ALBERT

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Notes: Observations are reported of the effect of the buffer gases He, Ne, and CF4, in the pressure range of 0-30 torr, on the branching ratio [HCl]/[DCl] of the unimolecular decomposition The ratio R = kH/kD has been measured in high-pressure thermal decomposition (670-1100 K) and was shown to give a unique measure of the internal energy of the decomposing molecules and hence, with RRKM theory and pressure fall-off data, a time scale for their decomposition.Applying the thermal data to the photolysis leads to the conclusion that excitation and decomposition are produced by the laser spike (high intensity, 70 ns FWHM) and also at a slower rate by the larger, less intense tail (1.6 μs). Added buffer gases quench the latter, leaving the former which, from measurements of R, is shown to correspond to excitations of 115 ± 15 kcal/mol and lifetimes of ∼30 ps. No bond breaking is seen despite the high energies, in accord with theoretical expectations. The results require an enhanced rate of photon absorption by the highly excited molecules, which are about hundredfold greater than that observed for 300 K molecules. Data are also reported for C2H2F2 and the secondary multiphoton photolysis of the ethylenes produced. Effects of beam geometry and wavelength are explored.

Notes: The rate constant for the reaction of ground-state oxygen atoms with methanol has been determined between 297 and 544 K by a phase-shift technique using mercury photosensitized decomposition of N2O to generate oxygen atoms. The relative oxygen atom concentration was monitored by the chemiluminescence from the reaction of oxygen atoms with nitric oxide. The results are accommodated by the Arrhenius expression k1 = (9.79 ± 2.71) × 1012 exp[(-2267 ± 111)/T]cm3/mol·s, where the indicated uncertainties are 95% confidence limits for 10 degrees of freedom. As an incidental part of this work, the third-body efficiency of CH3OH relative to N2O for the reaction O + NO + M → NO2 + M (M = CH3OH) was determined to be 3.1 at 298 K.

Notes: The flash photolysis resonance fluorescence technique was used to measure the rate constants of the reaction O + O2 + M → O3 + M (M = N2, O2, Ar, and He) as a function of temperature. The results for the rate constants are given by The activation energies with N2, O2, and Ar as third bodies are equal within the experimental error, (-1370 → 340 cal/mol), and the relative third-body efficiencies at 298 K for N2, O2, Ar, and He are 1.00, 0.99, 0.69, and 0.60, respectively.

Notes: The kinetics of the gas-phase thermal reaction between CF2(OF)2 and CO has been studied in a static system at temperatures ranging between 110 and 140°C. The only reaction products were CF2O and CO2, giving the following stoichiometry: \documentclass{article}\pagestyle{empty}\begin{document}$${\rm CF}_{\rm 2} {\rm (OF)}_{\rm 2} {\rm + 2CO = 2CF}_{\rm 2} {\rm O + CO}_{\rm 2} {\rm}\Delta n{\rm = 0}$$\end{document} The reaction is homogeneous. The rate is strictly second order in CF2(OF)2 and CO, and is not affected by the total pressure or by the presence of reaction products. Oxygen promotes a sensitized oxidation of CO and inhibits the formation of CF2O.The experimental results in the absence of oxygen can be explained by a chain mechanism similar to that proposed for the reaction between F2O and CO with an overall rate constant of \documentclass{article}\pagestyle{empty}\begin{document}$$k_1 = 1.45 \times 10^9 {\rm exp}(- 20,900/RT)L/mol \cdot s$$\end{document} From the experimental data obtained on the oxygen-inhibited reaction, the rate constant for the primary process can be calculated: \documentclass{article}\pagestyle{empty}\begin{document}$$\begin{array}{*{20}c} {({\rm I})} \quad {{\rm CF}_{\rm 2} ({\rm OF)}_{\rm 2} + {\rm CO} \to {\rm CF}_{\rm 2} (\mathop {\rm O}\limits^{\rm .}){\rm OF} + {\rm F}\mathop {\rm C}\limits^{\rm .} {\rm O}} \quad\quad {k_1 = 1.45 \times 10^9 {\rm exp}(- 20,900/RT)L/mol \cdot s} \\\end{array}$$\end{document} The chain length v = 2.5 is independent of the temperature. Taking for collision diameters σCF2(OF)2 = 6 Å and σCO = 3.74 Å, a value α = 5.3 × 10-3 for the steric factor is obtained.

