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Notes: The gas-phase dehydrogenation of cyclopentene to cyclopentadiene catalyzed by iodine in the range 178-283°C has been found to obey a rate law consistent with the slow rate-determining step, \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm I} + {\rm c} - {\rm C}_5 {\rm H}_8 \stackrel{4}{\rightarrow}{\rm HI} + {\rm c} - {\rm C}_5 {\rm H}_7 $\end{document}, log [k4/(1 mole-1 sec-1)] = 10.25 ± 0.08 - (12.26 ± 0.18)/θ, where θ = 2.303RT in kcal/mole. Surface effects are not important. This value of E4 leads to a value of DHf2980 = 82.3 ± 1 kcal/mole and ΔHf298 = 38.4 ± 1 kcal/mole. From difference in bond strengths in the alkane and the alkene, the allylic resonance stabilization in the cyclopentenyl radical is 12.6 ± 1.0 kcal/mole, in excellent agreement with the value for the butenyl radical.Arrhenius parameters for the other steps in the mechanism are evaluated. The low value of A4 (compared with A4 for cyclopentane) suggests a “tighter” transition state for H-atom abstraction from alkenes than from alkanes.

Notes: Reinvestigation of the gas phase thermal reaction of 1,1,2,2-tetramethylcyclopropane (699-759°K) gave for the unimolecular disappearance of reactant, k(TMC) = 1015.27-63.93/θ sec-1, in good agreement with the original results of Frey and Marshall. However, evidence for a high activation energy (E = 79 ± 5 kcal/mole), competitive unimolecular decomposition to 2,3-dimethyl-1 and -2-butenes was also obtained. It is proposed that the serious discrepancy noted [1] between the experimentally observed Arrhenius parameters for the overall reaction kinetics, and those predicted by transition state calculations assuming a biradical mechanism for the isomerization reactions (previously believed to be the only primary reaction mode) can be explained in terms of the increasing importance of the decomposition reactions at higher reaction temperatures.

Notes: Dilute mixtures of 4-methyl-l-pentyne have been pyrolyzed in a single-pulse shock tube. The decomposition process involves bond breaking: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm HC} \equiv {\rm C} - {\rm CH}_2 ({\rm i - C}_{\rm 3} {\rm H}_{\rm 7} )\stackrel{k_B}{\longrightarrow}{\rm HC} \equiv {\rm C} - {\rm CH}_2 \cdot ({\rm propynyl}) + {\rm i - C}_{\rm 3} {\rm H}_{\rm 7} \cdot $$\end{document} as well as a molecular reaction: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm HC}\equiv {\rm C} - {\rm CH}_{\rm 2} ({\rm \rm i - C}_{\rm 3} {\rm H}_{\rm 7} )\stackrel{k_M}{\rightarrow}{\rm C}_{\rm 3} {\rm H}_4 ({\rm allene}) + {\rm C}_3 {\rm H}_6 . $$\end{document} The rate parameters are: \documentclass{article}\pagestyle{empty}\begin{document}$$k_{\bf B} = 10^{15.56} \exp {\rm (} - 34,940/T){\rm (sec}^{ - {\rm 1}}) $$ $$\begin{array}{*{20}c} k_{\rm M} = 10^{13.1} \exp {\rm (} - {\rm 29,670/}T){\rm (sec}^{ - {\rm 1}})&{\rm 1100}^ \circ {\rm K, 1}{\rm .5} - 5{\rm atm}\end{array}$$\end{document} The heat of formation of propynyl radical is thus ΔHf300 = 338 kJ mol-1 (80.7 kcal mol-1)· This leads to a propynyl resonance energy of 40 kJ mol-1 (9.6 kcal mol-1).

Notes: O(1D), produced from the photolysis of N2O at 2139 Å, reacts with N2O in accord with: \documentclass{article}\pagestyle{empty}\begin{document} $$ \begin{array}{*{20}c} {(2)} & {{\rm O}(^1 D) + N_2 {\rm O} \to {\rm N}_2 + {\rm O}_2 } \\ \end{array} $$ \end{document}\documentclass{article}\pagestyle{empty}\begin{document} $$ \begin{array}{*{20}c} {(3)} & { \to 2{\rm NO}} \\ \end{array} $$ \end{document}We have used the method of chemical difference to obtain an accurate measure of k2/k3 = 0.59 ± 0.01. Furthermore, the quantum yield of production of O(3P), either on direct photolysis or on deactivation of O(1D) by N2O, is less than 0.02 and probably zero.

Notes: A method for constructing potential energy surfaces previously proposed by the author has been extended to hydrogen transfer reactions between halide, oxygen, and carbon atoms. A qualitative relation was found between the repulsive energy and the number of anti-bonding electrons. In general, the calculated kinetic isotope effect is in satisfactory agreement with observed values and the contributions by H-atom tunneling to the rate of reaction is smaller than that obtained from other surfaces.

