Chemical Rate Processes I n order to provide full coverage of the literature on chemical rate processes, the survey has been divided into four major topics. Each of these topics is reviewed by a member of the faculty of Princeton University. General rate principles and homogeneous reactions of various kinds are first considered. A review of recent developments i n heterogeneous catalysis follows. Transport properties of gases and liquids is the topic of a third section, and transport properties i n the solid state are represented by a review i n the important area of diffusion and oxidation of metals. A review of kinetics and equilibria was last presented i n 1951 (92A). This edition emphasizes papers published i n 1952, although many selected papers appearing i n 1951 are included. Numerical methods for obtaining rate constants for two consecutive second-order reactions, competitive consecutive second-order reactions with initiation reactants present i n equal quantities, and first- and second-order two-stage reactions have been reported. The decomposition of the oxides of nitrogen, the kinetics of the pyrolysis of hydrocarbons, and reactions between molecules and free radicals or atoms were the subject of numerous investigations. I n the field of heterogeneous catalysis, diversified answers have been evolved t o the questions of surface homogeneity or heterogeneity, the natura of slow sorption phenomena, the importance of geometrical or electronic factors, and catalyst selectivity. Experimental and theoretical work on viscosity, thermal conductivity, and diffusion coefficients were developments i n the study of the molecular transport properties of fluids. The intensity of investigations of diffusion i n metallic solid solutions and the oxidation of metals is reflected by a large number of papers i n this limited field. Diffusion coefficients have been obtained by a variety of experimental methods, including the utilization of radioactive gold and the unique properties of germanium.
Rate Theory and Homogeneous Reactions R. H. W I L H E L M
Princeton University, Princeton, N. J.
h-E nex general text., “Kinetics and Mechanism.” , bv ” Frost and Pearson (29A) was published. Homogeneous chemical reactions are treated, with attention to theories of gaseous and liquid phase reactions, catalyzed and uncatalyzed. A particularly extensive chapter on complex reactions is noted. Since the last review ( 9 2 A ) a monograph by Hougen ( 4 4 A ) appeared, “Reaction Kinetics in Chemical Engineering.” It is an amplification of an institute lecture before the American Institute of Chemical Engineers and presents material in the area of chemical kinetics and diffusional rate processes. An important nem source book is “Tables of Chemical Kinetics, Homogeneous Reactions” (67.4), published by the Kational Bureau of Standards, with Thon as editor. The purpose of these tables is a critically evaluated compilation of the available factual numerical data on rates and rate constants. Bowden and Yoffe ( 9 A ) have mitten a monograph summarizing their studies on the activation of explosions. The sequence of events when explosives are set off by mechanical blows or by frictional action are discussed. A monograph ( 2 0 A ) presents two symposia on ultrasonics, including application to chemical reactions. Reaction R a t e Calculations. Eyring and Smith (XX4)revien- the status of theoretical calculations of reaction rates. They note that the “semiempirical method” for calculating activation energies is seldom used now and draw attention to quantitative formulations of rate expressions starting with knowledge about charge distributions on organic compounds. Previous TT orkers, such as Ingold, Robinson, Lowry, and Lapa orth, proposed quali-
0
tative rules regarding reaction rate and equilibrium based on electronic interpretations. The latest in the series, by Smith and Eyring (79A), deals with charge distributions on halogensubstituted alkanes. The authors show that net charges may be simply correlated with the activation energies for reactions of halides with sodium atoms. In two papers, Hollingsworth (@?A,4911) discusses relations between chemical rate lam for reversible reactions and the necessary and sufficient conditions for equili6rium. Gilkerson, Jones, and Gallup (32A) consider chemical reactions near equilibrium. In seeming contradiction to usual deductions from thermodynamics, previous authors sho1Ted by purely mathematical reasoning or by methods of models that the rate of chemical reaction close to equilibrium is directly proportional to the free energy difference between reactants and products. Present aut.hors show the relation also to be the consequence of the theory of absolute reaction rates. A number of papers report mathematical formulations for complex reactions. Schwemer and Frost ( 7 6 A ) illustrate a numerical method for obtaining rate constants for two consecutive second-order reactions. Frost and Schwemer (SOd) extend their studies to integrations for competitive consecutive secondorder reactions with initiating reactants present in equal quantity. Experimental rate measurements on the saponification of ethyl adipate and ethyl succinate illustrate the use of the solutions. Widequist (9OA) presents formulas for the rate constants for first- and second-order tu-o-stage reactions, as do also Natta and Mantica (66A). The latter authors also consider operation on
