ALBERT

Notes: The flash photolysis-vacuum ultraviolet kinetic absorption spectroscopy technique has been used to measure the absolute rate constant for the reaction of ground state S(3P) atoms withnitric oxide,\documentclass{article}\pagestyle{empty}\begin{document}${\rm S}\left({^{\rm 3} P} \right) + {\rm NO}\mathop {\longrightarrow}\limits^{\rm M} {\rm SNO}\left({{\rm M} = {\rm CO}_2} \right)$\end{document} as a function of nitric oxide concentration and total pressure. The rateconstant was determined to be 1.9±0.1 × 1011 12/mol2.sec at 298°K, with a high-pressure limit of 9.3 ± 2.1×109 l/mol·sec-1. The observed kinetics are consistent with a termolecular energy transfer mechanism.

Notes: The title reaction has been investigated in the temperature range 667-715K. The only reaction products were trifluorosilyl iodide and hydrogen iodide. The rate law \documentclass{article}\pagestyle{empty}\begin{document}$$ - \frac{{d\left[{{\rm I}_2} \right]}}{{dt}} = \frac{{k\left[{{\rm I}_2} \right]^{1/2} \left[{{\rm F}_3 {\rm SiH}} \right]}}{{1 + k\prime \left[{{\rm HI}} \right]/\left[{{\rm I}_2} \right]}} $$\end{document} was obeyed over a wide range of iodine and trifluorosilane pressures. This expression is consistent with an iodine atom abstraction mechanism and for the step \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm I}^ \cdot + {\rm F}_3 {\rm SiH}\mathop {\longrightarrow}\limits^1 {\rm F}_3 {\rm Si}^\cdot + {\rm HI} $$\end{document} log k1(dm3/mol·sec) = (11.54 ± 0.17) - (130.5 ± 2.2 kJ/mol)/RT In 10 has been deduced. From this the bond dissociation energy D(F3Si—H) = (419 ± 5) kJ/mol (100.1 kcal/mol) is obtained. The kinetic andthermochemical implications of this value are discussed.

Notes: The recombination of bromine atoms at room temperature has been studied by flash photolysis in the range of 1-100 atm of the inert diluent He, leading to a value for the third-order rate constant of (1.5 ± 0.2) × 1015 cm6/mol2.sec. In the presence of NO the recombination is considerably accelerated. The falloff curve of the recombination Br + NO (+He) → BrNO (+He) was also measured resulting in a value for the limiting low-pressure rate constant of (3.4 ± 1.3) × 1015 cm6/mol2.sec. In experiments with excess NO, rate constants of (2.2 ± 1) × 1014 cm3/mol·sec for the reaction Br + BrNO → Br2 + NO, and (6.1 ± 0.4) × 109 cm6/mol2.sec for the reaction Br2 + 2NO → 2BrNO were obtained.

Notes: A previously developed model for active species concentration profiles in infinite cylindrical systems has been extended to include the spherical system. The model couples the processes of diffusion to and reaction at the wall. Predictions of time buildup under conditions of homogeneous production by light, and time decay after extinguishing the light source, are made for H atoms. Such predictions require a knowledge of the wall recombination coefficient and the binary diffusion coefficient for H in heat bath gas. The model is experimentally tested by measuring the first-order decay constants of H at room temperature in various pressures (10-1500 torr) of six heat bath gases. The atomic concentration is monitored by Lyman-α absorption photometry. The results show good agreement with model predictions in the various heat bath gases up to ∼400 torr and depend only on one parameter,γ, the recombi-nation coefficient. This should be contrasted with the earlier work where slight variation in γ was invoked. The rate constants at pressures higher than 400 torr are consistently higher than model predictions.

Notes: Atomic resonance absorption spectrometry with a nonreversed fluorine resonance lamp (∼95 nm) has been used to study the kinetics of elementary reactions of ground state F2PJ atoms in a discharge-flow system. The following rate constants (in cm3/molec·sec)All rate constants are given with 1.5 σ. were determined at 298° K: The reaction F 2PJ + HCl(1) was found to give J-excited Cl 2P1/2 atoms with a product branching ratio [Cl 2P1/2]/[Cl 2P3/2] = 0.10.

Notes: Equations for the average energy of chemically activated species are developed, and uncertainties in the various energy quantities involved are discussed. Various approaches to the energy distribution function of chemically activated species are discussed. Trial calculations on methylcyclobutane are presented.

Notes: The hydrogen-atom recombination reaction has been simulated using a molecular dynamics technique recently formulated by the authors [1]. The rate of recombination has been calculated over a range of temperatures and inert gas concentrations (He and Ar) and agrees well with available experimental data. The calculations reproduce the negative activation energy characteristic of an atom recombination process. Over the range of conditions studied recombination was found to proceed via the energy transfer mechanism only, no evidence of bound HAr or HHe species was observed. Recombination was found to occur through an intermediate metastable diatomic molecule which is in equilibrium with its environment and from which there is a bottleneck to the formation of a stable molecule. The initial formation of a metastable species is sensitive to the hydrogen-inert gas potential, but relaxation of the total energy is primary determined by the mass of the third-body and the collision frequency.

Notes: The title reaction, displaying peculiar characteristics as to relative rates and isomer distributions, has been studied in detail. Prior to this study, different mechanisms had been advanced by several groups. Kinetic features (isomer patterns, relative and absolute rates, reaction orders, influences of additives, H/D isotope effects) strongly point to a free-radical (chain) process, in which (1) is a crucial step. This abstraction reaction, endothermal by about 6 kcal/mol, apparently proceeds via a transition state closely resembling the free aryl radical. Relative rates and isomer distributions therefore reflect differences in stabilization energies, or in DH°(Ar - H). With high arene-Cl2 intake ratios or, more pronounced, with CCl4 as the reagent, aryl radicals also lead to biaryl, where arene successfully competes with the halogenating agent. This interpretation is quantitatively supported by our observation that “added,˝” recognizable aryl radicals yield the same chlorination-arylation product ratio, and by the results of competitive chlorination of benzene and chloroform over a temperature range of 200°C, where the latter study substantiates the value DH0(C6H5 - H) ≈ 109 kcal/mol.

Notes: The pyrolysis of ethylene-butene-2 mixtures has been studied in a static system over the temperature range of 689°-754°k and for initial pressures of each olefin of 20-200 torr. The two main addition products were cyclopentene and 3-methylpentene-1. Kinetic evidence indicated that cyclopentene was formed from radical processes while 3-methylpentene-1 was formed by the molecular “ene¨” addition of ethylene to butene-2 through a six-center transition state. The following rate constants were obtained: The pyrolysis of 3-methylpentene-1 has been studied over the same temperature range and for initial pressures of 20-100 torr. Kinetic evidence showed that the products ethylene and butenes were formed in both radical and molecular processes. Estimates of the rate constant k-1t and k-1c were, however, in reasonable agreement with the measurements of k1t and k1c. The mechanism of the ene reaction is discussed, and it is concluded that the transition state does not involve the formation of a biradical.

Notes: The pyrolysis of ethylene-propylene and ethylene-isobutene mixtures has been studied in a static system over the temperature range of 682°-754° K and for initial pressures of each olefin of 33-300 torr. The following molecular ene reactions were observed and the rate constants measured: Using thermodynamic data, rate constants for the corresponding retro-ene decomposition reactions were calculated and compared to kinetic data reported for similar compounds. Other products were formed by radical chain processes, the main higher molecular weight ones being cyclopentene and 1-methylcyclopentene. A mechanism involving addition of allyl radicals is suggested for the formation of these products.

Notes: A program system is described for the integration of the rate equations resulting from large systems of elementary reactions. The Gear integration method is used for this problem, which frequently may exhibit stiffness instability when other integration methds are employed. No usage of the quasi-steady-state approximation is necessary. Ease in varying the reaction mechanism and simplicity of input structure are coupled with efficient execution and minimal demands for program storage as key features. The input-output structure, method of operation, and implementation are summarized, along with core storage requirements and execution times for trials using an IBM 360/44 computer.

Notes: The thermal reaction of 2-pentene (cis or trans) has been performed in a static system over the temperature range of 470°-535°C at low extent of reaction and for initial pressures of 20-100 torr. The main products of decomposition are methane and 1,3-butadiene. Other minor primary products have been monitored: trans-2-pentene, trans- and cis-2-butenes, ethane, 1,3-pentadienes, 3-methyl-1-butene, propylene, 1-butene, hydrogen, ethylene, and 1-pentene. The initial orders of formation, 0.8-1.1 for most of the products and 1.5-1.8 for 1-pentene, increase with temperature. The formation of the products and the influence of temperature on their orders can be essentially explained by a free radical chain mechanism. But cis-trans or trans-cis isomerization and hydrogen elimination from cis-2-pentene certainly involve both molecular and free radical processes. The formation of 1-pentene mainly occurs from the abstraction of the hydrogen atom of 2-pentene by resonance stabilized free radicals (C5H9.).

Notes: The kinetics of the thermal bromination reaction have been studied in the range of 261°-391°C. The observed rate law is compatible with initiation by the step for which we obtain \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k_6 (cm^3 /mol - \sec) = (13.41 \pm 0.26) - (11700 \pm 700)/\theta $$\end{document} where Θ = 2.303RT cal/mol. Using the above value of E6, we have \documentclass{article}\pagestyle{empty}\begin{document}$$ {D(C}_{\rm 6} {\rm F}_{{\rm 5}^{\rm -}} {\rm I)} \to {\rm 53.5}\ {\rm kcal/mol} $$\end{document}This result disagrees with values of D(C6F5-I) obtained in other ways and we conclude that reaction (3) probably does not involve initiation by reaction (6). Instead, initiation may involve an addition of Br to the ring in C6F5I followed by decomposition of the adduct to give C6F5Br. If correct, this implies that the Arrhenius parameters above refer to the addition reaction rather than to reaction (6).

Notes: The rate constants and modes of reaction of NO2+ and C2H5ONO2NO2+ with aromatic compounds and alkanes have been determined in a pulsed ion cyclotron resonance mass spectrometer. Both ions undergo competing charge transfer and substitution reactions (NO2+ + M → MO+ + NO; C2H5ONO2NO2+ + M → MNO2+ + C2H5ONO2) with aromatic molecules. In both cases, the probability that a collision results in charge transfer increases with increasing exothermicity of that process. The C2H5ONO2NO2+ ion does not undergo charge transfer with molecules having an ionization potential greater than about 212 kcal/mol (9.2 eV); this observation leads to an estimate of 13 kcal/mol for the binding energy between NO2+ and C2H5ONO2. The importance of the substitution reaction depends on the number of substituents on the aromatic ring and the molecular structure, and, in the case of C2H5ONO2NO2+ ions, on the energetics of the competing charge transfer process. Both NO2+ and C2H5ONO2NO2+ undergo hydride transfer reactions with alkanes. For both these ions, k(hydride transfer)/k (collision) increases with increasing exothermicity of reaction, but in both cases the rate constants of reaction are unusually low when compared with other hydride transfer reactions of comparable exothermicity which have been reported in the literature. This is interpreted as evidence that the attack on the alkane preferentially involves the nitrogen atom (where the charge is localized) rather than one of the oxygen atoms of NO2+.

