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Notes: The rate of the perchloric acid hydrolysis of aqueous ethyl and butyl vinyl ethers at 25.0°C, in the presence of micellar aggregates [anionic, sodium dodecyl sulfate (SDS); cationic, cetyl trymethyl ammonium bromide (CTAB); and nonionic, polyoxyethylen—23—dodecanol, (Brij 35)], has been studied. Negligible effects were observed in the cases of cationic and nonionic micelles. Anionic micelles produce an enhancement in the reaction velocity, and the rate constants go through maxima with increasing SDS concentration. These maxima disappear in the presence of excess sodium perchlorate. All these facts are interpreted quantitatively by means of the pseudo-phase ion-exchange model.

Notes: The recent experiments on the chloride-assisted dealkylation of alkylcobalamins by a variety of oxidants (IrCl62-, AuCl4-, Fe(H2O)5Cl2+, and PtCl62-), which are scattered in several previous publications, and their general kinetic characteristics are summarized. The kinetic studies are also extended to include the dealkylations of (methylaquo)-3,5,6-trimethylbenzimidazolylcobamide and protonated base-off ethylcobalamin by IrCl62- (1.0M Cl-) and by Fe(III) ions at 0.1M Cl-, and the demethylation of (methylaquo)-3,5,6-trimethylbenzimidazolylcobamide by AuCl4- (1.0M Cl-). This extension is in an effort to substantiate the general mechanism which has been previously proposed for these oxidative dealkylations. The general kinetic characteristics are described in terms of a preassociation of the reactants, followed by a rate-determining electron-transfer process to yield the R-B12+ radical, which then undergoes further reactions to produce the products observed. The overall reactions are discussed within the framework of chlorine-bridging inner sphere electron-transfer reactions.

Notes: Hydrogen abstration from H2S by CF3 radicals, generated by the photolysis of both CF3COCF3 and CF3I, has been studied in the temperature range 314-434 K. The rate constant, based on the value of 1013.36 cm3/mol · s for the recombination of CF3 radicals, is given by \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log }\,k_2 \, = \,\left( {12.20\, \pm \,0.05} \right)\, - \,{{\left( {19,220\, \pm \,360} \right)} \mathord{\left/ {\vphantom {{\left( {19,220\, \pm \,360} \right)} {19.145T}}} \right. \kern-\nulldelimiterspace} {19.145T}} $$\end{document} with CF3COCF3 as the radical source, and \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log }\,k_2 \, = \,\left( {12.00\, \pm \,0.07} \right)\, - \,{{\left( {18,270\, \pm \,470} \right)} \mathord{\left/ {\vphantom {{\left( {18,270\, \pm \,470} \right)} {19.145T}}} \right. \kern-\nulldelimiterspace} {19.145T}} $$\end{document} with CF3I as the radical source, where k2 is in cm3/mol · s and E is in J/mol. These results resolve a previously existing controversy concerning the values of the rate constants for this reaction. They show that CF3 radicals are less reactive than CH3 radicals in attacking H2S, and this behavior indicates that polar effects play a significant role in the hydrogen transfer reactions of CF3 radicals.

Notes: A vacuum ultraviolet photolysis of C2H5Br at 147 nm was studied over a pressure range of 0.5-50 torr at 298 K. The effects of additives He and NO were also investigated.The principal reaction products were found to be C2H4 and C2H6, with lesser yields of CH4 and C2H2. With increasing pressure the product quantum yields Φi of C2H4, CH4, and CH2H6 remained constant, while that of C2H2 decreased from 0.03 to almost 0. The effect of He as an additive was found to be extremely small on the quantum yields of the major products. Addition of NO completely suppresses the formation of CH4, C2H2, and C2H6, and reduces partially the production of C2H4. The primary processes appear to involve two electronically excited states. One state mainly yields C2H4 by molecular elimination of HBr and is thought to be due to a Rydberg transition. The other state decomposes to C2H5 and Br radicals by C—Br bond fission. These two competitive reaction modes contribute to the photodecomposition in proportions of 50% and 50%. The extinction coefficient for C2H5Br at 147 nm and at 298 K has been determined as ∊ = (1/PL) In(Io/It) = 712 ± 7 atm-1 · cm-1.

Notes: Energetic H atoms produced by photolysis of gaseous HI react with CD3Br by D abstraction (1) and Br abstraction (2) and possibly also by the substitution reactions (4) and (5). Yields of HD and CD3H have been determined for several defined initial translational energies of H*. The phenomenological threshold energy of reaction (1) is 53 ± 5 kJ/mol. Over the range of initial energies of 76-109 kJ mol the integral probability of reaction (1) increases substantially, but the sum of the integral probabilities of reactions (2) and (5) shows little change. The ratio of the sum of the integral yields of reactions (2) and (5) to the integral yield of reaction (1), when normalized to equal numbers of Br and D atoms, is 69 ± 33 at an initial energy of 76 kJ/mol and 31 ± 6 at 109 kJ/mol.

Notes: Mixtures of cyanogen and nitrous oxide diluted in argon were shock-heated to measure the rate constants of A broad-band mercury lamp was used to measure CN in absorption at 388 nm [B2Σ+(v = 0) ← X2Σ+(v = 0)], and the spectral coincidence of a CO infrared absorption line [v(2 ← 1), J(37 ← 38)] with a CO laser line [v(6 → 5), J(15 → 16)] was exploited to monitor CO in absorption. The CO measurement established that reaction (3) produces CO in excited vibrational states. A computer fit of the experiments near 2000 K led to \documentclass{article}\pagestyle{empty}\begin{document}$$ k_2 \, = \,10^{11.70\left( { + 0.25, - 0.19} \right)} \,{{{\rm cm}^3 } \mathord{\left/ {\vphantom {{{\rm cm}^3 } {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}\, \cdot \,{\rm s} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k_3 \, = \,10^{13.26 \pm 0.26} \,{{{\rm cm}^3 } \mathord{\left/ {\vphantom {{{\rm cm}^3 } {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}\, \cdot \,{\rm s} $$\end{document} An additional measurement of NO via infrared absorption led to an estimate of the ratio k5/k6: with k5/k6 ≃ 103.36±0.27 at 2150 K. Mixtures of cyanogen and oxygen diluted in argon were shock heated to measure the rate constant of and the ratio k5/k6 by monitoring CN in absorption. We found near 2400 K: \documentclass{article}\pagestyle{empty}\begin{document}$$ k_4 \, = \,10^{12.68\left( { + 0.27, - 0.19} \right)} \,{{{\rm cm}^3 } \mathord{\left/ {\vphantom {{{\rm cm}^3 } {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}\, \cdot \,{\rm s} $$\end{document} and \documentclass{article}\pagestyle{empty}\begin{document}$$ {{k_5 } \mathord{\left/ {\vphantom {{k_5 } {k_6 }}} \right. \kern-\nulldelimiterspace} {k_6 }}\, = \,10^{2.68 \pm 0.28} $$\end{document} The combined measurements of k5/k6 lead to k5/k6 ≃ 10-3.07 exp(+31,800/T) (±60%) for 2150 ≤ T ≤ 2400 K.

Notes: Rate constants for the reactions of O3 and OH radicals with acetylene, propyne, and 1-butyne have been determined at room temperature. The rate constants obtained at 294 ± 2 K for the reactions of O3 with acetylene, propyne, and 1-butyne were (7.8 ± 1.2) × 10-21 cm3/molecule · s, (1.43 ± 0.15) × 10-20 cm3/molecule · s, and (1.97 ± 0.26) × 10-20 cm3/molecule · s, respectively. The rate constants at 298 ± 2 K and atmospheric pressure for the reactions with the OH radical, relative to a rate constant for the reaction of OH radicals with cyclohexane of 7.57 × 10-12 cm3/molecule · s, were determined to be (8.8 ± 1.4) × 10-13 cm3/molecule · s, (6.21 ± 0.31) × 10-12 cm3/molecule · s, and (8.25 ± 0.23) × 10-12 cm3/molecule · s for acetylene, propyne, and 1-butyne, respectively. These data are discussed and compared with the available literature rate constants.

Notes: The shock-induced thermal decompositions of vinylsilane and vinylsilane-d3 (0.2% on argon) have been studied in the temperature range of 1085-1275 K, and at total pressures of about 3100 torr. In systems without silylene traps, some induced decomposition occurs which is attributed to the silylene chain sequence VSiH → C2H2 + SiH2, S̈iH2 + VSiH3 ⇄ VSiH2SiH3 → VSiH2S̈iH + H2, VSiH2S̈iH → VSiH + S̈iH2. In the presence of silylene traps (butadiene and acetylene), the overall decomposition kinetics are log k(VSiH3, s-1) = 14.95 - 63,268 cal/2.303RT and log k(VSiD3, s-1) = 15.14 - 64,815 cal/2.303RT. Three primary processes contribute to the decomposition: 1,1-H2 elimination, 1,2-H2 elimination, and ethylene elimination. Two mechanisms are proposed, one for exclusive primary process formation of C2H4, and the other for both primary and secondary formation routes. Modeling studies are reported which show that both mechanisms can be made compatible with the rate and product yield data within experimental errors.

Notes: Relative rate constants for the reactions of hydroxyl radicals with a series of alkyl substituted olefins were measured by competitive reactions between pairs of olefins at 298 ± 2 K and 1 atmospheric pressure. Hydroxyl radicals were produced by the photolysis of H2O2 with 254-nm irradiation. The obtained rate constants were (× 10-11 cm3 molecule-1 s-1): 2.53 ± 0.06, propylene; 5.49 ± 0.17, cis-2-butene; 5.47 ± 0.1, isobutene; 6.46 ± 0.13, 2-methyl-1-butene; 6.37 ± 0.16, cis-2-pentene; 6.23 ± 0.1, 2-methyl-1-pentene; 8.76 ± 0.14, 2-methyl-2-pentene; 6.24 ± 0.08, trans-4-methyl-2-pentene; 10.3 ± 0.1, 2,3-dimethyl-2-butene; 9.94 ± 0.1, 2,3-dimethyl-2-pentene; 5.59 ± 0.07, trans-4,4-dimethyl-2-pentene. A trend in alkyl substituent effect on the rate constant was found, which is useful to predict kOH on the basis of the number of alkyl substituents on the double bond.