May 1953
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INDUSTRIAL AND ENGINEERING CHEMISTRY
batch and continuous basis, with and without recirculation of intermediates. Skrabal (77A) uses a matrix analysis among nine parallel reactions. An interesting system of coordinated reactions such as are presumed to occur in enzyme systems is the subject of a generalized treatment in a paper by Hinshelwood (40A). In sucb reactions, the rates of reactions are mutually dependent upon products of other component reactions. Steady state conditions are suggested t o be of interest in autosynthesis. Procedures for formulating complex chain reactions are proposed by Breitman ( 1 4 8 ) where direct solution of the system of differential equations is impracticable. Stepukhovich and Timonin ( 8 I A ) integrate and discuss the characteristics of a sequence of reactions of the type: 2A -c 2A1 -P Az. A number of papers report on the decomposition of nitric acid or its oxides. The classical reaction of the decomposition of nitrogen pentoxide in the presence of nitric oxide is studied by Mills and Johnston ( 6 8 A ) over a large pressure range. Previously proposed mechanisms for this reaction and for the decomposition of nitrogen pentoxide are confirmed. Anomalies long associated with the low pressure decompositions of the pentoxide are explained. The mechanisms of four kinetic systems involving nitrogen pentoxide are analyzed in a second paper by Johnston (46A). Eight reactions which are believed to be the elementary steps are set forth. The kinetics of the second-order reaction between nitryl chloride and nitric acid are studied over a wide range of conditions by Freiling, Johnston, and Ogg (B7A). This reaction is studied as an analog to the reaction between nitrogen trioxide and nitrogen dioxide in the nibrogen tetroxide decomposition. Wise and Frech (94A)measured the decomposition rates of nitric oxide between 872’ and 1275” K. They report second-order kinetics, but with a homogeneous mechanism a t higher temperatures and heterogeneous mechanism at lower temperatures and with different energies of activation in the two regions. Johnston (47A) examined the data extant on the thermal decomposition of nitrous oxide. Correction for a heterogeneous component brings the results in line with the theory of Rice and Ramsperger. The thermal decomposition of nitric acid is the topic of two papers. Johnston, Foering, Tao, and Messerly (48A) find the decomposition to be a heterogeneous reaction a t low temperatures and predominantly a fast, homogeneous, firstorder reaction above 400” C. Frejacques (28A) reports the initial reaction a t higher temperatures in this decomposition to be second-order reaction, At lower temperatures, a second-order reaction is observed between nitric acid and nitric oxide with a lower energy of activation than for the high temperature reaction. The reactions between carbon dioxide, carbon monoxide, and charcoal were studied in detail by Bonner and Turkevich ( 8 A ) , using GIpas the means for following the reaction. A rapid reaction of radioactive carbon dioxide with the surface to give an oxygenated surface and a radioactive carbon monoxide is postulated. A subsequent slow two-stage unimolecular decomposition of the oxygenated surface occurs, leading to nonradioactive carbon monoxide and charcoal. A novel method of determining the rates of very fast bimolecular reactions is reported by Garvin and Kistiakowsky (S1A). It is an adaptation of the “diffusion flame” method of Polanyi. The authors studied the kinetics of the reactions between boron trifluoride and various monomethyl amines. The kinetics of the vapor-phase dimerization reactions of tetrtlfluoroethylene and chlorotrifluoroethylene were measured by Lacher, Tompkin, and Park (51A). The reactions are homogeneous and second order. The Arrhenius constants are determined and discussed. Frequency factors of some bimolecular reactions are discussed by Rollefson (74A). Frequency factors of unimolecular dissociation reactions are reviewed and discuased by Szwarc (86A), by Szwarc and Leigh (87A), and by Steacie and Szwarc (80~4).The effects of substi?