Notes: A general competitive method is described for the study of the kinetics of the reactions of radicals with halogens and interhalogen compounds in the gas phase. The method is applied to the reactions over the temperature range of 48°-199°C. It is found that \documentclass{article}\pagestyle{empty}\begin{document}$$ \log (k_6 /k_5) = \log (k_\varepsilon /k_7) = (- 0.042 \pm 0.060) - (8700 \pm 1300)/\theta $$\end{document} where Θ=2.303RT J/mol. Also \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} A_7 = 1/2A_{5,} \,\,\,\,\,\,\,\,\,\,\,\,E_7 = E_5 \\ A_8 = 1/2A_{6,} \,\,\,\,\,\,\,\,\,\,\,\,E_8 = E_6 \\ \end{array} $$\end{document}There is no correlation between differences in activation energies and differences in enthalpy changes for these reactions, but polar effects may be important in reactions (7) and (8).

Notes: 4-Methylhexyne-1, 5-methylhexyne-1, hexyne-1, and 6-methylheptyne-2 have been decomposed in comparative-rate single-pulse shock-tube experiments. Rate expressions for the initial decomposition reactions at 1100°K and from 2 to 6 atm pressure are \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}s {\rm C}_{\rm 4} {\rm H}_{\rm 9} \to {\rm HC} \equiv {\rm CCH}_{\rm 2} \cdot + s{\rm C}_{\rm 4} {\rm H}_{\rm 9} \cdot) = 10^{15.9} \exp (- 35,000/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}i {\rm C}_{\rm 4} {\rm H}_{\rm 9} \to {\rm allene} + n{\rm C}_{\rm 4} {\rm H}_{\rm 8}) = 10^{12.9} \exp (- 28,000/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}i {\rm C}_{\rm 4} {\rm H}_{\rm 9} \to {\rm HC} \equiv {\rm CCH}_{\rm 2} \cdot + i{\rm C}_{\rm 4} {\rm H}_{\rm 9} \cdot) = 10^{16.1} \exp (- 36,700/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}i {\rm C}_{\rm 4} {\rm H}_{\rm 9} \to {\rm allene} + i{\rm C}_{\rm 4} {\rm H}_{\rm 8}) = 10^{2.3} \exp (- 27,500/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}n {\rm C}_{\rm 3} {\rm H}_{\rm 7} \to {\rm HC} \equiv {\rm CCH}_{\rm 2} \cdot + n{\rm C}_{\rm 3} {\rm H}_{\rm 7} \cdot) = 10^{15.9} \exp (- 36,300/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm HC} \equiv {\rm CCH}_{{\rm 2}^{{\rm -}}}n {\rm C}_{\rm 3} {\rm H}_{\rm 7} \to {\rm allene} + n{\rm C}_{\rm 3} {\rm H}_{\rm 6}) = 10^{12.7} \exp (- 28,400/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm CH}_3 {\rm C} \equiv {\rm CCH}_{2^{-}}i {\rm C}_4 {\rm H}_9 \to {\rm CH}_3 {\rm C}) \equiv {\rm CCH}_{\rm 2} \cdot + i{\rm C}_{\rm 4} {\rm H}_{\rm 9} \cdot) = 10^{16.2} \exp (- 36,800/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm CH}_3 {\rm C} \equiv {\rm CCH}_{2^{-}}i {\rm C}_4 {\rm H}_9 \to 1,2-butadiene + i{\rm C}_{\rm 4} {\rm H}_{\rm 8}) = 10^{12.3} \exp (- 28,700/T)\sec ^{- 1} $$\end{document} In combination with previous results, rate expressions for propargyl C—C bond cleavage are related to that for the alkanes by the expression \documentclass{article}\pagestyle{empty}\begin{document}$$ k_{\rm B} (alkyne) = \frac{1}{{3 \pm 1.5}}\exp (+ 4.25/T)k_{\rm B} (alkane) $$\end{document} These results yield a propargyl resonance energy of D(nC3H7-H) - D(C3H3-H) = 36 ± 2 kJ, in excellent agreement with a previous shock-tube study. They also lead to D(CH3C≡CCH2-H) - D(C3H3-H) = 0.6 ± 3 kJ, D(sC4H9-H) - D(iC3H7-H) = 0 ± 3 kJ, D(iC4H9-H) - D(nC3H7-H) = 2 ± 3 kJ, and D(nC3H7-H) - D(iC3H7-H) = 13.9 ± 3 kJ (all values are for 300°K). The systematics of the molecular decomposition process are explored.

Notes: The reactions of hydroxyl radicals with eight substituted aromatic hydrocarbons and four olefins were studied utilizing the flash photolysis-resonance fluorescence technique. The rate constants were measured at 298°K using either Ar or He as the diluent gas. The values of the rate constants (k × 1012) in the units of cm3/molec. sec are (a) OH + o-xylene → products: (12.9±3.8), 20 torr He; (13.0±0.3), 20 torr Ar; (12.4±0.1), 200 torr He;(b) OH + m-xylene → products: (15.6±1.4), 3 torr Ar; (19.4±0.8), 20 torr Ar; (21.4±0.2), 20 torr He; (20.3±1.9), 200 torr Ar; (20.6±1.3), 200 Torr He;(c) OH + p-xylene → products: (8.8±1.2), 3 torr Ar; (10.1±1.0), 20 torr He; (10.5±0.6), 200 torr He;(d) OH + ethyl benzene → products: (7.50±0.38), 3 torr He; (7.06±0.26), 20 torr He; (7.95±0.28), 200 torr He;(e) OH + n-propylbenzene → products: (6.40±0.36), 20 torr He; (5.86±0.16), 200 torr He;(f) OH + isopropylbenzene → products: (7.79±0.40), 200 torr He;(g) OH + hexafluorobenzene → products: (0.221±0.020), 20 torr He; (0.219±0.016) 200 torr He;(h) OH + n-propyl pentafluorobenzene → products: (2.52±0.54), 3 torr He; (3.01±0.76), 20 torr He; (3.06±0.24), 200 torr He;(i) OH + propylene → products: (25.6±1.2), 20 torr He; (26.3±1.2), 200 torr He;(j) OH + 1-butene → products: (29.6±1.9), 3 torr He; (29.4±1.4), 20 torr He;(k) OH + cis-2-butene → products: (43.2±4.1), 3 torr He; (42.6±2.5), 20 torr He;(l) OH + tetramethylethylene → products: (56.9±1.3), 20 torr He.

Notes: The first-order kinetics and hydrogen kinetic isotope effect of the decarboxylation of oxalic acid in acetophenone were studied in the temperature range of 109.6°-150.0°C. The rate constants, activation parameters, and hydrogen kinetic isotope effect were calculated. Detailed comparison and discussion of the results were made with the data reported in the literature. Kinetic isotope effects and solvent effects on rates should be considered similar in mechanistic and/or theoretical studies in the sense that kinetic isotope effects result from a small perturbation of the reaction coordinate, while the solvent effect causes a general overall variation on the potential energy surface (thereby resulting in a change in the reaction coordinate).

Notes: The mechanism of the photolysis of formaldehyde was studied in experiments at 3130 Å and in the pressure range of 1-12 torr at 25°C. The experiments were designed to establish the quantum yields of the primary decomposition steps (1) and (2), CH2O + hν → H + HCO (1): CH2O + hν → H2 + CO (2), through the effects of added isobutene, trimethylsilane, and nitric oxide on ΦCO and ΦH2. The ratio ΦCO/ΦH2 was found to be 1.01 ± 0.09(2σ) and (ΦH2 + ΦCO)/2 = 1.10 ± 0.08 over the range of pressures and a 12-fold change in incident light intensity. Isobutene and nitric oxide additions reduced ΦH2 to about the same limiting value, 0.32 ± 0.03 and 0.34 ± 0.04, respectively, but these added gases differed in their effects on ΦCO. With isobutene addition ΦCO/ΦH2 reached a limiting value of 2.3; with NO addition ΦCO exceeded unity. The addition of small amounts of Me3SiH reduced ΦH2 to 1.02 ± 0.08 and lowered ΦCO to 0.7. These findings were rationalized in terms of a mechanism in which the “nonscavengeable,” molecular hydrogen is formed in reaction (2) with φ2 = 0.32 ± 0.03, while the “free radical” hydrogen is formed in reaction (1) with φ1 = 0.68 ± 0.03. In the pure formaldehyde system these reactions are followed by (3)-(5): H + CH2O → H2 + HCO (3); 2HCO → CH2O + CO (4); 2HCO → H2 + 2CO (5). The data suggest k4/k5 ≅ 5.8. Isobutene reduced ΦH2 by the reaction H + iso-C4H8 → C4H9 (20), and the results give k20/k3 ≅ 43 ± 4, in good agreement with the ratio of the reported values of the individual constants k3 and k20.

Notes: The oxidative cleavage of some aliphatic ketoximes by thallium(III) acetate was studied in the temperature range of 20-40°C. The reactions were followed by determination of the rates of disappearance of thallium(III) acetate for variations in [substrate], [Tl(III)], [H+], ionic strength, temperature, etc. The reactions were found to be totally second order-first order with respect to each reactant. The second-order rate constants and thermodynamic parameters were evaluated and discussed. The mechanism proposed involves one-electron oxidation to the iminoxy radical followed by an another one-electron oxidation to the hydroxynitroso compound which dimerizes and decomposes to give the carbonyl compounds and hyponitrous acid.

Notes: The rate of disappearance of C2N2 in the presence of a large excess of H atoms has been measured in a discharge-flow system at pressures near 1 torr and temperatures in the range of 282-338 K. Under these conditions the reaction has a small negative temperature coefficient. A transition from second-order to third-order kinetics with decreasing pressure occurs at pressures near 1 torr. The results are discussed in terms of the mechanism where k7 = (1.5 ± 0.2) × 10-15 cm3/molec1·sec is found for the forward rate of reaction (7). The results also give k7k8/k-7 = 3.7 × 10-31 cm6/molec2·sec and k7k9/k-7 = 3.0 × 10-32 cm6/molec2·sec, the first being probably an upper limit and the second probably a lower limit; hence k8/k9 = 12 is found as an upper limit.