Notes: The kinetics of the nitric oxide catalyzed, homogeneous, gas-phase isomerization of 1,trans-3,trans-5-heptatriene have been studied for temperatures ranging between 130°C and 241°C. The very clean reaction involves exclusive geometrical isomerization about the 5,6-π-bond. The observed rate constants for \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm NO} + trans - {\rm 3,}trans{\rm - 5}\stackrel{1}{\rightarrow}trans - 3,cis - 5 + {\rm NO} $\end{document} can be represented (with standard errors) by log k1 = (7.18 ± 0.06) - (16.75 ± 0.12)/θ, where θ = 2.303 RT in kcal/mole. The consecutive-step reaction mechanism involves addition of NO to the double bond (Ka, b = ka/kb), followed by rotation of the 5,6-C—C bond in the adduct radical (kc.)Analysis of the observed activation parameters shows, that kc is rate-controlling and consequently k1 = kcKa, b. Estimates of kc and Ka, b lead to a value of k1 in good agreement with experiment.Comparing our data with those previously obtained for the similar 1,3-pentadiene system results in a value for the extra stabilization energy generated in the 1,3-heptadienyl radical of 18.5 ± 1.7 kcal/mole. This value is discussed in view of comparable data in the literature.

Notes: Arrhenius parameters have been measured for the abstraction of hydrogen from the C Si, Ge, and Sn tetramethyls: \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {k({\rm ml m}^{ - {\rm 1}} {\rm s}^{ - {\rm 1}} )} \hfill & {{\rm for}} \hfill & {{\rm C}({\rm CH}_{\rm 3} )_4 = {{10^{12.0} {\rm e}^{ - 8370} } \mathord{\left/ {\vphantom {{10^{12.0} {\rm e}^{ - 8370} } {RT}}} \right. \kern-\nulldelimiterspace} {RT}}} \hfill \\ {k({\rm ml m}^{ - {\rm 1}} {\rm s}^{ - {\rm 1}} )} \hfill & {{\rm for}} \hfill & {{\rm Si(CH}_{\rm 3} {\rm )}_{\rm 4} = {{10^{11.9} {\rm e}^{ - 7300} } \mathord{\left/ {\vphantom {{10^{11.9} {\rm e}^{ - 7300} } {RT}}} \right. \kern-\nulldelimiterspace} {RT}}} \hfill \\ {k({\rm ml m}^{ - {\rm 1}} {\rm s}^{ - {\rm 1}} )} \hfill & {{\rm for}} \hfill & {{\rm Ge(CH}_{\rm 3} {\rm )}_{\rm 4} = {{10^{11.7} {\rm e}^{ - 7370} } \mathord{\left/ {\vphantom {{10^{11.7} {\rm e}^{ - 7370} } {RT}}} \right. \kern-\nulldelimiterspace} {RT}}} \hfill \\ {k({\rm ml m}^{ - {\rm 1}} {\rm s}^{ - {\rm 1}} )} \hfill & {{\rm for}} \hfill & {{\rm Sn(CH}_{\rm 3} {\rm )}_{\rm 4} = {{10^{11.7} {\rm e}^{ - 7250} } \mathord{\left/ {\vphantom {{10^{11.7} {\rm e}^{ - 7250} } {RT}}} \right. \kern-\nulldelimiterspace} {RT}}} \hfill \\ \end{array} $$\end{document} The rate constants correlate with the proton chemical shift, which is related to a polar effect. In all cases except carbon, a hot-molecule β-fluorine rearrangement-elimination reaction occurs following radical combination: \documentclass{article}\pagestyle{empty}\begin{document} $$ ({\rm CF}_3 {\rm CH}_2 {\rm M}({\rm CH}_3 )_3 )^* \to {\rm CF}_2 {\rm CH}_2 + {\rm FM}({\rm CH}_3 )_3 $$ \end{document} We suggest the occurrence of a radical exchange reaction for the Si, Sn, and Ge systems, \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm CF}_3 + {\rm M}({\rm CH}_3 )_4 \rightleftharpoons{\rm CF}_3 {\rm M}({\rm CH}_3 )_4 \to {\rm CH}_3 + {\rm CF}_3 {\rm M}({\rm C}{\rm H}_3 )_3 $$\end{document} with kexchange (CF3 + Sn(Me)4) ∼ 107 ml m-1 s-1.

