摘要:

Journal of the Chemical Society, Faraday Transactions I ISSN 0300-959944 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) M ETAL(contd) prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 2. Thermodyn- amic parameters and theoretical models, 578-90 Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(III), molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds, 603-1 3 Kinetics of thermal decomposition of 4,4'-azobis-(kyanopentanoic acid) and its salts in aqueous solution, 935-41 Electrolyte solutions in liquid ammonia. Part 6. Mean mold activity coefficients and cation transference numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 Structural analysis of some molten materials by x-ray diffraction.Part 6. Manganese chloride-lithium or potassium chloride, 1 16 1-8 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl sulfoxide + water mixtures, 1780-7 Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part I . Equilibriums and affinities, 1969-83 Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 2. Heats and entropies, 1984-99 Infrared spectroscopic investigation of sorption of nitric oxide by fi-metal phthalocyanines of the first transition period, 2594-600 Influence of lithium, sodium, and potassium cation concentrations in X and Y zeolites on isotherms and heats of adsorption of propane and water, 2662-77 Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, dichromates and thiocyanates, 273543 Kinetics of reaction of magnesium vapor with silica, 2807-18 Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron spectroscopy and gravimetry, 2839-50 Mechanism for incorporation of oxygen in vapor-phase selective oxidation of isobutene, butadiene and furan over various bismuth catalysts, 1757-68 Effect of hexamethylphosphoric triamide on the living radical polymerization initiated by aged chromium ion + benzoyl peroxide, 1821-9 Photolysis of methyl- and acetyl-manganese pentacarbonyls in methyl methacrylate.Initiation of polymerization and formation of methyl(2-methyl4oxopentanoate), 2562-75 Methanation of carbon dioxide and carbon monoxide on supported nickel-based composite catalysts, y-Radiolysis of methane adsorbed on y-alumina. Part 3. Influence of additives on product desorption, Reaction between hydrogen atoms and methane, 83543 Structure and dynamics of ammonia adsorbed on graphitized carbon black. Part 2. Neutron diffraction, Viscosities of trichlorofluoromethane, chlorotrifluoromethane, dichlorofluoromethane, chlorodifluoro- Detection of formyl radicals in low pressure methane + oxygen flames, 2423-32 Evaluation of semiconducting tin oxide as an electrocatalyst support, 2 165-76 Characterization of electrooxidation catalysts prepared by ion-xchange of platinum salts with surface Study of the methanol oxidation reaction on platinum using the potential-step technique, 2535-43 Reaction between hydrogen atoms and methane, 83543 Effect of hexamethylphosphoric triamide on the living radical polymerization initiated by aged Kinetics and mechanism of the decomposition of hydrogen sulfide, methyl hydrogen sulfide and dimethyl Surface activity of methyl orange, 2077-82 Intermediates in catalytic hydrogenation of 1,3-butadiene, propadiene and methylacetylene on molybdenum disulfide catalyst, 1403-16 Ligand substitution processes at five coordinate copper( 11) centers in hydrophilic and hydrophobic environments, 123- 305-13 METHACRYLALDEHYDE METHACRYLATE METHANATION 787-802 METHANE 192-204 1542-52 methane, and trifluoromethane from 373 to 570 K, 1752-6 METHANOL oxide groups on carbon, 23 12-24 METHYL chromium ion + benzoyl peroxide, 1821-9 sulfide in a radio-frequency pulse discharge, 1868-75 METHYL ORANGE METHY LACETY LENE METHYLAMINOETHY LAMINOJ.C.S. FARADAY I SUBJECT INDEX VOL.75 (1979) 45 METHY LANILINE Association between polar molecules. Part 1. Nuclear magnetic resonance study of the dipole association Use of internal and external references in nuclear magnetic resonance determinations of association Use of internal and external references in nuclear magnetic resonance determinations of association Micellar catalysis of metal complex formation. Part 2. Kinetics of the reaction between aquated of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 constants of weak molecular complexes, 1222-5 constants of weak molecular complexes, 1222-5 nickel(2+ ) and various neutral bidentate ligands in the presence of sodium dodecylsulfate micelles in aqueous solution, 2395-405 METHY LBENZENE METHY LBIPY RIDY L METHYLBUTANE M ETHY LBUTY L METHY LCHLOROSILANE M ETHY LCYCLOALKANE METHYLDISILAZANE Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Catalytic reactions with deuterium of several polymethylcycloalkanes on cobalt films, 1320-9 Temperature dependence of the interdiffusion coefficient for very dilute solutions of 1,l ,2,2-tetrabromo- Temperature dependence of the interdiffusion coefficient for very dilute solutions of 1,1,2,2-tetrabromo- Photolysis of methyl- and acetyl-manganese pentacarbonyls in methyl methacrylate.Initiation of Photolysis of methyl- and acetyl-manganese pentacarbonyls in methyl methacrylate. Initiation of Analysis of thermal effects in adsorption rate measurements, 2404-22 Diffusion in viscous solvents. Part 3. Interdiffusion coefficients for planar and spherical solutes in 2- methylpentane-2,&diol and their relationship to diffusion coefficients derived from luminescence measurements, 141 7-32 Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane.Evidence for an intermediate radical cation, 9 14-20 Kinetics of the gas-phase reaction between iodine and trimethylsilane and the bond dissociation energy Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 Photolysis of thioketene S-oxides, 2624-7 Metal-ion complexation reactions in the presence of surfactants. Part I . Mechanism of pH4ependent reaction between nickel(I1) and murexide in aqueous solution and application of the reaction to study of micellar phenomena, 1 19-3 1 Characterization of water-containing reversed micelles by viscosity and dynamic light scattering methods Ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexide-containing micelles in the system aerosol-OT + water + heptane, 48 1-96 Absorption and emission studies of solubilization in micelles.Part 5. Pyrene-3-sodium sulfonate solubilized in didodecyldimethylammonium bromide inverted micelles in benzene, 55CL60 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous solutions. A pulse radiolysis study, 1674-87 Kinetic investigation and numerical analysis of a micelle-catalyzed metal complex formation, 2 199-2 10 Micellar catalysis of metal complex formation. Part 2. Kinetics of the reaction between aquated nickel(2 +) and various neutral bidentate ligands in the presence of sodium dodecylsulfate micelles in aqueous solution, 2395-405 Attachment of spherical particles to surface of a pendant drop and tension of the wetting perimeter, 1-6 Stability of metal uncharged ligand complexes in ion exchangers.Part 4. Hydration effects and stability ethane in hexamethyldisiloxane and hexamethyldisilazane, 1232-5 ethane in hexamethyldisiloxane and hexamethyldisilazane, 1 232-5 polymerization and formation of methyl(2-methyl4oxopentanoate), 2562-75 polymerization and formation of methyl(2-methyl-&oxopentanoate), 2562-75 METHY LDISILOXANE M ETHY LMANGANESE METHYLOXOPENTANOATE METHYLPENTANE METHYLPENTANEDIOL METHY LPHENYLENEDIAMINE METHYLSILANE of the hydrogen-silicon bond. Part 2, 1 126-3 I METHY LSILY L METHYLTHIOCARBONY LCYCLOHEXANE MICELLE 132-9 MICROSPHERE MINERAL changes of copper-ethylenediamine complexes in montmorillonite, 5 13-2446 J.C.S. FARADAY I SUBJECT INDEX VOL.75 (1979) MINERAL(contd) MIXED Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-91 Oxidation of propene on mixed oxides of copper and cobalt, I33745 Acidic and basic properties of mixed tin-antimony oxides, 2762-7 Study of the oxidative reactions of butenes over mixed tin-antimony oxides, 2768-80 Excess volumes of mixing of aniline + aromatic hydrocarbons, I 120-5 Correlation between topological features and molar volumes of n-alkanes and excess volumes of their MIXING binarv mixtures. 1132-41 Effects 6f molecular flexibility and shape on the excess enthalpies and heat capacities of alkane systems, 1700-7 Effects of molecular shape on the excess enthalpies and heat capacities of cycloalkane systems, 1708-14 Excess Gibbs function for a binary mixture obtained from dew point and bubble point pressure Thermodynamic properties and local structures of nonstoichiometry.Galvanic cell study of fused indium Magnetooptical rotation studies of liquid mixtures. Part 3. Specific interactions in mixtures of carbon Thermal diffusion in mixtures of helium with argon, neon, nitrogen and carbon dioxide and of neon with Calculation of excess enthalpies for binary mixtures of liquid metal halides, 729-35 Structural analysis of some molten materials by x-ray diffraction. Part 6 . Manganese chloride-lithium or Molecular reorientation and molecular association. Study by Raman depolarized light scattering, Excess Gibbs function for a binary mixture obtained from dew point and bubble point pressure Rhodium catalysts prepared by attachment of hexarhodium hexadecacarbonyl onto chemically modified Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Molecular reorientation and molecular association.Study by Raman depolarized light scattering, Thermodynamic functions of hydration of saturated uncharged organic compounds. Free energies, Binding of n-alkyl sulfates to lysozyme in aqueous solution, 173M4 Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Surface activity of methyl orange, 2077-82 Detection of formyl radicals in low pressure methane + oxygen flames, 2423-32 High resolution proton nuclear magnetic resonance studies of interaction between deoxyhemoglobin and Correlation between topological features and molar volumes of n-alkanes and excess volumes of their Partial molar enthalpies in the sodium chloride + calcium chloride + water system.Application of the Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Calculation of excess enthalpies for binary mixtures of liquid metal halides, 729-35 Mechanism for incorporation of oxygen in vapor-phase selective oxidation of isobutene, butadiene and furan over various bismuth catalysts, 1757-68 Dynamic properties of n-alkyl and s-alkyl intermediates in reactions of simple alkenes with hydrogen on molybdenum disulfide catalyst, 7-21 Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(III), molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds, 603-1 3 Intermediates in catalytic hydrogenation of 1,3-butadiene, propadiene and methylacetylene on molybdenum disulfide catalyst, 1403-1 6 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared by use of tetrakis-n-allylmolybdenum.1465-76 Kinetics of solid state thermal monomerization of 9-cyanoanthracene photodimer and photodimerizatio- n of 9+yanoanthracene, 11 11-19 measurements, 1940-9 iodide mixtures, 2303-1 1 tetrachloride with aliphatic, alicyclic and aromatic hydrocarbons and amines, 370-3 argon, 621-30 potassium chloride, 1 161-8 MIXT 1 179-83 measurements, 1940-9 silicas.Characterization of the infrared spectra in the carbonyl stretching region. 1888-99 MODIFICATION MOL 1 179-83 enthalpies and entropies at 25"C, 1 184-95 small molecules. Dithionite and diphosphoglycerate, 285 1-64 binary mixtures, 1 132-41 McKay-Perring method to thermal measurements, 1745-5 1 MOLAR MOLECULE MOLTEN MOLYBDATE MOLYBDENUM MONOM ERIZATIONJ.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 47 MONOXIDE Some unusual properties of activated and reduced silver-sodium A zeolites, 56-64 infrared study of carbon monoxide adsorption on magnesium oxide, 96-108 Chemisorption of carbon monoxide and hydrogen on silver-sodium mordenite, 109-1 8 Surface acidity of q-alumina. Part 2. Interaction of pyridine with other adsorbates, 289-304 Search for strongly adsorbed carbon monoxide at copper single crystal surfaces using ultraviolet Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Hydrogenation of acetylene over supported metal catalysts.Part 3. Carbon- 14 tracer studies of the effects of added ethylene and carbon monoxide on the reaction catalyzed by silica-supported palladium, rhodium and iridium, 1900-1 1 Mechanism and kinetics of the chain reaction in hydrogen peroxide + nitrous oxide + carbon monoxide systems, 2048-59 MONTMORILLONITE Stability of metal uncharged ligand complexes in ion exchangers. Part 4. Hydration effects and stability changes of copper-ethylenediamine complexes in montmorillonite, 5 13-24 Chemisorption of carbon monoxide and hydrogen on silver-sodium mordenite, 109-1 8 Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 Changes in mordenite upon various pretreatments.Part 1. Structural rearrangements, 1245-53 Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 1. Changes in mordenite upon various treatments. Part 2. Hydroxyl groups, 2366-76 Metal-ion complexation reactions in the presence of surfactants. Part 1. Mechanism of pHdependent reaction between nickel(I1) and murexide in aqueous solution and application of the reaction to study of micellar phenomena, I 19-3 1 ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexide-containing micelles in the system aerosol-OT + water + heptane, 48 1-96 Thermal diffusion in mixtures of helium with argon, neon, nitrogen and carbon dioxide and of neon with photoelectron spectroscopy, 984-6 1487-94 MORDENITE Equilibriums and affinities, 1969-83 MUREXIDE NEON argon, 62 1-30 neopentane, 161 9-22 NEOPENTANE Reaction probabilities and threshold energy in the abstraction reaction between hydrogen atoms and Structure and dynamics of ammonia adsorbed on graphitized carbon black.Part 2. Neutron diffraction, Structure and dynamics of ammonia adsorbed on graphitized carbon black. Part 3. Neutron quasielastic Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending Metal-ion complexation reactions in the presence of surfactants.Part 1. Mechanism of pHdependent Magnetic determination of metallic nickel particles dispersed on X and Y zeolite structures, 165-71 Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 Ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexide-containing micelles in the system aerosol-OT + water + heptane, 48 1-96 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 1. Experiment- al results, 561-77 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 2. Thermodyn- amic parameters and theoretical models, 578-90 Catalysis by highly diluted nickel-copper alloy foils.Hydrogen-deuterium equilibration and ethylene- deuterium exchange reaction, 1001-10 Intermediates in carbon monoxide and dioxide hydrogenation over nickel catalysts, 101 1-1 5 Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in zeolite Y, 1 1 96-206 Kinetics of heterogeneous electron transfer on dicyanobis(tertiary phosphine)nickel complexes, 1330-6 Synergistic effect of composite catalysts on the direct hydrogenation of carbon, 1495-506 Selectivity in ethylene dimerization over supported nickel oxide catalysts, 15 13--20 Distribution of nickel ions among octahedral and tetrahedral sites in nickel zinc aluminate spinel solid Surface characterization of nickel boride and nickel phosphide catalysts by x-ray photoelectron Kinetic investigation and numerical analysis of a micelle<atalyzed metal complex formation, 2 199-2 I0 Micellar catalysis of metal complex formation.Part 2. Kinetics of the reaction between aquated NEUTRON 1542-52 and inelastic spectra, 1553-69 and descending scanning curves within B-type hysteresis loops, 42-55 reaction between nickel(I1) and murexide in aqueous solution and application of the reaction to study of micellar phenomena, 1 19-3 1 NICKEL solutions, 1876-87 spectroscopy, 2027-3948 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) N ICKEL(contd) nickel(2 + ) and various neutral bidentate ligands in the presence of sodium dodecylsuifate micelles in aqueous solution, 2395405 Oxygen electrodes in fused salts. Potentiometric and x-ray photoelectron spectroscopic (ESCA) findings on the system nickel dicarbonyl + dioxygen/carbonate ion in molten nitrates, 2628-37 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and nickel bromide solutions.Part 1. Application of solvent nuclear magnetic relaxation, 2700-1 1 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and nickel bromide solutions. Part 2. Application of halide nuclear magnetic relaxation, 27 12-34 Electrolyte solutions in liquid ammonia. Part 6. Mean molal activity coefficients and cation transference numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 Composition dependence of fluidities and conductances of mixtures of hydrated melts, 13 12-1 9 Particle adhesion and removal in model systems.Part 2. Monodispersed chromium hydroxide on steel. 20 14-26 Electric conductivity of nitrate inclusion complexes of A and X zeolites, 2083-8 Oxygen electrodes in fused salts. Potentiometric and x-ray photoelectron spectroscopic (ESCA) findings Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlordtes, Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 Infrared spectrum and surface reaction of nitric oxide adsorbed on silicon films, 1788-97 Study of the interaction of nitric oxide with copper( 100) and copper( 1 1 1) surfaces using low energy Infrared spectroscopic investigation of sorption of nitric oxide by p-metal phthalocyanines of the first Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron NITRATE on the system nickel dicarbonyl + dioxygen/carbonate ion in molten nitrates, 2628-37 dichromates and thiocyanates, 2735-43 NITRIC electron diffraction and electron spectroscopy, 2 143-59 transition period, 2594-600 spectroscopy and gravimetry, 2839-50 NITRI LOACETATOCOBALTATE Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 NITRI LOTRIACETATOCOBALTATE Radiation chemistry of cobalt(1 I) nitrilotriacetate in aqueous solutions, 2089-99 NITROANISOLE Association between polar molecules.Part 1.Nuclear magnetic resonance study of the dipole association of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 Association between polar molecules. Part 1. Nuclear magnetic resonance study of the dipole association of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 Reactions of hydroxymethyl radicals with some solutes in methanol. Pulse radiolysis study, 446-62 Use of internal and external references in nuclear magnetic resonance determinations of association constants of weak molecular complexes, 1222-5 Scanning studies on capillary condensation and evaporation of nitrogen. Part I . Apparatus and calculation method, 36-41 Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending and descending scanning curves within B-type hysteresis loops, 42-55 Second virial coefficients of nitrogen at very low temperatures, 479-80 Thermal diffusion in mixtures of helium with argon, neon, nitrogen and carbon dioxide and of neon with argon, 621-30 Kinetics of chemiionization in atomic oxygen + nitrogen mixtures, 1301-1 1 Kinetics of chlorine oxide radicals using modulated photolysis.Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Mechanism and kinetics of the chain reaction in hydrogen peroxide + nitrous oxide + carbon monoxide systems, 2048-59 Anomalous adsorption kinetics. y-Nitrogen on the (1 10) plane of tungsten, 21W15 Kinetics of chlorine oxide radical reactions using modulated photolysis. Part 3.Pressure and temperature dependence of the reaction: chlorine monoxide + nitrogen dioxide (+ M) + nitryl hypochlorite (+ Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron spectroscopy and gravimetry, 2839-50 Acid ionization of solvent, acetic acid, and 4-substituted phenols in tert-butanol + water mixtures, Electron spin resonance studies of the radicals formed from carbon-nitroso compounds and olefins. Part NITROBENZENE NITROGEN M), 2649-61 NITROPHENOL 278 1-97 NITROSODURENEJ.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 49 NITROSODURENE(contd) NITROSOFLUOROALKENE 1. Nitrosodurene and fluoroolefins, 1521-30 Electron spin resonance studies of the radicals formed from carbon-nitroso compounds and olefins. Part Association between polar molecules.Part I . Nuclear magnetic resonance study of the dipole association Viscosity of gaseous carbon dioxide, sulfur hexafluoride and nitrous oxide at low temperatures, 892-7 Decomposition of nitrous oxide on high surface area calcium oxide-magnesium oxide catalysts Elcctron-transfer reactions of nitroxyl radicals with one-electron reduced quinones and viologens, Association between polar molecules. Part 1. Nuclear magnetic resonance study of the dipole association Ion condensation model and nuclear magnetic resonance studies of counterion binding in lyotropic liquid Use of internal and external references in nuclear magnetic resonance determinations of association Structure and dynamics of ammonia adsorbed on graphitized carbon black.Part 4. Nuclear magnetic High resolution proton nuclear magnetic resonance studies of interaction between deoxyhemoglobin and Thermal diffusion in mixtures of helium with argon, neon, nitrogen and carbon dioxide and of neon with Sound velocity and surface tension from Flory’s statistical theory, 2 1 6 W Methanation of carbon dioxide and carbon monoxide on supported nickel-based composite catalysts, Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Analysis of thermal effects in adsorption rate measurements, 2406-22 Chlorine-35 nuclear quadrupole resonance studies of pentachlorophenol-amine hydrogen-bonded Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Crystal growth kinetics of globular proteins.Lysozyme and insulin, 2753-6 1 Electrolyte solutions in liquid ammonia. Part 6. Mean molal activity coefficients and cation transference Structure of molten potassium chloride, 1477-86 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and Observation of weak interactions between polystyrene particles and textile fibers, 2386-92 Solubility of octacosane and hexatriacontanc in different n-alkane solvents, 1254-8 Effects of water on proton migration in alcoholic solvents. Part 6. Conductance of hydrogen chloride in Inorganic photophysics in solution. Part 3.Temperature activation of decay processes in the luminescen- Thermodynamics of aggregation of some tertiary n-alkylammonium picrates in benzene solutions, Infrared study of carbon monoxide adsorption on magnesium oxide, 96-108 Reactions involving electron transfer at semiconducting surfaces. Part 8. Room temperature activity of 1. Nitrosodurene and fluoroolefins, 152 1-30 NITROTOLUENE of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 NITROUS investigated by electron spin resonance spectroscopy, 2 177-87 NITROXYL 1912-18 NMR of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 crystals, 663-8 constants of weak molecular complexes, 1222-5 resonance spectra, 1570-86 small molecules. Dithionite and diphosphoglycerate, 285 1-64 argon, 621-30 NOBLE GAS NOBLE METAL 787-802 NONIONIC reactions.Volume changes on activation and reaction, 172-9 1 NON ISOTHERMAL NQR complexes, 1587-92 NTA NUCLEATION NUMBER numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 nickel bromide solutions. Part 1. Application of solvent nuclear magnetic relaxation, 2700-1 1 nickel bromide solutions. Part 2. Application of halide nuclear magnetic relaxation, 27 12-34 NYLON OCTACOSANE OCTAN OL hexan-1-01 and in octan-1-01 at 25°C. 1667-73 OSMIUM ce of tris(2,2’-bipyridine)osmium(II) and tris( 1,l 0-phenanthroline)osmium(II) ions, 353-62 OSMOTIC 825-34 OXIDE50 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1 979) OXIDE(contd) prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, 305-13 Electron spin resonance and volumetric investigations of oxygen adsorption on high surface area cobalt oxide-magnesium oxide, 533-44 Isotopic exchange of molecular oxygen with oxygen of vanadium oxide (V205) modified with oxides of alkaline earth metals, 691-4 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces.Part 1. Role of hydroxyl groups in photoadsorption, 736-47 Methanation of carbon dioxide and carbon monoxide on supported nickel-based composite catalysts, Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Viscosity of gaseous carbon dioxide, sulfur hexafluoride and nitrous oxide at low temperatures, 892-7 Search for strongly adsorbed carbon monoxide at copper single crystal surfaces using ultraviolet Intermediates in carbon monoxide and dioxide hydrogenation over nickel catalysts, 101 1-1 5 Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 Oxidation of propene on mixed oxides of copper and cobalt, 1337-45 Electrical behavior of powdered tin-antimony mixed oxide catalysts, 1346-55 ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Synergistic effect of composite catalysts on the direct hydrogenation of carbon, 1495-506 Selectivity in ethylene dimerization over supported nickel oxide catalysts, 1 5 13-20 Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 650"C, 1593-605 Electron spin resonance studies of formation of sulfur dioxide anion radicals on calcium oxide, 161 3-18 Kinetics of chlorine oxide radical reactions using modulated photolysis.Part 1. Disproportionation of chlorine monoxide in the dichlorine photosensitized decomposition of ozone, 1635-47 Kinetics of chlorine oxide radicals using modulated photolysis. Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Surface structure and surface states in magnesium oxide powders, 1769-79 Infrared spectrum and surface reaction of nitric oxide adsorbed on silicon films, 1788-97 Reactions involving electron transfer at semiconductor surfaces.Part 9. Oxygen labelling study of Mechanism and kinetics of the chain reaction in hydrogen peroxide + nitrous oxide + carbon monoxide Study of the interaction of nitric oxide with copper( 100) and copper( 1 1 1) surfaces using low energy Evaluation of semiconducting tin oxide as an electrocatalyst support, 2 165-76 Decomposition of nitrous oxide on high surface area calcium oxide-magnesium oxide catalysts Adsorption of 2,2'-bipyridyl on magnesium oxide and calcium oxide. Infrared spectra of neutral and Solid solution in the calcium oxide-manganese oxide system, 2285-94 Photoelectrochemical reactions of pigmentary titanium dioxide, 2507- 16 Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Photoelectrochemical reactions of pigmentary titanium dioxide with alcohols and aliphatic amines, 2526-34 Infrared spectroscopic investigation of sorption of nitric oxide by /%-metal phthalocyanines of the first transition period, 2594-600 Photolysis of thioketene S-oxides, 2624-7 Kinetics of chlorine oxide radical reactions using modulated photolysis.Part 3. Pressure and temperature dependence of the reaction: chlorine monoxide + nitrogen dioxide (+ M) + nitryl hypochlorite (+ Acidic and basic properties of mixed tin-antimony oxides, 2762-7 Study of the oxidative reactions of butenes over mixed tin-antimony oxides, 2768-80 Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron Hydrogen oxidation catalyzed by X zeolite containing transition metal ions, 3 14-22 Influence of activation conditions of cerium-X zeolite on its oxidizing properties as shown by infrared Kinetics of cobalt(II1) + iron(I1) reaction by platinum4ectrode chronoamperometry in aqueous Kinetics of oxidation of bromide ions by aquacerium(1V) ions in aqueous perchlorate media, 8 16-24 Adsorption of water vapor on evaporated germanium films.An infrared study, 962-70 Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in 787-802 photoelectron spectroscopy, 984-6 of alkylpyridines, 1356-70 1487-94 photooxidation of C3 and C4 alcohols over rutile, 2000-13 systems, 2048-59 electron diffraction and electron spectroscopy, 2 143-59 investigated by electron spin resonance spectroscopy, 21 77-87 anionic surface species, 2 188-98 M), 2649-61 spectroscopy and gravimetry, 2839-50 OXIDN and electron spin resonance spectroscopies, 335-4 1 perchloric acid solution, 473-6 zeolite Y, 1 19&206J.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) 5 1 OXIDN(contd) Oxidation of propene on mixed oxides of copper and cobalt, 133745 Electrical behavior of powdered tin-antimony mixed oxide catalysts, 134655 ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation Oxidation of isobutyraldehyde in aged boric-acidxoated vessels, 143346 Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 65OoC, 1593-605 Electron spin resonance studies of formation of sulfur dioxide anion radicals on calcium oxide, 16 I 3- 18 Mechanism for incorporation of oxygen in vapor-phase selective oxidation of isobutene, butadiene and Reactions involving electron transfer at semiconductor surfaces.Part 9. Oxygen labelling study of Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Evaluation of semiconducting tin oxide as an electrocatalyst support, 21 65-76 Characterization of electrooxidation catalysts prepared by ion-exchange of platinum salts with surface Chloride ion adsorption effects in the recombinationxontrolled kinetics of anodic chlorine evolution at Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Study of the methanol oxidation reaction on platinum using the potential-step technique, 2535-43 Oxygen electrodes in fused salts.Potentiometric and x-ray photoelectron spectroscopic (ESCA) findings on the system nickel dicarbonyl + dioxygen/carbonate ion in molten nitrates, 2628-37 Reactions of oxygenated radicals in the gas phase. Part 6. Reactions of isopropyl peroxy and isopropoxy radicals, 2678-87 Study of the oxidative reactions of butenes over mixed tin-antimony oxides, 2768-80 Effect of anion basicity in influencing the acidity function of concentrated aqueous solutions of oxyacids Some unusual properties of activated and reduced silver-sodium A zeolites, 56-64 Rate constants for hydrogen + oxygen system, and for hydrogen atoms and hydroxyl radicals + alkanes 140-54 Reactions involving electron transfer at semiconducting surfaces.Part 8. Room temperature activity of prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, Electron spin resonance and volumetric investigations of oxygen adsorption on high surface area cobalt oxide-magnesium oxide, 533-44 Isotopic exchange of molecular oxygen with oxygen of vanadium oxide (V205) modified with oxides of alkaline earth metals, 6914 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces. Part 1. Role of hydroxyl groups in photoadsorption, 736-47 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces. Part 2. Study of radical intermediates by electron paramagnetic resonance, 748-6 1 Kinetics of chemiionization in atomic oxygen + nitrogen mixtures, 1301-1 1 Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Reaction of tert-butyl radicals with hydrogen and with oxygen, 1458-64 Kinetics of chlorine oxide radicals using modulated photolysis.Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Oxygen chemisorption by y-alumina phthalocyaninato cobalt(I1). Influence of the surface species y- alumina on dioxygen ion fixation, 1857-67 Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Decomposition of nitrous oxide on high surface area calcium oxide-magnesium oxide catalysts investigated by electron spin resonance spectroscopy, 2177-87 Detection of formyl radicals in low pressure methane + oxygen flames, 2423-32 Oxygen electrodes in fused salts.Potentiometric and x-ray photoelectron spectroscopic (ESCA) findings on the system nickel dicarbonyl + dioxygen/carbonate ion in molten nitrates, 2628-37 Kinetics of chlorine oxide radical reactions using modulated photolysis. Part 1. Disproportionation of chlorine monoxide in the dichlorine photosensitized decomposition of ozone, 1635-47 Radiotracer study of adsorption during catalysis. Carbon-I 4 chlorobenzene hydrogenation over palladium on silica, 1798-14 Hydrogenation of acetylene over supported metal catalysts. Part 3. Carbon-I4 tracer studies of the effects of added ethylene and carbon monoxide on the reaction catalyzed by silica-supported palladium, rhodium and iridium, 1900-1 1 oxide groups on carbon, 23 12-24 of alkylpyridines, 135670 furan over various bismuth catalysts, 1757-68 photooxidation of C3 and C, alcohols over rutile, 2000-13 oxide groups on carbon, 23 12-24 platinum electrodes, 2454-72 OXY ACID 1606-12 OXYGEN 305- 1 3 * OZONE PALLADIUM PAPER Characterization of electrooxidation catalysts prepared by ion-exchange of platinum salts with surface52 J.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) PARTITION Behavior of simple ions in the system cellulose-anionic dye-simple electrolyte. Part 1. Partition Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 65OoC, 1593-605 Attachment of spherical particles to surface of a pendant drop and tension of the wetting perimeter, 1-6 Chlorine-35 nuclear quadrupole resonance studies of pentachlorophenol-amine hydrogen-bonded Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Hydrogenation of alkenes over supported gold, 385-94 Hydrogenation over boehmitely-alumina catalysts, 395-405 Effect of anion basicity in influencing the acidity function of concentrated aqueous solutions of oxyacids Free radical addition to olefins.Part 25. Addition of perfluoretert-butyl radicals to fluoro olefins, Ultraviolet spectrum of periodate ion, 63 1-5 Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and Effect of hexamethylphosphoric triamide on the living radical polymerization initiated by aged Mechanism and kinetics of the chain reaction in hydrogen peroxide + nitrous oxide -t carbon monoxide Reactions of oxygenated radicals in the gas phase.Part 6. Reactions of isopropyl peroxy and isopropoxy Enthalpy-entropy relations in the fluid kinetics of aqueous solvent systems, 942-52 Particle adhesion and removal in model systems. Part 1. Monodispersed chromium hydroxide on glass, Kinetics of polymorphic transitions in tetrahedral structures. Part 1. Experimental methods and the Inorganic photophysics in solution. Part 3. Temperature activation of decay processes in the luminescen- Acid ionization of solvent, acetic acid, and 4-substituted phenols in tert-butanol + water mixtures, Mixed ligand complexes in solution of manganous ion. Nuclear relaxation rate analysis, 2433-8 Sequence studies in liquid phase hydrocarbon oxidation.Part 5. Temperature dependence of hydroperox- Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane. Exchange of hydrogen deuteride with boron phosphate, 503-1 2 Charge determination at calcium salt/aqueous solution interface, 1034-9 Surface characterization of nickel boride and nickel phosphide catalysts by x-ray photoelectron Kinetics of heterogeneous electron transfer on dicyanobis(tertiary phosphine)nickel complexes, 1 330-6 High resolution proton nuclear magnetic resonance studies of interaction between deoxyhemoglobin and Effect of anion basicity in influencing the acidity function of concentrated aqueous solutions of oxyacids equilibriums, 705-1 6 PASSIVATION PENDANT PENTACHLOROPHENOL complexes, 1587-92 solutions.A pulse radiolysis study, 1674-87 PENTANOL PENTENE PERCHLORIC 1606-12 PERFLUOROBUTYL 1040-9 PERIODATE PEROXIDE mechanism of polymerization, 774-86 chromium ion + benzoyl peroxide, 1821-9 systems, 2048-59 radicals, 2678-87 PEROXY PERTURBATION PH 65-78 PHASE transition y +- fi lithium zinc silicate(LizZnSi04), 374-84 ce of tris(2,2'-bipyridine)osmium(II) and tris( 1 ,I 0-phenanthroline)osmium(II) ions, 353-62 PHENANTHROLINE PHENOL 278 1-97 PHENYLALANINE PHENYLETHYL ide-alcohol and hydroperoxide-ketone transitions, 236-46 Evidence for an intermediate radical cation, 914-20 PHENYLMETHYL PHOSPHATE PHOSPHIDE spectroscopy, 2027-39 PHOSPHINE PHOSPHOGLYCERATE small molecules. Dithionite and diphosphoglycerate, 285 1-64 PHOSPHORICJ.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1 979) 53 PHOSPHORIC(contd) PHOTOADSORPTION 1606-12 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces. Part Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces. Part Effect of F -+ M phototransformation on aquoluminescence from y-irradiated sodium chloride, 844-9 Photoinitiated pblymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and Kinetics of solid state thermal monomerization of 9-cyanoanthracene photodimer and photodimerizatio- Reactions involving electron transfer at semiconductor surfaces. Part 9. Oxygen labelling study of Reactions of oxygenated radicals in the gas phase.Part 6. Reactions of isopropyl peroxy and isopropoxy Kinetics of solid state thermal monomerization of 9kyanoanthracene photodimer and photodimerizatio- Photoelectrochemical reactions of pigmentary titanium dioxide, 2507-1 6 Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Photoelectrochemical reactions of pigmentary titanium dioxide with alcohols and aliphatic amines, Kinetics of chlorine oxide radical reactions using modulated photolysis. Part I . Disproportionation of chlorine monoxide in the dichlorine photosensitized decomposition of ozone, 163547 Kinetics of chlorine oxide radicals using modulated photolysis. Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Photolysis of thioketene S-oxides, 2624-7 Reactions of oxygenated radicals in the gas phase.Part 6. Reactions of isopropyl peroxy and isopropoxy Reactions involving electron transfer at semiconductor surfaces. Part 9. Oxygen labelling study of Reactions of oxygenated radicals in the gas phase. Part 6. Reactions of isopropyl peroxy and isopropoxy Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and Photolysis of methyl- and acetyl-manganese pentacarbonyls in methyl methacrylate. Initiation of Photoelectrochemical reactions of pigmentary titanium dioxide, 2507-1 6 Oxygen chemisorption by y-alumina phthalocyaninato cobalt(I1). Influence of the surface species y- Oxygen chemisorption by y-alumina phthalocyaninato cobalt(I1).Influence of the surface species y- Infrared spectroscopic investigation of the formation of adducts of iron-phthalocyanine with hydrazine Infrared spectroscopic investigation of sorption of nitric oxide by p-metal phthalocyanines of the first Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron Thermodynamics of aggregation of some tertiary n-alkylammonium picrates in benzene solutions, Dissociation kinetics of picric acid and dipicrylamine in methanol. Steric effect on a proton transfer rate, Dissociation kinetics of picric acid and dipicrylamine in methanol. Steric effect on a proton transfer rate, Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, 1.Role of hydroxyl groups in photoadsorption, 736-47 2. Study of radical intermediates by electron paramagnetic resonance, 748-6 1 PHOTOANNEALING PHOTOCHEM mechanism of polymerization, 774-86 n of 9-cyanoanthracene, 1 1 11-19 photooxidation of C3 and C4 alcohols over rutile, 2000-13 radicals, 2678-87 PHOTODIMERIZATION n of 9-cyanoanthracene, 1 I I 1-19 PHOTOELECTROCHEM 2526-34 PHOTOLYSIS radicals, 2678-87 photooxidation of C3 and C4 alcohols over rutile, 2000-1 3 radicals, 2678-87 mechanism of polymerization, 774-86 polymerization and formation of methyl(2-methyl-4-oxopentanoate), 2562-75 PHOTOOXIDN PHOTOPOLYMN PHOTOVOLTAIC PHTHALICYANINATOCOBALT alumina on dioxygen ion fixation, 1857-67 PI~THALOCYANINATOCOBALT alumina on dioxygen ion fixation, 1857-67 and ammonia, 2587-93 transition period, 2594-600 spectroscopy and gravimetry, 2839-50 PHTHALOCYANINE PICRATE 825-34 2 I 3742 PICRY LAMINE 2 13742 PIEZORESISTANCE54 J.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) PI EZORESISTANCE(contd) PITZER dichromates and thiocyanates, 2735-43 water at 298.15 K and correlations between Harned and Pitzer equations, 1371-9 Redetermination of mean ionic activity coefficients for the system hydrochloric acid-potassium chloride- Magnetic properties of supported platinum-iron alloys, 257-61 Hydrogenation of alkenes over supported gold, 385-94 Ring-disk electrodes. Part 19. Adsorption studies at low frequency a.c., 1623-34 Evaluation of semiconducting tin oxide as an electrocatalyst support, 2 165-76 Characterization of electrooxidation catalysts prepared by ion-exchange of platinum salts with surface Chloride ion adsorption effects in the recombination-controlled kinetics of anodic chlorine evolution at Diffusion of chemisorbed hydrogen in a platinum zeolite, 2473-80 Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Photoelectrochemical reactions of pigmentary titanium dioxide with alcohols and aliphatic amines, Study of the methanol oxidation reaction on platinum using the potential-step technique, 253543 Hydrogenation of ethylene on metal electrodes. Part 4.Electroreduction of ethylene at a platinum Synergistic effect of composite catalysts on the direct hydrogenation of carbon, 1495-506 Hydrogenation of acetylene over supported metal catalysts.Part 3. Carbon-14 tracer studies of the effects of added ethylene and carbon monoxide on the reaction catalyzed by silica-supported palladium, rhodium and iridium, 1900-1 1 Hydrogenation of acetylene over supported metal catalysts. Part 3. Carbon-14 tracer studies of the effects of added ethylene and carbon monoxide on the reaction catalyzed by silica-supported palladium, rhodium and iridium, 1900-1 I PLATINUM oxide groups on carbon, 23 12-24 platinum electrodes, 2454-72 2526-34 electrode, 2638-48 PLATINUM GROUP POISON POISONING POLARIZATION Effect of preadsorbed ammonia on isomerization of but-lkne on sodium hydrogen-Y zeolite, I 150-60 Thermoelectric effects and dielectric polarization in biopolymers, 323--34 Membrane polarization at high current densities, 463-72 Effect of sodium carboxymethylcellulose on the reactivity of hydrated electrons and cystamine anion Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Radiation chemistry of cobdlt(l1) nitrilotnacetate in aqueous solutions, 2089-99 On Ritchie’s equation for the analysis of kinetics of adsorption of gases on solids, 477-8 Effect of electrolytes on solution behavior of water soluble macromolecules, 993- 1000 Thermodynamic properties of polyelectrolyte solutions determined from studies of Donnan equilibriums Electrokinetic study of surfactant adsorption, 669-78 Effect of sodium carboxymethylcellulose on the reactivity of hydrated electrons and cystamine anion Thermodynamic properties of polyelectrolyte solutions determined from studies of Donnan equilibriums Electrokinetic study of surfactant adsorption, 669-78 Observation of weak interactions between polystyrene particles and textile fibers, 238&92 Precipitation of calcium surfactants.Part 3, 2 126-36 Catalytic reactions with deuterium of several polymethylcycloalkanes o n cobalt films, 1320-9 Correlation between topological features and molar volumes of n-alkanes and excess volumes of their Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and Effect of hexamethylphosphoric triamide on the living radical polymerization initiated by aged POLARON radicals in aqueous solutions, 1 142-9 POLEMIC POLY ACRY LAMIDE POLYACRY LATE 1207-1 4 POLYAMIDE POLYELECTROLYTE radicals in aqueous solutions, 1 142-9 1 207- 1 4 POLYESTER POLYMER POLYMETHY LCYCLOALKANE POLYMETHYLENE binary mixtures, 1 13241 mechanism of polymerization, 774-86 chromium ion + benzoyl peroxide, 182 1-9 POLYMNJ.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) 55 POLY MN(contd) POLYMORPHISM Theory of compartmentalized free-radical polymerization reactions. Part 3,2332-58 Kinetics of polymorphic transitions in tetrahedral structures. Part 1. Experimental methods and the Effect of electrolytes on solution behavior of water soluble macromolecules, 993-1000 Observation of weak interactions between polystyrene particles and textile fibers, 2386-92 Quaternization of poly(4-vinylpyridine). Kinetic and viscometric measurements, 1728-35 Effect of electrolytes on solution behavior of water soluble macromolecules, 993-1000 Scanning studies on capillary condensation and evaporation of nitrogen.Part 1. Apparatus and Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending Behavior of simple ions in the system cellulose-anionic dye-simple electrolyte. Part 1. Partition Catalytic deamination on solid surfaces. Part 1. Kinetic studies by stopped-flow gas chromatography, Electron transfer reactions involving porphyrins and carotenoids, 2869-72 Site group interaction effects in zeolite Y. Part 3. Potassium-sodium equilibrium, 440-5 Ultraviolet spectrum of periodate ion, 631-5 Electrokinetic study of surfactant adsorption, 669-78 Structural analysis of some molten materials by x-ray diffraction.Part 6. Manganese chloride-lithium or Structural analysis of molten potassium sulfate, 1 169-78 Redetermination of mean ionic activity coefficients for the system hydrochloric acid-potassium chloride- Structure of molten potassium chloride, 1477-86 Flame inhibition by potassium, 2377-85 Adsorption of ethylene glycol and its higher homologs at the mercury/aqueous solution interface, Kinetics of cobalt(II1) + iron(I1) reaction by platinum4ectrode chronoamperometry in aqueous Electrokinetic study of surfactant adsorption, 669-78 Charge determination at calcium salt/aqueous solution interface, 1034-9 Calorimetric enthalpy of reaction of cobalt(II1) with iron(I1) in aqueous perchloric acid, and the standard Viscosities of trichlorofluoromethane, chlorotrifluoromethane, dichlorofluoromethane, chlorodifluoro- Adsorption of ethylene glycol and its higher homologs at the mercury/aqueous solution interface, Oxygen electrodes in fused salts.Potentiometric and x-ray photoelectron spectroscopic (ESCA) findings Thermodynamic properties of polyelectrolyte solutions determined from studies of Donnan equilibriums Surface structure and surface states in magnesium oxide powders, 1769-79 Precipitation of calcium surfactants. Part 2,211625 Precipitation of calcium surfactants. Part 3,2126-36 Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Dependence of the viscosity of liquid 3,34iethylpentane, mercury, argon and methane on thermal Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, Kinetics of'reaction of magnesium vapor with silica, 2807-1 8 Reaction probabilities and threshold energy in the abstraction reaction between hydrogen atoms and transition y -+ fi lithium zinc silicate(Li2ZnSi04), 374-84 POLYOXY ETHYLENE POLYSTYRENE POLYVINY LPYRIDINE PO LY VINY LPY R ROLIDONE PORE calculation method, 36-41 and descending scanning curves within B-type hysteresis loops, 42-55 equilibriums, 705-1 6 POROUS 248 1-95 PORPHYRIN POTASSIUM potassium chloride, 1 161-8 water at 298.15 K and correlations between Harned and Pitzer equations, 1371-9 2576-86 POTENTIAL perchloric acid solution, 473-6 electrode potential for cobalt(IIIHII), 1268-9 methane, and trifluoromethane from 373 to 570 K, 1752-6 2576-86 on the system nickel dicarbonyl + dioxygen/carbonate ion in molten nitrates, 2628-37 POTENTIOMETRY 1 207- 1 4 POWDER PPTN PRESSURE reactions.Volume changes on activation and reaction, 172-9 1 pressure over a wide PVT range, 2295-302 dichromates and thiocyanates, 2735-43 PROBABILITY56 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) PROBABILITY(contd) PROBE neopentane, 1619-22 Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 Mixed ligdnd complexes in solution of manganous ion. Nuclear relaxation rate analysis, 2433-8 Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 Wall-less reactor studies. Part 2. Propane pyrolysis, 1 101-10 Influence of lithium, sodium, and potassium cation concentrations in X and Y zeolites on isotherms and heats of adsorption of propane and water, 2662-77 Role of acetaldehyde in propane combustion, 2798-806 Single-pulse shock tube studies of hydrocarbon pyrolysis.Part 7. Pyrolysis of propane, 28 19-26 Adsorption of ethylene glycol and its higher homologs at the mercury/aqueous solution interface, Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-9 I Dynamic properties of n-alkyl and s-alkyl intermediates in reactions of simple alkenes with hydrogen on Oxidation of propene on mixed oxides of copper and cobalt, 1337-45 Electrical behavior of powdered tin-antimony mixed oxide catalysts, 1346-55 Reactions of oxygenated radicals in the gas phase. Part 6.Reactions of isopropyl peroxy and isopropoxy Thermal ion-molecule reactions in oxygen-containing molecules. Inverse isotope effect in the ion- Reactions of oxygenated radicals in the gas phase. Part 6. Reactions of isopropyl peroxy and isopropoxy Thermodynamic properties for transfer of cations from propylene carbonate to dimethylsulfoxide and to High temperature pyrolysis of ethane and propylene. 13904 Ultrasonic behavior of the system water + propylene glycol, 2067-72 Crystal growth kinetics of globular proteins. Lysozyme and insulin, 2753-6 1 Quantum theory of kinetic isotope effects in proton transfer reactions, 205-26 Heat capacity changes in proton addition to the nitrogen of saturated organic molecules in water. Effects Dissociation kinetics of picric acid and dipicrylamine in methanol.Steric effect on a proton transfer rate, Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane. Pulse radiolysis study of acrylonitrile in aqueous solution, 1050-66 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Absorption and emission studies of solubilization in micelles. Part 5. Pyrene-3-sodium sulfonate PROMAZINE PROPANE PROPANEDIOL 2576-86 PROPANOL PROPENE molybdenum disulfide catalyst, 7- 21 PROPOXY radicals, 2678-87 PROPYL molecule dimerization reactions of protonated propyl acetate ions, 525-32 radicals, 2678-87 PROPYLENE propylene carbonate + dimethylsulfoxide mixtures, 227-35 PROTEIN PROTON PROTON ATION of solvation, 363--9 2 137-42 PULSE Evidence for an intermediate radical cation, 914-20 solutions.A pulse radiolysis study, 1674-87 PYRENESULFONATE solubilized in didodecyldimethylammonium bromide inverted micelles in benzene, 550-60 Surface acidity of q-alumina. Part 1. Pyridine chemisorption at room temperature, 27 1-88 Surface acidity of q-alumina. Part 2. Interaction of pyridine with other adsorbates, 289-304 Influence of activation conditions of cerium-)< zeolite on its oxidizing properties as shown by infrared Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-9 I Ultraviolet excitation of Raman spectra of pyridines adsorbed on oxides, 121 5-2 I Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation Selectivity in ethylene dimerization over supported nickel oxide catalysts, 15 13-20 Acidic and basic properties of mixed tin-antimony oxides, 2762-7 PYRIDINE and electron spin resonance spectroscopies, 33541 of alkylpyridines, 1356-70J.C.S. FARADAY I SUBJECT INDEX VOL.75 (1979) 57 PY RIDINE(contd) PY RIDINEAZODIMETHY LANILINE PYROLYSIS Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron Kinetic investigation and numerical analysis of a micelle-catalyzed metal complex formation, 21 99-2 10 Pyrolysis of benzene, 652-62 Wall-less reactor studies.Part 1. Ethane pyrolysis, 1089-100 Wall-less reactor studies. Part 2. Propane pyrolysis, 1 101-10 High temperature pyrolysis of ethane and propylene, 1390-4 Thermal unimolecular decomposition of I , I-difluorocyclobutane, 25566 1 Wall-less reactor studies. Part 3. n-Butane pyrolysis, 2688-93 Single-pulse shock tube studies of hydrocarbon pyrolysis. Part 7. Pyrolysis of propane, 28 19-26 Ion condensation model and nuclear magnetic resonance studies of counterion binding in lyotropic liquid Quantum theory of kinetic isotope effects in proton transfer reactions, 205-26 Ion-solvent interaction of some tetraalkylammonium ions in sulfolane from viscosity data, 2325-3 1 Quaternization of poly(4-vinylpyridine). Kinetic and viscometric measurements, 1728-35 Inorganic photophysics in solution. Part 2.Temperature activation of decay processes in the luminescen- Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and Mechanism of interactions between aquametallic complexes and the photoexcited aquouranyl(V1) ion, Chemical quenching by water of the photoexcited uranyl ion in aqueous acidic solution, 2273-84 Electron transfer reactions involving porphyrins and carotenoids, 2869-72 Structure of molten potassium chloride, 1477-86 Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 Reactions of hydroxymethyl radicals with some solutes in methanol.Pulse radiolysis study, 44662 Reaction between hydrogen atoms and methane, 835-43 Free radical addition to olefins. Part 25. Addition of perfluorwtert-butyl radicals to fluoro olefins, Effect of sodium carboxymethylcellulose on the reactivity of hydrated electrons and cystamine anion Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Reaction of tert-butyl radicals with hydrogen and with oxygen, 1458-64 Electron spin resonance studies of the radicals formed from carbon-nitroso compounds and olefins. Part 1. Nitrosodurene and fluoroolefins, 1521-30 Electron spin resonance studies of formation of sulfur dioxide anion radicals on calcium oxide, 16 1 3- 18 Kinetics of chlorine oxide radical reactions using modulated photolysis.Part 1. Disproportionation of chlorine monoxide in the dichlorine photosensitized decomposition of ozone, 163547 Kinetics of chlorine oxide radicals using modulated photolysis. Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Electron-transfer reactions of nitroxyl radicals with one-electron reduced quinones and viologens, Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Decomposition of nitrous oxide on high surface area calcium oxide-magnesium oxide catalysts Theory of compartmentalized free-radical polymerization reactions. Part 3,2332-58 Flame inhibition by potassium, 2377-85 Detection of formyl radicals in low pressure methane + oxygen flames, 2423-32 Kinetics of gas phase addition reactions of trichlorosilyl radicals.Part 2. Additions to I-olefins, 2617-23 Kinetics of chlorine oxide radical reactions using modulated photolysis. Part 3. Pressure and temperature dependence of the reaction: chlorine monoxide + nitrogen dioxide ( + M) + nitryl hypochlorite (+ Reactions of oxygenated radicals in the gas phase. Part 6. Reactions of isopropyl peroxy and isopropoxy radicals, 2678-87 Wall-less reactor studies. Part 3. n-Butane pyrolysis, 2688-93 Electron transfer reactions involving porphyrins and carotenoids.2869-72 Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 spectroscopy and gravimetry, 2839-50 QUADRUPOLE crystals, 663-8 QUANTUM QUATERNARY QUATERNIZATION QUENCHING ce of uranyl[U022+] ion, 342-52 mechanism of polymerization, 774-86 2252-72 RADIAL RADICAL 1040-9 radicals in aqueous solutions, 1142-9 1912-18 investigated by electron spin resonance spectroscopy, 2 177-87 M), 2649-61 RADIOLYSIS58 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) RADIOLYSIS(contd) y-Radiolysis of methane adsorbed on y-alumina. Part 3. Influence of additives on product desorption, Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexdne. Pulse radiolysis study of acrylonitrile in aqueous solution, 1050-66 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Molecular reorientation and molecular association.Study by Raman depolarized light scattering. Ultraviolet excitation of Raman spectra of pyridines adsorbed on oxides, 1215-21 Rotating disk dissolution rates of ionic solids. Part 3. Natural and synthetic ilmenite, 97 1-83 Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 Rate constants for hydrogen + oxygen system, and for hydrogen atoms and hydroxyl radicals + alkanes Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Proton exchange in aqueous solutions of glucose.Hydration of carbohydrates, 262-70 Reactions of hydroxymethyl radicals with some solutes in methanol. Pulse radiolysis study, 44662 Ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexidexontaining micelles in the system aerosol-OT + water + heptane, 48 1-96 Reaction between hydrogen atoms and methane, 835-43 Kinetics of the gas-phase reaction between iodine and trimethylsilane and the bond dissociation energy of the hydrogen-silicon bond. Part 2, 1 126-3 1 Effect of sodium carboxymethylcellulose on the reactivity of hydrated electrons and cystamine anion radicals in aqueous solutions, 1 142-9 Calorimetric enthalpy of reaction of cobalt(II1) with iron(I1) in aqueous perchloric acid, and the standard electrode potential for cobalt(III)-(II), 1268-9 Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Reaction of tert-butyl radicals with hydrogen and with oxygen, 1458-64 Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 650°C 1593-605 Reaction probabilities and threshold energy in the abstraction reaction between hydrogen atoms and neopen tane, 16 19-22 Kinetics of chlorine oxide radicals using modulated photolysis.Part 2. Chlorine monoxide and chlorine dioxide radical kinetics in the photolysis of dichlorine + dioxygen + dinitrogen mixtures, 1648-66 Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Mechanism and kinetics of the chain reaction in hydrogen peroxide + nitrous oxide + carbon monoxide systems, 2048-59 Reactions of oxygenated radicals in the gas phase.Part 6. Reactions of isopropyl peroxy and isopropoxy radicals, 2678-87 Kinetics of reaction of magnesium vapor with silica, 2807-18 Wall-less reactor studies. Part 1. Ethane pyrolysis, 1089-100 Wall-less reactor studies. Part 2. Propane pyrolysis, 1101-10 Wall-less reactor studies. Part 3. n-Butane pyrolysis, 2688-93 Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Reactions of positive halogen ions in flames and their rate coefficients, 195M8 Flame inhibition by potassium, 2377-85 Adsorption and recombination of hydrogen atoms on glass surfaces.Part 1. Method of study and the Chloride ion adsorption effects in the recombination-controlled kinetics of anodic chlorine evolution at Magnetic determination of metallic nickel particles dispersed on X and Y zeolite structures, 165-7 I Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in Calorimetric enthalpy of reaction of cobalt(II1) with iron(I1) in aqueous perchloric acid, and the standard Kinetics of heterogeneous electron transfer on dicyanobis(tertiary ph0sphine)nickel complexes, 1330-6 Temperature programmed reduction of copper ions in zeolites, 1688-99 Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Radiation chemistry of cobalt(I1) nitrilotriacetate in aqueous solutions, 2089-99 Spectrophotometric investigations in aqueous solution at elevated temperatures.Kinetics of the 192-204 Evidence for an intermediate radical cation, 914-20 solutions. A pulse radiolysis study, 1674-87 RAMAN 1179-83 RATE REACTION 140-54 reactions. Volume changes on activation and reaction, 172-9 1 REACTOR RECOMBINATION mechanism of recombination at 77 and 273 K, 2439-53 platinum electrodes, 2454-72 REDN zeolite Y, 1 196-206 electrode potential for cobalt(III)-(II), 1268-9J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 59 R EDN(contd) reduction of copper(I1) to copper(1) at 473 K in aqueous solution in the presence of 2,2'-bipyridyl, 2496-506 Hydrogenation of ethylene on metal electrodes.Part 4. Electroreduction of ethylene at a platinum Oxidation-reduction behavior of highly dispersed ruthenium in ruthenium sodium Y zeolite, 1395-402 Heterogeneous catalysis in solution. Part 17. Kinetics of oxidation-reduction reactions catalyzed by Magnetic determination of metallic nickel particles dispersed on X and Y zeolite structures, 165-7 I Magnetic determination of metallic nickel particles dispersed on X and Y zeolite structures, 165-7 I Use of internal and external references in nuclear magnetic resonance determinations of association Ultraviolet spectrum of periodate ion, 63 1-5 Surface structure and surface states in magnesium oxide powders, 1769-79 Structure and dynamics of ammonia adsorbed on graphitized carbon black.Part 4. Nuclear magnetic resonance spectra, 1570-86 Application of the environmental relaxation model to the ideal and nonideal glass-forming molten mixtures, 1830-6 Mixed ligand complexes in solution of manganous ion. Nuclear relaxation rate analysis, 2433-8 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and nickel bromide solutions. Part 1. Application of solvent nuclear magnetic relaxation, 2700-1 1 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and nickel bromide solutions. Part 2. Application of halide nuclear magnetic relaxation, 271 2-34 Molecular reorientation and molecular association. Study by Raman depolarized light scattering, Thermodynamics of ion exchange of trivalent cobalt or chromium complex with cerium( 111) ions on Characterization of water-containing reversed micelles by viscosity and dynamic light scattering methods Ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexide-containing micelles in the system aerosol-OT + water + heptane, 48 1-96 Rhodium catalysts prepared by attachment of hexarhodium hexadecacarbonyl onto chemically modified silicas.Characterization of the infrared spectra in the carbonyl stretching region, 1888-99 Hydrogenation of acetylene over supported metal catalysts. Part 3. Carbon-1 4 tracer studies of the effects of added ethylene and carbon monoxide on the reaction catalyzed by silica-supported palladium, rhodium and iridium, 1900- 1 1 electrode, 2638-48 REDOX electron transfer through the solid.An electrochemical treatment, 1507-1 2 REDUCED REDUCIBILITY REF constants of weak molecular complexes, 1222-5 REFLECTANCE RELAXATION REORIENTATION I 179-83 RESIN cation exchange resin, 925-34 REVERSE 132-9 RHODIUM RING DISK ROD Ring-disk electrodes. Part 19. Adsorption studies at low frequency ax., 1623-34 Capillary phenomena. Part 8. Bridge meridians for fluid interfaces with solids in cylindrical vessels: computation, Bessel approximations and comparison of properties of finite and infinite systems, 1929-39 ROTATING ROTATION Rotating disk dissolution rates of ionic solids. Part 3. Natural and synthetic ilmenite, 971-83 Magnetooptical rotation studies of liquid mixtures. Part 3.Specific interactions in mixtures of carbon Electron spin resonance studies of the radicals formed from carbon-nitroso compounds and olefins. Part Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 Oxidation-reduction behavior of highly dispersed ruthenium in ruthenium sodium Y zeolite, I 395-402 Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 1. Reactions involving electron transfer at semiconductor surfaces. Part 9. Oxygen labelling study of Particle adhesion and removal in model systems. Part I . Monodispersed chromium hydroxide on glass, tetrachloride with aliphatic, alicyclic and aromatic hydrocarbons and amines, 370-3 I .Nitrosodurene and fluoroolefins, 1521-30 RUTHENIUM Equilibriums and affinities, 1969-83 photooxidation of C3 and C4 alcohols over rutile, 2000-13 'RUTILE SALT60 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) S A LT(con td) 65-78 Behavior of simple ions in the system cellulose-anionic dye-simple electrolyte. Part 1. Partition Behavior of simple ions in the system cellulose-anionic dye-simple electrolyte. Part 2. Diffusion Calculation of excess enthalpies for binary mixtures of liquid metal halides, 729-35 Effect of F -+ M phototransformation on aquoluminescence from y-irradiated sodium chloride, 844-9 Kinetics of thermal decomposition of 4,4'-azobis-(4--cyanopentanoic acid) and its salts in aqueous Effect of electrolytes on solution behavior of water soluble macromolecules, 993- loo0 Charge determination at calcium salt/aqueous solution interface, 1034-9 Quaternization of poly(4-vinylpyridine). Kinetic and viscometric measurements, I 728-35 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl Application of the environmental relaxation model to the ideal and nonideal glass-forming molten Scaled-particle theory and salting constant of tetrahydrofuran in various aqueous electrolyte solutions at Ion-solvent interaction of some tetraalkylammonium ions in sulfolane from viscosity data, 2325-3 I Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, Scaled-particle theory and salting constant of tetrahydrofuran in various aqueous electrolyte solutions at Characterization of water-containing reversed micelles by viscosity and dynamic light scattering methods Structure and dynamics of ammonia adsorbed on graphitized carbon black.Part 3. Neutron quasielastic Thermoelectric effects and dielectric polarization in biopolymers, 323-34 Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Photoelectrochemical reactions of pigmentary titanium dioxide, 2507- 16 Electron-transfer reactions of nitroxyl radicals with one-electron reduced quinones and viologens, Scaled-particle theory and salting constant of tetrahydrofuran in various aqueous electrolyte solutions at Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 High temperature pyrolysis of ethane and propylene.I3904 Single-pulse shock tube studies of hydrocarbon pyrolysis. Part 7. Pyrolysis of propane, 28 19-26 Infrared studies of the adsorption of probe molecules onto the surface of hematite, 1259-67 Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending Heat of immersion of thermally treated silica gel, 646-5 I Ultraviolet excitation of Raman spectra of pyridines adsorbed on oxides, 1215-21 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared Radiotracer study of adsorption during catalysis. Carbon-14 chlorobenzene hydrogenation over Rhodium catalysts prepared by attachment of hexarhodium hexadecacarbonyl onto chemically modified Infrared study of the adsorption of cyclohexanone on silica immersed in 2,2,4-trimethylpentane, 22 1 1-20 Kinetics of reaction of magnesium vapor with silica, 2807-18 Infrared study of the adsorption of ethyl acetate on silica immersed in carbon tetrachloride, 2865-8 Kinetics of polymorphic transitions in tetrahedral structures. Part 1.Experimental methods and the equilibriums, 705-16 coefficients, 7 17-28 solution, 935-41 sulfoxide + water mixtures, 1780-7 mixtures, 1830-6 298.15 K, 1837-48 dichromates and thiocyanates, 273543 SALTING 298.15 K, 1837-48 SCATTERING 132-9 and inelastic spectra, 1553-69 SEEBECK SELF DIFFUSION SEMICONDUCTOR SEM IQUINONE 1912-18 SETCHENOV 298.15 K, 183748 SHEAR SHOCK SILANE SILICA and descending scanning curves within B-type hysteresis loops, 42-55 by use of tetrakis-n-allylmolybdenum, 1465-76 palladium on silica, 1 798- 14 silicas. Characterization of the infrared spectra in the carbonyl stretching region, 1888-99 SILICATE transition y + fi lithium zinc silicate(Li2ZnSi04), 374-84 SILICON Kinetics of the gas-phase reaction between iodine and trimethylsilane and the bond dissociation energyJ.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) 61 SILICON(contd) SILVER of the hydrogen-silicon bond. Part 2, 1 126-3 1 Infrared spectrum and surface reaction of nitric oxide adsorbed on silicon films, 1788-97 Some unusual properties of activated and reduced silver-sodium A zeolites, 56-64 Thermodynamic properties for transfer of electrolytes from water to acetonitrile and to acetonitrile + water mixtures, 8 6 9 5 Chemisorption of carbon monoxide and hydrogen on silver-sodium mordenite, 109-1 8 Thermodynamic properties for transfer of cations from propylene carbonate to dimethylsulfoxide and to propylene carbonate + dimethylsulfoxide mixtures, 227-35 Hydrogenation of alkenes over supported gold, 385-94 Electrolyte solutions in liquid ammonia.Part 6. Mean molal activity coefficients and cation transference numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl sulfoxide + water mixtures, 178&7 Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, dichromates and thiocyanates, 273543 Sorption of hydrocarbons and water in silanated and unsilanated partial hydrogen-forms of zeolite Y, Search for strongly adsorbed carbon monoxide at copper single crystal surfaces using ultraviolet Selectivity in ethylene dimerization over supported nickel oxide catalysts, 15 13-20 Scanning studies on capillary condensation and evaporation of nitrogen.Part 1. Apparatus and Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in Some unusual properties of activated and reduced silver-sodium A zeolites, 5 6 6 4 Thermodynamic properties for transfer of cations from propylene carbonate to dimethylsulfoxide and to Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 Site group interaction effects in zeolite Y.Part 3. Potassium-sodium equilibrium, 44&5 Ultraviolet spectrum of periodate ion, 631-5 Behavior of simple ions in the system cellulose-anionic dye-simple electrolyte. Part 2. Diffusion Effect of F + M phototransformation on aquoluminescence from y-irradiated sodium chloride, 844-9 Effects of dehydration and x-ray irradiation on the structure of manganese sodium zeolite, 898-906 Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in Thermodynamic properties of polyelectrolyte solutions determined from studies of Donnan equilibriums Oxidation-reduction behavior of highly dispersed ruthenium in ruthenium sodium Y zeolite, 1 395402 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Temperature programmed reduction of copper ions in zeolites, 1688-99 Binding of n-alkyl sulfates to lysozyme in aqueous solution, 173644 Partial molar enthalpies in the sodium chloride + calcium chloride + water system.Application of the McKay-Perring method to thermal measurements, 1745-5 1 Scaled-particle theory and salting constant of tetrahydrofuran in various aqueous electrolyte solutions at 298.15 K, 183748 Kinetic investigation and numerical analysis of a micelle-catalyzed metal complex formation, 2 199--2 10 Micellar catalysis of metal complex formation.Part 2. Kinetics of the reaction between aquated nickel(2 +) and various neutral bidentate ligands in the presence of sodium dodecylsulfate micelles in aqueous solution, 2395-405 pressure over a wide PVT range, 2295-302 SILY LATION 2221-34 SINGLE photoelectron spectroscopy, 984-6 SITE SIZE calculation method, 36-41 and descending scanning curves within B-type hysteresis loops, 42-55 zeolite Y, 1 196206 SODIUM propylene carbonate + dimethylsulfoxide mixtures, 227-35 coefficients, 7 17-28 zeolite Y, 1 196206 1207-14 solutions. A pulse radiolysis study, 1674-87 SOFT SPHERE Dependence of the viscosity of liquid 3,34iethylpentane, mercury, argon and methane on thermal Electrokinetic study of surfactant adsorption, 669-78 Charge determination at calcium salt/aqueous solution interface, 1034-9 Capillary phenomena.Part 8. Bridge meridians for fluid interfaces with solids in cylindrical vessels: computation, Bessel approximations and comparison of properties of finite and infinite systems, SOL SOLID62 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) SOL1 D(contd) SOLN 1929-39 Solid solution in the calcium oxide-manganese oxide system, 2285-94 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part I . Experiment- Absorption of hydrogen by titdnium-cobah and titanium-nickel intermetallic alloys. Part 2. Thermodyn- Thermodynamics of aggregation of some tertiary n-alkylammonium picrates in benzene solutions, Mechanism of production and growth inhibition of collision crystallites, 92 1-4 Enthalpy-ntropy relations in the fluid kinetics of aqueous solvent systems, 942-52 Electrolyte solutions in liquid ammonia.Part 6. Mean molal activity coefficients and cation transference Effect of sodium carboxymethylcellulose on the reactivity of hydrated electrons and cystamine anion Solubility of octacosane and hexatriacontane in different n-alkane solvents, 1254-8 lsentropic compressibility behavior of dilute solutions of alcohols in water, 2237-45 Relation between the temperature of maximum density and partial compressibility of aqueous solutions Solid solution in the calcium oxide-manganese oxide system, 2285-94 Composition of the first coordination sphere of nickel(2 +) in concentrated aqueous nickel chloride and Electrolyte solutions in liquid ammonia.Part 6. Mean molal activity coefficients and cation transference Solvation spectra. Part 60. Specific solvation of iodide ions, 191 9-28 Hydrophobic interaction in aqueous organic mixed solvents, 636-45 Molecular reorientation and molecular association. Study by Raman depolarized light scattering, Far-ultraviolet solution spectroscopy of hydrosulfide ion, 1380-9 Effect of hexamethylphosphoric triamide on the living radical polymerization initiated by aged Solvation spectra. Part 60. Specific solvation of iodide ions, 191 9-28 lon-solvent interaction of some tetraalkylammonium ions in sulfolane from viscosity data, 2325-3 I Photolysis of thioketene S-oxides, 2624-7 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys.Part 1. Experiment- Solubility of octacosane and hexatriacontane in different n-alkane solvents, 1254-8 Infrared studies of the adsorption of probe molecules onto the surface of goethite, 872-82 Scanning studies on capillary condensation and evaporation of nitrogen. Part 1. Apparatus and Scanning studies on capillary condensation and evaporation of nitrogen. Part 2. Analysis of ascending State of water in cellulose acetate membranes, 803-1 5 Infrared study of surface hydroxyl groups on hematite, 1073-88 Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 1. Sorption of hydrocarbons and water in silanated and unsilanated partial hydrogen-forms of zeolite Y, Analysis of thermal effects in adsorption rate measurements, 2406-22 Infrared spectroscopic investigation of sorption of nitric oxide by b-metal phthalocyanines of the first Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron Sound velocity and surface tension from Flory’s statistical theory, 21 60-4 Selectivity in ethylene dimerization over supported nickel oxide catalysts, 15 13-20 Absorption and emission studies of solubilization in micelles.Part 5. Pyrene-3-sodium sulfonate Ultraviolet spectrum of periodate ion, 631-5 Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part I . Kinetics and al results, 561-77 amic parameters and theoretical models, 578-90 825-34 numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 radicals in aqueous solutions, 1142-9 of nonelectrolytes, 2246-5 1 nickel bromide solutions.Part 2. Application of halide nuclear magnetic relaxation, 27 12-34 numbers of silver(l), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 SOLVATION SOLVENT 1179-83 chromium ion + benzoyl peroxide, 182 1-9 SOLY al results, 561-77 SORBED SORPTION calculation method, 36-41 and descending scanning curves within B-type hysteresis loops, 42-55 Equilibriums and affinities, 1969-83 222 1-34 transition period, 2594-600 spectroscopy and gravimetry, 2839-50 SOUND SPECIFICITY SPECTRA solubilized in didodecyldimethylammonium bromide inverted micelles in benzene, 550-60 mechanism of polymerization, 774-86J.C.S. FARADAY I SUBJECT INDEX VOL.75 (1979) 63 SPECTRA(contd) Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-9 1 Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane. Effect of preadsorbed ammonia on isomerization of but-l-ene on sodium hydrogen-Y zeolite, 1 1 5 0 6 0 Molecular reorientation and molecular association. Study by Raman depolarized light scattering, ' Ultraviolet excitation of Raman spectra of pyridines adsorbed on oxides, 1215-21 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared Solvation spectra. Part 60. Specific solvation of iodide ions, 191 9-28 Charge-exchange reactions. Part 1. Charge+xchange mass spectra of the polyfluorobenzenes, 2040-7 Detection of formyl radicals in low pressure methane + oxygen flames, 2423-32 Ion condensation model and nuclear magnetic resonance studies of counterion binding in lyotropic liquid Evidence for an intermediate radical cation, 914-20 1 179-83 by use of tetrakis-n-allylmolybdenum, 1465-76 SPECTROMETER SPLITTING crvstals.663-8 STAB~LITY Molecular complexes between substituted indoles and tetracyanoethylene, 497-502 Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-9 1 STAB1 LIZATION Stability of metal uncharged ligand complexes in ion exchangers. Part 4. Hydration effects and stability Particle adhesion and removal in model systems. Part 2. Monodispersed chromium hydroxide on steel, Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Particle adhesion and removal in model systems.Part 2. Monodispersed chromium hydroxide on steel, Dissociation kinetics of picric acid and dipicrylamine in methanol. Steric effect on a proton transfer rate, Photolysis of thioketene S-oxides, 2624-7 Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-9 1 High resolution proton nuclear magnetic resonance studies of interaction between deoxyhemoglobin and Infrared studies of water in complexes, 762-73 Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(II1). molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds, 603-1 3 of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 changes of copper-ethylenediamine complexes in montmorillonite, 5 13-24 STAINLESS 20 1 4-26 STANIA 1487-94 STEEL 20 14-26 STERIC 213742 STILBITE STOICH IOM ETRY small molecules.Dithionite and diphosphoglycerate, 285 1-64 STRETCH SUBLIMATION SUBSTITUENT Association between polar molecules. Part 1. Nuclear magnetic resonance study of the dipole association Acid ionization of solvent, acetic acid, and &substituted phenols in tert-butanol + water mixtures, Association between polar molecules. Part 1. Nuclear magnetic resonance study of the dipole association Ligand substitution processes at five coordinate copper(I1) centers in hydrophilic and hydrophobic Proton exchange in aqueous solutions of glucose.Hydration of carbohydrates, 262-70 Structural analysis of molten potassium sulfate, 1 169-78 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Binding of n-alkyl sulfates to lysozyme in aqueous solution, 173- Precipitation of calcium surfactants. Part 2,211625 Precipitation of calcium surfactants. Part 3,212636 Kinetic investigation and numerical analysis of a micelle-catalyzed metal complex formation, 2 199-210 Micellar catalysis of metal complex formation. Part 2. Kinetics of the reaction between aquated 278 1-97 SUBSTITUTED of hexamethylphosphoramide with psubstituted nitrobenzenes, 79-85 environments, 1236-44 SUBSTITUTION SUGAR SULFATE solutions. A pulse radiolysis study, 1674-8764 J.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) SULFATE(contd) SU LFATION SULFIDE nickel(2 + ) and various neutral bidentate ligands in the presence of sodium dodecylsulfate micelles in aqueous solution, 2395-405 Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 650C, 1593-605 Dynamic properties of n-alkyl and s-alkyl intermediates in reactions of simple alkenes with hydrogen on Intermediates in catalytic hydrogenation of 1,3-butadiene, propadiene and methylacetylene on Kinetics and mechanism of the decomposition of hydrogen sulfide, methyl hydrogen sulfide and dimethyl Ion-solvent interaction of some tetraalkylammonium ions in sulfolane from viscosity data, 2325-3 1 Viscosity of gaseous carbon dioxide, sulfur hexafluoride and nitrous oxide at low temperatures, 892-7 Kinetics of reaction of calcium carbonate with sulfur dioxide and oxygen below 65OoC, 1593-605 Electron spin resonance studies of formation of sulfur dioxide anion radicals on calcium oxide, 161 3- I8 Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Isomerization of n-butenes on type A zeolites studied by infrared spectroscopy.Part 2. n-Butene isomerization on zeolites containing alkali and alkaline earth cations cocatalyzed by sulfur dioxide, 2744-52 SULFURIC molybdenum disulfide catalyst, 7-2 1 molybdenum disulfide catalyst, 1403-1 6 sulfide in a radio-frequency pulse discharge, 1868-75 SULFOLANE SULFUR Rotating disk dissolution rates of ionic solids. Part 3. Natural and synthetic ilmenite, 971-83 Effect of anion basicity in influencing the acidity function of concentrated aqueous solutions of oxyacids Mechanism of production and growth inhibition of collision crystallites, 92 1 4 Selectivity in ethylene dimerization over supported nickel oxide catalysts, 15 13-20 Evaluation of semiconducting tin oxide as an electrocatalyst support, 2 165-76 Attachment of spherical particles to surface of a pendant drop and tension of the wetting perimeter, 1-6 Scanning studies on capillary condensation and evaporation of nitrogen.Part 2. Analysis of ascending Infrared study of surface hydroxyl groups on goethite, 59 1-602 Rotating disk dissolution rates of ionic solids. Part 3. Natural and synthetic ilmenite, 971-83 Catalysis by highly diluted nickelxopper alloy foils. Hydrogen-deuterium equilibration and ethylene- Infrared study of surface hydroxyl groups on hematite, 1073-88 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Surface structure and surface states in magnesium oxide powders, 1769-79 Surface activity of methyl orange, 2077-82 Crystal growth kinetics of globular proteins. Lysozyme and insulin, 2753-6 1 Meniscuses formed by a cone at a free liquid surface.An absolute method of surface tension measureme- Attachment of spherical particles to surface of a pendant drop and tension of the wetting perimeter, 1-6 Metal-ion complexation reactions in the presence of surfactants. Part I .Mechanism of pH-dependent reaction between nickel(I1) and murexide in aqueous solution and application of the reaction to study of micellar phenomena, 1 19-3 1 Characterization of water-containing reversed micelles by viscosity and dynamic light scattering methods Electrokinetic study of surfactant adsorption, 669-78 Rate of exchange of surfactant monomer radicals and long chain alcohols between micelles and aqueous Binding of n-alkyl sulfates to lysozyme in aqueous solution, 173644 Precipitation of calcium surfactants. Part 3-21 26-36 Infrared studies of water in complexes, 762-73 Methanation of carbon dioxide and carbon monoxide on supported nickel-based composite catalysts, Synergistic effect of composite catalysts on the direct hydrogenation of carbon, 1495-506 Synergistic effect of composite catalysts on the direct hydrogenation of carbon, 1495-506 1606-12 SUPERSATD SUPPORT SURFACE and descending scanning curves within B-type hysteresis loops, 42-55 deuterium exchange reaction, 1001-10 by use of tetrakis-7c-allylmolybdenum, 146576 1487-94 nt, 2827-38 SU RFACTANT 1 32-9 solutions.A pulse radiolysis study, 1674-87 SYMMETRY SYNERGISM 787-802 SYNERGISTICJ.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 65 SYSTEM Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 1. Experiment- al results. 561-77 TANTALUM Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(III), molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds.603-1 3 TCNE TEMP Molecular complexes between substituted indoles and tetracyanoethylene, 497-502 Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Relation between the temperature of maximum density and partial compressibility of aqueous solutions Effect of pressure on the electrical conductivities of some fused nitrates, chlordtes, perchlorates, Meniscuses formed by a cone at a free liquid surface. An absolute method of surface tension measureme- Attachment of spherical particles to surface of a pendant drop and tension of the wetting perimeter, 1-6 Meniscuses formed by a cone at a free liquid surface. An absolute method of surface tension measureme- Reaction of tert-butyl radicals with hydrogen and with oxygen, 1458-64 Diffusion in viscous solvents.Part 3. Interdiffusion coefficients for planar and spherical solutes in 2- methylpentane-2,kIiol and their relationship to diffusion coefficients derived from luminescence measurements, 141 7-32 TETRABROMOETHANE ethane in hexamethyldisiloxane and hexamethyldisilazane, 1232-5 polymerization and formation of methyl(2-methyl4oxopentanoate), 2562-75 reactions. Volume changes on activation and reaction, 172-91 of nonelectrolytes, 2246-5 1 dichromates and thiocyanates, 2735-43 nt, 2827-38 TENSION nt, 2827-38 TERTIARY TETRABROMIDE Temperature dependence of the interdiffusion coefficient for very dilute solutions of 1,1,2,2-tetrabromo- Photolysis of methyl- and acetyl-manganese pentacarbonyls in methyl methacrylate. Initiation of Precipitation of calcium surfactants.Part 2 , 2 1 16-25 Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Addition of 2,2,3,3-tetramethylbutane to slowly reacting mixtures of hydrogen and oxygen, 1447-57 Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane. Cation exchangers with several groups of sites. Validity of theoretical base, 247-51 Mechanism of production and growth inhibition of collision crystallites, 92 1-4 General approach t o processing of kinetic adsorption data. But- l-ene + zeolite sodium-hydrogen-Y Effect of anion basicity in influencing the acidity function of concentrated aqueous solutions of oxyacids Theory of compartmentalized free-radical polymerization reactions. Part 3,2332-58 Thermodynamic properties for transfer of electrolytes from water to acetonitrile and to acetonitrile + Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Thermodynamic properties for transfer of cations from propylene carbonate t o dimethylsulfoxide and to Heat capacity changes in proton addition to the nitrogen of saturated organic molecules in water.Effects Molecular complexes between substituted indoles and tetracyanoethylene, 497-502 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 1. Experiment- Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 2. Thermodyn- Thermodynamics of aggregation of some tertiary n-alkylammonium picrates in benzene solutions, Thermodynamics of ion exchange of trivalent cobalt or chromium complex with cerium(II1) ions on TETRACHLORIDE TETRADECYL TETRAMETHYLBUTANE TETRAMETHYLBUTY L TETRAMETHY LPHENYLENEDIAMINE Evidence for an intermediate radical cation, 914-20 THEORY system, 101 6-22 160612 THERMODN water mixtures, 86-95 reactions.Volume changes on activation and reaction, 172-91 propylene carbonate + dimethylsulfoxide mixtures, 227-35 of solvation, 363-9 al results, 561-77 amic parameters and theoretical models, 578-90 825-34 cation exchange resin, 925-3466 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) THERMODN(contd) Enthalpy-entropy relations in the fluid kinetics of aqueous solvent systems, 942-52 Thermodynamic functions of hydration of saturated uncharged organic compounds.Free energies, Calorimetric enthalpy of reaction of cobalt(II1) with iron(I1) in aqueous perchloric acid, and the standard Statistical molecular calculation of thermodynamic parameters of adsorption of aromatic hydrocarbons Statistical molecular calculation of thermodynamic parameters of adsorption of aromatic hydrocarbons Electron spin resonance studies of the radicals formed from carbon-nitroso compounds and olefins. Part Effects of molecular flexibility and shape on the excess enthalpies and heat capacities of alkane systems, Effects of molecular shape on the excess enthalpies and heat capacities of cycloalkane systems, 1708-1 4 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl Thermodynamic properties and local structures of nonstoichiometry.Galvanic cell study of fused indium Thermoelectric effects and dielectric polarization in biopolymers, 323-34 Scaled-particle theory and salting constant of tetrahydrofuran in various aqueous electrolyte solutions at Ion condensation model and nuclear magnetic resonance studies of counterion binding in lyotropic liquid Photolysis of thioketene S-oxides, 2624-7 Photolysis of thioketene S-oxides, 2624-7 Ring-disk electrodes. Part 19. Adsorption studies at low frequency a.c., 1623-34 Reaction probabilities and threshold energy in the abstraction reaction between hydrogen atoms and Electrical behavior of powdered tin-antimony mixed oxide catalysts, 1346-55 Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Evaluation of semiconducting tin oxide as an electrocatalyst support, 2 165-76 Acidic and basic properties of mixed tin-antimony oxides, 2762-7 Study of the oxidative reactions of butenes over mixed tin-antimony oxides, 2768-80 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces.Part Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Photoelectrochemical reactions of pigmentary titanium dioxide with alcohols and aliphatic amines, Reactions involving electron transfer at semiconducting surfaces. Part 8. Room temperature activity of prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys.Part 1. Experiment- al results, 561-77 Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 2. Thermodyn- amic parameters and theoretical models, 578-90 Photoadsorption and photodesorption of oxygen on highly hydroxylated titanium dioxide surfaces. Part 1. Role of hydroxyl groups in photoadsorption, 736-47 ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation of alkylpyridines, 1356-70 Reactions involving electron transfer at semiconductor surfaces. Part 9. Oxygen labelling study of photooxidation of C3 and C, alcohols over rutile, 2000-13 Photoelectrochemical reactions of pigmentary titanium dioxide, 2507-16 Photoelectrochemical and photocatalytic reactions of pigmentary titanium dioxide, 25 17-25 Photoelectrochemical reactions of pigmentary titanium dioxide with alcohols and aliphatic amines, Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 enthalpies and entropies at 25"C, 1184-95 electrode potential for cobalt(III)-(II), 1268-9 on graphite.Part 1. Condensed aromatic hydrocarbons, 1281-7 on graphite. Part 2. Polymethyl and monoalkyl benzenes, 1288-300 1. Nitrosodurene and fluoroolefins, 1521-30 1700-7 sulfoxide + water mixtures, 1780-7 iodide mixtures, 2303-1 1 THERMOELEC THF 298.15 K, 1837-48 crystals, 663-8 THICKNESS THIOCARBONYLCYCLOHEXANE THIOKETENE THIONINE THRESHOLD neopentane, 1619-22 TIN 1487-94 TITAN1 A 2.Study of radical intermediates by electron paramagnetic resonance, 748-61 2526-34 TITANIUM 305- 13 2526-34 TOLUENEJ.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 67 TOPOLOGY Correlation between topological features and molar volumes of n-alkanes and excess volumes of their binary mixtures, 1 132-41 Thermodynamic properties for transfer of electrolytes from water to acetonitrile and to acetonitrile + water mixtures, 86-95 Quantum theory of kinetic isotope effects in proton transfer reactions, 205-26 Thermodynamic properties for transfer of cations from propylene carbonate to dimethylsulfoxide and to propylene carbonate + dimethylsulfoxide mixtures, 227-35 Reactions involving electron transfer at semiconducting surfaces.Part 8. Room temperature activity of prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, 305- 13 Magnetooptical rotation studies of liquid mixtures. Part 3. Specific interactions in mixtures of carbon tetrachloride with aliphatic, alicyclic and aromatic hydrocarbons and amines, 370-3 Molecular complexes between substituted indoles and tetracyanoethylene, 497-502 Studies in ion solvation in nonaqueous solvents and their aqueous mixtures. Part 19. Free energies of transfer of alkali metal chlorides from water to DMSO + water mixtures up to 40% w/w DMSO, of sodium chloride to 60% w/w DMSO and of potassium bromide and iodide to 10% w/w DMSO. Comparison with silver chloride, 907-1 3 Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane.Evidence for an intermediate radical cation, 91 4-20 Heterogeneous catalysis in solution. Part 17. Kinetics of oxidation-reduction reactions catalyzed by electron transfer through the solid. An electrochemical treatment, 1507-1 2 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl sulfoxide f' water mixtures, 1780-7 Electron-transfer reactions of nitroxyl radicals with one-electron reduced quinones and viologens, Mechanism of interactions between aquametallic complexes and the photoexcited aquouranyl(V1) ion, Electron transfer reactions involving porphyrins and carotenoids, 2869-72 Electrolyte solutions in liquid ammonia. Part 6. Mean molal activity coefficients and cation transference Kinetics of polymorphic transitions in tetrahedral structures. Part 1.Experimental methods and the Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 1. Adsorption properties of zeolitic ruthenium and of chromium, iron and lanthanum mordenites. Part 2. Hydrogen oxidation catalyzed by X zeolite containing transition metal ions, 3 14-22 Hydrogenation of alkenes over supported gold, 385-94 Methanation of carbon dioxide and carbon monoxide on supported nickel-based composite catalysts, Infrared spectra of carbon monoxide adsorbed on transition metal cation exchanged tin(1V) oxide, Surface characterization of nickel boride and nickel phosphide catalysts by x-ray photoelectron Mechanism of interactions between aquametallic complexes and the photoexcited aquouranyl(V1) ion, Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Thermodynamics of aggregation of some tertiary n-alkylammonium picrates in benzene solutions, Kinetics of gas phase addition reactions of trichlorosilyl radicals.Part 2. Additions to I-olefins, 2617-23 Analysis of thermal effects in adsorption rate measurements, 2406-22 Kinetics of the gas-phase reaction between iodine and trimethylsilane and the bond dissociation energy Use of internal and external references in nuclear magnetic resonance determinations of association Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane. TRANSFER 1 9 1 2- 1 8 2252-72 TRANSFERENCE numbers of silver(I), copper(I1) and lead(I1) nitrates in liquid ammonia at 233 K, 1023-33 transition y -+ B lithium zinc silicate(Li$ZnSiOd), 374-84 Equilibriums and affinities, 1969-83 Heats and entropies, 198499 TRANSITION TRANSITION METAL 787-802 1487-94 spectroscopy, 2027-39 2252-72 TRANSPORT TRIALKY LAMMONIUM 825-34 TRICHLOROSILY L TRIMETHYLPENTANE TRI METHY LSI LANE of the hydrogen-silicon bond.Part 2,1126-3 1 constants of weak molecular complexes, 1222-5 Evidence for an intermediate radical cation, 914-20 TRINITROBENZENE TRIPHENYLMETHYL68 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) TRIPLE Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(III), molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds, 603-1 3 Evidence for an intermediate radical cation, 914-20 TRITY L Formation of the trityl cation in pulse irradiated solutions of triphenylmethyl chloride in cyclohexane.Mixed ligand complexes in solution of manganous ion. Nuclear relaxation rate analysis, 2433-8 Single-pulse shock tube studies of hydrocarbon pyrolysis. Part 7. Pyrolysis of propane, 28 19-26 Enthalpies of formation of homoleptic dimethylamido compounds of tantalum(V), molybdenum(III), molybdenum(IV), tungsten(II1) and tungsten(V1). Enthalpy contributions of metal-metal triple bonds, 603-1 3 Ultraviolet photoelectron spectroscopic study of chemisorption of halogens on tungsten( 1 OO), 850-62 Ultraviolet photoelectron spectroscopic study of the chemisorption and decomposition of formic acid on tungsten( IOO), 863-7 1 Anomalous adsorption kinetics.y-Nitrogen on the ( 1 10) plane of tungsten, 2100-15 TWIN CALORIMETER Twin calorimeter for the determination of enthalpies of vaporization of small samples from 300 to 420 K Ultrasonic behavior of the system water + propylene glycol, 2067-72 Ultraviolet photoelectron spectroscopic study of chemisorption of halogens on tungsten( loo), 850-62 Ultraviolet photoelectron spectroscopic study of the chemisorption and decomposition of formic acid on Search for strongly adsorbed carbon monoxide at copper single crystal surfaces using ultraviolet Study of the interaction of nitric oxide with copper( 100) and copper( 1 1 1) surfaces using low energy Inorganic photophysics in solution.Part 2. Temperature activation of decay processes in the luminescen- Mechanism of interactions between aquametallic complexes and the photoexcited aquouranyl(V1) ion, Chemical quenching by water of the photoexcited uranyl ion in aqueous acidic solution, 2273-84 Pulse radiolysis study of chlorpromazine and promazine free radicals in aqueous solution, 22-35 Molecular complexes between substituted indoles and tetracyanoethylene, 497-502 Absorption and emission studies of solubilization in micelles. Part 5. Pyrene-3-sodium sulfonate Ultraviolet spectrum of periodate ion, 63 1-5 Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1 . Kinetics and Pulse radiolysis study of acrylonitrile in aqueous solution, 1050-66 Ultraviolet excitation of Raman spectra of pyridines adsorbed on oxides, 121 5-21 Far-ultraviolet solution spectroscopy of hydrosulfide ion, 1380-9 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared Surface structure and surface states in magnesium oxide powders, 1769-79 Some free radical reactions of camphor in relation to the action of cytochrome P450, 1849-56 Solvation spectra.Part 60. Specific solvation of iodide ions, 1919-28 Optical spectroscopy of hydrated, dehydrated, and ammoniated cobalt( 11) exchanged zeolites X and Y, ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation Isotopic exchange of molecular oxygen with oxygen of vanadium oxide (V205) modified with oxides of ESCA investigation of vanadium pentoxide + titanium dioxide catalysts for the vapor phase oxidation Surface structure and surface states in magnesium oxide powders, 1769-79 Kinetics of reaction of magnesium vapor with silica, 2807-18 Ultrasonic behavior of the system water + propylene glycol.2067-72 TRYPTOPHAN TUBE TUNGSTEN 2359-65 ULTRASOUND UPS tungsten( IOO), 863-7 1 photoelectron spectroscopy, 984-6 electron diffraction and electron spectroscopy, 2143-59 ce of uranyl[U02* f ] ion, 342-52 URANYL 22 52-72 uv solubilized in didodecyldimethylammonium bromide inverted micelles in benzene, 550-60 mechanism of polymerization, 774-86 by use of tetrakis-z-allylmolybdenum, 1465-76 260 1 - 1 6 VANADATE of alkylpyridines, 1356-70 alkaline earth metals, 691-4 of alkylpyridines, 1356-70 VAN AD1 UM VAPOR VELOCITYJ.C.S.FARADAY I SUBJECT INDEX VOL. 75 (1979) 69 VELOCITY(contd) VIBRATION VI NY LCA RBAZOLE Sound velocity and surface tension from Flory's statistical theory, 2 160-4 Infrared studies of water in complexes, 752-73 Photoinitiated polymerization of N-vinylcarbazole in presence of benzoyl peroxide. Part 1. Kinetics and mechanism of Dolvmerization. 774-86 VINYLFLUORID'E Free radical addition to olefins. Part 25. Addition of uerfluoro-tert-butyl radicals to fluoro olefins, 1040-9 VIOLOGEN Electron-transfer reactions of nitroxyl radicals with one-electron reduced quinones and viologens, Second virial coefficients of nitrogen at very low temperatures, 479-80 Characterization of water-containing reversed micelles by viscosity and dynamic light scattering methods Viscosity of gaseous carbon dioxide, sulfur hexafluoride and nitrous oxide at low temperatures, 892-7 Effect of electrolytes on solution behavior of water soluble macromolecules, 993-1 000 Diffusion in viscous solvents.Part 3. Interdiffusion coefficients for planar and spherical solutes in 2- methylpentane-2.4-diol and their relationship to diffusion coefficients derived from luminescence measurements, 1417-32 Quaternization of poly(4-vinylpyridine). Kinetic and viscometric measurements, 1728-35 Viscosities of trichlorofluoromethane, chlorotrifluoromethane, dichlorofluoromethane, chlorodifluoro- Application of the environmental relaxation model to the ideal and nonideal glass-forming molten Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Dependence of the viscosity of liquid 3,3-diethylpentane, mercury, argon and methane on thermal lon-solvent interaction of some tetraalkylammonium ions in sulfolane from viscosity data, 2325-3 I Scanning studies on capillary condensation and evaporation of nitrogen.Part 2. Analysis of ascending Effect of pressure and temperature on some kinetic and thermodynamic parameters of nonionic Heat capacity changes in proton addition to the nitrogen of saturated organic molecules in water. Effects Absorption of hydrogen by titanium-cobalt and titanium-nickel intermetallic alloys. Part 2. Thermodyn- Excess volumes of mixing of aniline + aromatic hydrocarbons, I 120-5 Correlation between topologcal features and molar volumes of n-alkanes and excess volumes of their Self-diffusion and shear viscosity in dense molecular liquids, 2060-6 Effect of pressure on the electrical conductivities of some fused nitrates, chlorates, perchlorates, Wall-less reactor studies.Part I. Ethane pyrolysis, 1089-100 Wall-less reactor studies. Part 2. Propane pyrolysis, 110!-10 1912-18 VIRIAL VISCOSITY 132-9 methane, and trifluoromethane from 373 to 570 K, 1752-6 mixtures, 1830-6 pressure over a wide PVT range, 2295-302 VOL and descending scanning curves within B-type hysteresis loops, 42-55 reactions. Volume changes on activation and reaction, 172-91 of solvation, 363-9 amic parameters and theoretical models, 578-90 binary mixtures, 1 132-41 dichromates and thiocyanates, 273543 WALL LESS Wall-less reactor studies.Part 3. n-Butane pyrolysis, 2688-93 WATER Thermodynamic properties for transfer of electrolytes from water to acetonitrile and to acetonitrile + Proton exchange in aqueous solutions of glucose. Hydration of carbohydrates, 262-70 Surface acidity of q-alumina. Part 2. Interaction of pyridine with other adsorbates, 289-304 Ion reactivity in reversed-micellar systems. Kinetics of reaction between micelles containing hydrated nickel(I1) and murexide-containing micelles in the system aerosol-OT + water + heptane, 48 1-96 Infrared studies of water in complexes, 762-73 State of water in cellulose acetate membranes, 803-15 Studies in ion solvation in nonaqueous solvents and their aqueous mixtures. Part 19. Free energies of transfer of alkali metal chlorides from water to DMSO + water mixtures up to 40% wjw DMSO, of sodium chloride to 60% w/w DMSO and of potassium bromide and iodide to 10% w/w DMSO.Comparison with silver chloride, 907-1 3 water mixtures, 86-95 Enthalpy-entropy relations in the fluid kinetics of aqueous solvent systems, 942-52 Adsorption of water vapor on evaporated germanium films. An infrared study, 962-70 Infrared study of surface hydroxyl groups on hematite, 1073-88 Partial molar enthalpies in the sodium chloride + calcium chloride + water system. Application of the McKay-Perring method to thermal measurements, 1745-5 I70 J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) WATER(contd) Surface structure and surface states in magnesium oxide powders. 1769-79 Thermodynamic properties for transfer of electrolytes from water to dimethyl sulfoxide and to dimethyl Ultrasonic behavior of the system water + propylene glycol.2067-72 Sorption of hydrocarbons and water in silanated and unsilanated partial hydrogen-forms of zeolite Y, lsentropic compressibility behavior of dilute solutions of alcohols in water. 2237-45 Chemical quenching by water of the photoexcited uranyl ion in aqueous acidic solution, 2273-84 Influence of lithium, sodium, and potassium cation concentrations in X and Y zeolites on isotherms and Meniscuses formed by a cone at a free liquid surface. An absolute method of surface tension measureme- Rotating disk dissolution rates of ionic solids. Part 3. Natural and synthetic ilmenite, 971-83 Ultraviolet photoelectron spectroscopic study of chemisorption of halogens on tungsten( 100).850-62 Ultraviolet photoelectron spectroscopic study of the chemisorption and decomposition of formic acid on Kinetics of polymorphic transitions in tetrahedral structures. Part 1. Experimental methods and the Effects of dehydration and x-ray irradiation on the structure of manganese sodium zeolite. 898-906 Changes in mordenite upon various pretreatments. Part 1. Structural rearrangements. 1245-53 Spectroscopic study on the surface structure and environment of fixed molybdenum catalysts prepared Surface characterization of nickel boride and nickel phosphide catalysts by x-ray photoelectron Study of the interaction of nitric oxide with copper( 100) and copper( 1 1 1 ) surfaces using low energy Adsorption of nitric oxide and pyridine by various metal phthalocyanines studied by x-ray photoelectron Mordenite catalysts: influence of metal on disproportionation, dealkylation and isomerization, 434-9 Some unusual properties of activated and reduced silver-sodium A zeolites, 56-64 Magnetic determination of metallic nickel particles dispersed on X and Y zeolite structures, 165-71 Cation exchangers with several groups of sites.Validity of theoretical base, 247-5 1 Hydrogen oxidation catalyzed by X zeolite containing transition metal ions, 3 14-22 Influence of activation conditions of cerium-)< zeolite on its oxidizing properties as shown by infrared Electron spin resonance of copper-exchanged Y zeolites. Part 1. Behavior of the cation during Electron spin resonance studies of copper-exchanged Y zeolites. Part 2. Interactions with cerium( 111) Site group interaction effects in zeolite Y. Part 3. Potassium-sodium equilibrium, 440-5 Preparation and properties of hydrogen form of stilbite, heulandite and clinoptilolite zeolites, 883-91 Effects of dehydration and x-ray irradiation on the structure of manganese sodium zeolite, 898-906 General approach to processing of kinetic adsorption data. But- I -ene + zeolite sodium-hydrogen-Y Effect of preadsorbed ammonia on isomerization of but-l-ene on sodium hydrogen-Y zeolite, 1150-60 Redox behavior of transition metal ions in zeolites. Part 7. Characterization of a nickel metal phase in Oxidation-reduction behavior of highly dispersed ruthenium in ruthenium sodium Y zeolite, 1395-402 Temperature programmed reduction of copper ions in zeolites. 1688-99 Adsorption properties of zeolitic ruthenium and of chromium. iron and lanthanum mordenites. Part I . Adsorption properties of zeolitic ruthenium and of chromium. iron and lanthanum mordenites. Part 2. Alternative interpretation of infrared spectra of the zeolite sodium hydrogen Y + but-l-ene system, Electric conductivity of nitrate inclusion complexes of A and X zeolites, 2083-8 Sorption of hydrocarbons and water in silanated and unsilanated partial hydrogen-forms of zeolite Y. Changes in mordenite upon various treatments. Part 2. Hydroxyl groups. 236676 Analysis of thermal effects in adsorption rate measurements, 2406-22 Diffusion of chemisorbed hydrogen in a platinum zeolite, 2473-80 Study of the acidity of ZSM-5 zeolite by microcalorimetry and infrared spectroscopy, 2544-55 sulfoxide + water mixtures, 1780-7 222 1-34 heats of adsorption of propane and water, 2662-77 nt, 2827-38 WEATHERING WORK tungsten( 100), 86%7 I transition y + p lithium zinc silicate(LizZnSi04). 374-84 X RAY XPS by use of tetrakis-n-allylmolybdenum. 1465-76 spectroscopy, 2027-39 electron diffraction and electron spectroscopy, 2 143-59 spectroscopy and gravimetry, 2839-50 XYLENE ZEOLITE and electron spin resonance spectroscopies, 335-41 dehydration, 406-22 ions inside the zeolite framework, 423-33 system, 1016-22 zeolite Y, 1196-206 Equilibriums and affinities, 1969-83 Heats and entropies, 1984-99 2073-4 222 1-34J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 71 ZEOLITE(contd) Optical spectroscopy of hydrated, dehydrated, and ammoniated cobalt(I1) exchanged zeolites X and Y, Influence of lithium, sodium, and potassium cation concentrations in X and Y zeolites on isotherms and heats of adsorption of propane and water, 2662-77 Isomerization of n-butenes on type A zeolites studied by infrared spectroscopy. Part 2. n-Butene isomerization on zeolites containing alkali and alkaline earth cations cocatalyzed by sulfur dioxide, 2744-52 2601-16 ZETA ZINC Electrokinetic study of surfactant adsorption, 669-78 Reactions involving electron transfer at semiconducting surfaces. Part 8. Room temperature activity of prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, Kinetics of polymorphic transitions in tetrahedral structures. Part 1. Experimental methods and the transition y -+ Binary gaseous diffusion coefficients of zinc in hydrogen, helium and argon at 1 atmosphere and 72& 1120 K, 5459 Composition dependence of fluidities and conductances of mixtures of hydrated melts, 13 12-1 9 Distribution of nickel ions among octahedral and tetrahedral sites in nickel zinc aluminate spinel solid solutions, 1876-87 305-1 3 .lithium zinc silicate(Li2ZnSi04), 374-84J.C.S. FARADAY I SUBJECT INDEX VOL. 75 (1979) 71 ZEOLITE(contd) Optical spectroscopy of hydrated, dehydrated, and ammoniated cobalt(I1) exchanged zeolites X and Y, Influence of lithium, sodium, and potassium cation concentrations in X and Y zeolites on isotherms and heats of adsorption of propane and water, 2662-77 Isomerization of n-butenes on type A zeolites studied by infrared spectroscopy. Part 2. n-Butene isomerization on zeolites containing alkali and alkaline earth cations cocatalyzed by sulfur dioxide, 2744-52 2601-16 ZETA ZINC Electrokinetic study of surfactant adsorption, 669-78 Reactions involving electron transfer at semiconducting surfaces. Part 8. Room temperature activity of prereduced zinc oxide, titanium dioxide and magnesium oxide surfaces for oxygen isotope exchange, Kinetics of polymorphic transitions in tetrahedral structures. Part 1. Experimental methods and the transition y -+ Binary gaseous diffusion coefficients of zinc in hydrogen, helium and argon at 1 atmosphere and 72& 1120 K, 5459 Composition dependence of fluidities and conductances of mixtures of hydrated melts, 13 12-1 9 Distribution of nickel ions among octahedral and tetrahedral sites in nickel zinc aluminate spinel solid solutions, 1876-87 305-1 3 .lithium zinc silicate(Li2ZnSi04), 374-84

