ALBERT

Notes: The rate constant for tert-butyl radical recombination has been measured near 700°K by the very-low-pressure pyrolysis (VLPP) technique and was found to be 108.8±0.3 M-1·sec-1 with neglibible temperature dependence. The thermochemical parameters for tert—butyl radicals were varied within reasonable limits to bring into agreement the data for the decomposition of 2,2,3,3-tetramethyl butane and the recombination of tert-butyl radicals. The revised thermochemistry also makes the gas-phase results and liquid-phase results compatible.

Notes: The reaction of O(1D) with CH4 was studied to determine the efficiency of H2 production in a direct process, and it was found to be 0.11 ± 0.02. Thus the two channels which account for all of the reaction between O(1D) and CH4 in the gas phase are

Notes: The rate of decomposition of t-butyl nitrite (TBN) has been studied in a static system over the temperature range of 120-160°C. For low concentrations of TBN (10-5- 10-4M), but with a high total pressure of CF4 (∼0.9 atm) and small extents of reaction (∼1%), the first-order homogeneous rates of acetone (M2K) formation are a direct measure of reaction (1), since k3» k2 (NO): TBN . Addition of large amounts of NO in place of CF4 almost completely suppresses M2K formation. This shows that reaction (1) is the only route for this product. The rate of reaction (1) is given by k1 = 1016.3-40.3/θ s-1. Since (E1 + RT) and ΔH°1 are identical, both may be equated with D(RO-NO) = 40.9 ± 0.8 kcal/mole and E2 = O ± 1 kcal/mole. From ΔS°1 and A1, k2 is calculated to be 1010.4M-1 ·s-1, implying that combination of t—BuO and NO occurs once every ten collisions. From an independent observation that k2/k2′ = 1.7 ± 0.25 independent of temperature, it is concluded that k2′ = 1010.2M-1 · s-1 and k1′ = 1015.9-40.2/θ s-1; . This study shows that MeNO arises solely as a result of the combination of Me and NO. Since NO is such an excellent radical trap for t-Bu\documentclass{article}\pagestyle{empty}\begin{document}${\rm Me\dot O}$\end{document}, reaction (2) may be used in a competitive study of the decomposition of t—Bu\documentclass{article}\pagestyle{empty}\begin{document}${\rm Me\dot O}$\end{document} in order to obtain the first absolute value for k3. Preliminary results show that k3 (∞) = 1015.7-17.0/θ s-1. The pressure dependence of k3 is demonstrated over the range of 10-2-1 atm (160°C). The thermochemistry for reaction (3) implies that the Hg 6(3P1) sensitised decomposition of t-BuOH occurs via reaction (m): In addition to the products accounted for by the TBN radical split, isobutene is formed as a result of the 6-centre elimination of HONO: TBN \documentclass{article}\pagestyle{empty}\begin{document}$\mathop \to \limits^7 $\end{document} isobutene + HONO. The rate of formation of isobutene is given by k7 = 1012.9-33.6/θ s-1. t-BuOH, formed at a rate comparable to that of isobutene-at least in the initial stages-is thought to arise as a result of secondary reactions between TBN and HONO. The apparent discrepancy between this and previous studies is reconciled in terms of the above parallel reactions (1) and (7), such that k + 2k7 = 1014.7-36.2/θ s-1.

Notes: The flash photolysis of HN3 was studied by coordinated time-resolved spectroscopic measurements of HN3 NH(a1Δ), NH(X3Σ), NH(c1π), NH(A3π), NH2, and N3 following flash photolysis of mixtures of HN3 with argon or helium. The primary photolysis is complex, but when the wavelength distribution of the flash is limited to values greater than about 200 nm, the major reactive product is NH(1Δ), or states which quickly decay to NH(1Δ). Disappearance of NH(1Δ) occurs predominantly by the process \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{*{20}c} {{\rm NH(}^1 \Delta {\rm)} + {\rm HN}_{\rm 3} \to {\rm NH}_{\rm 2} + {\rm N}_{\rm 3}} & {k_2 = 9.3 \times 10^{- 11} {\rm cm}^{\rm 3} /{\rm mole} \cdot {\rm sec}} \\ \end{array}$$\end{document}The process has little, if any, energy of activation, and no detectable dependence on the pressure of inert gas below 1 atm. The rate of formation of NH2 in its ground vibrational state depends on the inert gas pressure in a way that can be accounted for by vibrational relaxation from initial excited vibrational states. The total amount of NH2 is roughly comparable with the amount of HN3 decomposed by primary photolysis. The observed N3 can be attributed to the NH(1Δ) + HN3 reaction, although a smaller amount could also be formed by primary photolysis. The value of k2 is revised upward from the value given in a preliminary report on the basis of a more careful consideration of the effects of Beer's law failure in absorption measurements involving narrow spectral lines.

Notes: The life times of chemically activated alcohols have been determined using the high-pressure unimolecular rate parameters for thermal decomposition of alcohols from shocktube studies and RRKM calculations. They are compared with literature numbers (from insertion of 0(1D) into hydrocarbons). It is suggested that in some cases singlet oxygen carries excess energy into the hydrocarbon. The consequences of such an assumption are explored and discrepancies with previously published conclusions discussed.

