ALBERT

Notes: The decomposition of dimethyl peroxide (DMP) was studied in the presence and absence of added NO2 to determine rate constants k1 and k2 in the temperature range of 391-432°K: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{rcl} {{\rm DMP}} & \stackrel{1}{\longrightarrow} & {2{\rm MeO}} \\ {{\rm MeO + DMP}} & \stackrel{2}{\longrightarrow} & {{\rm MeOH + CH}_{\rm 2} {\rm O} + {\rm MeO}} \\ \end{array} $$\end{document} The results reconcile the studies by Takezaki and Takeuchi, Hanst and Calvert, and Batt and McCulloch, giving log k1(sec-1) = (15.7 ± 0.5) - (37.1 ± 0.9)/2.3 RT and k2 ≈ 5 × 104M-1· sec-1. The disproportionation/recombination ratio k7b/k7a = 0.30 ± 0.05 was also determined: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{rcl} {{\rm MeO} + {\rm NO}_2 (+ {\rm M})} \stackrel{7a}{\longrightarrow} & {{\rm MeONO}_{\rm 2} (+ {\rm M})} \\ {{\rm MeO} + {\rm NO}_2} \stackrel{7b}{\longrightarrow} & {{\rm CH}_{\rm 2} {\rm O} + {\rm HONO}} \\ \end{array} $$\end{document}When O2 was added to DMP mixtures containing NO2, relative rate constants k12/k7a were obtained over the temperature range of 396-442°K: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{rcl} {{\rm CH}_{\rm 3} {\rm O} + {\rm O}_2} \stackrel{12}{\longrightarrow} & {{\rm CH}_2 {\rm O} + {\rm HO}_{\rm 2}} \\ \end{array} $$\end{document} A review of literature data produced k7a = 109.8±0.5M-1·sec-1, giving log k12(M-1·sec-1) = (8.5 ± 1.5) - (4.0 ± 2.8)/2.3 RT, where most of the uncertainty is due to the limited temperature range of the experiments.

Notes: By means of the technique of laser-induced fluorescence, the room-temperature vibrational relaxation of DF(v = 1) has been studied in the presence of several polyatomic chaperones. The rate coefficients obtained [in units of (μ;sec·torr)-1] are CH4, 0.22; C2H6, 0.61; C4H10, 1.26; C2H2, 4.0 × 10-2; C2H2F2, 1.86 × 10-2; C2H4, 0.175; CH3F, 0.36; CF3H, 1.95 × 10-2; CF4, 1.0 × 10-3; CBrF3, 5.6 × 10-4; NF3, 5.1 × 10-4; SO2, 1.27 × 10-2; and BF3, 7.1 × 10-3. Results are also reported for vibrational relaxation rate coefficients for HF(v = 1) in the presence of the following chaperones: CH4, 2.6 × 10-2; C2H6, 5.9 × 10-2; C3H8, 8.4 × 10-2; and C4H10, 0.128. A comparison of DF and HF results indicates that for deactivation by CnHn+2, rate coefficients for DF are approximately an order of magnitude larger than for HF. The deactivation rate coefficient of DF(v = 1) by CH4 was found to decrease with increasing temperature between 300 and 740°K.

Notes: The reactions between alizarin yellow G and six different bases B (including OH-) and between tropaeolin 0 and eight different bases have been investigated at 25°C and an ionic strength of 0.5M, using the temperature-jump method. From the form of the log kB versus ΔpK curves it is concluded that for alizarin yellow G the observed relaxation time is due chiefly to a diffusion-controlled reaction between the base and that fraction which is present in the “open” non-hydrogen-bonded form, whereas for tropaeolin 0 the base attacks the hydrogen bridge.The dissociation constants for the internally bound hydrogen have been measured under the same conditions of temperature and ionic strength, using a spectrophotometric method.

