values of AE,,,, and, from eqn (2) with eqn (9, values of the free energy of transfer AG&, are collected in table 2. The apparent absence of a negative sign in eqn ( 5 ) arises from the method of defining the transfer e.m.f. and follows the practice of Feakins and Voice.4 The approach by these workers to calculate free energies of transfer from their measurements on amalgam double cells is related to that outlined above except that Feakins employs the molal standard state (AG&$. The relation between the two free energies of transfer with respect to aquamolality or molality standard states can easily be shown to be where the second term arises from the ratio of molality to aquamolality.AG& = AG,",,,+2RT In 10 log(18.015/WA) (6) 640 590 520 450 340 I timelrnin FIG. 1 .-Transfer of an Orion Na+-responsive and an Orion Cl- responsive electrode between solutions 0.01 aquamolal in CsCl and increasing methanol content, cells 1-6. Fig, 2 compares results obtained with two types of cation-responsive glass electrode with the corrected results of Feakins and Voice for CsCf in methanol+ water mixtures. The results incorporated in table 3 are those obtained with the Orion Na+-responsive electrode. It is clear from fig. 2, that in high methanol content (x > 0.8) solutions the results using the E.I.L. K+-responsive glass electrode82 TRANSFER OF ALKALI METAL FLUORIDES are anomalous by comparison with both the Orion glass and the amalgam electrode results.The high potentials observed probably result from abstraction of water from the gel layer on the surface of the glass electrode, the thickness of which depends on the glass composition and may be thicker for the E.I.L. electrode and more sensitive to methanol than that of the Orion glass. A similar effect has been reported of MeOH 0 0.197 0.201 0.204 0.305 0.401 0.403 0.436 0.478 0.590 0.612 0.616 0.651 0.784 0.789 0.800 0.806 0.899 0.900 1 .ooo x 4 = CI 2 a G a 10 C!l TABLE 2.-vALUES OF AEt,xjmV FOR CELL (11) AND AQUAMOLAL FREE ENERGIES OF TRANSFER AGp(xl/kJ mol-1 FROM WATER TO MeOHf H20 mole fraction A@/ A G ~ ) I AGYl AGyl A G ~ I LiCl NaCl KCl RbCl CSCl AEtlmV kJ mol-1 AEtlmV kJ mol-1 AEt/mV kJ mol-1 AEt/mV kJ mol-1 AEtlmV kJ mol-1 0 0 0 0 47.9 4.72 31.1 3.11 89.1 8.82 60.8 6.12 125.0 12.29 83.8 8.48 0 0 0 0 0 43.8 4.33 47.9 4.72 67.3 6.65 62.3 6.17 84.2 8.35 107.8 10.68 101.8 10.10 121.8 12.13 140.2 13.95 130.3 12.99 167.4 16.66 164.0 16.34 155.2 15.28 107.7 10.86 134.6 13.68 178.8 18.00 182.1 18.09 178.5 17.97 197.2 19.78 187.5 18.72 I I I 1 Q2 OA 0.6 0.8 1-0 methanol mole fraction FIG.2.-Free energy of transfer for CsCl from water to methanol + water mixtures. - , amalgam electrode result^.^ + , Orion Na+-responsive electrode ; 0, x , E.I.L. K+-responsive electrodes.A . K . COVINGTON AND J . M. THAIN 83 previously. * Ivanovskaya, Gavrilova and Shults and Eisenman l 8 have studied certain electrode glass compositions in water, methanol +water and methanol solutions all saturated with NaCl or KCl.The Russian workers l7 refer to the specific effect of methanol on lithium aluminosilicate-based glass electrodes in solutions containing alkali metal ions not contained in the bulk electrode glass, and attributed the observed effects of 20-40 mV to diffusion potential effects in the glass. An alternative method of comparing the present results with those of Feakins and Voice is shown in fig. 3 as plots, against the mole fraction of methanol, of which is a deviation function from the straight line joining AG&( =0) and AG&= 11, a procedure which greatly magnifies the experimental differences. It can be seen from fig. 3, that up to x = 0.7, except for RbCl, all cells show agreement to within I - o x o x + + a r .I I 2 - 0 0 -0. - W rl I - 2 I 0.2 0.4 46 as methanol mole fraction LO LiCl NaCl 8 KCI RbCl CsCl FIG. 3.-Alternative method of comparison of results for methanol + water mixtures with those from amalgam cells in terms of a deviation function AG?(,) - X A G ~ ( ~ = l). 0, Amalgam result^.^ LiCl : +, 0, E.I.L. Na+-responsive electrodes ; x , Orion Na+-responsive electrode. NaCl : x , 0, E.I.L. Na+-responsive electrodes without Tris added to solutions ; 0, E.I.L. Na+-responsive electrode. KCl : 0, E.I.L. K+-responsive electrode ; x , Corning monovalent ion-responsive electrode. RbCl : 0, E.I.L. K+-responsive electrode ; x , Orion Na+-responsive electrode. CsCl : x , 0, E.I.L. K+- responsive electrode ; + , Orion Na+-responsive electrode.84 TRANSFER OF ALKALI METAL FLUORIDES + 5 mV.This is true over the whole mole fraction range if the anomalous results for E.I.L. K+-responsive electrodes are excluded (as has been done in table 2). These high values are observed for KCl, as well as for CsCl already referred to, and are greatly enhanced at low mole fractions for RbCl. In making this comparison with Feakins and Voice’s result^,^ it should be borne in mind that amalgam electrodes are reactive in aqueous solutions and that the e.m.f. results presented were obtained from plateau potentials held for only 10-20 min. Hence the agreement is judged satisfactory particularly since the present intention was to obtain results quickly of only moderate accuracy for comparison with spectroscopically-derived data.Attention should be drawn to points for NaCl in fig. 3 which were obtained by sub- stituting a conventional thermal electrolytic Ag I AgCl electrode for the Orion chloride-responsive ion-selective electrode ; these points confirm the correct function- ing of the latter in methanol + water mixtures. Transfer e.m.f. values for lithium, sodium and rubidium fluorides in methanol- water mixtures were obtained from measurements using cell (111) : alkali metal ion (M) F- responsive LaF, (111) responsive glass electrode 1 ion selective electrode. 78.3 87.9 7.82 8.89 TABLE 3.-\’ALUES OF AEt(,)/rnV FOR CELL (111) AND AQUAMOLAL FREE ENERGIES OF TRANSFER AG?(,)/kJ rnoP FROM WATER TO METHANOL+ WATER mole fraction LiF NaF RbF MeOH *E*(x)lmV A@) AEt(x)/mV 4) AEt(x)’mV A G g ) 0 0 0 0 0 0 0 0.201 53.8 5.89 0.212 48.3 4.77 0.305 88.1 8.66 0.403 107.2 10.57 0.410 94.9 9.39 0.456 0.478 134.0 13.21 0.503 121.6 12.02 0.603 157.7 15.59 135.3 13.42 0.612 161.1 15.92 0.651 166.0 16.43 0.789 207.1 20.49 0.825 177.1 17.60 0.837 116.9 11.85 185.2 18.43 0.900 199.0 19.82 1 .ooo 124.8 12.79 184.0 18.50 221.50 22.10 These and values of AG& obtained from eqn (5) are given in table 3.No previous data are available for comparison. For LiF, the E.I.L. and Orion Na+ responsive glass electrodes were used but for NaF only the former was available. For both salts, cells were well-behaved, reaching steady values on transfer quickly. For RbF, however, the results were less satisfactory, and only those obtained with Corning monovalent ion electrode were used in compiling table 3. E.I.L.K+-responsive electrodes again showed high values in the high methanol content solutions, but additionally the values were less reproducible and dependent on the time the electrodes were kept in high methanol content solutions. Table 4 shows tests for the additivity of single ion free energy of transfer values at four methanol mole fractions. Differences were taken between free energy values for pairs of salts and the discrepancies are expressed in terms of A, which should be zero. The new data meet the additivityA. K . COVINGTON AND J . M. THAIN 85 criterion to within - 3 kJ mol-l. Values in brackets in table 4 were derived using the amalgam cell data for chlorides of Feakins and Voice,4 which effects a slight improvement in the additivity test ( w 1 kJ mole-I). TABLE 4.