Notes: A check of the data from comparative rate single-pulse shock tube experiments have been carried out through the use of a new standard reaction, the decyclization reaction of ethylcyclobutane. The rate expressions for cyclohexene and 2,2,3-trimethylbutane have been found to be \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{l} k({\rm C}_{\rm 6} {\rm H}_{{\rm 10}} \to 1,3 - {\rm C}_4 {\rm H}_6 + {\rm C}_2 {\rm H}_4 ) = 10^{15.3} \exp ( - 33,690/T)\sec ^{ - 1} ,950^ \circ - 1100^ \circ {\rm K,2} - {\rm 6atm} \\ k(t{\rm C}_4 {\rm H}_9 - {\rm iC}_3 {\rm H}_7 \to t{\rm C}_{\rm 4} {\rm H}_{\rm 9} \cdot + {\rm iC}_{\rm 3} {\rm H}_7 \cdot ) = 10^{16.5} {{\exp ( - 36,830} \mathord{\left/ {\vphantom {{\exp ( - 36,830} T}} \right. \kern-\nulldelimiterspace} T})\sec ^{ - 1} ,1000^ \circ - 1100^ \circ {\rm K,2} - {\rm 6atm} \\ \end{array} $$\end{document} in excellent agreement with previously published results. Most of the small discrepancy that does exist is apparently due to the differences between the present and earlier (decomposition of isopropyl bromide) "standard" reaction. For the latter process, the present study yields \documentclass{article}\pagestyle{empty}\begin{document}$$ k({\rm iC}_{\rm 3} {\rm H}_7 {\rm Br} \to {\rm C}_{\rm 3} {\rm H}_6 + {\rm HBr}) \to 10^{13.73} {\rm exp(}{{ - 23,970} \mathord{\left/ {\vphantom {{ - 23,970} {T\sec ^{ - 1} ,800^ \circ - 1000^ \circ {\rm K,2} - {\rm 6}}}} \right. \kern-\nulldelimiterspace} {T\sec ^{ - 1} ,800^ \circ - 1000^ \circ {\rm K,2} - {\rm 6}}}{\rm atm)} $$\end{document} These results confirm the correctness of previously published comparative rate single-pulse shock tube experiments. They demonstrate once again that for the decomposition of paraffin hydrocarbons, calculated preexponential factors are at least an order of magnitude higher than the directly measured number and that the accepted value of the heat of formation of t-butyl radicals ΔHf300(tC4H9·) = 29 kJ (6.8 kcals) is at least 10 kJ too low. Finally, attention is called to recent studies on neopentane decomposition in flow and static systems which are in complete agreement with the present conclusions.

Notes: The kinetics of the gas-phase reaction CH3COCH3 + I2 ⇄ CH3COCH2I + HI have been measured spectrophotometrically in a static system over the temperature range 340-430°. The pressure of CH3COCH3 was varied from 15 to 330 torr and of I2 from 4 to 48 torr, and the initial rate of the reaction was found to be consistent with \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm CH}_3 {\rm COCH}_3 + {\rm I}^{\rm .} \stackrel{1}{\rightarrow}{\rm CH}_{\rm 3} {\rm COCH} + {\rm HI} $\end{document} as the rate-determining step. An Arrhenius plot of the variation of k1 with temperature showed considerable scatter of the points, depending on the conditioning of the reaction vessel. After allowance for surface catalysis, the best line drawn by inspection yielded the Arrhenius equation, log [k1/(M-1 sec-1)] = (11.2 ± 0.8) - (27.7 θ 2.3)/θ, where θ = 2.303 RT in kcal/mole. This activation energy yields an acetone C—H bond strength of 98 kcal/mole and δHf0 (CH3COĊH2) radical = -5.7 ± 2.6 kcal/mole. As the acetone bond strength is the same as the primary C—H bond strength in isopropyl alcohol, there is no resonance stabilization of the acetonyl radical due to delocalization of the radical site. By contrast, the isoelectronic allyl resonance energy is 10 kcal/mole, and reasons for the difference are discussed in terms of the π-bond energies of acetone and propene.

Notes: The absolute rate constants have been measured for several gas-phase chlorine atom-molecule reactions at 25°C by resonance fluorescence. These reactions and their corresponding rate constants in units of cm3 mole-1 sec-1 are: The effects of varying the substrate pressure, total pressure, light intensity and chlorine-atom source on the value of the bimolecular rate constants have been investigated for all these reactions. Conditions under which no competing side reaction occurs were established and the reported rate constants were measured under these conditions. For reactions (2), (5), (6), (7), and 8, there is a discrepancy of a factor of two between the rate constants measured in this work and values in the literature; it is suggested that this is due to an error in the previously measured value of kCH4/kH2 upon which the relative measurements in the literature ultimately depend.