Notes: The reaction of 1,3-cyclopentadiene (CPD) with ground-state atomic oxygen O(3P), produced by mercury photosensitized decomposition of nitrous oxide, was studied. The identified products were carbon monoxide and the following C4H6 isomers: 3-methylcyclopropene, 1,3-butadiene, 1,2-butadiene, and 1-butyne. The yield of carbon monoxide over oxygen atoms produced (φCO) was equal to the sum of the yields of C4H6 isomers in any experiment. φCO was 0.43 at the total pressure of 6.5 torr and 0.20 at 500 torr. We did not succeed in detecting any addition products such as C5H6O isomers.It was found that 3-methylcyclopropene was produced with excess energy and was partly isomerized to other C4H6 isomers, especially to 1-butyne. The excess energy was estimated to be about 50 kcal/mol.The rate coefficient of the reaction was obtained relative to those for the reactions of atomic oxygen with trans-2-butene and 1-butene. The ratios kCPD+O/ktrans-2-butene+O= 2.34 and kCPD+O/k1-butene+O = 11.3 were obtained.Probable reaction mechanisms and intermediates are suggested.

Notes: The kinetics of chlorine atom abstraction from the chloroethanes (EClH) 1,1,2-C2Cl3H3, 1,1-C2Cl2H4, and 1,2-C2Cl2H4 by radiolytically generated cyclohexyl radicals was studied in the liquid phase by a competitive method. The chlorine atom abstraction data were put on an absolute basis by comparing the rates of the metathetical reactions with the known rate of addition of cyclohexyl radicals to C2Cl4. The following Arrhenius parameters were obtained: TextE(EClH)-TemperatureA(ECLH)E(CCl4)log A(EClH)E(EClH)RangelogEClHA(C2Cl4)(kcal/mol)(1.mol·sec)(kcal/mol)(°C)CHCL2CH2Cl0.03 ± 0.083.87 ± 0.178.98 ± 0.1411.17 ± 0.27150 - 250CHCL2CH30.13 ± 0.134.63 ± 0.278.18 ± 0.1911.93 ± 0.37130 - 250CHCL2CH2Cl0.50 ± 0.177.57 ± 0.359.18 ± 0.2314.87 ± 0.45150 - 250The error limits are the standard deviations from least mean square Arrhenius plots.The α and ß activation effects on the kinetics of Cl atom abstraction from chloroalkanes are discussed. From the linear relation between the relative reactivities of cyclohexyl radicals toward the XCCl3 and XCHCl2 series, ECl(c-C6H11· + CHCl2CHCl2) = 10.2 ± 1 kcal/mol and ECl(c-C6H11· + CHCl22CCl3) = 9.7 ± 1 kcal is derived.

Notes: Reactions of OH(v = 1) with HBr, O, and CO have been studied at 295°K using a fast discharge flow apparatus: The reaction O + HBr → OH(v = 1) + Br was used as a source of OH(v = 1), and subsequent chemical reactions of the excited radical were followed using EPR spectroscopy. Rate constants for reactions (2b), (3b), and (6b) were measured as (4.5 ± 1.3) × 10-11, (10.5 ± 5.3) × 10-11, and 〈5 × 10-12 cm3/molec·sec, respectively. The rate constant for physical deactivation of OH(v = 1) by CO was determined as 〈4 × 10-13 cm3/molec·sec.

Notes: The reactions have been studied by a mass-balance method involving the photolysis of small amounts of biacetyl in the presence of a large excess of isobutane containing a small proportion of the unsaturated substrate. The following Arrhenius parameters have been derived: TextTemperatureElog ArangeReaction(kcal/mol)(1./mol·sec)(°K)ĊH3 + C2H4 → Ċ3H77.3 ± 1.08.32 ± 0.5350 - 500ĊH3 + C2H2 → Ċ3H57.7 ± 1.58.79 ± 0.8379 - 487ĊH3 + C6H6 → C7H97.6 ± 1.08.79 ± 0.5372 - 484The results for methyl addition to ethylene are based on previous determinations by other techniques as well as the present studies. The results for methyl addition to acetylene and benzene are derived solely from the present experiments and are calculated relative to a rate constant of log k2(l./mol·sec) = 7.42 - (7.1/θ) for the reference reaction (2), ·H3 + (CH3)3CH → CH4 + ·4H9.