Notes: Aquation rates forcis-CoCl(en)2(A)2+ (A = 3,5-lutidine, imidazole, N-methylimidazole, benzimidazole) have been determined by halide release titration in 1.0 M HNO3 at 50-80°C. Kinetic parameters are (in the above order of A) 107k298 (sec-1), 7.4, 5.7, 1.3, 9.7; Ea (kJ/mole), 103, 101, 130, 112; log PZ (sec-1), 11.89, 11.53, 16.04, 13.58; ΔS298

Notes: The Diels-Alder addition of acrolein to cyclohexa-1,3-diene has been studied between 486 and 571°K at pressures ranging from 55 to 240 torr. The products are endo- and exo-5-formylbicyclo[2.2.2]oct-2-ene (endo- and exo-FBO), and their formations are second order. The rate constants (in l./mole · sec) are given by \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_{{\rm endo}} = -(19,470 \pm 50)/4.576T + (5.65 \pm 0.02)$$\end{document}\documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_{{\rm exo}} = - (20,630 \pm 50)/4.576T + (5.51 \pm 0.02)$$\end{document}The retro-Diels-Alder pyrolysis of endo-FBO has also been studied. In the ranges of 565-638°K and 6-38 torr, the reaction is first order, and its rate constant (in sec-1) is given by \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_{{\rm p}} = - (46,390 \pm 110)/4.576T + (12.98 \pm 0.04)$$\end{document}The reaction mechanism is discussed. The heat of formation and the entropy of endo-FBO are estimated.

Notes: The isomerization of cycloheptatriene at high temperatures (800-1250°K) has been studied experimentally using a shock tube at high pressures and by very low-pressure pyrolysis in the intermediate-pressure region (the first direct use of the latter technique for an isomerization). Rate coefficients obtained are in accord with previous results at lower temperatures. The results are examined theoretically to account for weak-collision nonequilibrium effects. These corrections are found to be appreciable at the temperatures studied.

Notes: Rate constants of change transfer reactions kCT, involving C3—C9 alkanes and cycloalkanes, have been determined in an ion cyclotron resonance mass spectrometer. The rate constants are significantly lower than the corresponding rate constants for collision when the reaction is less than about 0.5 eV exothermic for linear alkane ions, or less than about 0.2 eV exothermic for cycloalkane ions. In this region of low reaction efficiency, the efficiency of reaction with linear or branched alkanes seems to depend primarily on reaction exothermicity. (The efficiencies of reaction of a given ion with cyclic alkanes also depend on ΔHrn, but are higher than for reactions with other compounds). Although the lowered reaction efficiencies probably result, at least in part, from unfavorable Franck-Condon factors in the energy range near the ionization onset, quantitative correlations between reaction efficiency and estimated relative Franck-Condon factors were not observed. When the enthalpy of reaction is small (less than about -0.15 eV), it is seen that the reverse charge transfer can also occur, and equilibrium is established under the conditions of these experiments. From the observed equilibrium constants, values for the standard free energy change are derived, and assuming that ΔS is small for electron transfer equilibria, values of ΔHrn are estimated. In the case of the equilibria involving cyclohexane ion, these values of ΔHrn lead to estimates of the ionization potentials of methylcyclopentane, 3-methylpentane, n-octane, 2,2-dimethylbutane, and 2,3-dimethylbutane, which are lower than the ionization potentials of cyclohexane, that is, 〈9.88 eV, although all these compounds had previously been reported to have ionization potentials above 10.03 eV. That the ionization potentials are indeed lower than 10.03 eV is confirmed by determining the quantum yields of ionization with 10.03-eV photons.It is pointed out that the conclusions reached here apparently also apply to the charge transfer reactions of alkane ions in the liquid phase.

Notes: Concentration-time profiles have been measured for hydroxyl radicals generated by the shock-tube decomposition of hydrogen peroxide in the presence of a variety of additives. At temperatures close to 1300°K the rate constants for the reaction \documentclass{article}\pagestyle{empty}\begin{document}$${\rm OH} + {\rm RH} \to {\rm R} + {\rm H}_{\rm 2} {\rm O}$$\end{document} are found to be in the ratio 0.18:0.19:0.59:1.00:2.33:2.88 for the additives CO:CF3H:H2:CH4:C2H4:C2H6, respectively.

Notes: Crossed molecular beam techniques have been used to study the endoergic reaction between F2 and I2. Above a threshold energy of 4 kcal/mole the observed products are I2F and F. At higher energies IF is also produced. Angular and velocity distributions indicate that the IF does not result from a four-center exchange reaction.

Notes: An approximate method of analyzing nonlinear reaction models in modulated molecular beam surface kinetic studies is developed. The exact method for treating nonlinear surface mechanisms is tedious and almost always requires computer analysis. The proposed approximate method is a simple extension of the Fourier expansion technique valid for linear surface reactions; it quickly provides analytical expressions for the phase lag and amplitude of the reaction product for any type of nonlinear surface mechanism, which greatly facilitates comparison of theory and experiment. The approximate and exact methods are compared for a number of prototypical adsorption-desorption reactions which include coverage-dependent adsorption and desorption kinetics of order greater than unity. Except for certain extreme forms of coverage-dependent adsorption, the approximate method provides a good representation of the exact solution. The errors increase as the nonlinearities become stronger. Fortunately, when the discrepancy between the two methods is substantial, the reaction product signal is so highly demodulated that reliable experimental data usually cannot be obtained in these regions anyway.

Notes: A competitive technique employing the SO2(3B1) photosensitized isomerization of cis-C2F2H2 to trans-C2F2H2 in the presence of selected fluorinated olefins has been used at 3712 Å and 22°C to determine the quenching rate constants of the reaction \documentclass{article}\pagestyle{empty}\begin{document}${\rm SO}_{\rm 2} ({}^3B_1){\rm M}\mathop \to \limits^{k_{_4}}$\end{document} removal. With PSo2 = 25.4 torr and Pcis-C2F2H2 = 0.239 torr Stern-Volmer plots for M = C2H4, C2H2F, 1,1-C2F2H2, C2F4, and C3F6 yielded k4 (units of 1010 l./mole · sec) values of 5.29 ± 0.16, 4.21 ± 0.53, 1.92 ± 0.23, 0.575 ± 0.060, and 0.0335 ± 0.0027, respectively. The results were consistent with the ability of an olefin to quench SO2(3B1) being inversely proportional to the polarizability of the olefin's π bond and the effect can be clearly noted as each H atom in C2H4 is individually replaced by an F atom.