-ADDITIVITY TESTS FOR FREE ENERGY OF TRANSFER DATA (AG/kJmOl-') FOR METHANOL AND WATER MIXTURES mole fraction MeOH 0.5 0.75 0.9 1 .o { NaCl-LiCl 3.7 (3.8) 3.35 (4.1) 4.2 (5.2) 4.1 (5.2) 1.85 (1.75) 4.15 (3.4) 2.1 (1.7) 2.5 (0.8) Na- Li NaF- LiF 5.55 7.5 6.9 6.6 NaCl-RbCl 0.3 (-3.9) -0.8 (-1.05) 1.1 (1.45) 0 (-1.6) Na- Rb NaF- RbF 2.15 1.4 1.9 - 2.4 1.85 (2.25) 2.2 (2.45) 0.8 (0.45) 2.4 (0.8) i l a rLiC1-RbCl -3.4 (-3.9) -5.15 (-5.95) 5.3 (6.65) -4.1 (6.8) 0 (0.5) 0.95 (0.15) 2.1 (0.75) 5.3 (2.6) Li-Rb -(LiFiRbF -3.4 - 6.1 7.4 - 9.4 LiCI- LiF - 1 .O 0.05 0.2 1.1 RbCl- RbF - 1 .O - 0.9 - 1.9 - 4.2 max A 1.75 3.15 2.7 5.3 Cl-F { NaCI- NaF - 2.75 - 3.1 -2.5 - 1.2 The preliminary study carried out in this laboratory using sodium-responsive glass and lanthanum fluoride F--responsive ion selective electrodes in aqueous hydrogen peroxide solutions of sodium fluoride l 2 has been extended to lithium, potassium and caesium fluoride solutions using cell (IV) alkali metal ion (M) I MF, ql_o;H;O, F- responsive LaF, (IV) responsive glass electrode For KF and CsF, ELL.K+-responsive electrodes and for LiF, E.I.L. Na+-responsive electrodes were used. Results are given in table 5 which includes those for sodium fluoride given previously l 2 in graphical form only. Cell e.m.f. values were steady after a few minutes even in the most concentrated hydrogen peroxide solutions. However, because of the attack of peroxide on the epoxy cement used to secure the lanthanum fluoride crystal in place, and the tendency to bubble formation on the electrode which caused variations in potential, the electrodes were left in the cells for a minimum possible time contingent upon reliable measurements being achieved. ion selective electrode.COMPARISON WITH SPECTROSCOPICALLY DERIVED DATA Elsewhere l9 it has been shown that the free energy of transfer AGFx) can be related to a free energy of preferential solvation AGE obtainable from n.m.r. or U.V. studies of the ions of the salt in the mixed solvent systems. For isodielectric solvent systems such as hydrogen peroxide + water mixtures these quantities were shown to be identical, i.e. AG& = AGg. Values of AGg for Li+, Na+, Cs+ and F- have been determined from n.m.r. measu~ements.~~ By combination of these values to obtain free energies of transfer for the appropriate salts, direct comparison is possible with the e.m.f.-derived data presented in table 5.This comparison has already been made elsewhere for NaF and is shown for LiF and CsF in fig. 4. As with NaF, the agreement is remarkably good and within 1 kJ mol-l in the free energy of transfer. For LiF, the agreement is particularly satisfying since Li+ is preferentially solvated by86 TRANSFER OF ALKALI METAL FLUORIDES water and F- by peroxide so there is a subtraction of contributions leading to a very small negative free energy of transfer as shown in fig. 4, and this system provides an exacting test of the theory. No. n.m.r. data are available for 39K chemical shifts in this mixed solvent system so comparison is not possible.TABLE 5.-vALUES OF AEt(x)/mV FOR CELL (IV) AND AQUAMOLAL FREE ENERGIES OF TRANSFER AGTx,/kJ mol-1 FROM H,O TO Hz02+ HzO LiF NaF KF CsF mole fraction - A G a -AGg)/ - 4 1 -A& of H 2 0 2 -AEt(x)/mV kJ mol-1 -AEt(x) /mV kJ mol-1 -AEt(&nV kJ mol-1 -AEt(x)/mV kJ mol-1 0 0.0310 0.0486 0.0909 0.125 0.133 0.140 0.180 0.282 0.322 0.330 0.376 0.404 0.416 0.427 0.484 0.554 0.569 0.584 0.730 0.753 0 0 0 0 7.6 0.73 13.5 1.30 19.7 1.90 19.3 1.86 0 0 0 0 52.7 5.09 34.2 3.3 8.1 0.78 28.4 2.74 35.3 3.41 17.8 1.72 22.7 2.19 43.9 4.24 101.5 9.79 76.7 7.4 113.8 10.99 97.7 9.427 54.8 5.29 28.0 2.19 42.0 4.05 70.0 6.75 139.8 13.494 155.6 15.086 126.2 12.177 149.2 14.396 I I I I 2 0 0>2 0.4 0.6 0.8 1.0 mole fraction peroxide FIG.4.-Free energy of transfer for CsF (0) and LiF (0) from water to hydrogen peroxide + water mixtures. Comparison of the e.m.f. data for methanol+water mixtures with n.m.r. data which requires an estimation of an electrostatic contribution in addition to the free energy of preferential solvation, in order to calculate the free energy of transfer, will be postponed until a later paper.A. K . COVINGTON AND J . M. THAIN a7 CONCLUSIONS It has been shown that ion-selective electrodes can be used to obtain free energies of transfer of alkali metal halides from water to methanol or to hydrogen peroxide+ water solutions, provided certain precautions are observed. It is necessary to use several cation-responsive glass electrodes of different glass compositions in order to ensure that no specific solvent effects are present.Few studies have been made of the response of these glass electrodes in solutions of pure electrolyte, aqueous or mixed aqueous, although selectivity studies have been reported.l* In the absence of a direct comparison with an established electrode responsive to the same ion in the same solvent mixtures, or of activity coefficient data, it is by no means easy to demonstrate that theoretical response is obtained in mixed solvent systems. However, if no hysteresis is shown when the electrodes are subjected to change of solvent composition and several different electrodes yield substantially similar results then confidence in the results is instilled. With these precautions and safeguards the method is useful for obtaining results of moderate accuracy (& 1 mV).We thank S.R.C. for a research studentship (to J. M. T.). The preliminary measurements on NaF+H,O,+H,O were made by Mr. Michael Wood. An account of this work was presented at the IUPAC-sponsored symposium on ion- selective electrodes held in Cardiff in April 1973. 0. Popovych, Crit. Rev. Anal. Chem., 1970,1,73. R. G. Bates, Determination ofpH(Wiley, New York, 2nd edn., 1973), pp. 211-234. G. Akerlof, J. Amer. Chem. Soc., 1930, 52, 5353. D. Feakins and P. J. Voice, J.C.S. Faraday I, 1972, 68, 1390. B. M. Lowe and D. G. Smith, Chem. Comm., 1972,989. M. S. Frant and J. W. Ross, Science, 1966,154,1553. A. K. Covington and J. M. Thain, J. Chem. Educ., 1972,49,554. K. Srinivasan and G. A. Rechnitz, Anal. Chem., 1968,40,509. J. N. Butler and R. Huston, Anal. Chem., 1970,42,1308. W. E. Bazzelle, Anal. Chim. Acta, 1971, 54, 29. lo J. J. Lingane, Anal. Chem., 1968,40,935. l2 A. K. Covington, K. E. Newman and M. Wood, Chem. Comm., 1972,1234. l 3 A. G. Mitchell and W. F. K. Wynne-Jones, Trans. Faraday SOC., 1955,51,1690; 1956,52,824. l4 A. K. Covington, K. E. Newman and T. H. Lilley, J.C.S. Farahy I, 1973,69,973. l5 R. A. Robinson and R. H. Stokes, ElectroZyte Solutions (Butterworth, London, 2nd. rev. edn., l6 A. K. Covington and T. Dickinson in Physical Chemistry of Organic Solvent Sys?ems (Plenum, l7 I. S. Ivanovskaya, V. I. Gavrilova and M. M. Shults, Souiet Electrochem., 1970, 6, 975. l 8 G. Eisenman in Aduances in Analytical Chemistry and Instrumentation (Wiley, New York, 1965), l9 A. K. Covington, T. H. Lilley, K. E. Newman and G. A. Porthouse, J.C.S. Faraday I, 1970), p. 30. London, 1973), p. 18. vol. 4, pp. 295-305 (reprinted in The Glass Electrode, Interscience Reprint, 1966). 1973, 69, 963.

ISSN:0300-9599    

DOI:10.1039/F19757100078

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Thermodynamics of Ionization of Amino Acids Part 6.t-The Second Ionization Constants of Some Glycine Peptides BY (THE LATE) EDWARD J. KING Department of Chemistry, Barnard College, Columbia University, New York, U.S.A. Received 30th May, 1974 Thermodynamic functions for the ionization of the ammonium groups of glycylglycine, glycyl-DL- alanine, glycyl-DL-a-amino-n-butyric acid, and glycyl-L-leucine were obtained from e.m.f. measure- ments at temperatures from 5 to 50°C on cells without liquid junction. The ionization functions of the carboxyl groups of six glycine peptides have been The ionization of the ammonium group reported in an earlier paper of this series.l +NHSCH2CONHCH(R)COO-+ HzO+NH2CH,CONHCH(R)COO-+ H30f has now been investigated for four peptides, namely those in which the substituent group R is hydrogen, methyl, ethyl or isobutyl.The ionization constants of the ammonium group have the form K2 = [H,O+][A-]/[HA*] where HA* represents the dipolar form of the peptide and A- the anionic form. The second ionization constants were obtained from measurements of the e.m.f. of cells of the type Pt, H,IHA* (mJ, NaA (m,), NaCl (m,)lAgCl, Ag (1) where m,, m2 and m3 are the stoichiometric molalities. Twenty-five buffer solutions were used for determination of the constants of glycylglycine. These were divided into three groups having stoichiometric buffer ratios, p = m,/m,, of approximately 2, 1, and 0.5. With each of the other peptides eight buffer solutions, all with p = 1, were used. Measurements were made at ten temperatures from 5 to 50°C in order to obtain accurate values of the enthalpies and entropies of ionization.EXPERIMENTAL Glycylglycine, glycylalanine, and glycyl-a-amino-n-butyric acid were the same prepara- tions as used in the earlier w0rk.l However, some of the earIier supply of glycyl-L-leucine [2-(N-glycyl)amino-4-methylpentanoic acid] was supplemented by a new sample of glycine- free material from the same source (H.M. Chemical Co., Santa Monica, California). The new material was found to be 99.94 % pure by form01 titration and contained less than 0.004 % each of ammonia, chloride, phosphate, iron and heavy metals. Sodium chloride containing less than 0.001 % bromide was prepared and fused according to the directions of Pinching and Bates.2 Carbonate-free sodium hydroxide solution was prepared and standardized against National Bureau of Standards potassium acid phthalate as in earlier Stock -f Part 5, E.J. King, J. Amer. Chem. SOC., 1960, 82, 3575. 