Notes: The thermal decomposition of propene in the presence of D2 was studied in a single-pulse shock tube in the temperature range of 1200-1400°K. The main decomposition products were methane, ethylene, allene, and propyne. Furthermore, deuterated species were observed of each product and of propene, with characteristic compositions that were dependent on propene conversion. Geometrical isomers of monodeuterated propene, as the result of H-D exchange, were analyzed by microwave spectroscopy. From these observations, the reactivities of n- and isopropyl radicals at high temperatures were determined. The former was found to be an intermediate of methane and ethylene and the latter was found to be responsible for the formation of the deuterated propene as follows: The rate constant ratio kn/ki was estimated to be 0.5-0.8, which was more than ten times greater than that obtained at room temperature. It was also found that allene or propyne was produced from allyl radicals and that acetylene was produced from vinyl radicals. In addition, the rate constant of the hydrogen abstraction by the hydrogen atom from C3H6 was found to be six times greater than that by the hydrogen atom from D2.

Notes: The oxidation of cerous ions by bromate ions in sulfuric acid medium was followed spectrophotometrically under various experimental conditions. The results were compared to the calculated predictions on the basis of a mechanism suggested by R. M. Noyes and collaborators. The computations were done by solving the complete set of the kinetic differential equations. The results of the computations show that the proposed mechanism explains adequately most of our and previous experimental data. In particular, the mechanism predicts the main features of the reaction, namely, the induction and the fast and slow reaction periods which occur during the oxidation.

Notes: The thermal decomposition of NO2 and its atom-transfer reactions with SO2 and CO have been studied behind incident shock waves using photometric detection methods. From the decomposition study it is possible to obtain information on the rate of the reaction 2NO2 → antisymmetric-NO3 + NO. The results on the reaction, NO2 + SO2 → NO + SO3 extend the earlier work of Armitage and Cullis to about 2000°K. The reaction with CO [NO2 +] [CO NO + CO2] at shock temperatures is somewhat faster than predicted from available low-temperature data and provides a modification of the rate-constant expression that is applicable over a wide temperature range.

Notes: The kinetics of the reaction of cis-(NO)2 with solid oxygen to form iso-N2O4 have been studied between 13 and 29 K. The overall reaction is pseudo first order in cis-(NO)2, and solid oxygen serves both as a reactant and the matrix. The pseudo-first-order rate constants are calculated to be k(14N) = 4.25 × 10-2 exp(-103/RT), and k(15N) = 3.00 × 10-2 exp(-105/RT) sec-1, based on temperature measurements from a thermocouple junction which may be at most three degrees lower than the actual reacting film. Most significantly, however, 14k/15k = 1.55 at∼13 K. The condensed phase reaction has been compared to that observed in the gas phase, and the extremely small pre-exponential factors and large isotope effects have been discussed in terms of tunneling corrections and orientational constraints. It is suggested that the form of the crystal plays an integral role in the observed process.

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: Computer calculations of the kinetics of the ferric ion catalyzed decomposition of H2O2 by Walling and Weil are not in contradiction to the “complex mechanism.” The examination of their results reveals that their simulations correspond to the terminal state of the system in which the secondary complex between Fe3+ and H2O2 becomes stabilized.

Notes: The unimolecular decomposition of chemically activated methylallylether (MAE) formed by the cross combination of methoxymethyl and vinyl radicals was studied in the gas phase. The experimentally determined rate constant was found to be 1.11 × 108 sec-1 at 9.6°C for the decomposition of MAE into propene and formaldehyde. The decomposition of MAE via the six-center retro-“ene” type transition state is analyzed by using the RRKM unimolecular reaction theory. For the molecular parameter assignments of energized MAE, a model which contains one internal rotational mode is supported, and MAE decomposition is characterized by a tight complex model. The best agreement between experimental and theoretical results was found when a critical energy of 40.1 kcal/mol was used.

Notes: The reaction H + CH3OOH was investigated under conditions of excess atomic hydrogen concentration using a flow reactor attached to a photoionization mass spectrometer. The rate coefficient of the reaction was determined as \documentclass{article}\pagestyle{empty}\begin{document}$$ k = (2.8 \pm 0.9) \times 10^{- 13} {\rm exp\,}[- (1860 \pm 190){\rm cal}/RT \cdot {\rm mol\,}]\,{\rm cm}^3 /{\rm molec} \cdot {\rm sec} $$\end{document} The three important reaction channels were found to be with the individual contributions determined as indicated. The product methoxy and methylperoxy radicals react mainly with atomic hydrogen under the employed experimental conditions according to where the estimates for the percentage contributions of the various channels were derived from the measured product yields.