Notes: Oxidative kinetics of diethyl ketone in perchloric acid media in the presence of mercuric acetate have been studied by using N-bromosuccinimide (NBS) as oxidant in the temperature range of 25°-50°C. It has been found that the order with respect to NBS is zero while with respect to diethyl ketone and [H+], it is unity. Succinimide, sodium perchlorate, and mercuric acetate have an insignificant effect on the reaction rate, while the dielectric effect was negative. A solvent isotope effect (k0D2O/k0H2O = 1.6-1.8) at 35°C has been observed. On the basis of the available evidences a suitable mechanism consistent with the experimental results has been proposed in which it is suggested that the mechanistic route for NBS oxidation in an acidic medium is through the enol form of the ketone. The magnitude of the solvent effect also supports the mechanism. Various activation parameters have been calculated, and the 1,2-dicarbonyl compound has been identified as the end product of the reaction.

Notes: Vinylacetylene was pyrolyzed at 300-450°C in a packed and an unpacked static reactor with a pinhole bleed to a quadrupole mass spectrometer. The reactant and C8H8 products were monitored continuously during a reaction by mass spectrometry. In some runs, the products were also analyzed by gas chromatography after the run. In these runs CH4, C2H6, C3H6, and C2H4 were also detected.The reaction for vinylacetylene removal and C8H8 formation is homogeneous, second order in reactant, and independent of the presence of a large excess of N2 or He. However, C8H8 formation is about half-suppressed by the addition of the free-radical scavengers NO or O2. The rate coefficient for total vinylacetylene removal is 1.7 × 106 exp(-79 ± 13 kJ/mol RT) L/mol · s. The major reaction for C4H4 removal is polymerization. In addition four C8H8 isomers, carbon, and small hydrocarbons are formed. The three major C8H8 isomers are styrene, cyclooctatetraene (COT), and 1,5—dihydropentalene (DHP).The C8H8 compounds are formed by both molecular and free-radical processes in a second-order process with an overall k ≃ 3 × 108 exp(-122 kJ/mol RT) L/mol · s (average of packed and unpacked cell results). The molecular process occurs with an overall k = 8.5 × 107 exp (-118 kJ/mol RT) L/mol · s. The COT, DHP, and an unidentified isomer (d), are formed exclusively in molecular processes with respective rate coefficients of 4.4 × 104 exp(-77 kJ/mol RT), 1.7 × 105 exp(-89 kJ/mol RT), and 3.1 × 109 exp(- 148 kJ/mol RT) L/mol · s. The styrene is formed both by a direct free-radical process and by isomerization of COT.

Notes: The product quantum yields in the photolysis of 2,2,4,4-tetramethyl-3-pentanone have been measured in homogeneous solvents of different viscosities, in micellar solutions of cetyltrimethylammonium chloride and sodium dodecyl sulfate, and in dioctadecyl ammonium chloride vesicles.The product quantum yield in n-heptane was found to be 1. This value decreases to 0.5 in paraffin oil as a consequence of geminate recombination. In the presence of free radical scavengers, the extent of geminate disproportionation can be evaluated from the yields of isobutene and 2,2-dimethyl propionaldehyde. From these yields and the geminate recombination yields the total amount of geminate processes and the disproportionation-to-combination ratio for caged radicals are estimated. It is found that micelles provide the most efficient cages. In these media only about 10% of the radicals avoid cage processes. The disproportionation-to-combination ratio of tert-butyl and pivaloyl radicals was found to be extremely media dependent. The measured values ranged from about 0.2 in paraffin oil to 0.8 in cetyltrimethylammonium chloride micelles.

Notes: Kinetics of the thermal decomposition of acetic acid vapor dilute in argon have been studied over the temperature range of 1300-1950 K in a single-pulse shock tube. The acid was found to decompose homogeneously and molecularly via two competing firstorder reaction channels at nearly equal rates, to form methane and carbon dioxide on the one hand, and ketene and water on the other. Fall-off behavior has been taken into account and limiting high-pressure rate constants for both channels have been derived. Ketene was found to decompose both unimolecularly to methylene radicals and carbon monoxide and also by a radical reaction with CH2 to form ethylene and carbon monoxide. The rate constant derived for the unimolecular reaction was found to be in good agreement with an earlier shock tube measurement by H. G. Wagner and F. Zabel [Ber. Bunsenges Phys. Chem., 75, 114 (1971)]. The bimolecular reaction of ketene to produce allene and carbon dioxide, important in lower temperature reaction systems, has been found to be unimportant under the present conditions. A computer model for the decomposition kinetics involving 46 reactions of 21 species has been found to simulate the experimental yield data substantially. Sensitivity analyses have been used to identify reactions which make important contributions to the overall mechanism and yields of major products. Methylene radicals play important roles in determining yields of major species.

Notes: The abstraction of hydrogen/deuterium from CH3CH2Cl, CH3CHDCl, and CH3CD2Cl by photochemically generated ground-state chlorine atoms has been investigated over the temperature range of 8-94°C using methane as a competitor. Rate constant data for the following reactions have been obtained:The temperature dependence of the relative rate constants ki/kj was found to conform to the Arrhenius rate law, where the stated error limits are one standard deviation:\documentclass{article}\pagestyle{empty}\begin{document}$$ k_1 /k_2 = (1.099 \pm 0.015)\exp [(429 \pm 2)/T] $$ $$ k_1 /k_r = (1.422 \pm 0.026)\exp [(1113 \pm 3)/T] $$ $$ k_2 /k_r = (1.295 \pm 0.029)\exp [(684 \pm 3)/T] $$ $$ k_3 /k_r = (1.177 \pm 0.025)\exp [(717 \pm 4)/T] $$ $$ k_4 /k_r = (1.115 \pm 0.023)\exp [(732 \pm 2)/T] $$ $$ k_5 /k_r = (0.978 \pm 0.020)\exp [(985 \pm 2)/T] $$\end{document} and kr is the rate constant for the reference reaction (CH4 + Cl → CH3 + HCl). The β secondary kinetic isotope effects (k2/k3/k4) are close to unity and show a slight inverse temperature dependence. Both preexponential factors and activation energies decrease as a result of deuterium substitution in the adjacent chloromethyl group. The trends are well outside the limits of experimental error.

Notes: The reaction between ozone and carbon monoxide was reinvestigated in the range of 80-160°C. The previously reported rate law -d[O3]/dt = ka[O3][CO] + kb[O3]2 was confirmed and simulated using a mechanism based on an impurity-initiated chain reaction. When the CO was sufficiently purified, kb tended to zero and ka reduced to the value expected for the thermal decomposition of O3. Subsequent reactions of O atoms with CO produced chemiluminescence which was used to measure k3 for \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm O}\,\, + \,\,{\rm CO}\,\,\mathop {\longrightarrow}\limits^{\rm 3} \,\,{\rm CO}_{\rm 2} \left( {^3 B_2 } \right) $$\end{document} as 10-14.0±0.3 exp[-(1630 ± 325)/T] cm3 molecule-1 s-1. The implications of this are discussed.

Notes: The thermal decomposition of azoisopropane (AIP) was studied by detailed product analysis in the temperature and pressure intervals 498-563 K and 0.67-5.33 kPa. Besides the predominant termination and hydrogen-abstraction reaction of the 2-propyl radical, the decomposition is characterized by a very short chain process. The following rate constants were determined from the measurements\documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{rcl} \log (k_1 {\rm /s}^{ - 1} ) &=& (16.3 \pm 0.2) - (199.9 \pm 1.6){\rm kJ mol}^{ - 1} /2.3RT \\ \log (k_4 /k_3^{1/2} {\rm dm}^{{3 \mathord{\left/ {\vphantom {3 2}} \right. \kern-\nulldelimiterspace} 2}} {\rm mol}^{{{ - 1} \mathord{\left/ {\vphantom {{ - 1} 2}} \right. \kern-\nulldelimiterspace} 2}} {\rm s}^{{{ - 1} \mathord{\left/ {\vphantom {{ - 1} 2}} \right. \kern-\nulldelimiterspace} 2}} ) &=& (4.1 \pm 0.3) - (52.5 \pm 3.0){\rm kJ mol}^{ - 1} /2.3RT \\ \log (k_5 /k_3^{1/2} {\rm dm}^{{3 \mathord{\left/ {\vphantom {3 2}} \right. \kern-\nulldelimiterspace} 2}} {\rm mol}^{{{ - 1} \mathord{\left/ {\vphantom {{ - 1} 2}} \right. \kern-\nulldelimiterspace} 2}} {\rm s}^{{{ - 1} \mathord{\left/ {\vphantom {{ - 1} 2}} \right. \kern-\nulldelimiterspace} 2}} ) &=& (2.4 \pm 0.1) - (27.6 \pm 1.3){\rm kJ mol}^{ - 1} /2.3RT \\ k_2 /k_3 &=& 0.51 \pm 0.02 \end{array}$$\end{document} for the following reactions:

Notes: Chlorocyclopropane has been produced by addition of CH2(1A1) and CH2(3B1) to chloroethene. CH2 was generated by the photolysis of ketene at 313 and 366 nm. Chlorocyclopropane was formed in a chemically activated state, had an energy content between 378 and 427 kJ/mol, and reacted in three parallel channels to 3-chloropropene, cis- and trans-1-chloropropene. As secondary reactions elimination of HCl from the chemically activated primary products occurred to form allene and propyne. The apparent rate constants for the isomerization and elimination reactions are reported. The results of RRKM calculations including distribution functions for the activated chlorocyclopropane and a stepladder model for the deactivation support the proposed reaction scheme.

Notes: The kinetics of OH reactions with furan (k1), thiophene (k2), and tetrahydrothiophene (k3), have been investigated over the temperature range 254-425 K. OH radicals were produced by flash photolysis of water vapor at λ 〉 165 nm and detected by timeresolved resonance fluorescence spectroscopy. The following Arrhenius expressions adequately describe the measured rate constants as a function of temperature (units are cm3 molecule-1 S-1): k1 = (1.33 ± 0.29) × 10-11 exp[(333 ± 67)/T], k2 = (3.20 ± 0.70) × 10-12 exp[(325 ± 71)/T], k3 = (1.13 ± 0.35) × 10-11 exp[(166 ± 97)/T]. The results are compared with previous investigations and their implications regarding reaction mechanisms and atmospheric residence times are discussed.

Notes: The reactions of ground-state S(3PJ) atoms with thiirane, methylthiirane, and trans-2,3-dimethylthiirane have been studied by flash photolysis-VUV kinetic absorption spectroscopy. From the analysis of the S(3PJ) decay plots the following rate constants were determined: (1.4 ± 0.2) × 1013, (2.7 ± 0.3) × 1013 and (4.0 ± 0.2) × 1013 (in cm3 mol-1 s-1 units) for thiirane, methylthiirane and trans-2,3-dimethylthiirane, respectively, showing an upward trend with increasing methylation.