895
tution on the thermal stability of a number of hydrocarbon radicals are evaluated by Szwarc (86A). The thermal decomposition of hydrocarbons is reviewed by Stubbs and Hinshelwood (8SA). Pyrolysis. Experimental studies of the kinetics of pyrolysis reactions have been numerous. Among the compounds whose decomposition kinetics were studied are normal paraffin hydrocarbons, branched-chain hydrocarbons, olefins, benzene derivatives, nitro hydrocarbons, peroxides, diborane, and hydrazine. Stubbs, Ingold, Spall, Danby, and Hinshelwood (84A) studied the decomposition of normal hydrocarbons by nitric oxide in higher pressure ranges than in former work. Complex changes in order of reaction with changes in pressure are discussed. Nonchain reactions are presumed to be occurring. By contrast, in the decomposition of branched-chain hydrocarbons studied by Peard, Stubbs, Hinshelwood, and Danby (71A ) , decomposition is stated to occur partly by radical chain reactions (suppressible by nitric oxide) and partly by molecular reactions. The thermal decomposition of propylene was studied in a static system by Ingold and Stubbs (@A). A complex series of steps is proposed to account for the distribution of products obtained. At lower temperatures the decomposition sequence of events is additionally complicated by polymerization reactions. In two papers by Leigh and Szwarc ( 5 9 4 6SA), the pyrolysis of propylbenzene and butylbenzene, respectively, were measured by means of the “toluene carrier” technique through which the formation of hydrogen radicals may be detected. From measurements of energy of activation and various previously known heats of formation, the heat of formation of ethyl radical and of propyl radical is estimated. The thermal decomposition of 1nitropropane and nitroethane was studied in static experiments by Cottrell, Graham, and Reid ( I I A ) . The reactions are substantially homogeneous and with unimolecular mechanisms leading initially to an olefin and nitrous acid, followed by oxidation of some of the olefin. Hillenbrand and Kilpatrick (S7A) verify the results of previous investigators for the decomposition of nitromethane using a flow method in this work. The kinetics of the pyrolysis of diborane are measured by Clarke and Pease (WIA) and by Bragg, McCarthy, and Norton (10A). Sequential reaction steps are suggested to account for the high fractional order of the reaction. Hanratty, Pattison, Clegg, and Lemmon ( 3 4 4 ) studied the decomposition of hydrazine vapor in a silica flow reactor. The reaction is first order and heterogeneous, the nature of the surface affecting the energy of activation substantially. The thermal decomposition of di-tertbutyl peroxide was investigated by Murawski, Roberts, and Szwarc (666)and by Brinton and Volman (15A), the former, in an excess of toluene, the latter in an excess of ethylenimine. Both sets of investigators are in substantial agreement with each other and with previous workers regarding the Arrhenius constants in the -0-0- splitting, first-order, homogeneous initiating Istctp. Rebbert and Laidler (73A) investigated the decomposition of diethyl peroxide in a flow system with an excess of toluene. The kinetic constants and mechanism of the initial scission are given. The heat of formation of CIH,O radical is calculated. Free Radicals. Certain papers during the last year emphasized reactions between molecules and free radibals or atoms. Broad-gaged papers include an address by Melville (59A), in which the author discusses new methods for the quantitative study of free-radical chemistry, and a review on the nature and importance of free radicals by Muller (6SA). A lecture by Hey ( 6 6 A ) considers recent developments in the chemistry of freeradical reactions in solutions. The interaction of atomic hydrogen and deuterium with cyclic and paraffin hydrocarbons is the subject of a paper by Schiff and Steacie (75A). Hydrogen and deuterium atoms, produced by the discharge method, were reacted a t room temperature with nine different hydrocarbons.