Notes: Several fluorine-containing ethanes (monofluoro, 1,1-difluoro, 1,1,1-trifluoro, 1,1,2-trifluoro, 1,1,2,2-tetrafluoro, and pentafluoro) and ethenes (1,1-difluoro and trifluoro) form hydrogen fluoride when irradiated with gamma rays in the gas phase at 25°C. Hydrogen fluoride is apparently formed from fluoroethanes by a mechanism which involves formation of an intermediate semiion pair. We observed identical HF yields both in the absence and presence of molecular oxygen, except for monofluoroethane. A reduction of G(HF) with increasing sample pressure, for example, of 1,1,2,2-tetrafluoroethane, indicates that collisional stabilization of excited fluoroethane molecules competes with the process of HF elimination. High G values for HF and CO2 in mixtures of CF2=CFH and O2 reveal the occurrence of a chain reaction.

Notes: The application of the RRKM theory to the reverse bimolecular reaction of the amine-borontrifluoride system has been made. The results are in good agreement with experiments.

Notes: The kinetics of gas-phase reaction of CH3CF2I with HI were studied from 496 to 549K and have been shown to be consistent with the following mechanism: A least squares treatment of the data gave \documentclass{article}\pagestyle{empty}\begin{document}$$\log k_1 (M^{- 1} \cdot \sec ^{- 1}) = (11.4 \pm 0.3) - \frac{{(15.7 \pm 0.8)}}{\theta}$$\end{document} where θ = 2.303 RT kcal/mole. The observed activation energy E1 was combined with E2 = 0 ± 1 kcal/mole to yield \documentclass{article}\pagestyle{empty}\begin{document}$$DH^ \circ ({\rm CH}_{\rm 3} {\rm CF}_2 - {\rm I}) = 52.1 \pm 1.0{\rm kcal}/{\rm mole}$$\end{document} The result, combined with data for several C—I bond dissociation energies, leads us to conclude that the C(sp3)—I bond is relatively insensitive to F for H substitution and that the C(sp2)-I bond has considerable double-bond character.

Notes: Tandem chemical laser techniques have been used to detect HF formed through photoelimination from 1,1,1,5,5,5-hexafluoroacetylacetone, HFAA. Information about the HF vibrational population ratios is deduced: N2/N1 〈 0.37 and N1/N0 〈 0.45. This work supports the deduction by Bassett and Whittle that photodecomposition of the enol form of HFAA results mainly in HF elimination and ring closure to a dihydrofuranone, a new type of reaction.

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 role of ethenoxy radicals in the pyrolysis of CH3CDO was studied by mass spectrometric analysis of the isotopic composition of the methane, ethane, and recovered aldehyde. Experimental evidence was obtained for the formation of ethenoxy radicals and for their reaction with acetaldehyde.Mixtures of CH3CHO and CH3CDO were pyrolyzed in order to minimize H-D scrambling in the methyl group of the aldehyde. A kinetic treatment of the methyl radical reactions and furnished the rate constant ratios (k2a + k2b)/k1a = 2.7 and k1b/k1a = 0.62 at 785°K. It is concluded that at the usual temperatures of CH3CHO pyrolysis the rate of alkyl hydrogen capture is comparable to that of formyl hydrogen abstraction.The results and conclusions are discussed and compared with previous work.

Notes: The gas-phase photochlorination (λ = 436 nm) of the 1,1,1,2-C2H2Cl4 has been studied in the absence and the presence of oxygen at temperatures between 360 and 420°K. Activation energies have been estimated for the following reaction steps: \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{*{20}c} {{\rm CCl}_3 {\rm CHCl}^{\rm .} + {\rm Cl}_2 \to {\rm CCl}_3 {\rm CHCl}_{\rm 2} + {\rm Cl}^{\rm .}} & {E_3 = (4.6 \pm 0.4){\rm kcal/mole}} \\ \end{array}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{*{20}c} {{\rm CCl}_3 {\rm CHCl}^{\rm .} \to {\rm CCl}_2 {\rm CHCl} + {\rm Cl}^{\rm .}} & {E_4 = (20.6 \pm 1.4){\rm kcal/mole}} \\ \end{array}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{*{20}c} {{\rm CCl}_3 {\rm CHClO}_{\rm 2} {\rm CHClCCl}_3 \to 2{\rm CCl}_3 {\rm CHClO}^{\rm .}} & {E_{15} = (33.5 \pm 3.0){\rm kcal/mole}} \\ \end{array}$$\end{document}The dissociation energy D(CCl3CHCl—O2) ± (24.8 ± 1.5) kcal/mole has also been estimated from the difference in activation energy of the direct and reverse reactions The mechanism is discussed and the rate parameters are compared to those obtained for a series of other chlorinated ethanes.

Notes: Hydroxyl radicals were prepared from the photolysis of N2O at 213.9 nm in the presence of excess H2. The O(1D) produced in the primary photolytic act reacts with H2 to produce OH radicals. If CO is also present, then OH can react either with H2 or CO: The competition between reactions (1) and (2) was measured by measuring the CO2 yield at various values of the ratio [CO]/[H2] at 217-298°K. At 298°K the ratio of the rate coefficients k1/k2 increased with pressure from a low-pressure limiting value of 14 to a high-pressure limiting value of 50. The low-pressure limiting value agrees well with the low-pressure values found by others. At lower temperatures our high-pressure values of k1/k2 were larger than deduced from the accepted low-pressure Arrhenius expression and could be fitted to the expression \documentclass{article}\pagestyle{empty}\begin{document}$$k_1 ^\infty /k_2 = 0.20{\rm exp (} + 3400/RT{\rm)}$$\end{document}The mechanism which seems to fit the results best is with k1° = kakb/k-a and k1∞ = ka.