Notes: NO Abstract.

Notes: An examination of the results of measurements of the forward and reverse rate constants for the reaction \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm H} + {\rm C}_{\rm 2} {\rm H}_{\rm 6} \mathbin{\lower.3ex\hbox{$\buildrel\textstyle\rightarrow\over {\smash{\leftarrow}\vphantom{_{\vbox to.5ex{\vss}}}}$}} {\rm H}_{\rm 2} + {\rm C}_{\rm 2} {\rm H}_{\rm 5} $$\end{document} shows that agreement between the kinetics and the thermochemistry is achieved only through use of a value of ΔHf(C2H5) = 28 kcal mol-1. This system therefore provides further support for the recent measurement of this quantity.

Notes: The very low pressure reactor (VLPR) technique has been used to measure the bimolecular rate constant of the title reaction at 300 K. The rate constant is given by log k1 (1/mol s) = (11.6 ± 0.4) - (5.9 ± 0.6)/θ the equilibrium constant has also been measured at the same temperature and is given by K1 = (5.6 ± 1) × 10-3 and hence log k-1 (1/mol s) = 9.5 ± 0.1. The results show that the reaction Br + t—C4H9 → HBr + i—C4H8 is unimportant under the present experimental conditions. Assigning the entropy of t-butyl radical to be 74 ± 2 eu which is in the possible range, the value of K1 gives ΔH°f (t-butyl) = 9.1 ± 0.6 kcal/mol-1. This yields for the bond dissociation, DH° (t-butyl-H) = 93.4 ± 0.6 kcal/mol. Both of these values are found to be in good agreement with recent VLPP studies.

Notes: Kinetics of the basic hydrolysis of glyceryl trinitrate (TNG) were investigated in CO2-free aqueous calcium hydroxide solutions. The hydrolysis reactions were carried out in a temperature controlled reactor vessel with provision for continuous N2 sparging of the reaction mixture. TNG hydrolyzed via second-order reaction at 25°C, 18°C, and 10°C. The activation energy of the hydrolysis reaction of TNG was calculated from the kinetic data and found to be equal to 27.53 kcal/mol. The major products of the hydrolysis of TNG in solution of calcium hydroxide were calcium nitrate and calcium nitrite, accounting for approximately 50% of the degradation products. The minor identified products such as calcium oxalate and nitrate esters amounted to approximatey 6% of the products. The remaining 30% of the isolated products was a mixture of calcium formate, a nitrate ester, and unidentified volatiles, polymerlike substances, and other organic residue.

Notes: Results obtained from the photolysis of ketene with acetylene strongly support the formation of C3H3 radicals in the title reaction. Stationary state studies are interpreted in terms of the reaction \documentclass{article}\pagestyle{empty}\begin{document}$${\rm C}_3 {\rm H}_{\rm 4}^{\rm *} \buildrel3\over\rightarrow{\rm C}_3 {\rm H}_3^ \cdot + {\rm H}^ \cdot$$\end{document} with a rate constant (109.8 s-1) which is compared to RRKM predictions. In pulsed laser induced decomposition experiments, recombination products involving C3H3 have been detected (some for the first time) and their formation modeled using step (3) with the same rate constant.

Notes: A readily applicable empirical formula is obtained for the collisional efficiency for energy transfer between a highly vibrationally excited reactant and a seasoned (usually quartz) wall, in terms of the molecular weight, potential well depth and dipole moment of the reactant. This expression is used to examine corrections due to nonunit wall collision efficiency in the high-pressure rate parameters obtained from very low-pressure pyrolysis experiments. It is found that these corrections are up to ca. ±5 kJ/mol in the high-pressure activation energy and a factor of ca. 2 in the high-pressure frequency factor, for molecules with molecular weight less than ca. 100 and where experiments are carried out at temperatures exceeding 1000 K.

Notes: The rate constants for reactions (4) and (5) were determined at room temperature by pulsed laser photolysis and time resolved mass spectrometry. A description of the experimental setup is given. CFCl2O2 radicals were generated by photolysis of CFCl3 at 193 nm in the presence of an excess of oxygen, using an excimer laser. The rate constant for reaction (4), determined under different experimental conditions is: \documentclass{article}\pagestyle{empty}\begin{document}$$k_4 = 1.6{\rm }(\pm 0.2) \times 10^{ - 11} {\rm cm}^{\rm 3} \cdot {\rm molecule}^{ - 1} \cdot {\rm s}^{ - 1}$$\end{document} The rate constant of reaction (5) was determined in the pressure range of 1-12 torr, using oxygen as the buffer gas. The reaction is in its fall-off region and the parameters determined by using the semiempirical method of Troe, taking Fc = 0.6 are: \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{l} k(0) = 3.5{\rm }(\pm 0.5) \times 10^{ - 29} {\rm cm}^{\rm 6} \cdot {\rm molecule}^{ - 2} \cdot {\rm s}^{ - 1} \\ k(\infty) = 6.0{\rm }(\pm 1.0) \times 10^{ - 12} {\rm cm}^{\rm 3} \cdot {\rm molecule}^{ - 1} \cdot {\rm s}^{ - 1} \\ \end{array}$$\end{document} The value of k(∞) is obtained from the low-pressure measurements and therefore the uncertainty on the actual high-pressure limit is higher than the error limits quoted above. The results are compared with those reported for similar reactions.

Notes: Potassium persulfate oxidizes triphenylphosphine to triphenylphosphine oxide in 60% aqueous acetonitrile. It has been suggested that the oxygen of the product, triphenylphosphine oxide, might originate from solvent water, following nucleophilic attack on an intermediate phosphonium ion. We have investigated the origin of the oxygen in the oxidation of triphenylphosphine by potassium persulfate in 60% aqueous acetonitrile containing 20% [18O]water. The product was analyzed by using the 18O isotope effect in 31P NMR spectroscopy. The magnitude of the 18O isotope-induced shift was determined by synthesizing triphenylphosphine [18O]oxide and was found to be 0.038 ppm upfield. The product of the oxidation reaction in 20% [18O]water displayed no 18O isotope effect. The origin of the oxygen in the oxidation reaction is the persulfate ion, consistent with an alternative mechanism involving nucleophilic attack by water at the sulfur atom of a phosphonium peroxysulfate intermediate.

Notes: It is shown how kinetic electron spin resonance spectroscopy with intermittent radical generation can be used to obtain rate constants of various simultaneous reactions in systems containing more than one kind of transient radicals. The technique is applied to reactions of tert-butyl [(CH3)3Ċ] and isopropylol [(CH3)2ĊOH] radicals generated by photolysis of di-tert-butyl ketone and acetone in 2-propanol/acetone mixtures. It yields the rates of generation of the two radicals, the rate constants for their self- and crossterminations and for the reaction of tert-butyl with 2-propanol. The extent of diffusion control of the termination constants is discussed.

Notes: Trifluoro-t-butoxy radicals have been generated by reacting fluorine with 2-trifluoromethyl propan-2-ol: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm \dot F} + {\rm CF}_3 {\rm C}({\rm OH})({\rm CH}_3)_2 \to {\rm HF} + {\rm CF}_3 {\rm C}({\rm \dot O})({\rm CH}_3)_2 $$\end{document} Over the temperature range 361-600 K the trifluoro-t-butoxy radical decomposes exclusively by loss of the —CF3 group [reaction (-2)] rather than by loss of —CH3 group [reaction (-1)]: The limits of detectability of the product CF3COCH3, by gas-chromatographic analysis, place a lower limit on the ratio k-2/k-1 of ca. 75. The implications of these results in relation to the reverse radical addition reactions to the carbonyl group are discussed along with the thermochemistry of the reactions.

Notes: The gas-phase reaction of CF3 with HCN has been examined over a wide conversion range using CF3I as a thermal and photolytic source of radicals. Quantitative and qualitative results show a significant increase of the specific rate constant for the hydrogen abstraction reaction relative to CF3 recombination when reaction is carried out under ultraviolet irradiation. This “extra” formation of the reaction product, CF3H, has been assigned to the participation of iodine in this system through the formation of a (I-HCN) intermediate. Arrhenius parameters obtained for the addition mechanism of I to HCN do not seem to conform to a single reaction step, on the contrary, they correspond to a more complex reaction scheme.

Notes: Small low residence time flow tube reactors made of alumina and used as molecular beam sources are described. In these reactors, gas mixtures are rapidly heated and brought to reaction. The composition of the gas leaving the reactor is analyzed by molecular beam mass spectroscopy. For quantitative simulation of the reacting gas flow, the theory of one-dimensional compressible flow with friction, heat transfer, and chemical reaction is brought into a form suitable for practical computation. The system has been applied to study the thermal decompositions of O3 and N2O. The experimental results on both reactions can be well modeled by homogeneous reaction mechanisms with accepted rate constants. Heterogeneous reaction steps are shown to be unimportant.

Notes: The yield of benzene in the reaction of 1,4- and 1,3-cyclohexadiene with OH radicals in the presence of oxygen was determined using H2O2 and CH3ONO as OH radical sources. Both in the H2O2 and the CH3ONO systems, the yield of benzene from 1,4-cyclohexadiene was 15.3% and the yield from 1,3-cyclohexadiene was 8.9%. On the basis of the obtained yields, the rate constant for allylic hydrogen abstraction per C—H in cyclohexadiene was determined to be 3.8 × 10-12 cm3 molecule-1 s-1. The branching ratio of the hydrogen abstraction to overall reaction for 1-butene and 1-pentene was estimated to be (25-14)% by applying the obtained rate constants. The result was in good agreement with the branching ratio determined directly by use of the discharge flow photoionization mass spectrometer by Biermann, Harris, and Pitts [4].

Notes: The photolysis of carbon tetrachloride in the presence of a number of organosilicon compounds has been investigated in the gas phase. The products obtained from the photolysis experiments were those expected from a chain reaction in which trichloromethyl radicals abstract hydrogen atoms from the organosilane. Arrhenius parameters for hydrogen atom transfer were determined relative to those for trichloromethyl radical combination. The activation energies for the reaction of methyl, trifluoromethyl, and trichloromethyl radicals with organosilicon compounds are compared and the results rationalized in terms of polar effects.

Notes: Calculations of low-pressure limit, third-order rate constants are presented for the association reactions A + O2 + N2 and A + OH + N2 (A = Li, Na, K) over the temperature range 200-2000 K and a comparison is made with the available experimental data.

Notes: The photooxidation of formaldehyde in CH2O—O2, oxygen-lean mixtures was studied in the temperature range 298-378 K. H2 and CO formation and the loss of O2 proceed by a chain mechanism, which between 328 and 378 K follows the previously suggested kinetics [1] with one modification. The reaction HO2 + CH2O ⇄ HO2CH2O (5) is now assumed to be reversible and ΔH5° is estimated to be between 14 and 19 kcal/mol. The relative yields of the chain formed H2 and CO and of the consumed O2 remained constant over the entire temperature range indicating that the relative efficiencies of the HO reactions: HO + CH2O → H2O HCO† (7), HO + CH2O → H2O + HCO (8) and HO + CH2O → HOCH2O (9) are temperature independent.