88E. J. KING 89 buffer solutions were made from weighed amounts of peptide, sodium chloride, sodium hydroxide and water. These stock solutions were diluted further to give the cell solutions which were boiled under reduced pressure and swept with hydrogen. Buoyancy correction for air or hydrogen was made as required. During preparation of the solutions and filling of the cells, exposure of the solutions to air was kept to a minimum to avoid contamination but none was found. The test was carried out by mixing about 5 cm3 of buffer solution with 3 cm3 each of 1.0 mol dnr3 barium chloride solution and 0.2 mol dm-3 carbonate-free, sodium hydroxide solution.The appearance of the mixture after 5-10 min was compared with turbidities developed from known amounts of carbonate. Solutions were viewed in a blackened comparator block and as little as 3 x mol dm-3 could sometimes be detected though the limit of detection was usually about twice this. Measurements over the temperature range required about ten hours and were made in the following sequence : 25, 20, 15, 10, 5, 30, 35, 40,45,50 and 25°C. An indication of the performance of the cells is given by comparing the initial and h a 1 readings at 25°C. The standard deviations between initial and final values were 0.12 mV for glycylalanine, 0.063 mV for glycyl-a-amino-n-butyric acid, and 0.065 mV for glycylleucine.For glycylglycine buffers with ionic strengths above 0.02 mol kg-' the standard deviation was 0.12 mV. The performance of cells containing more dilute glycyl- glycine buffers was generally much poorer, the standard deviation in EZ5 for such cells being 0.21 mV. A few results obtained from these dilute buffers appeared to be low and were neglected, as discussed later, in calculating the ionization constant of glycylglycine. Smoothed values of the e.m.f. corrected to 1 atm hydrogen gas partial pressure are given as parameters of the equation, where t is in "C, The apparatus has been described Et = E25+ ~(t-25)+ b(t- 25)' in table 1. The standard deviations between observed and calculated values were & 0.085 mV for glycylglycine, f 0.035 mV for glycylalanine, & 0.056 mV for glycylaminobutyric acid, and k0.026 mV for glycylleucine.Actual experimental values were used in the calculations. TABLE PARAMETERS OF THE EQUATION E' = EZ5+ a(t- 25)+ b(t- 25)' 10 3ml/moi kg-1 2.572 4.967 6.185 7.226 10.186 15.107 19.34 30.84 5.270 8.243 9.638 12.131 15.819 18.560 23.60 32.58 40.51 53.56 103ma/mol kg-' lo3 m3/mol kg-' E25/V glycylglycine ( p = 4, m1-4m2-4m3) 5.028 9.710 12.240 14.126 20.16 29.90 3 8.27 61.03 5.067 9.787 12.283 14.237 20.23 30.00 38.41 61.25 0.863 98 0.847 61 0,842 71 0.838 37 0.830 29 0.820 69 0.814 83 0.803 97 glycylglycine ( p = 1, ml = m2 = m3) 5.288 8.273 10.005 11 -893 15.507 18.627 23.68 32.58 40.51 53.56 5.286 8.269 10.004 12.01 3 15.664 18.618 23.67 32.63 40.56 53.63 0.846 17 0.835 12 0.831 17 0.825 37 0.818 73 0.815 10 0.809 27 0.801 51 0.796 37 0.789 89 - 10k/V K-' - 10Eb/V K-' -5 + 52 73 85 116 148 172 21 3 59 96 110 133 154 166 188 216 235 258 245 235 245 240 245 250 245 250 245 240 235 235 240 240 240 240 240 24090 IONIZATION OF AMINO ACIDS 103m1/mol kg-1 10.566 18.99 40.29 51.32 68.28 89.29 109.00 6.880 10.462 15.322 20.21 24.25 29.44 34.55 39.60 10.272 11.041 16.256 20.05 23.76 29.53 35.46 40.54 7.986 10.191 15.004 17.98 24.48 29.88 34.83 40.19 TABLE 1-continued 103rn2/mol kg-1 lo3 m3/mOl kg-1 E2 5/v - 106a/V K-' - 106b/V K-* glycylglycine (p = 2, ml = 2m2 = 2m3) 5.176 5.212 0.828 24 9.301 9.708 0.812 91 19.74 20.60 0.794 54 25.14 26.24 0.788 77 33.63 43.09 0.776 86 43.98 56.34 0.770 72 53.69 68.78 0.766 26 glycyl-~~-alanine (p = 1) 6.938 10,519 15.453 20.32 24.38 29.50 34.63 39.69 6.909 10.490 15.387 20.26 24.32 29.47 34.59 39.65 0.843 82 0.833 33 0.824 07 0.817 15 0.812 74 0.808 09 0.814 25 0.801 02 glycyl-DL-a-amino-n-butyric acid (p = 1) 10.235 11.086 16.320 19.98 23.67 29.26 35.43 40.1 6 7.913 10.070 14.867 17.82 24.19 29.4 1 34.28 39.56 10.252 11.062 16.286 20.02 23.71 29.39 35.29 40.35 0.833 59 0.831 89 0.822 36 0.817 23 0,813 03 0.807 72 0.803 34 0.800 15 glycyl-L-leucine ( p = 1) 7.950 10.129 14.936 17.90 24.33 29.64 34.56 39.87 0.839 55 0.833 51 0.823 95 0.819 62 0.812 01 0.807 15 0.803 48 0.800 09 113 176 236 261 298 327 344 110 146 178 202 220 236 248 26 1 154 161 192 209 226 244 258 27 1 115 140 170 185 21 5 23 1 244 256 250 255 240 240 240 225 225 236 225 248 225 225 235 240 230 245 225 225 245 250 240 250 240 235 258 235 245 235 245 245 245 CALCULATIONS AND RESULTS From the experimental data the quantity -log YYIHYHYC~, can be calculated from -log mHyHyC1- = (E - E")/k +log m3 (1) where k = (RT/F)ln 10 and E" is the standard e.m.f.s* This function is related to PK2 byE.J. KING 91 In eqn (2) the hydroxyl ion concentration, m&, which is very small in comparison with m, and m2 is given with sufficient accuracy by (3) The activity coefficient term in eqn (2) should depend on both the ionic strength (Z = m2 +m3) and the dipolar ion concentration (m,). Alternatively, the inde- pendent variables can be chosen as Z and p = m /m2 since Z is proportional to m, for the buffer solutions used in this investigation. If the behaviour of glycine+ potassium chloride solutions is used as a guide, the activity coefficient term in eqn (2) should be expressed by -log mbH = pK, +log mHyH.and taurine +hydrochloric acid solutions (4) Y C l - log ---?HA * = (a + a’p)l+ (p f p’p)12 + (6p + 8’p2)I2 YA - where the Greek letters represent constants. In dilute solutions only the first term on the right-hand side of eqn (4) should be important, with the result that A plot of pKi against Z will thus give pK2 at Z = 0 and have a slope which varies with the buffer ratio. Such graphs for glycylglycine at 25°C are illustrated in fig. 1 and those for the other peptides in fig. 2. If the unreliable results below Z = 0.02 are discounted, the plots are straight lines up to Z of about 0.06. The curvature at higher pK$ pK2+(a+a’p)Z.(5) 8.34 I I I I I 0 0 O / 4 8.30 - -- 0 /e O 0 826 - I I I I I 20 40 60 80 I00 103ml /mol kg-’ FIG. 1 .-Extrapolation of glycylglycine results at 25°C against acid molality (ml) according to eqn(6): $ , p = & ; @ , p = l ; @ , p = l ; O , p = 2 .92 IONIZATION OF AMINO ACIDS concentrations is to be expected because of neglect of higher terms of eqn (4). Alter- natively, one could choose a in eqn ( 5 ) by trial so that points of a plot of (pK5 - aZ) against m, all fall on the same curve irrespective of the buffer ratio. At 25°C for glycylglycine, the expression fits all the data except the three points at high ionic strengths. pK; -0.3801 = pK, + f ( m l ) (6) I I I I I I I t 8.28 I I I I I 40 6 0 80 1 0 0 I 2 0 8.24 x ) 1 031/mol kg-l FIG.2.-Extrapolation of glycylglycine results at 25°C using experimental values for activity co- efficients according to eqn (9) : o, p = 3 ; a, p = 1 ; 0, p = 2. A third method of treatment of the data for glycylglycine is possible utilizing data and in aqueous sodium * (NaC1). Combining these data, which are presented in the for the activity coefficient of glycylglycine YHAi chloride solutions,' form of polynomials, for the activity coefficient, then in water log YHA* (NaCl) - - -0.23296~~,+0.17976~~: -0.5216~~3 -0.3357mIms +0.7723~$ (7) YHA* (0) where terms in higher powers than rn: have not been included because they are negli- gible for the present purposes. Assuming log yA-/yCl- in eqn (2) is proportional to I, and any salting effect of A- on HA' is proportional to m, (and hence to I ) , it is therefore reasonable to define a new extrapolation function pK; -= pKL +log = pK2 +BI (8) YHA * (NaCl) YHA'(0) where B is an adjustable parameter.A plot of eqn (8) against Zfor the results at 25°C is shown in fig. 3. It may be noted that results for I = 0.02 remain low but a straight line can be drawn through all the remaining points for the three buffer ratios to give a value of pK, in close agreement with the value obtained using eqn (5).E. J. KING 93 in determining the ionization constants of glycine and its Datta and Grzybowski N-methyl derivatives used the relation This is similar to that derivable from eqn (4) except that B and C were not expressed as functions of the buffer ratio since Datta and Grzybowski assumed the activity co- efficient of the dipolar ions was unity.In applying this relation to their data and to the data of King for glycine they gave equal weight to all points including those at PK; = pK2-BI-CI%. (9) 8.36 8.34 8.32 8.3, 8.3 2 0 40 60 80 1 031/mol kg- FIG. 3.