Notes: The rate of CH4 disappearance in a shock-heated CH4:D2:Ar = 9.70:9.86:80.44 mixture was monitored by coincidence absorption of the 2948 cm-1 He-Ne laser line over the shockfront temperature range of 1900-2300K. Comparison with CH4 pyrolysis results by means of computer simulations suggested that atom and free radical chains are responsible for the homogeneous D/H exchange reaction on CH4.Additional simulations for the experimental conditions of previous single-pulse shock tube experiments led to the recognition of a high sensitivity of the exchange rate to trace amounts of hydrocarbon impurity and to the dissociation rate of CH4.

Notes: Direct determinations of the rate constants (cm3/molec · sec) k1, k2, and k3 from 298 to 299°K are reported, using atomic resonance fluorescence in discharge flow systems: One standard deviation.The rate constant k1, which has not been determined previously, was found to possess an insignificant temperature coefficient (EA = (0 ± 700) J/mole) in the range of 299 to 619°K.The present result for k2 agrees well with reinterpreted values from the one previous determination. Measurements of O atom consumption rates and Br atom production rates in the O + Br2 reaction are interpreted to give an estimate of the rate constant k4, which has not been reported previously, at 298°K: k3 has been measured at three temperatures between 299 and 602°K. The present and previous results for k3 were combined to give the following rate expression: \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_3 = (- 11.38 \pm 0.19) - (594 \pm 58)/T$$\end{document}

Notes: In order to describe the kinematical behavior of the bimolecular chmical reaction in a dilute solution of species characterized by a single diffusion coefficient and by some nonuniform spatial distribution of the active site, a diffusion equation with a simple sink term is derived by reducing the many-body problem to a one-body problem. The equation is normalized with the concentration of unreacted particles. The so-called second-order reaction rate constant can be calculated from the solution of the equation. The equation is applied to the intermolecular termination reaction of polymer radicals on the assumption of free draining. The reaction rate constant gradually decreases with time.

Notes: The chemical activation data for three- and four-centered hydrogen fluoride elimination from CH2FCDF2 have been analyzed to assign the energy released to the olefin fragment in the three-centered process and to estimate the threshold energies for elimination channels. Based upon the cis-trans isomerization rates of CHF = CHF, 78% of the total available energy was released to the olefin fragment for the αα channel. The analysis suggests the existence of an appreciable barrier (∼10 kcal/mole) for the reverse reaction, addition of the CH2FCF carbene to DF. The threshold energies for αα, αβ, and βα elimination from 1,1,2-trifluoroethane-1-d1 were assigned as 71, 68, and 68 kcal/mole, respectively. Analysis of the chemical activation data for 1,1,2,2,-tetrafluoroethane, without distinguishing between the three- and four-centered elimination channels, suggests a threshold energy of ∼75 kcal/mole.

Notes: The first-order rate constants for aquation of Co(NH3)5(DMF)3+ were determined in aqueous perchlorate media. The rate constants were independent of hydrogen ion concentration and ionic strength, but decreased with increasing perchlorate ion concentration. Proton magnetic resonance studies showed that dimethylformamide was not hydrolyzed to formic acid and dimethylamine in the aquation step. Mass-spectrometer studies showed that cobalt-oxygen bond breaking occurred in 98 percent, or more, of the aquation acts. Enthalpies and entropies of activation were determined. It was concluded that aquation occurred by an Id mechanism.

Notes: Our earlier work on the formation of particulate NH4NO3 in the NH3—O3 reaction at 25°C is extended to include air as a diluent and H2O vapor as an additive. More extensive data at different values of [NH3]/[O3]0 were obtained also, where [O3]0 is the initial O3 concentration. In our earlier work we concluded that NH4NO3 vapor was dissociated to NH3 + HNO3 and that the HNO3 was removed by diffusion to the walls with a rate coefficient kdiff = 0.4 min-1 or by condensation on the suspended particles. Particles were nucleated by 8 NH3—HNO3 pairs when their concentration product reached 5.8 × 1027 molec2/cm6 with a rate coefficient knucl of 6.2 × 10-224 cm45/min and removed by coagulation with a rate coefficient kcoag of 1.3 × 10-7 cm3/min. A corrected calculation modifies the number of pairs required to 6-7 with a correspondingly changed value of knucl.With the more extensive data of the present study the indications are that the vapor-phase NH4NO3 monomer is not dissociated and that its diffusion constant for loss to the walls varies between 0.3 and 0.9 min-1 for different reaction conditions. Nucleation occurs when the NH4NO3 vapor concentration reaches 1.0 × 1012 molec/cm3 via. \documentclass{article}\pagestyle{empty}\begin{document}$$r{\rm NH}_{\rm 4} {\rm NO}_{\rm 3} (g) \to ({\rm NH}_{\rm 4} {\rm NO}_{\rm 3})_r (s)$$\end{document}where r is 9 and the nucleation rate coefficient knucl is 3 × 10-108 cm24/min. With 5.0 or 9.5 torr of H2O vapor present, there is an excess of particles produced over that expected from this rate coefficient, indicating an additional nucleation step in which H2O vapor participates directly to produce a hydrated salt. The coagulation coefficient of (1.87 ± 0.14) × 10-7 cm3/min found here is in good agreement with that found previously.