Notes: The kinetics of OH reactions with 1-4 carbon aliphatic thiols have been investigated over the temperature range 252-430 K. OH radicals were produced by flash photolysis of water vapor at λ 〉 165 nm and detected by time-resolved resonance fluorescence spectroscopy. All thiols investigated react with OH at nearly the same rate; k(298 K) = 3.2-4.6 × 10-11 cm3 molecule-1 s-1, -Eact = 0.6-1.0 kcal/mol, A = 0.6-1.2 × 10-11 cm3 molecule-1 s-1. CH3SH and CH3SD react with OH at identical rates over the entire temperature range investigated. We conclude that the dominant reaction pathway is addition to the sulfur atom.

Notes: Mixtures of N2O, H2, O2, and trace amounts of NO and NO2 were photolyzed at 213.9 nm, at 245°-328°K, and at about 1 atm total pressure (mostly H2). HO2 radicals are produced from the photolysis and they react as follows:Reaction (1b) is unimportant under all of our reaction conditions. Reaction (1a) was studied in competition with reaction (3) from which it was found that k1a/k31/2 = 6.4 × 10-6 exp { z-(1400 ± 500)/RT} cm3/2/sec1/2. If k3 is taken to be 3.3 × 10-12 cm3/sec independent of temperature, k1a = 1.2 × 10-11 exp {-(1400 ± 500)/RT} cm3/sec. Reaction (2a) is negligible compared to reaction (2b) under all of our reaction conditions. The ratio k2b/k1 = 0.61 ± 0.15 at 245°K. Using the Arrhenius expression for k1a given above leads to k2b = 4.2 × 10-13 cm3/sec, which is assumed to be independent of temperature. The intermediate HO2NO2 is unstable and induces the dark oxidation of NO through reaction (-2b), which was found to have a rate coefficient k-2b = 6 × 1017 exp {-26,000/RT} sec-1 based on the value of k1a given above. The intermediate can also decompose via Reaction (10b) is at least partially heterogeneous.

Notes: Tertiary-amyl amine has been decomposed in single-pulse shock-tube experiments. Rate expressions for several of the important primary steps are \documentclass{article}\pagestyle{empty}\begin{document}$$ k(t{\rm C}_5 {\rm H}_{11 - {\rm NH}_2} \to t{\rm C}_5 {\rm H}_{11}\!\!\cdot + {\rm NH}_2\cdot) = 10^{15.9} \exp (-39,700/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm C}_2 {\rm H}_5- {\rm C}({\rm CH}_3)_2{\rm NH}_2 \to {\rm C}_2 {\rm H}_5 \cdot + \cdot{\rm C}({\rm CH}_3)_2{\rm NH}_2) = 10^{16.5} \exp (- 38,500/T)\sec ^{- 1} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ k(t{\rm C}_5 {\rm H}_{11} {\rm NH}_2 \to {\rm C}_5 {\rm H}_{10} + {\rm NH}_3) 〈10^{14.5} \exp (- 37,200/T)\sec ^{- 1} $$\end{document}This leads to D(CH3—H) - D(NH2—H) = -10.5 kJ and D[(CH3)3C—H] - D[(CH3)2NH2C—H] = + 6 kJ.The present and earlier comparative rate single-pulse shock-tube data when combined with high-pressure hydrazine decomposition results-(after correcting for fall off effects through RRKM calculations) gives \documentclass{article}\pagestyle{empty}\begin{document}$$ [k_r^2 (t{\rm C}_5 {\rm H}_{11} \cdot,{\rm NH}_2 \cdot)/k_r (t{\rm C}_5 {\rm H}_{11} \cdot,t{\rm C}_5 {\rm H}_{11} \cdot)k_r ({\rm NH}_2 \cdot,{\rm NH}_2 \cdot)]^{1/2} \sim 2\,{\rm at}\,1100^o {\rm K} $$\end{document} where kr(…) is the recombination rate involving the appropriate radicals. This suggests that in this context amino radical behavior is analogous to that of alkyl radicals. If this agreement is exact, then \documentclass{article}\pagestyle{empty}\begin{document}$$ k_\infty ({\rm N}_2 {\rm H}_4 \to 2{\rm NH}_2 \cdot) = 10^{16.25} \exp (- 32,300/T)\sec ^{- 1} $$\end{document} Rate expressions for the primary step in the decomposition of a variety of primary amines have been computed. In the case of benzyl amine where data exist the agreement is satisfactory. The following differences in bond energies have been estimated: \documentclass{article}\pagestyle{empty}\begin{document}$$ D(i{\rm C}_3 {\rm H}_7 {-\!-} {\rm H}) {-\!-} D[{\rm CH}_3 ({\rm NH}_2){\rm CH} {-\!-} {\rm H}] = 14.3\,{\rm kJ} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ D({\rm C}_2 {\rm H}_5 {-\!-} {\rm H}) - D({\rm NH}_2 {\rm CH}_2 {-\!-} {\rm H}) = 15.9\,{\rm kJ} $$\end{document}

Notes: The thermal decomposition of F5SOOSF5, P, in the presence of CO has been investigated between 130.1° and 161.9°C at total pressures between 50 and 600 torr. The reaction is homogeneous, and the only final products formed are CO2 and S2F10. The rate of reaction is proportional to the pressure of P. The partial pressures of CO and O2 and the total pressure have no influence on the course of reaction: \documentclass{article}\pagestyle{empty}\begin{document}$$ - \frac{{d\left[P \right]}}{{dt}} = k\left[P \right] $$\end{document}The results are explained by the following mechanism:

Notes: The unimolecularity of the thermal dehydrogenation of cyclopentene has been confirmed using the technique of very low-pressure pyrolysis (VLPP). Application of RRKM theory shows that the experimental unimolecular rate constants obtained over the temperature range of 942°-1152°K are consistent with the high-pressure Arrhenius parameters given by\documentclass{article}\pagestyle{empty}\begin{document}$$ \log (k/\sec ^{-1}) = 13.35 - 61/\theta $$\end{document} where θ = 2.303 RT kcal/mol. These parameters are in good agreement with static and shock tube studies. No firm evidence could be found for any side reactions or reversibility under the experimental conditions used.

Notes: The thermal decomposition of isopropyl peroxide in solvents is shownto be a competition between unimolecular homolysis (kr) and an electrocyclic reaction yielding H2 plus acetone (kH). The value of kH increases markedly with increasing solvent polarity: PhCH3 〈 i-PrOH 〈 MeOH 〈 H2O. Log kH correlates linearly with Et (or Z). The effect on kr is similar but less pronounced. Log kr's for both i-Pr2O2 and t-Bu2O2 are shown to correlate with the molar solubility parameter δ. The negligible contribution of the electrocyclic reaction to thermal decompositions of peroxides in the vapor phase is thus not due to peculiar behavior of that reaction, which clearly has a polartransition state. It is argued that the transition state for unimolecular homolysis has no polar character despite the apparent sensitivity of homolysis rates to solvent polarity.

Notes: The kinetics of the gas-phase decomposition of 1-methylcyclohex-1-ene has been studied over the temperature range of 1000°-1180° K in a single-pulse shock tube using a comparative-rate technique. The retro Diels-Alder reaction to ethylene and isoprene accounts for the bulk of the products with a rate constant given by \documentclass{article}\pagestyle{empty}\begin{document}$$ k = 10^{15.57} \exp \left({- 35,000/T} \right)\sec ^{- 1} $$\end{document}

Notes: The decomposition of acetonyl bromide, isopropenylmethylether, and hexanedione-2,5 was studied using the very-low-pressure pyrolysis (VLPP) technique. The acetonyl radical is a product of each reaction. Arrhenius parameters determined are or acetonyl bromide ← acetonyl + Br: \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k\left({\sec ^{- 1}} \right) = 16.0 - 62.5/\theta\,{\rm at}\,300^{\rm o} {\rm K} $$\end{document} and for isopropenylmethylether ← acetonyl + CH3: \documentclass{article}\pagestyle{empty}\begin{document}$$ \log k\left({\sec ^{- 1}} \right) = 15.8 - 66.3/\theta\,{\rm at}\,300^{\rm o} {\rm K} $$\end{document} These lead to values of acetonyl stabilization energy (SE) of 0.8 and -4.0 kcal/mol, respectively. Comparison of the pyrolyses of hexanedione-2,5 and 2,5-dimethylhexane indicate a value of SE ∼ 2 kcal/mol. The total of these results is taken, along with previous work, to indicate that 0 ≲ SE ≲ 2 kcal/mol.

Notes: Relative Arrhenius parameters have been determined for the addition of trifluoromethyl radicals to vinyl chloride (X=Cl), vinyl bromide (X=Br), acrylonitrile (X=CN), methyl vinyl ketone (X=COCH3), 2-methylbut-2-ene, and to iso-butene, taking the rate of addition to ethylene as standard. Addition occurred virtually exclusively at the unsubstituted end of all the olefins. The relative rates of addition at 164°C show very little variation. The preexponential terms for addition to the CH2 ends are constant within experimental error, and the small variations in rate are due principally to the activation energy term. The present results are put on an absolute scale using the previously determined Arrhenius parameters for addition to ethylene and are compared with data for other similar monomers determined previously.

Notes: The kinetics of the deprotonation of tropaeolin O by OH- ions was investigated between 9° and 30°C, and by OD- ions at 24.7°C. The pH range was 10.7-12.5, and the ionic strength 0.1M throughout. All results were obtained by the temperature jump method. On the basis of a mechanism suggested earlier, rate constants k31 for the reaction between OL- and the internally bonded weak acid and k32 for the opening of the internal hydrogen bond were evaluated. The activation energies in ordinary water were found to be ΔH≠31 = 3.6 kcal/mol, ΔS≠31 = -19 eu, and ΔH≠32 = 27 kcal/mol, ΔS≠32 = 46 eu. The kinetic isotope effect was k31H2O/k31D2O ∼ 1.5 and k32H2O/k32D2O ∼ 0.9. The unusual results for reaction path are discussed in terms of solvent participation.

Notes: The autoinhibiting reaction of ozone with dimethyl sulfide (DMS), has been studied at 296°K and 1.1 kPa (8 torr) as a function of the concentrations of both reactants. The major products of the reaction are H2CO, H2O, CO, and SO2. The specific rate of primary attack of O3 on DMS is immeasurably slow. It is suggested that the rapid overall rate observed for this reaction is due to a chain reaction initiated by the very slow primary reaction. It is concluded that reaction (1) cannot be important under atmospheric conditions and that the major loss process for DMS in the atmosphere is probably reaction with photochemically generated free radicals.