ISSN:0300-9599    

DOI:10.1039/F197975BA001

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Dynamic Properties of n-Alkyl and s-Alkyl Intermediates in Reactions of Simple Alkenes with Hydrogen on MoS2 Catalyst BY TOSHIO OKUHARA AND KEN-ICHI TANAKA" Research Institute for Catalysis, Hokkaido University, Sapporo, Japan Received 16th December, 1977 \ / / \ The hydrogen exchange between (Z)-[l-2Hl]propene, C=C , and [2H6]propene taking H H place on MoS2 in the presence of hydrogen yielded a variety of deuteropropene isomers in the relative proportions 7.5 % [l,l-2Hz], 4.5 % (Z)-[1,2-2H2], 34.0 % (E)-[l,2-2H2], 51.0 % (E)-[1-2H1], 0 % [2-2H1] and 3 % [2Ho]propene in the initial stage of the exchange reaction at room temperature. This distribution was in good agreement with calculations based on a relative contribution of 70 % n-propyl and 30 % isopropyl intermediates determined by a separate experiment, and the relative contributions of the n-alkyl and s-alkyl intermediates were found to be the same for the hydrogen exchange reaction of but-1-ene.The hydrogen addition and elimination processes in the isomeriza- tion reaction as well as in the hydrogen exchange of olefins are strictly controlled by a cis-stereo R \ / specificity of the site which has one coordination vacancy, Mo(S)4. The s-alkyl species formed on this site are restricted from free rotation around the coordination bond by steric interaction with the surrounding sulphur ions. Such dynamic properties of the s-alkyl intermediates give unexpectedly low values for [3-2Hl]propene : [l-2Hl]propene of c 1/15, and also cause two distinguishable confor- mations for the s-butyls, one from but-l-ene and the other from but-2-ene.depending on the orienta- R \ tion of the vacant coordination site. The alkyl species in the form of Mo(S)4 are intermediates for the isomerization and/or the hydrogen exchange but do not undergo consecutive hydrogenation, because dissociation of H2 requires two vacancies. / Reactions catalysed on solid surfaces may be classified phenomenologically into " structure sensitive and structure insensitive " reacti0ns.l Furthermore, some catalytic reactions such as ammonia synthesis (N2 + 3H2 3 2NH3) and methanation (CO + 3H2 + CH4 + H20) occur readily on heterogeneous catalysts but hardly proceed in a homogeneous catalyst system. This strongly suggests that structural requirements must be fulfilled and/or particular arrangements of sites must exist in order to promote smooth chemical reactions.Such structural prerequisites may depend on the type of chemical reaction. However, at present, in a few cases mechanistic bases have been established by which the structural requirement of the sites in heterogeneous catalysis can be deduced. we demonstrated that the alkyl intermediates in the hydrogen exchange of alkenes and/or their isomerization reactions on molybdenum disulphide and on sulphided nickel catalysts cannot be hydrogenated to alkanes, 7 In our previous papers,2*8 ALKENE REACTIONS WITH H ON MOS, CATALYST and the hydrogenation reaction on these catalysts occurs on different active sites. These discrete properties of the active sites for catalysis were explained by the different degrees of coordinative unsaturation of the sites,,.4* i.e., the active sites for hydrogenation require three degrees of coordinative vacancies such as are possessed by Wilkinson’s hydrogenation catalyst which exhibits activity through loss of one phosphine ligand.6 In contrast, while the alkyl species formed on sites with two degrees of coordinative unsaturation undergo isomerization and/or hydrogen exchange, they are not hydrogenated to alkanes, as shown in the following reaction scheme ; A3MH2- site) 1 (3M- site 1 SCHEME 1 .-Hydrogenation, S, & ./ Hz s’ I ‘s %M’H + ;c=‘c< - +h - s‘ I ‘s S (*MH-s i te 1 S SCHEME 2.-Isomerization and hydrogen exchange. where ,M and 3M denote the sites which originally have two or three degrees of coordinative vacancies, and which correspond to the B- and C-sites in our previous papers.If these sites are bonded to hydrogen atoms, they may be expressed as ,MH, 3MH and 3MHz, depending on the number of hydrogen atoms. The MoS, powder used here has a layer structure comprising a trigonal prismatic unit cell, the molybdenum ion being surrounded by six sulphur ions. However, the molybdenum ions exposed on the edge surface of a sandwich-like layer structure of MoS, are surrounded by fewer sulphur ions, and a hard sphere model suggests a probable structure where molybdenum ions are surrounded by four sulphur ions. Actually, the authors have succeeded in demonstrating that only the edge surface of MoS, single crystal wafers catalyses the isomerization and hydrogen exchange in olefins in the presence of hydr~gen.~.This has clarified the place and/or the sites on which the reactions occur, and our interest now focuses on the dynamic properties of the intermediates formed on these sites during catalysis correlated with the structure of the active sites. EXPERIMENTAL Reactions were carried out at room temperature in a conventional circulation system with a total volume of about 300 cm3, and the analysis of the olefins was performed with an on-line gas chromatograph. The material was obtained and purified as follows: [2H8]but-l-ene ; by deuteration of perdeuterobutadiene on a ZnO catalyst at room tempera- ture. [2H6]propene; by hydrogen exchange of propene with D2 on a MgO catalyst atT. OKUHARA AND K . - I . TANAKA 9 room temperature.(Z)-[l-2Hl]pr~pene ; [l-2Hllmethyl acetylene, obtained by shaking methyl acetylene in an alkaline deuterium oxide solution, was hydrogenated on a Pdlcarbon catalyst. All deutero-olefins obtained by the above methods were purified by gas chromato- graphic separation. The surface area of the MoS2 powder was 15 m2 g-l by the B.E.T. method with a nitrogen adsorbent, and X-ray diffraction indicated a hexagonal layer structure (2H-structure) composed of a trigonal prismatic unit cell. The impurities in the MoS2 analysed by atomic absorption analysis were ; Fe 0.02, Mg 0.0015, Ca 0.0077, Na 0.012, Mn 0.0003, Cr < 0.0001 and K < 0.1 %. The MoS2 powder was activated by evacuation at >400°C for such reactions as the hydrogenation, the equilibration of H2 and D2, the isomerization and/or the hydrogen exchange in olefins and the metathesis of olefins.8 A mass spectroscopic analysis was made using parent peaks obtained from 12 eV of the ionization voltage. A microwave spectroscopic analysis of the deutero-olefins was carried out at dry ice temperature using a spectrometer with 110 kHz sinusoidal Stark modula- t i ~ n , ~ ~ lo and multiple reflectance of microwaves in the absorption cell was prevented by setting a pair of ferrite isolators in the cell.RESULTS n-BUTYL AND S-BUTYL INTERMEDIATES In the absence of hydrogen, there was no appreciable double bond migration or cis-trans isomerization on MoS, catalyst; however, when hydrogen was added, these reactions accelerated greatly, as shown in fig. 1. This remarkable promoting effect of hydrogen seems to indicate the participation of an associative mechanism through butyl intermediates.The isomerization of but-1-ene in the presence of D2, however, gave > 90 % nondeuterated but-2-ene, and in contrast, the hydrogenation reaction taking place simultaneously predominantly yielded [1,2-2H2]butane.3 TABLE 1 .-GEOMETRICAL ISOMERS OF [2HJBUT-I-ENE AND C~S-[~H~]BUT-~-ENE FORMED DURING REACTION OF BUT-~-ENE WITH D2 (1 : 1) AT ROOM TEMPERATURE [ZH llbut- 1-ene ck+H ~Ibut-Z-ene runs (Z)-l-[ZH1] (E)-1-[2H1] 2-[2H1] 3-[2H1] 4-[2H1] 1-[2H1] 2-[2H11 - - run 1 a 29.6 % 34.3 36.1 0 0 run 2 b 28.0 % 31.0 35.0 6.0 0 88.0 % 12.0 a Conversions in hydrogenation and isomerization were 9.5 and 10.6 %, respectively and 12Hl]but-l-ene was NN 9.0 % of the total but-1-ene.Conversions in hydrogenation and isomerization were 42.0 and 34.0 %, respectively and the concentration of [2Hl]but-l-ene in the total but-1-ene was 32.3 % and of ~is-[~H~]but-2-ene was 29.8 % in the total cis-but-2-ene. To elucidate this phenomena, [2H,]but-l-ene and ci~-[~H,]but-2-ene which were formed as the minor products in the isomerization of but-1-ene in the presence of D2 were analysed by microwave spectroscopy. As shown in table 1, a deuterium atom H H \ / in [2H,]but-l-ene is equally distributed throughout the vinyl group, C=C , / \ H and a deuterium atom in ci~-[~H,]but-2-ene is located mainly on the outer carbon atoms. As will be discussed later, the formation of [2H,]but-l-ene by hydrogen exchange between but-1-ene and D2 is inevitably slow compared with the reshuffling of the vinyl hydrogens, which results in the random distribution of a deuterium atom in the vinyl group.10 ALKENE REACTIONS WITH H ON MoS, CATALYST TABLE 2.-HYDROGEN EXCHANGE BETWEEN [2Ho]BUT-I -ENE AND ['HslBUT-l-ENE IN THE PRESENCE OF AN Hz+D2 MIXTURE (1 : 1) AND IN THE ABSENCE OF HYDROGEN AT ROOM MENTAL RUNS.TEMPERATURE. RESULTS AT 3min AND 5min WERE OBTAINED FROM DIFFERENT EXPERI- (1) deuterium distribution of but-1-ene/ % time/ min 0 3 5 (120) 2Ho 2Hi 2Hz 2H3 2H4 2H5 2H6 2H7 2Hs H 2 HD D2 50.4 0.6 49.0 48.1 0.5 51.4 (50.4) 24.2 j 19.2 / 5.5 0.8 0.1 0.1 6.4 19.7 23.1 48.1 3.2 48.7 9.2 ! 19.5 15.4 4.9 0.6 5.0 16.3 20.0 9.1 41.9 19.2 39.0 (14.1) k21.8) _-----_---- i (10.7) (2.3) (-) (2.3) (11.9) (23.0) (14.0) - - - (0.5) (49.1) - (2) distribution of geometrical isomers of [2H1]but-l-ene/ % time/ [ZHllbut-1-ene isomers rnin (Z)-l-ZHl (E)-l-2H1 2-2H1 3-2H1 4-2H1 3 15.9 16.0 68.1 0 0 5 24.9 29.1 46.1 0 0 (120) (15.7) (18.5) (65.8) (0) (0) Results at 3 and 5 min were the separate experimental runs in the presence of hydrogen (H2/D2 = l), and the values in parentheses are the results in the absence of hydrogen.Simultaneous hydrogenation and double bond migration gave butane 1.8 % (4.9), cis-but-2-ene 2.7 % (5.6) and t-but-2-ene 1.5 % (3.7) at 3 min ( 5 rnin). In order to determine which is the primarily exchanged hydrogen in but-1-ene, intermolecular hydrogen exchange between [,H,]but-l-ene and [2H8]bUt-l-ene, and coisomerization of ~is-[~H,]but-2-ene and ~is-[~H,]but-2-ene were performed with the MoS, catalyst at room temperature.When a mixture of [2H,]but-l-ene (2 mmHg) and [,H,]but-l-ene (2 mmHg) was admitted with 3 mmHg of an H2 +D, mixture onto the MoS, catalyst, intermolecular hydrogen exchange occurred rapidly as shown in table 2. [2Hl]but-l-ene appearing in the initial stage of this exchange reaction should hold the deuterium atom at its originally exchanged position if its desorption is sufficiently rapid. The results indicate that the vinyl hydrogens in H H c=c \1 2/ / \ but-1-ene, , are particularly exchangeable on MoS, in the presence of H hydrogen, and that the relative ratio of hydrogen exchange on C-1 and C-2 is about 3/7 calculated from the ratio of [l-2Hl]but-l-ene/[2-2€€l]but-l-ene. This may indicate that the relative importance of the s-butyl and n-butyl intermediates in the inter- molecular hydrogen exchange is 3 to 7.Note that the geometrical distribution of [2Hl]but-l-ene appearing in the absence of hydrogen (run 120min in table 2) is similar to that in the presence of hydrogen. This suggests that the intermolecular hydrogen exchange reaction of but-1-ene taking place slowly in the absence of hydrogen might proceed uia the butyl intermediates over the small fraction of sites which contain a hydrogen atom which can produce the intermediates. As noted in our previous paper,l the coisomerization of ~is-[~H,]but-2-ene and ci~-[~H,]but-2-ene gave a value of 0.5 as the average number of exchanged hydrogens per isomerized molecule, trans-but-2-ene. This proves that a s-butyl intermediate from cis-but-Zene formed by picking up a hydrogen atom, H or D, changes to trans-T.OKUHARA AND K . - I . TANAKA 11 but-2-ene, if a hydrogen atom other than that picked up returns to the active site. That is, the cis-trans isomerization reaction on MoS, catalyst necessarily occurs with the hydrogen exchange. The presence of a s-butyl intermediate in the cis-trans isomerization is also supported by microwave spectroscopic analysis of cis-['H I]- but-2-ene which is nearly 100 % of the ci,~-[2-~H~]but-2-ene formed in the coisomeriza- tion. 10 x" 1 ri 2 0 g 5 a (4 0 50 I00 150 200 250 timelmin 0 50 100 I50 time/min FIG. 1.-Hydrogen promoting effects on isomerization reactions at room temperature. (a) But-1-ene : 0, 1-C4Hs; 0, C4HI0; A, t-2-C4Hs; 0, cis-2-C4H8. (b) cis-but-2ene: 0, 1-C4Hs; 0, C4 Hlo; A, C ~ ~ - Z C ~ H ~ ; C ) , ~ - Z C ~ H ~ .As described in scheme 2, the s-butyl species formed on the 2M-sites do not easily undergo isotopic exchange with D2, because the D2 molecule is hard to dissociate either on 2MH or s-butyl-2M which has only one coordinative vacancy. As a result, the isomerization of but-1-ene and/or cis-but-2-ene in the presence of D, gives nondeutero isomerized products ;3 the formation of ,MH-site is discussed in the following section.12 ALKENE REACTIONS WITH H ON MOS, CATALYST n-PROPYL AND ISOPROPYL INTERMEDIATES Similarly to the hydrogen exchange between but-1-ene and D2, the hydrogen exchange between propene and D2 is very slow although the intermolecular hydrogen exchange of propene is markedly enhanced by hydrogen.A microwave spectro- scopic analysis of [2H,]pr~pene slowly formed by hydrogen exchange between D2 and propene is shown in fig. 2. It is known that the vinyl hydrogens in propene are especially active in the exchange, as observed for but-1-ene in table 1; that is, the deuterium atom of [2Hl]propene is randomly distributed within the vinyl group, and [3-2H,]propene formation is unexpectedly slow provided that [ l-2Hl]propene is formed via the isopropyl intermediate. 6 i = 1 C idif6 FIG. 2.-Relative proportion of [2Hl]propene isomers formed in the reaction of propene with D2 at room temperature. + , CH2=CH-CH2D (3-[2H1]) ; 0, CHZ=CD--CH3 (2-[’H11) ; A, H CH3 D CH3 \ / \ / C=C ((E)-1-[2H111; A, C=C {(Z)-1-[2Hll 1.D ’ ‘H H ’ ‘H As hydrogen exchange between D2 and propene is slow, the initial exchange position of the deuterium atom is completely erased by a rapid reshuffle of the vinyl hydrogens through the intermolecular hydrogen exchange reaction. Consequently, to estimate the initial exchanged position, intermolecular hydrogen exchange between [2H,]propene and [2H6]propene was performed in the presence of hydrogen. As shown in table 3, the intermolecular hydrogen exchange proceeds stepwise. Accord- ingly, [2H,]propene appearing in the initial stage of the exchange reaction might not be reshuffled. Note that the ratio of [2-2H,]propene to [l-2H,]propene is about 713 in fig. 3, and that this ratio is approximately equal to that of [2-2H,]-but-l-ene to [l-2H,]but-l-ene in table 2.As mentioned above, if the hydrogen exchange of propene occurs via n- and/or isopropyl intermediates, the unusually low value of the ratio of [3-2H,]propene to [l-2H,]propene should be considered carefully. To confirm the contribution of theT . OKUHARA AND K . - I . TANAKA 13 alkyl intermediates in the hydrogen exchange and/or the isomerization reactions, hydrogen exchange between (Z)-[l-2Hl]propene and [2H6]propene, which is a method proposed by Kondo et aZ.,9, has been adopted. The (Z)-[l-2H,]propene used as a reactant gas in this experiment contained impurities, as shown in table 4(a), and hence is termed “crude (Z)-[l-2H,]propene” in this paper. A mixture of TABLE 3.-INTERMOLECULAR HYDROGEN EXCHANGE BETWEEN r2HO] PROPENE (3.9 m m g ) AND [’H6] PROPENE (3.1 mmHg) IN THE PRESENCE OF A MIXTURE OF H2 AND D2 (7 mmHg) ON MoS2 (0.lg) AT ROOM TEMPERATURE 3 X GHif i = O (6-i)zHi time/ 2Ho 2H 1 2H2 2H3 2H4 2H5 2H6 i = 4 min 0 55.2 3 50.2 6 47.5 12 41.8 24 35.0 36 30.3 48 26.9 130 16.2 - 2.8 5.3 0.3 - 0.7 7.2 (8.0)t 0.5 - 1.0 9.2 12.5 1.6 0.3 2.3 12.9 (16.3)T 3.5 0.9 4.7 15.6 19.3 5.3 2.0 6.8 17.2 20.9 6.7 2.5 8.1 18.4 (23.1)t 12.5 7.3 18.9 18.0 I - - t Subject to the microwave spectroscopic analysis shown in fig.concurrently occurring; 5.1 % at 55 min and 10.0 % at 130 min. 42.2 2.8 36.3 14.8 33.6 20.2 28.6 34.1 23.9 51 .O 19.2 66.2 16.6 76.4 8.7 116.4 eq. 190 5. Hydrogenation reaction 5 mmHg crude (Z)-[1-2Hl]propene and 5 mmHg [2H6]propene (96 %) containing 5 mmHg HD gas (98 %) was admitted to 0.1 g MoS, catalyst at room temperature.The primarily exchanged products of this exchange reaction are [2H2]propene and [ H,]propene. The [2Hl]isomers other than (Z)-[1-2Hl]propene are formed by a conformational change around the double bond of propene with no isotopic exchange. The results TABLE 4.-INTERMOLECULAR HYDROGEN EXCHANGE BETWEEN ( Z ) - [ 1 - 2 H l ] ~ ~ ~ ~ ~ ~ ~ (5 m H g ) AND [’HtiIPROPENE (5 mmHg) IN THE PRESENCE OF 5 m H g HD AT ROOM TEMPERATURE (a) composition of crude (Z)-[1-2Hl]propene deuterium distribution/ % geometrical isomers/ % [‘Ho]propene 3.6 [2Hl]propene 93.6 (2)-[1-’H1] 92 (E)-[l-2H1] 8 [’H2]propene 2.8 (2)-[1,2-’H,] 49 (E)-[l,2-2H2] 0 (l,l-2H2) 51 (b) deuterium distribution change in the reaction of (Z)-[1-2Hl]propene and E2H6]propene time/ min 2Ho 2H1 2H2 2H3 2H4 2H5 2H6 HZ HD D2 0 1.8 46.7 1.4 - - 2.0 48.1 1.9 98.1 - 2.5 2.0 44.1 4.0 - - 4.7 45.3 2.5 97.2 0.3 5.0 2.1 41.7 6.0 0.1 0.1 6.2 43.6 10 2.4 37.9 9.9 0.8 0.8 10.0 38.4 20 2.5 30.6 15.0 2.1 2.0 14.3 32.8 2.5 96.0 1.514 x \ 3 6 0 - .?= .- i3 8 40- 20 ALKENE REACTIONS WITH H ON MoS, CATALYST h " - Y 4 l o o k 80 0- 0 20 40 60 80 100 120 140 I( 3 6 i - 1 i = 4 C id+ C (6-i)di 3 FIG.3.-Relative proportions of [2Hl]propene isomers formed by the intermolecular hydrogen exchange between [2Ho]propene and [2H6]propene at room temperature. 0, (2-[2H11) ; 0 , (3-[2H~1> ; A, ((E)-1-[2H11} ; A, {(Z)-~-['HII}- time /min I H. X 40t ,c = C\ D H I 1 I I 5 10 15 20 time/min FJG. 4.-Intermolecular hydrogen exchange between (Z)-[1 -*Hl]propene and [2H,]propene at room temperature.(a) Relative proportion of [2Hl]propene isomers. (b) Relative proportion of [2H2]propene isomers.T . OKUHARA A N D K . - I . TANAKA 15 of the exchange reaction are summarized in table 4(6), and fig. 4(a) and (b) show the time dependence of concentration of the [2H,]propene and [2H,]propene isomers. The initial composition of the [2H2]propene isomers estimated by extrapolating the curves in fig. 4(b) is 71 % (E)-[1,2-2H,]propene, 16 % [1,1-2H,]propene and 13 % (2)-[ 1 ,2-2H,]propene. By making corrections for the impurities included in the starting (Z)-[l-2H,]propene, the [2H,]isomers formed directly from (Z)-{l-2H,]propene on MoS, are estimated as 81.6 % (E)-[1,2-2H,]propene and 18.4 % [l,l-2H,]propene. About 13 % (Z)-[1,2-2H2]propene observed in fig.4(b) in the initial stage is formed mainly from the (E)-[l-2H,]propene which was included in the starting materials as an impurity. The initial compositions of [2H,]propene and [2H,]propene are also estimated, as listed in table 5, by extrapolation. The values obtained by extrapolation for the [,H,]-, [,HI]- and [2H2]-isomers are in excellent agreement with the calculated values assuming an associative mechanism under strict cis-stereochemistry ; this is discussed in the next section. DISCUSSION A noticeable hydrogen promoting effect such as observed in fig. 1 has been accepted as evidence for formation of alkyl intermediates in the isomerization and/or the hydrogen exchange of olefins. In such a case, it has been tacitly assumed that the intermediates in the isomerization and/or the hydrogen exchange of olefins overlap with those of the hydrogenation reaction, that is, the reverse process of alkyl inter- mediate formation in the hydrogenation reaction causes the isomerization and/or the hydrogen exchange of olefins.12 Such an approach, however, is incompatible with the results on oxides and sulphide~.~-~* l3 A typical case is the reaction of but-1-ene with D2 on MoS, catalyst.By adding D2 to but-1-ene, both the isomerization and the hydrogenation of but-1-ene begin simultaneously, but the products are [1,2-,H2]- butane and [2H,]but-2-ene, respectively, even though both reactions could occur via half-hydrogenated butyl intermediate^.^. This indicates that a s-butyl intermediate if formed during the isomerization reaction, is not hydrogenated to butane.The hydrogen exchange reaction as well as the isomerization of olefins most probably proceeds via alkyl intermediates. Recent experiments in our laboratory on MoS, single crystal wafers have proved that the active sites are located on the edge surface of the MoS, cry~tal,~. upon which exposed molybdenum atoms may be in coordinative unsaturation. Accordingly, the hydrogenation reaction may proceed on sites with three degrees of coordinative unsaturation, while the isomerization occurs on sites with two degrees of coordinative unsaturation as described by scheme 1 and 2., The presence of the hydrogen H \ / promoting effect observed in fig. 1 suggests the formation of an 2MH-site, Mo(S),, on which adsorbed olefin is converted to alkyl species by picking up a hydrogen atom.The ,MH-site formation process remains undetermined but is assumed to be uia heterolytic dissociation of an H2 molecule onto an ,M-site (,MH) and sulphur atom (SH), and this dissociation seems to be irreversible at room temperature. Accordingly, if a mixture of D2 and but-1-ene is introduced onto the catalyst, ,MD-sites are formed first but are changed to ,MH-sites in forming alkyl intermediates during the isomeriza- tion reaction. That is, the sites used in the isomerization reaction are almost ,MH- instead of ,MD- during isomerization because the isotopic exchange of H-atoms on16 ALKENE REACTIONS WITH H ON MoS, CATALYST 2MH-sites with D2 is slow for the following reasons; (i) the 2MH-site formation process is irreversible and (ii) the reversible dissociation of D, cannot take place on 2MH-sites because the dissociation of D2 requires two coordinative vacancies.Therefore, the isomerization of but-1-ene in the presence of D2 yields >90 % C2H,]but-2-ene. Investigation of the coisomerization reaction of ci~-[~H,]but-2-ene and cis- [2H,]but-2-ene clarifies the processes of hydrogen addition and elimination taking place on the 2MH-sites,11 Taking the conformation of the s-butyl intermediate into account, the cis to trans rotation of but-2-ene is necessarily accompanied by hydrogen exchange as described in reaction scheme 3 on the 2MH-site. H3C, ,W3 S, 1 /D cis-elim. S' 1's cis-add. C=C, + Mo Mo H' H S SCHEME 3.--cis to trans rotation regulated bv cis-stereochemical hydrogen-transfer on the 2MH-site.This mechanism is realized by nearly 100 % ci~-[2-~H~]but-2-ene formation in the coisomerization reaction. Either the double bond migration or the cis-trans rotation reactions occur via s-butyl intermediates, while hydrogen exchange of but-1-ene may take place via both n-butyl and s-butyl intermediates, which is similar to the reactions of a-olefins. As described in scheme 4, hydrogen exchange of a-olefin via an n-alkyl intermediate gives [2-2Hl]-olefin, while exchange via a s-alkyl intermediate gives [1-*H,]-olefin. f 4 S, ! ,D S S'Y'S 7 /H HzC" 'D s, I ,' M i s' g s =i H, ,R / c = C\ H D H .R D H ;c =c, or S, i ,H s0 's + Mo S, I,H I- Mo s' A'S SCHEME 4.-Hydrogen exchange via n-alkyl and s-alkyl intermediates. The ratio of the two geometrical isomers accordingly represents the relative contribution of the two intermediates, n-alkyl and s-alkyl, in the hydrogen exchange reaction.As shown in table 2 and fig. 3, the ratios of [1-2H,]-olefin/[2-2H,I-defin are 3/7 for propene as well as for but-I-ene. This result may indicate that the a- olefins adsorbed on the 2MH-sites take the n-alkyl form more readily than the s-alkylT. OKUHARA AND K . - I . TANAKA 17 form, i.e., anti-Markovnikoff hydrogen addition predominates on the 'MH-sites. In contrast, the intermediates in the hydrogenation reactions which proceed on 3M-sites or on 3MH-sites, as shown in scheme 1, are s-alkyl~.~. l4 The fact that the n-alkyls predominate on the 2MH-sites which have tight spacing and the s-alkyls are more prevalent on the 3M-sites which have wider openings is comparable with a result of Wilke and coworkers,15 who observed that the size of the ligands of nickel complexes determines the proportion of n-propyl to isopropyl species in the dimerization of propene.If the contribution of the isopropyl intermediate is ~ 3 0 % total hydrogen exchange of the a-olefins over MoS, as estimated in this paper, the question arises as to why [3-2Hl]propene formation is so slow compared with [ l-2Hl]propene formation, as shown in fig. 2 and 3. If the isopropyl intermediate formed on the 2MH-site, (isopr o pyl) Mo(S),, would have an equal probability of losing a hydrogen atom from either of the two methyls, the ratio of [3-2Hl]propene/[l-2Hl]propene would be 3/2. If [l-2H,]propene is formed via an isopropyl intermediate, abnormally slow [3-2Hl]- propene formation indicates inequality of the two methyl groups in the isopropyl species formed on the 2MH-sites.In order to establish isopropyl intermediate formation and to estimate the contribution of dissociative type hydrogen exchange, hydrogen exchange between (Z)-[l-2H,]propene and [2H,]propene was performed. \ / ( I I Assotiative Mechanism cis-od_d. - CH3Dq: Mo (n-propyl) c:%fH M o l (iso-propyl) & 1, F H 3 cis-elim. Ic='\ D H SCHEME 5.-Hydrogen exchange by an associative and a dissociative mechanism. As shown in scheme 5, under the experimental condition of an equal population of ,MH-(Z = H) and ,MD-(Z = D) sites, the proportion of each geometrical isomer can be estimated for the two mechanisms.As the hydrogen addition and elimination processes on ,MH- or ,MD-sites are regulated by the cis-stereo chemistry, the reaction routes for the all isomers formed from (Z)-[l-2H,]propene can be described by scheme 6 ;18 ALKENE REACTIONS WITH H O N MOS, CATALYST ‘[HI--+ - D 2Ho SCHEME 6.-Reaction routes from (Z)-[l-2Hl]propene. where the loops indicate that the rotation of a methyl group is occurring necessarily in the hydrogen exchange reaction under cis-stereochemical regulation. n-propyl from (Z)-[l-2Hl]propene may have equal probability of yielding (E)-[1 ,2-2H2]propene and (E)-[l-2Hl]propene providing there is an equal population of 2MH- and 2MD- sites. Similarly, iso-propyl from (Z)-[l-2Hl]propene gives (E)-[l-2H,]propene, [2H,]propene and [l,l-2H2]propei~e in the ratio 2 : 1 : 1.If the relative contribution of the n-propyl and isopropyl intermediates is assumed to be 70 and 30 %, respect- ively, the [2H,]propene isomers formed from (Z)-[l-2H,]propene may be evaluated as 82.4 % (E)-[l,2-2H2]propene and 17.6 % [l,l-2H,]propene, which is in excellent agreement with the experimental values of 81.6 % (E)-[1,2-2H2]propene and 18.4 % [l,l-2H2]propene when a correction had been made for the contribution of (E)- [1-2H,]propene contained in the starting materials. The relative proportions of TABLE 5.-PROPORTION OF EXCHANGED PRODUCTS. EXCHANGE REACTION BETWEEN (2)- [ 1 -2H 1lPROPENE AND [2H 6IPROPENE ON MOSz AT ROOM TEMPERATURE time/ mm O f 2.5 5.0 10 20 calc. [2H~1, PH11 and [2Hzl~ro~ene/% 12H01 (E)-[l-’H11 [2-2H1I (E>-[1,2-2Hd (2)-[1,2-2H2] [l,l-2H2] 3.0 51 .O 0 34.0 4.5 7.5 3.8 46.5 0 32.1 7.7 9.8 3.0 50.6 0 27.5 8.5 10.4 3.5 42.2 4.4 26.1 10.6 13.1 2.9 36.0 5.6 26.0 16.0 13.3 7.5 50.0 0 35.0 0 7.5 t A graphical extrapolation to zero time was made.[2H,]propene isomers and [2H,]propene can be similarly estimated. As shown in table 5, on the calculation run the relative compositions calculated for all isomers are in excellent agreement with the experimental values extrapolated to zero time. These results prove that the hydrogen exchange reaction of propene proceeds entirely via n-propyl and isopropyl intermediates on MoS, catalyst. Accordingiy, the iso- propyl species coordinated to the 2M-sites is unsymmetrical in elimination of hydrogen from the two methyls. A slight inequality between the two methyls of the iso-propyl for hydrogen elimination has been reported on the EDA-complexes of phthalocyanine with alkali metals;I6 however, such large values (> 15) for the ratio ~f[l-~H,]propene/ [3-2H,]propene have not been observed in the past.The structure of the isopropyl isopropyl Mo(S),, suggests that one of \ intermediate coordinated to an 2M-site, /T. OKUHARA AND K . - I . TANAKA 19 the two methyls is closely adjacent to a vacant coordination site. Accordingly, if the isopropyl species bonded to the 2M-site is restricted from free rotation around the coordination bond, the two methyls undergo hydrogen atom elimination with different probabilities, as described in scheme 7 : , CH3-CH-f: H2 CHj-CH=CHD s, i ,D - H3C-CH=CH2 + Mo I ,,D, s’ I ‘s Mo’ [‘HJ prop-1-ene S Rotation 1 C H2 D-CH- CH, I ,-A - CHzD-CH=CHZ Mo‘ FHA prop-3- ene SCHEME 7.-Restricted rotation of the isopropyl intermediate formed on an *M-site.where the abstraction of a hydrogen atom from a methyl group occurs at a vacant coordination site which appears after donation of a hydrogen atom to adsorbed propene. If the rotational motion of the isopropyl species bonded to the 2M-sites is restricted either by the spacing and/or the configuration of sites, a similar restriction might be expected to hold for other s-alkyl species bonded to the sites. If the s-butyl is restricted from rotating around the coordination bond, two distinguishable conformations may be expected on the 2M-sites depending on the direction of the vacant site, as described in scheme 8.Rotation CHzD-CH-CH-CH3 s, i ,H - (sec- b u t y 1 -II 1 Mo 7 I -l!l Mo’. But-2-ene + s’ 1‘s S SCHEME 8.-Two distinguishable s-butyls formed on 2M-site. If the energy barrier for rotational motion of the s-butyl species is high, the vinyl hydrogen atoms of but-1-ene undergo rapid hydrogen exchange via s-butyl-I and n-butyl intermediates, and the cis-trans rotation of but-Zene also occurs rapidly via s-butyl-11. However the double bond migration reaction of but-1-ene to but-2-ene is expected to be slow because this isomerization reaction must pass over the rotational barrier from s-butyl-I to s-butyl-11. This is the case for the reactions of but-1-ene and but-2-ene on MoS, catalyst. Such slow rotation around a single bond in carbon-molybdenum may not be surprising, because it is known that some bulky groups have high rotational barriers around a single b0nd.l’ Thus it may be possible that the restricted rotation of the s-alkyl species coordinated to an Mo(S), site which has a narrow opening may be caused by the steric interaction of the s-alkyl species with the sulphur ions.These results remind us that allene and methylacetylene are hydrogenated to propene on \ /20 ALJSENE REACTIONS WITH H ON MoS, CATALYST MoS, catalyst via different intermediates, o-ally1 intermediate and n-propenyl intermediate respectively.’ They prove that the reactions over MoS, catalyst prefer intermediates taking lesser strain. As demonstrated here and in our previous work, the degree of coordinative unsaturation of active sites is a structural prerequisite for the catalytic hydrogenation reaction as well as for the isomerization and/or the hydrogen exchange reactions taking place via the alkyl intermediates.In this sense, these reactions should be structure sensitive. Recently, the isomerization of 2-methylbut-1-ene was found to proceed without hydrogen over MoS, catalyst although the isomerization of but-1-ene and of but-Zene requires the assistance of hydrogen.lg By using a single crystal of MoS,, it has been shown that the isomerization of 2-methylbut-1-ene takes place on the basal plane of the MoS, single crystal which is composed of a sulphur sheet, whereas the isomerization of but-1-ene or of but-Zene is promoted only on the edge surface of the crystal in the presence of hydr~gen.~ These facts indicate that the isomerization of 2-methylbut- 1-ene proceeds via a carbonium ion intermediate instead of the alkyl intermediates elucidated here.It is reasonable to deduce that isomerization via carbonium ion intermediates is virtually structure insensitive, because carbonium ion formation is entirely controlled by the proton activity of the surface. The coordinative unsaturation number is undoubtedly a structural prerequisite of the active sites upon which alkyl species are formed and consecutive hydrogenation of the alkyls requires one additional coordinative vacancy. It should be emphasized here that more complex catalytic reactions, such as ammonia synthesis and the methanation reaction of carbon monoxide, should have different structural prerequisites. Recently, Demitras and Muetterties” have shown an interesting example suggesting the existence of such complex structural requirements including ensemble operation of several sites using metal clusters.The authors thank Dr. T. Kondo of Sagami Chemical Research Centre for providing the microwave spectroscopic analysis, and they are also indebted to Prof. K. Miyahara of our Institute for his contribution to the mass spectrometry. M. Boudart, A. Aldag, J. E. Benson, N. A. Dougerty and C. G. Harkins, J. Catalysis, 1966 6, 92; D. W. Blakely and G. A. Somorjai, J. Catalysis, 1976, 42, 181. A. Takeuchi, K. Tanaka and K. Miyahara, Chem. Letters, 1974, 171, 411 ; A. Takeuchi, K. Tanaka, I. Toyoshima and K. Miyahara, J. Catalysis, 1975,40,94 ; A. Takeuchi, K. Tanaka and K. Miyahara, J. Catalysis, 1975, 40, 101. K. Tanaka, T. Okuhara, S. Sat0 and K. Miyahara, J. Catalysis, 1976, 43, 360. S. Siegel, J. Catalysis, 1973, 30, 139 ; R. L. Burwell, Jr., G. L. Haller, K. C. Taylor and J. F. Read, Advances in Catalysis (Academic Press, N.Y., 1969), vol. 20, p. 1 ; K. Tanaka, J. Catalysis 1975, 37, 558. K. Tanaka and T. Okuhara, Catalysis Rev. Sci. and Eng., 1977, 15, 249. H. Arai and J. Halpern, Chem. Comm., 1971, 1571. T. Okuhara and K. Tanaka, J. Phys. Chem., 1978, 82, 1953. T. Okuhara and K. Tanaka, J. Catalysis, 1976, 42, 474. T. Kondo, S. Saito and K. Tamaru, J. Amer. Chem. SOC., 1974,96,6857. ’ T. Okuhara, K. Tanaka and K. Miyahara, 41st Catalysis Suc. Meeting (Toyama, 1977); lo S. Saito, J. Mol. Spectr., 1969, 30, 1. l 1 T. Okuhara and K. Tanaka, J. Amer. Chem. SOC., 1976,98, 7884. l 2 CJ G. C. Bond, Catalysis by Metals (Academic Press, London and N.Y., 1962). l3 K. Tanaka, J. Catalysis, 1975, 37, 558. l4 T. Okuhara, T. Kondo, K. Tanaka and K. Miyahara, J. Phys. Chem., 1977,81,90 ; T. Okuhara, l5 B. Bogdanovic, B. Henc, H.-G. Karmann, H. C. Nussel, D. Walter and G. Wilke, Ind. and T. Kondo and K. Tanaka, Chem. Letters, 1977,119 ; 1976,717. Eng. Chem., 1970,62,34.T. OKUHARA A N D K . - I . TANAKA 21 l6 S. Naito, M. Ichikawa, S. Saito and K. Tamaru, J.C.S. Farahy I, 1973, 69, 685. l7 Dynamic Magnetic Resonance Spectroscopy, ed. L. M. Juckman and F. A. Cotton (Academic l 8 T. Okuhara, T. Kondo and K. Tanaka, Chem. Letters, 1977, 119. l9 T. Okuhara, K. Tanaka and K. Tanabe, J.C.S. Chem. Comm., 1977, 180. 2o G. C. Demitras and E. L. Muetterties, J. Amer. Chem. SOC., 1977, 99, 2796. Press, 1975) ; G. Natile, L. Cattalini and F. Gasparrini, J.C.S. Chem. Comm., 1977, 89. (PAPER 7/2202)