Notes: The reactions of O3 with CH3ONO and C2H5ONO were studied using infrared absorption spectroscopy in a static reactor at temperatures between 298 and 352K. Both reactions followed simple second-order kinetics forming the corresponding nitrate: The rate coefficients are given by \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_1 ({\rm cm}^{\rm 3} /{\rm molc} \cdot {\rm sec}) = (- 12.17 \pm 0.23) - (\frac{{5315 \pm 172}}{{2.303{\rm}T}})$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} {\rm}k_2 ({\rm cm}^{\rm 3} /{\rm molc} \cdot {\rm sec}) = (- 15.50 \pm 0.16) - (\frac{{2351 \pm 116}}{{2.303{\rm}T}})$$\end{document}.

Notes: Quantum yield measurements for the SO2(3B1) photosensitized isomerization of cis-1,2-difluoroethylene have been made at 3712 Å and 22°C. The [SO2]/[cis-C2F2H2] ratio was varied from 47.4 to 455 and the quantum yield measurements over this variation of concentration ratios were consistent with a mechanism in which SO2(3B1) molecules and the cis isomer form a collision intermediate which decomposes with a probability of 0.42 ± 0.17 and 0.58 ± 0.17 of producing trans- and cis-1,2-difluoroethylene, respectively. When SO2 was subjected to prolonged irradiations in the presence of initially either pure cis- or pure trans-1,2-difluoroethylene, a photostationary composition, [cis]/[trans] = 1.0 ± 0.2, was obtained. The rate constant at 22°C for removal of SO2(3B1) molecules by cis-1,2-difluoroethylene was estimated to be (1.72 ± 0.72) × 1010 1./mole · sec.

Notes: The photolysis of SO2 at 3080 Å, FWHM = 150 Å, and 22°C has been investigated in the presence of cis- and trans-C2F2H2. Quantum yield measurements for the photosensitized isomerization of cis-C2F2H2 to trans-C2F2H2 have been made for a variation in the [SO2]/[cis-C2F2H2] ratio from 0.992 to 253. The results fit a mechanism which is consistent with the SO2(3B1) state being the reactive excited state of sulfur dioxide. A mechanism employing only the SO2(1B1) and SO2(3B1) excited states is quite satisfactory to rationalize the data. A value for the SO2 collisionally induced intersystem crossing efficiency from SO2(1B1) to SO2(3B1) of 0.35 ± 0.14 was estimated while the cis-C2F2H2 efficiency was found to be 0.030 ± 0.012. The rate constant at 22°C for the removal of SO2(3B1) molecules by cis-C2F2H2 was found to be (1.43 ± 0.13) × 10101./mole · sec. A photostationary composition, [cis]/[trans] = 1.0 ± 0.1, was found from prolonged irradiations of SO2 in the presence of the cis and trans isomers.

Notes: A combined flash photolysis and pulse radiolysis experiment was carried out to produce triplet pyrene (P) molecules in micelles of cetyltrimethylammonium bromide and Br2- in the surrounding aqueous medium. The reaction 3Pmic + Br2aq- → Pmic+ + 2 Br- was followed by optical absorption measurements in the 10-8-10-4-sec range. This reaction possesses a “fast” and a “slow” component with respective rate constants of 2.3 × 106 sec-1 and 1 × 109M-1 · sec-1.The fast component is related to the probability of a Br2- radical meeting a triplet pyrene containing micelle on the first encounter (only 16% of the micelles contained a triplet molecule). Reactions involving more than one Br2- radical-micelle encounter are ascribed to the slow component. The presence of two components reflects the fact that the residence time of a Br2- radical in the vicinity of a cationic micelle is substantially longer than the diffusion time of the radical between micelles. Thus the conditions met in micellar chemistry differ dramatically from those in ordinary solution kinetics where the encounter time is generally much shorter than the time between encounters. Some considerations on the energetics of this electron transfer reaction are also presented.

Notes: The decomposition of ethane sensitized by isopropyl radicals was studied in the temperature range of 496-548°K. The rate of formation of n-butane, isopentane, and 2,3-dimethylbutane were measured. The expression k1/k2½ was found to be \documentclass{article}\pagestyle{empty}\begin{document}$$\log {\rm}k_1 /k_2 ^{1/2} ({\rm dm}^{{\rm 3/2}} /{\rm mol}^{{\rm 1/2}} \cdot \sec ^{1/2}) = (3.2 \pm 0.4) - (54 \pm 3{\rm kJ/mol})/2.3RT$$\end{document} where k1 and k2 are rate constants of The decomposition of propylene sensitized by isopropyl radicals was studied between 494 and 580°K by determination of the initial rates of formation of the main products. The ratio of k13/k21/2 was evaluated to be \documentclass{article}\pagestyle{empty}\begin{document}$$\log {\rm}k_{13} /k_2 ^{1/2} ({\rm dm}^{{\rm 3/2}} /{\rm mol}^{{\rm 1/2}} \cdot \sec ^{1/2}) = (1.9 \pm 0.7) - (32 \pm 7{\rm kJ/mol})/2.3RT$$\end{document} where k13 is the rate constant for The isomerization of the isopropyl radical was investigated by studying the decomposition of azoisopropane. The decomposition of the iso-C3H7 radical into C2H4 and CH3 was followed by measuring the rate of formation of C2H4. On the basis of the experimental data, obtained in the range of 538-666° K, k15/k2½ was found: \documentclass{article}\pagestyle{empty}\begin{document}$$\log {\rm}k_{15} /k_2 ^{1/2} ({\rm mol}^{{\rm 1/2}} /{\rm dm}^{{\rm 3/2}} \cdot \sec ^{1/2}) = (9.3 \pm 0.8) - (152 \pm 9{\rm kJ/mol})/2.3RT$$\end{document} where k15 is the rate constant of