Notes: Methyl radical reactions with matrix molecules in glasses C2H5OH, (CH2OH)2, n- and i-C3H7OH, n- and i-C4H9OH, n- and i-C5H11OH, C2D5OH, and i-C3D7OD, and the reactions of Ċ2H5, Ċ3H7, Ċ4H9, Ċ5H11 with methanol glasses have been studied. Alkyl radicals were produced by photolysis of diphenylamine-alkylhalide-alcohol mixtures using ultraviolet light. In all cases the alkyl radical decay follows the law c = c0 exp(-k √t). The √t law should not be associated with alkyl radical diffusion in a matrix. A method of processing the kinetics of those reactions in which one paramagnetic species changes into another with the total concentration being constant and the electron spin resonance spectra of both species overlapping, is described.

Notes: A kinetic study has been made of the 3130-Å photolysis of CH2O (8 torr) in O2-containing mixtures (0.02-8 torr) and in the presence of added CO2 (0-300 torr) at 25°C. Quantum yields of formation of H2, CO, and CO2 and the loss of O2 were measured. ΦH2 and ΦCO were much above unity. In an explanation of these unexpected results, a new H-atom-forming chain mechanism was postulated involving HO2 and HO addition to CH2O: CH2O + hν → H + HCO (1) H + CH2O → H2 + HCO (3) H + O2 + M → HO2 + M (6) HCO + O2 → HO2 + CO (8) HO2 + CH2O → (HO2CH2O) → HO + HCO2H (15) HO + CH2O → H2O + HCO† (16); HCO† → H + CO (19) HO + CH2O → H2O + HCO (17) and HO + CH2O → HCO2H + H (18). When the results are rationalized in terms of this mechanism, the data suggest k16 ≫ k17 and k16/k18 ≅ 0.5. The data require that a reassessment of the relative rates of reactions (7) and (8) be made, since in the previous work HCO2H formation was used as a monitor of the rate of reaction (7) HCO + O2 + M → HCOO2 + M (7). The present data from experiments at PCH2O = 8 torr and PO2 = 1-4 torr give k7[M]/(k7[M] + k8) ≥ 0.049 ± 0.017. These data coupled with the k8 estimates of Washida and coworkers give k7 ≥ (4.4 ± 1.6) × 1011 l2/mol2·sec for M = CH2O. The reaction sequence proposed here is consistent with the observed deterimental effect of O2 addition on the laser-induced isotope enrichment in HDCO. In additional studies of CH2O-O2-isobutene mixtures it was found that ΦH2 was equal to φ2 as estimated in O2-free CH2O-isobutene mixtures. These results suggest that the increase in CO (ν = 1) product observed with O2 addition in CH2O photolysis does not result from perturbations in the fragmentation pattern of the excited CH2O, but it is likely that it originates in the occurrence of the exothermic reaction HCO + O2 → HO2 + CO (ν = 1).

Notes: Using the ECR technique the electron scavenging properties of the halomethanes CHF3, CHClF2, and CHCl2F were investigated. A brief discussion is given on the evaluation of experimental data for inefficient polar scavengers. Absolute attachment rate constants were obtained from the decrease of the cyclotron resonance signal upon addition of the above scavengers. For the molecules CHF3 and CHCl2F the activation energies for the capture processes are given. Possible capture mechanisms are discussed.

Notes: A nitrogen laser pumped tunable dye laser has been used to observe the three-photon ionization of NO through a two-photon resonance with the C2II state. Fluorescence is also observed from this state. The wavelength dependence of both signals have been measured. A reaction mechanism is postulated, which includes the initial two-photon excitation of the C2II state as the rate-limiting step. This mechanism predicts the observed second-order intensity dependence of the ionization signal and shows that the simple rate equation treatment is valid in this system.

Notes: It is shown that the activation energies E oF chlorine atom abstraction by cyclohexyl radicals and hydrogen atom abstraction by Cl atoms from polychloroalkanes can be correlated with the bond dissociation energies D and the Taft polar and steric substituent constants σ* and Es by the expression: \documentclass{article}\pagestyle{empty}\begin{document}$$ \Delta E = \alpha \Delta D + \rho \sigma ^* + \delta E_{\rm s} $$\end{document} where ΔE and ΔD represent the differences between the E and D values of a given substrate and those of a reference compound (CH3 substituted substrate) and α, ρ, and δ are the corresponding correlation coefficients. The use of this expression allows quantitative evaluation of the relative contribution of the various factors affecting the activation energies of these reactions and estimation of related thermochemical data.

Notes: The reactions of ozone with propene and isobutene have been studied in the gas phase at 298°K and 530 Pa (4 torr) using a stopped-flow reactor coupled to a photoionization mass spectrometer. Reactant and product concentrations were followed as a function of reaction time. The major reaction products monitored were CH2O, CH3CHO, CO2, and H2O from the propene reaction, and CH2O, (CH3)2CO, CO2, and H2O from the isobutene reaction. The observations were interpreted on the basis of the Criegee mechanism for ozonolysis in solution: for which we find kA≈kB.In the gas phase the carbene peroxy radical is postulated to isomerize and decompose into molecular and free-radical products.

Notes: Rate constants for the reaction O(3P) + SO2 + M have been determined over the temperature range of 299°-440°K, using a flash photolysis-NO2 chemiluminescence technique. For M=Ar, the Arrhenius expression \documentclass{article}\pagestyle{empty}\begin{document}$$ k_2 ^{{\rm Ar}} {\rm = 3}{\rm .1 \times 10}^{{\rm - 32}} e^{{\rm - (2005 \pm 300)/RT}} {\rm cm}^{\rm 6} {\rm /molec}^2 \cdot \sec $$\end{document} was obtained. At room temperature k2Ar = (1.05 ± 0.21) × 10-33 cm6/molec2·sec. In addition, the rate constants k2N2 = (1.37 + 0.27) × 10-33 cm6/molec2·sec, k2SO2 = (9.5 ± 3.0) ± 10-33 cm6/molec2·sec, k3N2 = (1.1 ± 0.2) ± 10-31 cm6/molec2·sec, and k3SO2 = (2.6 - 0.9) ± 10-31 cm6/molec2·sec were obtained at room temperature where k3M is the rate constant for the reaction O + NO + M → NO2 + M. The rate data are compared and discussed with literature values.

Notes: The reactions where Y = CH3 (M), C2H5 (E), i—C3H7 (I), and t—C4H9 (T) have been studied between 488 and 606 K. The pressures of CHD ranged from 16 to 124 torr and those of YE from 57 to 625 torr. These reactions are homogeneous and first order with respect to each reagent. The rate constants (in L/mol·s) are given by \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm NMBO}} = - {{\left( {26530 \pm 80} \right)} \mathord{\left/ {\vphantom {{\left( {26530 \pm 80} \right)} {4.576T + \left( {6.05 \pm 0.03} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {6.05 \pm 0.03} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm XMBO}} = - {{\left( {28910 \pm 130} \right)} \mathord{\left/ {\vphantom {{\left( {28910 \pm 130} \right)} {4.576T + \left( {6.32 \pm 0.05} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {6.32 \pm 0.05} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm NEBO}} = - {{\left( {26150 \pm 120} \right)} \mathord{\left/ {\vphantom {{\left( {26150 \pm 120} \right)} {4.576T + \left( {5.85 \pm 0.05} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {5.85 \pm 0.05} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm XEBO}} = - {{\left( {28560 \pm 120} \right)} \mathord{\left/ {\vphantom {{\left( {28560 \pm 120} \right)} {4.576T + \left( {6.07 \pm 0.05} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {6.07 \pm 0.05} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm NIBO}} = - {{\left( {26560 \pm 80} \right)} \mathord{\left/ {\vphantom {{\left( {26560 \pm 80} \right)} {4.576T + \left( {5.57 \pm 0.03} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {5.57 \pm 0.03} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm XIBO}} = - {{\left( {28350 \pm 100} \right)} \mathord{\left/ {\vphantom {{\left( {28350 \pm 100} \right)} {4.576T + \left( {5.47 \pm 0.04} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {5.47 \pm 0.04} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm NTBO}} = - {{\left( {28920 \pm 50} \right)} \mathord{\left/ {\vphantom {{\left( {28920 \pm 50} \right)} {4.576T + \left( {5.86 \pm 0.02} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {5.86 \pm 0.02} \right)}} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} k_{{\rm XTBO}} = - {{\left( {32890 \pm 120} \right)} \mathord{\left/ {\vphantom {{\left( {32890 \pm 120} \right)} {4.576T + \left( {6.19 \pm 0.05} \right)}}} \right. \kern-\nulldelimiterspace} {4.576T + \left( {6.19 \pm 0.05} \right)}} $$\end{document} The Arrhenius parameters are used as a test for a biradical mechanism and to discuss the endo selectivity of the reactions.

Notes: The technique of laser photolysis of alkyl and perfluoroalkyl iodides at 266 nm followed by time-resolved detection of the 1.3-μm emission from I*(2P1/2) has been used to measure the rate constants for deactivation of I* by CH3I, C2H5I, CF3I, and CH4. The recommended values are (2.76± 0.22) × 10-13, (2.85 ± 0.40) × 10-13, (3.5 ± 0.5) × 10-17, and (7.52 ± 0.12) × 10-14, respectively, in units of cm3 molecule-1 S-1.

Notes: The reactions of labeled N15NO+ with CO, NO, O2, 18O2, N2, NO2, and N2O have been investigated using a tandem ICR instrument. In each case the total rate coefficient, product distribution, and kinetic energy dependence were measured. The results indicate that very specific reaction mechanisms govern these reactions. This conclusion is suggested by the lack of isotopic scrambling in many cases and by the complete absence of energetically allowed products in almost all of the systems. The kinetic energy studies indicate that most of the reaction channels proceed through an intermediate complex at low energies and via a direct mechanism at higher kinetic energies. Such direct mechanisms include long range charge transfer and atom or ion transfer.