Notes: The absolute rate constants for reversible addition of the substituted phenylthio radicals to vinylpyridines in solution at 23°C have been determined by flash photolysis method. For each phenylthio radical, the reactivities of vinylpyridines is 3-5 times lower than that of styrene. The relative equilibrium constants estimated with flash photolysis indicate that the stabilities of the adduct carbon-centered radicals generated from vinylpyridines are similar to that from styrene. The Hammett plots suggest that the difference in the reactivities originates in the polar transition state; the nitrogen atom in 4-vinylpyridine withdraws the electrons of vinylic double bond.

Notes: The formation of nitrous acid (HONO) in the dark from initial concentrations of NO2 of 0.1-20 ppm in air, and the concurrent disappearance of NO2, were monitored quantitatively by UV differential optical absorption spectroscopy in two different environmental chambers of ca.4300- and 5800-L volume (both with surface/volume ratios of 3.4 m-1). In these environmental chambers the initial HONO formation rate was first order in the NO2 concentration and increased with the water vapor concentration. However, the HONO formation rate was independent of the NO concentration and relatively insensitive to temperature. The initial pseudo-first-order consumption rate of NO2 was (2.8 ± 1.2) × 10-4 min-1 in the 5800-L Teflon-coated evacuable chamber and (1.6 ± 0.5) × 10-4 min-1 in a 4300-L all-Teflon reaction chamber at ca.300 K and ca.50% RH. The initial HONO yields were ca.40-50% of the NO2 reacted in the evacuable chamber and ca.10-30% in the all-Teflon chamber. Nitric oxide formation was observed during the later stages of the reaction in the evacuable chamber, but ca.50% of the nitrogen could not be accounted for, and gas phase HNO3 was not detected. The implications of these data concerning radical sources in environmental chamber irradiations of NOx- organic-air mixtures, and of HONO formation in polluted atmospheres, are discussed.

Notes: Rate constants for the gas phase reactions of O3 and OH radicals with 1,3-cycloheptadiene, 1,3,5-cycloheptatriene, and cis- and trans-1,3,5-hexatriene and also of O3 with cis-2,trans-4-hexadiene and trans -2,trans -4-hexadiene have been determined at 294 ± 2 K. The rate constants determined for reaction with O3 were (in cm3 molecule-1s-1 units): 1,3-cycloheptadiene, (1.56 ± 0.21) × 10-16; 1,3,5-cycloheptatriene, (5.39 ± 0.78) × 10-17; 1,3,5-hexatriene, (2.62 ± 0.34) × 10-17; cis-2,trans-4-hexadiene, (3.14 ± 0.34) × 10-16; and trans -2, trans -4-hexadiene, (3.74 ± 0.61) × 10-16; with the cis- and trans-1,3,5-hexatriene isomers reacting with essentially identical rate constants. The rate constants determined for reaction with OH radicals were (in cm3 molecule-1 s-1 units): 1,3-cycloheptadiene, (1.31 ± 0.04) × 10-10; 1,3,5-cycloheptatriene, (9.12 × 0.23) × 10-11; cis-1,3,5-hexatriene, (1.04 ± 0.07) × 10-10; and trans 1,3,5-hexatriene, (1.04 ± 0.17) × 10-10. These data, which are the first reported values for these di- and tri-alkenes, are discussed in the context of previously determined O3 and OH radical rate constants for alkenes and cycloalkenes.

Notes: Kinetics of the basic hydrolysis of 1,2-glyceryl dinitrate (1,2-DNG) and 1,3-glyceryl dinitrate (1,3-DNG) esters were investigated in CO2-free aqueous calcium hydroxide solutions. The hydrolysis reactions were carried out in a temperature controlled reactor vessel with provision for continuous N2 sparging of the reaction mixture. Both glyceryl dinitrate esters hydrolyzed via second-order reaction at 25°C. 1,2-DNG in basic solutions isomerized to 1,3-DNG which subsequently hydrolyzed to yield products. The main hydrolysis product of 1,3-DNG was identified as glycidyl nitrate. Other products formed during the basic hydrolysis of DNGs were nitrites and nitrates.

Notes: From the results of order of magnitude analyses, it is concluded that during the oxidation of toluene, radical-atom and radical-radical reactions (1) and (3) play an unusually important and approximately equal role in the formation of benzaldehyde, an intermediate that leads eventually to the complete removal of the side chain. An additional radical-radical system, reaction (2), is shown to be the most likely source of benzyl alcohol observed during toluene oxidation.

Notes: The kinetics of reactions involving halogen atom abstraction from haloalkanes by methyl radicals have been studied in the gas phase. Arrhenius parameters for halogen atom transfer were determined relative to those for methyl radical combination: TextRXlog10A2(L/mol · s)E2(kcal/mol)CFCl38.3 ± 0.210.7 ± 0.4CF3CCl37.9 ± 0.39.7 ± 0.6CF2Cl29.1 ± 0.411.3 ± 0.7CF3Cl8.8 ± 0.511.8 ± 1.0CF3CF2Cl8.3 ± 0.310.9 ± 0.7CF3Br8.6 ± 0.29.3 ± 0.5CF3I8.1 ± 0.14.3 ± 0.2CH3CH2I8.9 ± 0.37.4 ± 0.6The rate data obtained are used to provide information on the importance of polar effects for halogen abstraction processes.

Notes: The advantages and disadvantages of various methods of parametric sensitivity analysis in chemical kinetic modeling are discussed. Particular attention is given to estimates of computational labor for realistic problems, and quantitative comparisons are made utilizing a 52-reaction, 11-species CO oxidation mechanism. The authors′ CHEMSEN/AIM program compares favorably to other techniques in many circumstances, and provides the additional convenience of accepting input information in familiar chemical notation. This paper also reviews recent developments in theory of sensitivity analysis, relevant to chemical kinetic modeling.

Notes: The kinetic study of alkaline hydrolysis of 5, 5-diethyl barbituric acid has been carried out at various [OH] and different temperatures ranging from 85-95°C. The reaction follows an irreversible first-order consecutive reaction path of the type \documentclass{article}\pagestyle{empty}\begin{document}${\rm A}\buildrel{{k1{\rm obs}}}\over\longrightarrow{\rm B}\buildrel{{k2{\rm obs}}}\over\longrightarrow$\end{document} X under pseudo-first-order kinetic conditions. A, B, and X represent for 5, 5-diethyl barbituric acid, diethyl malonuric acid, and ammonia, respectively. The pH-rate profiles obtained at three different temperatures reveal distinct phases which are attributed to a change in rate-determining step with change in [OH]. On the basis of the observed data, a possible mechanism has been discussed.

Notes: The effect of particle size, type of crucible, and heating rate on the thermal curves obtained simultaneously for CuSO4 · 5H2O were discussed. The dissociation steps were confirmed. Thermogravimetric techniques for determining the rate-controlling processes and kinetic parameters were applied for the dehydration steps and the calcination of CuSO4 and CuSO4 · CuO. For the dehydration of the monohydrate one mechanism operates but the activation energy and preexponential factor vary over wide ranges. Differentiating between various mechanisms using the same technique was sometimes difficult giving completely different values for the kinetic parameters. In view of such difficulties the various methods were assessed, the best techniques to treat similar results were recommended and the operating mechanisms and kinetic parameters for the various steps were thus established.

Notes: The kinetics of the gas phase reactions of NO2 with a series of organics have been studied at 295 ± 2 K. It was observed that only 2,3-dimethyl-2-butene and the conjugated dialkenes studied reacted at observable rates, with rate constants which ranged from 1.5 × 10-20 cm3 molecule-1 s-1 for 2,3-dimethyl-2-butene to 1.3 × 10-17 cm3 molecule-1 s-1 for α-phellandrene. These rate constants are compared with the available literature data and the mechanisms of these reactions are discussed.

Notes: The reactions of some perfluoroalkyl radicals with carbon tetrachloride have been studied using the photolysis of the corresponding perfluoroalkyl iodide as the free radical source. The Arrhenius parameters, based on the value of 2.3 × 1013 cm3 mol-1 s-1 for the self-combination rate constant of all radicals are: TextReactionlog(A/cm3 mol-1 s-1)E/kcal mol-1CF3 + CCl412.811.3C2F5 + CCl412.811.6n-C3F7 + CCl412.912.0

Notes: The decomposition kinetics of ethylsilane under shock tube conditions (PT ca. 3100 torr, T ≃ 1080-1245 K), both in the absence and presence of silylene trapping agents (butadiene and acetylene) are reported. Arrhenius parameters under maximum butadiene inhibition are: log k(C2H5SiH3) = 15.14-64,769 ± 1433 cal/2.303 RT; log k(C2H5SiD3) = 15.29-66,206 ± 1414/2.303 RT. The uninhibited reaction is subject to silylene induced decomposition (63% lowest T -- 24% highest T). Major reaction products are ethylene and hydrogen, consistent with two dominant primary dissociation reactions: C2H5SiD3 → C2H5SiD + D2, φ ≃ 0.66; C2H5SiD3 → CH3CH = SiD2 + HD, φ ≃ 0.30. Minor products suggest several other less important primary processes: alkane elimination, φ ≃0.02, and free-radical production via simple bond fission, φ ≃0.02. An upper limit for the activation energy of the decomposition, C2H5SiH → C2H4 + SiH2, of E 〈 30 ± 4 kcal is established, and speculations on the mechanism of this decomposition (concerted or stepwise) with conclusions in favor of the stepwise path are made. Computer modeling studies for the reaction both in the absence and presence of butadiene are shown to be in good agreement with the experimental observations.