Notes: The photolysis of SO2 in the presence of cis- and trans-2-pentene has been investigated at 3660 Å and 22°C. Quantum yield measurements of the SO2 photosensitized conversion of one isomer into the other are consistent with a mechanism in which the only participating excited electronic state of SO2 is the SO2(3B1) state. Quantum yield measurements were made for a variation in PSO2/Pisomer reactant ratios of 4.01-283 and 57.5-351 for the cis and trans isomers, respectively. The data are consistent with a mechanism in which a (SO2-olefin)3 collision intermediate is the precursor to the photosensitized isomeric products. The intermediate undergoes unimolecular decay to yield the cis and trans isomers with probabilities of 0.26 ± 0.05 and 0.69 ± 0.04, respectively. Estimates of the quenching rate constants at 22°C for removal of SO2(3B1) molecules by cis- and trans-2-pentene are (0.633 ± 0.125) × 1011 l./mole/sec and (1.00 ± 0.27) × 1011 l./mole/sec, respectively. An experimentally determined photostationary composition, [trans-2-pentene]/[cis-2-pentene] = 2.3 ± 0.1 was in fair agreement with that of 1.7 ± 0.7 as predicted from kinetic data derived in this study.

Notes: The photochemistry of azoethane and hexafluoroazomethane at 366 nm has been reinvestigated up to 1 atm pressure, and over a range of temperature from 27 to 150°C. The Stern-Volmer type quenching plots primarily demonstrate the decomposition of a single electronic and vibrationally excited state for azoethane, but competitive photodissociation from two different electronic and vibrationally excited states, which was previously postulated for hexafluoroazomethane and azoisopropane, is confirmed for hexafluoroazomethane. It is concluded, however, that two different electronic and vibrationally excited photodissociating states are present in azoethane photolysis, but that one of them is difficult to detect, at least by the present approachPhotosensitization with biacetyl at 436 nm also causes the dissociation of azoethane, and this is probably from the vibrationally equilibrated first triplet state. The energy barrier for this process was found to be 5.0 kcal/mol.

Notes: Measurements of hydroxyl radical (HO) concentrations in ambient air by the technique of laser-induced fluorescence have been recently reported. The present study was undertaken to provide an independent test of the validity of those measurements. A photochemical reactor was used to provide a source of HO, and the concentration of HO in the reactor was determined by the laser-induced fluorescence technique. The HO concentration was also deduced from measured hydrocarbon decay rates in the reactor. There was agreement between the HO concentrations obtained by these two different methods, thus providing further validation of the fluorescence method. Some studies of HO fluorescence efficiency as well as of possible interferences with the fluorescence measurements are reported.

Notes: New experimental data have been obtained for H + C2H2, D + C2H2, H + C2D2, and D + C2D2 at room temperature. Two previously described apparatus were used in order to measure the pressure dependence of the reactions. The absolute rate constants are compared to results from other laboratories. The present results and those of Payne and Stief are used to obtain the high-pressure limiting rate constant at room temperature. When the activation energy from the work of Payne and Stief is considered, it is shown that the A factor for H + C2H2 is too low by a factor of ∼20. If a transmission coefficient is introduced which is constant for all isotopic variations, the pressure dependence can be explained in terms of the randomly energized radicals. RRKM theory is then invoked to explain the observed statistical nonequilibrium kinetic isotope effects.

Notes: The flash photolysis of azomethane in a quartz reaction vessel produces mainly ethane (〉75%) plus smaller quantities of methane, ethylene, and acetylene. The minor products are interpreted quantitatively in terms of methyl radical photolysis at 216 nm to give CH2 and H. This interpretation is substantiated by the dependence of the minor products on flash intensity. The reduction of the ethane yield on adding NO is employed to obtain a rate constant for CH3 + NO as a function of total pressure, based on a value for methyl radical recombination of 4.2 × 10-11 cm3/molec · sec. An RRKM analysis is used to extrapolate the data to give a limiting high-pressure rate constant for CH3 + NO of (1.2 ± 0.1) × 10-11 cm3/molec · sec at 298°K.

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: An analytical and kinetic study of the thermal reaction of cis- or trans-2-butene has been performed in a static system over the temperature range of 480-550°C and at a low extent of reaction and initial pressures of 10-100 torr.The rate constant of the unimolecular cis-trans isomerization of cis-2-butene, determined under the conditions\documentclass{article}\pagestyle{empty}\begin{document}$$k_{ct} = 10^{13.6 \pm 0.3 - 62,000 \pm 1000/2.3{\rm RT}} {\rm sec}^{- 1} $$\end{document}(2.3 RT in cal/mole) is in good agreement with previous measurements made at lower pressures.A comparison between the formation rates of hydrogen from the thermal reactions of cis- and trans-2 butene around 500°C leads to the rate constant value\documentclass{article}\pagestyle{empty}\begin{document}$$k_m = 10^{13 \pm 0.5 - 65,500 \pm 2000/2.3{\rm RT}} {\rm sec}^{- 1} $$\end{document}(2.3 RT in cal/mole) for the unimolecular 1,4—hydrogen elimination from cis—2—butene.

Notes: The hydrogen abstraction from asymmetrically fluorinated and chlorofluorinated ethanes by chlorine atoms has been investigated in the gas phase between 264 and 333°K using the competition method. Arrhenius parameters for the reaction on both sites of the molecules are discussed.