Notes: The title reaction has been investigated in the temperature range of 490-573 K. Initial reactant pressures were varied in the range of 0.2-5.2 torr (I2) and 2-20 torr (C6H5SiH3). The rate of iodine consumption, monitored spectrophotometrically, was found to obey \documentclass{article}\pagestyle{empty}\begin{document}$$ - \frac{{d[{\rm I}_{\rm 2}]}}{{dt}} = \frac{{k_{3/2} [{\rm I}_{\rm 2}]^{{\raise0.7ex\hbox{$1$} \!\mathord{\left/ {\vphantom {1 2}}\right.\kern-\nulldelimiterspace} \!\lower0.7ex\hbox{$2$}}} [{\rm C}_{\rm 6} {\rm H}_{\rm 5} {\rm SiH}_{\rm 3}]}}{{1 + k'[HI]/[I_2]}}$$\end{document} both by initial rate and integrated equation fitting procedures. The effect of added initial HI conformed to this expression. The data are consistent with a conventional I-atom propagated chain reaction, and for the step \documentclass{article}\pagestyle{empty}\begin{document}$${\rm I}^{\rm .} + {\rm C}_{\rm 6} {\rm H}_{\rm 5} {\rm SiH}_{\rm 3} \to {\rm C}_{\rm 6} {\rm H}_{\rm 5} \mathop {\rm S}\limits^{\rm .} {\rm iH}_{\rm 2} + {\rm HI}$$\end{document} the rate constant is given by \documentclass{article}\pagestyle{empty}\begin{document}$${\rm log}k_1 (dm^3 /mol \cdot s) = (11.52 \pm 0.08) - (76.8 \pm 0.8{\rm kJ/mol})/RT{\rm ln}10$$\end{document} From this is derived the bond dissociation energy value C6H5SiH2—H = 374 kJ/mol(88 kcal/mol). A comparison with other Si—H dissociation energy values indicates that the “silabenzyl” stabilization energy is small, ≈7 kJ/mol.

Notes: Diethylhydroxylamine, (C2H5)2NOH, was oxidized by NO2 at 25°C in a long-path-length infrared gas cell. The measured products of the reaction were HONO and CH3CHO. The reaction scheme which explains the reaction is was oxidized by NO2, and the reaction was found to be very rapid with k1 〉 10-16 cm3/s. The products of the reaction were verified by both infrared absorption (CH3CHO, C2H5NO) and gas chromatography (CH3CHO, NO).

Notes: The rates of oxidation of phenylthioacetic acid (PTAA) and several substituted phenylthioacetic acids by potassium peroxodiphosphate (PP) in 50% (v/v) aqueous acetic acid have been studied in detail. The rate of oxidation is expressed as \documentclass{article}\pagestyle{empty}\begin{document}$$\frac{{- d[{\rm PP]}}}{{{\rm dt}}} = k[{\rm PP][PTAA][H}^{\rm +}]$$\end{document}An analysis of the dependence of the rate on [H+] reveals that H3P2O8- is the active oxidizing species in the oxidation. The effect of ring substituents on the rate gives a ρ+ value of -0.45 ± 0.03 (r = 0.998, s = 0.02 at 40°C), pointing to the development of an electron-deficient center in the transition state. The results are discussed in terms of a mechanism involving the rate-determining formation of an intermediate between PP and phenylthioacetic acids, followed by the decomposition of the intermediate. These kinetic results are compared with those obtained in the oxidation of phenylthioacetic acids by peroxodisulfate.

Notes: A kinetic study on the oxidation of V(IV) by chloramine-T (CAT) at pH 6.85 by N-bromo succinimide (NBS) in aqueous acetic acid-perchloric acid media and by N-iodo succinimide (NIS) in aqueous perchloric acid medium has been carried out. In all the systems studied the order with respect to the oxidant is unity. NBS and CAT oxidation reactions exhibited Michaelis-Menten type kinetics, and the NIS study indicated unit dependence on [substrate]. Independence on acidity has been observed in the case of CAT and NBS reactions, but NIS reactions exhibited inverse unit dependence on [acid]. Novel solvent influences have been noticed in the case of CAT reactions, but with NIS and NBS reactions retardation in the rate has been observed with an increase in the percentage of acetic acid. Plausible mechanisms consistent with the results have been postulated, and suitable rate laws in consonance with the postulated mechanisms have been derived.