Notes: The kinetics and mechanism of the thermal decomposition of n-propylsilane have been studied by the single pulse shock tube-comparative rate technique at pressures around 4700 torr between 1095-1240 K. The primary dissociation processes are 1,1 and 1,2 H2 elimination with ø1,1 ⋍ 0.75 and ø1,2 ⋍ 0.25, respectively. Subsequent decompositions of the primary process product, n-propylsilylene, to propylene and ethylene is complete even in the presence of excess butadiene. Possible mechanistic paths for these decompositions are discussed and an activation energy range of 30 ± 4 kcal is established for both processes. Induced decomposition via silylene chains accounts for 36-46% of the overall reaction in the uninhibited decomposition of n-propylsilane. The silylene chains are quenched in excess butadiene, and studies under maximum inhibition give overall decomposition kinetics of, log k(nPrSiD3, s-1) = 15.26-65,300 ± 1950 cal/2.303. Computer modeling results of the overall reaction both in the absence and presence of butadiene are also presented and shown to be in acceptable agreement with the experimental observations.

Notes: An efficient process for converting the energy stored in electronically excited NF(a1Δ) into blue-green BiF (A—X) band emissions has been discovered. Bismuth atoms are converted to BiF and then repeatedly pumped to the (AO+) state in two steps via collisions with NF(a1Δ). A model has been formulated that predicts BiF emission intensities that are in excellent agreement with the observed values. Some of the rate coefficients required for this model are estimated by means of charge-transfer theories. The cyclic nature of this system, along with its other basic properties, suggests that it has a strong potential to be an efficient blue-green laser system.

Notes: Rate constants have been determined at 296 ± 2 K for the gas phase reaction of NO3 radicals with a series of aromatics using a relative rate technique. The rate constants obtained (in cm3 molecule-1 s-1 units) were: benzene, 〈2.3 × 10-17; toluene, (1.8 ± 1.0) × 10-17; o—xylene, (1.1 ± 0.5) × 10-16; m—xylene, (7.1 ± 3.4) × 10-17; p—xylene, (1.4 ± 0.6) × 10-16; 1,2,3-trimethylbenzene, (5,6 ± 2.6) × 10-16; 1,2,4-trimethylbenzene (5.4 - 2.5) × 10-16; 1,3,5-trimethylbenzene, (2.4 ± 1.1) × 10-16; phenol, (2.1 ± 0.5) × 10-12; methoxybenzene, (5.0 ± 2.8) × 10-17; o-cresol, (1.20 ± 0.34) × 10-11; m-cresol, (9.2 ± 2.4) × 10-12; p-cresol, (1.27 ± 0.36) × 10-11; and benzaldehyde, (1.13 ± 0.25) × 10-15. These kinetic data, together with, in the case of phenol, product data, suggest that these reactions proceed via H-atom abstraction from the substituent groups. The magnitude of the rate constants for the hydroxy-substituted aromatics indicates that the nighttime reaction of NO3 radicals with these aromatics can be an important loss process for both NO3 radicals and these organics, as well as being a possible source of nitric acid, a key component of acid deposition.

Notes: The kinetics and equilibria of the reaction: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm Br} + {\rm CH}_{\rm 3} {\rm CH}_{\rm 2} {\rm OCH}_{\rm 2} {\rm CH}_{\rm 3} {\rm}\mathop {{\rm \rightleftharpoons}}\limits^{\rm 1} {\rm HBr} + {\rm CH}_{\rm 3} {\rm CHOCH}_{\rm 2} {\rm CH}_{\rm 3} $$\end{document} have been studied in the temperature range 298-333 K by using the very low pressure reactor (VLPR) technique. Combining the estimated entropy change of reaction (1), ΔS°1 = 8.1 ± 1.0 eu, with the measured ΔG°1, we find ΔH°1 = 4.2 ± 0.4 kcal/mol; ΔH°f(CH3CHOC2H5) = -20.2 kcal/mol, and DH° [Et OCH(Me)-H] = 91.7 ± 0.4 kcal/mol. We find: \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k_1 (1/{\rm mol s)} = 10.8(\pm 0.7) - (3.9 \pm 4.0)/\theta $$\end{document} where θ = 2.3 RT in kcal/mol. It has been shown that the reaction proceeds via a loose transition state and the “contact TS” model calculation gives a very good agreement with the observed value.

Notes: Chemiluminescence from the a1Δ and b1Σ+ excited electronic states of nitrogen halide diatomics is observed when HN3 is allowed to react with mixtures of halogen atoms in a discharge-flow apparatus. Excited NF (a1Δ) is produced by the F + HN3 reaction, and NCl (a1Δ, b1Σ+) and NBr (a1Δ, b1Σ+) are produced by the F, Cl, + HN3 and F, Br + HN3 reactions, respectively. In the low-density limit, the yield of NF(a1Δ) was found to be near unity. The yields of the b1Σ+ states of NCl and NBr were determined to have a lower limit of ca. 10%. A number of results from these experiments, including direct observation of N3 radicals in the flow, support a hypothetical mechanism in which N3 acts as an intermediate. A second possible mechanism proceeding via an HNF intermediate cannot be ruled out.

Notes: A mechanism for the formation in a chain of H2, CO, and HCOOH in the photooxidation of formaldehyde is proposed. This mechanism is initiated by the addition of HO2 to formaldehyde. Hydrogen atoms are produced by the thermal dissociation of the HOCH2O radical: HOCH2O → H + HCOOH; ΔH = + 3.2 kcal/mol [5]. Photolysis of CH2O—O2—NO mixtures and product analysis were carried out in conjunction with kinetic simulation yielding an estimate for the activation energy of the dissociation reaction : E5 = 14.9 ± 1.0 kcal/mol. Previous observations of this chain process are considered in view of this mechanism.

Notes: The yields of C5 and C6 alkyl nitrates from neopentane, 2-methylbutane, 2-methylpentane, 3-methylpentane, and cyclohexane have been measured in irradiated CH3ONONO-alkane-air mixtures at 298 ± 2 K and 735-torr total pressure. Additionally, OH radical rate constants for neopentyl nitrate, 3-nitro-2-methylbutane, 2-nitro-2-methylpentane, 2-nitro-3-methylpentane, and cyclohexyl nitrate, relative to that for n-butane, have been determined at 298 ± 2 K. Using a rate constant for the reaction of OH radicals with n-butane of 2.58 × 10-12 cm3 molecule-1 s-1, these OH radical rate constants are (in units of 10-12 cm3 molecule-1 s-1): neopentyl nitrate, 0.87 ± 0.21; cyclohexyl nitrate, 3.35 ± 0.36; 3-nitro-2-methylbutane, 1.75 ± 0.06; 2-nitro-2-methylpentane, 1.75 ± 0.22; and 2-nitro-3-methylpentane, 3.07 ± 0.08. After accounting for consumption of the alkyl nitrates by OH radical reaction and for the yields of the individual alkyl peroxy radicals formed in the reaction of OH radicals with the alkanes studied, the alkyl nitrate yields (which reflect the fraction of the individual RO2 radicals reacting with NO to form RONO2) determined were: neopentyl nitrate, 0.0513 ± 0.0053; cyclohexyl nitrate, 0.160 ± 0.015; 3-nitro-2-methylbutane, 0.109 ± 0.003; 2-nitro-2methylbutane, 0.0533 ± 0.0022; 2-nitro-2-methylpentane, 0.0350 ± 0.0096; 3- + 4-nitro-2-methylpentane, 0.165 ± 0.016; and 2-nitro-3-methylpentane, 0.140 ± 0.014. These results are discussed and compared with previous literature values for the alkyl nitrates formed from primary and secondary alkyl peroxy radicals generated from a series of n-alkanes.

Notes: By photolyzing azomethane over the temperature range 331-491 K in the presence of trifluoroacetone the kinetics of the addition reaction (1), ĊH3 + CF3COCH3 → CF3C(Ȯ)(CH3)2 have been studied. Detailed analyses have shown that the principal product of the adduct radical, CF3C(Ȯ)(CH3)2, is CH3COCH3 from reaction (-2), CF3C(Ȯ)(CH3)2 → CH3COCH3 + ĊF3. The rate constant of the addition reaction has been determined to be k1(dm3/mol s) = (4.5 ± 1.4) × 107 exp(-(3370 ± 120)/T) over the temperature range 331-491 K, based on the value k3 = 2.2 × 1010 dm3/mol s for the reaction (3), 2ĊH3 → C2H6. The results are discussed in relation to existing data for radical additions to groups.

Notes: The photolysis of trans-3,4-dimethylcyclopentanone has been studied in the gas phase, principally at 313 nm. However, a few experiments have also been performed using laser sources at 308 and 325 nm. Additionally, experiments were also carried out using the cis isomer. The major products produced by all three wavelengths and in the temperature range 100 to 150°C were propene, 1,2-dimethylcyclobutane, carbon monoxide, and 3,4-dimethylpent-4-en-al. The formation of the cyclobutane was stereospecific and the effects of temperature, pressure, and wavelength on the relative product yields could be rationalized in terms of a mechanism involving the formation of a vibrationally excited triplet state which could yield both hydrocarbons and the aldehyde and a nonexcited triplet yielding only the aldehyde. Some high intensity experiments with an exciplex laser at 308 nm gave results which could be due to the occurrence of some two photon absorption by the cyclopentanone or absorption by an excited intermediate. The results are compared with those previously reported for other substituted cyclopentanones.

Notes: The use of the KMS method for the evaluation of the exponent by the experimentalist requires that a lag or retardation be chosen. The best choice of lag is discussed quantitatively. Its dependence upon the decade range of the data, the relative magnitude of the background, the level and character of the noise, the number and distribution of the data, and the weighted least-squares method of analysis are all detailed. A table of lags and error factors is given, for various ranges, backgrounds, and kinds of noise. Bias in the extracted rate (k), due to noise and treatment of the data, is also discussed.