ISSN:0300-9599    

DOI:10.1039/F19797500007

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Pulse Radiolysis Study of Chlorpromazine and Promazine Free Radicals in Aqueous Solution BY A. KEITH DAVIES Department of Chemistry and Applied Chemistry, University of Salford EDWARD J. LAND Paterson Laboratories, Christie Hospital and Holt Radium Institute, Manchester AND SUPPIAH NAVARATNAM, BARRY J. PARSONS" AND GLYN 0. PHILLIPS School of Natural Sciences, North E. Wales Institute, Connah's Quay, Clwyd Received 18th January, 1978 .OH radicals may react with chlorpromazine in four different ways, viz. (i) electron transfer to produce the cation radical, (ii) addition to the sulphur atom followed by acid-catalysed OH- elimination to yield the cation radical, (iii) addition to the aromatic rings to produce cyclohexadienyl type radicals, and (iv) abstraction of hydrogen atoms from the -CH2- which is in the cc position to the ring nitrogen.Electron transfer and addition to the sulphur atom each account for 40 % of the *OH radicals. Similar considerations apply to promazine. The hydrated electron may attach itself to either the sulphur atom in promazine or its aromatic system with equal probability. Addition to the sulphur atom probably leads to immediate cleavage of the carbon-sulphur bond whereas addition to the aromatic system probably produces an anion radical. The lifetime of this anion radical is sufficiently long to enable electron transfer between itself and another promazine molecule to occur to produce the promazine cation radical. With chlorpromazine, the hydrated electron probably reacts solely with the aromatic system to eliminate chloride ions by the end of the reaction.The promazine radical which is produced simultaneously may add on to another chlorpromazine molecule to produce a cyclohexadienyl type radical. Chlorpromazine (I) and promazine (CH2)3 1 /*\ CH3 CH3 ( I 1 are known to photoionise monophotonically in aqueous solution when excited by light of 347 nm wavelength. Photoionisation of the parent compound, phenothiazine has also been shown to occur in aqueous micellar and methanolic solutions.2 These compounds find use as drugs in psychiatric treatment where side effects include photosensitisation of the skin and eye ti~sues.~ In micellar systems, they also 22DAVIES, LAND, NAVARATNAM, PARSONS A N D PHILLIPS 23 provide model systems for photosynthesi~.~ A study of the fates of the cation radicals and solvated electrons produced on photoionisation of phenothiazine-based derivatives can be most easily undertaken using pulse radiolysis techniques where difficulties due to overlapping absorption by the triplet excited state of the drug are absent.Radical cations may be generated in aqueous solution using free radicals such as SO;- and *OH radicals as oxidants.6 Hydroxyl radicals, for example, have been shown to react with methoxylated benzenes to produce initially hydroxycyclo- hexadienyl radicals which upon protonation and water elimination yield radical catiom6 In this work, the possibility of direct electron transfer between the drug and either *OH or Bri- was investigated. The reaction of the hydrated electron with halogenated aromatics has been the subject of several studies, and in these cases, has resulted in the cleavage of the carbon-halogen bond to yield halide In view of this, clear differences may be expected in the reactions of promazine and chlorpromazine. EXPERIMENTAL Pure samples of chlorpromazine hydrochloride (ClP) and promazine hydrochloride (PH) were kindly donated by May and Baker and Cohn Wyeth and Brothers, respectively.All other reagents were of AnalaR grade. Solutions were prepared using water which was distilled from alkaline permanganate. The pH of chlorpromazine and promazine solutions was normally in the range 5-6. pH values outside this range were obtained by addition of either perchloric acid or sodium hydroxide solutions. When required, the ionic strength of solutions was adjusted by addition of sodium perchlorate.Solutions were saturated with either argon (Air Products) or nitrous oxide (Air Products). The pulse radiolysis experiments were carried out with the electron linear accelerator at the Paterson laboratories lo- Where appropriate, cut-off filters were used to minimise photolysis of the solutions by the analysing light. In all experiments, a single Bausch and Lomb monochromator was used with bandwidths of 10 nm. An E.M.I. 95584 photomultiplier was used in the majority of experiments but for monitoring at wavelengths up to 1000 nm, a U.D.T. PIN10 photo- diode was also necessary. The response times of the two amplifier circuits employed were ~7 and 100 ns. Doses were monitored by a secondary emission chamber calibrated against aqueous potassium thiocyanate l2 taking the yield of the species absorbing at 500 nm to be 2.9 molecules per 100 eV and its extinction coefficient to be 7.1 x lo2 m2 mol-l.The optical path length was 2.5 cm. For kinetics measurements, photographs of the oscilloscope traces were magnified four times and, by using an appropriate computer program, data from these were used to obtain the first-order plots. using pulse lengths up to 100 ns. RESULTS REACTIONS OF HYDROXYL RADICALS WITH PROMAZINE A N D CHLORPROMAZINE Fig. 1 shows the absorption spectra obtained immediately, and 2 5 0 p after an electron pulse was delivered to a nitrous oxide saturated solution of chlorpromazine mol dm-3). These spectra have been corrected for loss of chlorpromazine absorption assuming G(0H) = 5.5. The end of pulse spectrum shows maxima at 850, 770, 525, 440 and <350 nrn (measurements at wavelengths below this were not possible because of the strong absorption of chlorpromazine).The absorptions at 850,770 and 525 nm increased by almost a factor of two over 250 ,us before decaying, whereas the absorptions at 440nm and that at 350nm decreased. In addition a growth in absorption at 950nm was observed over 2ms. These spectra were not affected by changing the chlorpromazine concentration from 5 x to 5 x mol dm-3.24 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS The spectrum observed immediately ( ~ 1 ps) after the pulse is due to products formed in the reaction of *OH radicals with chlorpromazine. By making observations at shorter time scales, the growth in this absorption was found to follow first-order kinetics and be directly dependent on the chlorpromazine concentration over the range 5 x 10-5-5 x mol dm-3.The second-order rate constant was determined as (8.3 k0.4) x lo9 dm3 mol-l s-l. At 525nm, the subsequent growth over about 250ps also followed first-order kinetics with a rate constant of (1.4k0.2) x lo4 s-l [fig. 2(a)]. Similar observations were also made at 770 and 850 nm. The initial absorption at 350 nm decayed by first-order kinetics yielding a similar rate constant of (1.5f0.2) x lo4 s-l [fig. 2(b)]. 3.c 2.0 m X s I .c 0 I ioo 400 5 0 0 6 0 0 700 800 900 h/nm FIG. 1 .-Transient absorption spectra (after correction for loss of chlorpromazine) at various times after pulsing a nitrous oxide saturated solution of chlorpromazine mol dm-3), (-0-) end of pulse ; (-U--) after 250 ps.Dose = 225 rad, path length 2.5 cm. Neither variation of dose per pulse (80 rad to 1.0 krad) nor chlorpromazine con- centration (5 x mol dm-3) was found to affect the first-order rate constant measured at either 350 or 525 nm. The first-order rate constants at both of these wavelengths were however identically affected by pH. Table 1 summarises these results for 525 nm where the measurements are most accurate and shows that a decrease in the pH leads to an increase in the rate constant. At 440nm, the absorption measured immediately after the pulse showed no dependence on chlorpromazine concentration and decayed according to first-order kinetics but over a longer time-scale than that for the 350 and 525 nm absorptions [fig.2(c)]. The first-order rate constant was found to be (2.850.3) x lo3 s-l. Similarly, at 950 nm, the increase in absorption followed first-order kinetics [fig, 2(d)] yielding a rate constant of (2.8k0.3) x lo3 s-I. Variation of dose per pulse, (200 rad to 1.5 krad) chlorpromazine concentration (5 x to 5 x mol dm-3) and pH had no effect on the first-order rate constant. to 5 xDAVIES, LAND, NAVARATNAM, PARSONS AND PHILLIPS 25 After allowance for the absorption of the cation radical, the spectrum measured at the end of this reaction showed peaks around 440 and 900 nm, respectively. By comparing the absorption spectra obtained after chemical oxidation l3 and photolysis of aqueous chlorpromazine solutions, it would appear that a major portion of the spectra in fig.1 could be assigned to the chlorpromazine cation radical. I -I 0.c 0 0.5 ID - 8 . 0 ~ 100 200 0 I 0 0 5 I .o I .5 I 0 5 0 100 time/p timelps timelms timelms FIG. 2.-First-order plots for changes in absorption after pulsing nitrous oxide saturated solutions of chlorpromazine (lW4 mol dm-3) in a 2.5 cm path length : (a) h = 525 nm, dose = 250 rad ; 6) h = 350 nm, dose = 250 rad; (c) h = 440 nm, dose = 800 rad; (d) h = 950 nm, dose = 800 rad.26 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS In order to provide further evidence for this, the radical anion, Bri-, was used for the oxidation. The species is a milder oxidising agent than -OH and might be expected to yield a single oxidised species from chlorpromazine by electron transfer.On pulsing a nitrous oxide saturated solution of KBr (1 x mol d ~ a - ~ ) in the presence of mol dm-3 chlorpromazine, the absorption immediately after the pulse at 360 nm (attributable to Br;-)14 disappeared by first-order kinetics with a rate constant of (5.0k0.3) x lo5 s-l. Simultaneously, a first-order growth in absorption was observed at 525 nm with a rate constant of 5.4k0.3 x lo5 s-I. The TABLE EFFECT OF pH ON FIRST-ORDER RATE CONSTANTS FOR GROWTH IN ABSORPTION AT 525 nm k l x 10-4 pH Is-1 (525 nm) 3.75 6.1 4.04 4.3 4.66 3.4 5.46 2.9 Nitrous oxide saturated chlorpromazine solutions, 5 x mol dm-3 ; all solutions were adjusted to an ionic strength of rnol d~n-~). Dose = 275 rad, path length 2.5 cm. second-order rate constant for the reaction of Bri- with chlorpromazine was thus determined as (5.0k0.3) x lo9 dm3 mol-1 s-l.In this solution, no subsequent grow- in of absorption over -250 ps at 525 nm was observed. The spectrum at the end of this reaction is shown in fig. 3. The decay rate of this spectrum was wavelength independent suggesting it corresponds to a single species, the chlorpromazine cation radical. The spectrum of fig. 3 yields an extinction coefficient for the chlorpromazine cation radical at 525 nm of (1.00+0.08) x lo3 m2 mol-l. Using this value, the G- value for the production of the cation radical in N20 saturated solutions at the end hlnm FIG. 3.-Transient absorption spectrum measured 14ps after pulsing a nitrous oxide solution of KBr (1 x mol dm-3) containing rnol dnr3 chlorpromazine.Dose = 215 rad, path length 2.5 cm.27 of the -OH reaction is estimated as 2.2. A further yield of the cation radical (G = 2.2) is found at the end of the slower pH dependent step (lasting 250 zps after the pulse). The reaction of hydroxyl radicals with promazine produces species whose absorp- tion spectra are similar to those obtained with chlorpromazine (fig. 4). The end of pulse (-1 ps) spectrum shows peaks at 850, 760, 515, 440 and ~ 3 5 0 nm. The initial absorptions at 850, 760 and 515 nm, as for chlorpromazine, almost doubled in intensity (following first-order kinetics) over 150 p s after the pulse, whereas those at DAVIES, LAND, NAVARATNAM, PARSONS AND PHILLIPS 2 X 8 h/nm FIG. 4.Transient absorption spectra (after correction for loss of promazine) at various times after pulsing a nitrous oxide saturated solution of promazine mol dm-3), (-0-) end of pulse ; (-Cl-) after 150 ps.Dose = 250 rad, path length 2.5 cm. 440 and 350 nm decreased at a much slower rate. As found for chlorpromazine, the decay in absorption at 440 nm was first-order and matched a growth in absorption at 950 nm. All these first-order rate constants were unaffected by dose per pulse or promazine concentration. The rate constant for the growth in absorption lasting -150 p s at 515 nm was found to be (5.8k0.5) x lo4 s-l. This was measured at pH 5.5 only. The rate constant for the reaction between .OH radicals and promazine was determined as (3.7k0.1) x lo9 dm3 mol-1 s-l. REACTION OF HYDRATED ELECTRONS WITH PROMAZINE AND CHLORPROMAZINE The rate of reaction of the hydrated electron with promazine was followed at 720 nm by pulsing argon saturated solutions containing promazine and t-butanol (0.5 mol dm-3).The decay of the hydrated electron absorption followed first-order kinetics for the range of promazine concentrations 2.5 x to 5 x mol dm-3. After allowance for other modes of hydrated electron decay at low promazine concentration, the first-order rate constant was found to be directly proportional to promazine concentration yielding a second-order rate constant for the reaction between hydrated electrons and promazine of (5.3 k0.3) x lo9 dm3 mol-' s-l. The spectrum of the product(s) formed at the end of this reaction was measured28 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS in a rnol dm-3 promazine solution and is shown in fig.5. The effect of proma- zine concentration on this spectrum was investigated and it was found that the absorption band at 440 nm was affected. At mol dm-3 promazine the spectrum at the end of the pulse was approximately half that measured in a mol dm-3 solution. Also shown in fig. 5 is the spectrum at 90 ,us after the pulse which has maxima at 515, 760 and 850 nm. In addition, there is a shoulder at about 450 nm. The growth in absorption at 515nm which occurs for 9Ops after the end of the hydrated electron reaction follows first-order kinetics giving a rate constant of (3.5k0.4) x lo4 s-l at mol dm-3 promazine. Identical kinetics were found for the decays of the absorptions at 350, 925 and 975 nm.Experiments were also carried out using sodium formate instead of t-butanol as a scavenger for OH radicals. These yielded essentially similar results. 16 m 8. X n 0 rr. h/nm FIG. 5.-Transient absorption spectra at various times after pulsing an argon saturated solution of promazine mol dm-3) containing 0.5 mol dm-3 t-butanol, (-0-) end of hydrated electron reaction (- 1 ps) ; (-El-) 90 ps after pulse. Dose = 410 rad, path length 2.5 cm. It is clear from fig. 5 that a major portion of the spectrum which is observed 90 ps after the pulse is attributable to the promazine cation radi~a1.l~ From the shape of the promazine cation radical spectrum observed in the chemical oxidation of promazine l3 and also by analogy with the chlorpromazine cation radical spectrum generated in this work using Bri- (fig. 3), it is also apparent that the promazine cation radical absorption should contribute little ( N 10 %) to the 440 nm absorption observed 90 p s after the pulse. However, since there is relatively little change in absorption at 440 nm between the end of the hydrated electron reaction and 90 ,us later, it would, therefore, appear that the kinetic behaviour at this wavelength is different from that observed at 350, 525, 925 and 975 nm.In order to investigate further the conditions governing the formation of the promazine cation radical, the effect of promazine concentration on its yield and rate of formation was studied. Fig. 6(a) shows the effect of promazine concentration (5 x 10-5-5.0 x rnol dm-3) on the first-order rate constant for the cation radicalDAVIES, L A N D , NAVARATNAM, PARSONS A N D PHILLIPS 29 formation.It is apparent from this figure that an increase in the promazine con- centration results in an increase in the rate constant, and, that there is a small but detectable value when the plot is extrapolated linearly to zero promazine concentration. From a consideration of all possible errors, this value was estimated as (3.5 10.5) x lo3 s-l. The second-order rate constant obtained from the slope of this plot was determined as 3.3 x lo7 dm3 mol-1 s-l. The dose per pulse was = 100 rad in these experiments which was considered to be sufficiently low to avoid unwanted radical- radical recombination. From the same experiments, it was also observed that the amount of cation radical formed at the end of this reaction also increased with n I X E: [promazine] x 103/mol dm-3 X n 0 Y [promazinel-I x 10-4/dm3 mol-I FIG.6.-(a) Plot of the first-order rate constant for the formation of the promazine cation radical against promazine concentration. Dose 100 rad, path length 2.5 cm. (b) Plot of the reciprocal optical density (measured at 515 nm at the end of the cation radical formation) against reciprocal promazine concentration. Dose = 100 rad, path length 2.5 cm. increase in promazine concentration. Fig. 6(b) summarizes this data in the form of a plot of the reciprocal optical density due to the promazine cation radical at 515 nm against reciprocal promazine concentration. The slope of this plot was determined to be 1.1 x mol dm-3.From fig. 6(b) it can be deduced that the maximum optical density due to the cation radical at 515 nm that could be expected under these conditions at high promazine concentration would be 3.7 x Assuming similar extinction coefficients for the promazine and chloropromazine cation radicals at their respective wavelength maxima, it can be calculated that the maximum G-value for its production is 1.43. By assuming therefore that the remainder of the hydrated electron yield (G = 2.75-1.43 = 1.32) accounts for the species formed at the end of the hydrated electron reaction which absorbs at 440 nm, and by assuming also that30 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS the decay of this species is slow compared with cation radical formation and allowing for the contribution of the cation radical to the absorption at 440 nm, the extinction coefficient of the 440 nm species can be calculated as 4.1 x lo2 m2 mol-l. The reaction of the hydrated electron with chlorpromazine is substantially different from that for promazine.The rate of the reaction was studied by following the decay of the hydrated electron absorption at 720 nm in aqueous, deaerated solutions A/nm FIG. 7.-Transient absorption spectra measured at various times after pulsing an argon saturated solution of chlorpromazine mol dm-3) containing 0.5 mol dm-3 t-butanol ; (-0-1 2 ps ; (-El--) 20 ps after pulse. Dose = 930 rad, path length 2.5 cm. containing chlorpromazine and 0.5 mol d r r 3 t-butanol. The decay followed first- order kinetics over the range of chlorpromazine concentration 5 x 10-5-2.5 x mol dm-3 yielding a second-order rate constant of (2.2k0.1) x 1O1O dm3 mol-1 s-l.It was difficult to measure the spectrum at the end of the hydrated electron reaction because of the electron’s relatively large absorption and because of a subsequent build-up in absorption which occurred over several microseconds. Fig. 7 therefore shows the spectrum at 2 p s after the pulse, at which time there may be a substantial contribution from subsequent reactions. Also shown in fig. 7 is TABLE Z-EFFECT OF CHLORPROMAZINE CONCENTRATION ON OPTICAL DENSITY AT 440llm MEASURED AT 20pS AFTER PULSING ARGON SATURATED SOLUTIONS OF CHLORPROMAZINE CONTAINING 0.5 mOl dm-3 f-BUTANOL O.D. (440 nm) [ClPI /mol dm-3 / x 102 5x 1.77 2~ 10-4 1.32 1 x 10-4 0.79 Dose = 1.8 had, path length 2.5 cm.DAVIES, LAND, NAVARATNAM, PARSONS A N D PHILLIPS 31 the spectrum measured 20 ps after the pulse when the subsequent build-up is complete.Both of these spectra have essentially the same shape with one prominent maximum at 440 nm in the range 350-1000 nm. It was observed that an increase in chlorpromazine concentration led to an increase in the amount of absorption measured 20 ,us after the pulse at 440 nm. The dose was considered low enough to minimise radical-radical reactions on this time scale and table 2 summarises the data. It was difficult, for the reasons mentioned above, to observe the effect of increasing chlorpromazine on the rate of formation of this transient and this aspect cannot therefore be discussed.DISCUSSION REACTIONS OF HYDROXYL RADICALS WITH PROMAZINE A N D CHLORPROMAZINE From this work in which chlorpromazine has been oxidised using Bri-, previous chemical l3 and photolytic oxidation experiments, it has been shown that the chlorpromazine cation radical exhibits absorption maxima at 525, 770 and 850 nm. A major portion of the end of pulse spectrum in fig. 1 is therefore attributed to this species, the yield, as argued earlier, being given by G = 2.2. The reaction can be written thus : Fig. 1 also shows that a further yield (G = 2.2) of chlorpromazine cation radical is obtained at 250 ps after the pulse. This is most probably achieved uia addition of the hydroxyl radical to the sulphur atom followed by an acid-catalysed elimination of OH- : *OH + C1P 3 ClP" + OH-.(1) OH I. OH I R I R A similar mechanism has been proposed in the formation of the radical cation of simple aliphatic and cyclic monosulphides. In those studies, dimer cation radicals were also formed at sulphide concentrations in the range 10-4-10-2 mol dm-3. In this work, no such dependence on chlorpromazine concentration (5 x 5 x mol d w 3 ) could be found for the species having absorption maxima at 440 and 525 nm. However, for either steric or resonance reasons, dimer cation radicals are not formed for R2S when R is a t-butyl group.' Stabilisation of the monomeric cation radical may also be achieved intramolecularly via electron pairs on nitrogen and oxygen.16 It would seem, therefore, that coordination of the cation radical of chlorpromazine to chlorpromazine itself is not necessarily a requisite for its stability.32 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS Our results also suggest that the *OH adduct to the sulphur atom absorbs principallj at 350nm and below.The conversion of this species into the radical cation can occur via two pathways depending upon the pH (see table 1). From a plot of the rates given in table 1 against hydrogen ion concentration, the acid independent pathway, presumably elimination of OH- has a rate constant of (2.9k0.5) x lo4 s-I whereas the acid catalysed step occurs with a rate constant of (1.2k0.3) x lo8 dm3 mol-I s-l . An alternative pathway for the production of CIP+' during the 250 p s after the pulse would be the addition of *OH to the aromatic rings followed by OH- elimination.However, the production of cation radicals via *OH addition to the sulphur atom at diffusion controlled rates appears to be a universal reaction for monosulphides whereas cation radical formation from hydroxycylohexadienyl radicals has only been demonstrated for a few aromatics having electron-donating groups which are expected to stabilise the cation radicak6 However, addition of *OH radicals at the hydrogen containing positions in the aromatic ring is likely to account for the fate of some of the remaining -OH radicals. It is also likely that some will abstract hydrogen atoms from the -CH2- group a to the ring nitrogen from comparison with studies on triethy1amine.l' The total G-value for radicals produced from these latter two processes is expected to be the residue, 1.1.The slower rates of reaction of =OM radicals with protonated tertiary amines compared with the unprotonated forms also makes attack at any other side chain position unlikely. It is proposed therefore that the species absorbing at 440 nm and formed at the end of the *OH reaction with chlorpromazine are hydroxycyclohexadienyl type radicals, e.g. The decay of this absorption at 440 nm follows first-order kinetics and is not dependent on either pH or chlorpromazine concentration. A possible explanation for this observation would be the elimination of water (and possibly HC1 as a minor process) to yield a phenyl type radical, e.g. : This reaction would therefore have a rate constant of (2.8 k0.3) x lo3 s-l. The production of cyclohexadienyl radicals absorbing at 440nm has also been observed in preliminary experiments on the reaction of hydrogen atoms with chlor- promazine. Similarly, this type of radical is produced in the reactions of the hydrated electron with promazine and chlorpromazine which are discussed below. In these instances, the amount of absorption is dependent on the substrate concentration. In order to maintain consistency therefore, with the hydrated electron reaction schemesDAVIES, LAND, NAVARATNAM, PARSONS A N D PHILLIPS 33 discussed below, it would appear that the phenyl type radical produced in reaction (5) must add onto another chlorpromazine molecule to produce a dimeric chlorpromazine cyclohexadienyl type radical.Since the spectrum at the end of reaction (5) also shows an absorption maximum at 440 nm, the rate constant for the production of the dimeric species must be > 3 x lo8 dm3 mol-1 s-l.A rate constant for the analogous reaction of the phenyl radical with iodobenzene of 7 x lo7 dm3 mol-1 s-l has been determined and provides support for this mechanism. Further support is provided by other work,21 in which it was concluded that phenyl radicals add onto halogenated benzenes to form dimeric cyclohexadienyl type radicals. The reaction of .OM radicals with promazine shows identical features to those observed for chlorpromazine and hence the above discussion would also apply to this system. REACTIONS OF HYDRATED ELECTRONS WITH PROMAZINE A N D CHLORPROMAZINE From the kinetic behaviour of the absorption spectra shown in fig.5, it is apparent from the kinetics already detailed that two species contribute to the end of pulse spectrum in the range 350-1000 nm, i.e. one absorbing principally at 440 nm whose amount is promazine concentration dependent, the other having detectable maxima at 925 and 975 nm and accounting also for most of the absorption at 350 nm. It was concluded from the .OH reaction scheme that cyclohexadienyl type radicals of chlorpromazine and promazine absorb at 440 nm. To produce such a radical at the end of the hydrated electron reaction whose amount depends on promazine con- centration, therefore, it is necessary to generate a precursor which can add on to the aromatic positions of promazine itself. A suitable precursor would again be a phenyl type radical. It is suggested that this species is formed via addition of the hydrated electron to the sulphur atom followed by cleavage of the carbon-sulphur bond... .. 0 0 1 R I R :. 0 .. (6J;B - (7) I R I R Reactions (6) and (7) might be considered analogous to the elimination of halide ion after addition of the hydrated electron. In fact, the electronegativety of sp3 hybridised sulphur atoms is greater than, for example, chlorine atoms while the bond strength of a carbon-sulphur bond is some 55 kJ mol-1 less than that of a carbon-chlorine b0nd.l The phenyl type radical produced in reaction (7) is expected to add on to promazine itself to produce a cyclohexadienyl type radical. The rate of this reaction and 1-234 PULSE RADIOLYSIS OF PROMAZINE FREE RADICALS reaction (7) must be of the same order (or larger) as the initial reaction of the hydrated electron if the rate-determining step in the production of the 440 nm species is the rate of loss of the hydrated electron itself.The remainder of the end of pulse spectrum of fig. 5 is therefore presumably attributable to a promazine anion radical where the hydrated electron has added on to the aromatic system of the molecule. Absorption spectra extending as far as 1000 nm are a common feature of aromatic and heterocyclic radical anions 2o and this fact provides support for the assignment. The decay of the absorption due to this species matches that of the growth of the cation radical absorption. The dependence on promazine concentration of the rate of the growth and yield of the cation radical shown in fig.6(a) and (b) indicate that at low promazine concentrations a first-order reaction begins to compete with the direct reaction between the radical anion and promazine itself. The latter reaction has been shown to generate the promazine cation radical, a plausible reaction being, e.g. : I R I R 1 R + I R The possibility that t-butanol readicals can oxidise promazine to form the cation radical was ruled out from the results of experiments in which formate was used instead of t-butanol as a scavenger for -OH radicals. The first-order reaction of the radical anion is presumably hydrolysis, e.g. : R R The rate constants for reactions (8) and (9) can be determined from fig. 6(a) and are (3.3k0.5) x lo7 dm3 mol-1 s-' and (3.5k0.5) x lo3 s-l, respectively.The slope of the plot in fig. 6(b) gives the ratio k9/k8 (= 1.1 x mol dm-3) whose value is in agreement with that obtained from fig. 6(a). From the maximum yield of the promazine cation radical obtained at high promazine concentrations, it can be estimated that ~ 5 0 % of the initial hydrated electron reaction involves addition to the sulphur while the other 50 % add on to the aromatic s ys tem . The rate constant for the reaction between the hydrated electron and chlor- promazine is approximately four times greater than that for promazine and obviously reflects the presence of the chlorine atom in the molecule. Although the spectrum of the product of this reaction cannot be given with any certainty, it is apparent that the spectrum measured at 20 ps after the pulse is due to a subsequent reaction of thisDAVIES, LAND, NAVARATNAM, PARSONS A N D PHILLIPS 35 product.Again, the spectrum shows a maximum at 440 nm which is dependent on chlorpromazine concentration. A scheme which takes account of these facts and maintains consistency with the other mechanisms detailed above would involve addition of the hydrated electron to the aromatic system followed by elimination of a chloride ion. Such eliminations are often found with halogenated aromatic and heterocyclic The reaction can be described as : CIP- + P* + Cl- and probably occurs at a faster rate than 4 x lo5 s-l since no absorption similar to that attributed to the radical anion shown in reaction (8) could be detected. A reaction analogous to reaction (8) does not occur for chlorpromazine presumably due to this short lifetime of ClP-*.The promazine radical, P-, probably has little absorption in the range 350-1000 nm and is expected to add on to chlorpromazine itself to produce a cyclohexadienyl type radical. This radical probably accounts for the absorption at 440nm formed 2 0 p after the pulse. The dependence of this absorption on chlorpromazine concentration suggests that either an alternative pathway for the decay of the promazine radical exists at low chlorpromazine concentration or that the promazine radical and the cyclohexadienyl radical are in equilibrium. Since the only other feasible reaction of the promazine radical, P-, would be protonation to yield the cation radical, the equilibrium suggestion is considered more likely.(10) This work was supported by grants from the Cancer Research Campaign and the Medical Research Council. The authors thank Dr. A. J. Swallow for helpful discussions. S. Navaratnam, B. J. Parsons, G. 0. Phillips and A. K. Davies, J.C.S. Faraday I, 1978,74,1181. S . A. Alkaitis, G. Beck and M. Gratzel, J. Amer. Chem. SOC., 1975, 97, 5723. L. W. C. Massey, Canad. Med. Assoc. J., 1965, 92, 186. J. A. Bassara, J. C. Newton and J. C. Saunders, J. Amer. Med. ASSOC., 1965, 193, 10. G. Porter and M. D. Archer, Interdisciplinary Sci. Rev., 1976, 1, 119. P. O’Niell, S. Steenken and D. Schulte-Frohlinde, J. Phys. Chem., 1975, 79, 2773. ’ K. M. Bansal, L. K. Patterson and R. H. Schuler, J. Phys. Chem., 1972, 76,2386. L. K. Patterson and K. M. Bansal, J. Phys. Chem., 1972,76,2392. R. Koster and K.-D. Asmus, J. Phys. Chem., 1973,77,749. lo J. P. Keene, Quaderni dell’ Area di Ricerce dell’ Emilia-Romugna, 1972, 1, 49. J. P. Keene, Quaderni dell’ Area di Ricerca dell’ Emilia-Romagna, 1972, 1, 63. l2 G. E. Adam, J. W. Boag, J. Currant and B. D. Michael, Pulse Radiolysis, ed. M. Ebert, J. P. Keene, A. J. Swallow and J. H. Baxendale (Academic Press, London, 1965), p. 117. l 3 D. C. Borg and G. C. Cotzias, Proc. Nat. Acad. Sci., 1962,48,617. 14B. Cercek, M. Ebert, C. W. Gilbert and A. J. Swallow, in Pulse Radiolysis, ed. M. Ebert, l5 M. BonifgciC, H. Mockel, D. Bahnemann and K.-D. Asmus, J.C.S. Perkin I l , 1975, 675. l6 K.-D. Asmus, D. Bahnemann, M. BonifiiciC and H. A. Gillis, Favaday Disc. Chem. SOC., 1977, 63, 213. A. J. Swallow, in Free Radical Reactions, Organic Series 1, ed. D. H. Hey and W. A. Waters (Butterworth, MTP International Science), vol. 10. J. P. Keene, A. J. Swallow and J. H. Baxendale (Academic Press, London, 1965). l8 B. Cercek and M. Kongshaug, J. Phys. Chem., 1970,4,4319. l9 J. E. Huheey, Inorganic Chemistry (Harper and Row, London), p. 161 and 699. 2o A. Habersbergerova, I. Janovsky and P. Kourim, Radiation Res. Rev., 1972,4, 123. 21 A. MacLachlan and R. L. McCarthy, J. Amer. Chem. SOC., 1962,84,2519. (PAPER 8/085)