Notes: The application of modern theories of energy transfer to unimolecular reactions taking place at very high temperatures is discussed. It is shown that the efficiency of energy transfer for both reactant-reactant and reactant-inert diluent collisions may be substantially smaller than the values determined experimentally at lower temperatures. Consequently at high temperatures unimolecular falloff effects, particularly in some shock-tube measurements, may be greater than has been believed hitherto.The application of these calculations to the unimolecular reactions of cyclopropane, cyclobutane, and cyclohexene at temperatures around 1300°K is discussed, and it is shown that under shock-tube conditions the apparent first-order rate coefficient may be at least ten times less than the high-pressure limiting value.

Notes: The kinetics of the gamma radiation induced free radical chain decomposition of BrCH2CN in liquid cyclohexane (RH) was investigated over the temperature range of 60-170°C. In addition competitive experiments in the presence of CCl4 were carried out between 80 and 180°C. For the reactions the following Arrhenius expressions were derived: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} {\rm log}k_3 /(2k_4)^{1/2} ({\rm l}^{{\rm 1/2}} /({\rm mol} \cdot {\rm sec})^{1/2}) = 4.07 \pm 0.35 - (11.96 \pm 0.63)/\theta \\ {\rm log}k_2 /k_5 = - 0.699 \pm 0.167) + (2.69 \pm 0.31)/\theta \\ \end{array} $$\end{document} where θ = 2.303RT in kcal/mol.The effect of CN substitution on the activation energies of reactions (2) and (3) was evaluated based on the present and previously published results. The CN group effect on halogen atom abstraction [reaction (2)] is discussed in terms of inductive and enthalpic factors.The differences E3 - E(CH3 + RH) and E(CCl2CN + RH) - E(CCl3 + RH), which yield a value of about 5.5 kcal/mol, are considered to reflect the cyano stabilization effect at the radical center confirming D(CH2(CN)-H) ∼ 93 kcal/mol.

Notes: Nanosecond flash photolysis of 1,2- and 1,8-dinitronaphthalenes (1,2-DNO2N; 1,8-DNO2N) in nonpolar and polar solvents shows transient species with absorption maxima and lifetimes dependent on solvent polarity. In deaerated n-hexane the absorption maxima and lifetimes (1/K) are 490 nm and 1.0 μsec for 1,2-DNO2N and 550 nm and 2.5 μsec for 1,8-DNO2N. In deaerated ethanol the corresponding values are 550 nm and 4.3 μsec for 1,2-DNO2N and 590 nm and 5.3 μsec for 1,8-DNO2N. The transient absorptions are attributed to the lowest triplet excited states T1 of the 1,2- and 1,8-DNO2N. The observed red shifts in the absorption maxima of the T1 states are indicative of the extent to which electronic charge is transferred intramolecularly during the T1 → Tn transitions. Furthermore, the increased lifetime of the T1 states with increasing solvent polarity indicates the intramolecular charge transfer character of the T1 states. Changes of dipole moments accompanying the T1 → Tn transitions as well as rate constants for electron or proton transfer and hydrogen abstraction reactions involving the T1 states of 1,2- and 1,8-DNO2N and tributyl tin hydride (Bu3SnH) as the hydrogen donor were determined together with the activation energy of the hydrogen abstraction reaction for the case of 1,2-DNO2N. The spectroscopic and kinetic data obtained in this work demonstrate that the triplet states of 1,2- and 1,8-DNO2N behave like n → π* states in nonpolar media while in polar solvents the n → π* character of these states is reduced with a simultaneous increase in their intramolecular charge transfer character.

Notes: Using the technique of molecular modulation spectrometry, we have measured directly the rate constants of several reactions involved in the oxidation of methyl radicals at room temperature: k1 is in the fall-off pressure regime at our experimental pressures (20-760 torr) where the order lies between second and third and we obtain an estimate for the second-orderlimit of (1.2 ± 0.6) × 10-12 cm3/molec · sec, together with third-order rate constants of (3.1 ± 0.8) × 10-31 cm6/molec2 · sec with N2 as third body and (1.5 ± 0.8) × 10-30 with neopentane; we cannot differentiate between k2a and k2c and we conclude k2a + (k2c) = (3.05 ± 0.8) × 10-13 cm3/molec · sec and k2b = (1.6 ± 0.4) × 10-13 cm3/molec · sec; k3 = (6.0 ± 1.0) × 10-11 cm3/molec · sec.