Notes: The average downward collisional energy transfer (〈ΔEdown〉) is obtained for highly vibrationally excited tert-butyl chloride, both undeuterated and per-deuterated, with Kr, N2, CO2, and C2H4 bath gases, at ca. 760 K. Data are obtained using the technique of pressure-dependent very low-pressure pyrolysis. Reactant internal energies to which the data are sensitive are in the range 200-250 kJ mol-1. For C4H9Cl, the 〈ΔEdown〉 values (cm-1) are 255 (Kr), 265 (N2), 440 (CO2), and 585 (C2H4), and for C4D9Cl, 245 (N2), 370 (CO2), and 540 (C2H4). The uncertainties in these values are ca. 20% (40% for Kr); the uncertainties in the deuteration ratios are 10-15%. The value for Kr is in agreement with theoretical predictions of a biased random walk model for internal energy change in monatomic/substrate collisions. The effect of deuteration of 〈ΔEdown〉 is also in accord with that predicted by a modification of the theory. Extrapolated highpressure rate coefficients for the thermal decomposition of reactant are 1013.6 exp(-187 kJ mol-1/RT) s-1 (C4H9Cl) and 1014.2 exp(-196 kJ mol-1/RT) s-1 (C4D9Cl), in accord with other studies and the expected isotope effect.

Notes: Flash photolysis technique has been used to obtain the rate and thermodynamic parameters of the reversible dimerization reactions of a range of ten phenoxy radicals (I-X) in a toluene-dibutylphthalate mixture (0.6 cP ≤η≤18.4 cP): \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm R}^{.} + {\rm R}^{.} {\mathop{{\buildrel{-\!\!\longrightarrow}\over{\longleftarrow}}}\limits_{k_{-1}}^{k_1}}{\rm D} $$\end{document} The main reason for the difference in the k1 values are the different steric hindrances in radicals. It has been found that the values of k1 for 2,6-diphenyl-4-methoxy- (I), 2-phenyl-(III), and 2-methoxyphenoxy (IV) radicals are 3-5 times smaller than the respective diffusion constants calculated according to the Debye formula with regard to the spin-statistical factor: \documentclass{article}\pagestyle{empty}\begin{document}$$ k_{diff} = \sigma \frac{{8{\rm RT}}}{{3000{\rm \eta }}} $$\end{document} The resultant ΔH1≠values for these radicals in toluene and dibutylphthalate are close to the activation energies of the viscous flow of the solvents B. Linear relationships with a slope equal to unity are observed between log k1 and log(T/η). The recombination of radicals I, III, and IV is limited by translational diffusion. The k1 values for 2,6-diphenyl- (VII), 2,6-di-tert-butyl- (IX), and 2,6-di-tert-butyl-4-methylphenoxy (X) radicals are 10-60 times smaller than kdiff and Δ H≠ B. In the case of radical X in toluene ΔH1≠ 0. The recombination of these three radicals includes an intermediate step of complex formation: \documentclass{article}\pagestyle{empty}\begin{document}$${{\rm R}^\cdot+{\rm R}^\cdot}{\mathop {{\scriptstyle\longleftarrow}^{\hskip-13pt\longrightarrow}}}{\rm R^\cdot}\ldots {\rm R}^\cdot \rightarrow {\rm D}$$ \end{document} For 4-phenyl- (II), 2,6- dimethoxy- (V), 2,4-diphenyl- (VI), and radicals VII, IX, and X the linear relationships between log k1 and log (T/η) have a slope of from 0.5 ± 0.05 to 0.8 ± 0.05. The k1-1 versus η relationships for these radicals are not straight lines. The recombination of these six radicals is limited by translational and rotational diffusion. With the aid of theoretical models, the k1 versus η relationships have been used to derive the steric factor f in radical recombination and the angle θ between the axis and the solid angle generatrix. The solid angle defines the reaction spot on the radical-sphere surface. The recombination of the 2,6-diphenyl-4-diphenylmethylphenoxy radical (VIII) takes place in the region intermediate between the diffusion and the kinetic ones, and the relationship between log k1 and log (T/η) for this radical has a plateau portion. The log k-1 versus log (T/η) relationships have precisely the same form as the corresponding k1 relationships, which is quite in line with the theory of diffusion-controlled reversible recombination reactions.

Notes: The decomposition of CH3CD2CH3 was studied from 713 to 853 K at pressures of 98-466 torr. The values of k1/k2 = 2.08 ± 0.05 and k3/k4 = 2.04 ± 0.66 were found independent of temperature by measuring the ratios of CH4/CH3D and CH3CHD2/CH3CD3, respectively, for the following reactions:. Isomerization of CH3CDCH3 was detected by measuring CHDCH2 formed from the isomerized radical. The expression of k21/k22 was found to be where k21 and k22 are the rate constants of. The results and conclusions are discussed and compared with previous works.

Notes: A detailed investigation of the mechanism of cyanogen oxidation is presented. Recent induction time measurements of ignition in cyanogen-oxygen-argon mixtures behind reflected shocks are computer modeled to obtain an agreement between the experimental and calculated values. A 15-step reaction scheme is suggested to reproduce the parameters E and βi in the experimental parametric relation: τ = 10αexp(E/RT)IICiβi. An explanation is offered to the very strong dependence of the induction time on the cyanogen concentration and the very weak dependence on the oxygen concentration. The sensitivity spectrum shows that the induction time is highly dependent on the O + C2N2 → NCO + CN and NCO + M → N + CO + M reactions (shortened) and the O + NCO → CO + NO and N + NCO → N2 + CO reactions (increased).

Notes: Atmospheric photodissociation rate coefficients and photodissociation lifetimes for nitromethane, methyl nitrite, and methyl nitrate were calculated as a function of altitude from their measured visible and near ultraviolet photoabsorption cross sections at 298 K. The lifetime of methyl nitrite is nearly independent of altitude and is approximately 2 min. From 0 to 50 km the lifetime of nitromethane varies from 10 to 0.5 hr, while that of methyl nitrate changes from 5.3 to 0.09 days, respectively.

Notes: Ethyl N—methylcarbamate decomposes thermally over the temperature range of 600-650 K by competing first-order reactions, one forming methylamine, carbon dioxide, and ethylene, the other forming methyl isocyanate and ethanol. The first-order rate constants are described in S—1 units by the equations where R = 1.986 cal/deg mol. The appareance of sym—dimethylurea among the products raised the possibility of gas-phase transesterifications. These were ruled out by the study of the reactions of sym-dimethylurea at 604 K which showed its behavior to be well explained by the rapid decomposition in the gas phase which is reversed in the condensation stage in the analysis.

Notes: The law c = c0 exp(—K √t) in the alkyl radical abstraction reaction is affected by neither the matrix annealing nor the way of the radical generation. If the reaction runs at varying temperature, a dimensionless time τ which does determine unambiguously the degree of conversion can be introduced.

Notes: The kinetics of the gamma-radiation induced free radical reactions in carbon tetrachloride solutions of 2,3-dimethylbutane (DMB) were studied in the temperature range 32-118 °C. The kinetics of the following reactions were measured:\documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} {\rm CCl}_3 + {\rm RH} \to {\rm CHCl}_3 + {\rm R} \\ {\rm CCl}_{\rm 3} + {\rm CCl}_{\rm 3} \to {\rm C}_2 {\rm Cl}_6 \\ \end{array} $$\end{document} and the following rate constants expression was obtained:\documentclass{article}\pagestyle{empty}\begin{document}$$ \log k_2 /k_3^{1/2} (M^{ - 1/2} \sec ^{ - 1/2} )\, = \,(2.40 \pm 0.17) - (6.96 \pm 0.26)/{\rm \theta } $$\end{document}. The activation energy and the A factor obtained are in good agreement with the values obtained at the gaseous phase, considering the activation energy for self-diffusion of CCl4.

Notes: Measurement of the rate of the reaction is reported. The measurements were made in a flow tube apparatus. The result is based on data for the absolute density of OH(v = 0) obtained from laser-induced fluorescence measurements in the (0-0) band of the OH(A2Σ+ → X2II) system. The density of oxygen atoms was varied by changing the flow rate of NO which is consumed in the reaction N + NO → O + N2. We find that k1 (298 K) = (5.5 ± 3.0) × 106 cm3/mol sec. This result was obtained with consideration and control of the effect of reaction (2): for which vibrationally excited hydrogen is created by energy transfer in the presence of active nitrogen. It was found that the addition of N2 or CO2 effectively suppressed the excitation of H2(v = 1). Measurements of the density of H2(v = 1) were made by VUV absorption in the Lyman band system of H2. All of the reports of low-temperature measurements and some recent theoretical calculations for k1 are discussed. The present result confirms and extends the growingevidence for significant curvature in the low-temperature end of a modified Arrhenius plot of k1 (T).

Notes: The thermal decomposition of ammonia was studied by means of the shock-tube and vacuum ultraviolet absorption spectroscopy monitoring the concentration of atomic hydrogen. The rate constants of both the initiation reaction and the consecutive reaction were determined directly as \documentclass{article}\pagestyle{empty}\begin{document}$$ k_1 = 10^{16.14} \exp (- 90.6{\rm kcal}/RT){\rm cm}^{\rm 3} /{\rm molsec} $$\end{document} and \documentclass{article}\pagestyle{empty}\begin{document}$$ k_{II} = 10^{14.30} \exp (- 23.29{\rm kcal}/RT){\rm cm}^{\rm 3} /{\rm molsec} $$\end{document} respectively.

Notes: O2 in the A3Σu+ state has been prepared in a discharge flow system by recombining oxygen atoms on a nickel surface. The decay of this excited state was followed by observing the emission between 280 and 400 nm. The wall deactivation was observed to approach unit efficiency. Rate constants were determined to be 0.9 × 10-11, 2.9 × 10-13, and 8.6 × 10-16 cm3/molecule sec for the quenching of O2(A3Σu+) by O, O2, and Ar, respectively.

Notes: The rate constant for the reaction CH3O2 + NO2 → (products) has been measured directly by flash photolysis and kinetic spectroscopy. At room temperature and at total pressures between 53 and 580 Torr, k3 = (9.2 ± 0.4) × 108 liter/mole sec so that the rate of formation of the probable primary product peroxymethyl nitrate (CH3O2NO2) may be significant in urban atmospheres.