-Extrapolation of results at 25°C for glycylalanine (0), glycyl-a-amino-n-butyric acid ( 0 ) and glycylleucine (0). the lowest ionic strengths. Since pK, values at ionic strengths below 0.02 tend to be low, to include these points makes the extrapolations show pronounced curvatures at low ionic strengths instead of the simple linear behaviour. If the lowest point in their extrapolations (in their fig. 1) is ignored, good straight lines could be drawn through the rest of the points up to I = 0.05.Above this ionic strength, the lines curve over as do those for glycylglycine (fig. 1). This is demonstrated using extra- polation based on eqn (5) in fig. 4. The failure to obtain useful results from cells containing very dilute buffers may be94 IONIZATION OF AMINO ACIDS caused by a number of sources of error which are trivial for more concentrated solu- tions. The silver-silver chloride electrodes are porous and may trap water or oxygen which are not removed by mixing when the cells are filled. Traces of carbonate could alter the buffer ratio. At the pH values of these buffers about 99 % of the carbonate would be converted to hydrogen carbonate ion. Then mHA* = m1-k mbH + mHc0; m A - = nz2 - mbH - mHCOj and the term log[(m, +mbH)/(m2 - m 3 ] in eqn (2) would be too small, making pK4 too small.Yet, to account for the errors in pK4 in very dilute solutions would require about 5 x mol dm-3 HCO:, and even one-tenth of this could not be detected in the cell solutions. It is readily shown also that corrections for the solu- bility of silver chloride in the buffer solutions l * l 2 are entirely negligible. Another possibility would be the malfunctioning of the hydrogen gas electrode in these very dilute buffer solutions of poor buffer capacity. Further investigations would be necessary to decide whether the effect is a real one, or an artefact of the electrodes in these solutions. 1 I I I 20 40 6 0 80 1031/mol kg-' x , ref. (11). FIG. 4.-Comparison of extrapolation of results for glycine : 0, ref. (4) ( p > 1) ; @, ref.(4) ( p < 1) ; The values of pK2 obtained by graphical extrapolation are summarized in table 2. Those for glycylglycine are the averages of the three values obtained by independent extrapolations of results for the three buffer ratios (fig. 1). The deviations in pKz listed in table 2 give an indication of the average precision of the extrapolations. Table 2 also includes the parameters A, B, and G of the Harned-Robinson equation l3 pK, = (A/T)-B+CT. (10)E. J. KING 95 From these parameters the changes in various thermodynamic properties for the ionization reaction in the standard state, can be calculated. Values at 25°C are given in table 3 for AGq, the change in free energy; AH;, the change in enthalpy or the heat of ionization; AS;, the change in entropy.TABLE 2.-vALUES OF pK2 FOR GLYCINE PEPTIDES AND THE PARAMETERS OF EQN (8) glycyl-a-amino-n- temp./K glycylglycine glycylalanine butyric acid glycy'leucine 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 dev. AIK B C1K-l 8.813 8.668 8.529 8.394 8.265 8.140 8.018 7.901 7.788 7.680 k0.002 2 412.46 0.316 3 0.001 6420 8.896 8.746 8.602 8.462 8.331 8.202 8.077 7.958 7.843 7.732 f 0.001 2 580.60 1.168 2 0.002 828 4 8.899 8.748 8.604 8.464 8.331 8.200 8.073 7.952 7.837 7.725 + 0.002 2 563.10 0.985 6 0.002 410 1 8.887 8.740 8.597 8.459 8.327 8.199 8.074 7.955 7.840 7.729 *O.Ool 2 452.31 0.354 3 0.001 5298 The values of the ionization constants can be compared with the few values available in the literature. Neuberger l4 using cells with liquid junction and a double extrapolation, obtained for glycylglycine pK2 = 8.265 at 25"C, in precise agreement with that recorded in table 2.Smith and Smith l 5 obtained a lower value, pK2 = 8.255. They used the same method as the present investigation, but they measured the electromotive forces of only five cells to the nearest 0.1 mV and did not exclude oxygen from the solutions. Enthalpies and entropies of ionization derived from their data are also not in agreement with those reported here. Of numerous less precise determinations of ionization constants by potentiometric titration that of Monk is typical. The only values His value of pK2 for glycylglycine at 25°C is 8.23 4 0.01. TABLE 3.-THERMODYNAMIC PROPERTIES FOR THE IONIZATION OF NH3CH2CONHCH(R)COO- AT 298.15 K AG"/kcal mol-1 AH"/kcal mol-1 AS"/cal mol-1 K-1 glycine N-gl ycylgl ycine alanine 21 N-gl ycyl alan r ne a-amino-n-butyric acid N-gl ycyl-a-amino-n- bu t yric acid leucine 21 N-gl yc ylleucine N-methylglycine NN-dimethylglycine bicine a* 22 tricine b* 22 13 338 11 275 13 460 11 365 13 410 11 364 13 295 11 360 13 916 13 561 11 369 11 099 10 550 10 370 10 980 10 660 10 696 10 750 10 815 10 600 9 681 7 654 6 279 7 520 - 9.36 - 3.03 - 8.32 -2.37 -9.10 - 2.07 - 8.32 -2.55 - 14.20 - 19.81 - 17.07 - 12.00 u bicine is NN-di-(2-hydroxyethy1)glycine ; b tricine is N-tris(hydr oxymethy1)methylglycine.96 IONIZATION OF AMINO ACIDS available for glycylalanine and glycylleucine pK, were also determined by potentio- metric titration.Perkins l7 reports 8.21 and 8.17 respectively at 25°C.Ellenbogen l8 found 8.23 for glycylalanine. Salakhutdinov et aL1’ give 8.28k0.01 at 25°C and Dobbie and Kermack 2o report 8.41 at 20°C for glycylleucine. All of these values are lower than those reported in table 2. Table 3 gives a comparison of AG;, AH;, and AS; at 25°C for each of the four N-glycyl substituted amino acids and corresponding data for the parent unsubstituted amino acids.4* 21 The effect of N-glycyl substitution on AGZ is to lower it by between 1935 and 2095 cal mol-l, 2035 cal mol-l on average. The effect on AH; is much smaller with an average diminution of 165 cal mo1-l with a small increase of 54 cal mo1-1 for N-glycyl-a-amino-n-butyric acid. Consequently, changes in AS; on N-glycyl substitution depend mainly on changes in AC;.The average value of AS; is - 8.78 cal mol-1 K-I for the substituted acids representing a significant difference. The effect of N-glycyl substitution is almost independent of the nature of the R substituent for the four amino acids now investigated. However, with R = H, much more variable effects are found for substituents other than N-glycyl. Table 3 illustrates this for four different substituents.ll9 22* 23 There is some variation between the AG; values but very marked variations in the AH; values and consequently the AS; values. The experimental work was initiated at Barnard College in 1959 supported by a research grant from the National Heart Institute of the National Institute of Health, U.S. Public Health Service. This paper was prepared from Prof.King’s incomplete manuscript and notes by Dr. A. K. Covington in collaboration with Dr. Grace W. King and Prof. R. A. Robinson of the Department of Physical Chemistry, University of Newcastle upon Tyne, United Kingdom, where Prof. King was spending study leave at the time of his death. E. J. King, J. Amer. Chem. SOC., 1957, 79, 6151. G. Pinching and R. G. Bates, J. Res. Nat. Bur. Stand., 1946, 37, 311. E. J. King, J. Amer. Chem. SOC., 1954, 76, 1006. E. J. King and G. W. King, J. Amer. Chem. SOC., 1956,78,1089. E. J. King, J. Amer. Chem. SOC., 1956, 78, 6020. E. J. King, J. Amer. Chem. SOC., 1953, 75, 2204. H. D. Ellerton, G. Reinfelds, D. E. Mulcahy and P. J. Dunlop, J. Phys. Chem., 1964, 68,398. 4E. J. King, J. Amer. Chem. SOC., 1951, 73, 155. ’ R. M. Roberts and J. G. Kirkwood, J. Amer. Chem. SOC., 1941, 63, 1373. lo E. E. Schrier and R. A. Robinson, J. Biol. Chem., 1971, 246, 1179. l 1 S. P. Datta and A. K. Grzybowski, Trans. Faraduy SOC., 1958, 54, 1179. l2 C. B. Monk, Trans. Faraday SOC., 1951, 47,292. l3 H. S. Harned and R. A. Robinson, Trans. Faraday SOC., 1940, 36, 973. l4 A. Neuberger, Proc. Roy. SOC. A, 1937, 158, 68. l5 E. R. B. Smith and P. K. Smith, J. Biol. Chem., 1942, 146, 157. l 6 C. B. Monk, Trans. Faraday SOC., 1951, 47, 285. l7 D. J. Perkins, Biochem. J., 1954, 57, 702. l8 E. Ellenbogen, J. Amer. Chem. SOC., 1952, 74, 5198. l9 U. I. Salakhutdinov, A. P. Borisova, Y. V. Granovskii, I. A. Savich and V. I. Spitsyn, Dokludy 2o H. Dobbie and W. 0. Kermack, Biochem. J., 1955,59,246. 21 P. K. Smith, A. C. Taylor and E. R. B. Smith, J. Biol. Chem., 1937, 122, 109. 22 S. P. Datta, A. K. Grzybowski and R. G . Bates, J. Phys. Chem., 1962, 68, 275. 23 R. N. Roy, R. A. Robinson and R. G. Bates, J. Amer. Chem. SOC., 1973, 95,2831. Akad. Nauk S.S.S.R., 1967, 177, 365.