Notes: The experimental behavior of the cerium- and manganese-catalyzed Belousov Zhabotinskii oscillating reaction with ethyl acetoacetate as organic substrate has been investigated. Under certain conditions the system displays two types of temporal oscillations. Damped highfrequency oscillations appear immediately after the addition of potassium bromate solution to complete the reaction mixture. These high-frequency oscillations may be regarded as being superimposed on an induction period of the type found in the reaction using malonic acid. After the induction period, low-frequency oscillations of the normal type are obtained. Both the high-frequency and the low-frequency oscillations can be monitored with a platinum redox or with a bromide specific ion electrode.

Notes: Relative rate constants for the reaction of OH radicals with a series of ketones have been determined at 299 ± 2 K, using methyl nitrite photolysis in air as a source of hydroxyl radicals. Using a rate constant for the reaction of OH radicals with cyclohexane of 7.57 × 10-12 cm3 molecule-1 s-1, the rate constants obtained are (× 1012 cm3 molecule-1 s-1): 2-pentanone, 4.74 ± 0.14; 3-pentanone, 1.85 ± 0.34; 2-hexanone, 9.16 ± 0.61; 3-hexanone, 6.96 ± 0.29; 2,4-dimethyl-3-pentanone, 5.43 ± 0.41; 4-methyl-2-pentanone, 14.5 ± 0.7; and 2,6-dimethyl-4-heptanone, 27.7 ± 1.5. These rate constants indicate that while the carbonyl group decreases the reactivity of C—H bonds in the α position toward reaction with the OH radical, it enhances the reactivity in the β position.

Notes: A kinetic spectrophotometric investigation of the reaction of the hydrogen peroxide anion with methyl p-nitrophenyl sulfate in methanol solvent resulted in the evaluation of the pKa of HOOH in methanol at 25°C as 15.8 ± 0.2. Since normal kinetic procedures for the determination of the equilibrium constant K for the process CH3O- + H2O2 ⇄ CH3OH + HO2- were found to be associated with high uncertainty, another procedure was devised to establish the magnitude of K. This method is based on an analysis of the changing slopes of plots of pseudo-first-order rate constants against the total base concentration as the stoichiometric amount of hydrogen peroxide is varied. The method is applicable to any system in which anionic nucleophiles generated in situ compete with solvent anions. Such a corroboration of kinetically determined equilibrium constants is believed essential. The kinetic data allow the specific rate constant kHOO-for the reaction of methyl p-nitrophenyl sulfate with hydrogen peroxide anions to be evaluated and yield the rate constant ratio kHOO-/kMeO- = 8.8 ± 2.2. This confirms the existence of an α effect at saturated carbon in this system.

Notes: The rate law for the demetallation of the title indium(III)-porphin complex in aqueous acidic thiocyanate media at 3.00M ionic strength was found to be of the form \documentclass{article}\pagestyle{empty}\begin{document}$$+ \frac{{d{\rm [H}_{\rm 4} {\rm P}^{{\rm 2} - } {\rm]}}}{{dt}} = \frac{{ab[{\rm H}^{\rm + }]^2 [{\rm NCS}^ -]^2 }}{{1 + b[{\rm H}^{\rm + }]^2 }}[{\rm InP]}_{\rm t}$$\end{document} where [H4P2-] is the concentration of the diacid product formed, [InP]t is the total concentration of all forms of indium(III)-porphin complex present, and a and b are constants. The constant a is a pseudo-third-order rate constant with the value (0.057 ± 0.005)M-2 s-1 and b has the value 0.704M-2 at 50.5°C. If the mechanism for demetallation involves ringpuckering with the attachment of two H+ ions, then 1/b can be identified with the product K1K2 for the stepwise dissociation of two protons from two ring pyrrolic nitrogen atoms of H2InP-. In the sulfonated tetraphenylporphin used for these studies the ring pyrrolic nitrogen atoms seem to be the most probable sites for protonation. If this identification is correct, the value of 1.42 ± 0.13 found for the product K1K2 shows the enormous effect that the presence of the In3+ center has on the ionization constants of these two protons. That the kinetic studies show saturation effects with respect to proton addition to InP3- may result from the fact that In3+ sits about 0.6 Å above the porphin ring.