Notes: The reactions between SO2 and O2 were carried out in the presence of TiO2 and NiO under various partial pressures of SO2 and O2 at temperatures from 240 to 330°C. TiO2 and NiO were pretreated by applying an annealing effect from which the catalysts would have the different activity. The rates are the highest for TiO2 pretreated at high temperature in the region of 400 to 600deg;C in vacuum, 1.21 × 10-4 mmHg. In contrast, the rates are the lowest for NiO pretreated at high temperaturefrom 350 to 550°C. The data have been correlated with 1.4 and first-order kinetics for TiO2 and NiO, respectively. A reaction mechanism to explain the data was suggested.The quantities of anionic vacancies in TiO2 surfaces and of positive holes in NiO appeared to be paramount in determining the type of kinetics.

Notes: The general problem of eliciting reliable rate constants from experimental data is considered in detail for consecutive reactions. Practical aspects are emphasized.

Notes: The room temperature photolysis of 1,1-dichloroethane at 147 nm in the pressure range of 1.34-196.2 torr is characterized almost entirely by the molecular elimination of HCl, Cl2, and small quantities of H2. Acetylene is also produced. While it is possible that the C2H2 arises, in part, from the decomposition of vibrationally excited ground states of C2H3Cl and/or C2H4, in this particlar case serious consideration has to be given to alternative explanations where the products of the primary processes are formed in electronically excited states. The ±, elimination of molecular chlorine is not inconsistent with an increased degree of Cl—Cl interaction predicted for a «Rydberg «state of 1,1-C2H4Cl2. Varying small yields of CH4 are observed in the presence and absence of NO. The effect of large pressures of CF4 on the quantum yields of the major products is extremely small. The extinction coefficient for 1,1-C2H4Cl2 at 147 nm and 296°K is 246 ± 29 cm-1 ± atm-1.

Notes: NO Abstract.

Notes: 1,2-Fluorochloroethane was photolyzed at 147 nm in the pressure range of 3.8-20.9 torr. The effects of added NO, H2S, and large pressures of CF4 were also investigated. The exponential extinction coefficient at 147 nm and 296°K was found to be 147 ± 4 atm-1 · cm-1.The photochemistry in some respects is similar to that of ethyl chloride. The primary processes again appear to involve at least two excited states. One of these yields ethylene by FCl elimination (Φ ≃ 0.3) and has a lifetime of ∼3.2 × 10-10 sec, with respect to an internal conversion to the vibrationally excited ground state or, more probably, a collisionally induced crossover to a state decomposing mainly by carbon—halogen bond fission. The molecular elimination of HCl, H2, and small amounts of HF also occurs but not apparently from the same state as does FCl. The quantum yields of products with radical precursors, however, are not large, and hence little, if any, of the FCl and probably none of HCl, H2, and HF subsequently dissociates. The vinyl fluoride and vinyl chloride, accompanying the elimination of HF and HCl, are postulated as possible sources of the secondary production of acetylene.

Notes: Electron pulse radiolysis at ⋍298°K of 2 atm H2 containing 5 torr O2 produces HO2 free radical whose disappearance by reaction (1), HO2 + HO2 →H2O2 + O2, is monitored by kinetic spectrophotometry at 230.5 nm. Using a literature value for the HO2 absorption cross section, the values k1 = 2.5×10-12 cm3/molec·sec, which is in reasonable agreement with two earlier studies, and G(H) G(HO2) ⋍13 are obtained. In the presence of small amounts of added H2O or NH3, the observed second-order decay rate of the HO2 signal is found to increase by up to a factor of ⋍2.5. A proposed kinetic model quantitatively explains these data in terms of the formation of previously unpostulated 1:1 complexes, HO2 + H2O ⇋ HO2·H2O (4a) and HO2 + NH3⇋ HO2·NH3 (4b), which are more reactive than uncomplexed HO2 toward a second uncomplexed HO2 radical. The following equilibrium constants, which agree with independent theoretical calculations on these complexes, are derived from the data: 2×10-20≲K4a≲6.3 × 10-19 cm3/molec at 295°K and K4b = 3.4 × 10-18 cm3/molec at 298°K. Several deuterium isotope effects are also reported, including kH/kD = 2.8 for reaction (1). The atmospheric significance of these results is pointed out.