ISSN:0300-9599    

DOI:10.1039/F19757100088

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Changes in the Sieving Action and Thermal Stability of Zeolite A produced by Ion-exchange BY T. TAKAISHI" Institute for Atomic Energy, Rikkyo (St. Paul's) University, Yokosuka, Japan 240-01 AND Y. YATSURUGI, A. YUSA AND T. KURATOMI Komatsu Electronic Metals Co., Shinomiya, Hiratsuka City, Japan 254 Received 8th April, 1974 Na-A and K-A zeolites have been ion-exchanged with Ca2+ and Zn'f. The molecular sieving properties of the ion-exchanged zeolites were measured as functions of their composition. Abrupt changes in sieving action occur at 75 % exchange in the (K, Ca)-A series, and at 75 and 83 % exchange in the (K, Zn)-A series, as measured with chemically unreactive molecules. The sieving action of (K, Zn)-A zeolite for some polar molecules and unsaturated hydrocarbons changes gradually with its composition.The thermal stability of melocular sieve 3A is greatly improved by introducing a large number of divalent cations into this K-A zeolite. These properties are discussed in terms of site-selectivities of the exchangeable cations. From these, are derived diagrams of cation distribution on three kinds of site in the crystal. The ideal composition of Na-A zeolite, per unit cell, is Na 2( A102) ( SO2) *xH, 0. * According to recent X-ray structural analysis studies of dehydrated Na-A zeolite,2 three kinds of site are available to the twelve sodium ions. The first is near the centre of the six-membered oxygen ring and called the P-site, since the six-membered oxygen ring constitutes a window for the B-cage. The second is near the centre of the eight- membered oxygen ring which lies on a {loo) face of the unit cube. This site is named the a-site, since the eight-membered oxygen ring is a window for the central cavity (a-cage).The third is in the central cavity, near the centre of the four-membered oxygen ring, the y-site. Each of the (100) faces of the unit cube contains four a-sites, but each face contains only one cation ; that is, three of the twelve a-sites in a unit cell are occupied by Na+. On the other hand, the eight p-sites in a unit cell are occupied by eight Na+ ions. The twelfth Na+ ion in the unit cell occupies one of the twelve y-sites, to which the affinity of a cation is very weak. Sites a and B are preserved in hydrated zeolite A, although their coordinates are slightly changed.In contrast, y-sites cannot be defined in the hydrated crystal; that is, the position of the twelfth Na+ cannot be 10cated.~ The molecular sieving property of a dehydrated zeolite A is primarily determined by the kind of ion occupying the a-sites. Rees and Berry4 studied the blocking effect of pre-sorbed NH3 which is localized upon a cation on an a-site, and concluded that K+ in (Na, K)-A Ions on a-sites partially block the window aperture to the a-cage. 1-4 9798 SIEVING ACTION OF ZEOLITE A preferentially occupies the a-site. However, most existing data on ion-exchanged zeolite A concern those containing Na+ ions. In view of the difference in site select- ivities of Na+ and K+, more systematic studies of the distribution of cations on a-sites are desirable to obtain insight into the chemistry of zeolite A.In the present work, we have studied the differences between Na-A and K-A zeolites when they are ion- exchanged with divalent cations such as calcium and zinc ions. EXPERIMENTAL CHEMICAL ANALYSIS Zeolite A was dissolved in 2 equiv. dnr3 HCI, and A1 and Si contents were determined by the usual method^.^ The alkali metal ion content was measured with an atomic absorption spectrometer. The content of divalent cation in the ion-exchanging solution was determined by ethylenediaminetetra-acetic acid titration, while that in the ion-exchanged zeolite was, in most cases, calculated from the contents in the solution and material balance. In some cases, however, direct analysis was performed with an atomic absorption spectrometer.Radio- active tracers were also used to obtain greater accuracies for low concentrations of the species in solution. Tracers were obtained by irradiating reagents with the TRIGA-I1 reactor of our institute. MATERIAL The starting material was Linde molecular sieve 4A (lot. no. 480017) powder. After the standard treatment of the sample,6 its composition was determined. The sum of the com- ponents amounted to 98.8 % of the original sample. The formula was 0.94Na20-1.00 Al2O3=1.86 Si02*4.50 H20. The Na20/AI2O3 ratio of less than 1 may be ascribed partially to protons, which had replaced alkali metal ions and were not analyzed. The composition of a real crystal of Na-A zeolite differs from that of an ideal one and is 6~ 1 in our crystal. All chemical reagents used were S.P.grade and common adsorbate gases contained in glass cylinders (Takachiho Chemical Co.) had nominal purities better than 99.9 % as determined by mass spectroscopic analysis. Less common gases, such as SiH4, PH3, ASH,, and B2H6, were obtained from the Nihon Oxygen Co., and were of the special high purity used in the semiconductor industry. Na12(A102)12(Si02)12(NaA102)~ with 0 < 6 < 1.6 ADSORPTION MEASUREMENTS Adsorption measurements were made with a McBain type quartz spring balance which had a sensitivity of 0.5 m 8-l and an accuracy of +O.l mg. Prior to adsorption experiments, adsorbents were baked at 400°C for 5 h under a vacuum pressure lower than 1 x N m-2. RESULTS K-A AND Na-A ZEOLITES EXCHANGED WITH Zn2+ The ion exchange was carried out at 80+0.5"C with solutions of various composi- tions all of 0.2 total metal ion normality.Solutions of ZnCl, + NaCl and ZnC1, + KCl were used for Na-A and K-A, respectively. Although the reaction proceeded very rapidly,' the exchange was carried out for 20 h in order to obtain a true equilibrium. The forward exchange (Na+ or K+-+Zn2+) and the reverse exchange (Zn2++Na+ or K+) were measured to check the attainment of true equilibrium and the reversibility of the exchange reaction. After the reaction, the zeolite was filtered off and the Na, K and Zn contents in the solid and liquid phases were determined. The ion-exchange isotherms obtained are shown in fig. 1, in which (Zn), and (Zn), denote the equivalentT. TAKAISHI, Y . YATSURUGI, A. YUSA AND T.KURATOMI 99 fraction of zinc ion (2[Zn]/(2[Zn] + "a]) or 2[Zn]/(2[Zn] + [K])) in zeolite and in liquid, respectively. The isotherms show a steep rise up to 66 % exchange, and then become rather flat. The isotherm of (K, Zn)-A has a distinct step, while that of (Na, Zn)-A has the usual form. This difference may be ascribed, in part, to the differ- ence in the site-selectivities of potassium and sodium ions. After heating the zinc- exchanged zeolites in air at 400"C, their sorptive capacities and X-ray diffraction patterns were compared with those of unheated ones, and no change was introduced by such treatment. This indicates that the zinoexchanged zeolites are stable, in contrast to the earlier conclusion of Barrer and Meier.6 FIG. 1.-The ion-exchange isotherms for zeolite A at 0.2 totaf metal ion normality Zn/Na-A system ; (0) Zn/K-A system.and 80°C: (A) First, the sorptive abilities of the above zeolites for non-polar gases were studied. The rate of sorption was not studied in detail, and only the amounts of sorption at equilibrium are discussed herein. Sorption was measured at fixed temperatures and pressures for : nitrogen and argon at - 195°C ; carbon dioxide, n-butane, silane and diborane at 0°C. Curves of amount sorbed against composition of zeolite are shown in fig. 2, in which a distinct difference between the two series (Zn, Na)-A and (Zn, K)-A is apparent. The curves for the (Zn, Na)-A series for nitrogen and n-butane sorption show steep rises at 33 % exchange in a manner similar to (Ca, Na)-A systems.l* In the (Zn, K)-A series, on the other hand, the curve for n-butane rises sharply at 75 % exchange, and those for nitrogen and argon at 83 %.The window size of (Zn, K)-A zeolite with a composition 0.75 < (Zn), .c 0.85 may be larger than that of Na-A zeolite but smaller than that of Ca-A, since the Zn-exchanged zeolites sorb n-butane but do not sorb argon or nitrogen. Hence, we shall tentatively call these 4.5A. We must not conclude, from the above sieving action of 4SA, that the molecular diameters of argon and nitrogen are larger than that of n-butane, because the cation on the a-site can be displaced from its equilibrium position to widen the window for a visiting gaseous molecule. Such cation displacement, however, has an activation energy and does not easily occur at liquid nitrogen temperatures at which sorption of nitrogen and argon was measured.8 The steep rise is also seen in diborane adsorption at 83 % exchange in the (Zn, K)-A series, at 0°C.This result is reasonable, since diborane has a bridged structure and its molecular diameter is larger than that of ethane.100 SIEVING ACTION OF ZEOLITE A 0 0.2- 0.4 0.6 6.8 1.0 - 0 0.2 0.4 0.6 0.8 1.0 (4 (b) znz FIG. 2.-(a) Adsorption of non-polar gases on (Zn, K)-A zeolites. (A) nitrogen, 200 Torr, - 195°C ; (m argon, 50 Torr, - 195°C ; (0) carbon dioxide, 210 Torr, 0°C ; (e) n-butane, 200 Torr, 0°C ; (0) monosilane, 160 Tom, 0°C ; (m) diborane, 50 Torr, 0°C. (6) Adsorption of non-polar gases on (Zn, Na)-A zeolites. (A) nitrogen, 200 Torr, - 195°C ; (0) carbon dioxide, 210 Torr, 0°C ; (e) n-butane, 220 Torr, 0°C ; (0) monosilane, 160 Torr, 0°C.1 Torr = (101 325/760) N m-2. 0.20 0 c) .- c( 0.15 B c1 (d 0 (d .C) c, =!! 0.10 d 0 0 a .- c, .C( 0.05 .* c) a 5 4 0 znz FIG. 3.