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Evidence is presented for the belief, based upon the appearance of methane as the major product, that hydrogen atoms react with hydrocarbons to produce hydrocarbon radicals, which then undergo rapid "atomic cracking" reactions. Berlie and LeRoy ( 6 A ) studied the reaction of atomic hydrogen with ethane. A proposed mechanism of the exchange between hydrocarbon radicals and molecular deuterium involving the whole deuterium (Dz) molecule with an intermediate complex is given by Voevodskii, Lavrovskaya, and Mardaleishvili (89A). Blades and Winkler ( 7 A ) measured rates of reaction between nitrogen atoms and methane and ethane. Activation energies and steric factors are given and a mechanism is outlined. The rate of recombination a t room temperature of methyl radicals was studied by Miller and Steacie (61A)through photolyzing dimethyl mercury in the presence of nitric oxide. This work was extended by Durham and Steacie ( M A ) , using a radioactive tellurium mirror technique. The formation of trifluoromethyl radicals in the gas phase is reported by Hodgins and Haines ( 4 I A ) . The reaction is performed in a diffusion flame with sodium vapor and with trifluoro derivatives of iodine, bromine, or chlorine. Direct reactions of methyl radicals with hydrogen and deuterium are reported by Majury and Steacie (55A),Davison and Burton ( M A ) ,and Anderson, Davison, and Burton (1A). Raal and Steacie ( % A ) prepared methyl radicals by the photolysis of acetone and studied their reactions with a series of halogenated methanes. Energies of activation and steric factors are reported. Marcotte and Noyes (56A) reacted similarly prepared methyl radicals with oxygen and proposed a mechanism involving HCO as an intermediate in the formation of carbon monoxide or carbon dioxide from the indirect reaction of methyl radicals and oxygen. Marcus, in two papers (57A, 68A), discusses and extends previously presented theories regarding relations between the reverse reactions of unimolecular dissociations and free radical recombination reactions, particularly for methyl radical recombinations. The kinetics of the removal of hydrogen from acetaldehyde by methyl radicals, produced by the decomposition of di-tert-butyl peroxide, are studied by Brinton and Volman (16A). Characteristics of the reaction of ethyl radicals with deuterium were determined by Wijnen and Steacie (91A ) . Burns and Dainton (18.4) investigated the retardation of the photochemical formation of phosgene from carbon monoxide and chlorine by small amounts of nitrosyl chloride. Reactions by which nitrosyl chloride combines with intermediates of the prime reaction, chlorine atoms, and COCl radicals, are investigated. The potential energy diagram for the reaction of chlorine and nitrosyl chloride is calculated. Melville, Robb, and Tutton (60A),by a modified dilametric method, studied the photochemically induced interaction of trichloromethyl radicals with cyclohexene. Preliminary rate constants are given for the chain-propagating and chainterminating reactions as well as the over-all energy of activation. Johnston and Libby ( @ A ) 6nd the exchange reaction between hydrogen chloride and chlorine in the gaseous state to be heterogeneous a t room temperature. The homogeneous exchange is slow, but a rapid photochemical exchange takes place. Photochemical exchange, it is suggested, occurs through chlorine atoms exchanging rapidly with both hydrogen chloride and chlorine. I n a study of the kinetics of photochlorination of chloroform vapor, Winning (9SA) confirmed the previous rate expression of Schumacher and Wolff. Jones (60A)measured the relative rates of the photochemical reaction of hydrogen and of tritium hydride with chlorine, and the data are considered from the viewpoint of the theory of the absolute rates. The classical photolysis of acetaldehyde was studied by Dodd (&$A)in the high temperature range between 156' and 450' C. Various values for the over-all activation in the literature may be resolved, it is suggested, by correction for the effect of the unimolecular chain termination, Slow Oxidations. In the area of slow oxidations, Stephens and Pease ( H A ) studied the noncatalytic oxidation of ammonia.