Notes: The kinetics and absolute rate constants for the free-radical chain reaction of tri-n-butyltin hydride with di-t-butyl disulfide have been measured in cyclohexane at 30°. The rate controlling step for chain propagation involves the cleavage of the disulfide bond by an attacking tributyltin radical. The rate constant for this bimolecular homolytic substitution at sulfur is ∼8 × 104 Mole-1 sec-1. Chain termination involves the self-reaction of two tributyltin radicals.The rate constants for attack of tributyltin radicals on some other disulfides and on elemental sulfur have also been measured. The results are compared with literature data for homolytic substitutions on these compounds by a variety of radicals which have their unpaired electron centered on carbon.

Notes: The kinetics of the ethane pyrolysis have been studied at temperatures from 550 to 596°C and with 0 to 62% of added nitric oxide. The rates of production of various products were studied by gas chromatography; ethylene, hydrogen, methane, nitrogen, water, nitrous oxide and acetonitrile were found as primary products, with hydrogen cyanide, carbon monoxide, acetaldehyde, n-butane, 1-butene, cis- and trans-2-butene and 1,3-butadiene as secondary products. For all the primary products the orders with respect to C2H6and NO were determined, as were the activation energies at two different percentages of NO (15.7 and 45.5%).Nitric oxide was found to be rapidly consumed with a finite initial rate, and the rate of production of H2O was close to that of C2H4 at higher nitric oxide pressures.A mechanism is proposed which gives good agreement with all of the observed results. Its main features are: (1) Initiation takes place mainly by the unimolecular dissociation of ethane; there is no evidence for or against the process NO + C2H6 → HNO + C2H5; (2) NO scavenges ethyl radicals to form acetaldoxime which decomposes, and in this way the breakdown of C2H5 is hastened; (3) termination takes place mainly by the unimolecular decomposition of acetaldoxime to give inactive products. Some of the relevant rate parameters are evaluated. Reactions are proposed to account for the formation of the secondary products observed.

Notes: The kinetics of the gas-phase dehydrogenation of cyclopentane to cyclopentene is found to be consistent with a slow attack by an I atom (step 4, text) on cyclopentane in the range 282-382°C. The measured rate constants fit the Arrhenius equation, log k4 = 11.95 ± 0.08 - (24.9 ± 0.23)/θ 1 mole-1 sec-1, where θ = 2.303RT in kcal/mole. This leads to a value of ΔHf,2980 = 24.3 ± 1 kcal/mole and a bond dissociation energy DH = 94.9 ± 1 kcal/mole. The latter value is identical with DH0(i-Pr-H) = 95 ± 1 kcal/mole and signifies that cyclopentane and the cyclopentyl radical have the same strain energy. Arrhenius parameters are deduced for all six steps in the reaction mechanism. Surface reactions are shown to be unimportant.Cyclopentyl iodide is an unstable intermediate in the reaction and the rate constant for its bimolecular formation from HI + cyclopentene is found to be log k6 = 8.40 ± 0.29 - (26.9 ± 0.8)/θ 1 mole-1 sec-1. Together with the equilibrium constant, this yields for the unimolecular elimination of HI from cyclopentyl iodide, the rate constant, log k5 = 13.3 ± 0.3 - (42.8 ± 1.2)/θ sec-1.

Notes: The thermal decomposition of 1,2-dichloropropane at atmospheric pressure has been studied in the temperature range 227-590°C, in a flow system. Above 450°C, the reaction is homogenous and unimolecular with a rate constant: \documentclass{article}\pagestyle{empty}\begin{document}$$k = 10^{12.95 \pm 0.15} \exp (- 53,070 \pm 500/RT)\sec ^{ - 1}$$\end{document}Below 450°C, a low activation energy, probably heterogenous process competes with the gas phase reactionThe primary reaction products are HCl and the monochloropropene isomers; the relative amounts of each isomer depend on the temperature in the low but not in the high temperature region. The direction of the HCl elimination is discussed in terms of substituent effects at the α- and β-carbon positions and compared with literature data on similar reactionsSecondary products are formed principally by further pyrolysis of allyl chloride. The first-order rate constant of this reaction is given by: \documentclass{article}\pagestyle{empty}\begin{document}$$ k = 10^{8.54 \pm 0.2} \exp (- 37,275 \pm 700/RT)\sec ^{ - 1} $$\end{document}.

Notes: A record of the time dependence of the difference between two signals, one proportional to the concentration of a reactant or product in one reaction mixture and the other proportional to the concentration of the same or a corresponding substance in another mixture in which the reaction is initiated at the same time as the first, makes it possible to obtain not only the ratio, but also the individual values, of the rate constants for the two reactions. The effects of the experimental variables on a number of measurable parameters are examined, the errors associated with a number of different ways of evaluating the rate constants and their ratio are discussed, and it is shown how conditions can be selected that should provide values whose precisions compare favorably with those attainable by other techniques.