Notes: Rate constants have been estimated as a function of temperature for seven reactions of the type W + XYZ = WX + YZ, where W, X, Y, and Z are H and O atoms. From transition state theory and estimates of the heat capacities of activation, \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm int\, k/[cm}^{\rm 3} {\rm (mol} \cdot {\rm sec)] = 10}^{{\rm 15}{\rm .87}} \exp (\Delta S_{298}^{\circ _ \ne} /R)T^{0.75} \exp [- (\Delta H_{298}^{\circ _ \ne} + 0.74)/RT] $$\end{document} where int k is the rate constant per transferable atom for the forward and reverse reactions in the exothermic direction, and where ΔH°≠298 is in kcal/mol. Values of ΔS°≠298 and ΔH°≠298 were obtained from the above equation and previously measured and evaluated rate constants at 298°K. The results are summarized in a table. Rate constants were calculated at temperature from 250 to 2000 K. The estimated rate constants were compared with recommended values. The results for ΔH°≠298 for reactions (15), (16), (17), and (19), in which a stable intermediate may precede the transition state, together with similar results previously found for reactions X + YZ = XY + Z, suggests that many such reactions may have values of ΔH°≠298 that are close to zero. The result for the reaction O + O3 = O2 + O2 is however, an exception to the foregoing perhaps because it is the reaction of a singlet with a triplet.ΔS°≠298 for the same reaction is unexpectedly low.

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: The kinetics of the gas-phase thermal decomposition of 1,1,2-trifluoro-2-trifluoromethyl-cyclopropane to 2,3,3,3-tetrafluoropropene and CF2 was studied in the temperature range of 539-604°K and with a total pressure of 17 to 300 torr of a 1:20 mixture of reactant and perfluoroethylene. Some runs were also made with the reactant only, with initial pressures of 0.2 to 20 torr. The reaction is first order, homogeneous, and the rate constants fit the following Arrhenius relationship: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}\,k/{\rm sec}^{- 1}) = (15.23 \pm 0.12) - (50,557 \pm 1100)/4.576T $$\end{document} .

Notes: The thermal and methyl radical initiated decomposition of biacetyl-d6 were investigated at conversions 5% over the temperature range of 387-412°C and 251-292°C, respectively. The deuterated sample contained 10.5% biacetyl-d5. The reaction products were analyzed by mass spectrometry.The following data could be determined from product ratios: where E is in kcal/molWith reasonable assumptions and literature data Arrhenius parameters for elementary steps were derived and discussed in particular with respect to isotopic effects in H/D-abstraction reactions.

Notes: Rate constants for hydrogen transfer reactions from C - H bonds to O(3P) in the gas phase have been calculated by the bond-energy bond-order method. The use of a single or double triplet repulsion term is discussed. Computational results indicate that with carefully selected input parameters, it is unnecessary to use a double-repulsion term. Tunneling corrections were carried out by two different procedures. Results indicate that in this case there is not much difference between the simple Wigner's treatment and more complicated methods.

Notes: Arrhenius parameters are frequently evaluated incorrectly by applying the least squares method to the logarithmic form of the Arrhenius equation without simultaneously transforming the statistical weights as required for the change of variable from k to ln k. This has been mentioned briefly in an earlier paper. In the present communication the correct procedures are discussed and illustrated by several examples of evaluations. In particular, Arrhenius parameters calculated by the Taylor series expansion of the exponential form of the Arrhenius equation are compared with those calculated from the logarithmic form, using an exact and two approximate transformations of the statistical weights. The comparisons indicate thatthe preferred procedure for obtaining Arrhenius parameters is either the Taylor series method or the logarithmic method with proper transformation of the experimentally determined statistical weights of the rate constants ki. The common approximation of assuming equal statistical weights of ln ki when the logarithmic form of the Arrhenius expression is used is shown not to be always appropriate, and reasons forthis are given.