Vol. 45, No. 5
For ammonia-rich mixtures, rates are roughly proportional to the product of ammonia and oxygen concentrations. I n ammonialean mixtures the square of the oxygen concentration is involved. Experimental evidence for each of a number of steps is detailed in the work of Bell, Raley, Rust, Seubold, and Vaughan ( I A )for free radical reactions associated with low temperature oxidations of paraffins. Mechanisms of slow thermal oxidation of organic substances are discussed also by Iiiclause (69A), with acetaldehyde oxidation as an illustrative example. The author suggests a bimolecular initiating reaction rather than a unimolecular decomposition as do the investigators above. Hinshelwood (39A) provides a general discussion of the influence of substituents on the oxidation of hydrocarbons. The homogeneous oxidation of ethylene oxide was studied by Burgoyne and Kapur (17A)above and below the optimum condition for cool flame formation. The progressive changes in products and reactants were followed and were related to changes in total pressure. Liquid-Phase Reactions. Liquid-phase reactions in this review are restricted to acid-base catalyzed organic reactions. The Hammett acidity function in six formal perchloric acidsodium perchlorate mixtures are given by Harbottle (S6A). In a theoretical paper, WIulliken (64A) generalizes previous ideas t o give a quantum-mechanical picture of molecular complexes, thereby giving insight into possible stages through which a weak Lewis acid may be transformed into a functioning proton acid by a basic ionizing solvent. A series of studies on aromatic sulfonation has been continued by Brand and coworkers. I n one paper, Brand and Rutherford ( 1 S A ) determined between 20" and 80" C. the vapor pressure of sulfur trioxide above oleum solutions containing 0 to 35% free sulfur trioxide. The objective was to obtain more precise information regarding the activity of sulfur trioxide in these solutions. I n a second paper, Brand and Horning (114) found that the velocity of sulfonation of mononitrobenzene and p-halogen-CoHi-N( CH3)sf was a function of the cation, HOSOs+. A linear relation of the type expected was obtained between the logarithm of the rate constant and the J function of sulfuric acid, where J = H o - log pS03. The Hammett acid function is H , and p S O 3 is the partial pressure of sulfur trioxide as determined in the earlier paper (1SA). Brand and Paton ( I d A ) measured relative nitration rates of various derivatives of p-RCBHPN02 and p-RCaH4N(CHa)a+, where R is the same substituent in both compounds. Relative velocities are discussed in terms of the electromeric polarizability of the substituents. Blackall, Hughes, and Ingold ( 6 A ) , in a series of papers on kinetics and mechanism of aromatic nitration, studied the nitration of p-CIC~H~OCHsby nitric acid in acetic acid with nitrous acid as catalyst. The catalysis is believed to operate by the preliminary nitrosation of the aromatic molecule predominantly by the nitrosonium ion and to a minor extent by the