Notes: The kinetics of cleavage of 3-hydroxybicyclo[4.2.0]octa-1,3,5-trien-7-ones in aqueous sodium hydroxide, and of the alkoxy and acetoxy analogues in methanolic sodium methoxide solution, were examined under pseudo-first-order reaction conditions. The dependence of the rate upon the basicity of the solvent, whether measured by H- or by [OR-], reflects the possible structure of the transition state. The deduced mechanism is also supported by the effects of substituents upon the reaction rate. The relative amounts of the volatile reaction products derived from o-toluic acid and from phenylacetic acid are understood in terms of the substituent effect upon the relative stabilities of the carbanions.

Notes: The isomerization reaction of cholest-5-en-3-one has been studied in a solution of cyclohexane using trichloroacetic acid as catalyst. At the same time a general reaction scheme is proposed to be valid for all the cases assayed in which the monomer form of the acid is considered as the only effective catalyst. The experimental results agree with these hypotheses and with the calculation of the individual rate constant together with the reaction order with respect to the catalyst. Semiquantitative studies have been carried out with other catalysts and solvents, confirming the validity of the reaction scheme. The thermodynamic activation parameters have also been calculated, and a comparative study was made with the results of the evaluation of the reaction when it takes place in amphiprotic solvents. A reaction mechanism is proposed based on all the kinetic information obtained.

Notes: Kinetics of the basic hydrolysis of 1-glyceryl mononitrate (1-MNG) and 2-glyceryl mononitrate (2-MNG) were investigated in CO2-free aqueous calcium hydroxide solutions. The hydrolysis reactions were carried out in a temperature-controlled reactor vessel with provision for continuous N2 sparging of the reaction mixture. Both glyceryl nitrate esters hydrolyzed via second-order reaction at 25°C. 2-MNG in calcium hydroxide solution isomerized to 1-MNG, which subsequently hydrolyzed to form NO3-. In strongly basic aqueous solutions of NaOH (30%), 2-MNG is converted to glycidol and NO3-.

Notes: Relative rate constants for the gas-phase reactions of OH radicals with a series of alkyl nitrates have been determined at 299 ± 2 K, using methyl nitrite photolysis in air as a source of OH radicals. Using a rate constant for the reaction of OH radicals with cyclohexane of 7.57 × 10-12 cm3/molec·s, the rate constants obtained are (× 1012 cm3/molec·s): 2-propyl nitrate, 0.18 ± 0.05; 1-butyl nitrate, 1.42 ± 0.11; 2-butyl nitrate, 0.69 ± 0.10; 2-pentyl nitrate, 1.87 ± 0.12; 3-pentyl nitrate, 1.13 ± 0.20; 2-hexyl nitrate, 3.19 ± 0.16; 3-hexyl nitrate, 2.72 ± 0.22; 3-heptyl nitrate, 3.72 ± 0.43; and 3-octyl nitrate, 3.91 ± 0.80. These rate constants, which are the first reported for the alkyl nitrates, are significantly lower than those for the parent alkanes, and a formula, based on the numbers of the various types of C—H bonds in the alkyl nitrates, is derived for rate constant estimation purposes.

Notes: The literature results for the pyrolysis of bis trifluoromethyl peroxide are reexamined and compared with those for dimethyl peroxide. The thermochemistry yields the result that the π-bond energy in carbonyl fluoride is 96 ± 10 kcal/mol compared to 74 kcal/mol for that in formaldehyde. Thermodynamic additivity contributions are derived for the C—(F)3(O) and O—(C)(F) groups. Some conclusions are drawn in relation to the oxidation of halogeno methyl radicals and the chemistry of the atmosphere.

Notes: Aqueous iodination of trans-2-butenoic acid proceeds via hydrolysis of I2 to form HOI and I-, then rapid addition of HOI across the double bond to form the iodohydrin product. In the presence of iodate to keep iodide concentration low, the reaction proceeds at a conveniently measurable rate. The rate for the addition reaction \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm HOI + CH}_{\rm 3} {\rm CH=\!=CHCOOH} \to {\rm CH}_{\rm 3} {\rm CH(OH)CHICOOH}$$ \end{document} is -d[C4H6O2]/dt = 5900 [H+][C4H6O2][HOI]M/s at 25.0°C when [IO3-] = 0.025M and ionic strength = 0.3. The overall rate law in the presence of iodate is \documentclass{article}\pagestyle{empty}\begin{document}$$ -d[{\rm I}_{\rm 2}]/dt = 3.2 \times 10^{ - 3} \times 10^{ - 3} [{\rm H}^{\rm + }][{\rm IO}_{\rm 3}^ -]^{0.65} [{\rm C}_{\rm 4} {\rm H}_{\rm 6} {\rm O}_{\rm 2}]^{1/2} [{\rm I}_{\rm 2}]^{1/2} M/{\rm s}$$ \end{document} where [H+] and [IO3-] are total concentrations used to prepare the solution.