-Adsorption of polar gases and unsaturated hydrocarbons on (Zn, K)-A zeolites. (A) ammonia, 20 Torr, 0°C ; (& phosphine, 21 Torr, 0°C ; (W) arsine, 10 Torr, 0°C ; (0) but-1-ene, 50 Torr, 0°C ; (a) trans-but-2-ene, 50 Torr, 0°C ; (El) cis-but-Zene, 50 Torr, 0°C.T . TAKAISHI, Y . YATSURUGI, A . YUSA AND T . KURATOMI 101 Next, sorptive abilities for unsaturated hydrocarbons and polar gases were studied. Curves of the sorbed amount against composition of zeolite are shown in fig. 3. The curves have rather gentle slopes in comparison with those in fig. 2.In these cases molecules can be sorbed by zeolites whose window apertures are expected to be smal- ler than the molecular diameters. It is probable that the temporary displacement of the ion on an a-site is facilitated by the polarity or unsaturated bond of the visiting molecule. K-A AND Na-A ZEOLITES EXCHANGED WITH Ca2+ Ion-exchange isotherms and sorption properties of these zeolites are shown in fig. 4 and 5. The difference between the (K, Ca)-A and (Na, Ca)-A series is also apparent. The ion-exchange isotherm rises steeply up to 33 % exchange in the (Ca, Na)-A series, and in the (Ca, K)-A series up to 66 % exchange, both having no plateau. The change in the window size occurs at 33 % exchange in the (Ca, Na)-A series and at 75 % exchange in the (Ca, K)-A series, but never at 83 % exchange.The difference between the (K, Ca)-A and (Na, Ca)-A series may be attributable to a difference between the site-selectivities of potassium and sodium ions, and the existence of the 4.5A phase depends on the nature of the exchanging divalent cations. FIG. 4.4011-exchange isotherms zeolite A at 0.2 total metal ion normality and at 80°C : (0) Ca/K-A system : (0) Ca/Na-A system. THERMAL STABILITY OF ION-EXCHANGED MOLECULAR SIEVE 3 A Generally zeolite A is thermally not so stable as other zeolites, say, X or Y. Commercial molecular sieve 3A is used as a powerful sorbent for dehydrating organic vapours and oils. In the reactivation process, it is unavoidable that the mol- ecular sieve 3A is heated in an atmosphere of steam generated from itself, thus losing its sorptive ability more easily than in vacuum if the temperature is as high as 350°C.9 On the other hand, it is well-known that the (Na, Ca)-A zeolite becomes thermally more stable.It may be that thermally weak points in zeolite A are alkali metal ions or protons on y-site. Such weak points disappear if 17 % of the K+ is ion-exchanged102 SIEVING ACTION OF ZEOLITE A with CaZ+.I0 According to the results in the preceding section, K-A zeolite retains its sieving character as 3A, even if 66 % of its K+ is replaced by divalent cations. Hence, an investigation of the thermal stability of K-A, heavily ion-exchanged with divalent cations, promised to be interesting. " 0 0.2 0.4 0.6 0.8 1.0 Caz FIG. 5.-Adsorption of non-polar gases on (Ca, K)-A zeolites.(0) carbon dioxide, 200 Torr, 0°C ; (a) oxygen, 160 Torr, - 183°C ; (A) nitrogen, 200 Torr, - 195°C ; (0) n-butane, 200 Torr, 0°C ; (El) methane, 90 Torr, - 183°C ; (0) monosilane, 160 Torr, 0°C. FIG. 6.-The rate of adsorption of water by A zeolites in pellet form ; (A) KI2-A, untreated ; (E) K,,-A, treated ; (0) (Ca3K6)-A, untreated ; (0) (CasKs)-A, treated.T . TAKAISHI, Y . YATSURUGI, A . YUSA A N D T . KURATOMI 103 Thermal treatment was carried out at 450°C by passing air containing 25 Torr (3340 N m-2) water through a zeolite column for about 100 h. The thermal stabilities of the zeolities were checked by measuring and comparing the rate and equilibrium amount of sorption of water vapour at O’C, for untreated and treated samples. The results obtained are shown in fig.6. It is concluded that the thermal stability of K-A zeolite is appreciably improved by introducing a large number of divalent cations in place of the potassium ions occupying P-sites. DISCUSSION Let us consider the site-selectivity of alkali metal ions in dehydrated A zeolites. When Na-A zeolite is ion-exchanged with lithium, n-hexane is sorbed at about 66 % exchange.ll This is interpreted as follows : up to this composition, a-sites are occupied by three Na+ ions and the introduced Li+ ions occupy P- or y-sites selectively. In other words, the affinity of Li+ ions to P-sites is stronger than that of Na+ ions, for dehydrated samples.l1. l2 From X-ray structural analysis, Ca2+ and Na+ ions in dehydrated Ca,Na,-A occupy exlusively @-sites.Hence Ca2+ and Na+ prefer the P-site to the a-site. On the other hand, when Na-A is ion-exchanged with potassium, the (4A43A) transition takes place at 15-25 % exchange.’. Also, when K-A zeolite is ion-exchanged with zinc or calcium, the transition occurs at about 75 % exchange. This means that K+ selectively occupies the a-site and reduces the size of the aperture to the a-cage ; thus, the affinity of K+ to the a-site is stronger than to the P-site. The above site-selectivities of cations are qualitatively tabulated in table 1. This table is deduced from a limited number of two components systems, and gives no knowledge on such systems as (Li, Zn, Ca)-A or (Li, Na, Ca)-A which have not yet been studied. TABLE AF AFFINITY OF CATIONS TO THE THREE KINDS OF SITE IN DEHYDRATED ZEOLITE A Li+ ++ ++++ +(?I Na+ ++ +++ + K+ +++ ++ + ++++ -(?) Ca2+ - Zn2+ + ++++ -(?) cation u-site &site y-site The magnitude of the affinity is qualitatively expressed by the number of + .The sign - means no affinity, and (?) “probably but not verified”. A cation on a given site interacts with the charged anion-framework and other cations via long-range coulombic forces. Hence, its affinity to a site is a function of the composition of the crystal. The determination of such a composition dependence may be carried out, in principle, by structural analysis, but this is extremely tedious and difficult. Molecular sieving properties of ion-exchanged zeolites provide clues to the problem, but their effectiveness is limited. Here, we construct diagrams of distribution of cations on sites, mainly from sorption data.Some parts of the diagrams are supported by X-ray data, but others are more or less speculative. Furthermore, there exists some descrepancy between sorption and X-ray data. For instance, Rees and Berry’s sorption data are interpreted by the model where one a-site per unit cell of Na-A zeolite is occupied by a pair of Naf ions, although X-ray data deny such a model. This discrepancy may be attributed to a difference in states of the two samples, say, degrees of dehydration or amount of H30+ which had replaced Na+. The following diagrams refer to completely dehydrated zeolites with ideal1 04 SIEVING ACTION OF ZEOLITE A compositions. Such diagrams may be used as a guide-line for future structural studies.The ion distribution diagram for (Na, Ca)-A zeolites is shown in fig. 7(a). For the (K, Ca)-A zeolites, there are no X-ray structural data, and therefore only the sieving characteristics can be used to deduce the provisional distribution diagram of cations, as shown in fig. 7(b). The fifth and sixth Ca2+ may be arranged in two ways, which are shown by broken and dotted lines. A choice between these can not be made from data on the sieving characteristics, but should be possible from X-ray structural analysis. 2 1 8 I_ 6 4 2 0 - 2 1 8 - 6 4 2 0 0 2 4 6 8 1 0 1 2 0 2 4 6 8 1 0 1 2 no. of Na+ exchanged no. of K+ exchanged (4 (6) FIG. 7.-(a) The distribution of cations in the (Ca, Na)-A series. V indicates vacancy. (6) The dis- tribution of cations in the (Ca,K)-A series.V indicates vacancy. m +a .r( z 2 1 s 6 4 2 0 0 2 4 6 8 1 0 1 2 no. of I(+ exchanged FIG. 8.-The distribution of cations in the (Zn, K)-A series. V indicates vacancy. The series of (K, &)-A zeolites is unique in that the 4.5 A phase appears. A plausible explanation for such a window of intermediate size is as follows. The ionic radius of Zn2+ (7.4 x 10-l1 m) is smaller than that of Na+ (9.5 x 10-l1 m) and of K+ (1.33 x 10-lo m).I4 Hence, if Zn2+ occupies the a-site, the aperture in the window to the a-cage will be larger than in those occupied by K+ or Na+, but smaller than the aperture in the vacant window. Thus, the fifth Zn2+ may occupy an mite.T . TAKAISHI, Y . YATSURUGI, A . YUSA A N D T . KURATOMI 105 Further verification of the position of the fifth Zn2+ ion may be made by X-ray structural analysis and application of the theory of per~olation,~ which are future problems.The site potential experienced by a cation depends on the composition of the zeolite, that is, the curvature and the depth of the potential surface are functions of the composition. If the curvature is gentle, the displacement of a cation is easily induced by a visiting molecule. The abnormal sieving properties, seen in fig. 3, may be ascribed to such a situation. The gentle curvature brings about a large Debye- Waller factor in the X-ray diffraction pattern. For (K, Zn)-A zeolite, the verification of the present model is under way by Seff.ls The characteristic properties of (Zn, K)-A, given in fig. 3, are of practical use in the separation of some gas mixtures.We succeeded, for example, in ultra-high purifica- tion of SiH4 containing a trace of PH3.16 Detailed results will be published in a separate paper. A possible distribution diagram is shown in fig. 8. The authors thank Messrs. Y. Kaneko and K. Itabashi for their assistance in chemical analysis and Prof. G. S. Lehman for language correction. This work was partially supported by a Grant-in-Aid for Research from the Ministry of Education of the Japanese Government. D. W. Breck, W. G. Eversole, R. M. Milton and T. B. Reed, J. Amer. Chem. Suc., 1956, 78, 5963. R. Y. Yanagida, A. A. Amaro and K. Seff, J. Phys. Chern., 1973, 77, 805. V. Gramlich and W. M. Meier, 2. Krist., 1971, 133, 134. L. V. C. Rees and T. Berry, Pruc. Cunf.Molecular Sieves (SOC. Chem. Ind., London, 1968), p. 149. F. P. Treadwell and W. T. Hall, Analytical Chemistry, vol. I1 (Wiley, New York, 9th edn., 1959), p. 409. R. M. Barrer and W. H. Meier, Trans. Faruday Suc., 1958, 54, 1074; 1959, 55, 130. H. Hoinkins and H. W. Levi, 2. Naturforsch. A , 1967, 22, 226; 1968, 23, 813; 1969,24, 1672; Pruc. Cunf. MoZecular Sieves (SOC. Chem. Ind, London, 1968), p. 339. N. Nagase, Sekiyu Gakkai Shi, 1970, 14, 101 (in Japanese). * D. W. Breck and J. V. Smith, Sci. Amer., 1959,200, 85. lo C. R. Allenbach and F. M. O’Conner, U. S. Pat. 3,506,593/1970. l 1 P. Colline and R. Wey, Cumpt. rend., 1970, 270, 1069. l 2 R. M. Barrer, L. V. C. Rees and D. J. Ward, Proc. Roy SOC. A, 1963, 273, 180. l3 K. Seff and D. P. Shoemaker, Acta Cryst., 1967, 22, 162.l4 L. Pauling, The Nature of the Chemical Bond (Cornell University Press, 3rd edn., 1970). l6 T. Takasihi, A. Yusa and Y. Yatsurugi, Pruc. 3rd Int. Con$ Molecular Sieves (Leuven University K. Seff, personal communication. Press, 1973), p. 