Notes: The thermal decomposition of trifluoroacetaldehyde at temperatures f 471 to 519°C has been studied by measuring the rates of formation of CF3H and C2F6. It is concluded that the high-pressure reaction mechanism involves a Rice-Herzfeld reaction scheme with first-order initiation and second-order termination via CF3 combination. However, a falloff in reaction rates is observed at pressures below 100 mmHg. The Arrhenius parameters of the three rate constants corresponding to the overall reaction, the initiation reaction, and an abstraction reaction have been evaluated.

Notes: A kinetic investigation on the reaction has been carried out in HClO4 medium under different conditions. A spectrophotometric method of estimation of nitrous acid at various time intervals has been employed. The results are interpreted on the basis ofthe following mechanism: The absolute rate constant value of 39.7 M-1 plusmn; s-1 for k4 and the equilibrium constant Keq = 116M-1 for reaction (2) have been evaluated. The activation energy of the overall reaction has also been determined as Ea = 13.2 kcal/mol.

Notes: The thermal gas-phase decomposition of vinyl chloride has been studied behind shock waves over the temperature range of 1350-1900°K and the density range of 7 × 10-7-1.5 × 10-3 mol/cm3 (at 1600°K) in mixtures of C2H3Cl highly diluted with argon. The ultraviolet absorption of C2H3C was recorded at 230 nm as a function of time. The decomposition proceeds via molecular elimination of HCl. The unimolecular dissociation rate is pressure dependent at all but the highest pressures applied. Application of modified HKRR theory results in the rate expression \documentclass{article}\pagestyle{empty}\begin{document}$$ k_\infty = 10^{14.0 \pm 0.2}\, {\rm exp}[= (290 \pm 12){\rm kJ}/{\rm mol}/RT]\,{\rm s}^{- 1},\,\quad 1400^ \circ {\rm K} \le T \le 1800^ \circ {\rm K} $$\end{document} for the limiting high pressure rate constant, and in a collision efficiency of \documentclass{article}\pagestyle{empty}\begin{document}$$ \beta _c = \frac{{k_0 ({\rm experiment})}}{{k_{\rm 0} {\rm (theoretical, strong collisions})}} \cong 0.04,\quad {\rm for M} = {\rm Ar},T = 1600^ \circ {\rm K} $$\end{document} derived from the limiting low-pressure rate constant.

Notes: The rate of decomposition of ethyl nitrite (EN) has been studied in a static system over the temperature range of 162-218°C. The main products are formaldehyde, acetaldehyde, ethanol, and nitrous oxide. For low concentrations of EN (10-5-10-4M), but with a high total pressure of CF4 (∼0.9 atm) and small extents of reaction (2-6%), the first-order homogeneous rates of CH2O formation are a direct measure of reaction (1), since k3bk2(NO): Addition of large amounts of NO(∼0.9 atm) completely suppressed CH2O formation in agreement with the observed value for k3b.The rate of reaction (1) is given by k1 = 1016.0-41.8/θ-1. Since (E1 + RT) and ΔH±1 are identical, both may be equated with D(EtO-NO) = 42.4 ± 0.9 kcal/mol and E2 = O± 1 kcal/mol. The thermochemistry leads to the result ΔHDelta;f(EN) = -24.5 ± 1 kcal/mol. From ΔS1 and A1, k2 is calculated to be 1010.3M-1θ-1. From an independent observation that k6/k2 = 0.3 ± 0.05 independent of temperature \documentclass{article}\pagestyle{empty}\begin{document}$$ {{\rm EtO + NO}} \stackrel{6}{\longrightarrow} {{\rm AcH} + {\rm HNO}} $$\end{document} it is concluded that k6 = 109.8M-1Δ-1.The addition of NO has no effect on the AcH yields. Although the yields of AcH are affected by the surface-to-volume ratio of different reaction vessels, it is concluded that in a spherical reaction vessel, the AcH arises as the result of an essentially homogeneous elimination of HNO from EN(5): \documentclass{article}\pagestyle{empty}\begin{document}$$ {{\rm EN}} \stackrel{5}{\longrightarrow} {{\rm AcH} + {\rm HNO}} $$\end{document} and reaction (6). The rate of AcH formation is given by kobs = 1013.7-37.5/θ-1. By using isobutane (t-BuH) as a radical trap for EtO (4), \documentclass{article}\pagestyle{empty}\begin{document}$$ {{\rm EtO} + t - {\rm BuH}} \stackrel{4}{\longrightarrow} {{\rm EA} + (t - {\rm Bu})} $$\end{document} a value for k3b was determined to be 1015.0-21.6/θ s-1.From an independent observation that k2:k2:k6:k6 was 1: 0.4: 0.3: 0.18 we find k2θ = 109.9M-1→ s-1, k1θ = 1016.0-40.0/θ s-1, and k6± = 109.6M-1 · s-1.