Notes: At 540°C, H2S (0.1 to 5 mm Hg) diminishes the initial rate of pyrolysis of C2H6 (50 mm Hg) into C2H4 + H2 and, even more strongly, the rate of appearance of the traces of nC4H10; on the contrary, the initial rate of formation of the traces of CH4 is practically not modified.A mechanism is proposed in order to interpret these experimental facts.

Notes: Nitrogen quantum yields are reported for the photolysis of C2F5N=NC2F5 at 3660 Å over the pressure range 2-10 cm from 25° to 150°c. The Stern-Volmer plots obtained are discussed and compared with those obtained with azoethane and azoisopropane.

Notes: The kinetics of the decomposition of benzotrifluoride was studied from 720°c to 859°c in a flow system with and without carrier gas. Consideration of the product distribution made possible the study of the decomposition into CF3 and C6H5 radicals, which appeared to be truly homogeneous in character. The first-order rate constant of the C—C bond fission, log k (sec-1) = (17.9 ± 0.5) (99.7 ± 2.5)/θ, did not change with change of initial concentration, pressure of the carrier gas, or contact time. The Arrhenius parameters have been related to the appropriate thermodynamic data. Assumption of 0 kcal/mole for the activation energy of the reverse combination reaction yielded DH298°(C6H5—CF3) = 103.6 ± 2.5 kcal/mole and ΔHf298°(C6H5) = 77.1 ± 3.0 kcal/mole.Applicability of the simple first-order formula to calculation of the rate constant has been also dealt with.

Notes: Combination reactions of the methyl radical have been studied by following the decay of the absorbance of the methyl radical during the course of the reaction by means of kinetic spectroscopy. The limiting values of the second-order rate constants at high pressure were determined for two reactions at room temperature: \documentclass{article}\pagestyle{empty}\begin{document}$ \begin{array}{*{20}c} {{\rm CH}_{\rm 3} + {\rm CH}_{\rm 3} ( + {\rm M}) \to {\rm CH}_{\rm 3} {\rm CH}_{\rm 3} ( + {\rm M}):k_1 = (2.6 \pm 0.3) \times 10^{10} 1{\rm mole}^{ - {\rm 1}} \sec ^{ - 1} } \\ {{\rm CH}_{\rm 3} + {\rm NO }( + {\rm M}) \to {\rm CH}_{\rm 3} {\rm NO }( + {\rm M}):k_2 = (2.4 \pm 0.2) \times 10^9 1{\rm mole}^{ - {\rm 1}} \sec ^{ - 1} } \\ \end{array} $\end{document} The extinction coefficient of the methyl radical was found to have a maximum value of (1.02 ± 0.06) X 104 1 mole-1 cm-1 at 216.4 nm. Integration of the extinction coefficient over the absorption band of the methyl radical gave an oscillator strength of 1.0 X 10-2.

Notes: A simple electrostatic model of point dipoles is used which permits direct calculation of the activation energies for the addition of the molecules H2O, H2S, H3N, and H3P to olefins. These calculated values agree with the known experimental data to within ±2 kcal/mole on the average. It was found that the best fit could be obtained with a polar transition state that corresponded to a reduction in bond order from 1 to ½ for the bond-breaking coordinates and an increase in bond order from 0 to 0.18 for the bond-forming coordinates. The replacement of a hydrogen atom of the species H2O, H2S, H3N, or H3P by a polarizable methyl group is expected to stabilize the charge on the central atoms. The following stabilization energies for the pairs H2O—CH3OH, H2S—CH3SH, H3N—CH3NH2, H3P—CH3PH2 were calculated: -4.8 kcal/mole, -0.7 kcal/mole, -1.9 kcal/mole, -0.8 kcal/mole, respectively.

Notes: The reaction of trifluoromethyl radicals with ammonia in the gas phase has been studied in the temperature range 30-352°. Product formation is not explicable simply in terms of the reactions: \documentclass{article}\pagestyle{empty}\begin{document}$\begin{array}{*{20}c} {{\rm CF}_{\rm 3} + {\rm NH}_{\rm 3} \to {\rm CF}_3 {\rm H} + {\rm NH}_2 } \\ {2{\rm CF}_3 \to {\rm C}_2 {\rm F}_6 } \\\end{array}$\end{document} and curvature of the Arrhenius plot at high and low temperatures suggests that there are additional sources of fluoroform.