nitrogen tetroxide molecule. Bell and Goldsmith ( 4 A ) studied the kinetics of the base-catalyzed halogenation of 2-carbethoxycyclohexanone. The reaction is of the first order with respect to the ester, zero order with respect to chlorine, and is catalyzed by water and the anions of carboxylic acids. Bell and Caldin ( S A ) measured and analyzed the kinetics of the base-catalyzed decomposition of the inorganic compounds, nitramide and deuteriated nitramide, in anisole. The catalyst was dimethylanilhe. Results do not deviate from the simple Arrhenius expression. Relative rate values between the two compounds are attributed to differences of zero point energy in the energy of activation. The kinetic salt effect on the acid-catalyzed decomposition of trioxane is reported by Paul (70A). The rates are correlated with the Hammett indicator acidity function. The acids used were perchloric acid mixed with sodium perchlorate to obtain the salt effect (35A). Experimental measurements of acid or base hydrolysis reactions have been reported. Glew and Moelwyn-Hughes (SSA) studied the kinetics of the acid and alkaline hydrolysis of methyl fluoride in water and McKinley-McKee and Moelwyn-Hughes
May 1953
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INDUSTRIAL AND ENG INEERING CHEMISTRY
Rate Theory and Homogeneous Reactions
(39A) Hinshelwood, C. N., Discussions Faraday SOC., No. 10, 266 (1951). (40A) Hinshelwood, C. N., J . Chem. SOC.,1952, p. 745. (41A) Hodgins, J. W., and Haines, R. L., Can. J . Chem., 30, 473 (1952). (42A) Hollingsworth, C. A., J. Chem. Phys., 20, 921 (1952). (43A) Zbid.,,lO, 1649 (1952). (44A) Hougen, 0. A,, Chem. Eng. Progr., Monograph Ser., No. 1 (1951). (45A) Ingold, K. U., and Stubbs, F. J., J . Chem. SOC.,1951, p. 1749. (46A) Johnston, H. S.,J . Am. Chem. SOC.,73, 4542 (1951). (47A) Johnston, H. S.,J . Chem. Phys., 19, 663 (1951). (48A) Johnston, H. S., Foering, L., Tao, Yu-Sheng, and Messerly, G. H., J . Am. Chem. Soe., 73,2319 (1951). (49A) Johnston, W. H., and Libby, W. F., Ibid., 73, 854 (1951). (50A) Jones, W. M , J . Chem. Phys., 19,78 (1951). (51A) Lacher, J. R., Tompkin, G. W., and Park, J. D., J . Am. Chem. SOC.,74, 1693 (1952). (52A) Leigh, C. H., and Szwarc, ;LI , J . Chem. Phys., 20,403 (1952). (53A) Ibid., p. 407. (54A) McKinley-McKee, J. S.,and Moelwyn-Hughes, E. A., Tranr. Faraday Soe., 48,247 (1952). (55A) Majury, T. G., and Steacie, E . W. R., J . Chem. Phys., 20, 197
(1A) Anderson, R. D., Davison, S., and Burton, M., Discussions Faraday SOC.,No. 10, 136 (1951). (2A) Bell, E. R., Raley, J. H., Rust, F. F., Seubold, F. H., and Vauahan, W. E..Ibid.. No. 10.242 (1951). (3A) Bell, R. P., and Caldin, E. F., Trans. Faraday SOC.,47, 50 (1951). (4A) Bell, R. P., and Goldsmith, H. L., Proc. Roy. SOC.