246. Changes in the Sieving Action and Thermal Stability of Zeolite A produced by Ion-exchange BY T. TAKAISHI" Institute for Atomic Energy, Rikkyo (St. Paul's) University, Yokosuka, Japan 240-01 AND Y. YATSURUGI, A. YUSA AND T. KURATOMI Komatsu Electronic Metals Co., Shinomiya, Hiratsuka City, Japan 254 Received 8th April, 1974 Na-A and K-A zeolites have been ion-exchanged with Ca2+ and Zn'f. The molecular sieving properties of the ion-exchanged zeolites were measured as functions of their composition.Abrupt changes in sieving action occur at 75 % exchange in the (K, Ca)-A series, and at 75 and 83 % exchange in the (K, Zn)-A series, as measured with chemically unreactive molecules. The sieving action of (K, Zn)-A zeolite for some polar molecules and unsaturated hydrocarbons changes gradually with its composition. The thermal stability of melocular sieve 3A is greatly improved by introducing a large number of divalent cations into this K-A zeolite. These properties are discussed in terms of site-selectivities of the exchangeable cations. From these, are derived diagrams of cation distribution on three kinds of site in the crystal. The ideal composition of Na-A zeolite, per unit cell, is Na 2( A102) ( SO2) *xH, 0. * According to recent X-ray structural analysis studies of dehydrated Na-A zeolite,2 three kinds of site are available to the twelve sodium ions.The first is near the centre of the six-membered oxygen ring and called the P-site, since the six-membered oxygen ring constitutes a window for the B-cage. The second is near the centre of the eight- membered oxygen ring which lies on a {loo) face of the unit cube. This site is named the a-site, since the eight-membered oxygen ring is a window for the central cavity (a-cage). The third is in the central cavity, near the centre of the four-membered oxygen ring, the y-site. Each of the (100) faces of the unit cube contains four a-sites, but each face contains only one cation ; that is, three of the twelve a-sites in a unit cell are occupied by Na+. On the other hand, the eight p-sites in a unit cell are occupied by eight Na+ ions.The twelfth Na+ ion in the unit cell occupies one of the twelve y-sites, to which the affinity of a cation is very weak. Sites a and B are preserved in hydrated zeolite A, although their coordinates are slightly changed. In contrast, y-sites cannot be defined in the hydrated crystal; that is, the position of the twelfth Na+ cannot be 10cated.~ The molecular sieving property of a dehydrated zeolite A is primarily determined by the kind of ion occupying the a-sites. Rees and Berry4 studied the blocking effect of pre-sorbed NH3 which is localized upon a cation on an a-site, and concluded that K+ in (Na, K)-A Ions on a-sites partially block the window aperture to the a-cage. 1-4 9798 SIEVING ACTION OF ZEOLITE A preferentially occupies the a-site.However, most existing data on ion-exchanged zeolite A concern those containing Na+ ions. In view of the difference in site select- ivities of Na+ and K+, more systematic studies of the distribution of cations on a-sites are desirable to obtain insight into the chemistry of zeolite A. In the present work, we have studied the differences between Na-A and K-A zeolites when they are ion- exchanged with divalent cations such as calcium and zinc ions. EXPERIMENTAL CHEMICAL ANALYSIS Zeolite A was dissolved in 2 equiv. dnr3 HCI, and A1 and Si contents were determined by the usual method^.^ The alkali metal ion content was measured with an atomic absorption spectrometer. The content of divalent cation in the ion-exchanging solution was determined by ethylenediaminetetra-acetic acid titration, while that in the ion-exchanged zeolite was, in most cases, calculated from the contents in the solution and material balance.In some cases, however, direct analysis was performed with an atomic absorption spectrometer. Radio- active tracers were also used to obtain greater accuracies for low concentrations of the species in solution. Tracers were obtained by irradiating reagents with the TRIGA-I1 reactor of our institute. MATERIAL The starting material was Linde molecular sieve 4A (lot. no. 480017) powder. After the standard treatment of the sample,6 its composition was determined. The sum of the com- ponents amounted to 98.8 % of the original sample. The formula was 0.94Na20-1.00 Al2O3=1.86 Si02*4.50 H20.The Na20/AI2O3 ratio of less than 1 may be ascribed partially to protons, which had replaced alkali metal ions and were not analyzed. The composition of a real crystal of Na-A zeolite differs from that of an ideal one and is 6~ 1 in our crystal. All chemical reagents used were S.P. grade and common adsorbate gases contained in glass cylinders (Takachiho Chemical Co.) had nominal purities better than 99.9 % as determined by mass spectroscopic analysis. Less common gases, such as SiH4, PH3, ASH,, and B2H6, were obtained from the Nihon Oxygen Co., and were of the special high purity used in the semiconductor industry. Na12(A102)12(Si02)12(NaA102)~ with 0 < 6 < 1.6 ADSORPTION MEASUREMENTS Adsorption measurements were made with a McBain type quartz spring balance which had a sensitivity of 0.5 m 8-l and an accuracy of +O.l mg.Prior to adsorption experiments, adsorbents were baked at 400°C for 5 h under a vacuum pressure lower than 1 x N m-2. RESULTS K-A AND Na-A ZEOLITES EXCHANGED WITH Zn2+ The ion exchange was carried out at 80+0.5"C with solutions of various composi- tions all of 0.2 total metal ion normality. Solutions of ZnCl, + NaCl and ZnC1, + KCl were used for Na-A and K-A, respectively. Although the reaction proceeded very rapidly,' the exchange was carried out for 20 h in order to obtain a true equilibrium. The forward exchange (Na+ or K+-+Zn2+) and the reverse exchange (Zn2++Na+ or K+) were measured to check the attainment of true equilibrium and the reversibility of the exchange reaction. After the reaction, the zeolite was filtered off and the Na, K and Zn contents in the solid and liquid phases were determined.The ion-exchange isotherms obtained are shown in fig. 1, in which (Zn), and (Zn), denote the equivalentT. TAKAISHI, Y . YATSURUGI, A. YUSA AND T. KURATOMI 99 fraction of zinc ion (2[Zn]/(2[Zn] + "a]) or 2[Zn]/(2[Zn] + [K])) in zeolite and in liquid, respectively. The isotherms show a steep rise up to 66 % exchange, and then become rather flat. The isotherm of (K, Zn)-A has a distinct step, while that of (Na, Zn)-A has the usual form. This difference may be ascribed, in part, to the differ- ence in the site-selectivities of potassium and sodium ions. After heating the zinc- exchanged zeolites in air at 400"C, their sorptive capacities and X-ray diffraction patterns were compared with those of unheated ones, and no change was introduced by such treatment.This indicates that the zinoexchanged zeolites are stable, in contrast to the earlier conclusion of Barrer and Meier.6 FIG. 1.-The ion-exchange isotherms for zeolite A at 0.2 totaf metal ion normality Zn/Na-A system ; (0) Zn/K-A system. and 80°C: (A) First, the sorptive abilities of the above zeolites for non-polar gases were studied. The rate of sorption was not studied in detail, and only the amounts of sorption at equilibrium are discussed herein. Sorption was measured at fixed temperatures and pressures for : nitrogen and argon at - 195°C ; carbon dioxide, n-butane, silane and diborane at 0°C. Curves of amount sorbed against composition of zeolite are shown in fig.2, in which a distinct difference between the two series (Zn, Na)-A and (Zn, K)-A is apparent. The curves for the (Zn, Na)-A series for nitrogen and n-butane sorption show steep rises at 33 % exchange in a manner similar to (Ca, Na)-A systems.l* In the (Zn, K)-A series, on the other hand, the curve for n-butane rises sharply at 75 % exchange, and those for nitrogen and argon at 83 %. The window size of (Zn, K)-A zeolite with a composition 0.75 < (Zn), .c 0.85 may be larger than that of Na-A zeolite but smaller than that of Ca-A, since the Zn-exchanged zeolites sorb n-butane but do not sorb argon or nitrogen. Hence, we shall tentatively call these 4.5A. We must not conclude, from the above sieving action of 4SA, that the molecular diameters of argon and nitrogen are larger than that of n-butane, because the cation on the a-site can be displaced from its equilibrium position to widen the window for a visiting gaseous molecule.Such cation displacement, however, has an activation energy and does not easily occur at liquid nitrogen temperatures at which sorption of nitrogen and argon was measured.8 The steep rise is also seen in diborane adsorption at 83 % exchange in the (Zn, K)-A series, at 0°C. This result is reasonable, since diborane has a bridged structure and its molecular diameter is larger than that of ethane.100 SIEVING ACTION OF ZEOLITE A 0 0.2- 0.4 0.6 6.8 1.0 - 0 0.2 0.4 0.6 0.8 1.0 (4 (b) znz FIG. 2.-(a) Adsorption of non-polar gases on (Zn, K)-A zeolites. (A) nitrogen, 200 Torr, - 195°C ; (m argon, 50 Torr, - 195°C ; (0) carbon dioxide, 210 Torr, 0°C ; (e) n-butane, 200 Torr, 0°C ; (0) monosilane, 160 Tom, 0°C ; (m) diborane, 50 Torr, 0°C.(6) Adsorption of non-polar gases on (Zn, Na)-A zeolites. (A) nitrogen, 200 Torr, - 195°C ; (0) carbon dioxide, 210 Torr, 0°C ; (e) n-butane, 220 Torr, 0°C ; (0) monosilane, 160 Torr, 0°C. 1 Torr = (101 325/760) N m-2. 0.20 0 c) .- c( 0.15 B c1 (d 0 (d .C) c, =!! 0.10 d 0 0 a .- c, .C( 0.05 .* c) a 5 4 0 znz FIG. 3.