Notes: High-intensity pulse radiolysis, using a Febetron, has made it possible to obtain in a few microseconds homogeneous H2O2 concentrations of about 10-4 to 10-3M without any other reactive radiolytical product. The kinetics of the reaction of this H2O2 with a solute have been studied as a function of time for periods ranging from tens of microseconds to tens of milliseconds.An analytical model of this method is presented and then compared with the results of experiments in which the H2O2 concentrations are measured. The rate constant for the peroxidation of pentavalent vanadium is determined to be k = (5.7 ± 0.5) × 103M-1 plusmn; s-1.

Notes: The rate of demetallation of α, β, γ,δ-tetra(p-sulfophenyl)porphineiron (III), Fe(TPPS)3-, was determined in sulfuric acid-ethanol-water media for 8.5-10.65M sulfuric acid at different temperatures. The overall reaction was the conversion of the complex Fe(TPPS)3- into the diacid species H4TPPS2- without other spectrophotometrically important species being formed to an appreciable extent, as shown by three isosbestic points at 418, 462, and 563 nm. The rate was first order in the Fe(TPPS)3- concentration. The pseudo-first-order rate constants k were exponentially dependent on the sulfuric acid concentration, and log k was linearly dependent on the Hammett acidity function -H0. The average ΔH

Notes: The reaction between ozone and thiophene was studied from 30 to 125°C over a pressure range of 0.005-0.3 torr ozone and 0.1-1 torr thiophene. The most abundant product was O2 with smaller amounts of CO2 and SO2. The mass balance was 100% for oxygen and approached 100% for sulfur at the higher values of [O3]0. The carbon balance, however, was only 25% and no H-containing products were found, suggesting that the missing product is a hydrocarbon which may be a polymer. The rate law found was -d[O3]/dt = k1[Th] [O3] + k2 [Th] [O3]2 where log k1(M-1 · sec-1) = 7.8 ± 0.5 - (8400 ± 700)/2.3RT, and log k2(M-2 · sec-1) = 12.4 ± 0.4 - (4700 ± 400)/2.3RT. Added O2 had no effect on k1 but reduced k2 to a limiting value. It is thus not possible to measure the primary rate constant in this system by measuring the overall rate in the presence of oxygen, and this restriction may also apply to other ozone systems. A mechanism is postulated involving two chain sequences, one of which is inhibited by added O2. A comparison with other ozone systems is made, and the chain lengths are far greater for ozone + thiophene than other systems, under the conditions employed. Possible intermediates in the mechanism are discussed.

Notes: The mechanism of the interaction of Cu+-α,α-dipyridyl complex (Cu+L2) with O2 in both neutral and acid media was studied by the stopped-flow method. The dependence of the mechanism on the acidity of the medium was established. In an acid medium H+ participated in a direct O2 reduction to HO2 by interaction with an oxygen adduct L2Cu+O2 formed without displacement of ligand molecules. In a neutral medium the reaction rate was limited by inner sphere charge transfer from Cu+ to O2 to form an oxygen “charge transfer” complex L2CuO+2. The latter interacted either with the second ion Cu+L2 or with the free ligand, or else it dissociated, reversibly or irreversibly, to form a radical anion O-2. The bimolecular rate constants of the oxygen “adduct” and “charge transfer” complex formation appeared to be kbi = (1.0 ± 0.1) × 105 and (1.5 ± 0.2) × 104M-1ċsec-1, respectively. The effective termolecular rate constants of O2 reduction to HO2 in an acid medium (with contribution from H+) and to O-2 in a neutral medium (with contribution from α,α-dipyridyl) were kter = 2.7 × 108 and 107M-2ċsec-1. The rate constants of the elementary steps were estimated. The auto-oxidation mechanism of the aquoion and complexes of Cu+ is discussed in terms of the results obtained.

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}.