Notes: 3,3-Dimethylbutanol-2 (3,3-DMB-ol-2) and 2,3-dimethylbutanol-2 (2,3-DMB-ol-2) have been decomposed in comparative-rate single-pulse shock-tube experiments. The mechanisms of the decompositions are The rate expressions are \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\rm B} (2,3{\rm - DBM - ol - 2}) = 10^{16.24} {\rm exp}(- 37,400/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\rm EP} (2,3{\rm - DBM - ol - 2}) = 10^{14.17} {\rm exp}(- 32,300/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\rm ET} (2,3{\rm - DBM - ol - 2}) = 10^{13.66} {\rm exp}(- 32,700/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\rm B} (3,3{\rm - DBM - ol - 2}) = 10^{16.33} {\rm exp}(- 37,500/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\rm EP} (3,3{\rm - DBM - ol - 2}) = 10^{14.0} {\rm exp}(- 34,200/T)\sec ^{- 1}$$\end{document} They lead to D(iC3H7—H) - D((CH3)2(OH) C—H) = 8.3 kJ and D(C2H5—H) - D(CH3(OH) CH—H) = 24.2 kJ.These data, in conjunction with reasonable assumptions, give \documentclass{article}\pagestyle{empty}\begin{document}$$k(t{\rm C}_{\rm 4} {\rm H}_{\rm 9} {\rm OH} \to {\rm CH}_{\rm 3} \cdot + \cdot {\rm C(CH}_{\rm 3} {\rm)}_{\rm 2} {\rm OH}) = 10^{16.8} {\rm exp}(- 40,900/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k(i{\rm C}_{\rm 3} {\rm H}_{\rm 7} {\rm OH} \to {\rm CH}_{\rm 3} \cdot + \cdot {\rm CH(CH}_{\rm 3} {\rm)OH}) = 10^{16.5} {\rm exp}(- 41,100/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k(n{\rm C}_{\rm 3} {\rm H}_{\rm 7} {\rm OH} \to {\rm CH}_{\rm 3} \cdot + \cdot {\rm CH}_{\rm 2} {\rm CH}_{\rm 2} {\rm OH}) = 10^{16.2} {\rm exp}(- 41,100/T)\sec ^{- 1}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k({\rm C}_{\rm 2} {\rm H}_{\rm 5} {\rm OH} \to {\rm CH}_{\rm 3} \cdot + \cdot {\rm CH}_{\rm 2} {\rm OH}) = 10^{16.4} {\rm exp}(- 42,500/T)\sec ^{- 1}$$\end{document}andThe rate expressions for the decomposition of 2,3-DMB-1 and 3,3-DMB-1 are \documentclass{article}\pagestyle{empty}\begin{document}$$ k(2,3{\rm - DMB - 1} \to {\rm CH}_3 \cdot + {\rm H}_2 {\rm C} = {\rm C}({\rm CH}_3 ) - \mathop {\mathop {\rm C}\limits^{\rm .} {\rm H}({\rm CH}_3 )) = 10^{16.0} \exp ( - 35,700/T)\sec ^{ - 1} } $$\end{document} and \documentclass{article}\pagestyle{empty}\begin{document}$$ k(3,3{\rm - DMB - 1} \to {\rm CH}_{\rm 3} \cdot + {\rm H}_2 {\rm C = CH} - \mathop {\rm C}\limits^{\rm .} ({\rm CH}_3 )_2 ) = 10^{16.2} \exp ( - 35,500/T)\sec ^{ - 1} $$\end{document}

Notes: The kinetics of the oxidation of iodide by hydrogen peroxide catalyzed by acidic molybdate have been studied by a spectrophotometric stopped-flow method. The results are interpreted in terms of the mechanism and the implied rate law \documentclass{article}\pagestyle{empty}\begin{document}$$ - d[{\rm H}_{\rm 2} {\rm O}]/dt = \frac{{k_4 k_1 k_2 [{\rm H}_2 {\rm O}_2]^2 [{\rm mol}][{\rm I}^ -]}}{{1 + k_1 [{\rm H}_2 {\rm O}_2] + k_1 k_2 [{\rm H}_2 {\rm O}_2]^2 + {\rm K}_{\rm 3} [{\rm I}^{\rm -}]}}$$\end{document}where [mol] is total analytical concentration of molybdate. The values obtained for the rate and equilibrium constants are k4 = (3.3 ± 1) × 102 1./mole · s, K1 = (1.2 ± 0.6) × 104 1./mole, K2 = (1.3 ± 0.7) × 103 1./mole, and K3 = (4 ± 3) × 102 1./mole at 298°K.

Notes: The Co(NH3)5OH23+ ion reacts with malonate to form Co(NH3)5O2CCH2CO2H2+ or Co(NH3)5O2CCH2CO2+, depending on the pH of the reaction solution. The kinetics of this anation reaction have been studied as a function of [H+] for the acidity range 1.5 ≤ pH ≤ 6.0 in the temperature range of 60 to 80°C, the [total malonate] ≤ 0.5 M, and the ionic strength 1.0M. The anation by malonic acid follows second-order kinetics, the rate constant being 8.0 × 10-5 M-1·sec-1 at 70°C, and the anations by bimalonate (Q1, k1) and malonate ion (Q2, k2) are consistent with an Id mechanism. Typical values at 70°C for the ion pair formation constants are Q1 = 1.3, Q2 = 5.4M-1; and for the interchange rate constants k1 = 5.3 × 10-4; k2 = 7.3 × 10-4 sec-1. The activation parameters for the various rate constants are reported and the results discussed with reference to previously reported data for similar systems.