Notes: Existing data on the self-reactions of tertiary peroxy radicals RO2 has been reanalyzed and corrected to deduce Arrhenius parameters for both termination and nontermination paths. For R = t-Butyl, these are logkt(M-1sec-1) = 7.1 - (7.0/θ) and logknt(M-1sec-1) = 9.4 - (9.0/θ), respectively, different from those recommended by other authors. The higher magnitudes observed for termination processes of tertiary peroxy radicals like those of cumyl and 1,1-diphenylethyl have been discussed in terms of a much greater cage recombination of cumyloxy radicals as contrasted with t-butoxy radicals. It is shown that for benzyl peroxy radicals, the R - O·2 bond dissociation energy is sufficiently low (18-20 kcal) that reversible dissociation into R· + O2 opens a competing second-order path to fast recombination R· + RO·22 → ROOR. This path is probably not important for cumyl peroxy radicals under usual experimental conditions but can become important for 1,1-diphenyl ethyl peroxy radicals at (O2) 〈 10-3M. At very low RO·2 concentrations (〈10-5M), in the absence of added O2, an apparent first-order disappearance of RO·2 can occur reflecting the rate determining breaking of the cumyl - O·2 bond followed by the second step above. The thermochemistry of RO·n is used to show that the reaction of R2O4 → 2RO + O2 must be concerted and cannot proceed via RO·3 which is too unstable and cannot form even from RO· + O2.

Notes: Demetallation rates of α,β,γ,δ-tetrakis(p-sulfophenyl)porphiniron(III) in hydrochloric acid-ethanol-water, perchloric acid-ethanol-water, and sulfuric acid-alcohol-water media were determined. For a given acidity value H0 the order of the rates for the three acids was HCl 〉 H2SO4 〉 HClO4. This is also the order for complex formation between acid anion and iron(III). Consequently ligands as well as protons are involved in the breaking of bonds between the metal and the porphyrin leading to the formation of the activated complex. The log k values for HCl and HClO4 media were not linearly related to the Hammett acidity function as they were for sulfuric acid-ethanol-water media. The average ΔH

Notes: A method is presented for the prediction of rate coefficients and Arrhenius parameters for bimolecular hydrogen atom transfer reactions A + BC → AB + C. The treatment sets out from structural considerations of the complex A ⃛ B ⃛ C and calculates the energy of the complex along the reaction path from empirical functions for a bonding energy term and an endgroup contribution. The treatment proceeds by assuming ultrasimple transition state models and assigning the force constants and vibrational frequencies. Finally the rate coefficient and Arrhenius parameters are obtained on the basis of separable activated complex theory. Application of the method requires known properties of reactant and product molecules and does not demand the use of adjustable parameters. The relation and differences between this method and the BEBO treatment as well as Zavitsas' method are dealt with.

Notes: The rate of the reaction was determined in an isothermal discharge flow reactor with a combined ESR-LMR detection under pseudo-first-order conditions in HO2. The rate constant was identical in experiments with two different HO2 sources: F + H2O2 and H + O2 + M. The absolute rate constant at T = 293 K was measured as \documentclass{article}\pagestyle{empty}\begin{document}$$ k_1 (293K) = (4.6 \pm 1)10^{12} cm^3 /mol\sec $$\end{document} In the range 2 ≤ p mbar ≤ 17 no pressure dependence for k1 was found.

Notes: The addition reactions of CCl3 radicals with cis-C2Cl2H2, trans-C2Cl2H2, and C2Cl3H in liquid cyclohexane-CCl4 mixtures were studied between 323 and 448 K. The Arrhenius parameters of these reactions were competitively determined versus H-atom transfer from cyclohexane and addition to C2Cl4. The present data and the data obtained in previous liquid and gas phase studies show that the reactivities displayed in addition reactions of different radicals with chloroethylenes reflect primarily variations in activation energies rather than in A factors. The activation energies for the addition of CCl3, CF3, and CH3 radicals to chloroethylenes appear, to a large extent, to be determinedby the stability of the adduct radicals. Comparison of the reactivity trends in the addition reactions of chloro- and fluoro-substitutedethylenes indicates that these two electron-withdrawing substituentshave a converse effect on the reactivity of electrophilic radicals. This behavior is ascribed to the strong mesomeric effect of vinylic chlorosubstituents.

Notes: The titration of chemisorbed oxygen by carbon monoxide to form carbon dioxide has been studied from 373 to 673 K over polycrystalline platinum. The pressure transients for CO and CO2 have been measured and simulated numerically. A complex Langmuir-Hinshelwood mechanism is found which fits all the data, and it is not necessary to invoke Eley-Rideal kinetics. The results fall into two temperature regimes, above and below 473 K, which are characterized by different Arrhenius parameters. A change in activation energy with oxygen coverage is also found below 473 K.

Notes: The kinetics of dimethyl sulfoxide (DMSO) oxidation by peroxomonophosphoric acid (PMPA) in aqueous medium at 308 K and I = 0.4 mol/dm3 follow the rate expressions In the pH range from 0 to 2, where k1 and k2 are 5.092 × 10-1 dm3/mol sec and ≃ 0, respectively; in the pH range from 4 to 7, where k2 = 8.127 × 10-3 and k3 = 2.90 × 10-3 dm3/mol sec; and in the pH range from 10 to 13.6, where k4 ≃ 0, and k5 = 3.08 × 10-2 dm3/mol sec.The reaction is interpreted in terms of mechanisms involving an electrophilic and a nucleophilic attack of the peroxomonophosphoric acid species, respectively, in acid and alkaline regions, on the sulfur atom of the sulfoxide molecule giving rise to SN2-type transition states followed by oxygen-oxygen bond fission to form the products.

Notes: The oxidations of ferrocene (FcH) and n-butylferrocene (FcBu) by ferric salts (nitrate or bromide) are strongly inhibited by aqueous cetyltrimethylammonium bromide and nitrate (CTABr and CTANO3, respectively). The kinetics of inhibition fit a model in which the substrates are distributed between water, and the micelles and binding constants Ks to the micelle can be estimated. The oxidations are strongly catalyzed by micelles of sodium lauryl sulfate (NaLS), and the kinetics can be fitted to a model in which the reaction rate depends upon the concentration of both reactants in the micellar pseudophase and the rate constants in that pseudophase, which for both substrates are very similar to those in water. Some added salts reduce the micellar catalysis by excluding ferric ions from the micelle. The oxidations of FcH and FcBu by ferricyanide ions are too fast to be followed in water, but they are inhibited by anionic micelles of NaLS. By analyzing the rate surfactant profiles using independently measured values of Ks the second-order rate constants in water have been estimated.

Notes: The addition of ethene to cyclohexa-1,3-diene has been studied between 466 and 591 K at pressures ranging from 27 to 119 torr for ethene and 10 to 74 torr for cyclohexa-1,3-diene. The reaction is of the “Diels-Alder” type and leads to the formation of bicyclo[2.2.2]oct-2-ene. It is homogeneous and first order with respect to each reagent. The rate constant (in l./mol sec) is given by\documentclass{article}\pagestyle{empty}\begin{document}$$ \log _{10} k_a = - (25,970 \pm 50)/4.576{\rm T + }(6.66 \pm 0.02) $$\end{document}The retron-Diels-Alder pyrolysis of bicyclo[2.2.2]oct-2-ene has also been studied. In the ranges of 548-632 K and 4-21 torr the reaction is first order, and its rate constant (in sec-1) is given by\documentclass{article}\pagestyle{empty}\begin{document}$$ \log _{10} k_p = - (57,300 \pm 100)/4.576{\rm T + }(15.12 \pm 0.04) $$\end{document}The reaction mechanism is discussed. The heat of formation and the entropy of bicyclo[2.2.2]oct-2-ene are estimated.

Notes: The kinetics of ethane oxidation was studied at 320, 340, 353 and 380°C, mixture composition 2 C2H6 + 1 O2, and total pressure 609 torr. It was found that at 320°C CH2O and CH3CHO were branching agents. A series of experiments was conducted on 2C2H6 + O2 oxidation in the presence of 0.7% 14C-labeled ethylene. The ethylene oxide was found to form only from C2H4, formaldehyde formed from C2H4 and C2H6; and CH3CHO, C2H5OH, and CH3OH formed only from ethane. The formation rates of C2H4, C2H4O, and CH2O were calculated by the kinetic tracer method. At 320°C the fraction of oxygen-containing products formed from C2H4 was 16-18%, and at 353 and 380°C it was 30-40%.

Notes: The thermal unimolecular decomposition of bromocyclobutane has been investigated over the temperature range of 791-1224 K using the technique of very low-pressure pyrolysis (VLPP). HBr elimination is the sole mode of decomposition under the experimental conditions. No evidence could be found for the ring-cleavage pathway to ethylene and vinyl bromide. Assuming a four-center transition state and an Arrhenius A factor the same as that for HCl elimination from chlorocyclobutane, RRKM calculations show that the experimental unimolecular rate constants are consistent with the Arrhenius expression \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k(\,{\rm sec}^{{\rm - 1}} {\rm )}\, = \,(13.6 \pm 0.3) - (52.0 \pm 1.0)/{\rm \theta }\,$$\end{document} where θ = 2.303RT kcal/mol. The activation energy is higher than that for the open-chain analog, 2—bromobutane. This finding is consistent with the results for the corresponding chloro and iodo compounds.