ISSN:0300-9599    

DOI:10.1039/F19797500022

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Scanning Studies on Capillary Condensation and Evaporation of Nitrogen Part 1.-Apparatus and Calculation Method BY JOHAN c. P. BROEKHOF" AND WIM P. VAN BEEK Unilever Research, Vlaardingen, The Netherlands Received 17th February, 1978 For collection of extensive sets of data on scanning behaviour within nitrogen sorption hysteresis- loops at 77.6 K, an apparatus has been described by which a preprogrammed sequence of adsorption- desorption cycles within the hysteresis region can be recorded fully automatically. This has enabled the determination of up to 3000 datum points distributed over some 60 primary scanning curves, in a single run, without interruption of the measurement. Subsequently, a method has been described for obtaining quantitative information on capillary evaporation and condensation phenomena from scanning data.It has been shown that the slope and intercept of the tangent lines to a set of primary scanning curves at the point of pressure reversal can be directly related to the actual porous structure. This leads to a new independent method for determining pore size distributions. The common method for assessing the pore size distribution in the mesopore range is based on the interpretation of the adsorption-desorption hysteresis loop for physically adsorbed vapours in terms of capillary condensation. In practice, nitrogen at its normal boiling point is commonly used as adsorbate, The method is based upon a number of (largely unproven) assumptions, such as the validity in the mesopore range of the Kelvin relation between meniscus curvature and vapour pressure of the capillary-condensed phase, the predominance of one simple geometrical pore shape (e.g.either regular tubular pores or equidistant slits) and the neglect of effects related to the occurrence of networks of interconnected pores. The validity of these and similar assumptions is of great interest for actual porous systems and commonly used adsorbates. The potential use of scanning experiments within the hysteresis region to gain a better insight into the fundamentals of physical adsorption in the hysteresis region has been pointed out by a number of author~.l-~ Up to now the most extensive set of experimental data has been presented by Everett and c o ~ o r k e r s , ~ who have also provided a detailed formal framework for inter- pretation and correlation of scanning behaviour in the independent domain t h e ~ r y .~ - ~ In spite of the practical importance the number of scanning studies with nitrogen as adsorbate has been very small. In general, scanning studies are tedious and time- consuming and require very accurate measurements. Some years ago, a fully automated volumetric nitrogen adsorption apparatus was developed in our laboratory by Osinga, Wildschut and van Duyn to meet the need for a large number of nitrogen isotherms. The complete adsorption-desorption sequence was directly controlled through on-line coupling to an IBM-1800 process computer. After optimization of accuracy and reliability, this apparatus has been adapted for the automatic recording of a preprogrammed series of scanning curves throughout the hysteresis-region.A significant extension of the application of scanning curves can be made by combining the independent domain theory with the t-plot method of a n a l y s i ~ , ~ ~ lo 36J . C . P . BROEKHOFF AND W. P . VAN BEEK 37 in order to obtain estimates of the cumulative surface area and volume of the pore domains emptied or filled in each particular state of the system.ll In Part 1 of this series of two papers, we will discuss the essentials of the apparatus and the computational procedure for a t-plot analysis of the scanning curves. In Part 2 we will discuss information obtained from this analysis on the breakdown of the capillary-condensed state below a critical relative pressure along the desorption branch and on the mechanism of filling narrow mesopores along the adsorption branch.EXPERIMENTAL DESCRIPTION OF APPARATUS A constant-volume gas reservoir (V) at a fixed temperature was connected through a two-way dosing pump to a small sample vessel ( S ) immersed in a liquid nitrogen bath (fig. 1). The pressure in the gas reservoir and that in the sample vessel (pg) were recorded with Texas Instruments digitalised Bourdon gauges (resolution < 10 Pa). A third Bourdon gauge recorded the vapour pressure of pure liquid nitrogen ( p o ) at the temperature of the cryostat bath. The signals of the three gauges were converted into binary numbers through a pulse- counting technique in the read-out unit.I2 The dosing pump and read-out unit were operated by the command unit which is on-line to an IBM 1800 computer.The piston dosing pump of adjustable stroke-length was fitted with two magnetically ,operated vacuum-tight valves. Its dead space in the neutral position is small whereas the dead space is thermostatted by water flowing through a channel in the metal pump housing. Gas burette Equilibrium Reference pressure pressure - - pressure reading reading reading \ / Dosing pimp SC. C T " I- L 323 K 78 K Thermostat Liquid ndrogen 298 K bath I / IBM le00 / FIG. 1 .-Block-diagram of automatic volumetric sorption apparatus suitable for recording scanning curves. S = sample holder ; T = liquid-nitrogen thermometer ; V = constant-volume nitrogen burette, The level in the liquid nitrogen bath was controlled by a capacitative buoyancy level sensor which operates a magnetic valve between the bath and a large storage dewar vessel.However, a periodic fluctuation in the liquid nitrogen level could not be avoided. Therefore, the connection between the sample vessel and the rest of the system was water-jacketed and this jacket was surrounded by a vacuum chamber reaching into the liquid nitrogen bath. Variations in the effective dead space in the sample holder with the liquid nitrogen level were hereby largely suppressed. The dead space of the rest of the system was minimized; moreover the temperature of the working volumes of the Bourdon gauges was kept constant. The liquid nitrogen bath was open to the atmosphere, in such a way that the possibility of condensing atmospheric oxygen was minimized, even during long runs.The value of p o was still influenced by variations in the barometric pressure (as well as by the purity of the liquid nitrogen), but almost always fell within the range 102-106 kPa (= 765-795 mmHg),38 CAPILLARY CONDENSATION BY SCANNING which corresponds to 77.4-77.8 K. This temperature range was found acceptable for the present purpose, since no systematic influence of the temperature on the relation between amount adsorbed and pg/po in this range was found. STANDARD PROCEDURE First, the gas reservoir was carefully calibrated with nitrogen from a burette. For the determination of the gas volume in the sample holder space a complete blank isotherm was run under the measuring conditions, but without sample. In the actual measurement, the amount adsorbed was calculated from the pressure decrease in the gas reservoir as compared with that obtained in the blank run.An additional correction was made for the material volume of the solid, in the way described by Lippens.13 The optimum amount of sample to be used is dependent on its specific surface area, S (x 70 mg for 100 < (S/m2 g-') < 300). As many materials of interest are sensitive to heat, the standard sample pretreatment chosen was drying in air at 120 "C for 24 h, followed by degassing in sit& at 100 "C until a vacuum of w Pa had been reached. This left oxidic surfaces covered with one mono- layer of strongly bound water.14 Under these conditions non-porous oxides conform to the t-curve of nitrogen adsorption of De Boer et aZ.15 At the start of an experiment the pressure in the gas reservoir was read by the computer and stored.Then the dosing pump was activated uia the command unit and some nitrogen dosed to the evacuated sample. The pressure in the sample vessel was read by the computer at regular time intervals and compared with its preceding value. Equilibrium was regarded to be established if at least 50 pairs of readings are identical. This was found to be satisfactory if a time interval of 2 s in between readings was allowed. At equilibrium the pressure readings were stored for future processing and the '' distance " from this experi- mental point on the isotherm to its predecessor was calculated in terms of the amount adsorbed against the relative nitrogen pressure (pg/po).The dosing pump was then activated to transfer a calculated number of nitrogen doses in order to ensure a smooth spacing of points along the isotherm. During routine determination of the isotherm the dosing was automatically reversed at saturation adsorption, and the desorption branch of the isotherm was traced down top,/po = 0.25. SCANNING PROCEDURE The IBM 1800 process computer was preprogrammed to reverse the direction of the sorption sequence at certain specified relative pressures. In this way ascending as well as descending scanning curves could be recorded, or any desired sorption pathway within the hysteresis region. At present, up to 200 reversal points can be assigned to a single run. Upon approaching a reversal point, the dosing rate was automatically diminished in order to minimize " overshoot " and to assure an accurate recording of the start of the scanning curve.The number of experimental points to be recorded in a single run was not restricted: when the assigned memory storage capacity of the computer had been surpassed, its contents were automatically dumped on punched cards for further processing. In practice long runs were usually terminated only for maintainance of the mechanically moving parts of the system or in the case of leakage in stopcocks or greased joints. The actual number of experimental points in a single uninterrupted run has now exceeded some 3000, which corresponds to a few weeks' continuous operation. RESULTS t-PLOT ANALYSIS OF PRIMARY SCANNING CURVES : THE v-s CURVE The principle of t-plot analysis of scanning curves is based upon a theorem of the independent domain theoryY4 which states that directly after any direction reversal along a scanning pathway, sorption procedures correspond to reversible changes inJ .C. P. BROEKHOFF AND W . P. VAN BEEK 39 the state of the system. Such reversible changes in the state of the system will be largely due to variations with relative pressure of the thickness t of the adsorbed layers at the walls of the part of the porous system which was not filled with capillary condensate at the point of direction reversal, although reversible capillary pore filling cannot be excluded. True reversible capillary pore filling in an actual mesopore system is a rare phenomenon : no case of capillary condensation without hysteresis is known to us.For capillary condensation to be reversible, pores should possess closed ends and the pore geometry should be such that filling can take place without an increase in curvature of the adsorbate/vapour interface at any stage of the filling process. Reversible capillary pore filling should be discernible as an upward deviation from linearity in the t-plot of the initial part of the scanning curve following a direction reversal. In the absence of reversible capillary pore filling, the initial slope of a t-plot of any scanning curve in the vicinity of a direction reversal point will be proportional to the surface area of the walls of the pore domains not filled with capillary condensate at the direction reversal point. This proportionality takes a particularly simple form in cases where the radius of the wall curvature of the domains is large in comparison with the thickness of the adsorbed layer.For a scanning curve i, we may then write : where V, = the total volume of the adsorbate at a certain point along a scanning route directly after a pressure reversal, Vpi = the volume of all domains completely filled (and thus not accessible to changes in adsorbed volume), Sci = the surface area of the walls of all " empty " domains, and t = the statistical thickness of the adsorbed layer at the corresponding p g / ~ o (assumed to be independent of the domain geometry), in which case it can be derived from a suitable standard t-curve for the class of materials under investigati0n.l * Application of eqn (1) requires the assumptions that the density of the adsorbed layer equals that of the bulk liquid at the same temperature, and that changes in free liquid-vapour meniscus area (e.g.at pore mouths) between neighbouring states are negligible. The importance of eqn (1) lies in the possibility of characterizing quantitatively changes in the state of the domain system between two neighbouring reversal points in terms of pore volume and surface area. In particular the VPi-Sci relation for subsequent reversal points along either the adsorption or the desorption boundary curve (denoted here by the V-S curve), yields direct information about the " pore size " ri of domains filling or emptying along the boundary curves : For ideal tubular pores ri would equal Rp/2, for equidistant slits Op/2 and for spheroidal cavities Rp/3.In principle this will enable us to assess directly the characteristic dimensions of domains in pore space without recourse to any relation of the Kelvin-type 17* l8 and thus to make a qualitative and quantitative check upon the validity of the capillary condensation-evaporation model for hysteresis of physical adsorption in mesopores. More extensive sets of data taken in different parts of the hysteresis region will lead to an insight into the validity of the assumption of indepen- dent behaviour of domains in pore space, as in that case all Vp,-Sci points should lie on a single V-S curve which contains the essential information about the size distri- bution in the porous system. As stated, applicability of eqn (1) requires a negligible contribution of reversible capillary pore filling as well as negligible influence of pore wall curvature upon the40 CAPILLARY CONDENSATION BY SCANNING thickness of the adsorbed layer.If the t-plot of the adsorption boundary curve of the hysteresis loop is predominantly linear or shows a downward inflexion for increasing t-values, then it is reasonable to assume for a mesopore system that both conditions are well satisfied : such behaviour is generally associated with open slit-shaped pores. Even so, it must be stressed that eqn (1) can only be expected to hold for states in the direct vicinity of a direction reversal point. At larger separations from a direction reversal point, irreversible pore filling (ascending scanning curves) or emptying (descending scanning curves) is bound to occur.It would in principle be possible to determine experimentally the extent of the reversible part of a scanning curve after a direction reversal point by measuring, from each point of the scanning curve, a secondary scanning curve in the direction of the reversal point. In a few particular cases such measurements were made, and it was then found that in those cases the initial part of primary scanning curves was indeed reversible. In view of the large number of primary scanning curves required for the present analysis, such a procedure did not prove feasible as standard practice. Therefore, we were faced with the necessity of selecting the maximum number of consecutive experimental points on a scanning curve after direction reversal which is compatible with eqn (l), with due regard for the unavoidable experimental error in each of the datum points.The following procedure was adopted. A set of three or (preferably) more consecutive datum points, including the reversal point on the boundary curve of the hysteresis loop, was fitted to eqn (1) by means of linear regression. The regression coefficient and the 90 % confidence intervals for slope and intercepts were calculated according to ref. (19). The size of the ultimate set chosen was such that the correlation coefficient was maximal and the confidence intervals were minimal. There was usually a well-defined number of consecutive datum points that could be included in the set in order to satisfy these criteria.Inclusion of more datum points then led to larger confidence intervals and less satisfactory correlations, due to systematic deviations from linearity at larger distances from the reversal point at the boundary curve. This choice of the optimum number of datum points to be included in the regression procedure was checked by visual inspection for linearity of the initial part of the t-plot selected for the regression procedure from the complete t-plot of the scanning curve. The optimum set usually consisted of some five to ten experimental points. Application of this t-plot method to the analysis of scanning data, will be discussed elsewhere.20 We thank Mr. J. van Duijn and Mr. J. Siebesma for solving the problems in the electronics of the process control, Mr.B. H. van Wijngaarden and Mr. J. Bohm for programming the process control and Mr. E. Berends, Mr. J. H. Beun and Mr. A- Lukas for the development and construction of parts of the apparatus. K. S. Rao, J. Phys. Chem., 1941,45, 500, 506,513,517. P. H. Emmett and M. Cines, J. Phys. Colloid Chem., 1947, 51, 1260. H. W. Quinn and R. McIntosh, Cunud. J. Chem., 1967,35,745. D. H. Everett, in The Solid-Gas Interface, ed. E. A. Flood (New York, 1967), vol. 11, p. 1055. D. H. Everett and F. S. Smith, Trans. Furuday Soc., 1954,50,187. D. H. Everett, Trans. Furaday SOC., 1954, 50, 1077. ' E. Ekedahl and I. G. Sillkn, Actu Chem. Scund., 1965, 19,2323. * Th. J. Osinga, J. Wildschut and J. van Duyn, to be published. B. C. Lippens and J. H. de Boer, J. Catalysis, 1965,4, 319. lo K. S. W. Sing, Chem. and Ind, 1968, 1520. l1 J. C. P. Broekhoff, L. F. Brown and W. P. van Beek, Proc. Int. Symp. on Pore Structure and Properties of Materials, ed. S . Modry (Prague, 1973), vol IVY C-85.J . C . P . BROEKHOFF AND W. P . VAN BEEK 41 l2 J. van Duyn, Electronics, Dec. 21, 1970, p. 55. l3 B. C. Lippens, 23esis (University of Technology, Delft, 1961). l4 J. H. de Boer, J. M. H. Fortuin, B. C. Lippens and W. H. Meys, J. Catalysis, 1963, 2, 1. l5 J. H. de Boer, B. G. Linsen and Th. J. Osinga, J. Catalysis, 1965,4,643. l6 A. Lecloux, J. Colloid Interface Sci., to be published. l7 S. J. Gregg and K. S. W. Sing, in Adsorption, Surface Area and Porosity (Academic Press, London, 1967), chap. 3. J. C. P. Broekhoff, Thesis (University of Technology, Delft, 1969), chap. 3 and 4. I9 F. S. Acton, in Analysis of Straight-line Data (Dover, New York, 1959), pp. 19-26 and 119. 2o J. C. P. Broekhoff and W. P. van Beek, J.C.S. Faruday I, 1979,75,42. (PAPER 8/283)

ISSN:0300-9599    

DOI:10.1039/F19797500036

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Scanning Studies on Capillary Condensation and Evaporation of Nitrogen Part 2.-Analysis of Ascending and Descending Scanning Curves within B-Type Hysteresis Loops BY JOHAN C . P. BROEKHOFF* AND WIM P. VAN BEEK Unilever Research, Vlaardingen, The Netherlands Received 17th February, 1978 The mechanism of emptying and filling of pore domains along the adsorption as well as the desorption boundary curve of the hysteresis loop was studied by the determination of a V-S curve derived from the tangents to primary scanning curves immediately after a pressure reversal at the boundaries of the hysteresis region. In all three cases studied, the V-S curve for the desorption process exhibits a discontinuity aroundp,/po = 0.5, pointing to a sudden breakdown of the capillary- condensed state, which is not directly related to the detailed domain properties of the system.This breakdown of the capillary-condensed state invalidates in a number of cases the usual procedures for obtaining pore size distributions. The filling of slit-shaped pores was found to proceed most probably by additional sorption of one to two nitrogen molecules between adjacent adsorbed layers. Hysteresis of vapour adsorption in rigid porous systems is not commonly observed below a certain typical value of the relative vapour pressure. The magnitude of this " critical " relative pressure is apparently related to the physical properties of the adsorbate rather than to the actual size distribution of the porous system. Several explanations for this curious phenomenon have been put forward. Cohan has assumed that capillary condensation in tubular pores occurs by Kelvin instability of the adsorbed layer inside the pores, while capillary evaporation results from Kelvin instability of the meniscus at the pore entrance. If it is assumed that the radius of curvature of the surface of the adsorbed layer during adsorption is equal to the pore radius minus the adsorbed layer thickness, while during desorption the radius of the curvature of the meniscus equals that of the pore, then it follows that capillary sorption becomes reversible in pores with radii smaller than about twice the thickness of the adsorbed layer.Neglect of the presence of an adsorbate layer remaining at the pore walls after capillary evaporation renders this line of reasoning invalid.Foster has included the influence of adsorption forces upon the Kelvin instability of the adsorbed film during adsorption on the basis of arguments similar to those of Cohan and thus leading to qualitatively the same conclusion. The neglect of both the remaining adsorbed layer after capillary evaporation and the influence of adsorp- tion forces upon the capillary evaporation p~ocess,~ renders this argument equally invalid [for a fuller discussion, see ref. (5)]. More relevant is the remark of Everett and Haynes that capillary condensation is only expected to occur after the adsorbed film has reached a sufficient thickness for the development of liquid-like properties, specifically of a liquid-vapour interfacial tension. This would explain the absence of capillary condensation at the lower end of the relative-pressure range.On the other hand, it has been pointed out by several authors 7* that the curvature of the liquid meniscus at the pore entrances increases 42J. C. P . BROEKHOFF AND W. P. VAN BEEK 43 along the desorption branch of the isotherm in the direction of decreasing relative pressures. This increasing curvature is accompanied by an increase in magnitude of the negative hydrostatic pressure (tension) acting on the capillary condensed phase. At a certain point, this tension may well surpass the tensile strength of the capillary- held liquid, resulting in spontaneous nucleation and subsequent evaporation. This point has been elaborated by Everett 9s10 and co-workers and by Dubinin 11*12 and co-workers who presented evidence for a relation between the estimated tensile strengths of different adsorbates and the corresponding lower closing points of the hysteresis loops.It is commonly observed for systems exhibiting B-type hysteresis loops according to the classification of De Boer l3 that the desorption branch becomes very steep at or near the “ critical ” relative pressure. This is rather suggestive of an actual forced breakdown of the capillary-held phase, as it is improbable that nature should have equipped us with an abundance of porous systems all exhibiting a peak in their distribution around 2.5 nm, while the corresponding narrower pore range would be absent. We have recently demonstrated l4 that the behaviour of ascending and descending scanning curves in a system exhibiting B-type hysteresis, violates Everett’s consistency rule for ascending and descending scanning curves,l 9 while in other cases these rules were found to be closely followed.Thus, although in the former case all primary ascending curves starting from the desorption boundary curve converge towards the upper closing point of the hysteresis region, the descending scanning curves emanating from the adsorption boundary curve intersect the steep part of the desorption branch of the isotherm at an acute angle, instead of converging towards the lower closing point of the hysteresis loop. This unusual behaviour seems to suggest that this type of scanning behaviour at least in part is not determined by domain properties, but rather by some other phenomenon occurring in the capillary- condensed phase.Apparently, the run of the descending scanning curves is disrupted discontinuously around the “ critical ” relative pressure by the onset of an intrinsic instability of the capillary-held phase of the type predicted by S~hofield,~ Flood and EverettY9 rather than by the mechanism envisaged by Cohan and F ~ s t e r . ~ In the present paper, the application of the method we have published previously l7 to an analysis of the phenomena occurring around the “ critical ” relative pressure is discussed. At the same time the mechanism of filling of pores during adsorption will be studied. EXPERIMENTAL MATERIALS For this particular study, three samples were chosen from a large collection of nickel- on-silica catalysts, all three exhibiting pronounced B-type hysteresis with a very steep part in the desorption branch around pg/po = 0.5 (for nitrogen as the adsorbate, sorption hysteresis in a rigid porous sytem is hardly ever observed below pJp0 = 0.42).TABLE 1 .-PHYSICAL CHARACTERISTICS OF CATALYSTS sample surface area pore volume porosity mean pore size b mean particle size 0 code /lo3 m2 kg-1 /lO-3 m3 kg-1 /% /iO-9 m /10-9 m NS 25a 113.5 0.219 37 3.9 6.5 NP 113 200.8 0.63 1 68 6.3 3.0 NZ269 263.5 0.444 54 3.4 2.9 a Differences from the data cited for this preparation in ref. (14) are apparently to be attributed Defined here as 2 x pore volume/surface area. to heterogeneity within the bulk of the powder. C Defined here as 2/(solid density x surface area).44 CAPILLARY CONDENSATION BY SCANNING A summary of the relevant physical properties of these samples is presented in table I, NS 25 is a hydrothermally synthesized nickel silicate, apparently consisting of thin, two- dimensionally extended sheets.An electron microphotograph is presented in fig. l(a). NP 113 is a nickel silicate catalyst produced by direct precipitation of nickel hydroxide upon a support under alkaline conditions. The sheets of nickel silicate are less robust and far less extended than in the case of NS 25 [fig. l(b)]. The higher surface area indicates a finer dispersion. NZ 269 has a structure of a different morphology and was obtained by a homogeneous precipitation procedure. Small platelets may be observed [fig. l(c)], with certain fibrous elements. The high surface area indicates an even finer dispersion.RESULTS AND DISCUSSION Prior to the determination of the scanning curves, the boundary curves of the hysteresis region were traced. Sometimes a certain permanent hysteresis occurred in the low relative pressure region after the first adsorption-desorption cycle. Upon repeated cycling, the hysteresis loop usually settled down to a well-reproducible shape and position, provided the adsorption branch was always taken up to the same amount adsorbed at saturation prior to desorption. For the isotherms presented here, the adsorption branch had shifted upwards by some 5 % around ps/po = 0.5 after the first adsorption-desorption cycle, for both NS 25 and NZ 269, whereas for NP 113 the shift was negligible. The cause of this curious initial drift could not be ascertained.As there was no apparent reason for suspecting an instrumental artefact, the possibility of irreversible swelling in the stacking of nickel silicate sheets during the first sorption cycle cannot be excluded. Primary descending scanning curves were pursued from their origin on the adsorption boundary curve down to beyond the lower closing point of the hysteresis loop. The return curve then provides a check upon the reproducibility of the adsorption branch during the scanning process. This was found to be within experimental accuracy in most cases. Ascending scanning curves were not pursued all the way up to saturation, but traced up to a relative pressure exceeding 0.9. In previous studies it had been found that the exact position of the desorption boundary curve of the hysteresis loop is sometimes dependent on the actual value of the volume taken up at saturation (this points to the presence of very large pore domains which are only accessible through relatively narrow entrances).In the present study, process control is based on relative pressure, which makes it difficult to reproduce the exact uptake at saturation. For our purpose, it is sufficient to determine the initial slope of the ascending scanning curves. After direction reversal these scanning curves were found to return exactly to their origin at the desorption branch of the isotherm, although a secondary hysteresis loop with the primary ascending curve pointed to some domain filling at the high relative pressure end.It has not been possible to record both primary ascending and primary descending scanning curves in a single run, due to the necessity of mechanical maintenance or to the occurrence of an apparatus breakd0~n.l~ In general, care was taken to use the same individual sample for different runs, but sometimes the sample was lost during handling operations, and another sample of the same batch of material had to be used. NS 25 Ascending and descending scanning curves are presented in fig. 2(a) and (b). The corresponding t plot transforms are presented in fig. 2(c) and (d). The part of the primary ascending scanning curves in the direct vicinity of a direction reversal point could be successfully fitted to eqn (1) of Part l . 1 7 Typically, five to tenh 9, [To face page 44(4 FIG.1.-Transmission electron microphotos of catalysts used in the scanning study : (a) NS 25 ; (b) NP 113 ; (c) NZ 269.J . C . P . BROEKHOFF AND W. P . VAN BEEK consecutive experimental points, including the direction reversal point, were found to satisfy this relation. The relative pressure range for which eqn (I) was found to hold varied from 0.07 for scanning curves situated near the closing points of the hysteresis loop, to 0.25 for scanning curves traversing the full width of the hysteresis region. Correlation coefficients for the selected set of points for all of the scanning curves of NS 25 were between 0.99 and 1.0, lending support to the adopted analytical procedure. This satisfactory behaviour was not unexpected, as the t plot of the adsorption boundary curve does not show any upward deviation, suggesting that both appreciable reversible capillary pore filling and appreciable curvature of the pore walls are absent.It is remarkable in fig. 2(c) that, although the initial part of the t plot could satisfactorily be fitted to a straight line, a distinct upward curvature can often be discerned before the scanning curve approaches the t plot transform of the adsorption 0.2t I 1 s ~ " " ' * ' ' 0 02 0 4 06 0.8 1 .o (4 PgIPo 1.0r (6) :L PgIPo 0.8 1 .o FIG. 246 5 r I CAPILLARY CONDENSATION BY SCANNING &/lo3 m2 kg-' FIG. 2.-contd.J . C . P. BROEKHOFF AND W . P. VAN BEEK 1.0 - 47 0 0.2 0.4 0.6 Q8 1.0 0 PglPo FIG. 2.-Scanning behaviour on NS 25: (a) primary ascending scanning curves; (6) primary descending scanning curves ; (c) t plots of primary ascending curves ; (d) t plots of primary descendlng curves ; (e) characteristic V-S curves from ascending and descending scanning curves ; (f) sizes of domains emptying or filling along the boundary curves of the hysteresis region as a function of relative pressure. As a reference the thickness of the adsorbed layer on a free surface is also presented.boundary curve. Such a curvature is usually taken as an indication of capillary condensation processes, which are unexpected for systems where the straight-line behaviour of the adsorption boundary curve is commonly interpreted as the absence of pore wall curvature and capillary condensation phenomena. Fig. 2(d) demon- strates rather dramatically the acute intersection between the descending scanning curves and the steep part of the desorption boundary curve.The descending scanning curves near the lower closing point of the hysteresis loop are very limited in extent. In the others a distinct downward curvature can be discerned, pointing to a certain degree of pore emptying due to capillary evaporation at a decrease in relative pressure. Massive capillary evaporation does not set in before a certain, apparently critical, relative pressure has been reached. The relation between Vp and S, derived from the respective intercepts and slopes of the tangents to the descending and the ascending scanning curves (the V-S curve) is given in fig. 2(e). The size of the crosses indicates the 90 % confidence intervals resulting from independent estimations of slope and intercept.l Quite remarkable is the discontinuity in the V-S curve obtained from the ascending scanning curves.No corresponding discontinuity is found in the V-S curve derived from descending scanning curves. This distinction in behaviour is even more dramatically demon- strated by a plot of the slope of the V-S curves according to eqn (2) of ref. (17) against the relative pressure of the corresponding origins of each primary scanning curve situated at the boundary curves of the hysteresis region [fig. 2(f)]. Around a pe/po value of 0.495 (the relative pressure at the start of the very steep part in the desorption branch) very large formal domain diameters are derived from the V-S curve. A further decrease in relative pressure results in the emptying of domains48 CAPILLARY CONDENSATION BY SCANNING with progressively smaller dimensions, until finally around pg/po = 0.42 the size of the domains emptying just equals the thickness t of the adsorbed layer at that relative pressure. Apparently, over a quite restricted range of relative pressures, emptying occurs of all domains which were previously filled with a capillary-condensed phase, apart from the permanent presence of an adsorbed layer.This seems to happen to a first approximation independently of the size of the domain, but it is clear that the largest-size domains empty first, followed by increasingly smaller domains as the relative pressure is lowered towards the final closing point of the hysteresis loop.This behaviour is highly suggestive of a mechanism for hysteresis breakdown caused by surpassing of the tensile strength of the capillary-condensed phase below a certain critical relative pressure, as predicted by Schofield,’ Flood * and Everett. According to simple thermodynamics,s 9 the formal hydrodynamic pressure in the capillary condensed phase at a distance t from the nearest pore wall can be written as : (1) PL = 1/ VLrPg VL - RT In (PO/PA + Wl where VL is the liquid molar volume and F(t) is a correction term for the influence o adsorption forces emanating from the pore walls, for the sake of simplicity taken here to be independent of the pore geometrya4* According to this equation pL equals - 12.9 MPa (- 127 atm) at pg/po = 0.5 in very wide pores and -12.5 MPa (-123 atm) at ps/po = 0.4 in pores with a radius of 1.0 nm, if F ( t ) is approximated by the relation F(t) = 2.3 RT (0.14/t2-0.034), a fairly accurate representation of the De Boer-Lippens t curve.4* Thus, the stabilising effect of adsorption forces may at least qualitatively account for the sequential breakdown of the capillary condensed state in increasingly narrower pores with decreasing relative pressures.In narrower pores, disturbances in adsorbate packing may also lead to changes in VL, as well as in tensile strength, so we may not expect more than a qualitative agreement with eqn (1). The presence of very wide domains which empty around pg/po = 0.5, in itself is no more than an indication that certain wide domains exist, which are only connected to their environment through narrower-sized pathways (e.g.as in the classical bottle- neck theory of hysteresis) and thus are unable to empty at relative pressures corres- ponding to their proper size. Rather the fact that the emptying of these wide domains is closely followed by emptying of all other ones, even those which on the basis of a Kelvin-relation would be expected to empty at significantly lower relative pressures, is the most powerful indication of the occurrence of a sudden breakdown of the capillary condensed state as such. In this respect, it is particularly significant that below pg/po = 0.5 the slope of the characteristic curve in fig. 2(f) does not return to the extrapolation of the first branch of the curve, but definitely intersects it. The filling of pore domains along the adsorption boundary curve of the hysteresis loop can only be followed with the present technique in the region above the closing point of the hysteresis loop.Thus, on the basis of the foregoing discussion, we should not expect any discontinuity in the corresponding characteristic curve, as indeed is found. Remarkably, the relation between slope of the characteristic curve and the relative pressure of domain filling [fig. 2(f)] closely follows the t curve for the thickness of the adsorbed layer as a function of relative pressure, but is shifted upward over a distance of 0.2 to 0.4 nm. The basis of the t and MP method for assessing the size of micropores, is formed by the implicit assumption that pores should fill if the thickness of the adsorbed layer equals half the pore diameter for slit-shaped pores.In that case, the slope of the characteristic curve at all relative pressures wouId coincide with the t curve. The most simple interpretation of the present results is, that at pressures around pn/po = 0.5 pore filling occurs as soon asJ. C. P. BROEKHOFF AND W. P. VAN BEEK 49 one nitrogen molecule fits between opposite adsorbed layers, whereas in wider pores, and thus at higher relative pressures, filling occurs if two nitrogen molecules can bridge the distance between the adsorbed layers. The results obtained here are particularly significant in this context as the electron micrograph suggests that we are indeed dealing with sheetlike structures, where the absence of pore wall curvature will prevent the premature occurrence of capillary condensation.Nevertheless, this proof of the occurrence of additional sorption between adjacent adsorbed layers cannot be regarded as definite, as the presence of tapered pore structures might lead to the same behavi0ur.l In any case, there is ample reason to review the quantitative basis of the t and MP method. Returning to the analysis of the V-S curve from the desorption branch of the hysteresis loop, it is furthermore remarkable that the variation of the slope of the characteristic curve with relative pressure in the region above pg/po = 0.5 is slight and definitely less than is predicted by relations of the Kelvin type. Thus the quantitative validity of the Kelvin equation or any of its sophistications 5 9 l6 for the present case, is not confirmed.Whether this is the rule or an exception, can only be ascertained on the basis of more extended experimental work. Qualitatively, domains empty sequentially in the direction of smaller diameters at lower relative pressures, as is predicted by the Kelvin theory of capillary evaporation. The interference of swelling and shrinking phenomena of the porous system during the adsorption-desorption cycle cannot be completely excluded. However, the perfect reproducibility of the final adsorption and, under certain restrictions, the desorption branches of the system, even after partially filling or emptying, and the straightness of the t plots of the ascending scanning curves over a fairly large range of relative pressures, is in our opinion an indication that extensive swelling does not occur.In a swelling system, the volume adsorbed in completely filled domains would increase with relative pressure according to a capillary condensation mechanism rather than remain constant or follow the t curve relation as found in the present study. NZ 269 AND NP 113 The ascending and descending scanning behaviour in the system NZ 269 [fig. 3(a)- (f)] essentially confirms the findings for NS 25. Above a relative pressure of 0.65, I 02 04 06 0.8 1 .o50 5- 4 - 3- 2 - 1- CAPILLARY CONDENSATION BY SCANNING I I I I 1 , , i I , I , , , ( # ID- 0.0 - - 0.6 - s c \ - 0.4 - 0 0.2 0.4 0.6 0.e 1.0 PglPo "I Ll I ' 1 1 I " L 0 0.2 0.4 0.6 0.8 1.0 1.2 ' 1.4 1.6 (d) t/= FIG. 3.-contd.J . C. P . BROEKHOFF AND W.P. VAN BEEK &/lo3 m2 kg-I , I 0.2 0.4 0.6 0.8 1.0 PgIPo 51 FIG. 3.-Scanning behaviour on NZ 269 : (a) primary ascending scanning curves ; (b) primary descending scanning curves; (c) t plots of primary ascending curves; (d) t plots of primary descending curves ; (e) characteristic V-S curves from ascending and descending scanning curves ; (f) sizes of domains emptying or filling along the boundary curves of the hysteresis region as a function of relative pressure. As a reference the thickness of the adsorbed layer on a free surface is also presented.52 CAPILLARY CONDENSATION BY SCANNING the t plot of the adsorption boundary curve shows a small but distinct upward curvature, suggesting that pore wall curvature may not be completely negligible in this case.Nevertheless, the initial part of all scanning curves could be fitted in eqn (1) of Part 1 with correlation coefficients of 0.99 or better. Also here, there is a clearly distinguishable discontinuity just below pg/p0 = 0.5. In this case, the breakdown occurs near the lower closing point of the hysteresis loop, corresponding to the second steep part of the desorption boundary curve. The onset of a first steep part in this desorption curve, which occurs aroundpg/po = 0.61, is not reflected as a discontinuity in the V-S curve, and thus is caused by the presence in this particular structure of a large number of pores with a certain size emptying by straightforward capillary desorption around pg/po = 0.61, and not by a breakdown of the capillary condensed phase.This seems to reinforce strongly the point of view presented in the previous section, that final breakdown is related to relative pressure and not to details in the domain structure of the system. As to the relation between the slope of the V-S curve derived from the descending curves and the relative pressure of their original state, it is remarkable that this line virtually coincides with the one obtained for NS 25, and thus seems to be insensitive to the details of the porous structure. This may be considered as a reinforcement of 0.e I I c 2 0.6 - 1 s - 0.4 - 0.2 - i 0 0.2 0.4 0.6 0.8 I .o (4 PglPo (b) t/nm FIG. 4J . C. P. BROEKHOFF AND W. P . VAN BEEK 53 II 4.0 - - z 0" - 3.0 - Fl c; s 20- a a U 1.0 - &/lo3 m2 kg-l 0.2 0.4 06 00 1.0 (4 PgIPo FIG.4.-Scanning behaviour on Np 113 : (a) ascending scanning curves ; (b) t plots of ascending curves ; (c) characteristic V-S curve from ascending curves ; (d) size of domains emptying along the desorption branch of the hysteresis loop.54 CAPILLARY CONDENSATION BY SCANNING the view that during pore filling, additional sorption between adjacent sorbed layers occurs rather than capillary condensation in a tapered slit-shaped porous structure. As to NP 113, we have succeeded in obtaining data of sufficient accuracy only for the behaviour of the ascending scanning curves. In the case of the descending scanning curves, difficulties were encountered in obtaining sufficient reproducibility of the adsorption branch upon repeated cycling, either due to experimental difficulties or to some intrinsic property of the system (swelling?).The data for the ascending scanning curves are presented in fig. 4(a)-(d), because the V-S curve shows the same discontinuity around ps/po = 0.5 as observed for the other two systems, and thus is a confirmation of the (probably general) trend, for the slit-shaped porous systems. CONCLUSIONS A V-S curve for the relation between pore size, pore volume and pore surface area for the domains either emptying along the desorption boundary curve or filling along the adsorption one, can be obtained from a careful estimation of the position equation of the tangent lines to either ascending or descending primary scanning curves immediately after a pressure reversal at the boundary curves of the hysteresis region.In practice, this is possible for systems with sufficiently small curvature of the domain walls, such as slit-shaped pore systems, by application of the well-known t method. In the three cases studied, the characteristic curve for the desorption boundary curve showed a very pronounced discontinuity around ps/po = 0.5, corresponding to the commonly observed steep part of the desorption boundary curve. Steep parts in the desorption branch at higher relative pressures do not lead to such discontinuity. The present work seems to confirm the notion that around a certain critical relative pressure complete breakdown of the capillary condensed state occurs due to surpassing of the tensile strength of the capillary condensed liquid. Evidence for the hypothesis was hitherto mostly of an indirect nature.l0-l2+ 22 Confirmation of the quantitative validity of the Kelvin mechanism for the desorption process from the V-S curves, could not be obtained and should be a subject for further study.The sudden breakdown of the capillary condensed state around aps/po value of 0.5 for nitrogen at 78 K, invalidates the usual procedures for obtaining pore size distri- butions from the desorption branch of the hysteresis loop foi all samples which exhibit hysteresis in this pg/po region. In practical terms, this means that the peak in the pore distribution often observed in the region of 2.5 nm sizes has to be considered as an artefact of the method. Along the adsorption branch of the hysteresis loop, the dimensions of the pores filling at any relative pressure exceed the statistical thickness of the adsorbed layer on their pore walls by one or two adsorbate molecular diameters.This is suggestive of a filling mechanism by way of adsorption of one or two additional adsorbate molecules in between adjacent adsorbed layers, rather than a simple merging of adjacent adsorbed layers at the appropriate statistical thickness. In principle, the determination of the V-S curve from the study of ascending and descending primary scanning curves, should enable assessment of the structure of a porous system as well as study of the mechanism of capillary filling and emptying of pores. A stimulating discussion with Prof. D. H. Everett, University of Bristol, on the properties of scanning curves within the hysteresis loop is gratefully acknowledged.J .C. P . BROEKHOFF AND W . P. VAN BEEK 55 LIST OF SYMBOLS correction term for the influence of the adsorption forces emanating from the pore walls equilibrium gas pressure formal hydrodynamic pressure in the capillary- condensed phase saturation vapour pressure gas constant cumulative surface area from t plots thickness of the adsorbed layer temperature adsorbed volume liquid molar volume total pore volume reference for scaling purposes adsorbed volume at pg/po 3 1 J mol-I Pa Pa Pa J mol-' K-I m2 kg-' nm K m3 kg-I m3 mol-' m3 kg-I m3 kg-I m3 kg-l L. H. Cohan, J. Amer. Chem. Soc., 1944,66,98. D. H. Everett and J. M. Haynes, J. Colloid Interface Sci., 1972, 38, 125. A. G. Foster, J. Chem. Soc., 1952, 1806. B. V. Deryagin, Acta Physicochim. U.R.S.S., 1940, 12,139. J. C. P. Broekhoff, 27zesis (University of Technology, Delft, 1969), chap. 1 and 3. D. H. Everett and J. M. Haynes, Colloid Science (Spec. Period. Rep., Chem. SOC., London, 1973), vol. 1, p. 136. E. A. Flood, The Solid-Gas Interface, ed. E. A. Flood (Marcel Dekker, New York, 1969, vol. 1, chap. 1. D. H. Everett, The Solid-Gas Interface, ed. E. A. Flood (Marcel Dekker, New York, 1967), vol. 2, p. 1055. ' R. K. Schofield, Disc. Fiaday SOC., 1948,3, 105. l o C. G. V. Burgess and D. H. Everett, J. Colloid Interface Sci., 1970, 33, 611. 'l M. M. Dubinin, Pure and Appl. Chem., 1965, 10, 309. l2 0. Kadlec and M. M. Dubinin, J. Colloid Interface Sci., 1968, 31,479. l3 J. H. De Boer, The Structure and Properties of Porous Materials, ed. D. H. Everett and F. S. l4 J. C. P. Broekhoff, L. F. Brown and W. P. van Beek, Proc. Int. Symp. on Pore Structure and l5 D. H. Everett, Trans. Faraday Soc., 1954, 50, 1077. l6 D. H. Everett and F. S. Smith, Trans. Faraday Soc., 1954,50, 187. '' J. C. P. Broekhoff and W. P. van Beek, J.C.S. Farahy I, 1979,75, 36. l 8 F. S. Acton, in Andysis of Straight-line Data (Dover, New York, 1959). Stone (Butterworth, London, 1958), p. 68. Properties of Materials, ed. S . Modry (Prague, 1973), vol. IV, C-85. S. J. Gregg and K. S. W. Sing, Adsorption, Porosity and Surface Area (Academic Press, London, 1967), chap. 4. 2o J. H. De Boer, B. G. Linsen, Th. van der Plas and G. J. Zondervan, J. Catalysis, 1965,4, 649. 21 R. G. Mikhail, S. Brunauer and E. E. Bodor, J. Colloid Interface Sci., 1969, 26, 45. 22 R. G. Avery and J. D. F. Ramsay, J. Colloid Interface Sci., 1973, 42, 597. (PAPER 8 /282)