Notes: The room temperature reactions between oxygen atoms and methanethiol, ethanethiol, and methylsulfide have been studied in crossed jets to directly detect and identify the free-radical and stable products they produce. Knowledge of the products was used to assign reactive routes. The overall rate constants for all three O-atom reactions were also measured at 300 ± 2°K using a fast-flow reactor. They are 1.9 (CH3SH), 2.8 (C2H5SH), and 63 (CH3SCH3) × 10-12cm3/mol · sec. The identity of the detected products and the trend in rate constants in these reactions support an electrophilic addition mechanism followed by decomposition of the excited adduct by S-R bond cleavage (R = H, CH3, or C2H5).

Notes: The kinetics of the gamma-radiation induced free-radical chain reaction in solutions of carbon tetrachloride in cyclohexane (RH) has been investigated in the temperature range of 303-383°K. Trichloromethyl radicals were produced by the reaction of radiolytically generated cyclohexyl radicals with carbon tetrachloride. The kinetics of the following reactions were investigated: The following rate expression was obtained: \documentclass{article}\pagestyle{empty}\begin{document}$$\log {\rm}k_3 /(k_4)^{{1 \mathord{\left/ {\vphantom {1 2}} \right. \kern-\nulldelimiterspace} 2}} ({\rm cm}^{{1 \mathord{\left/ {\vphantom {1 2}} \right. \kern-\nulldelimiterspace} 2}} \cdot {\rm sec}^{{1 \mathord{\left/ {\vphantom {1 2}} \right. \kern-\nulldelimiterspace} 2}}) = 4.78 \pm 0.08 - (8.81 \pm 0.12)/\theta ^1$$\end{document} The error limits are the standard deviation from the least mean square Arrhenius plots. Effects of phase on the kinetics of reactions (3) and (4) are considered.

Notes: Metastable N2(A3Σu+), υ = 0, υ = 1, molecules are produced by a pulsed Tesla-type discharge of a dilute N2/Ar gas mixture. Rate coefficients for quenching these metastable levels by O2, O, N, and H were obtained by time-resolved emission measurements of the (0, 6) and (1, 5) Vegard-Kaplan bands. In units of cm3/mole · sec at 300°K and with an experimental uncertainty of ±20%, these rate coefficients for N2(A3Σu+) are \documentclass{article}\pagestyle{empty}\begin{document}$$k_{{\rm O2}} = 2.9 \times 10^{- 12}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{{\rm O}} = 1.5 \times 10^{- 11}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{{\rm N}} = 4.8 \times 10^{- 11}$$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$k_{{\rm H}} = 3.5 \times 10^{- 12}$$\end{document}Within the limits of error these coefficients apply to quenching N2(A3Σu+) υ′ = 1 as well.

Notes: Nanosecond flash photolysis of b-nitronaphthalene (b-NO2C10H7) in nonpolar and polar solvents shows a transient species with maximum absorption and lifetime dependent on solvent polarity. In deaerated n-hexane the absorption maximum and lifetime (1/k) are 425 nm and 530 nsec, while in deaerated ethanol the corresponding values are 470 nm and 1.7 ·sec. This transient absorption is attributed to the triplet excited state of b-NO2C10H7, and the observed red shift as well as its longer lifetime in polar solvents are indicative of the intramolecular charge transfer character of this state. The change of dipole moment accompanying the transition T1 → Tn, as well as rate constants for electron and proton transfer reactions involving the T1 state of b-NO2C10H7, were determined. The spectroscopic and kinetic data obtained in this work indicate that the triplet state of b-NO2C10H7 behaves like a n-π* state in nonpolar media, while in polar solvents the n-π* character of the state is reduced with a simultaneous increase in the charge transfer character.

Notes: The oxidation of α-hydroxy acids and α-hydroxy ketones by Br(V) follows the rate-law \documentclass{article}\pagestyle{empty}\begin{document}$$\frac{{- d[{\rm Br(V)}]}}{{dt}} = k_2 [{\rm Br(V)}][{\rm substrate}]$$\end{document}However, the former reaction exhibits a second-order dependence on hydrogen ion concentration while the latter reaction has a third-order dependence. A mechanism involving a slow formation of a bromate ester of the α-hydroxy acid followed by a fast decomposition is proposed. A rate-determining formation of a bromate ester from the conjugate acid of benzoin, followed by a rapid decomposition of the bromate ester, explains the kinetic data for the oxidation of benzoin.

Notes: Dichloroethylene (DCE), either cis or trans, was reacted with O3 at 23°C in both N2 and O2 buffered mixtures. Both reactant consumption and product formation were monitored by infrared spectroscopy and, in some cases, O3 consumption was monitored by ultraviolet absorption. For thoroughly dried mixtures, the initial products were only HCClO and O2, but geometrical isomerization also occurred. The stoichiometry of the overall reaction always was The HCClO was unstable and disappeared slowly in a first-order reaction which was, at least in part, heterogeneous. The products were CO and HCl so that the stoichiometric reaction was The rate law was complex. The rate was always faster in N2 than in O2. In the N2 buffered reaction, inhibition occurred as the reaction progressed and O2 was produced. From the reactant and product decay curves, the following rate behavior was established: where high and low concentrations are relative terms for the initial pressure ranges covered ([DCE]0 = 0.21-78.4 torr, [O3]0 = 0.30-6.76 torr). The rate coefficients k2, k3, and k4 were larger for the trans-DCE than the cis-DCE, and for each isomer they were larger in N2 than in O2 buffered reactions.The ozonolysis can be explained in terms of the mechanism where R2 is DCE, RO is HCClO, and RO2 is HCClO2. Rate ceofficients are computed.The isomerization is first order in [O3] and approximately first order in [DCE] for the limited kinetic data we were able to obtain. The isomerization does not appear to be explained by the reverse reactions of reactions (6), (7), and (9). Presumably isomerization occurs through some other route.