Notes: The vibrationally excited molecules CF3CH3* and CF3CD3* have been synthesized by radical combination (produced by ketone photolysis), and HF and DF elimination from them studied as a function of temperature and pressure. Using RRK theory many calculations have recently been made of critical energies for the decomposition of "hot" fluoroethane molecules. Taking CF3CH3* as an example, it is concluded that the empiricism involved in such calculations renders results of doubtful significance. The non-equilibrium kinetic isotope effect is kH/kD = 3.1 at 470°K. Arrhenius parameters are also presented for radical abstraction reactions from the ketone source molecules.

Notes: Rate constants for the bimolecular reactions of ethylene and propylene to form radicals and to form cyclobutane or its derivatives have been calculated using thermodynamic and kinetic data. Comparison of these rates with the kinetics of the thermal reactions of ethylene and propylene show that cyclobutane and its derivatives are probably not important intermediates in the processes forming radicals.

Notes: The gas phase reaction of iodine (2.8-43.3 torr) with methyl ethyl ketone (MEK) (7.4-303.4 torr) has been studied over the temperature range 280-355°C in a static system. The initial rate of disappearance of I2 is first order in MEK and half order in I2. The rate-determining step is the abstraction of a secondary hydrogen atom by an iodine atom: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm CH}_3 {\rm COCH}_2 {\rm CH}_2 {\rm } + {\rm I}^{\rm .} \stackrel{1}{\longrightarrow}{\rm CH}_{\rm 3} {\rm COCHCHL3 + HI} $$\end{document} where k1 is given by \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}_{{\rm 10}} ({{k_1 } \mathord{\left/ {\vphantom {{k_1 } {{\rm M}^{ - 1} \sec ^{ - 1} }}} \right. \kern-\nulldelimiterspace} {{\rm M}^{ - 1} \sec ^{ - 1} }}) = {{(10.76 \pm 0.15) - (22.0 \pm 0.4)} \mathord{\left/ {\vphantom {{(10.76 \pm 0.15) - (22.0 \pm 0.4)} \theta }} \right. \kern-\nulldelimiterspace} \theta } $$\end{document} and θ = 2.303RT in kcal/mole. This activation energy is equivalent to a secondary C—H bond strength of 92.3 ± 1.4 kcal/mole and ΔHf 2980 of the methylacetonyl radical = -16.8 ± 1.7 kcal/mole. By comparison with 95 kcal/mole for the secondary C—H bond strength, when delocalization of the unpaired electron with a pi bond is not possible, the resonance stabilization of the methylacetonyl radical is calculated to be 2.7 ± 1.7 kcal/mole. This value is 10 kcal/mole less than the stabilization energy of the isoelectronic methylallyl radical. The difference in pi bond energies in the canonical forms of the methylacetonyl radical is shown to account for the variation in stabilization energies.

Notes: The reactions of CF3 radicals with the C3 to C7 cyclanes and spiropentane were studied and the following Arrhenius parameters were obtained for the reaction CF3 + c-RH → CF3H + c-R: Textc-RHlog A (cm3mole-1sec-1)E (kcal/mole)D(c-R - H) (kcal/mole)Cyclopropane11.548.73100.7Cyclobutane11.666.4895.7Cyclopentane12.306.1894.3Cyclohexane12.126.2694.9Cycloheptane12.435.8994.0Spiropentane11.918.1298.8The CF3 radicals were generated by photolysis of hexafluoroacetone or CF3I and a comparison is made of the utility of the two compounds as radical sources. The Arrhenius parameters are compared with those for corresponding reactions of CH3 radicals with cyclanes, and the general reactivity of cyclic compounds toward free radicals is discussed. An Evans-Polanyi treatment is used to derive C—H bond dissociation energies in cyclanes; and these results, based on the reactions of CF3 with cyclanes, agree well with those previously obtained using CH3 plus cyclanes. The final mean values are shown above.

Notes: The competitive reactions of Br atoms with CH4 and CD4 were studied over the temperature range of 562° to 637°K. Over this temperature interval, the kinetic isotope effect, kH/kD, varied from 3.05 to 2.47 for the reactions \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {{\rm Br} + {\rm CH}_{\rm 4} \stackrel{k_H}{\rightarrow}{\rm HBr} + {\rm CH}_3 } \\ {{\rm Br} + {\rm CD}_{\rm 4} \stackrel{k_D}{\rightarrow}{\rm DBr} + {\rm CD}_3 } \\ \end{array} $$\end{document} The rate constant ratio kH/kD, expressed in Arrhenius form, was found to equal (1.10 ± 0.05) exp (1030 ± 60/RT). A comparison is presented between the experimental result and the result obtained theoretically from absolute rate theory using the London-Eyring-Polanyi-Sato (LEPS) method of constructing the potential energy surface of the reaction. The agreement between theory and experiment is very poor, and this is believed to arise from the highly unsymmetrical nature of the potential energy surface involved in these reactions. A comparison is also presented between the kH/kD values obtained in the Br + CH4-CD4 experiments and the available data on the corresponding Cl + CH4-CD4 reactions.