Notes: The kinetics and mechanism of ascorbic acid (DH2) oxidation have been studied under anaerobic conditions in the presence of Cu2+ ions. At 10-4 ≤ [Cu2+]0 〈 10-3M, 10-3 ≤ [DH2]0 〈 10-2M, 10-2 ≤ [H2O2] ≤ 0.1M, 3 ≤ pH 〈 4, the following expression for the initial rate of ascorbic acid oxidation was obtained: where χ2 (25°C) = (6.5 ± 0.6) × 10-3 sec-1. The effective activation energy is E2 = 25 ± 1 kcal/mol. The chain mechanism of the reaction was established by addition of Cu+ acceptors (allyl alcohol and acetonitrile). The rate of the catalytic reaction is related to the rate of Cu+ initiation in the Cu2+ reaction with ascorbic acid by the expression where C is a function of pH and of H2O2 concentration. The rate equation where k1(25°C) = (5.3 ± 1) × 103M-1 sec-1 is true for the steady-state catalytic reaction. The Cu+ ion and a species, which undergoes acid-base and unimolecular conversions at the chain propagation step, are involved in quadratic chain termination. Ethanol and terbutanol do not affect the rate of the chain reaction at concentrations up to ≈0.3M. When the Cu2+-DH2-H2O2 system is irradiated with UV light (λ = 313 nm), the rate of ascorbic acid oxidation increases by the value of the rate of the photochemical reaction in the absence of the catalyst. Hydroxyl radicals are not formed during the interaction of Cu+ with H2O2, and the chain mechanism of catalytic oxidation of ascorbic acid is quantitatively described by the following scheme.Initiation: Propagation: Termination:

Notes: Using published data on the kinetics of pyrolysis of C2Cl6 and estimated rate parameters for all the involved radical reactions, a mechanism is proposed which accounts quantitatively for all the observations:The steady-state rate law valid for after about 0.1% reaction is and the reaction is verified to proceed through the two parallel stages suggested earlier whose net reaction isA reported induction period obtained from pressure measurements used to follow the rate is shown to be compatible with the endothermicity of reaction A, giving rise to a self-cooling of the gaseous mixture and thus an overall pressure decrease.From the analysis, the bond dissociation energy DH0(C2Cl5—Cl) is found to be 70.3 ± 1 kcal/mol and ΔHf3000(·C2Cl5) = 7.7 ± 1 kcal/mol. The resulting π—bond energy in C2Cl4 is 52.5 ± 1 kcal/mol.

Notes: Spectrophotometric methods were used to investigate the rate of the reaction of Br2 with HCOOH in aqueous, acidic media. The reaction products are Br- and CO2. The kinetics of this reaction are complicated by both the formation of Br3- as Br- is formed and the dissociation of HCOOH into HCOO- and H+. Previous work on this reaction was carried out at acidities lower than the highest used here and led to the conclusion that only HCOO- reacts with Br2. It is agreed that this is by far the principal reaction. However, at the highest acidity experiments, an added small component of reaction was found, and it is suggested that it results from the direct reaction of Br2 with HCOOH itself. On this assumption, values of the rate constants for both reactions are derived here. The rate constant for the reaction of HCOO- with Br2 agrees with values previously reported, within a factor of 2 on the low side. The reaction involving HCOOH is more than 2000 times slower than the reaction involving HCOO-, but it does contribute to the overall rate as [H+] approaches 1M. These derived rate constants are able to simulate quantitatively the authors' absorbance-versus-time data, demonstrating the validity of their data treatment methods, if not mechanistic assignments. Finally, activation parameters were determined for both rate constants. The values obtained are: ΔE

Notes: The reaction of carbon monoxide with ozone was studied in the range of 75-160°C in the presence of varying amounts of CO2 and, for a few experiments, of O2. At room temperature the reaction was immeasurably slow, but in a flow system it showed chemiluminescence with undamped oscillations. In a static system above 75°C the emission showed damped oscillations when O2 was present. In the absence ofadded O2 the emission showed a slow decay with a half-life of 1 hr. The luminescence consisted of partially resolved bands in the range of 325-600 nm, and the source was identified as CO2(1B2) → CO2(1Σg+) + hv. The kinetics were complex, and the observed rate law could be accounted for bya mechanism involving the chain sequence \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm O(}^{\rm 3} P{\rm ) + CO( + M)}\mathop {{\rm rightarrow}}\limits^{\rm 3} {\rm CO}_{\rm 2} {\rm (}^{\rm 3} B_{\rm 2} {\rm ) ( + M), CO}_{\rm 2} {\rm (}^{\rm 3} B_{\rm 2} {\rm ) + O}_{\rm 3} {\rm }\mathop {{\rm rightarrow}}\limits^{\rm 7} {\rm CO}_{\rm 2} {\rm (}^{\rm 1} \sum\nolimits_{\rm g}^{\rm + } {} {\rm ) + O}_{\rm 2} {\rm + O} $\end{document}. From measurements of -d[O3]/dtand relative emission, rate constant ratios were obtained and estimates of k3were made.

Notes: The reaction NO + O3 → NO2 + O2 has been studied in a 220-m3 spherical stainless steel reactor under stopped-flow conditions below 0.1 mtorr total pressure. Under the conditions used, the mixing time of the reactants was negligible compared with the chemical reaction time. The pseudo-first-order decay of the chemiluminescence owing to the reaction of ozone with a large excess of nitric oxide was measured with an infrared sensitive photomultiplier. One hundred twenty-nine decays at 18 different temperatures in the range of 283-443 K were evaluated. A weighted least-squares fit to the Arrhenius equation yielded k = (4.3 ± 0.6) × 10-12 exp[-(1598 ± 50)/T] cm3/molecule sec (two standard deviations in brackets). The Arrhenius plot showed no curvature within experimental accuracy. Comparison with recent results of Birks and co-workers, however, suggests that a nonlinear fit, as proposed by these authors, is more appropriate over an extended temperature range.

Notes: The gas-phase free radical displacement reaction has been studied in the temperature range of 240-290°C and at 140°C with the thermal decomposition of azomethane (AM) and di-tert-butylperoxide (DTBP), respectively, as methyl radical sources. The reaction products of the CD3 radicals were analyzed by mass spectrometry. Assuming negligible isotope effects, Arrhenius parameters for the elementary radical addition reaction were derived: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_1 (cm^3 /mol\sec) = (10.5 \pm 0.4) - (11,500 \pm 1100)/4.576T $$\end{document}The data are discussed with respect to the back reaction and general features of elementary addition reactions.

Notes: The redox potential and iodine concentration behavior of the title reaction and component reactions have been examined. The effect of hydrogen peroxide, potassium iodate, manganese (II) sulfate, sulfuric acid, and acetone concentration on the time period and redox potential behavior is reported. Iodine production and consumption rates for the component reactions are given, and some mechanistic suggestions, involving iodine dioxide as the one electron oxidant, are made.

Notes: The production of both the b1Σ+ and a1Δ states of NCl has been observed from the reaction of HN3 with flowing streams of Cl and F atoms. The results suggest that a two-step reaction sequence is responsible for the production of excited NCl, as follows: The rate contant (all products) for the first step is k(F + HN3) 〉 1 × 10-11 cm3/molecule sec. Comparison of this value to results obtained in a previous study of the F + HN3 system yields a value k(F + N3) = 2 × 10-12 cm3/molecule sec. The rate constant for the reaction of chlorine atoms with HN3 was determined to be k(Cl + HN3) 1 × 10-12 cm3/molecule sec. The difference between the Cl + HN3 and F + HN3 rates is interpreted in terms of an addition-elimination mechanism.

Notes: The kinetics of oxidation of benzyl alcohol and substituted benzyl alcohols by sodium N-chloro-p-toluenesulfonamide (chloramine-T, CAT) in HClO4 (0.1-1 mol/dm3) containing Cl- ions, over the temperature range of 30-50°C have been studied. The reaction is of first order each with respect to alcohol and oxidant. The fractional order dependence of the rate on the concentrations of H+ and Cl- suggests a complex formation between RNCl- and HCl. In higher acidic chloride solution the rate of reaction is proportional to the concentrations of both H+ and Cl7hyphen;. The observed solvent isotope effect (kD2O/kH2O) is 1.43 at 30°C. The reaction constant (p = -1.66) and thermodynamic parameters are evaluated. Rate expressions and probable mechanisms for the observed kinetics have been suggested.

Notes: The kinetics and mechanism of the silver(II) oxidation of methanol, ethanol, 1-propanol, 1-methyl- ethanol, 1-butanol, 2-methyl-1-propanol, 2-butanol, 2-methyl-2-propanol, D4-methanol, and D6-methanol have been investigated at 8.0 and 20.0°C in aqueous perchloric acid media (1.00 ≤ [HClO4] ≤ 4.00M; μ = 4.0M). The kinetics were monitored by following the disappearance of Ag(II) with a spectrophotometric stopped-flow technique. The reactions are first order in each reactant and involve both Ag2+ and AgOH+ species. No kinetic or spectroscopic evidence for complex formation between reactants was obtained. The results are discussed with reference to electron density on the —OH or αC-H substrate sites and to the isotopic hydrogen/deuterium rate quotients found for methanol and ethanol.

Notes: The carbon kinetic isotope effect in the reaction of CH4 with OH has an experimentally measured value of 1.003. The measurement was performed using a static system in which the source of OH was the gas-phase photolysis of H2O2 with ultraviolet light produced by a high-pressure mercury arc lamp. Implications for the tropospheric cycle of CH4 are considered briefly.

Notes: It was found that the rates of absorption of oxygen by pyridine-aqueous sodium hydroxide emulsions and the same emulsions containing benzil were catalyzed by the addition of quaternary salts and followed the same rate law:\documentclass{article}\pagestyle{empty}\begin{document}$$ - \frac{{dPo_2}}{{dt}} = k_1 \left[{{\rm Po}_{\rm 2}} \right]\left[{^ -} \right]\left[{{\rm benzil}} \right]^0 + k_2 \left[{{\rm Po}_2} \right]\left[{OH^ -} \right]\left[{{\rm Et}_{\rm 4} {\rm NCl}} \right][benzil]^0 $$\end{document}. It was concluded that the autooxidation of benzil in pyridine-aqeuous sodium hydroxide emulsions has as its rate-determining step a step in the autooxidation of pyridine. Possibly the superoxide formed in the autooxidation of pyridine is the oxidizing agent in the oxidation of benzil.

Notes: It is shown how electron spin resonance spectroscopy with modulated radical initiation can be used to analyze by purely spectroscopic means the second-order termination kinetics of systems containing two different kinds of radicals. The technique is applied to species generated by photoreduction of acetone in tetraethoxy silane. The bimolecular self- and cross reactions of \documentclass{article}\pagestyle{empty}\begin{document}${\rm (CH}_{\rm 3} {\rm CH}_{\rm 2} {\rm O)}_{\rm 3} {\rm SiO\dot CHCH}_{\rm 3} (\dot R_1 )\,and\,(CH_3 )_2 \dot COH(\dot R_2 )$\end{document} are found to be encounter-controlled processes. For the cross termination the often used relation k12 = (4 k1k2)1/2 is verified experimentally.