Notes: The kinetics of gas-phase decomposition of methyl isocyanate have been investigated in the range of 427-548°C. Two decomposition routes are followed; the predominant one is a radical-chain process giving CO, H2, and HCN as major products, which has an order of 1.5 and an Arrhenius equation given by log k(L1/2/mol1/2·s) = (13.12 ± 0.06) - (56,450 ± 1670) cal/mol/2.303 RT. The minor route is the bimolecular formation of N,N′-dimethylcarbodiimide and CO2, which from the low activation parameters Ea = 31.6 kcal, A = 105.30 L1/2/mol1/2·s, and the reaction order of 1.57 appears to be heterogeneous.

Notes: Pyridiniumchlorochromate (PCC) oxidizes aniline and substituted anilines except nitro anilines smoothly in chlorobenzene-nitrobenzene mixtures in the presence of dichloroacetic acid. The reaction has unit dependence on each of the aniline, PCC, and dichloroacetic acid concentrations. Electron-releasing substituents accelerate the reaction, whereas electronwithdrawing groups retard the reaction, and the rate data obey Hammett's relationship. The reaction constant ρ is -3.75. Azobenzene and p-benzoquinone have been obtained as products. The observed experimental data have been rationalized in terms of the formation of an intermediate complex involving PCC-amine undergoing a rapid decomposition to products.

Notes: The reactions of stabilized carbonium ions of setoglaucin, methyl violet, and ethyl violet with cyanide ions are largely catalyzed by the cationic micelles of cetyltrimethylammonium bromide (CTAB) in aqueous media. Added counterions (anions in this case) have strong inhibitory effects on the CTAB-catalyzed reactions in the following order: N3- 〉 NO3- 〉 Br- 〉 Cl- 〉 F- 〉 no salt. The inhibitory effects of the counterions have been attributed to the exchange between added anions and reagent (CN-) in the micellar media. The data have been analyzed by the model schemes, and mathematical formulations were developed. Various parameters associated with the exchange process, such as equilibrium exchange constant, number of surfactant molecules per substrate molecule, number of added anions, and a factor related to the binding of additives to the catalytic micellar aggregates, have been evaluated.

Notes: The kinetics and mechanism of oxidation of iminodiacetic acid and N-methyliminodiacetic acid by aquasilver(II) and Ag(II)-2,2′-bipyridine complexes has been investigated. The results are discussed with reference to the active reaction pathways, the equilibrium quotient of the title reactions, the protolytic equilibria which involve the oxidizing complex, and the intrinsic self-exchange rates of the oxidants.

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Notes: Recent experimental results on the thermal decomposition of N2O5 in N2 are evaluated in terms of unimolecular rate theory. A theoretically consistent set of fall-off curves is constructed which allows to identify experimental errors or misinterpretations. Limiting rate constants k0 = [N2] 2.2 × 10-3 (T/300)-4.4 exp(-11,080/T) cm3/molec·s over the range of 220-300 K, k∞ = 9.7 × 1014 (T/300)+0.1 exp(-11,080/T) s-1 over the range of 220-300 K, and broadening factors of the fall-off curve Fcent = exp(-T/250) + exp(-1050/T) over the range of 220-520 K have been derived. NO2 + NO3 recombination rate constants over the range of 200-300 K are krec,0 = [N2] 3.7 × 10-30 (T/300)-4.1 cm6/molec2·s and krec,∞ = 1.6 × 10-12 (T/300)+0.2 cm3/molec·s.