Notes: The reversible isomerization of cis-hepta-1,3-diene to cis-2-trans-4-heptadiene via a 1,5 hydrogen shift has been investigated kinetically at nine temperatures in the range of 475° to 531°K. Equilibrium is reached near 94% reaction. Some cis-2-cis-4-heptadiene is also formed, but at a rate some 60 times slower than the cis,trans isomer. A least-squares analysis of the data yielded the Arrhenius equation for the isomerization of the cis-hepta-1,3-diene: \documentclass{article}\pagestyle{empty}\begin{document}$$ {{\log k_1 } \mathord{\left/ {\vphantom {{\log k_1 } {\sec ^{ - 1} = {{(11.110 \pm 0.087) - ([33,210 \pm 200]{\rm cal mole}^{ - {\rm 1}} )} \mathord{\left/ {\vphantom {{(11.110 \pm 0.087) - ([33,210 \pm 200]{\rm cal mole}^{ - {\rm 1}} )} \theta }} \right. \kern-\nulldelimiterspace} \theta }}}} \right. \kern-\nulldelimiterspace} {\sec ^{ - 1} = {{(11.110 \pm 0.087) - ([33,210 \pm 200]{\rm cal mole}^{ - {\rm 1}} )} \mathord{\left/ {\vphantom {{(11.110 \pm 0.087) - ([33,210 \pm 200]{\rm cal mole}^{ - {\rm 1}} )} \theta }} \right. \kern-\nulldelimiterspace} \theta }}} $$\end{document} Possible errors in the equilibrium constant measurements are discussed, and employing an equilibrium constant calculated by using group additivity estimates together with the values of k1, we obtained for the reverse reaction \documentclass{article}\pagestyle{empty}\begin{document}$$ {{\log k_2 } \mathord{\left/ {\vphantom {{\log k_2 } {\sec ^{ - 1} = (11.1 \pm 0.15) - ([35,800 \pm 300])}}} \right. \kern-\nulldelimiterspace} {\sec ^{ - 1} = (11.1 \pm 0.15) - ([35,800 \pm 300])}}{{{\rm cal mole}^{ - {\rm 1}} )} \mathord{\left/ {\vphantom {{{\rm cal mole}^{ - {\rm 1}} )} \theta }} \right. \kern-\nulldelimiterspace} \theta } $$\end{document} where \documentclass{article}\pagestyle{empty}\begin{document}$$ \theta = 2.303RT $$\end{document}.

Notes: The kinetics of the acetaldehyde pyrolysis have been studied at temperatures from 450° to 525°C, at an acetaldehyde pressure of 176 torr and at 0 to 40 torr of added nitric oxide. The following products were identified and their rates of formation measured: CH4, H2, CO, CO2, C2H4, C2H6, H2O, C3H6, C2H5CHO, CH3COCH3, CH3COOCH=CH2, N2, N2O, HCN, CH3NCO, and C2H5NCO. Acetaldehyde vapor was found to react with nitric oxide slowly in the dark at room temperature, the products being H2O, CH3COOCH3, CO, CO2, N2, NO2, HCN, CH3NO2, and CH3ONO2. The rates of formation of N2 and C2H5NCO depend on how long the CH3CHO-NO mixture is kept at room temperature before pyrolysis; the rates of formation of the other products depend only slightly on the mixing period.The pyrolysis of “clean” CH3CHO-NO mixtures (i.e., the results extrapolated to zero mixing time, which are independent of products formed in the cold reaction) are interpreted as follows: (1) There are two chain carriers, CH3 and CH2CHO, their concentrations being interdependent and influenced by NO in different ways: the CH3 radical is both generated and removed by reactions directly involving NO, whereas CH2CHO is generated only indirectly from CH3 but is also removed by direct reaction with NO. (2) An important mode of initiation by NO is its addition to the carbonyl group with the formation of which is converted into ; this splits off OH with the formation of CH3NCO or CH3 + OCN. (3) Important modes of termination are \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {{\rm CH}_3 + {\rm NO} \to {\rm CH}_3 {\rm NO} \to {\rm CH}_2 {\rm NOH} \to {\rm HCN} + {\rm H}_2 {\rm O}} \\ {{\rm CH}_{\rm 2} {\rm CHO} + {\rm NO} \to {\rm CH}_2 ({\rm NO}){\rm CHO} \to {\rm CH}({\rm NOH}){\rm CHO} \to {\rm HCN} + {\rm H}_2 {\rm O} + {\rm CO}} \\ \end{array} $$\end{document} The steady-state equations derived from the mechanism are shown to give a good fit to the experimental rate versus [NO] curves and, in particular, explain why there is enhancement of rate by NO at higher CH3CHO pressures and, at lower CH3CHO pressures, inhibition at low [NO] followed by enhancement at higher [NO].The cold reaction is explained in terms of chain-propagating and chain-branching steps resulting from the addition of several NO molecules to CH3CHO and the CH3CO radical. In the “unclean” reaction it is found that the rates of N2 and C2N5NCO formation are increased by CH3NO2, CH3ONO, and CH3ONO2 formed during the cold reaction. A mechanism is proposed, involving the participation of α-nitrosoethyl nitrite, CH3CH(NO)ONO.It is suggested that there are two modes of behavior in pyrolyses in the presence of NO: (1) In the paraffins, ethers, and ketones, the effects are attributed to the addition of NO to a radical with the formation of an oxime-like compound. (2) In the aldehydes and alkenes, where there is a hydrogen atom attached to a double-bonded carbon atom, the behavior is explained in terms of addition of NO to the double bond followed by the formation of an oxime-like species.