(London), A210,322 (1952). (5A)Berlie, M. R., and LeRoy, D. J., J . Chem. Phys., 20, 200 (1952). (6A) Blackall, E. L., Hughes, E. D., and Ingold, C. K., J . Chem. Soe., 1952, p. 28. (7A) Blades, H., and Winkler, C. A,, Can. J . Chem., 29, 1022 (1951). (SA) Bonner, F., and Turkevich, J., J . Am. Chem. SOC., 73, 561 (1951). (9A) Bowden, F. P., and Yoffe, A. D., “Initiation and Growth of Explosions in Liquids and Solids,” Cambridge Monographs on Physics, London, Cambridge University Press, 1952. (10A) Bragg, J. K., McCarthy, L. V., and Norton, F. J.,J.Am. Chem. SOC.,73, 2134 (1951). (11A) Brand, J. C. D., and Homing, W. C., J . Chem. SOC.,1952, p. 3922. (12A) Brand, J C. D., and Paton, R. P., Ibid., 1952, p. 281. (13A) Brand, J. C. D., and Rutherford, A., Ibid., 1952, p. 3916. (14A) Breitman, IM.V., Doklady Akad. N a u k S.S.S.R., 78, 1153 (1951). (15A) Brinton, R. K., and Volman, D. H., J . Chem. Phys., 20, 25 (1952). (16A) Brinton, R. K., and Volman, D. H., Zbid., 20, 1053 (1952). (17A) Burgoyne, J. H., and Kapur, P. K., Trans. Faraday SOC.,48, 234 (1952). (18A) Burns, W. G., and Dainton, F. S., Zbid., 48,52 (1952). 74,1462 (1952). (19A) Burwell, R. L., Jr., J . Am. Chem. SOC., (20A) Chem. Eng. Ppogr., Symposium Ser., No. 1 (1951). (21A) Clarke, R. P., and Pease, R. N., J . Am. Chem. SOC.,73, 2132 (1951). (22A) Cottrell, T. L., Graham, T. E., and Reid, T. J., Trans. Faraday SOC., 47, 1089 (1951). (23A) Davison, S., and Burton, M., J . Am. Chem. Soc., 74, 2307 (1952). (24A) Dodd, R. E., Trans. Faraday SOC.,47,56 (1951). (25A) Durham, R. W., and Steacie, E. W. R., J . Chem. Phys., 20, 582 (1952). (26A) Eyring, H., and Smith, R. P., J . Phys. Chem., 56,972 (1952). (27A) Freiling, E. C., Johnston, H. S., and Ogg, R. A., J . Chem. Phys., 20,327 (1952). (28A) Frejacques, Claude, Compt. rend., 232,2206 (1951). (29A) Frost, A. A., and Pearson, R. G., “Kinetics and Mechanism,” John Wiley & Sons, New York, 1953. (30A) Frost, A. A., and Schwemer, W. C., J . Am. Chem. SOC.,74, 1268 (1952). (31A) Garvin, D., and Kistiakowsky, G. B., J . Chem. Phys., 20, 105 (1952). (32A) Gilkerson, W. R., Jones, M. M., and Gallup, G. A,, Ibid., 20, 1182 (1952). (33A) Glew, D. N., and Moelwyn-Hughes, E. A,, Proc. Roy. 8 0 0 . (London),A211,254 (1952). (34A) Hanratty, T. J., Pattison, J. N., Clegg, J. W., and Lemmon, A. E., Jr., IND. ENO.CHEM.,43, 1113 (1951). (35A) Harbottle, G., J. Am. Chem. boc., 73,4024 (1951). (3611) Hey, D. H., J . Chem. SOC.,1952, p. 1974. (37A) Hillenbrand, L. J., Jr., and Kilpatrick, M. L., J . Chem. Phys., 19, 381 (1951). (38A) Him, Jack, J . Am. Chem. SOC.,72, 2438 (1950).
(56A) ill&o;ie, F. B., and Noyes, W. A4., Jr., Discussions Faraday SOC., No. 10,236 (1951). (57A) iMarcus, R. A., J . Chem. Phys., 20, 359 (1952). (58A) Ibid.,20, 364 (1952). (59A) Melville, H. W., J . Chem. SOC.,1952, p. 1547. (60A) Melville, H. W., Robb, J. C., and Tutton, R. C., Discussions Faraday Soc., No. 10,154 (1951). (61A) Miller, D. M., and Steacie, E. W. R., J . Chem. Phys., 19, 73 (1951). (62A) Mills, R. L., and Johnston, H. S., J . Am. Chem. SOC.,73, 938 (1951). (63A) Muller, E., Angew. Chem., 64,233 (1952). (64A) Mulliken, R. S., J . Chena. Phys., 19, 514 (1951). (65A) Murawski, J., Roberts, ,J S., and Szwarc, M., J . Chem. Phys., 19, 698 (1951). (66A) Natta, G., and hlantica, E., J . Am. Chem. SOC.,74, 3152 (1952). (67A) Natl. Bui. Standards (U. S,),Circ. 510, “Tables of Chemical Kinetics, Homogeneous Reactions” (1951). (68A) Nemtsov, M. S.,and Trenke, K. M., Zhur. Obshchei Khim., 22, 415 (1952). (69A) Niclause, Michel, J. Chem. Phys., 49, 157 (1952). (70A) Paul, M. A , , J . Am. Chem. Soc., 74, 141 (1952). (71A) Peard, M. G., Stubbs, F. J., Hinshelwood, C. N., and Danby, C . J.,Proc. Roy. SOC. (London),A214,330 (1952). (72A) Rad, F. -k, and Steacie, E. W. R., J . Chem. Phys., 20, 578 (1952). (73A) Rebbert, R. E.,and Laidler, K. J.,J. Chem. Phys., 20,574 (1952). (74A) Rollefson, G . K., J . Phys. Chem., 56,976 (1952). (75A) Schiff, H. I., and Steacie, E. W. R., Can. J . Chem., 29,l (1951). f76A) Schwemer. W. C.. and Frost. A. A.. J. A m . Chem. SOC..73, 4541 (1951). (77A) Skrabal, A., Monatsh., 83,530 (1952). (78A) Smith, Hilton A , , Conley, J. B., and King, W. H., Ibid. 73, 4633 (1951). (79A) Smith, R. P., and Eyring, Henry, J . Am. Chem. SOC.,74, 229 (1952). (80A) Steacie, E. W. R., and Szwarc, M., J . Chem. Phys., 18, 1309 (1951). (81A) Stephens, E. R., and Pease, R. N., J . Am. Chem. SOC.,74,3480 (1952). (82A) Stepukhovich, A. S., and Timonin, L. M.,Zhur. Fas. Khim., 25, 143 (1951). (83A) Stubbs, F. J., and Hinshelwood, C. N., Discussions Faraday SOC.,No. 10, 129 (1951). (84A) Stubbs, F. J., Ingold, K. U., Spall, C. B., Danby, C. J., and Hinshelwood, C. N., Proc. Roy. SOC.(London),A214,20 (1952). (85A) Szwarc, M., Discussions Faraday SOC.,No. 10, 143 (1951). (86A) Szwarc, M., J. Phys. & Colloid Chem., 55,939 (1951). (87A) Szwarc, M., and Leigh, C. H., Nature, 167,486 (1951). (88A) VBne, J., Tirouflet, J., and Carrib, R., Compt. rend., 234,2074 (1952). (89A) Voevodskii, V. V., Lavrovskaya, G. K., and Mardaleishvili, R. E., Doklady Akad. Nauk S.S.S.R., 81,215(1951). (90A) Wideauist, Siavard, Acta Chem. Scand., 4, 1216 (1950). (91A) Wijne;, M. H. $., and Steacie, E. W. R., J . Chem. Phys., 20, 205 (1952). (92A) Wilhelm, R. H., and Toner, R. K., IND.ENO.CHEM.,43, 1893 (1951). (93A) Winning. I. H., Tvans. Faraday S O C . , 47, 1084 (1951). (94A) Wise, Henry, and Frech, M. F., J . Chem Phys., 20, T22 (1952).
(64A) studied the uncatalyzed and the acid-catalyzed hydrolysis of methyl nitrate and methyl acetate in aqueous solution. Smith, Conley, and King ( 7 8 A ) measured and discussed the kinetics of acid-catalyzed esterification and of the catalytic hydrogenation of the furoic and furanacetic acids. Results of the alkaline hydrolysis of phthalides are reported by VBne, Tirouflet, and Carrie ( M A ) . Hine (%A) considers carbon dichloride as an intermediate in the basic hydrolysis of chloroform. ,Burwell ( 1 9 A ) contributes to the theory of hydrolysis mechanisms in basic and acidic solutions through studies in the hydrolysis of optically active secondary butyl hydrogen sulfate in which racemization measurements give added information about the process. The kinetics of the acid-catalyzed decomposition of formaldehyde is discussed by Nemtsov and Trenke (68A).
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