ISSN:0300-9599    

DOI:10.1039/F19797500042

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Some Unusual Properties of Activated and Reduced AgNaA Zeolites BY PETER A. JACOBS* AND JAN B. UYTTERHOEVEN Centrum voor Oppervlaktescheikunde en Colloi'dale Scheikunde, Katholieke Universiteit Leuven, De Croylaan 42, B-3030 Leuven (Heverlee), Belgium AND HERMANNK. BEYER Central Research Institute of Chemistry, Hungarian Academy of Sciences, 11 Pustaszeri-6t 57-69, Budapest, Hungary Received 2nd March, 1978 Carbon monoxide, oxygen and hydrogen were found to be chemisorbed on dehydrated AgA zeolites. This was investigated in detail using volumetric sorption and temperature programmed desorption techniques. Also i.r. and mass spectrometry were used to characterize the solid and the desorbed molecules. It was found that as a result of an auto-reductive process, colour centres are created upon de- gassing of the zeolite.These centres sorb hydrogen and oxygen dissociatively, while one molecule of carbon monoxide was chemisorbed per Ag ion available in the supercage. It is proposed that linear Ag': clusters are formed upon activation, the ends of which constitute chemisorption sites for hydrogen and oxygen. Ralek et al.' reported the use of AgNaA zeolites as moisture indicators. This application was based on the observation that the brick red colour of thoroughly dehydrated AgA zeolite was very sensitive to traces of moisture. Upon addition of increasing amounts of water, the sample colour was found to change towards orange, yellow and white. At that time no explanation for the phenomenon was advanced. have rediscovered these colour changes upon activation.The authors show that upon activation silver atoms are formed. They claim that uncharged silver clusters (Ag, molecules) are formed within cubes of eight (or six) Ag+ ions. The process of auto-reduction of transition metal ions in zeolites is not unique for Ag+, since it previously has been reported for cupric ions in Y ~ e o l i t e . ~ The silver zeolite system still shows other peculiarities. Upon thorough degassing of reduced samples, it was shown that hydrogen to a limited extent can be formed by oxidative thermal desorption from zeolites Y and m~rdenite.~ This property was combined with the ease of photochemical reduction of hydrated silver ions in zeolite. In this way, the silver+zeolite system could be used for the cleavage of water into oxygen and hydrogen using a photochemical and thermochemical step.6 It also was recently found that upon addition of ammonia to AgA the new mole- cules triazane (N3H5) and cyclotriazane (N3H3) were formed in supercages of zeolite A.3 In the present work, we report on other unusual properties of silver exchanged A zeolite. The development of bright colours upon activation could be related to a process of auto-reduction.We also attempt to explain the unexpected chemi- sorptive properties for hydrogen, oxygen and carbon monoxide in such a system. In recent single crystal X-ray diffraction studies of the same system, Seff et aL2* 56P. A . JACOBS, J. B . UYTTERHOEVEN AND H . K. BEYER 51 EXPERIMENTAL MATERIALS A commercial zeolite Na-A from Union Carbide Corporation, Linde Division, was purified as described earlier for Silver A zeolites with different silver cation content were obtained after exchanges in 0.005 mol dm-3 AgN03 solution.Afterwards the samples were washed until complete disappearance of any anions in the washing waters and dried at ambient temperature in the dark. Only freshly prepared and white samples were used in further work. Sample notation and anhydrous unit cell composition are given in table 1 . TABLE AN ANHYDROUS UNIT CELL COMPOSITION OF A ZEOLITES For particular applications other compositions were prepared in the same way. A ZK-4 zeolite was synthesized according to known procedures 9s lo with the following unit cell composition : where TMA represents tetramethylammonium cations. The sample was calcined in flowing oxygen at 773 K, and saturated with ammonia before contacting it with water. The silver zeolite A was obtained after repeated exchanges in an excess of 0.005 mol dm-3 AgN03 solution with a ratio of zeolite to exchange solution of 0.5 gdm3.Chemical analysis showed a sample with the following anhydrous unit cell composition : The X-ray diffraction pattern showed that the samples in all cases remained highly crystalline The number following the sample name corresponds to the degree of silver exchange. In a few cases, the number in brackets following the degree of cation exchange, denotes the degassing temperature in U ~ C U O in K. Labelled oxygen (I8O2) with 99.9 % isotopic purity was from L.C.B. Na8 TMA (A102)9 (Si02)15 Ag9 (A102)!3 (Si02)15.PROCEDURES AND METHODS Gas uptake and desorption measurements were performed in a low volume circulation system. Using a sample of kg, the accuracy is better than 3-0.005 mmol. Tempera- ture programmed desorption (t.p.d.) measurements were carried out in a Hewlett Packard 5992A gas chromatographjmass spectrometer combination, using single ion monitoring techniques. The mass scale of the mass spectrometer unit was calibrated automatically. 1.r. measurements were taken in situ with a Beckman IR12 grating spectrometer in the double beam/absorption mode, using a sample and reference cell attached to a vacuum sys tem. Sorption of hydrogen was carried out at 195 K. The " chemisorbed hydrogen " cor- responds to that amount of hydrogen adsorbed at 195 K that cannot be desorbed after a 10 min degassing at 293 K.Carbon monoxide is considered to be chemisorbed, when after room temperature adsorption (293 K) it cannot be desorbed after degassing in vacuo (1.33 mN m-") for 1 h. RESULTS AND DISCUSSION ACTIVATION OF AgNaA ZEOLITE A typical t.p.d. experiment for hydrated AgA-100 is shown in fig. 1. At low temperatures, water is the main desorption product, but above 400 K non-negligible amounts of oxygen can also be desorbed from the sample. Using a volumetric58 PROPERTIES OF AgNaA ZEOLITES system with cold trap (195 K) the oxygen desorption curve can be perfectly reproduced. It should also be noted, that when the desorption of water has ceased completely (525 K), oxygen desorption restarts again. Fig. 2 shows clearly that the rate of oxygen desorption declines in the 500-520 K region and shows two definite regions characterized by high rates of oxygen loss.In the low temperature region (c 500 K) desorption temperature/K FIG. 1.-T.p.d. experiment on hydrated AgA-100 with the desorption curve for water (a) and oxygen (b), respectively. the sample turns from white to a deep yellow, golden colour. At this stage almost complete dehydration is reached and a first quantity of oxygen is released. Con- sequently, the sample turns to a bright red colour, releasing a slightly higher quantity of oxygen. The same solid state reactions established for the auto-reduction of CuY4 zeolites can be used here: As+ As+ 0 0 0 0 0 \-/\ A-N \/\-/\ / /\ /\ /\ /\ /\ /\ A /\ Al Si A1 Si -+ Ag"+Si A1 Si+ Al+O- (1) Ag+ 0 0 v \--/ v \-/ /\ A /\ /\ Si Al+O--+Ag"+O-+ Si A1 (2) 2 0 .4 0 2 . (3) (4) The overall stoichiometry of these reactions is where ZO- represents the zeolite lattice and Z+ a Lewis site, written in the conven- tional form. Nothing changes in the reaction stoichiometry when alternative re- action schemes for Lewis acid site formation are used. The use of this stoichio- metry allows to estimate the degree of auto-reduction. Fig. 2 shows that at 600 K, the degree of Ag+ reduction amounts to almost 8 %. After contacting dry and degassed AgA-100 with I6O2 or 1 8 0 2 at 673 K and cooling to ambient temperatures, t.p.d. curves are given in fig. 3. In every case, the sample remains bright red. Only half of the amount of oxygen initially desorbed can be recovered during a second desorption experiment from a dry reoxidized sample.Upon chemisorption, these molecules are dissociatively adsorbed as revealed by the experiment with labelled oxygen. 2(Ag+ZO-) 3 *O, + Agi + 20- + Z+P . A . JACOBS, J . B . UYTTERHOEVEN AND H . K . BEYER 59 Upon oxygen desorption from a AgA-100, the initial state cannot be restored. Oxygen deficient sites and reduced silver ions are irreversibly formed. Therefore, it is not probable that chemisorbed oxygen fills up again oxygen deficient sites, but rather is dissociatively chemisorbed upon silver species containing both reduced and non-reduced silver ions, as proposed by Seff et aL2* X desorprtion temperature/K FIG. 2.-Rate of oxygen desorption from the corresponding degree of auto-reduction of silver ions in AgA-100, together with the corresponding colour changes in the sample.These colour changes and the concomitant phenomena are typical for the A zeolite structure. Indeed, these phenomena are not observed in Ag chabasite, Ag stilbite or Ag clinoptilolite, zeolites with comparable small pores. They neither are observed in AgA zeolite with higher Si/Al ratio, (AgZK-4 zeolite also does not show the colour changes), nor in germanium substituted faujasite with a &/A1 ratio equal to one. However, they are found on AgNaA zeolite with variable degree of Ag+ exchange. All this allows us to conclude that the A zeolite structure is needed, containing equal amounts of Si and A1 tetrahedra. This structure is typified in that desorp tion temperat ure/K FIG.3.-T.p.d. curve of oxygen from AgA-100: (a) desorption from wet sample; (6) desorption from dry sample, cooled from 673 K in oxygen to ambient temperature ; (c) same as (b) but cooled in labelled oxygen ('*02).60 PROPERTIES OF AgNaA ZEOLITES it contains more cations than sites are available per unit cell. Therefore it has ions that can be considered as zero-coordinated by a distance criterion. This property may well be at the origin of the unusual properties. However, an Ag6+(Agg) cluster 2 s in the sodalite cage is rather improbable, since its formation in low ex- changed forms would result in segregation of cations : parts of the crystal should con- tain only Na+ cations, while aggregation of charged silver should occur at a few other spots. Therefore, the existence of isolated linear species in the cubo-octahedra (Ag+ .. . Ago . . . Ag+) in order to explain X-ray diffraction data l3 has a higher probability. The existence of such clusters will be used further as a working hypo- thesis. There is little doubt that the red colour is due to charged clusters, Me,"+ I d 7' M 0.301 24 Ag+ exchange/ % FIG. 4.-Desorption of oxygen from freshly prepared AgNaA, degassed at 673 K for 1 h. (y > x) since NaA upon sodium vapour treatment at 673 K showed the same colour changes. The red colour of sodium vapour treated NaY could indeed be explained by the presence of Nad+ centred4 It is also well known l5 that, in silver halides, the colour is intensified going from F- to I-. At the same time, there is appreciable increase in the covalent character of the Ag .. . X interactions. In compounds with chain structure (as AgCN) the bond is predominantly covalent. All this supports our model. Fig. 4 shows that the oxygen-donating capability of the A zeolite lattice strongly depends on the degree of Ag+ exchange. The phenomenon is strongly enhanced at the highest degrees of exchange. CHEMISORPTION ON RED AgNaA Upon addition of H2 (at 195 K) and CO (at 293 K) on a red AgA-100, no colour However, increasing amounts of H20 cause gradual dis- changes are visible. appearance of the colour. co CHEMISORPTION Fig. 5 shows carbon monoxide sorption for AgNaA zeolites with different Ag+ content. The amount of CO chemisorbed at 293 K increases with the degree of exchange of Ag+ for Naf. However, the deviation from linearity indicates that both properties are not directly related.The total amount adsorbed also is enhanced by the presence of Ag+ and increases almost linearly with the degree of exchange up to values of 30 %. At higher degrees of exchange, this increase is less pronounced. Fig. 6 shows a linear relation between the amount of CO chemisorbed and the intensity of its i.r. stretching vibration. This constitutes firm proof that supercageP . A . JACOBS, J . B . UYTTERHOEVEN AND H. K . BEYER 61 Ag+ cations are the sites for CO chemisorption. Strong CO sorption has also been observed on AgY l6 and AgZSM-5 zeolites. Using an absorption coefficient determined for AgX, a proportionality factor of 1.10 0.13 was obtained between the amount of CO chemisorbed and Ag+ ions available for interaction.The supple- mentary amount of physisorbed CO most probably corresponds to some organization Ag+ exchange/ % intensity vcolarbitrary units FIG. 5.-Carbon monoxide sorption at 293 K on AgNaA zeolites with different Ag+ content; (a) amount irreversibly sorbed at 293 K, total amount sorbed at equal pressures : (b) 13.3, (c) 26.6, (d) 33.33 kN m-2. FIG. 6.-Relation between the amount of CO sorbed and the amount of available Ag ions, as measured by the i.r. intensity of CO at 2180k 5 cm-’ ; (a) amount chemi- sorbed, (b) amount sorbed at 13.3 kN m-’. in the second layer. The relative decrease of this amount at higher exchange levels could be due to some steric hindrance at the sorption site. This is entirely consistent with our hypothesis that mainly at high degrees of exchange Ag+--AgO-Ag+ species are formed.The location of these Ag+ ions is necessarily at site 11’ inside the cubo- octahedron. Direct coordination with one CO remains possible, while organization in the second layer will be weaker. HYDROGEN CHEMISORPTION Silver metal on support is not known to retain hydrogen at ambient temperature. However, AgNaA zeolite does, as shown in fig. 7. It should be stated that in each of the following cases reduction of Ag+ ions by hydrogen can be excluded, since in no case lattice hydroxyls are formed. The phenomenon becomes important for degrees of silver ion exchange above 70 %. Physisorption of hydrogen (at 195 K) is also observed and shows the same overall but less marked changes.Comparison with fig. 4 shows that these properties are related to the ability of the lattice to release oxygen, and in our mind to the formation of Ag+-Ag-Ag+ species. In fig. 8 it is shown for AgA-100 that the amount of hydrogen retained is strongly dependent on the outgassing temperature of the sample. Hydrogen is only retained when the sample is completely degassed and further increases as more and more lattice oxygen comes off. When water is adsorbed, the hydrogen chemisorption capacity of the AgA-lOO(673) sample decreases linearly, at least for small amounts adsorbed (fig. 9). The slope of the straight line equals 1. This indicates that water is preferentially sorbed at the62 PROPERTIES OF AgNaA ZEOLITES ends of the Agf-Ag-Agf species. Each water molecule is able to replace two hydrogen atoms.This is also strong evidence for the dissociative nature of hydrogen chemisorption on the silver agglomerates, just as was shown for oxygen. The differ- ence in chemisorption capacity in each case (Hz/Oz = 6) may be due to the fact that oxygen molecules cannot enter the six rings of the sodalite unit, while hydrogen can. 0 50 I0 0 Ag+ exchange/ % 13.33 kN m-2, and (b) amount adsorbed at 195 K, not desorbable at 293 K. FIG. 7.-Sorption of hydrogen on AgNaA-(673) ; (a) amount reversibly adsorbed at 195 K and For carbon monoxide, which cannot enter either, the Ag/CO ratio also equals 6. This indicates that only one out of five silver cations are accessible from the supercage. This also may reflect the average availability of the ends of the Agf-Ag-Agf species, although for the moment it seems highly speculative to propose two different locations for this species.I 5 00 6 00 outgassing temperature/][( FIG. 8.-Infiuence of degassing temperature of AgA-100 on the hydrogen retention capability ; (a) and (6) same as in fig. 7.P. A . JACOBS, J . B . UYTTERHOEVEN AND H . K . BEYER 63 0.2 / / b 24 water adsorbedlmol kg-’ FIG. 9.-Decrease of the amount of chemisorged hydrogen when increasing amounts of water are sorbed on AgA-lOO(673). In fig. 10 a straight line relation is shown between the amount of Ago atoms obtained through auto-reduction and the amount of irreversibly held hydrogen molecules. Its slope equals one. Samples with different degree of exchange and outgassed at different temperatures fit this relation.Since hydrogen is dissociatively adsorbed (fig. 9), this relation clearly shows that each Ago atom formed by auto- reduction, is associated with two Ag+ ions, each capable of chemisorbing one atom of hydrogen or one water molecule. 0.3 ..( 1 M A4 “0 0.2 E si \ a E g 0.1 s Y 0 0 0.15 0.30 amount Ag+ reduced/equiv. kg-I FIG. 10.-Relation between the degree of auto-reduction (abscissa) of AgNa zeolites and the amount chemisorbed hydrogen (ordinate) ; (a) AgNaA-10(673), (b) AgNaA-30(673), (c) AgNaA-70(673), (d) AgA-100(523), (e) AgA-100(593), cf) AgA-100(673). CONCLUSIONS This work contains strong experimental evidence for the existence of partly reduced silver agglomerates in AgNaY zeolites. The following main observations are on the basis of our hypothesis that these species are isolated and linear Ag+-Ag-Ag+ clusters located in the cubo-octahedra and are responsible for the chemisorption64 PROPERTlES OF AgNaA ZEOLITES properties of the solid : (1) a process of auto-reduction occurs together with intense colouration of the sample, the former implies the appearance of Ago atoms, the latter is indicative of covalent bond formation of silver; (2) comparison with other Ag zeolites different in structure and %/A1 ratio shows that the initial presence of zero coordinated ions requiring A zeolite structure and Si/Al equal to one is related to the formation of coloured centres; (3) the phenomenon occurs at different degrees of Ag+ exchange, making the existence of compact (Agg A$) clusters thermodynamic- ally rather improbable in the present case; (4) oxygen adsorbed dissociatively on the colour centres as evidenced by the tracing with I8O2; (5) there is proportionality between the degree of auto-reduction and the chemisorption of hydrogen, indicating that chemisorption occurs on the colour centres; (6) one molecule of water is able to replace two hydrogen atoms.This shows that hydrogen is dissociatively cheini- sorbed; (7) for every Ago formed, two H atoms are chemisorbed on the colour clusters. Carbon monoxide is chernisorbed in the Ag+ ions available in the supercages, involved in the cluster formation or not. P. A. Jacobs acknowledges a research position as “ Bevoegdverklaard Navorser ” from N.F.W.O. (Belgium). The experimental help of Mrs.I. Szaniszlo (Budapest) and J.-Ph. Linart (Leuven) is appreciated. Stimulating discussions with Dr. W. J. Mortier are also acknowledged. We thank the Belgian Government (Dienten Wetenschapsbeleid) for financial help. M. Ralek, P. Jiru, 0. Grubner and H. Beyer, Coll. Czech. Chem. Comm., 1962, 27, 142. Y . Kim and K. Seff, J. Amer. Chem. SOC., 1977,99,7055. Y. Kim, J. W. Gilje and K. Seff, J. Amer. Chem. Soc., 1977,99,7057. P. A. Jacbos, W. De Wilde, R. A. Schoonheydt, J. B. Uytterhoeven and H. K. Beyer, J.C.S. Furuday I, 1976, 72, 1221. P. A. Jacobs, J. B. Uytterhoeven and H. K. Beyer, J.C.S. Furuduy I, 1977, 73, 1755. P. A. Jacobs, J. B. Uytterhoeven and H. K. Beyer, J.C.S. Chem. Comm., 1977, 128. ’ R. A. Schoonheydt, L. J. Vandamme, P. A. Jacobs and J. B. Uytterhoeven, J. Catalysis, 1976, 43, 292. P. A. Jacobs, M. Tielen, J.-Ph. Linart and J. B. Uytterhoeven, J.C.S. Faraduy I, 1976,72,2793. G. T. Kerr, J. Phys. Chem., 1962, 66,2271. lo G. T. Kerr, US. patent 3,314,752. l1 P. A. Jacobs, Curboniogenic Activity of ZeoZites (Elsevier, Amsterdam, Oxford, New York, l2 R. L. Firor and K. Seff, J. Amer, Chem. SOC., 1977,99, 1112. l3 L. Gellens, W. Mortier and J. €3. Uytterhoeven, to be published. l4 J. A. Rabo, C. L. Angell, P. H. Kasai and V. Schomaker, Disc. Furuduy Soc., 1966,41, 328. 1977), p. 54-55. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (Interscience, New York, London, Sydney, 2nd edn, 1966), p. 1041. l6 H. Beyer, P. A. Jacobs and J. B. Uytterhoeven, J.C.S. Fapaday I, 1976, 72,674. l7 J. A. Rabo, J. N. Francis and C. L. Angell, U.S. Patent 4,019,880. (PAPER 8/382)

ISSN:0300-9599    

DOI:10.1039/F19797500056

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Particle Adhesion and Removal in Model Systems Part 1 .-Monodispersed Chromium Hydroxide on Glass BY JAN E. KOLAKOWSKI~ AND EGON MATIJEVIC" Institute of Colloid and Surface Science and Department of Chemistry, Clarkson College of Technology, Potsdam, New York 13676, U.S.A. Received 7th March, 1978 The interactions of monodispersed chromium(m) hydroxide particles of 0.28 pm diameter with glass were studied using the packed column technique. The particles were first adsorbed by passing the sol through a bed containing glass beads at pH - 3, at which conditions the particles and the beads are of opposite charge. Desorption was then studied as a function of pH by rinsing the column with solutions containing different electrolytes in varying concentrations. Optimum removal at low ionic strength occurred at pH 11 5, which is well above the point of zero charge of the colloidal chromium hydroxide particles.When the particles were adsorbed at pH 4 and - 6, respectively, the removal was substantially less under otherwise identical conditions. The addition of NaN03, NaF, Ca(NO& and C~(dipy)~(ClO& at the pH of optimum removal (11.5) of the rinse solution caused a decrease in desorption and, depending on ionic strength, eliminated altogether particle separation. The greater the counterion charge, the less salt was necessary to suppress desorption. These resuIts are best explained in terms of double layer interactions between the particles and the substrate and no indication of chemical bonding could be detected. The removal of adhered colloidal particles from substrates immersed in liquids or solutions is important in many processes, such as in the cleaning of soiled materials, in the elimination of oxide scale from metals, etc.Most of the work, both experi- mental and theoretical, involving fine particle adhesion is concerned with the de- position (or adsorption) of such particles rather than their removal. Particle deposition has been frequently studied with the rotating disc apparatus. The technique, the theory of which was given by Levich,l was originally used in the investigations of the transport of ions.2 The first application of the rotating disc principle to particle adhesion was given by Marshall and Kit~hener,~ who followed the deposition of carbon black from dilute aqueous suspensions on glass and other substrates.Later, Hull and Kitchener deposited polystyrene latex onto a glass disc coated with either cationic or anionic polymers, and Clint et aL5 examined the adhesion of polystyrene latex on films of polystyrene cast onto glass discs. When the latex particles and the polymer were of opposite charge, the deposition rate followed the Levich equation but the existence of a potential energy barrier to adsorption required a modification of the theory. Using the same technique, Tewari and Camp- bell studied the adhesion of chromium hydroxide spheres and of rod-like P-FeOOH particles onto stainless steel discs, both in the presence and in the absence of a potential barrier. Clayfield and Lumb developed a packed column method to investigate the detach- ment of adhered colloidal particles.The technique was applied to the system carbon black/metal surface in non-aqueous and aqueous surfactant solutions,* and to deposition of monodispersed latexes on glass. p This work is part of a ms. thesis by J. E. K., Clarkson College, 1977. 1-3 6566 PARTICLE ADHESION AND REMOVAL Particle removal may be difficult to interpret because of surface irregularities, the possibility of chemical bond formation, mechanical deformation, hydrodynamic effects, and other phenomena that can occur between bodies in contact. The problem is considerably simplified if only electrical double layer forces and/or diffusion play a role. In this work a well defined model system has been studied. The removal of mono- dispersed spherical colloidal chromium hydroxide particles deposited on glass beads in a packed column was measured as a function of pH, ionic strength and counter- ion charge.The chosen system is well suited for this type of investigation because the particle charge can be readily changed and even reversed by varying the pH. Furthermore, since the particles are uniform in size, their concentration is easily monitored. The results obtained are interpreted in terms of the existing theory of the electrical double layer. EXPERIMENTAL MATERIALS The glass beads were Ballotini type 0, of which > 85 % were in the form of true spheres. The fraction of a sample, passing a no. 230 and retained by a no. 325 A.S.T.M. sieve (resulting in beads with a diameter between 44 and 62pm), was collected and cleaned by stirring in aqua regia followed by a thorough rinse with doubly distilled water.The glass powder was used after being dried overnight in an oven at 95°C and then cooled in a vacuum desiccator. The preparation of the hydrous chromium(1Ir) oxide sols was described in detail else- where.l0 The sols were obtained by ageing a solution 4x mol dm-3 in CT(NO~)~ and 6x mol dm-3 in KzS04 at 75&3"C for 72 h. This gave a modal particie radius of 0.14pm. The sol was then filtered through a 0.1 pm Nuclepore membrane, washed with a 0.001 mol dm-3 HN03 until the excess Cr3+ and SO$- were removed, and resuspended in 0.001 rnol dnr3 HN03. All chemicals were of analytical reagent grade and were used without further purification. Water was doubly distilled, the second distillation being carried out in an all-Pyrex apparatus.All glassware, except the glass beads and plastic containers, were cleaned with chromic acid. METHODS PARTICLE DEPOSITION AND REMOVAL The packed column (powder-bed) procedure was similar to the kind developed by Clay- field and L ~ m b . ~ The apparatus consisted of a two-piece Pyrex chromatographic column with an i.d. of 9 mm and a reservoir capacity of 50 an3. To support the beads, a stainless steel 170x 700 mesh disc was clamped between the two column sections and sealed by two " 0 " rings. The column reservoir, located above the beads, was adapted to be pressurized slightly by nitrogen gas for flow rate control. The effluent was discharged from the packed bed into a 50 cm3 test tube which was attached to the column with a vent to the atmosphere (fig.1). The glass beads (2.00 g) were fed into the column as a slurry suspended in the same super- natant solution as the Cr(OH)3 sol. This gave a bed height of 1.75 cm once all the beads settled. Next, a measured volume of the chromium hydroxide sol of known number con- centration was forced through the bed. The pH of this sol was adjusted to - 3 which resulted in a stable suspension of positively charged particles (fig. 2). At the same pH the glass beads are negatively charged (table 1) onto which the particles of chromium hydroxide were quantitatively deposited, as ascertained by the absence of a Tyndall beam in the column effluent. The bed was then rinsed with 10 cm3 of the sol supernatant solution to remove any particles not truly adhered.The column was washed free of acid using 20 cm3 of the same solution to which sufficient amount of NaOH was added to raise pH to 8. No particles were desorbed during this washing step.J . E. KOLAKOWSKI AND E. MATIJEVIC 67 Finally, additional volumes (10 and 20 cm3) of the particle-free rinse solution of the same pH (- 8) were passed through the bed, but no particle removal was affected. Thus, one can assume that the particles, deposited as described, did adhere to the glass beads and were not mechanically filtered by the bed. The conclusion was substantiated by the observation that an adjustment of the sol to pH 11.5 prior to passing through the column resulted in no particle adhesion; i.e., the entire amount of chromium hydroxide was recovered in the effluent.FIG. 1.-Packed column used in adsorption and desorption studies. (a) N2 inlet ; (6) column reservoir ; (c) column, 9 m i.d. ; (d) glass beads ; (e) Viton “0 ” rings ; (f) stainless steel 170 x 700 mesh disc ; (g) bleed ; (h) collection tube. To investigate the removal of the adsorbed particles, the bed was washed with 10 cm3 of a rinse solution of a given pH and electrolyte concentration. The procedure was repeated as often as necessary and the particle concentration was determined in each effluent sample. Desorption was always carried out at the same salt concentration as present during particle deposition. Thus, the adsorption and removal steps differed only in the pH of the sol and rinse solutions used. All pH adjustments were made with either HN03 or NaOH.The time of every rinse stage was recorded using a stopwatch and the throughput rates varied between 0.75 and 1.0 cm3 min-l. The rinsing of the column for particle removal was done within an hour following deposition unless otherwise stated. PARTICLE NUMBER CONCENTRATION The particle size distributions were obtained by light scattering using the polarization ratio method.’l* l2 The measurements were carried out with a Brice-Phoenix Model 2000 Universal light scattering photometer. The solid content of the chromium hydroxide stock sol was determined by dry weight analysis for which purpose a deionized sol was kept in a vacuum oven at 75°C until no change in weight was observed. Knowing the average particle size, as obtained by light scattering,68 PARTICLE ADHESION AND REMOVAL and the density of the material (2.42k0.02 g ~ m - ~ ) the particle number concentration of the sol could be readily e~tab1ished.l~ The ratio of scattering intensities, measured at 45 and 0" angles, for vertically polarized light of wavelength 436 nm, was found to be directly pro- portional to the number concentration of the chromium hydroxide sols for dilutions of interest in this work.These optical measurements were, therefore, used to determine the particle number concentration of the effluents in the desorption studies. ELECTROKXNETICS Electrophoretic mobilities of the sol particles were determined with a Rank Brothers Microelectrophoresis Apparatus Mark I1 (Bottisham, Cambridge) using a van Gils cell.Fig. 2 gives the mobilities of the chromium hydroxide particles as a function of pH. Different symbols indicate measurements carried out by four different investigators using different samples of the sols prepared by the same procedure. I I I I I 1 I I 2 4 6 8 10 12 PH FIG. 2.-Electrophoretic mobility of a monodispersed chromium hydroxide sol consisting of spherical particles (modal diameter 0.28 pm) as a function of pH at 25°C. Various symbols show data obtained by four different investigators on two different samples. Zeta potentials, 5, were calculated by means of the Henry equation : c = 1.5vlPeld1 + m a ) ] (1 1 where ,ue is the particle mobility in ,um s-l/V cm-', q the viscosity, E the permittivity of the medium, and the values of the correction term [1+3.(rca)] were taken from tables by Smith.14 In this factor a is the particle radius and K is the reciprocal of the Debye-Huckel distance.Some of the 5 potentials for systems of interest in this work are given in table 1. The zeta potentials of the glass beads were obtained from electro-osmosis measurements carried out with an apparatus described in detail by Mirnik et aE.15 using platinum electrodes, one of which was coated with AgCl. The values of 5 were calculated by means of the expression in which Y is the snecific conductivitv of the solution and h is the done of the nlot obtained 5 = 4 n q ~ b / ~ (2) of the current. for the chromium hydroxide particles. Table 1 gives zeta potentials of glass under the same conditions as reportedJ . E . KOLAKOWSKI A N D E. MATIJEVIC 69 TABLE 1 .-ZETA POTENTIALS OF CHROMIUM HYDROXIDE PARTICLES AND OF GLASS PH 3.1 4.0 5.8 9.6 10.3 11.0 11.5 11.5 11.5 11.5 11.5 11.5 11.5 zeta potential, </mV electrolyte(s) Cr(OH)3 glass HN03 HN03 NaOH NaOH NaOH NaOH HNO, NaOH, lo-, mol dm-3 Ca(N03)2 NaOH, mol dm-, Ca(N03)2 NaOH, lod5 mol dm-3 Ca(N03)2 NaOH, mol dm-3 C~(dipy),(ClO~)~ NaOH, mol dm-3 C~(dipy),(ClO~)~ NaOH, mol dm-3 C~(dipy),(ClO~)~ 46 37 28 - 32 - 37 - 45 - 47 - 23 - 27 - 34 - 19 - 37 - 38 - 36 - 63 - 18 - 48 - 89 - 137 - 43 - 65 - 88 - 43 - 99 - - RESULTS Fig. 3 is a plot of the fraction of chromium hydroxide particles remaining on the glass beads in the column after each of several rinse cycles.Each curve is for a different pH of the wash solutions (adjusted with NaOH).The abscissa is given in time of elution calculated from the known volume of rinse liquid and the flowrate. In all cases approximately the same number of particles was deposited on glass beads (- 1.3 x lolo). Because this number varied from experiment to experiment, data are presented on a relative basis. 1 .o 0.8 i I I I I t 0 2 4 6 8 10 time x 10-3/s FIG. 3 .-Fraction of monodispersed spherical chromium hydroxide particles (modal diameter 0.28 pm) desorbed from glass on repeated elution with rinse solutions of different pH: (0) 9.6, (0) 10.3, (0) 11.0, (0) 12.6 and (A) 11.5. The abscissa is calculated from the known volume of rinse liquid (- 150 cm3) and the flowrate. In all cases, approximately the same number of particles was deposited on glass (- 1.3 x 1O'O) at pH - 3.70 PARTICLE ADHESION AND REMOVAL Although the zero point of charge of the chromium hydroxide particles used is - 8.0, the pH of the rinse solution had to be adjusted to - I0 before any particle desorption occurred.At still higher pH the removal increased, reached a maximum at pH 11.5, and then decreased again when the rinse solution was of pH 12.6. Even at optimum removal conditions a measurable fraction of the particles remained on the beads. The effect of the pH of the sol used in particle deposition on their subsequent removal is illustrated in fig. 4(a). The three curves refer to systems adsorbed at pH 3.0 (as shown before in fig. 3), 4.0 and 5.8, respectively, whereas elution in all cases was carried out with rinse solutions of pH 11.5.Obviously, there was a marked decrease in removal of the particles adsorbed at higher pH values. 1.0 0.8 0.6 0.4 0.2 1 .O 0.8 0.6 0.4 0.2 0 0 - 0 2 4 6 I I 1 I 2 4 6 8 time x 10-3/s RG. 4.-(a) Effect of the pH of the sol during the adsorption stage [( 0) 3.0, (0) 4.0 and (A) 5.8, respectively] on the desorption of chromium hydroxide particles from glass at pH 11.5 for the same system as described in fig. 3. (b) Effect of the ageing period of the chromium hydroxide particles on glass adsorbed at pH 3 prior to rinsing with a solution of pH 11.5 for the same systems as de- scribed in fig. 3. (0) unaged, (0) 21 h, (A) 1 week. The time of ageing after the particles were adsorbed also greatly affected their subsequent separation. Chromium hydroxide was deposited at pH 3.0 and left on the beads, which were in contact with supernatant solution at room temperature for varying periods of time [fig.4(b)]. The column was then repeatedly rinsed with solu- tions of pH 11.5. Ageing for 21 h prior to elution caused a large reduction in de- sorption, and after ageing for 1 week, the removal was negligible. To study the effect of ionic strength and of the nature of the counterions on the desorption of chromium hydroxide particles from glass, several electrolytes in different concentrations were added to the investigated systems. Fig. 5 and 6 show the results obtained when NaN03, NaF, C~I(NO~)~, and Co(dipy)3(C104)3, respec- tively, were present in the rinse solution, the pH of which was always adjusted to 11.5. As described in the experimental section, in each case the adsorption stepJ .E . KOLAKOWSKI AND E. MATIJEVIC (4 (b) 71 0 0 2 4 6 8 time x 10-3/s FIG. S.-Effect of different concentrations of NaN03 (a) and of NaF (b) on desorption of chromium hydroxide from glass for the same system as described in fig. 3. Adsorption pH3.0, desorption pH 11.5. (a) (0) 0.2, (0) 0.1, (A) 0.04, (0) 0.02 and (0) 0.0 mol dm-3; (b) (0) 0.1, (A) O.Ol,!(tl) 0.001 and (0) 0.0 mol dm-3. (4 (6) 1.0 0.8 S 0.6 ' 0.4 z \ 0.2 0 1.0 0.8 0.6 0.4 0.2 A-A-A- I I I 0 2 4 6 0 2 4 6 time x 10-3/s FIG. 6.-Effect of different concentrations of Ca(N03)2 (a) and of Co(dipy)3(C104)3 (b) on desorption of chromium hydroxide from glass for the same system as described in fig. 3. Adsorption pH 3.0, desorption pH 11.5. (a) (A) lW3, (0) lW4, (0) (0) 0.0mol dm-3; (b) (A) (0) (0) (0) 0.0 mol dmd3.72 PARTICLE ADHESION AND REMOVAL was also carried out with chromlum hydroxide sols containing the same concentration of a given salt.The addition of salts can considerably decrease particle desorption and, at suffi- ciently high electrolyte concentrations, the removal can be completely suppressed. The efficiency of a given salt depends strongly on the charge of the counterions. Thus, only mol dm-3 of tris(2,2'-dipyridyl)cobalt(r11)-ion, Co(dipy)i+ , suffices to completely inhibit particle desorption. Calcium ion requires a ten times higher concentration to achieve the same effect (fig. 6), whereas in the presence of 0.2 mol dm-3 NaN03 some removal still occurs (fig. 5). The chelated Co"' ion was used in order to test a highly charged counterion a t high pH.No polyvalent uncomplexed cation is available at pH 11.5. In fig. 5 the two sets of data refer to the addition of NaN03 and NaF, respectively. The latter salt was used to test the possible effect of the fluoride ion. If adsorption of the chromium hydroxide particles on glass was due to chemical bond formation between the silanolic and the =CrOH groups, fluoride ions should greatly affect the removal process. DISCUSSION The rate of particle desorption depends on a number of factors including external forces (e.g., hydrodynamic forces), chemical bonding, particle diffusivity and the interaction potentials between the substrate and the adsorbed material. It is there- fore necessary to first establish which of these parameters may play a role in a given sys tem.The Reynolds number, Re, for the flow through a packed bed is given by l6 where R is the bead radius, G the superficial mass flowrate in g cm-2 s-l (the flow in the absence of the beads), and Po is the bed porosity (void volume per total bed volume). The Reynolds number for the flow rate of 1.0 cm3 min-l, used in this work, was found to be 0.015. This value indicates that hydrodynamics played an insignificant role in the particle desorption. This is further confirmed by the finding that a signifi- cant increase in the rate of flow of the rinse liquid had no effect on the rate of de- sorption. Clayfield and Lumb' showed for their systems that, under similar con- ditions, a variation in flow rates had negligible effect in changing the extent of particle removal.A cursory inspection of the data would seem to indicate that chemical bonding between the acidic silanolic groups of the glass surface and the basic groups on the chromium hydroxide particles may play a significant role in the described adsorption/ desorption phenomena. For example, a sizeable fraction of the particles remains adsorbed even under the most favorable conditions as seen in fig. 3. Furthermore, the higher the pH of the sol at the adsorption step (however, lower than the P.z.c.), the more difficult becomes the particle removal [fig. 4(a)]. One could postulate that the larger number of hydroxyl groups per unit area on the chromium hydroxide particles at higher pH provide more bonding sites for the silanolic groups on glass.Finally, the particle escape decreases with ageing time of adsorbed chromium hydroxide on glass. Considerable effort was made to establish the formation of the =Cr-0-Si= bond between the adsorbent and the adsorbate, but without success. In a series of ex- periments the glass beads were stored overnight at room temperature and at 100°C in 10-l mol dm-3 Cr(NO& solution, the pH of which was adjusted with NaOH to 4, Re = 2RG/y(l -Po) (3)J . E. KOLAKOWSKI AND E. MATIJEVIC 73 Xn neither case could a change in the c-potential of the glass be detected when com- pared to the untreated beads. Under the described conditions, chromium ions are strongly hydrolysed (particularly at the elevated temperature) to give polynuclear cationic complexes. Any adsorption of these species through bonding to silanolic groups would greatly affect the surface charge of glass, ultimately reversing the sign of it.A chromium hydroxide-silanolic condensation bond would also be attacked by fluoride ions, resulting in a much more efficient removal of the adsorbed particles in the presence of NaF than in the presence of NaNO,. Yet with both electrolytes rather similar effects were observed (fig. 5). It would, therefore, appear that the desorption phenomena of chromium hydroxide from glass should be interpreted only in terms of particle diffusion in a region of interacting potentials. In the absence of chemical bonds the interaction between two solids in an ionic medium is due to the contribution of London-van der Waals and electrical double layer potentials.The unretarded attraction potential between a sphere and a flat plate at a distance of separation x < 250 A is 17* l8 x(x+2a) 1 ' A132 x+2a 2a(a+x) 4*(x) = ?Pn - - where A l S 2 is the overall Hamaker constant for the system which can be approximated by A 1 3 2 ( d ~ l - d ~ 3 ) ( d ~ 2 - d ~ 3 ) * (4) sphere-medium-plate ( 5 ) A l l , A22, and A33 are the individual Hamaker constants for interaction between the materials composing the spheres, plates and the aqueous medium (all in vacuo), respectively. The double layer potential between a sphere and a flat plate is given by the follow- ing solution of the linearized Poisson-Boltzman equation in which the upper sign is for the condition of constant surface potentials (t,bl and $ 2 ) and the lower sign for the condition of constant surface charge densities, whereas IC is the Debye-Hiickel reciprocal length.Eqn (6) was derived from the expression of Hogg, Healy and Fuerstenau 2o for the interaction of two dissimilar spheres, taking the radius of one sphere as infinite. The total interaction potential between a sphere and a plate is then given by The potential energies of interaction as a function of separation were calculated for different solutions employed in both deposition and removal of particles in the system chromium hydroxide + glass using the electrokinetic data in table 1 and eqn (4), (6) and (7). In eqn (6) the condition of constant potential was assumed. The overall Hamaker constant, A132 was calculated to be 0.8 kT from eqn (5).This value was based on using = 10.6 kT, for the water-water constant, as ob- tained from the Lifshitz theory and tabulated by Visser.l* For Cr(OH), particles A l = 14.9 kT was determined experimentally from measured coagulation rates of the corresponding sol 21 and for glass, the value A22 = 20.9 kT was calculated by Biittner and Gerlach 22 for the silica-silica interaction. 4 ( 4 = 4 ' 4 ( X ) + 4 R ( X ) . (7)74 PARTICLE ADHESION AND REMOVAL Fig. 7* gives in the insert the total potential energy as a function of distance for adsorption of chromium hydroxide spheres on glass at two different pH values. Obviously, few adsorbed particles are expected to be found many Angstrom units away from the surface due to the fact that 4(x) approaches zero at large separations.The considerable difference in the attraction at two different pH values may be attributed to the thickness of the hydration layer between the substrate and the particles which determines the distance of closest approach, xo. At pH 4.0 the particles penetrate more deeply into the hydration layer and, therefore, are more strongly adsorbed. As a result their removal at higher pH is less efficient as indeed 4c 30 b4 % 2a 3 n -8- 10 0 -I 0 . _ 0 so 100 I50 distance xlA FIG. 7.--Calculated total potential energy curves as a function of distance using the sphere-plate model [eqn (4) and (611 for the chromium hydroxide+glass systems described in fig. 3. Desorption pH : (a) 11.5, (6) 11.0, (c) 10.3 and (d) 9.6. Insert corresponds to systems in iig.4(a) ; adsorption pH : (i) 3.0, (ii) 4.0. observed [fig. 4(a)]. Similar explanation may apply to the ageing effect [fig. 4(b)]. With increasing time of adsorption the particles approach the substrate more closely. Alternately, these effects may be due to surface roughness of the glass beads. On ageing, more particles may move into the surface crevices and thus become more difficult to remove. When solution conditions are changed desorption may occur. In this respect, the pH of the rinse solution plays a major role. Fig. 7 gives total potential energy curves as calculated for the chromium hydroxide+glass system at four different pH values used in the desorption studies (fig. 3). At pH 9.6 attraction still persists * See note added in proof, p. 78.J .E. KQLAKOWSKI AND E . MATIJEVIC 75 and the particles remain adsorbed. Similarly, at pH 10.3 there is no significant potential drop at x > x1 (x, being the location of the maximum) and, therefore, relatively few particles can escape. At higher pH values all particles at sufficient distance xo, which enables them to overcome the energy barrier, will desorb, but will be unable to readsorb. Fig. 8 shows the calculated total potential energy curves for the systems containing Ca2+ and Co(dipy)$+ ions in varying concentrations corresponding to the desorption experiments illustrated in fig. 6. Again these diagrams explain well the observed particle removal phenomena. A comparison of the total potential curves shows that these are quite similar for mol dm-3 Ca2+ and mol dm-3 Co(dipy)z+ and again for mol dm-3 Ca2+ and mol dm-3 Co(dipy)$+. Each of these pairs of salts shows comparable desorption phenomena at cited concentrations (fig.6). Thus, the analogy is striking. 40 - 30 - Q 20- n -8. 3 10- 0- 40 3 0 , 20 10,. .- 0 50 100 IS0 0 100 I so distance x/A FIG. &-Same as fig. 7 for systems described in fig. 6. (a) Ca(NO& : (i) 0, (3) lo-’, (iii) (iii) moi d ~ n - ~ . (iv) lov3 mol dm-3 ; (b) Co(dipy):+ : (i) 0, (ii) The diffusion coefficient, D, for a particle moving normal to a plane surface can be given by where k is the Boltzmann constant, T the absolute temperature and f is the friction coefficient which for a sphere of radius a in a fluid of viscosity q is given by the Stokes law D = kT/f (8) f = 6nqa (9) For the diffusional escape of particles, physically bound to substrates : dN(t) - - I -pN(t) dt76 PARTICLE ADHESION AND REMOVAL where N(t) is the number of particles adhered to the substrate at time t andp is the escape probability. If p is time independent, it becomes a rate constant, and integra- tion of eqn (10) gives : N(t) = N(0) exp [-pi] (1 1) Dahneke 23 gives two solutions for constant p , one describing '' equilibrium " and the other " non-equilibrium " desorption.The equilibrium condition means that for every particle which diffuses away from the substrate and escapes another diffuses toward the substrate and replaces the lost particle. This can happen when there is no potential barrier present ['(x) -+ 01; i.e., when &(x) < 0. To escape the particle must have sufficient thermal energy to leap out of the potential well in order to diffuse away from the substrate.For equilibrium desorption, p is given by where C = [8kT/nrnIt m being the particle mass, and y = 11: exp[ -'(x)/kT] dx. (14) The upper limit of the integral, x l , is the distance at which ' ( x ) reaches a maximum value or, effectively, an asymptotic value. In the case of equilibrium desorption, the maximum tends to zero. The lower limit xo is the distance of closest particle approach to the substrate. Particles within these limits are considered adsorbed and those whose separation exceeds x1 are considered to be "free". The value of '(x,) is taken arbitrarily to be zero so that the potential barrier will be equal to '(x,). In non-equilibrium desorption a greater number of particles diffuse out of the region between xo and x1 than into it.In this case the function 4(x) has a true maximum at xl. Dahneke gives the following expression for the probability of particle escape in the non-equilibrium situation This expression is valid if 4(x) can be represented as a parabola in the vicinity its maximum as follows : 5 ) of where x x1 and w2 is the 2nd derivative or curvature of ' ( x ) evaluated at its maximum. Eqn (15) was found to be valid for +(xl) 2 10 kT.24 This means that particle desorption has to occur at a slow enough rate such that the distribution of particle energies remains unchanged. The non-equilibrium solution does not take into account particle readsorption into a secondary minimum of ' ( x ) .If the peak of 4(x) is sharp enough such that mo21f2 $ 1 then the factor in brackets in eqn (15) approaches unity and eqn (15) reduces to eqn (12), the equilibrium solution. The physical interpretation of particle escape over a sharp maximum is that the particles have so short a distance x to travel in order to be free of the substrate that friction has a negligible role in impeding their escape.J . E . KOLAKOWSKI AND E . MATIJEVIC 77 The function -In [N(t)/N(O)] was plotted against time for all desorption condi- tions studied. In several cases the curves were linear indicating that p is time inde- pendent [fig. 91. Using the experimentally determined rate constants in eqn (ll), xo was determined for four of the five desorption conditions. For this purpose, different values for xo were substituted into the following expression : which is a rearranged eqn (15), until the root of F(xo) was found.Eqn (9) was used to calculate the coefficient$ The results are given in table 2. The above method cannot be used to calculate xo for desorption data at pH 9.6, since the function 4(x) at its peak cannot be represented by a parabola. 1.5 1.0 ;;2 \ m c1 z u a 0.5 I 0 2 4 6 8 10 time x lO-”/s FIG. 9.-Plot according to eqn (11) of the data obtained for the desorption of chromium hydroxide spherical particles from glass under several different conditions illustrated in fig. 3 and 4. Adsorption pH: (V) 3.0 (21 h ageing), (A) 3.0 (1 wk ageing), (0) 3.0, (0) 3.0, (0) 5.8; desorption pH: (V) 11.0, (A) 11.0, (0) 9.6, (0) 11.0, (0) 11.5.It is also recognized that eqn (9) is an approximation. Calculations, based on expressions for the frictional coefficient, which consider particles moving normal to a planar surface,23 have shown that the error in using eqn (9) does not substantially affect the results given in table 2. TABLE 2.-vARIOUS PARAMETERS CALCULATED FOR THE DESORPTION OF CHROMIUM HYDROXIDE PARTICLES FROM GLASS desorption conditions PH desorption p x 1051s-1 ~01.4 (b(x1)lkT PH adsorption 3.0 11.0 13.7 6.3 26.5 3.0 9.6 1.1 3.0, aged 1 w 11.5 1 .o 9.1 29.4 3.0, aged 21 h 11.5 9.1 9.5 27.4 5.8 11.5 9.8 9.5 27.4 - -78 PARTICLE ADHESION A N D REMOVAL It is reasonable to expect that the adsorbed particles will populate the region of lowest energy, i.e., at the distance xo from the substrate.However, the high energy barriers which correspond to the values of xo (table 2) indicate that the modal particle separation cannot be at xo and that the hydration layer on the glass and/or on the particles must greatly affect the particle separation distance. Otherwise no desorption would be likely to occur, owing to the high energy barriers. The apparent in- consistency is probably due to errors in the calculation of the 4(x), which arise from the uncertainty in the value of the Hamaker constants and the application of the linearized solution of the Poisson-Boltzmann equation, eqn (6), beyond its range of validity (Le., $ < 25 mv). Both of these factors can greatly alter the 4(x) against x curves. Finally, calculations based on the constant charge model gave much higher values of #(x).Note added in proof.- The total potential energies (+/kT) as plotted in fig. 7 and 8 are about one order of magnitude too low. This error was introduced in the calculation of the double layer repulsion energies +R(x) [eqn (6)] which were underestimated by a factor of 4n. The latter factor must be introduced when S.I. units are used, Fortunately, this error in calculation does not affect any of the conclusions as derived from the particle desorption studies. E. M. is indebted to Drs. J. G. Maroto and M. A. Blesa of CNEA, Buenos Aires, Argentina, for pointing out the erroneous magnitude of the total potential energy values. The authors are indebted to Profs. Eric Clayfield (Shell Research Centre, Thornton, England), Barton Dahneke (University of Rochester) and Eytan Barouch (Clarkson College) for many valuable discussions.The assistance of Mr. Terry A. Ring in various computations is greatly appreciated. We acknowledge the helpful remarks of one referee. This work was supported by a grant from N.S.F. V. Levich, Acta Physicochim. U.S.S.R., 1942, 17, 257. A. C. Riddiford, Adv. Electrochem. Electrochem. Eng., 1966,4,47. J . K. Marshall and J. A. Kitchener, J. Colloid Interface Sci., 1966, 22, 342. M. Hull and J. A. Kitchener, Trans. Faraday Soc., 1969, 65, 3093. ’ G. E. Clint, J. H. Clint, J. M. Corkill and T. Walker, J. Colloid Interface Sci., 1973, 44, 121. P. H. Tewari and A. B. Campbell, Abstracts 172 A.C.S. Nat. Meeting (San Francisco, Sept. 1976) ; also personal communication by P. H. Tewari. E. J. Clayfield and E. C. Lumb, Disc. Faraday Soc., 1966, 42,285. E. J. Clayfield and A. L. Smith, Enuiron. Sci. Technol., 1970, 4, 413. J. A. Fitzpatrick and L. A. Spielman, J. Colloid Interface Sci., 1973, 43, 350. l o R. Demchak and E. MatijeviC, J. Colloid Interface Sci., 1969, 31, 257. l1 M. Kerker, E. MatijeviC, W. Espenscheid, W. A. Farone and S . Kitani, J. Colloid Sci., 1964, l2 M. Kerker, The Scattering of Light (Academic Press, New York, 1969). l3 A. Bell and E. MatijeviC, J. Phys. Chem., 1974, 78, 2621. l4 A. L. Smith, in Dispersions of Powders in Liquids, ed. G. D. Parfitt (Elsevier, New York, 1969), l5 M. Mirnik, V. PravdiC and F. Matijevac, Croat. Chem. Acta, 1958, 30,207. l6 R. Bird, W. Stewart and E. Lightfoot, Transport Phenomena (J. Wiley, New York, 1960), p. 196. l7 H. C. Hamaker, Physica, 1937, 4, 1058. l8 J. Visser, in Surface and Colloid Science, ed. E. MatijeviC (Wiley-Interscience, New York, l9 D. C. Prieve and E. Ruckenstein, J. Colloid Interface Sci., 1977, 60, 337. 2o R. Hogg, T. W. Healy and D. W. Fuerstenau, Trans. Faraday Suc., 1966,62,1638. 21 A. Bleier and E. MatijeviC, J. Colloid Interface Sci., 1976, 55, 510. 22 H. Buttner and E. Gerlach, Chem. Phys. Letters, 1970,5, 91. 23 B. Dahneke, J. Colloid Interface Sci., 1975, 50, 89. 24 S. L. Zimmer and B. Dahneke, J. Colloid Interfuce Sci., 1976, 54, 329. 19, 213. p. 129. 1976), vol. 8, p. 21. (PAPER 8/419)