Notes: The kinetics of the gas-phase elimination of several chloroesters were determined in a static system over the temperature range of 410-490°C and the pressure range of 47-236 torr. The reactions in seasoned vessels, and in the presence of a free-radical inhibitor, are homogeneous, unimolecular, and follow a first-order law. The temperature dependence of the rate coefficients is given by the following Arrhenius equations: for methyl 3-chloropropionate, log k1(s-1) = (13.22 ± 0.07) - (231.5 ± 1.0) kJ/mol/2.303RT; for methyl 4-chlorobutyrate, log k1(s-1) = (13.31 ± 0.25) - (221.5 ± 3.4) kJ/mol/2.303RT; and for methyl 5-chlorovalerate, log k1(s-1) = (13.12 ± 0.25) - (221.7 ± 3.2) kJ/mol/2.303RT. Rate enhancements and lactone formation reveal the participation of carbonyl oxygen of the carbomethoxy group. The order COOCH3-5 〉 COOCH3-6 〉 COOCH3-4 in assistance is similar to the sequence of group participation in solvolysis reactions. The partial rates for the parallel eliminations to normal dehydrohalogenation products and lactones have been estimated and reported. The present results lead us to consider that an intimate ion-pair mechanism through participation of the carbomethoxy group may well be operating in some of these reactions.

Notes: Many experiments in chemical kinetics are initiated by a fast pulse, such as electric discharge, shock wave, flash lamp, or laser. After this pulse one observes the production and subsequent decay of a reactive intermediate. One then postulates a mechanism and adjusts the associated rate constants so as to minimize the difference between the results of the experiment and the prediction of the mechanism. The parameters to be estimated are usually strongly correlated, so that it is not possible to determine them separately. These estimated parameters are of little value unless we can also estimate statistically valid confidence limits for them. The difficulties are discussed which frequently arise in estimating parameters and confidence limits for a kinetic mechanism which is widely used in interpreting laser excitation and fluorescence measurements, that is, first-order production and decay. These difficulties, and methods for dealing with them, are illustrated with realistic data. The estimation problem is particularly ill conditioned when the production and loss rates are nearly equal. In some experimental systems this can be avoided, but in others it is inevitable.

Notes: A method is proposed whereby the orders and rate constants for processes obeying the rate law -dA/dt = kAn may be determined. The method is illustrated in two ways. First, simulated data for processes of various orders are treated, and the treatment is shown to be capable of reproducing orders and rate constants to a high degree of accuracy. The factors affecting the accuracy with which n and k can be determined are considered. These are inaccuracy in the determination of concentration values, irregularity of the time intevals between concentration determinations, and the length of those time intervals. It is shown that if concentrations are determined at times that are close together, the effect of the other two factors is small, but if the time intervals are made longer, the errors due to the other two factors affect the calculated values of n and k much more seriously. Second, the method was applied to two homogeneous reactions, of which one was first-order and one was second order, and three heterogeneous reactions, of which one was found by the original workers to be first order, one to be zero order, and one to vary between zero and first order, depending on the initial pressure. The present method gives results in agreement with these conclusions and reproduces the rate constants to within ±5% in all cases.

Notes: Results are reported from moderated nuclear recoil 18F experiments with the 273 K CHF3/C3F6/C2F6 system. Although the measurement sensitivity is only about ±12%, there is no evidence to support the occurrence of nonthermal F-to-HF reactions at 95 mol % C2F6 moderator concentration.

Notes: The main difference between the simple RRK theory and the better based but more complex RRKM theory is explained. Starting from the premise that the classical versus quantum mechanical estimation of the density of states is the major source of the difference, earlier attempts to incorporate the quantum effects in an effective value for the number of oscillators s are noted. By examining the expression for the RRKM rate coefficient it is found that a single effective s value will generally not suffice, but a much better representation of the quantum effects can be obtained if it is recognized that the problem inherently contains two different effective s values. A theory based on this analysis is constructed. It reproduces RRKM results to much improved accuracy, removing difficulties found earlier with single-s-value theories.