-Adsorption of polar gases and unsaturated hydrocarbons on (Zn, K)-A zeolites. (A) ammonia, 20 Torr, 0°C ; (& phosphine, 21 Torr, 0°C ; (W) arsine, 10 Torr, 0°C ; (0) but-1-ene, 50 Torr, 0°C ; (a) trans-but-2-ene, 50 Torr, 0°C ; (El) cis-but-Zene, 50 Torr, 0°C.T . TAKAISHI, Y .YATSURUGI, A . YUSA AND T . KURATOMI 101 Next, sorptive abilities for unsaturated hydrocarbons and polar gases were studied. Curves of the sorbed amount against composition of zeolite are shown in fig. 3. The curves have rather gentle slopes in comparison with those in fig. 2. In these cases molecules can be sorbed by zeolites whose window apertures are expected to be smal- ler than the molecular diameters. It is probable that the temporary displacement of the ion on an a-site is facilitated by the polarity or unsaturated bond of the visiting molecule. K-A AND Na-A ZEOLITES EXCHANGED WITH Ca2+ Ion-exchange isotherms and sorption properties of these zeolites are shown in fig. 4 and 5. The difference between the (K, Ca)-A and (Na, Ca)-A series is also apparent.The ion-exchange isotherm rises steeply up to 33 % exchange in the (Ca, Na)-A series, and in the (Ca, K)-A series up to 66 % exchange, both having no plateau. The change in the window size occurs at 33 % exchange in the (Ca, Na)-A series and at 75 % exchange in the (Ca, K)-A series, but never at 83 % exchange. The difference between the (K, Ca)-A and (Na, Ca)-A series may be attributable to a difference between the site-selectivities of potassium and sodium ions, and the existence of the 4.5A phase depends on the nature of the exchanging divalent cations. FIG. 4.4011-exchange isotherms zeolite A at 0.2 total metal ion normality and at 80°C : (0) Ca/K-A system : (0) Ca/Na-A system. THERMAL STABILITY OF ION-EXCHANGED MOLECULAR SIEVE 3 A Generally zeolite A is thermally not so stable as other zeolites, say, X or Y.Commercial molecular sieve 3A is used as a powerful sorbent for dehydrating organic vapours and oils. In the reactivation process, it is unavoidable that the mol- ecular sieve 3A is heated in an atmosphere of steam generated from itself, thus losing its sorptive ability more easily than in vacuum if the temperature is as high as 350°C.9 On the other hand, it is well-known that the (Na, Ca)-A zeolite becomes thermally more stable. It may be that thermally weak points in zeolite A are alkali metal ions or protons on y-site. Such weak points disappear if 17 % of the K+ is ion-exchanged102 SIEVING ACTION OF ZEOLITE A with CaZ+.I0 According to the results in the preceding section, K-A zeolite retains its sieving character as 3A, even if 66 % of its K+ is replaced by divalent cations.Hence, an investigation of the thermal stability of K-A, heavily ion-exchanged with divalent cations, promised to be interesting. " 0 0.2 0.4 0.6 0.8 1.0 Caz FIG. 5.-Adsorption of non-polar gases on (Ca, K)-A zeolites. (0) carbon dioxide, 200 Torr, 0°C ; (a) oxygen, 160 Torr, - 183°C ; (A) nitrogen, 200 Torr, - 195°C ; (0) n-butane, 200 Torr, 0°C ; (El) methane, 90 Torr, - 183°C ; (0) monosilane, 160 Torr, 0°C. FIG. 6.-The rate of adsorption of water by A zeolites in pellet form ; (A) KI2-A, untreated ; (E) K,,-A, treated ; (0) (Ca3K6)-A, untreated ; (0) (CasKs)-A, treated.T . TAKAISHI, Y . YATSURUGI, A . YUSA A N D T . KURATOMI 103 Thermal treatment was carried out at 450°C by passing air containing 25 Torr (3340 N m-2) water through a zeolite column for about 100 h.The thermal stabilities of the zeolities were checked by measuring and comparing the rate and equilibrium amount of sorption of water vapour at O’C, for untreated and treated samples. The results obtained are shown in fig. 6. It is concluded that the thermal stability of K-A zeolite is appreciably improved by introducing a large number of divalent cations in place of the potassium ions occupying P-sites. DISCUSSION Let us consider the site-selectivity of alkali metal ions in dehydrated A zeolites. When Na-A zeolite is ion-exchanged with lithium, n-hexane is sorbed at about 66 % exchange.ll This is interpreted as follows : up to this composition, a-sites are occupied by three Na+ ions and the introduced Li+ ions occupy P- or y-sites selectively. In other words, the affinity of Li+ ions to P-sites is stronger than that of Na+ ions, for dehydrated samples.l1.l2 From X-ray structural analysis, Ca2+ and Na+ ions in dehydrated Ca,Na,-A occupy exlusively @-sites. Hence Ca2+ and Na+ prefer the P-site to the a-site. On the other hand, when Na-A is ion-exchanged with potassium, the (4A43A) transition takes place at 15-25 % exchange.’. Also, when K-A zeolite is ion-exchanged with zinc or calcium, the transition occurs at about 75 % exchange. This means that K+ selectively occupies the a-site and reduces the size of the aperture to the a-cage ; thus, the affinity of K+ to the a-site is stronger than to the P-site. The above site-selectivities of cations are qualitatively tabulated in table 1.This table is deduced from a limited number of two components systems, and gives no knowledge on such systems as (Li, Zn, Ca)-A or (Li, Na, Ca)-A which have not yet been studied. TABLE AF AFFINITY OF CATIONS TO THE THREE KINDS OF SITE IN DEHYDRATED ZEOLITE A Li+ ++ ++++ +(?I Na+ ++ +++ + K+ +++ ++ + ++++ -(?) Ca2+ - Zn2+ + ++++ -(?) cation u-site &site y-site The magnitude of the affinity is qualitatively expressed by the number of + . The sign - means no affinity, and (?) “probably but not verified”. A cation on a given site interacts with the charged anion-framework and other cations via long-range coulombic forces. Hence, its affinity to a site is a function of the composition of the crystal. The determination of such a composition dependence may be carried out, in principle, by structural analysis, but this is extremely tedious and difficult.Molecular sieving properties of ion-exchanged zeolites provide clues to the problem, but their effectiveness is limited. Here, we construct diagrams of distribution of cations on sites, mainly from sorption data. Some parts of the diagrams are supported by X-ray data, but others are more or less speculative. Furthermore, there exists some descrepancy between sorption and X-ray data. For instance, Rees and Berry’s sorption data are interpreted by the model where one a-site per unit cell of Na-A zeolite is occupied by a pair of Naf ions, although X-ray data deny such a model. This discrepancy may be attributed to a difference in states of the two samples, say, degrees of dehydration or amount of H30+ which had replaced Na+.The following diagrams refer to completely dehydrated zeolites with ideal1 04 SIEVING ACTION OF ZEOLITE A compositions. Such diagrams may be used as a guide-line for future structural studies. The ion distribution diagram for (Na, Ca)-A zeolites is shown in fig. 7(a). For the (K, Ca)-A zeolites, there are no X-ray structural data, and therefore only the sieving characteristics can be used to deduce the provisional distribution diagram of cations, as shown in fig. 7(b). The fifth and sixth Ca2+ may be arranged in two ways, which are shown by broken and dotted lines. A choice between these can not be made from data on the sieving characteristics, but should be possible from X-ray structural analysis.2 1 8 I_ 6 4 2 0 - 2 1 8 - 6 4 2 0 0 2 4 6 8 1 0 1 2 0 2 4 6 8 1 0 1 2 no. of Na+ exchanged no. of K+ exchanged (4 (6) FIG. 7.-(a) The distribution of cations in the (Ca, Na)-A series. V indicates vacancy. (6) The dis- tribution of cations in the (Ca,K)-A series. V indicates vacancy. m +a .r( z 2 1 s 6 4 2 0 0 2 4 6 8 1 0 1 2 no. of I(+ exchanged FIG. 8.-The distribution of cations in the (Zn, K)-A series. V indicates vacancy. The series of (K, &)-A zeolites is unique in that the 4.5 A phase appears. A plausible explanation for such a window of intermediate size is as follows. The ionic radius of Zn2+ (7.4 x 10-l1 m) is smaller than that of Na+ (9.5 x 10-l1 m) and of K+ (1.33 x 10-lo m).I4 Hence, if Zn2+ occupies the a-site, the aperture in the window to the a-cage will be larger than in those occupied by K+ or Na+, but smaller than the aperture in the vacant window.Thus, the fifth Zn2+ may occupy an mite.T . TAKAISHI, Y . YATSURUGI, A . YUSA A N D T . KURATOMI 105 Further verification of the position of the fifth Zn2+ ion may be made by X-ray structural analysis and application of the theory of per~olation,~ which are future problems. The site potential experienced by a cation depends on the composition of the zeolite, that is, the curvature and the depth of the potential surface are functions of the composition. If the curvature is gentle, the displacement of a cation is easily induced by a visiting molecule. The abnormal sieving properties, seen in fig. 3, may be ascribed to such a situation. The gentle curvature brings about a large Debye- Waller factor in the X-ray diffraction pattern. For (K, Zn)-A zeolite, the verification of the present model is under way by Seff.ls The characteristic properties of (Zn, K)-A, given in fig. 3, are of practical use in the separation of some gas mixtures. We succeeded, for example, in ultra-high purifica- tion of SiH4 containing a trace of PH3.16 Detailed results will be published in a separate paper. A possible distribution diagram is shown in fig. 8. The authors thank Messrs. Y. Kaneko and K. Itabashi for their assistance in chemical analysis and Prof. G. S. Lehman for language correction. This work was partially supported by a Grant-in-Aid for Research from the Ministry of Education of the Japanese Government. D. W. Breck, W. G. Eversole, R. M. Milton and T. B. Reed, J. Amer. Chem. Suc., 1956, 78, 5963. R. Y. Yanagida, A. A. Amaro and K. Seff, J. Phys. Chern., 1973, 77, 805. V. Gramlich and W. M. Meier, 2. Krist., 1971, 133, 134. L. V. C. Rees and T. Berry, Pruc. Cunf. Molecular Sieves (SOC. Chem. Ind., London, 1968), p. 149. F. P. Treadwell and W. T. Hall, Analytical Chemistry, vol. I1 (Wiley, New York, 9th edn., 1959), p. 409. R. M. Barrer and W. H. Meier, Trans. Faruday Suc., 1958, 54, 1074; 1959, 55, 130. H. Hoinkins and H. W. Levi, 2. Naturforsch. A , 1967, 22, 226; 1968, 23, 813; 1969,24, 1672; Pruc. Cunf. MoZecular Sieves (SOC. Chem. Ind, London, 1968), p. 339. N. Nagase, Sekiyu Gakkai Shi, 1970, 14, 101 (in Japanese). * D. W. Breck and J. V. Smith, Sci. Amer., 1959,200, 85. lo C. R. Allenbach and F. M. O’Conner, U. S. Pat. 3,506,593/1970. l 1 P. Colline and R. Wey, Cumpt. rend., 1970, 270, 1069. l 2 R. M. Barrer, L. V. C. Rees and D. J. Ward, Proc. Roy SOC. A, 1963, 273, 180. l3 K. Seff and D. P. Shoemaker, Acta Cryst., 1967, 22, 162. l4 L. Pauling, The Nature of the Chemical Bond (Cornell University Press, 3rd edn., 1970). l6 T. Takasihi, A. Yusa and Y. Yatsurugi, Pruc. 3rd Int. Con$ Molecular Sieves (Leuven University K. Seff, personal communication. Press, 1973), p. 246.

ISSN:0300-9599    

DOI:10.1039/F19757100097

出版商:RSC    

年代:1975

数据来源: RSC