Notes: Equilibrium constants for the reaction CH3COCH2CH3 + I2 ⇌ CH3COCHICH3 + HI have been computed to fit the kinetics of the reaction of iodine atoms with methyl ethyl ketone. From a calculated value of S2980(CH3COCHICH3) = 93.9 ± 1.0 gibbs/mole and the experimental equilibrium constants, ΔHf0(CH3COCHICH3) is found to be -38.2 ± 0.6 kcal/mole. The Δ(ΔHf2980) value on substitution of a hydrogen atom by an iodine atom in the title compound is compared with that for isopropyl iodide. The relative instability of 2-iodo-3-butanone (3.4 kcal/mole) is presented as further evidence for intramolecular coulombic interaction between partial charges in polar molecules. The unimolecular decomposition of 2-iodo-3-butanone to methyl vinyl ketone and hydrogen iodide was also measured in the same system. This reaction is relatively slow compared to the formation of the above equilibrium. Rate constants for the reaction over the temperature range 281°-355°C fit the Arrhenius equation: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log(}{{k_{\rm 3} } \mathord{\left/ {\vphantom {{k_{\rm 3} } {{\rm sec}^{ - 1} }}} \right. \kern-\nulldelimiterspace} {{\rm sec}^{ - 1} }}{\rm )} = {\rm 13}{\rm .4} - {{(41.9 \pm 0.5)} \mathord{\left/ {\vphantom {{(41.9 \pm 0.5)} \theta }} \right. \kern-\nulldelimiterspace} \theta } $$\end{document} where θ = 2.303RT kcal/mole. The stability of both the ground and transition states is discussed in comparing this activation energy with that reported for the unimolecular elimination of hydrogen iodide from other secondary iodides. The kinetics of the reaction of hydrogen iodide with methyl vinyl ketone were also measured. The addition of HI to the double bond is not rate controlling, but it may be shown that the rate of formation of 1-iodo-3-butanone is more rapid than that for 2-iodo-3-butanone. Both four- and six-center transition complexes and iodine atom-catalyzed addition are discussed in analyzing the relative rates.

Notes: Numerical Methods were used to solve the differential equation for diffusion of a trace gas into a flowing carrier gas having a parabolic velocity profile in a cylindrical tube. Steady state solutions are given in the form of contour diagrams of constant trace gas concentration.

Notes: The reaction of nitrogen atoms with methylacetylene has been studied using a fast-flow low-pressure reactor coupled to a mass spctrometer by a nozzle-beam sampling system. Hydrogen atom concentrations were measured by ESR analysis. Experimental second-order rate constants for the consumption of N atoms, of C3H4, and for the formation of N2 were determined in the temperature range of 283° to 485°K. Product profiles of all stable species and of hydrogen atoms and methyl radicals were obtained for different initial concentrations of the reactants. Two different reaction pathways can be distinguished: one provides for recombination of N atoms, and the second leads to the formation of cyano compounds and other hydrocarbons. Only the latter process is influenced by the addition of hydrogen atoms. Mechanisms for the two pathways are discussed.

Notes: The quantitative kinetics of pyrolysis of some cyclic and polycyclic compounds are examined from the point of view of biradical intermediates. Transition state methods for estimating the Arrhenius parameters of the lowest free energy biradical pathway to products in cyclic and polycyclic compound reactions are described and illustrated. A large number of the polycyclic reactions are found to have Arrhenius parameters consistent with the biradical mechanism estimates. Other reactions are found to have much faster experimental rates and must therefore be concerted. The value of activation energy and activation entropy estimates as discriminatory tests of mechanism, i.e., single step (concertedness) or consecutive step processes, is discussed.