Notes: The homogeneous gas-phase thermal decomposition kinetics of germane have been measured in a single-pulse shock tube between 950 and 1060 K at pressures around 4000 torr. The initial decomposition is GeH4 → GeH2 + H2 in its pressure-dependent regime, with log kGeH4(4000) = 13.83 ± 0.78 - 50,750 ± 3570 cal/2.303RT. RRKM calculations suggest that the high-pressure Arrhenius parameters are log k GeH4(M → ∞) = 15.5 - 54,300 cal/2.303RT. Extrapolations to static system pyrolysis conditions (T ∼ 600 K, P ∼ 200 torr) give homogeneous reaction rates which are much slower than those observed, hence the static system pyrolysis of germane must be predominantly heterogeneous. Shock-initiated pyrolysis reaction stoichiometry is 2 mol H2 per mole GeH4, suggesting that the subsequent decomposition of germylene is essentially quantitative. Investigations of the hydrogen product yields for pyrolysis of GeD4 in øCH3 further indicate that the germylene decomposition reaction is mainly GeH2 → H2 + Ge, but that a small amount of reaction to H atoms may also occur.

Notes: The photooxidation of chloral was studied by infrared spectroscopy under steady-state conditions with irradiation of a blackblue fluorescent lamp (300 nm 〈 λ 〈 400 nm, λmax = 360 nm) at 296 ± 2 K. The products were hydrogen chloride, carbon monoxide, carbon dioxide, and phosgen. The kinetic results reveal that the reaction proceeds via chain reaction of the Cl atom:The results lead to the conclusion that mechanism (B) is confirmed to be more likely than mechanism (A), which was favored at one time by Heicklen for the mechanism of the oxidation of trichloromethyl radicals by oxygen molecules: The ratio of the initial rates of CO and CO2 formation gave k7/k6 = 4.23M-1, and the lower limit of reaction (5) was found to be 3.7 × 108M-1 sec-1.

Notes: The kinetics of bromination of several aromatic compounds (anilides, anisoles, and phenols) was investigated in 80% aqueous acetic acid (v/v) in the temperature range 20-50°C using N-bromosuccinimide (NBS) as the reagent. The reaction was found to be first order in the aromatic substrate (ArH), and zero order in NBS, the overall order being 1. Stoichiometry of the reaction was 1:1. An increase in solvent polarity increased the reaction rate, and chloride ions were found to be specific catalysts for the reaction. Arrhenius activation energy remained almost constant for all the substrates. A probable mechanism explaining all these observed facts was proposed. The mechanism involved an attack by Br+ or more probably by a solvated Br+ ion on the aromatic substrate.

Notes: The photooxidation of the 1,3-butadiene-NO-air system at 298 ± 2 K was investigated in an environmental chamber under simulated atmospheric conditions. The irradiation gave rise to the formation of acrolein in a 55% yield, based on 1,3-butadiene initial concentration for all the experimental runs.The rate of formation of acrolein was the same as that of 1,3-butadiene consumption, indicating that acrolein is the major product of the 1,3-butadiene oxidation in air.The dependence of acrolein concentration on irradiation time showed thata secondary process, identified as an oxidation of acrolein by ⋅OH radicals, was occurring during the photochemical runs. The rate constant of this secondary process was determined by measuring the relative rates of disappearance of acrolein and n-butane during the irradiation of acrolein-n-butane-NO-air mixtures. The so obtained relative rate constant value was placed on an absolute basis using a reported rate constant for the n-butane + ⋅OH reaction; a value of (1.6 ± 0.2) × 1010 M-1 sec-1 was obtained.

Notes: The Cl atom-initiated oxidation of CH2Cl2 and CH3Cl was studied using the FTIR method in the photolysis of mixtures typically containing Cl2 and the chlorinated methanes at 1 torr each in 700 torr air. The results obtained from product analysis were in general agreement with those reported by Sanhueza and Heicklen. The relative rate constant for the Cl atom reactions of CH2Cl2 and CH3Cl was determined to be k(Cl +CH3Cl)/k(Cl + CH2Cl2) = 1.31 ± 0.14 (2σ) at 298 ± 2 K.

Notes: A chain mechanism is proposed to account for the very rapid termination reactions observed between alkyl peroxy radicals containing α-C - H bonds which are from 104 to 106 faster than the termination of tertiary alkyl peroxy radicals. The new mechanism is with termination by. \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm R}\overline {{\rm CHOO}} $\end{document} is the zwitterion originally postulated by Criegee to account for the chemistry of O3-olefin addition. Heats of formation are estimated for \documentclass{article}\pagestyle{empty}\begin{document}$ \overline {{\rm CH}_2 {\rm OO,}} {\rm }\overline {{\rm RCHOO}} $\end{document}, and \documentclass{article}\pagestyle{empty}\begin{document}$ ({\rm C}\overline {{\rm H}_3 )_2 {\rm COO}} $\end{document} and it is shown that all steps in the mechanism are exothermic. The second step can account for (1Δ)O2 which has been observed. k1 is estimated to be 109-2/θ liter/M sec where θ = 2.303RT in kcal/mole. The second and third steps constitute a chain termination process where chain length is estimated at from 2 to 10. This mechanism for the first time accounts for minor products such as acid and ROOH found in termination reactions. Trioxide (step 3) is shown to be important below 30°C or in very short time observations (〈10 s at 30°C). Solvent effects are also shown to be compatible with the new mechanism.

Notes: The effect of copper (II) and chloride ions on the manganese (II) catalyzed iodate-peroxide reaction has been examined with reference to the hydrogen peroxide-iodic acid-manganese (II)-organic species oscillatory reaction. The observations are considered to provide evidence for iodine dioxide as the key intermediate in the manganese (II) catalyzed reaction. Kinetic data for the copper (II) catalyzed reaction are reported.

Notes: The pyrolysis of di-tert-butyl sulfide has been investigated in static and stirred-flow systems at subambient pressures. The rate of consumption of the sulfide was measured in some experiments, and the rate of pressure increase was followed in others. The results suggest that the reaction is essentially homogeneous in a seasoned reactor and proceeds through a free radical mechanism. In the initial stages, the decomposition rate follows first-order kinetics, and the rate coefficient in the absence of an inhibitor is given by \documentclass{article}\pagestyle{empty}\begin{document}$$ k_{^u} (\sec ^{ - 1}) = 10^{15.1 \pm 0.6} \exp \left[{\left({ - 229 \pm 8} \right){\rm kJ/mol/RT}} \right] $$\end{document} between 360 and 413°C. The stoichiometry of the uninhibited reaction at 380°C and 50% decomposition is approximately \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm t}^{\rm \_} {\rm C}_{\rm 4} {\rm H}_{{\rm 9}^{\rm -}} {\rm S}_{\rm -} {\rm t}_{\rm -} {\rm C}_{\rm 4} {\rm H}_{\rm 9} = 1.72i_ - {\rm C}_{\rm 4} {\rm H}_{\rm 8} + 0.88{\rm H}_{\rm 2} {\rm S} + 0.29i_ - {\rm C}_{\rm 4} {\rm H}_{{\rm 10}} + 0.11t_ - {\rm C}_{\rm 4} {\rm H}_{\rm 9} {\rm SH} $$\end{document} between 360 and 413°C. The stoichiometry of the uninhibited reaction at 380°C and 50% decomposition is approximately.

Notes: The bimolecular reaction is shown to proceed via a simple, nonchain, radical mechanism:with the net reaction the same as (1). Rate constants are estimated for each step and for each possible competing reaction and shown to yield minor or negligible side reactions in agreement with the observations of Lalonde and Back. Estimated and observed rate constants (1) and (1′) are in excellent agreement with the assumption that k'-1 is a typical radical disproportionation with zero activation energy.From the reported data a best value for k′1 is \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k'_1 [l./mol\,\sec ] = 10.3 - 44.3/\theta $$\end{document} where θ = 2.303RT kcal/mol.

Notes: Data on the liquid-phase oxidation of isobutane at 50 and 100°C have been reexamined, using a modified mechanism to take into account the termination by isobutylperoxy radicals. Algebraic expressions are derived from steady-state methods. Using Arrhenius parameters fitted by transition-state A factors and activation energies derived from observed “best” rate constants, new sets of parameters are derived for the rate constants for propagation by t—BuO2 + t—BuH → t-BuO2H + t—Bu⋅: \documentclass{article}\pagestyle{empty}\begin{document}$$ k_4 \, = \,1 \times 10^{8 - 14.5/{\rm \theta }} {\rm M}^{{\rm - 1}} \sec ^{ - 1} $$\end{document} where θ = 2.303RT in kcal/mol. This, together with new values for the termination parameters and rates of i-butyl production by k4B, is shown to give good agreement with the published data. An important reaction:\documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm R}'{\rm O}_{2}^{.} + {\rm RO}_{2} {\rm H}\mathop{{\buildrel{-\!\!\longrightarrow}\over{\longleftarrow}}}\limits^{{\rm 12}}{\rm R'O}_{\rm 2} {\rm H} + {\rm RO}_{2}^{.} $$\end{document} is shown to quench the possible contributions to termination of adventitious radicals such as CH3O⋅2.

Notes: The oxidation of butane 2,3-, propane 1,2-, ethane diol and 2-methoxy ethanol in aqueous alkaline medium by Os(VIII) has been studied. The reaction is base catalyzed and shows first-order kinetics in Os(VIII), whereas the order is less than 1 in butane 2,3-diol [BD]. The rate of oxidation is BD 〉 propane 1,2 〉 ethane diol ≈ 2-methoxy ethanol. The change in ionic strength has no effect on the rate of reaction. Activation parameters ΔE, PZ, and ΔS* have been evaluated.