Notes: The thermal unimolecular decomposition of pent-2-yne has been studied over the temperature range of 988-1234 K using the technique of very low-pressure pyrolysis (VLPP). The main reaction pathway is C4—C5 bond fission producing the resonance-stabilized 3-methylpropargyl radical. There is a concurrent process producing molecular hydrogen and penta-1,2,4-triene presumably via the intermediate formation of cis-penta-1,3-diene. The 1,4-hydrogen elimination from cis-penta-1,3-diene is the rate-determining step in the molecular pathway. This is supported by an independent VLPP study of cis- and trans-penta-1,3-diene. RRKM calculations show that the experimental rate constants for C—C bond fission are consistent with the following high-pressure rate expression at 1100 K: \documentclass{article}\pagestyle{empty}\begin{document}$$\log k_1 = \left({s^{ - 1}} \right) = \left({16.0 \pm 0.3} \right) - \left({72.6 \pm 2.0} \right)/\theta $$\end{document} where θ = 2.303RT kcal/mol and the A factor was assigned from the results of shock-tube studies of related alkynes. The activation energy leads to ΔHf,3000[CH3C≡CĊH2] = 70.3 and DH3000[CH3CCCH2—H] = 87.4 kcal/mol. The resonance stabilization energy of the 3-methylpropargyl radical is 10.6 ± 2.5 kcal/mol, which is consistent with previous results for this and other propargylic radicals.

Notes: An iterative method has been devised for the simulation of chemiluminescence data during the oxidative decomposition of αα′ azobisisobutyronitrile in the presence of ethylbenzene. From this simulation the cross termination rate constant of the two types of peroxy radicals present has been estimated.

Notes: The kinetics of gamma-radiation-induced free-radical reactions in carbon tetrachloride solutions of ethanol and n-pentanol were studied in the range of 0.05-0.80M and 25-170°C. The rate constant for the reaction \documentclass{article}\pagestyle{empty}\begin{document}$${\rm CCl}_{\rm 3} + {\rm R} - CH_2 - {\rm OH}\mathop \to \limits^{k1} {\rm CHCl}_{\rm 3} + {\rm R} - {\rm CH} - {\rm OH}$$\end{document} was found as \documentclass{article}\pagestyle{empty}\begin{document}$$k1(M^{- 1} \cdot s^{- 1}) = 10^{8.6 \pm 0.4} \exp - (\frac{{9900 \pm 600{\rm cal}}}{{RT}})$$\end{document} The activation energy is larger by 0.8 kcal/mol than for secondary alcohols, while the A1 factors are about the same.

Notes: Dark-phase experiments between isoprene and O3 are discussed. UNC outdoor chamber experiments have shown that in high-concentration systems of isoprene and O3 (5 ppm C and 1 ppm) approximately 75% of the reacted carbon can be observed in the product formation of HCHO, CO, methacrolein, methylvinylketone, methylglyoxal, acetaldehyde, and propylene. Mechanisms were developed which gave reasonable fits to dark-phase chamber experiments of MACR, MVK, isoprene, and O3. Experimental data and modeling results were used to generate O3 rates of attack on MVK and MACR. An isoprene-O3 rate of 1.67 × 10-2 ppm-1·min-1 was used and is consistent with other rates reported in the literature. Dark isoprene-O3 systems appear to form homogeneously nucleated aerosol. Most of these particles appear and remain at diameters well below the optical cutoff region (0.3-0.5 μm), as opposed to the particles from similar α-pinene-O3 systems, which also form at smaller sizes but then grow into the optical size range (0.5 μm). Lower concentrations of α-pinene and O3 (0.2 ppm C and 0.12 ppm) still generated substantial aerosol, but by comparison, rapid CN nucleation was not observed during a similar side-by-side system of isoprene and O3.

Notes: Time-resolved absorption spectra for a reaction mixture of p-methoxystyrene and tetracyanoethylene (TCNE) are found to have a band maximum at 325 nm which is assigned to the 1,4-cycloadduct. The reaction in chloroform at 15, 20, and 25°C is followed by the charge-transfer band at 600 nm. The 1,4-cycloadduct, besides the so far known 1,2-cycloadduct and EDA complex, is taken into account to derive the rate equation for the EDA complex that is a linear second-order differential equation. The rate constants for the elementary steps involved in the reaction are obtained. The 1,4-cycloaddition has an activation entropy of -63 J/K·mol for the cycloreversion and a reaction constant ρ of -4.7, both of which indicate the polar transition state. On the other hand, activation entropy of the 1,2-cycloaddition is 73 J/K·mol more negative than that of the 1,4-cycloaddition, supporting the zwitterionic mechanism for the 1,2-cycloaddition.