ISSN:0300-9599    

DOI:10.1039/F19797500065

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Association Between Polar Molecules Part 1 .-Nuclear Magnetic Resonance Study of the Dipole Association of Hexamethylphosphoramide with p-Substituted Nitrobenzenes BY HIDEAKI FUJIWARA,* TOSHIKAZU TAKABA, YUTAKA YAMAZAKI Faculty of Pharmaceutical Sciences, Osaka University, Yamadakami, Suita, Osaka, Japan AND YOSHIO SASAKI Received 17th March, 1978 Dipole association between hexamethylphosphoramide (HMPA) and p-substituted nitrobenzenes was investigated in non-polar solvents. N,N-dimethyl, 0-methyl, methyl and chloro groups were considered as substituents, and carbon tetrachloride and methylcyclohexane were used as solvents. The data are analysed assuming the formation of a 1 : 1 associate between HMPA and nitrobenzenes. A significant substituent effect on both the equilibrium constant and the induced shift of association is confirmed, and these results are discussed in relation to the dipole moment or the linear combination of the substituent constants ci and u,f.Association between polar molecules in solution is of interest in the study of the solvent and substituent effects on the reactivity of organic compounds. Although such association, often called dipole association, has been investigated by several physico-chemical methods of analyses,1-8 it is too weak for its character to have been clarified in detail. For example, although dipole association of DMSO and DMF with substituted benzonitriles has been investigated by analysing the concentration dependence of integrated intensities of i.r. bands in carbon tetrachloride,2 relatively large experimental errors were considered to have masked the expected substituent effects on the equilibrium constant.Thermodynamic parameters of the self- association of dipoles have been estimated with some aliphatic nitriles, but the values obtained by different groups of investigators 4* are incompatible with each other. N.m.r. studies of the self-association of some ketones, nitriles and nitro compounds reveal that the association constant is very In the present study we considered hexamethylphosphoramide (HMPA), which is 9 p lo the most remarkable of the aprotic polar solvents ; it has a dipole moment of 5.37 D and is soluble in many organic solvents. We studied the association of this substance with p-substituted nitrobenzenes in non-polar solvents by n.m.r. spectroscopy.The results are discussed assuming the formation of a 1 : 1 associate. EXPERIMENTAL Commercial methylcyclohexane was shaken with 10 vol % of the mixed acid to remove a small amount of toluene, and washed successively with water, 20 % sodium hydroxide, and water (three times). It was then dried over Pz05 and % 30 % was discarded prior to distillation, to remove a trace amount of cyclohexane. Commerical carbon tetrachloride of J.I.S. G.R. grade was dried over PzOs and distilled. HMPA was dried over barium oxide and distilled under reduced pressure. p-substituted nitrobenzenes purified by recrystallization were stored in a desiccator containing silica gel and the solvents and HMPA were stored over molecular sieves 4A. 7980 ASSOCIATION BETWEEN POLAR MOLECULES [2,4,6-2H3]Nitrobenzene (isotope purity 95 %) was synthesized from [2,4,6-2H3]aniline by the decomposition of the diazonium salt with Na3Co(N0&.l [2,4,6-2H3]Aniline was obtained by the acid catalysed exchange reaction of H with D in D20.12 N.m.r.spectra were observed with a Hitachi R-22 spectrometer operating at 90 MHz and at 34.loC, and the shift was measured by a frequency counter within an error of +O.l Hz. As an internal reference 0.02 vol % TMS was added to the solvent. No discernible differ- ence was observed in the parameter calculated below, when a solvent peak (a low field peak of the methyl doublet of methylcyclohexane) was used as the reference. RESULTS AND DISCUSSION Charge transfer and hydrogen bonding l3 interactions have been suggested to explain the interaction between HMPA and nitrobenzenes, in addition to the electro- static attraction between dipoles.However, negligible contribution from hydrc en bonding may be proved from the fact that addition of pyridine and triethylar.L ie, which are stronger bases than HMPA, to dilute solutions of p-nitroanisole causes far smaller changes in the l H shift than does HMPA (table 1). If a charge transfer force is dominant in the present case, high-field 13C shifts of the nitrobenzenes, TABLE EFFECT OF ADDITION OF THREE SUBSTANCES ON THE lH SHIFT OF p-NITROANLSOLE IN METHYLCYCLOHEXANE' added substance A2.6b A3.P AOCH3 HMPA 0.046 0.338 0.158 pyridine - 0.01 5 0.024 - 0.010 triethylamine 0.008 0.014 0.006 a Observed changes (p.p.m.) in the shift of 0.03 rnol dm-3 p-nitroanisole on addition of the specified substances by 0.7 mol dm-3 are listed.Plus sign denotes the low field shift on addition. b The numbering corresponds to that for nitrobenzene. A = S(HMPA added)- S(HMPA free). electron acceptors, would occur on association with HMPA, an electron donor. l4 However, this is not the case, since addition of HMPA to dilute solutions of p- substituted nitrobenzenes causes down-field shifts for both lH and 13C shifts in the latter (table 2). Furthermore, no new bands which might be assigned to charge transfer were observed in the electronic spectra of HMPA solutions of the nitro- benzenes. From these results, we may conclude dipole association to be dominant in the interaction between HMPA and the nitrobenzenes.This is also supported by the facts that tetramethylurea and nitroethane, with smaller dipole moment than TABLE 2.-EFFECT OF ADDITION OF HMPA ON THE 'H AND l3C SHIFTS OF p-NITROANISOLE AND P-NITROTOLUENE IN METHYLCYCLOHEXANEa compound citeb A1Hlp.p.m." A*3C/p.p.mec p-nitroanisole 296 0.076 0.18 3, 5 0.350 1.08 OCH3 0.148 1 .oa p-nitro t oluene 2,6 0.071 - 0.01 3, 5 0.258 0.83 CH3 0.064 d a Observed changes in the shift of 0.2 mol dm-3 substituted nitrobenzenes on addition of 1 mol dm-3 HMPA are listed. 13C n.m.r. spectra were measured with a Hitachi R-22 CFT n.m.r. spectro- meter using methyl signal of the solvent as an internal reference. Plus sign denotes the low field shift on association. C A = S(HMPA The numbering corresponds to that for nitrobenzene.added) - S(HMPA free). d Methyl carbon signal was masked by the solvent peak.H. FUJIWARA, T . TAKABA, Y . YAMAZAKI AND Y . SASAKI 81 HMPA, cause smaller changes in the shifts, and that molecules with no appreciable dipole moment, such as benzene, xylene and hydroquinone dimethyl ether, undergo only a small change in the l H shift (< 0.02 p.p.m.) on addition of HMPA under similar conditions to those in table 1. Although association between polar molecules is often represented by an anti- parallel pair of dipoles, determination of the composition of the associate is necessary. The Job plot, a well-known method of determining the composition of the complex, supports the existence of a l(HMPA):l(nitrobenzene) associate in all cases, as is shown in fig.1. 8 a X CA+ CB FIG. 1 .-Job plots for the p-nitroanisole + HMPA system in methylcyclohexane. A = &,sd- 8 ~ . 0, Experimental point for the meta protons, with regard to the nitro group, of p-nitroanisole. A, Experimental point for the methoxy protons of p-nitroanisole. To determine the association constant, concentration shifts were measured, maintaining the concentration of the nitrobenzene as low as possible (x 0.03 mol dm-3) and varying that of HMPA from 0 to 0.7moldm-3; this gave enough changes in shifts for data processing. Carbon tetrachloride and methylcyclohexane were used as an inert solvent. The results are shown in fig. 2 for the p-nitroanisole + HMPA system as an example. In the interpretation of these results, equilibrium of eqn (1) was assumed and root mean square deviations of the calculated [eqn (2)] and observed shifts were minimized by the curve fitting method.? Assumption of eqn (1) may be reasonable because we have treated a dilute solution, where binary collision is dominant, and because of the results of the Job method mentioned earlier.R A+B = AB In eqn (2), dcalcd = calculated shift of A, dA = shift of A in the free state, J A B = shift of A in the associated state, AAB = JAB-aA, C, = initial concentration of A, CB = initial concentration of B and K = equilibrium constant. t The calculation was performed with a NEAC 2200 model 700 computer at the Osaka University Computer Centre. A library program for the minimization of functions using derivatives was employed.82 ASSOCIATION BETWEEN POLAR MOLECULES As for the value of aA, data for samples without HMPA were used.K and aAB were reproducible to within M 5 % and f 1 Hz, respectively, on repeated runs, and the root mean square deviation was always < 0.2 Hz. The shift of ortho protons with regard to the nitro group exhibited insufficient concentration dependence to allow data processing, therefore, the data discussed below are for meta protons unless otherwise specified. In the cases of p-nitroanisoIe and N,N-dimethyl-p- nitroaniline, exceptionally, the methyl protons showed such a large concentration dependence (x 10 Hz) in methylcyclohexane as to make the calculation reliable, 7.0- 6.81 , , 0 0.4 0.8 CHMpA/mOI dm--3 FIG. 2.-Concentration shifts of the meta protons, with regard to the nitro group, of p-nitroanisole.0, Experimental point; -. calculated curve with K = 3.50 dm3 mol-l, 6m = 7.302 p.p.m., and 6~ = 6.827 p.p.m. and the K values thus determined were comparable with those obtained for the meta protons (table 3). The K and AAB values for each substituted nitrobenzene were almost equal in carbon tetrachloride and methylcyclohexane. Thus, the following equations were obtained by the least squares calculation based upon linear relation- ship, K(Cc1,) = 0.3 K(c6Hl1CH3), A.4B(CC14) = O.78 AAB(C6H1 1CH3), r.m.s.d. = 0.12 dm3 mol-l, r.m.s.d. = 0.04p.p.m., where r.m.s.d. denotes root mean square deviation between the observed and the calculated values. Small values of K in carbon tetrachloride, compared with those TABLE 3.-ASSOCIATION CONSTANT (K) AND BOUND SHIFT (AAB) DETERMINED FOR THE DIPOLE ASSOCIATION OF HMPA WITH P-SUBSTITUTED NITROBENZENES IN CARBON TETRACHLORIDE AND METHYLCYCLOHEXANEu K A+B K AAB substituent (in carbon tetrachloride) (in methylcyclohexane) [2,4,6-'H31 0.47 0.388 1.65 0.448 P-CH, 0.57 0.362 1.43 0.481 p-c1 0.66 0.460 2.52 0.542 p-OCH3 1.10 0.350 3.40 0.482 (3.37 0.2 1 4) p-N(CH312 0.97 0.252 2.46 0.392 (2.30 0.206)b a K and A m are determined from the data of the meta protons with regard to nitro group, the These values are estimated units being dm3 mol-I for K and p.p.m.for AAB. Am = SAB- 8 ~ . from the data for the methyl protons in the substituent.H. FUJIWARA, T . TAKABA, Y . YAMAZAKI AND Y . SASAKI 83 in methylcyclohexane, suggest the preferred solvation of polar molecules with the solvent molecules in carbon tetrachloride.As shown in fig. 3, AAB was approximately linearly related to the dipole moment p of p-substituted nitrobenzenes, AAB being small for the strong electron donating group. This suggests that the substituent partly compensates for the decrease in 0.6 d 0 . 4 - 2 -5 a 0.2- 0.61 1 ' I I d 4 a a -2 O.Zl*, 0 3 6 PIDebYe FIG. 3.-Correlation between the induced shift (Am) and the dipole moment (p). p is cited in ref. (17) ; the value in benzene is adopted consistently. Am = 8m- SA. Solvent : 0 and -, methylcyclohexane ; A and - - -, carbon tetrachloride. Substituent for nitrobenzenes : (1) [2,4,6-2H31, (2) P-CH3, (3) P-CL (4) P-OCH3, (5) p-N(CH3)2. electron density induced by the association. Plots of AAB against a$ and oi l 5 (fig.4) show that AAB approximately increases with increasing o$ and oi values. The two parameter treatment of AAB given by eqn (3) is summarized in table 4 (3) AAB = acr, +boz +c. The small values of the fitting parameter show that the treatment applies satisfactorily. This result suggests both inductive and mesomeric effects for AAB. / / / / / A' 0.6 E' 0.4 4 3 a d 1 3.; A 0 0.2 0.4 ai FIG. 4.-Plots of Am against (a) 0: and (b) oi. The numbering corresponds to that in fig. 3. Solvent : 0 and -, methylcyclohexane ; A and - - -, carbon tetrachloride.84 ASS 0 CIA T I 0 N BETWEEN P 0 LA R MO I, E C UL E S TABLE 4.-TWO PARAMETER TREATMENT OF THE VALUES BY THE EQUATION Am= aai+ba,++c solvent U fa carbon methyl- tetrachloride 0.25 0.29 0.39 0.01 cyclohexane 0.22 0.16 0.48 0.04 root mean square deviation root mean square of the data' a Fitting parameter f = Three possible explanations of the association shift have been so far pointed out :6* (1) electric field effect of the partner molecule, (2) magnetic anisotropy of the partner molecule and (3) variation in the magnetic anisotropy of the intramolecular group by the association.In the present case, (3) may be excluded, since AAB is 1 -0.6 -0.3 0 0: FIG. 5.-Plots of In K against a t The numbering corresponds to that in fig. 3. Solvent : 0 and -, methylcyclohexane ; A and - - -, carbon tetrachloride. measured at the meta position of the nitro group. If (2) is a dominant factor, associ- ation shift will be governed by the relative orientation of the two molecules participat- ing in the association, and a correlation such as eqn (3) may not be realized.In this manner the low field shift of AAB and its substituent effect are explained by the TABLE 5.-TWO PARAMETER TREATMENT OF THE 1nK VALUES BY THE EQUATION h K = pai+qa,+ + r solvent seriesa P r f b carbon I 0.80 - 1.34 -0.68 0.30 tetrachloride I1 0.43 -2.73 -0.75 0.01 methyl- I 1.58 -0.87 0.44 0.17 cyclohexane I1 1.27 - 2.01 0.38 0.1 1 a Data for N,N-dimethyl-p-nitroaniline are excluded from series 11, but are included in I. root mean square deviation root mean square of the data' b Fitting parameter f =H . FUJIWARA, T . TAKABA, Y . YAMAZAKI A N D Y . SASAKI 85 electron density decrease in the phenyl ring, which is induced by the electric field effect of HMPA, and is partly compensated for by electron donation from the sub- stituent.In contrast to AAB, there is no discernible correlation between K and p, but a In K against a,+ plot suggests a linear correlation. These facts indicate a local contribution to the association, since the Hammett 0 constants are related to the polarization of the nitro group.16 As a model of the associate, an antiparallel pair of dipoles, such as those below, is consistent with the results obtained above ; a parallel pair model seems improbable, since a 31P n.m.r. study l 8 does not prove high field shift by association which is expected for the model. Further study on the structure of the associate and the substituent effect is in progress by the measurement of 3C n.m.r. spectra, including several p-substituted benzonitriles.This work was supported by a grant-in-aid for scientific research from the Ministry of Education, Japan. B. J. Bulkin, Helv. Chim. Acta, 1969, 52, 1348. C. D. Ritchie, B. A. Bier1 and R. J. Honour, J. Amer. Chem. Soc., 1962, 84,4687. A. V. Sechkarev and G. E. Trostentsova, Optics and Spectroscopy, 1973, 34, 707. W. Dannhauser and A. F. Flueckinger, J. Phys. Chem., 1964,68, 1814. T. A. Renner and M. Blander, J. Phys. Chem., 1977,81,857. T. Yonezawa and I. Morishima, Bull. Chem. SOC. Japan, 1966,39,2346. J. A. Riddick and W. B. Bunger, Organic Solvents (John Wiley, New York, 1970). ti H. Saito, Y. Tanaka, S. Nagata and K. Nukada, Canad. J. Chem., 1973,51,2118. * R. W. Taft, G. B. Klingensmith and S. Ehrenson, J. Amer. Chem. Soc., 1965, 87, 3620. lo H. Normant, Bull. SOC. chim. France, 1968, 791. l 1 F. Lagenbucher, R. Mecke and E. D. Schmid, Lieb. Ann. Chem., 1963, 669, 11. l2 Y. Sasaki, A. Takahata, M. Yoritaka, H. Kawaki and Y. Okazaki, Chem. Pharm. Bull., 1974, l 3 A. J. Dale, Acta Chem. Scand., 1970, 24, 3403. l4 I. Prins, J. W. Verhoeven and Th. J. de Boer, Org. Magnetic Resonance, 1977,9, 543. l 5 Y. Yukawa and Y. TSU~O, Nippon Kagaku Zasshi, 1965, 86, 873. 22, 50. M. Avram and GH. Mateescu, Infvared Spectroscopy Applications in Organic Chemistry, trans- lated by L. Birladeanu (Wiley, New York, 1972), p. 308. l7 A. L. McClellan, Tables of Experimental Dipole Moments (W. H. Freeman, San Francisco, 1963). '* H. Fujiwara, unpublished results. (PAPER 8/500)