Notes: The kinetic isotope effect (KIE) for carbon and oxygen in the reaction CO + OH has been measured over a range of pressures of air and at 0.2 and 1.0 atm of oxygen, argon, and helium. The reaction was carried out with 21-86% conversion under static conditions, utilizing the photolysis of H2O2 as a source of OH radicals. The value of the KIE for carbon varies with pressure and the kind of ambient gas; for air the ratio of the reaction rates 12k/13k has the value 1.007 at 1.00 atm and decreases to 0.997 at 0.2 atm; for oxygen and argon over the same pressure range the values are 1.002-0.994 and 1.000-0.991, respectively. The value of the KIE for the CO oxygen atom is 16k/18k = 0.990 over the pressure range 0.2-1.0 atm and is independent of the kind of ambient gas. No exchange of the oxygen atoms in the activated complex, followed by decomposition to the starting molecules, was observed. From the mechanistic standpoint the normal KIE observed for carbon at the high pressure is attributed to the initial formation of the activated HOCO radical, whereas the inverse KIE observed at low pressures is a result of the KIE for the reverse reaction HOCO† → CO + OH being greater than that for the forward reaction HOCO† → CO2 + H. The derived isotopic equilibrium constant for HOCO ⇄CO favors the enrichment of 13C in the more strongly bound HOCO.

Notes: At 495°C and a low extent of reaction, ethanal pyrolysis is slightly inhibited by the addition of small quantities of butadiene-1,3, whereas it is accelerated by more important quantities. The inhibiting effect is interpreted in terms of a free-radical chain mechanism in which the main chain carriers of ethanal pyrolysis (CH3.free radicals) reversibly add to butadiene-1,3 and yield penten-2-yl (R.) free radicals. These free radicals either react in a metathetical step: or in terminating steps. But butadiene-1,3 also gives rise to new initiation steps: which account for the accelerating effect. Process (i3) seems to be more important than process (i2) in the experimental conditions, but its nature could not be identified. The results are consistent with literature data and the following value of k6: \documentclass{article}\pagestyle{empty}\begin{document}$$ k_6 = 10^{12 - 12,000/4,57T} cm^3 /mol\sec $$\end{document}(4.57T in cal/mol).

Notes: The decadic extinction coefficient of the methyl radical at 216.4 nm and the rate constant for mutual combination were redetermined as: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} \varepsilon (216.4) = (9.5 \pm 0.4) \times 10^3 1./mol\,cm \\ k_2 = (3.2 \pm 0.4) \times 10^{10} 1./mol\sec \\ \end{array} $$\end{document}. The application of the Beer-Lambert law to these measurements was justified experimentally. The absorption spectrum of the methylperoxy radical was characterized as a weak, broad, structureless band, having a maximum at 240 nm with ∊(240) = 1.55 × 103 l./mol cm. The mutual interaction of methylperoxy radicals leads to the generation of methoxy and hydroperoxy radicals as a consequence of the nonterminating interaction\documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} 2{\rm CH}_{\rm 3} {\rm OO}^{\rm .} \to 2{\rm CH3O}^{\rm .} + {\rm O}_{\rm 2} \\ {\rm CH}_{\rm 3} {\rm O}^{\rm .} + {\rm O}_2 \to {\rm HCHO + HOO}^{\rm .} \\ \end{array} $$\end{document}. Each derivative radical may consume a significant fraction of the methylperoxy radicals, and either of these cross interactions may be made predominant by a suitable choice of oxygen pressure. The mutual interaction was studied under both conditions. The overall mechanism was analyzed by a precise computational method, and the rate constant of the total mutual interaction \documentclass{article}\pagestyle{empty}\begin{document}$$ 2{\rm CH3O}^{\rm .} \to {\rm all}\,{\rm products} $$\end{document} was estimated as \documentclass{article}\pagestyle{empty}\begin{document}$$ k_4 = (3.5 \pm 0.3) \times 10^8 1./mol\sec $$\end{document}.

Notes: Mixtures of up to 14% azomethane in propane have been photolyzed using mainly 366 nm radiation in the ranges of 323-453 K and 25-200 torr. Detailed measurements were made of the yields of nitrogen, methane, and ethane. Other products observed were isobutane, n-butane, ethene, and propene. A detailed mechanism is proposed and shown to account for the observed variation of product yields with experimental conditions. The quantum yield of the molecular process \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm CH}_3 {\rm N}_2 {\rm CH}_3 + h\nu \to {\rm C}_2 {\rm H}_6 + {\rm N}_2 $$\end{document} is found to be given by the temperature-independent equation\documentclass{article}\pagestyle{empty}\begin{document}$$ \phi _2 = 0.0092 \pm 0.0005] $$\end{document} The values of rate constants obtained are \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log} k_3 ({\rm cm}^3 /{\rm mol}\,{\rm sec}) = 12.03 \pm 0.15 - 40.8 \pm 1.1{\rm kJ}/{\rm mol}/(2.3RT) $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}k_4 ({\rm cm}^3 /{\rm mol}\,{\rm sec}) = 11.42 \pm 0.15 - 40.7 \pm 1.1{\rm kJ}/{\rm mol}/(2.3RT) $$\end{document} where the reactions are and it is assumed that the rate constant for the reaction is given by \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}k_5 ({\rm cm}^3 /{\rm mol}\,{\rm sec}) = 13.4 $$\end{document}

Notes: In recent publications from this laboratory, we have shown that the fragmentation of photoexcited olefinic molecules in the vacuum UV region leads mainly to the cleavage of a C - C bond located in the ß position relative to the double bond. The allyl fragment bears away part of the excess energy of the photon. At low pressure, this excited radical is capable of undergoing further decomposition. From the pressure effect, we were able to measure the first order rate constant for this secondary fragmentation. In this paper we shall use RRKM calculations in order to get a better idea on how the energy is distributed among the primary fragments. In cases where α- and β;-methallyl radicals were involved, the results show that an important part of the excess energy is located in the methallyl fragment in the 7.1 and 7.6 eV photolysis of 3-methyl-1-butene, 2-methyl-1-butene, and cis-2-pentene.

Notes: The kinetics of the slow oxidation of CO in the presence of H2 have been studied above the second explosion limit for the mixture 2CO + O2 + X% H2 at the temperature range of 530-570°C, pressures from 300 to 530 torr, and hydrogen contents of 1.1, 2.8, and 5.7%. The second explosion limit has been experimentally determined for the mixture of 2CO + O2 containing 1.0, 3.0, and 5.7% H2. On the basis of the oxidation scheme of CO in the presence of H2, which includes the accepted mechanism of oxidation of hydrogen supplemented by the reactions in which CO takes part, the second explosion limit and the profiles of the slow reaction are calculated by computer methods. The agreement found between experimental and calculated values allows one to conclude that the scheme under consideration rather completely described the slow reaction above the second limit and the occurrence of the second explosion limit in the mixture CO-O2-H2. The rate constant for the reaction HO2 + CO → OH + CO2 was calculated from the experimental data and was found to agree with previous determinations.

Notes: An example of a sequence of competing first-order reactions, leading to a common product, has been found for the sulfodeacylation of acetylpropionylmesitylene. The basic reaction scheme is Scheme l and a kinetic analysis of the component rate constants allows an estimate to be made of the concentrations of the reactant and products as a function of time. In the present system the total concentration of the intermediates B, C, and D never exceeds 0.26% at 25°C, and therefore the kinetics, which were followed spectrophotometrically, were essentially of the conversion of substrate A to the final product E.The kinetics of protiodeacylation have been measured, over a range of temperatures, of propionyl-, dipropionyl-, and acetylpropionylmesitylenes in 89.8% (w/w) sulfuric acid.

Notes: The rate of HCN evolution during the inert pyrolysis of pyridine in the temperature range of 900-1000°C was determined in a flow system using a stirred-flow reactor. The data indicated that HCN was formed through a sequence of reactions rather than during the initial step(s) involving the disappearance of pyridine. The Arrhenius parameters for the first-order step yielding HCN were 39.5 kcal/mol and 6.8 for the activation energy and log frequency factor, respectively. The mechanistic implications of the rate data are discussed, and these are related to the overall pyridine pyrolysis mechanism.

Notes: The previously reported extensive mechanism for the pyrolysis of propane and n-butane around 800 K is reexamined in the light of a recent reevaluation of the rate constant data base, and the sensitivity of model simulations to variations in the rate parameters is studied. The pyrolysis rates of butane and the product distribution of propane remain in good agreement with the available experiments, while the rate of propane and the product distribution of butane now show significant differences. The linear sensitivity analysis of the reaction matrix demonstrates an intimate coupling between initiation, hydrogen abstraction, radical decomposition, and recombination reactions as primarily responsible for the overall behavior of the mechanism. The role of unsaturated radicals in the self-inhibition of the pyrolysis process is quantitatively established. The study of the sensitivity coefficients for butane product formation has permitted pinpointing those specific reaction steps in the mechanism which are most likely responsible for the remaining discrepancies between model and experiment. This particular example demonstrates the usefulness of sensitivity calculations for the isolation of reactions for which improvements in rate parameter values are needed.

Notes: Di-tert-butylnitroxide (DTBN) is the simplest of the stable nitroxide radicals and is only consumed at temperatures higher than 90°C or in the presence of very reactive substrates. The pyrolysis of DTBN in solution gives, at least at low conversion, 2-methyl-2-nitrosopropane and di-tert-butylnitroxide-tert-butyl ether. The reaction involves, as the rate-limiting step, the cleavage of the C—N bond. This reaction takes place with an activation energy of 33 kcal/mol. DTBN is stable in the presence of styrene, aldehydes, hydrogen peroxide, α-methyl-N-ethyl nitrone, phenol, and triphenylmethane. On the other hand, it reacts readily with diethylhydroxylamine, ascorbid acid, ethanethiol, and hexanethiol. For the two former compounds the reaction involves a hydrogen transfer as the rate-determining step, and the reaction proceeds, a low conversion, with simple second-order kinetics. The reaction with the thiols is complex and shows a clear inductiontime.