Notes: The reaction of CH4 + Cl2 produces predominantly CH3Cl + HCl, which above 1200 K goes to olefins, aromatics, and HCl. Results obtained in laboratory experiments and detailed modeling of the chlorine-catalyzed polymerization of methane at 1260 and 1310 K are presented. The reaction can be separated into two stages, the chlorination of methane and pyrolysis of methylchloride. The pyrolysis of CH3Cl formed C2H4 and C2H2 in increasing yields as the degree of conversion decreased and the excess of methane increased. Changes of temperature, pressure, or additions of HCl had little effect. In the absence of CH4 C2H4 and C2H2 are formed by the recombination of ĊH3 and ĊH2Cl radicals. With added CH4 recombination of ĊH3 forms C2H6, which dehydrogenates to C2H4 + H2. C2H4 in turn dehydrogenates to C2H2 + H2. While HCl, C, CH4, and H2 are the ultimate stable products, C2H4, C2H2, and C6H6 are produced as intermediates and appear to approach stationary concentrations in the system. Their secondary reactions can be described by radical reactions, which can lead to soot formation. ĊH3 - initiated polymerization of ethylene is negligible relative to the Ċ2H3 formation through H abstraction by Cl. The fastest reaction of Ċ2H3 is its decomposition to C2H2. About 20% of the consumption of C2H2 can be accounted for by the addition of Ċ2H3 to it with formation of the butadienyl radical. The addition of the latter to C2H2 is slow relative to its decomposition to vinylacetylene. Successive H abstraction by Cl from C4H4 leading to diacetylene has rates compatible with the experimental values. About 10% of Ċ4H5 abstracts H from HCl and forms butadiene. Successive additions of Ċ2H3 to butadiene and the products of addition can account for the formation of benzene, styrene, naphthalene, and higher polyaromatics. The following rate parameters have been derived on the basis of the experimentally measured reaction rates, the estimated frequency factors, and the currently available heat of formation of the Ċ2H3 radical (69 kcal/mol): \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {\mathop {{\rm C}_{\rm 2} }\limits^. {\rm H}_{\rm 3} \mathop {\longrightarrow}\limits_{\left( {\rm M} \right)}^{39} {\rm H}\,\, + \,\,{\rm C}_{\rm 2} {\rm H}_{\rm 2} } & {\log k\left( {1\,{\rm atm,}\,{\rm 1300}\,{\rm K}} \right)\, = \,5.2\, + \,0.3\,s^{ - 1} } \\ \end{array} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {{\rm C}_{\rm 2} {\rm H}_{\rm 4} \, + \,\mathop {{\rm C}_{\rm 2} }\limits^. \,\mathop {\longrightarrow}\limits^{17} \,\mathop {{\rm C}_{\rm 4} }\limits^. {\rm H}_{\rm 7} } \hfill & {E\, \ge \,2\, \pm \,2\,{{{\rm kcal}} \mathord{\left/ {\vphantom {{{\rm kcal}} {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}\,} \hfill \\ {\mathop {{\rm C}_{\rm 2} }\limits^. {\rm H}_{\rm 5} \, + \,{\rm C}_{\rm 6} {\rm H}_{\rm 6} \,\mathop {\longrightarrow}\limits^{40} \,\mathop {{\rm C}_{{\rm 12}} }\limits^. {\rm H}_{{\rm 11}} } \hfill & {E\, = \,11\, \pm \,2\,{{{\rm kcal}} \mathord{\left/ {\vphantom {{{\rm kcal}} {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}} \hfill \\ \end{array} $$\end{document}

Notes: The kinetics of fast elementary recombination of neutral ketyl radicals of benzophenone and its four derivatives (BPH⋅), the dismutation of benzophenone radical anions, the disproportionation between BPH⋅ and stable nitroxyl radicals, (), and the electron transfer have been investigated in both individual solvents and binary mixtures of different viscosities. Reaction (1) for unsubstituted BPH in water, water glycerol, and n-hexane is controlled by diffusion with 2k1 ≃ kdiff. In aliphatic alcohols and toluene, which form solvation complexes with BPH⋅, reaction (1) is diffusion-enhanced and activation-controlled, respectively, with 2k1 〈 kdiff. In a viscous solvent such as 1-propanol-glycerol mixture (100 ≲ η ≲ 450 cP) reaction (1) is diffusion-controlled. Reaction (2) in alkaline 1-propanol and alkaline 1-propanol-glycerol mixture is activation controlled. The rates of reactions (3) and (4) for benzophenone radicals and nitroxyl radicals of the imidazoline series decrease as the viscosity of the water-glycerol and 1-propanol-glycerol mixtures is increased. The reactions are molecular mobility limited; nevertheless, the numerical values of k3 (k4) are 2-6 times as small as the corresponding kdiff values due to the low steric factor of the reactions (therefore called pseudodiffusion-controlled reactions). The theoretical estimates of k3 (k4) are in good agreement with the experimental results. The elimination of spin forbiddance in the process of radical recombination in viscous solvents is discussed.

Notes: The approximations developed to determine the energy distribution function of molecules activated above energy decomposition threshold, from experimental data, have been tested. The approach involved the theoretical (RRKM) calculations of “pseudoexperimental” data for a variety of activated energy distributions. (Single or double Gaussian representations were used in all cases.) Subsequently the algorithms mentioned were applied in order to recuperate the original (i.e., input) energy distributions from these pseudoexperimental data. The results obtained provide strong evidence in favor of the validity of the algorithms and illustrate the necessary requirements for their applications. A trend toward lower accuracy as the energy distributions move to higher energies has been observed. Evidence of the influence of the distribution width is also reported. The origins of the approximation errors have been studied, and ways for further improvement are suggested.

Notes: The autooxidation of retinyl acetate and methyl retinoate was investigated in chlorobenzene at 45°C. The rates of thermal initiation in the retinyl acetate solutions were measured, and a value was determined of the rate constant for the reaction of oxygen with retinyl acetate (RH + O2 → R· + HO2·): kio = (1.3 ± 0.2) × 10-5 L/mol · s. The number of moles of oxygen absorbed per mole of polyene depends on the substrate concentration. A kinetic scheme for the methyl retinoate autooxidation was proposed which takes into account the isomerization of primary peroxy radicals, and the rate constants for different elementary reactions were estimated. The partial rate constant for “allylic” hydrogen abstraction from retinyl acetate was estimated to be ≥ 1.65 × 103 L/mol · s. A probable propagation sequence was proposed for the autooxidation of retinyl acetate.

Notes: Rate constants for the self- and cross-termination of the isopropylol radical [(CH3)2ĊOH] and its anion [(CH3)2ĊO-] in aqueous solution are determined by kinetic electron spin resonance. Whereas the self-termination of the neutral radical occurs close to the diffusion-controlled limit, the cross- and self-terminations involving the anion are slower and reflect effects of charge repulsion and steric constraints by solvation.

Notes: The unimolecular decomposition of methyl nitrite in the temperature range 680-955 K and pressure range 0.64 to 2.0 atm has been studied in shock-tube experiments employing real-time absorption of CW CO laser radiation by the NO product. Computer kinetic modeling using a set of 23 reactions shows that NO product is relatively unreactive. Its initial rate of production can be used to yield directly the unimolecular rate constant, which in the fall-off region, can be represented by the second-order rate coefficient in the Arrhenius form: \documentclass{article}\pagestyle{empty}\begin{document}$$k_1 = 10^{17.90 \pm 0.21} \exp (- 17200 \pm 400/T){\rm cm}^{\rm 3} {\rm mol}^{ - 1} {\rm s}^{ - 1}$$\end{document} A RRKM model calculation, assuming a loose CH3ONO≠ complex with two degrees of free internal rotation, gives good agreement with the experimental rate constants.

Notes: Using a relative rate technique, rate constants for the gas phase reactions of the OH radical with n-butane, n-hexane, and a series of alkenes and dialkenes, relative to that for propene, have been determined in one atmosphere of air at 295 ± 1 K. The rate constant ratios obtained were (propene = 1.00): ethene, 0.323 ± 0.014; 1-butene, 1.19 ± 0.06; 1-pentene, 1.19 ± 0.05; 1-hexene, 1.40 ± 0.04; 1-heptene, 1.51 ± 0.06; 3-methyl-1-butene, 1.21 ± 0.04; isobutene, 1.95 ± 0.09; cis-2-butene, 2.13 ± 0.05; trans-2-butene, 2.43 ± 0.05; 2-methyl-2-butene, 3.30 ± 0.13; 2,3-dimethyl-2-butene, 4.17 ± 0.18; propadiene, 0.367 ± 0.036; 1,3-butadiene, 2.53 ± 0.08; 2-methyl-1,3-butadiene, 3.81 ± 0.15; n-butane, 0.101 ± 0.012; and n-hexane, 0.198 ± 0.017. From a least-squares fit of these relative rate data to the most reliable literature absolute flash photolysis rate constants, these relative rate constants can be placed on an absolute basis using a rate constant for the reaction of OH radicals with propene of 2.63 × 10-11 cm3 molecule-1 s-1. The resulting rate constant data, together with previous relative rate data from these and other laboratories, lead to a self-consistent data set for the reactions of OH radicals with a large number of organics at room temperature.

Notes: The kinetics of the oxidation of lactic and atrolactic acids by ceric sulfate have been studied in the medium HClO4-Na2SO4-NaClO4 at 25.0°C and ionic strength 2.0 mol dm-3 over a wide range of organic substrate (HL), hydrogen and bisulfate ion concentrations. The redox reactions proceed significantly through three simultaneous paths involving intermediate complexes between the reactive cerium(IV) species and the organic substrate according to the following expression \documentclass{article}\pagestyle{empty}\begin{document}$$k_{{\rm obs}} = \frac{{(b[{\rm HSO}_4^ -] + c[{\rm HSO}_4^ -]^2 + [{\rm H}^ +]){\rm [HL]}}}{{\{ f_1 [{\rm HSO}_4^ -]^3 + d_1 [{\rm HSO}_4^ -] + e_1 [{\rm HSO}_4^ -]^2){\rm }[{\rm H}^ +]\} + A'[{\rm HL}]}}$$\end{document} where kobs indicates the observed pseudo-first-order rate constant, b and c are rate constants relative to that for the path associated with the term [H+] in the numerator, and A' is a quantity depending on the [H+] and [HSO4-] concentrations. Moreover, three equilibria involving cerium(IV) and HSO4- (or SO42-) ions are important from a kinetic point of view, the cumulative equilibrium constants being in the ratios β1: β2: β3 = d1: e1: f1. The present data are compared with those obtained previously for the cerium(IV) oxidation of glycolic acid and the substituent effects discussed.

Notes: A kinetic study of oxidation of hydroxylamine by bromate ion in acid sulfate solution using spectrophotometric and potentiometric methods is reported. Oxidation of hydroxylamine to nitrate is quantitative and followed competitive, consecutive, and auto catalytics steps characterized by induction periods. In the slow rate limiting step, hydroxylamine on reaction with HOBr (k1′) forms an intermediate I, which further reacts fast with second molecule of HOBr (k2′) giving nitrite. Nitrite reacts with HOBr (k3′) yielding the final product nitrate. Nitric acts as an autocatalyst also and its initial addition decreased the induction periods. In excess of hydrogen ion concentration all the reaction steps follow second-order kinetics. All the second-order rate constants are reported and the reaction mechanism is proposed.