ISSN:0300-9599    

DOI:10.1039/F19797500079

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Thermodynamic Properties for Transfer of Electrolytes fro m Water to Acetonitrile and to Acetonitrile+ Water Mixtures BY BRIAN G. Cox* AND RAJAGOPALAN NATARAJAN Department of Chemistry, University of Stirling, Stirling FK9 4LA, Scotland W. EARLE WAGHORNE* AND Department of Chemistry, University College, Belfield, Dublin 4, Ireland Received 22nd March, 1978 The free energies, enthalpies and entropies of transfer of a variety of electrolytes, including alkali metal halides, silver halides, tetraphenylarsonium and tetraphenylboride salts, from water to acetonitrile and water + acetonitrile mixtures have been measured. The enthalpies and entropies of transfer show a complex dependence upon solvent composition, which is discussed in terms of the effects of solvent sorting in the immediate neighbourhood of the ions, and of the effect of the solvated ions on the bulk solvent properties.The large structural effects occurring when ions are dissolved in water do not, however, appear to make a net significant contribution to the free energies of solution (and transfer) of electrolytes. Thus the overall free energies of transfer and their dependence upon solvent composition may be very simply interpreted in terms of the assumption that alkali metal and halide ions interact more strongly with water, and silver ions with acetonitrile, without reference tolany special structural properties of water. The entropies of solution of simple electrolytes in water are considerably more positive than those in a variety of dipolar aprotic s01vents.~'~ The differences, which are of the order of 40 cal K-l mol-1 (corresponding to a TAS term of some 12 kcal mol-1 at 25"C), may be ascribed primarily to effects resulting from the hydrogen bonded structure of water.Entropy losses resulting from the " freezing " of solvent molecules in the vicinity of the ions should be lower in water, and it has been suggested that the disruption of the water structure by the solvated ions may cause a significant increase in entropy in a region further from the ion?* It is not clear, however, to what extent these " structural " effects contribute to the free energies of solution of the ions, because of the tendency of the resulting enthalpy and entropy changes to compensate for one another. Indeed, much available evidence suggests that changes in structure of aqueous solutions frequently occur with little change in free energy.'-l We have recently studied the changes in thermodynamic properties accompanying the transfer of electrolytes (Ag+, Na+ and Li+ salts) between two dipolar aprotic solvents, dimethylsulphoxide and propylene carbonate, and their mixtures.The observed free energies (AG;), enthalpies (AHtr) and entropies (Astor) of transfer of the cations can be very simply interpreted in terms of the interactions between the ions and their immediate neighbour solvent molecules. In particular, a sharp minimum in ASFr occurs in the mixtures, resulting from large differences in the solvent composition of the coordination sphere of the ions, and of the bulk solvent. In the present paper, we report a comprehensive study of AG,", AH,", and ASFr for transfer of a variety of electrolytes from water to acetonitrile and their mixtures.86B. G. COX, R. NATARAJAN AND W. E. WAGHORNE 87 The results are discussed with reference to those obtained in mixtures of dipolar aprotic solvents. Acetonitrile +water mixtures are particularly interesting in relation to studies of ion-solvent interactions, because of the contrasting behaviour of alkali metal cations, which interact more strongly with water, and silver (and cuprous) cations which are more strongly solvated by acetonitrile.' EXPERIMENTAL CHEMICALS Acetonitrile was purified by successive distillation from phosphorus pentoxide and calcium hydride, after initial drying with calcium hydride.13 Alkali metal salts were Specpure grade (Johnson Mathey).They were dried at 100°C under vacuum prior to use. AnalaR silver nitrate was used without further purification. Bu4NBr (Cambrian Chemicals) was dried under vacuum. Ph4AsI was prepared from Ph4AsC1 and KI. It was purified by recrystallization from water. KBPh4 was prepared from KCI and NaBPh4, and purified by recrystallization from acetone+ H20. NaBPh4 (B.D.H.) was purified by recrystallization from acetone+ toluene as previously described.14 Ph4AsC1 (Aldrich) was purified by recrystallization from dichloromethane, and dried at 100°C under vacuum. This was necessary to prevent contamination of the sample with Ph4AsCl 2H20. FREE ENERGIES OF TRANSFER SOLUBILITY MEASUREMENTS The standard free energies of transfer AGPr(MX),from water to solvent S of AgBr, AgI, P u s 1 and KBPh4 were calculated from eqn (l), where Ksp is the solubility product referred to infinite dilution in the appropriate solvent (S or H20) Ksp values for AgCl are available in the 1iterat~re.l~ Ksp for silver halides were determined from potentiometric titrations of AgN03 into standard tetra-alkylammonium halide solutions, AGZ(:(MX) = RT In [&pCH20)IKsp(~)I.(1) TABLE 1 .-SOLUBILITIES OF ELECTROLYTES IN ACETONITRILE+ WATER MIXTURES AT 25°C PKS a JAN^ AgBr &I Ph&I KBPh4 0 12.21 16.03 5.26 7.34= 0.10 11.i6 14.8, 4.4, 0.20 10.84 14.2, 3.60 6.12 0.40 10.27 13.69 2.22 4.61 0.60 10.2~ 13.41 1.67 3.10 0.80 10.78 13.5, 1.53 2.74 0.90 11.51 14.00 1.00 13.90f 14.49 g 2.69' 3-11 a Solubility products corrected to zero ionic strength ; estimated uncertainty kO.1 log unit ; b volume fraction of acetonitrile; Cref.(4); d c f : pKs = 5.1, ref. (12); e c f . pK, = 7.5, ref. (12); fref. (23) ; g ref. (12) ; h cf. pK, = 2.72 ref. (23) ; 1 cf. pKs = 3.24 ref. (23). The difference arises from the activity coefficients used to correct to zero iomc strength. the titrations being monitored with an Ag/Ag+ electrode. KBPh4 and PLAsI solubilities were determined by analysing saturated solutions of the salts for K+ by atomic absorption spectroscopy, and by titration against standard AgN03 solutions to determine I-, respectively. The solubility products were corrected to infinite dilution using activity coefficients (yk) calculated from the Davies eqn (2),16 where A is the Debye-Hiickel parameter88 ELECTROLYTES I N WATER AND ACETONITRILE which depends upon the solvent dielectric constant and temperature, and I is the ionic strength.Values of A may be calculated from the known dielectric constants of acetonitrile+ water mixture~.'~ The results are listed in table 1. ELECTROCHEMICAL MEASUREMENTS Free energies of transfer of alkali metal halides, MX, were determined from e.m.f. measurements on cell A.18s21 In cell A, the glass electrode was either a sodium ion-selective electrode (Activation) or a general cationic Ag-AgX 1 M x ~ ~ ~ ~ ~ = m, I glass electrode electrode (Beckman 39137). Earlier studies have shown that such electrodes give results in good agreement with those of conventional amalgam electrodes,' 8-20 the results for Cs+, however, being generally less reliable than for the other cations.Ag-AgX electrodes were prepared. by electrolysing silver electrodes in MX solutions.21 The free energy of transfer of MX from water to solvent S is given by eqn (3),22 where I& refers to the e.m.f. of cell A, (3) with electrolyte MX of concentration ms, and similarly for EH20. In eqn (3, y; represen s the mean ionic activity coefficient of electrolyte MX referred to infinite dilution in solvent S. The activity coefficients were calculated from eqn (2). Electrolyte concentrations in cell A varied between and mol dm-3. AGS(MX) = F(Es-E~20)+2RTln (mHZ0/mS)+2RTln (ypo/y$) TABLE 2.-FREE ENERGIES OF TRANSFER (AGt",)' OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C A AN 0 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 0.95 1 .oo LiCl 0.00 0.30 0.69 3.12 4.28 7.4 10.4 17.2f AGtJkcal mol-1 b (molar scale) NaCl NaBr KCl KBr KI CsCl AgBreld 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.27 0.l9 0.l9 0.09 0.03 0.l9 -1.40 0.51 0.39 0.45 0.29 0.05 0.37 -1.92 0.82 0.68 0.73 0.52 0.19 0.61 -2.40 1.24 1.00 1.13 0.82 0.30 0.89 -2.70 1.68 1.38 1.59 1.19 0.52 1.18 -2.81 2.84 2.39 2.68 2-20 1.25 2.21 -2.40 3.94 3-22 3.59 3.06 1.94 3.48 -2.00 7.02 5.49 6.28 5.57 3.19 6.08 -1.00 9.74 7.79 8.68 7.28 4.20 8.41 0.30 2.23 1.81 2.09 l.60 0.81 l.60 -2.65 13.4g 17.0h 12.0i 9.2g 6.4h 11.6g 2.3g KBPh4 c*e 0.00 -0.8 - 1.8 - 2.9 - 4.0 - 5.2 - 6.0 - 6.4 - 6.5 - 6.4 - 6.2 - 6.0 g Ph&I '1' 0.00 - 0.9 - 2.3 - 3.4 - 4.0 - 4.4 - 4.7 - 4.9 - 4.9 - 4.4 - 3.9 -3.3 i a Estimated error i- 0.05 kcal mol-l unless otherwise stated ; 1 cal = 4.184 J ; C from table 1 ; d estimated error & 0.1 kcal mol-' ; e estimated error rt 0.2 kcal mol-1 ; f M.Solomon, J. Phys. Chem., 1970,74, 2519 ; g ref. (23) ; h J. Pavlopoulous and H. Strehlow, 2. phys. Chem., 1954, 202, 474 ; i ref. (12). Cell A was not used for solvent mixtures containing > -85 vol % acetonitrile, be- cause of the tendency of AgX to dissolve (giving AgX,) when the anion activity becomes too high. However, AG,", values for transfer to pure acetonitrile have been previously determined from solubility meas~rements,~~ and AGP, values for the region between 85 vol % and pure acetonitrile were obtained by interpolation. AGP, values for Na+ salts have been shown to vary smoothly and monotonically over this region, using AG,", ma+- Ag+) values obtained from measurements on cell B, together with values of AG: (AgX) obtained from solubility measurement~.~~~ 24 The combined results, listed at selected volume fractions of acetonitrile are given in table 2.B .G . COX, R . NATARAJAN A N D W . E. WAGHORNE 89 In determining AGg values from solubility measurements [eqn (l)] and electrochemical measurements [eqn (3)], it is assumed that all of the salts are fully dissociated at the con- centrations used. The only measurements carried out in pure acetonitrile, where ion-pair formation should be strongest, involved solutions of tetra-alkylammonium, tetraphenylboride or tetraphenylarsonium salts, which are known to be strong electrolytes in acetonitrile.l2? 23 All other measurements were carried out in solvent mixtures containing at least 20vol % water, and at electrolyte concentrations < mol drr3. Under these conditions complete dissociation of the salts may be safely assumed.l2. 23 TABLE 3.-ENTHALPIES OF TRANSFER OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C (kcal mol-l)b $AN 0 LiCl d 0 0.0 0.10 0.20 -0.5 0.30 0.40 -0.9 0.50 0.60 -1.8 0.70 0.80 -2.5 0.90 -4.6 0.95 -7.0 1.00 +4.9 0 0.0 0.10 -0.4 0.20 -1.0 0.30 -1.6 0.40 -1.9 0.50 -2.4 0.60 -3.0 0.80 -4.6 0.90 -6.6 1.00 -1.3e & A X C CScI 0.70 0.95 NaCl 0.0 - 0.5 -0.8 - 1.0 - 1.5 - 2.2 - 3.0 -4.5 + 1.6f AgCl 0.0 - 5.9 - 7.4 - 8.5 -9.1 - 11.0 - 11.4 - 7.9 g NaBr 0.0 - 0.3 -0.6 - 1.0 - 1.4 - 1.9 - 2.4 - 2.9 - 3.6 - 5.2 - 5.6 - 0.9 f 0.0 - 6.0 - 8.0 -9.1 - 10.1 - 10.6 -11.2 -11.9 - 10.9 g AgBr d KCI 0.0 - 0.9 - 1.4 - 1.6 -2.1 - 2.7 - 3.2 - 4.0 - 4.9 - 0.7 f AgI 0.0 - 7.0 - 8.6 - 9.3 - 9.5 - 10.7 - 11.7 - 12.5 - 13.7 - 14.6 KBr 0.0 - 1.3 - 1.9 - 2.2 -2.8 - 3.3 -4.1 - 4.6 - 5.3 - 3.4f 0.0 3.6 6.5 6.4 6.0 5.9 BudNBr 6.4 g KI RbCl 0.0 0.0 - 1.4 - 1.1 -2.8 - 1.8 - 4.2 - 4.7 - 5.7 -7.1 -7.1 f Ph4AsCl 0.0 4.7 5.7 4.9 2.5 - 2.8 - 4.0 - 5.8 -4.8 -1.5f NaBPhdd 0.0 4.2 2.1 0.5 - 2.0 1.3 - 5.0 0.8 - 6.0 1.9 - 5.6 g a Estimated uncertainty +_ 0.1 kcal mol-' unless otherwise stated ; enthalpies of solution in water :25 LiCI, - 8.85 ; NaCl -0.14 ; NaBr -0.14 ; KC1 4.16 ; KBr 4.75 ; KI 4.86 ; RbCl4.13 ; CsC14.25 ; AgCI4 15.7 ; AgBr4 20.2 ; Ag14 26.9 ; Bu4NBr4 -2.2 ; Ph4AsC14 - 2.6 ; NaBPh, -4.8 ; b 1 cal = 4.184 J ; C volume fraction of acetonitrile ; destimated uncertainty k 0.5 kcal mol-I ; e Y .4 . Choi and C. M. Criss, J. Chem. Eng. Data, 1977, 22, 297 and ref. (4) ; fref. (3) and (4) ; g ref. (4). ENTHALPIES OF TRANSFER The enthalpies of solution, AHs, (or precipitation in the case of silver halides) of electro- lytes in acetonitrile+ water mixtures were determined by standard calorimetric techniques using a Tronac 450-4 calorimeter. Electrolyte concentrations were in the range w mol dm-3. The values have not been extrapolated to infinite dilution, as the variations with concentration observed within this range were negligible within experimental error. Thus values are taken to be standard state values. Except for the silver halides, HCl and NaBPh4, AH: values were in the range - 5 kcal mol-1 < AH: < 5 kcaI mol-1 and were reproducible to +O.l kcal mol-l.In water, AH? values agreed to within kO.1 kcal mol-1 of literature values.25 Values of AH: for the Ag halides, HCI and NaBPh4 varied between - 5 and 27 kcal rnol-l, and consequently the errors in AH: values for these electrolytes may be as high as k0.5 kcal rno1-l. The collected values are given in table 3.90 ELECTROLYTES IN WATER AND ACETONITRILE ENTROPIES OF TRANSFER Standard entropies of transfer, AS:, have been obtained by the application of eqn (4) to the data in tables 2 and 3. The values are listed in table 4. To facilitate comparisons AG;' = AH,", - TAS,O, (4) of the relative effects of AH" and AS" in determining the total free energy changes, they are quoted as values of 298ASP, (corresponding to 25°C).TABLE 4.-ENTROPIES OF TRANSFER (AS;)' OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C values of 298 AS,,/kcal mol-I (molar scale) AN C LiCl 0 0.0 0.10 -0.5 0.20 -0.9 0.30 -1.6 0.40 -2.3 0.50 -3.0 0.60 -4.1 0.70 -5.5 0.80 -7.6 0.90 -12.1 0.95 -16.3 1.00 -12.3 +AN c RbCl 0 0.0 0.10 -0.6 0.20 -1.4 0.30 -2.2 0.40 -3.0 0.50 -3.9 0.60 -4.9 0.70 -6.1 0.80 -8.0 0.90 -12.1 0.95 -13.4 1.00 -13.2 NaCl 0.0 - 0.5 - 1.0 - 1.6 - 2.2 - 2.9 - 3.7 - 5.0 - 7.0 -11.5 - 14.7 -11.8 CSCl 0.0 - 0.5 - 1.3 -2.1 - 2.8 - 3.6 - 4.6 - 5.8 -8.1 - 12.7 - 14.9 - 12.9 NaBr 0.0 - 0.5 - 1.0 - 1.7 - 2.4 - 3.2 - 4.2 - 5.3 - 6.8 - 10.6 - 13.4 - 11.9 AgCl'* f 0.0 - 4.6 - 5.6 - 5.9 - 6.1 - 6.5 - 7.2 - 8.2 - 9.6 - 12.0 - 13.5 - 12.8 KCl 0.0 -0.6 - 1.3 - 2.0 - 2.8 - 3.7 - 4.8 - 5.9 - 7.5 -11.2 - 13.9 - 12.7 AgBr f 0.0 - 4.5 - 5.9 - 6.3 - 6.5 - 6.9 - 7.5 - 8.2 - 9.2 - 10.9 - 12.0 - 13.2 KBr 0.0 - 0.6 - 1.5 - 2.4 - 3.1 - 4.0 - 5.0 - 6.2 - 7.5 - 10.9 - 13.0 - 12.6 &I f 0.0 - 5.2 - 6.0 - 6.4 - 6.7 - 6.7 - 7.7 - 8.5 - 9.4 -11.3 - 11.9 - 12.6 KI 0.0 - 0.8 - 1.5 - 2.4 - 3.1 - 4.0 - 4.9 - 6.0 - 7.7 - 10.3 - 11.8 - 13.6 Bu4NBr 0.0 4.1 6.8 7.9 8.6 9.2 9.7 9.8 9.4 8.8 aValues obtained by application of eqn (4) to the data in tables 2 and 3 ; estimated error 1 cal = 4.184 J ; C volume fraction of acetonitrile ; dAGtr values from ref.(19) ; e AGpr values from ref. (15), (24); festimated uncertainty t0.5 kcal mol-' ; B ref. (26). 0.2 kcal mol-1 unless otherwise stated ; DISCUSSION FREE ENERGIES OF TRANSFER The results in table 2 show that, with the exception of AgI, all of the simple electrolytes are more soluble in water than in acetonitrile.KBPh4, Ph4AsI and B u , N B ~ , ~ ~ on the other hand, are all more soluble in acetonitrile. More informative, however, is the way in which the free energies vary with solvent composition in the mixtures.15* 27-29 Thus if an ion is strongly solvated by one of the solvent com-B . G . COX, R . NATARAJAN A N D W . E . WAGHORNE 91 ponents, addition of the second, more weakly interacting, component will have little effect on its free energy until the last few percent of the " better '' solvent is removed, when the free energy will increase rapidly. Conversely, addition of small amounts of a "good'' solvent to an ion in a poorly solvating medium will result in a sharp decrease in its free energy.Strong solvent-solvent interactions may complicate the situation, but water and acetonitrile do not interact strongly with one another (their mixtures show positive deviations from Raoult's law).30 The results in table 2 can be readily understood on this basis, if it is assumed that the alkali metal cations and halide anions are more strongly solvated by water, and Ag+ and the " organic " ions, Ph,As+, BPh; and Bu,N+, more strongly solvated by acetonitrile. I . I GI- AgCl Na+ t ao 0.2 0.4 0.6 0.8 1.0 ~ C H ~ C N FIG. 1.-Free energies of transfer of NaCl and AgCl from water to acetonitrile+water mixtures at 25°C. This behaviour may be clearly seen in fig. 1, in which the free energies of transfer of NaCl and AgCl are plotted against the volume fraction of acetonitrile.Also included are the single ion values for Ag+, Na+ and C1-, based on the assumption AG,",(Ph,As+) = AG,"(BPh;) applied to the results in table 2. The free energy of NaCl increases slowly with increasing acetonitrile content of the solvent, until = 80 vol % acetonitrile, beyond which it increases rapidly to its value in pure acetonitrile. The free energy of AgCl, on the other hand, passes through a minimum. This is typical of the behaviour expected for a binary electrolyte, of which one ion is bound more strongly by one solvent compound and the other ion prefers the second component.2 9 s 31 We have shown earlier that in mixtures of dipolar aprotic solvents, where specific solvent-solvent interactions are small, free energies of transfer ions can be very satisfactorily accounted for in terms of a simple model in which the solvents are regarded as competing ligands for coordination sites of the ions.11* 15* 2 8 Excellent agreement is obtained between experimentally estimated Act", values, and values calculated from measured stability constants for coordination of the ions with the solvents involved.The present results are clearly in qualitative agreement with such a concept, although quantitative calculations are difficult because of the lack of stability constant data, and the non ideality of the solvent mixtures. However, it is known from potentiometric titrations 32* 33 that the complex Ag(CH,CN)Z is quite stable in aqueous solutions, and this, together with the absence of similar complexes with alkali metal cations, supports the above interpretation.92 ELECTROLYTES IN WATER A N D ACETONITRILE ENTHALPIES A N D ENTROPIES The relatively simple behaviour of the free energies of transfer is clearly not reflected in the enthalpies and entropies of transfer (tables 3 and 4).The transfer of both silver and alkali metal halides to acetonitrile results in very large decreases in entropy, the total entropy loss (TASP, = - 12.5 & 1.0 kcal mol-l) being relatively independent of the electrolyte involved. The variation of the entropy with solvent composition is, however, very different for silver and alkali metal salts. This may be seen in fig. 2, where AH; and -TAS,", for NaCl and AgCl are plotted against the volume fraction of acetonitrile.1 0 : : $1 a 0.0 0.2 0.4 0.6 0.8 1.0 k H X N FIG. 2.-Enthapies and entropies of transfer of NaCl(0) and AgCl (+) from water to acetonitrile+ water mixtures at 25°C. In discussing the entropy results, it is useful to consider firstly the effects expected in the immediate neighbourhood, or coordination sphere, of the ions. In all cases, there will be a loss in entropy on dissolving the electrolyte, resulting from the loss of translational entropy of the solvent molecules. More important in the present context, is that in a mixed solvent, where the ions are preferentially solvated by one of the components, there will be an unfavourable entropy term (relative to the ions in either of the pure solvents) because the composition of the coordination sphere will be different from that of the bulk solvent mixture, i.e.there will be an entropy decrease resulting from solvent sorting. If the composition of the solvation sphere of the ion is known, e.g. from n.m.r. measurements 29 or calculated from known stability constants for complex formation,ll' 15* 28 then the entropy losses may be readily calculated from simple statistical thermod ynamics.l We have shown, for example, that the entropy of transfer of Ag+ salts from propylene carbonate to dimethyl- sulphoxide (AG& large, -ve, AStf z 0) passes through a sharp minimum at 3 vol % dimethylsulphoxide (with ASFr falling to z - 20 cal K-l mol-1 or 29SAS,O, = - 6 kcal mol-l), directly attributable to solvent sorting around Ag+.l The effect is discussed in detail elsewhere, but model calculations show that as the strength of interactionB .G . COX, R . NATARAJAN AND W . E . WAGHORNE 93 with the more strongly solvating component increases, the minimum in AS& increases in magnitude, and moves to lower fractions of this component.ll In the present system, the free energy data (table 2 and fig. 1) suggest that Agf will be preferentially coordinated to acetonitrile and the alkali metal and halide ions will be preferentially coordinated to water. This should result in a large loss in entropy (increase in -TAS,") at low volume fractions of acetonitrile for Ag+ salts and at high fractions of acetonitrile (low water content) for the other ions. This is shown schematically in fig.3, and such an effect can clearly be seen in the results in fig. 1, 2 and table 4. Single ion values, obtained as before from the Ph4Ag+, BPh; assumption, suggest that these entropy losses are comparable for the alkali metal and chloride ions, but much smaller for bromide and iodide ions. 0.0 0.2 0.4 0.6 0.8 1.0 k H C N FIG. 3.-Entropy changes resulting from preferential solvation of ions in acetonitrile + water mixtures. Superimposed on the minima in the entropies resulting from solvent sorting, is the large net decrease in entropy accompanying the transfer of electrolytes from water to pure acetonitrile, which may be associated with bulk solvent properties. The more extensive structure of water compared with acetonitrile means that the loss in translational entropy on coordination to an ion will be smaller for water molecules than acetonitrile (cf.the considerably lower entropy losses on freezing of water compared with dipolar aprotic In addition, more positive entropies may result from the disruption of the water structure by the hydrated ions, whose bound water molecules are not correctly orientated to fit into the hydrogen bonded structure of ti What is perhaps surprising, however, is that the effect is important even at high fractions of acetonitrile. In fact, at least half of the observed total increase in entropy on transfer from acetonitrile to water has occurred upon addition of only 20 vol % of water, where the average distance apart of the water molecules is five times that in pure water. This may be seen for example from the results for KI, where solvent sorting effects should be smaller, or from the results for the alkali metal chlorides and bromides, when reasonable allowance is made for entropy losses arising from solvent sorting.This suggests that long range order in water is not of major importance in determining entropy effects. Rather, similar effects may result from considerably smaller, but presumably still strongly hydrogen bonded, groups of water molecules. The entropies of transfer of Bu,NBr and Ph4As+ and BPh; salts (values for which may be obtained by appropriate combinations from the results in tables 2 and 3) differ from the alkali metal salts in two ways : first transfer to acetonitrile leads to a considerable increase in entropy, and secondly the major change occurs at low fractions of acetonitrile.This behaviour is very similar to that of neutral organic94 ELECTROLYTES I N WATER AND ACETONITRILE molecule^,^ and is consistent with the notion that the addition of such molecules increases the hydrogen bonded structure of water.' The ions are presumably preferentially solvated by acetonitrile in the mixtures, so that the effect will disappear after the addition of relatively small amounts of acetonitrile. The enthalpies of transfer may be discussed in an analogous way to the entropies. In mixtures of dipolar aprotic solvents, these are determined almost entirely by the relative strengths of the interactions with the two solvent cornp0nents.l This effect would also account for the rapid changes in AH,, of the alkali metal halides at low water concentrations, and for the silver halides at low acetonitrile concentrations. It is also noticeable that diflerences between the heats of transfer of pairs of alkali metal and silver halides from water to acetonitrile agree closely with differences between the corresponding free energies of transfer, e.g. AG,,(NaCl) - AG,,(AgCl) = 8.9 kcal mol-1 ; AHt,(NaC1) -AHt,(AgC1) = 9.5 kcal mol-l.The absolute enthalpies of transfer, however, bear no simple relationship to the free energies of transfer. This is not surprising, as the entropy results show that large effects associated, with solvent reorganisation are involved, and these must have corresponding enthalpy effects, e.g., disruption of the water structure leads to an increase in entropy, but also to an unfavourable increase in enthalpy, because of the breaking of hydrogen bonds.The apparently simple behaviour of the free energies of transfer of electrolytes, strongly suggests that the contributions of solvent structural effects to ASFr and AH; terms largely cancel, thus making little overall contribution to the free energies of the ions. In conclusion, the enthalpies and entropies of transfer of electrolytes from water to acetonitrile show a complex dependence upon the solvent composition, reflecting the interactions between the ions and the neighbouring solvent molecules, as well as the effect of the solvated ions on the bulk solvent properties. The free energies of transfer on the other hand may be very simply interpreted in terms of the assumption that alkali metal halide ions interact more strongly with water molecules, and silver ions with acetonitrile, without reference to any special structural properties of water.The large (in terms of entropy and enthalpy) structural effects occurring when ions (and non electrolytes) are dissolved in water do not make a significant net contribution to the free energies of solution. We thank the S.R.C. for a research grant. C. M. Criss and E. Held, J. Phys. Chem., 1968,72, 2966. B. G. Cox and A. J. Parker, J. Amer. Chem. SOC., 1973,95,6879. M. H. Abraham, J.C.S. Farday I, 1973, 69, 1375. H. S. Franck and M. W. Evans, J. Chem. Phys., 1945,13,507. H. S . Franck and W. Y. Yen, Disc. Faraday Soc., 1957,24, 133. 4B. G. Cox, Ann, Rep. Chem.SOC. A, 1973,249. ' W. P. Jencks, Catalysis in Chemistry andEnzymo2og.y (McGraw-Hill, New York, 1969), chap. 8. * E. M. Arnett and D. R. McKelvey, Rec. Chem. Progr., 1965,26, 185. B. G. Cox, J.C.S. Perkin 11, 1973, 607. B. G. Cox, W. E. Waghorne, and C. K. Pigott, J.C.S. Faraday I, 1979, 75, 227. 94, 1148. 1962,34,1139. lo M. Ruseman and W. P. Jencks, J. Amer. Chem. Sac., 1975, 97, 631. l2 R. Alexander, A. J. Parker, J. H. Sharp and W. E. Waghorne, J. Amer. Chem. SOC., 1972, l3 J. F. Coetzee, G. P. Cunningham, D. K. McGuire and G. R. Podmanashan, Analyt. Chem , l4 B. G. Cox, G. R. Hedwig, A. J. Parker and D. W. Watts, Austral. J. Chem., 1974, 27, 477. l5 B. G. Cox, A. J. Parker and W. E. Waghorne, J. Phys. Chem., 1974,78,1731. l6 C . W. Davies, Ion Association (Butterworth, London, 1962), eqn (3.14). l7 C. Moreau and G. Douheret, J . Chern. Thermodynamics, 1976, 8,403.B . G. COX, R . NATARAJAN AND W . E. WAGHORNE 95 l 8 Y. Pointud, J. Juillard, J.-P. Morel and L. Avedikian, Electrochim. Actu, 1974, 19, 229. l9 R. Smits, D. L. Massart, J. Juillard and J.-P. Morel, Electrochim. Actu, 1976, 21, 425, 431 2o T. A. Clune, D. Feakins and P. J. McCarthy, J. Electronulyt. Chem., 1977, 84, 199. 21 G. J. Janz, in Reference Electrodes, ed. D. G. J. Ives and G. J. Janz (Academic Press, New York, 22 H. P. Benetto, D. Feakins and K. G. Lawrence, J. Chem. SOC. A, 1968,1493. 23 I. M. Kolthoff and M. K. Chantooni, J. Phys. Chem., 1972,76,2024. 24 W. E. Waghorne, PhB. Thesis (Australian National University, 1972). 2 5 V. B. Parker, Thermal Properties of Aqueous Univalent Electrolytes (National Bureau of 26 C. Treiner, P. Tzias and M. Chemla, J.C.S. Faruduy I, 1976, 72,2007. 27 A. K. Covington and A. D. Covington, J.C.S. Furuduy I, 1975,71, 831, and references therein. 28 G. Clune, W. E. Waghorne and B. G. Cox, J.C.S. Furuduy I, 1976,72,1294. 29 H. Schneider, Electrochim. Actu, 1976, 21, 711. 30 A. M. von Vierk, 2. unorg. Chem., 1950,261,283. 31 H. Strehlow and H. Schneider, J. Chim. phys. (Numero Special), 1969,118. 32 H. Strehlow and H.-M. Koepp, 2. Electrochem. Ber. Bunsenges. phys. Chem., 1958, 62, 373. 33 S. E. Manahan and R. T. Iwamoto, J. Electrounulyt. Chem., 1967, 14,213. 34 Selected Values of Chemical iVzermodymmic Properties (National Bureau of Standards Circular and 437. 1961), p. 179. Standards, NSRDS-NBS 2, Washington, D.C., 1965). 500). (PAPER 8/546)

ISSN:0300-9599    

DOI:10.1039/F19797500086

出版商:RSC    

年代:1979

数据来源: RSC

摘要:

Infrared Study of CO Adsorption on Magnesium Oxide BY EUGENIO GUGLIELMINOTTI, SALVATORE COLUCCIA, EDOARDO GARRONE, Istituto di Chimica Fisica dell'Universit8 di Torino, Corso M. d'Azeglio 48, 10125 Torino, Italy LUIGI CERRUTI AND ADRIANO ZECCHINA * Received 11th April, 1978 The adsorption of CO at room temperature on well outgassed specimens of MgO gives rise to a large number of bands in the 2200-1000 cm-' range, which can be divided into two main groups. The bands of the former group are destroyed by oxygen at room temperature : some react instan- taneously and are associated with a marked pink colour of the sample ; the others are less reactive as they require prolonged contact time in order to be completely oxidized at room temperature (r.t.). The bands of the latter group, far from being destroyed by oxygen, grow when the oxygen-sensitive species are depleted.The oxygen-sensitive species are thought to be negatively charged polymeric CO structures (CO clusters) of the type (CO)",-, where x = 2 or 4 and n is > 2. The simplest CO clusters (dimers) can be transformed into larger polymers by further CO addition. Under the correct conditions the reverse process can also be carried out. The oxygen-insensitive species have a car- bonate-like structure and are present on the surface in fairly constant ratios with respect to the former group species. A chemisorption mechanism leading both to oxidized (carbonate-like) and to reduced (negative CO polymers) species is proposed. The active centres for CO chemisorption consist of groups of ions (both positive and negative) in strongly uncoordinated situations.1.r. studies on the CO/MgO system have been carried out by Little et a2.l and by Jiru et aL2 Both studies deal with samples outgassed at moderately low temperatures, i.e., when still covered by a fair number of OH groups. Under these conditions, CO is adsorbed as carbonate-like species only if 0, is present. In contrast, other studies carried out by u.v.-vis. reflectance spectroscopy on completely dehydrated samples 9 have shown that CO chemisorption gives rise to peculiar polymeric species character- ized by electronic transitions in the visible and near ultraviolet. The aim of this investigation is to obtain further information on the structures revealed by reflectance spectroscopy.It will be shown that on the basis of vibrational spectra alone the presence of polymeric CO species is confirmed, in agreement with the preliminary results discussed in ref. (5). EXPERIMENTAL The samples were prepared by decomposing magnesium hydroxide pellets under vacuum N m-2) in the i.r. cell at 523 K. Two hydroxide samples of different purity were used. The purest one was the same as that used in the reflectance study,4 whereas the other was prepared following Lunsford and Jayne from Carlo Erba Reagent Grade ACS MgO (Fe impurity 50 p.p.m.). In both cases, specimens were obtained with indistinguishable pro- perties (B.E.T. specific surface area - 250 m' g-'), which do not sinter appreciably in vacuo even under the severe thermal treatments at 1073 K required for the complete surface dehydration.Pellets of % 30mgcm-2 were used. The i.r. spectra were measured on a Beckman I.R. 12 spectrometer using a standard in situ silica cell. High purity (Matheson) CO and O2 were used. On 1073 K outgassed samples the CO coverage, estimated following the procedure described in ref. (6), was 0.01-0.015. 96GUG LIELMI N OTTI, COLU C CIA, GARRONE, CERRUTI , Z E C C HI N A 97 RESULTS EFFECTS OF PRETREATMENT TEMPERATURE The MgO sample formed by decomposition of a magnesium hydroxide pellet was outgassed at 773, 923, 1073 and 1173 K for 4 h. After each outgassing step, the pellet was cooled to r.t. and contacted with 13.3 kN m-2 CO. The sample outgassed at 773 K shows no (or very limited) activity, in agreement with the 1iterature.l For higher pretreatment temperatures, the same number of bands with constant relative intensities are observed but the overall activity increases up we 1073 K, then declines at 1173 K.We ascribe the increase in activity to the increasing surface dehydration and the decrease to sintering of the particles. For the above reasons, only the spectra taken after outgassing at 1073 K will be reported and discussed. co ADSORPTION ON FULLY DEHYDRATED SAMPLE Due both to the exceptional complexity of the spectra and to the fact that the correlations among the bands are only established on the basis of the whole set of spectra reported, we are unable to make the usual detailed description of the results and we anticipate, for the sake of clarity, the correlation table 1 of the bands.The following discussion will, hopefully, justify the correlations proposed. TABLE 1 .-ADSORBED SPECIES number of observed frequencies/cm-1 species modes features A 1 2200 weak and reversible B 2 2064, 13 18 transient species, oxygen sensitive 1480, 1275 C 4 desorbed first 1197, 1066 grow slowly upon ads., very sensitive to oxygen D1 2108, 1358 D1 grows upon C depletion D2 and D3 grow upon ads., 2097,1365 D{D2 2 resistant to desorption D3 2084, 1392 all oxygen sensitive 1582, 1160 1548, 1160 2 1574, 1160 as D1 species F see text oxygen insensitive, not studied in detail In fig. 1 the i.r. spectrum of adsorbed CO is shown at increasing coverages: the dotted curve was obtained by contacting the surface with 650 N m-2 CO for 15 min, while the broken one refers to the same sample contacted for the same time with 2.6 kN m-2 CO.The solid curve is the spectrum of the sample after 15 h at 13.3 kN rn-2. The last spectrum does not yet correspond to the maximum coverage, as small but definite intensity increases are observed for longer contact times. This shows that CO chemisorption is a slow and complex process, strongly pressure and time dependent. High CO pressures considerably increase the rate of adsorption; I 498 CO ADSORPTION ON MgO however the same situation can be reached with lower pressures and longer contact times. Fig. 1 shows that : (1) a weak band at 2200 cm-l is formed immediately upon CO contact with time-independent intensity. This unusual behaviour enables it to be assigned to a single species A.(2) The bands located at 2064 and 1318 cm-' FIG. 1.-Spectra of CO adsorbed at increasing coverages. . . ., 650 N m-' for 15 min; - - -, 2.6 kN m-2 for 15 min ; -, background ; 13.3 kN m-2 for 15 h (nearly maximum coverage). 1 OOy 3/cm-1 FIG. 2.-CO desorption up to 373 K. -, nearly maximum coverage (as last curve in fig. 1) ; , . ., sample outgassed 5 min at beam temperature ; - - -, sample outgassed 5 min at 373 K. The bands due to C species are shadowed.GUGLIELMINOTTI, COLUCCIA, GARRONE, CERRUTI, ZECCHINA 99 initially grow in intensity and then disappear at high coverages. As this behaviour is unique we ascribe them to a single species B. (3) The bands at 2097, 2084, 1392 and 1365-72cm-' (D species) appear immediately (dotted curve), whereas those at 1480,1275, 1197 and 1066 cm-' (C species) require longer contact times and/or higher pressures.(4) The apparent maxima of the bands of the C species shift with the coverage to higher frequencies (Av = + 8 cm-I). Both this fact and their large half width suggest that a family of species with very similar structure absorb at the quoted frequencies. co DESORPTION AND READSORPTION In fig. 2 and 3 the desorption of CO from the same sample is illustrated. Due to complexity of the spectrum, two figures are needed for a complete description of the experiment. In fig. 2 the modifications caused by outgassing in the 310-373 K range are shown, whereas in fig. 3 the remaining part of the experiment is illustrated (tem- perature range 423-473 K). 0 Vlcm-' FIG. 3.-CO desorption in the range 373-473 K.- - -, after outgassing 5 min at 373 K (as last curve in fig. 2) ; - - * -, sample outgassed 15 min at 423 K ; -, sample outgassed 30 min at 473 K. The solid curve in fig. 2 is the starting spectrum, and actually coincides with the last spectrum in fig. 1. Outgassing for 5 min at the beam temperature (b.t.) (fig. 2 dotted curve) decreases the intensity of the bands at 1480, 1275, 1197 and 1066 cm-1 and increases that of the bands at 2108, 1574,1548,1358 and 1160 cm-I. Moreover, the 2200 cm-I band completely disappears. The relative intensities of the former group of bands are not basically modified, so confirming that they are due to the same type of surface species C . The lability of C species upon outgassing is surprising, as fig. 1 has shown that their formation is definitely activated.This fact indicates that the depletion pathway is different from the formation one. In fact, new species are created on the surface.100 CO ADSORPTION ON MgO The classification of the bands growing upon outgassing at b.t. is not easy. Some bands are already present in the initial spectrum (1574, 1548 and 1160 cm-l), whereas the others are nearly new (2108 and 1358cm-l). Thus, different surface species cause these absorptions. In particular, the bands at 2108 and 1358 cm-l (clearly coupled together) are both ascribed to species D1, whereas the other bands belong to another family (E species). Evacuation at 373 K (fig. 2 dashed curve) leads to a further dramatic decrease in the C species, while the D1 and E bands reach their maximum intensity.The residual components of the C species are now shifted to higher frequency, so confirming the composite nature of these bands. By outgassing in the 310-373 K range the intensity of a large number of bands does not change at all. They are ascribed to other less reactive F species. 7 S/cm-l FIG. 4.-CO readsorption after outgassing at 423 K. - * - a, spectrum of a different sample after outgassing 15 min at 423 K ; . . ., 13.3 k N m-2 CO at beam temperature. Successive evacuation at 423 K (fig. 3, dotted curve) causes the disappearance of a couple of weak bands at 2084 and 1392 cm-l characteristic of the D3 species, while the bands of the F species now change in intensity in a complex way. The exposure of the sample to 13.3 kN m-2 CO after each outgassing step de- scribed so far restores the full coverage spectrum. This is shown in fig.4 after desorption at 423 K, the dotted curve being the final spectrum. The C species are regenerated at the expense of D1 and E species in a fast (i.e. non-activated) process, in contrast to the slow formation of C species on the unreacted sample. Finally by outgassing at 473 K the D1, D2 and E species completely disappear, while the F species evolve to a final simplified situation. CO/O2 INTERACTION In fig. 5 and 6 the reactivity of CO surface species towards gaseous oxygen is illustrated. This experiment has been carried out on a lighter sample: as a con- sequence the bands are less intense than those shown in fig. 1-4. The solid curve inGUGLIELMINOTTI, COLUCCIA, GARRONE, CERRUTI, ZECCHINA 101 fig.5 refers to full CO coverage: the broken one has been taken immediately after contact with 2.6 kN r r 2 CO. The same curve has been reported in fig. 6 for com- parison: the other two were taken after 1 and 20 h, respectively. The following important features are observed : (a) the C species are extremely sensitive to oxygen and an immediate reaction takes place leading to F species. In the meantime the I I I I I I I 2 9 It300 .woo 1200 10 ' I 11 ij/cm- 0 FIG. 5.-CO/02 interaction ; first stage. -, nearly total coverage ; - - -, immediately after contact with 2.6 kN m-2 02. 100 FIG. 6.-CO/O2 interaction ; further stages. - - -, as the last curve in fig. 5. - - -, after 1 h; . . ., after 20 h.102 co ADSORPTION ON MgO sample colour turns white (fig.5). (b) The D and E species react at a much lower rate with formation of species of the F family (fig. 6). This experiment clearly shows that the C , D and E species are oxygen sensitive, whereas the F species are not. The transient B species is also oxygen sensitive. In fact if a sample showing the broken or dotted spectra of fig. 1 is contacted with 02, a slow reaction takes place leading to the complete destruction of the bands at 2064- 1318 cm-l and to the formation of F species. Similar experiments carried out with N,O as oxidizing agent (not reported here for sake of brevity) gave the following results : (a) the C species were oxidized leading to species of the F family; (b) the D and E species did not react at r.t.In conclusion, species and/or families of species which show different behaviour on outgassing also have different reactivity towards oxidizing agents. DISCUSSION STRUCTURE OF ADSORBED SPECIES A SPECIES The presence of a single band indicates that the diatomic CO structure is retained. The reversibility of this species and the positive shift with respect to the gaseous molecule stretch (Av = +47 cm-l) are in agreement with a CO molecule weakly adsorbed on cationic sites. Similar positive shifts have been found for CO inter- acting with exposed Mg2+ ions in magnesium spinel and exchanged zeolites (MgX and MgY).' B SPECIES These are characterized by bands at 2064 and 1318 cm-'. As these frequency values are very near to those characteristic of the D species and a similar reactivity toward oxygen and N20 is also observed, we are led to assume that B transient entities and the D species have similar structures, which will be discussed together in the following section.c SPECIES These are characterized by four bands in the 1500-800 cm-I range and hence must have a more than diatomic nature. The highest mode of the C species is at N 1475 cm-I . In compounds containing carbon4arbon or carbon-oxygen bonds, bands are found : in the 2100-1900 cm-I range in the case of triple or double cumulated bonds (C=C, C=C=O), in the 1850-1600 cm-' range in the case of double bonds (C=O in ketones and esters, C=C in olefins), and in the 1200-900 cm-l range in the case of single bonds (C-0 in alcohols and esters, etc., C-C in hydrocarbons), so that none of these structures can account for the experimental results.A C-0 bond order of between one and two is suggested by the above considera- tion, so that we are forced to consider less common structures, like anionic, resonance stabilized, CO clusters with the general formula : (C0);- where n 2 2 and x = 2, 4 . . .. This conclusion based on pure i.r. data completely confirms the reflectance results.3* Cyclic, resonance stabilized carbanions (a)-(d) are well known : 2-GUGLIELMINOTTI, COLUCCIA, GARRONE, CERRUTI, ZECCHINA 103 other ring sizes (e.g. 3, 7 membered) can also occur. The i.r. spectra of the solid alkali derivatives of all these compounds are indeed characterized by two i.r. bands in the 1500-800 cm-l range,9 which is half the number of observed bands of the C species.However, due to the reduced symmetry caused by the strong interaction with the surface, the two Raman modes in this range could also become i.r. active. For simplicity we refer to the most common six-membered twice-charged ring (c). One reason for this choice is that the electronic transition of the rhodizonate com- pounds lo is close to the 21 500 cm-l absorption shown by C species in the visible region.3 We propose that a substantial interaction occurs with exposed Mg ions, e.g., as in structure (e), where three Mg ions in a triangular array are considered. The n (el reason for such a choice is given later. The interaction between carbon atoms and Mg ions will induce a partial sp3 character in the carbon atoms and cause a slight distortion of the rhodizonate structure from planarity, so increasing the number of i.r.active modes to four without much loss of aromaticity of the molecule. Model (e) is tentative, as many others can be conceived, e.g., with x = 4 or interacting with only two Mg ions. Moreover, as already observed,l resonance stabilized anions with open structure can also have the low bond order characteristic of C species. Open, resonance stabilized (C0):- clusters can be represented as in structure (f) 1 df) where n >, 2 and x = 2 or 4. In our opinion more than diatomic clusters with open structure have to be ruled out because : (a) compounds with similar structure are unknown in the homogeneous phase and (b) the presence of only four i.r. active modes in the 1500-800 cm-l range restricts the n value to 3 or 4, as for higher degrees of polymerization the i.r.spectrum should be more complicated. However, a low degree of polymerization (i.e., a conjugated system of limited extension) does not explain the 21 500 cm-1 band observed by reflectance spectroscopy. D SPECIES Three types of D species have been found each characterized by two bands at : 2108 and 1358 cm-l (D1), 2097 and 1365 cm-l (Dz) and 2084 and 1392 cm-l (D3). Two bands in the 2150-1300cm-1 region indicate that again more than diatomic structures are involved and undoubtedly suggest that ketenic groups are present in the adsorbed species.12 Moreover, as the D3 species can be transformed into C species by adding CO from the gas phase, they can be considered as less complicated versions of the C species.104 CO ADSORPTION ON MgO The two requirements are completely fulfilled by the dimeric CO structures :13 0 I! C II C -0’ ‘Mg’ (s) In fact they are the smallest (more than diatomic) charged fragment of the C structures containing a ketenic chromophore.Moreover, the two observed bands are typical of in-phase and out-of-phase vibrations of cumulated double bonds. The C-D3 transformation process can consequently be represented as in scheme (1) Of course reaction (l), leading to resonance stabilized structure, can occur only if suitable arrangements of the counter Mg2+ ions are present (at least a pair of Mg2+ ions in a suitable position being required). If this favourable arrangement is missing, structure (g) cannot show any tendency to incorporate CO to give a cyclic species, This could explain the presence of D1 and D2 species, which as a consequence are assumed to differ from D3 species only by the number and geometry of surrounding ions.Finally, owing to the spectroscopic similarity between D and B species, we con- clude that this transient species also has a D-like ketenic structure. However, the reason for its transient nature cannot be easily understood in the light of this dis- cussion. E SPECIES These absorb in two different regions, namely 1585 and 1160 cm-l. In this case more than diatomic species are also necessarily involved. We think that, as in the case of D species, we are dealing with three components (El, E2, E3). In fact the absorption at high frequency is clearly the superposition of three bands which are affected by outgassing in different ways.For instance the El component does not change during the desorption process in the 373-423 K range, unlike the others. To make a reasonable hypothesis of their structure the following have to be taken into account : (i) the frequencies of the observed bands imply bond orders of < 2; (ii) the E2 and E3 species are less complicated versions of the C species as they can be transformed into each other by addition or removal of CO. These requirements are met by the dimeric fragments(C0);‘ in a steric arrangement different from the one of D species, e.g., involving three magnesium ions :GUGLIELMINOTTI, COLUCCIA, GARRONE, CERRUTI, ZECCHINA 105 Similar bands have been found by Buchner13b in the spectra of the reaction products of CO with alkali metals in liquid ammonia.Moreover, the structure is very close to the cis-hyponitrite NzOi- found by some of us on NO/Mg0,14 which shows two infrared absorptions in a similar position. The E2, E3-C transformation process can be consequently illustrated by scheme (2) It is noticeable that schemes (1) and (2) are very similar : in fact they represent parallel ways for formation and destruction of C species. Species D1, E2 and E3 are present in fairly fixed ratios at all outgassing stages : hence a sort of equilibrium between the two types of species can be hypothesized following scheme (3) During the last outgassing stage at 473 K both types of species disappear simul- taneously leaving on the surface only F (carbonate-like) species.This is further evidence that all these species have similar stability and similar destruction patterns during the outgassing procedure. F SPECIES The band position l5 and stability towards oxygen demonstrate that F species have a carbonate-like structure. Moreover, as in the previous cases, we are dealing with families of similar species (organic and bidentate-type carbonates) rather than single well-defined structures. Owing to the complexity of the spectrum, we shall not attempt a full assignment. The organic carbonates (bands at 1774, 1740, 1710 and 1270-1220 cm-l) in any case disappear at 473 K, whereas some bidentate car- bonates (1665-1650, 1325 cm-l) are left on the surface as the only residual species. We stress here that during the adsorption process the overall intensity of the F species parallels the overall intensity of species C , D and E.This implies that upon CO adsorption both F species (on one hand) and C, D and E species (on the other hand) are simultaneously formed in fairly constant ratios. ADSORPTION MECHANISM The large number of different species and the slow formation rate indicate that CO adsorption is a very complicated process leading to both reduced and oxidized species at the same time, whose relative concentrations are fairly constant. A suitable overall mechanism can be schematized as follows :4 (n +x/2)CO + xO& -+ x/2CO$- + (CO):- (4)106 CO ADSORPTION ON MgO where x = 2 or 4 and the involved oxygen ions are considered to be in low coordina- tion state (and hence have a reduced Madelung potential and an enhanced instability).Scheme (4) indicates that a disproportionation reaction of CO is taking place. Thus, we neglect the possible contribution to the reaction from electron transfer processes from the solid, which have been evidenced by other studies.16 Although the CO coverage is not too far from the reported concentration of electron releasing centres, the constant presence of both oxidized and reduced species strongly favours a disproportionation scheme like (4). Moreover the CO reaction on CaO and SrOY1' leading to much higher coverages, further supports this idea. Therefore, we assume that each reduced species is formed along with an oxidized partner. An adsorption/desorption scheme taking into account most of the results is reaction (5) : -co t +"+ F,C slow s1 .Cq F,B 1 .CQ 16 I I It tI Three different sites (Sly S2, S , ) are assumed, arbitrarily, as no definite evidence is available about the independence of the three mechanisms.The transient nature of the B species is taken into account, as well as different pathways for C formation and depletion. The possible equilibrium between D1 and E is indicated by dashed arrows. In the above reaction scheme, both reduced and oxidized species are envisaged as being depleted together by reversal of reaction (4), with complete restoration of the surface. This is probably an over-simplification. It has been shown (fig. 3) that the last desorption step leaves on the surface small amounts of bidentate carbonates without any reduced partner.In our opinion, this suggests that the reduced dimeric species E and D can be, to some extent, also depleted in a different way, which leads to the formation of carbon atoms (not visible in i.r.), according to eqn (6) : (C0)2 - -+. co,,, + c + 0::. (6) We observed a moderate darkening of the pellets after several adsorption/ desorption cycles and Lunsford and Jayne revealed the presence of traces of C 0 2 in the gas phase during desorption. In conclusion, a Boudouard reaction seems to occur to a small extent at 473 K in agreement with old observations.18 The CO disproportionation shown in scheme (4) would be the first step towards Boudouard reaction. ADSORBING SITES Scheme (4) has the merit of pointing out that basic oxygen ions on the surface act as initiators of the reaction, but it does not take into account the contribution ofGUGLIELMINOTTI, COLUCCIA, GARRONE, CERRUTI, ZECCHINA 107 the positive surface ions to the stabilization of negative CO polymers.It can thus generate the incorrect idea that the reaction centres consist oiily of single uncoordin- ated oxygen ions instead of uncoordinated oxygen ions surrounded by a suitable environment of Mg ions as the proposed assignment of species E, C and D suggests. It has been also shown that the negative CO clusters and the carbonate-like species are formed together, so that both kinds of species should lie in the region. same surface - \-co Fm. 7.Tentative model for the mechanism and site of reaction. A possible model of the site which seems to explain much of experimental data is given in fig. 7.The initial attack of the CO molecule occurs on a coordinatively unsaturated 02- ion, e.g., in a corner or kink position [fig. 7(a)]. The formed car- bonate-like species can be located in such a way that three Mg ions become strongly uncoordinated and available to stabilize a hexagonal dinegative ring [fig. 7(b)], formed most likely through the transient intermediate B. It is noteworthy that if the C-C distance in the rhodizonate ion is used,g the ring of this model compound exactly fits the oxygen vacancy created in the first adsorption step.108 co ADSORPTION ON MgO Desorption at moderate temperatures leads to ring destruction and to the forma- tion of E and D species : a plausible location of these fragments is illustrated fig.7(c) and (d). CO elimination from both dimeric fragments and carbonate-like species restores the initial situation. The Boudouard-like reaction is shown in fig. 7(e) and (f). The dimeric species release one CO molecule leaving one carbon atom on the surface, according to scheme (6) [part (e)] : finally the surface carbonates are desorbed as C 0 2 CO attack on the oxygen ion located at the crystal edge is probably also occurring. However, due to the less favourable arrangement of the ions, the stabilization of large CO clusters is more difficult and the process thus stops at low degrees of polymer- ization. In agreement with this hypothesis, we observed that some D species (D1, DJ and the El one do not show any tendency to incorporate CO to give larger clusters.[part ff >I - CONCLUSIONS The i.r. evidence gives further support to the assignment of U.V. bands of CO adsorbed on MgO to dimeric and polymeric structures. Unexpected species are revealed by vibrational spectroscopy, as well as species like carbonates and molecular CO, which are not seen by U.V. spectroscopy. A detailed assignment has been attempted (except for the carbonates), although this assignment is difficult due to the peculiar reactions occurring at the surface, which do not have an exact correspondence in ordinary chemistry. The tentative nature of the assignments must be acknowledged. An overall disproportionation reaction scheme is shown to hold, and several steps in the reaction network have been understood. The overall trend is to yield polymeric reduced species, and this coincides with what was already found for NO chemisorption on the same samples.4* l4 In the case of NO, the degree of polymer- ization is necessarily smaller and only dimers are found.A further interesting coincidence is the existence in both cases of a depletion pathway which does not coincide with the reversal of the formation reaction, leading in the case of CO to carbon atoms and to N,O in the case of NO. R. St. C. Smart, T. Slager, L. H. Little and R. G. Greenler, J. Phys. Chem., 1973, 77, 1019. H. Kolbel, M. Ralek and P. Jiru, Z. Naturforsch., 1970, 25a, 670. A. Zecchina and F. S. Stone, J.C.S. Chem. Comm., 1974, 582. F. S. Stone and A. Zecchina, Proc. Sixth Znt. Congr. Catalysis (Chem. SOC., London, 1976), p. 162. A. Zecchina, Proc. Sixth Znt. Congr. Catalysis Discussion (Chem. SOC., London, 1976), p. 179. J. H. Lunsford and J. P. Jayne, J. Chem. Phys., 1966,44,1492. ’ (a) C. Morterra, G. Ghiotti, F. Boccuzzi and S. Coluccia, J. Catalysis, in press; (b) C. L. Angel1 and P. C. Schaffer, J. Phys. Chem., 1966,70,1413. R. West, H. Y. Niu, D. L. Powell and M. V. Evans, J. Amer. Chem. SOC., 1960,82,6204. R. West in Non-benzenoid Aromatics, ed. J. P. Snyder (Academic Press, New York, 1969), vol. 1 , p. 31.. lo B. Eistert and G. Bock, Angew. Chem., 1958,70,595. l1 A. Zecchina and F. S. Stone, J.C.S. Faraduy Z, 1978, 74, 2278. l2 J. H. Wotiz and W. D. Celmer, J. Amer. Chem. SOC., 1952, 74, 1860. l3 (a) W. Buchner, Helv. Chim. Acta, 1963, 46, 2111 ; (6) 1966, 49, 907. l4 L. Cerruti, E. Modone, E. Guglielminotti and E. Borello, J.C.S. Faraday I, 1974, 70, 729. l5 L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, London, 1966), p. 74. l6 A. J. Tench and R. L. Nelson, Trans. Faraday Soc., 1967, 63,2254. l7 A. Zecchina et al., unpublished data. l8 J. Cleminson and H. V. A. Briscoe, J. Chem. Soc., 1926, 2148. (PAPER 8/685)

ISSN:0300-9599    

DOI:10.1039/F19797500096

出版商:RSC    

年代:1979

数据来源: RSC