Notes: Mixtures of NH3 and N2O dilute in Ar were heated behind incident shock waves in the temperature range 1750-2060 K. A cw ring dye laser, tuned to the center of an OH absorption line in the ultraviolet, was used to monitor OH concentration profiles by absorption spectroscopy. Infrared emission was used to follow N2O (at 4.5 μm) and NH3 (at 10.5 μm) concentration - time histories. The early-time NH3 and OH concentration profiles were sensitive to the rate constants of the reactionsleading to the following best-fit expressions for k2 and k3:k2 = 1013.34±0.3 exp(-4470/T) and k3 = 1013.91±0.2 exp(-4230/T) cm3 mol-1 s-1. The results of this study combined with previous low-temperature data suggest a significant non-Arrhenius behavior for both k2 and k3.

Notes: The thermal decomposition of azoisopropane (AIP) was studied in the presence of various quantities of propylene in the temperature and pressure intervals 498-553 K and 3.33-5.33 kPa. The inhibition functions relating to formation of the products were determined; these proved a good basis for interpretation of the formation of the secondary decompositon products of AIP. The experimental data support the conception that the βμ radical - radical reactionoccurs. The product of this is not stable; its decomposition is one of the sources of the secondary products. The ratio of the rate constants was determined for the following reactions:

Notes: The photooxidation of acrylonitrile, methacylonitrile, and allylcyanide in the presence of NO was studied in parts per million concentration using the long-path Fourier transform IR spectroscopic method. The stoichiometry of the OH radical initiated oxidation of methacrylonitrile was established as \documentclass{article}\pagestyle{empty}\begin{document}$ \left( {{\rm OH}} \right) + {\rm CH}_{\rm 2} = {\rm C}\left( {{\rm CH}_{\rm 3} } \right){\rm CN + 2NO + 2O}_{\rm 2} \mathop {\hbox to 20pt{\rightarrowfill}}\limits^{1.0} {\rm HCHO + CH}_{\rm 3} {\rm COCN + 2NO}_{{\rm 2}} + \left( {{\rm OH}} \right) $\end{document}. The yield of HCHO for acrylonitrile and allylcyanide was found to be ca. 100 and 80%, and the stoichiometric reactions were assessed to proceed, \documentclass{article}\pagestyle{empty}\begin{document}$ \left( {{\rm OH}} \right) + {\rm CH}_{\rm 2} = {\rm CHCN + 2NO + 2O}_{\rm 2} \mathop {\hbox to 20pt{\rightarrowfill}}\limits^{1.0} {\rm HCHO + HCOCN + 2NO}_{\rm 2} + \left( {{\rm OH}} \right) $\end{document} and \documentclass{article}\pagestyle{empty}\begin{document}$ \left( {{\rm OH}} \right) + {\rm CH}_{\rm 2} = {\rm CHCH}_{\rm 2} {\rm CN + 2NO + 2O}_{\rm 2} \mathop {\hbox to 20pt{\rightarrowfill}}\limits^{0.8} {\rm HCHO + HCOCH}{\rm 2} {\rm CN + 2NO}_{\rm 2} + \left( {{\rm OH}} \right) $\end{document}, respectively. These results revealed that the reaction mechanism for these unsaturated organic cyanides are analogous to that of olefins.

Notes: Rates and thermodynamic data have been obtained for the reversible self-termination reaction: \documentclass{article}\pagestyle{empty}\begin{document}$${\rm R}^ \cdot + {\rm R}^ \cdot \mathop{\buildrel\longleftarrow\over\longrightarrow}^{2k1}_{2k_{-1}}D $$\end{document} Involving aromatic 2-(4′dimethylaminophenyl)indandione-1,3-yl (I), 2-(4′diphenylaminophenyl)indandione-1,3-yl (II), and 2,6 di-tert-butyl-4-(β-phthalylvinyl)-phenoxyl (III) radicals in different solvents. The type of solvent does not tangibly affect the 2k1 of Radical(I), obviously due to a compensation effect. The log(2k1) versus solvent parameter ET(30) curves for the recombination of radicals (II) and (III) have been found to be V shaped, the minimum corresponding to chloroform. The intensive solvation of Radical (II) by chloroform converts the initially diffusion-controlled recombination of the radical into an activated reaction. The log (2k-1) of the dimer of Radical (I) has been found to be a linear function of the Kirkwood parameter (ε - 1)/(2ε + 1), the dissociation rate increasing with the dielectic constant of the solvent. The investigation revealed an isokinetic relationship for the decay of the dimer of Radical (I), an isokinetic temperature β = 408 K and isoequilibrium relationship for the reversible recombination of Radical (I) with β° = 651 K. For Radical (I) dimer decay In(2k-1) = const + 0.8 In K, where K is the equilibrium constant of this reversible reaction. The transition state of Radical (I) dimer dissociation reaction looks more like a pair of radicals than the initial dimer. The role of specific solvation in radical self-termination reactions is discussed.

Notes: The thermal decomposition of SO2 and of the primary dissociation product SO have been studied in shock waves by the uv absorption technique. The controversy about SO2 dissociation data from uv absorption signals was resolved and attributed to the extensive overlap of SO2 and SO uv absorption spectra. The derived rate coefficients are k1/[Ar] = 1015.6 exp(-420 kJmol-1/RT) cm3mol-1 s-1 (temperature range 3000-5000 K) for SO2 dissociation, and k3/[Ar] = 1014.6 exp(-448 kJmol-1/RT) cm3 mol-1 s-1 (temperature range 4000-6000 K) for SO dissociation. Anomalously high values of the apparent collision efficiencies βc in SO2 dissociation are attributed to marked contributions from excited electronic states.

Notes: 5-Methyl-hexanone-2, 3-methyl-pentanone-2, and hexanone-2 have been decomposed in comparative rate single pulse shock tube experiments. The mechanism of decomposition involves the breaking of carbon-carbon bonds as well as molecular processes involving 6-center complexes. The following rate expressions at 1100 K have been obtained: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{rcl} k(3{\rm - methyl - pentanone - 2} \to {\rm CH}_{\rm 3} {\rm \dot CO} + {\rm sC}_4 {\rm H}_9 .) = 10^{16.4} \exp (- 38,300/T)/{\rm s} \\ k(5{\rm - methyl - hexanone - 2} \to {\rm CH}_{\rm 3} {\rm \dot CO} + {\rm iC}_4 {\rm H}_9 .) = 10^{16.6} \exp (- 40,600/T)/{\rm s} \\ k(5{\rm - methyl - hexanone - 2} \to {\rm CH}_{\rm 3} {\rm COCH}_{\rm 3} + {\rm iC}_4 {\rm H}_8) = 10^{12.56} \exp (- 31,600/T)/{\rm s} \\ k({\rm hexanone - 2} \to {\rm CH}_{\rm 3} {\rm COCH}_{\rm 3} + {\rm C}_3 {\rm H}_6) = 10^{13.28} \exp (- 32,400/T)/{\rm s} \\ \end{array} $$\end{document} These results lead to ΔHf(CH3ĊO) = - 13.8 kJ and ΔHf(CH3COCH2·) = - 12.6 kJ at 300 K. They are compared with existing literature values and some generalizations are made with regard to the stability of carbonyl compounds.

Notes: Kinetics of the incorporation of mercury(II) ion in tetra (p-trimethylammoniumphenyl)porphine have been investigated in aqueous solution at 30.0°C and 0.2 M (NaNO3) ionic strength. The reaction was found to be first order each in mercury(II) and the porphyrin. The forward (formation) and the reverse (dissociation) rate constants were found to be 1.9 ± 0.2 × 103 M-1 s-1 and 7 ± 2 × 106 M-1 s-1, respectively. Kinetics of zinc(II) incorporation in tetra(p-trimethylammoniumphenyl)porphine catalyzed by mercury(II) were also investigated. This catalysis is explained in terms of steady-state formation of mono mercury(II) porphyrin followed by zinc(II) displacement of mercury(II) ion from the porphyrin. Such a mechanism also illustrates the importance of porphyrin core deformation to metal incorporation.

Notes: The absolute rate constant for the OH + HCl reaction has been measured from 240 to 295 K utilizing the techniques of laser/flash photolysis-resonance fluorescence. The HCl concentrations were monitored continuously by ultraviloet and infrared spectrophotometry. The results can be fit to the following Arrhenius expression: \documentclass{article}\pagestyle{empty}\begin{document}$$k_1 = (4.6{\rm } \pm {\rm }0.3){\rm } \times {\rm }10^{ - 12} \exp [- (500{\rm } \pm {\rm }60)/T{\rm cm}^3 /{\rm molecule} \cdot {\rm s}$$\end{document} The rate constant values obtained in this study are 20-30% larger than those recommended previously for modeling of stratospheric chemistry.

Notes: The reaction \documentclass{article}\pagestyle{empty}\begin{document}${\rm Br} + {\rm CH}_3 {\rm CHO}\buildrel1\over\rightarrow{\rm HBr} + {\rm CH}_3 {\rm CO}$\end{document} has been studied by VLPR at 300 K. We find k1 = 2.1 × 1012 cm3/mol s in excellent agreement with independent measurements from photolysis studies. Combining this value with known thermodynamic data gives k-1 = 1 × 1010 cm3/mol s. Observations of mass 42 expected from ketene suggest a rapid secondary reaction: \documentclass{article}\pagestyle{empty}\begin{document}$${\rm Br} + {\rm CH}_3 {\rm CO}\buildrel2\over\rightarrow[{\rm CH}_3 {\rm COBr}]^* \buildrel3\over\rightarrow{\rm HBr} + {\rm CH}_2 {\rm CO}$$\end{document} in which step 2 is shown to be rate limiting under VLPR conditions and k2 is estimated at 1012.6 cm3/mol s from recent theoretical models for radical recombination. It is also shown that 0 ≤ E1 ≤ 1.4 kcal/mol using theoretical models for calculation of A1 and is probably closer to the lower limit. Reaction -1 is negligible under conditions used.

Notes: The effect of pressure on the rate constant of the OH + CO reaction has been measured for Ar, N2, and SF6 over the pressure range 200-730 torr. All experiments were at room temperature. The method involved laser-induced fluorescence to measure steady-state OH concentrations in the 184.9 nm photolysis of H2O-CO mixtures in the three carrier gases, combined with supplementary measurements of the CO depletion in these same carrier gases in the presence and absence of competing reference reactants. The effect of O2 on the pressure effect was determined. A pressure enhancement of the rate constant was observed for N2 and SF6, but not for Ar, within an experimental error of about 10%. The pressure effect for N2 was somewhat lower than previous literature reports, being about 40% at 730 torr. For SF6 a factor of two enhancement was seen at 730 torr. In each case it was found that O2 had no effect on the pressure enhancement. The roles of the radical species HCO and HOCO were evaluated.