Notes: The reaction of acetylene (A) with cyclohexa-1,3-diene (CHD) has been studied between 450 and 592 K. The pressures of A ranged from 25 to 112 torr and those of CHD from 8 to 62 torr. The reaction yields only ethene (E) and benzene (B) instead of bicyclo[2.2.2]octa-2,5-diene (BOD), the product that is expected for a 1,4,1′,2′ addition of the Diels-Alder type. It is first order with respect to each reagent. The rate constant (in L/mol·s) is given by \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} k = - (27,150 \pm 120)/4.576T + (7.49 \pm 0.05)$$\end{document} The thermal decomposition of BOD has also been studied. In the ranges of 354-435 K and 0.5-6 torr, the reaction is first order and results in the formation of equal amounts of B and E as the reaction of A with CHD does. Its rate constant (in s-1) is given by \documentclass{article}\pagestyle{empty}\begin{document}$$\log _{10} k_d = - (32,520 \pm 40)/4.576T + (14.06 \pm 0.02)$$\end{document} The following consecutive reactions are proposed for the reaction between A and CHD: where BOD is the primary product that is too unstable to be detected. This implies that the rate constant k is equal to ka. The reaction mechanisms and the strain energy in BOD are discussed.

Notes: Energetic hydrogen atoms generated by photolysis of HBr or HI react with CDCl3 by abstracting either a deuterium atom (1) or a chlorine atom (2): The integral probability of reaction (2) has been measured for several defined initial translational energies of H*, and the phenomenological threshold energy is 31 ± 14 kJ/mol. For initial translational energies in the range of 66-121 kJ/mol, the ratio of the integral probabilities of Cl abstraction and of D abstraction, when normalized to equal numbers of Cl and D atoms, is 2.4 ± 0.3. The interpretation of the integral reaction probabilities in terms of the excitation functions of reactions (1) and (2) is discussed. Measurements of the moderating effect of CO2 on reactions (1) and (2) show that CDCl3 is slightly more effective than CO2 as a moderator of H atoms in the energy range of 90-30 kJ/mol.

Notes: Results are reported from moderated nuclear recoil 18F experiments with the CH4/C3F6/C2F6 mixture system. At a 99.5% confidence level measurement precision of ±3.4%, non-thermal F-to-HF reactions are phenomenologically suppressed at C2F6 moderator concentrations in the range of 95.0-99.95 mol-%. Effectively equilibrium reaction conditions can be established in well-designed experiments of this type.

Notes: The reaction 2NO2 + ROH = RONO + HNO3 (R = CH3 or C2H5) has been studied using the FTIR method at reactant pressures from 0.1 to 1.0 torr at 25°C. The termolecular rate constant for the forward reaction was determined to be (5.7 ± 0.6) × 10-37 cm6/molec2·s for CH3OH and (5.7 ± 0.8) × 10-37 cm6/molec2·s for C2H5OH, that is, d[RONO]/dt = k[NO2]2[ROH]. The corresponding equilibrium constants were measured as 1.36 ± 0.06 and 0.550 ± 0.025 torr-1, respectively. These results are consistent with those of a previous study based on the NO2 decay measurements at reactant pressures from 1 to 10 torr.

Notes: The kinetics of oxidation of arginine, histidine, and threonine by chloramine-T (CAT) have been investigated in alkaline medium at 35°C. The rates are first order in both [CAT] and [amino acid] and inverse fractional order in [OH-] for arginine and histidine. The rate is independent of [OH-] for threonine. Variation of ionic strength and addition of the reaction product, p-toluenesulfonamide, or Cl- ions had no effect on the rate. A decrease of the dielectric constant of the medium by adding methanol decreased the rate with arginine, while the rates increased with histidine and threonine. The solvent isotope effect was studied using D2O. (kobs)D2O/(kobs)H2O was found to be 0.55 and 0.79 for arginine and histidine, respectively. The reactions were studied at different temperatures, and activation parameters have been computed. The oxidation process in alkaline medium, under conditions employed in the present investigations, has been shown to proceed via two paths, one involving the interaction of RNHCl (formed rapidly from RNCl-), with the amino acid in a slow step to form monochloroamino acid, which subsequently interacts with another molecule of RNHCl in a fast step to give the products, p-toluenesulfonamide (RNH2), and the corresponding nitrile of the amino acid (R'CN). The other path involves the interaction of RNCl- with the amino acid in a similar way to give RNH2 and R'CN. Mechanisms proposed and the derived rate laws are consistent with the observed kinetics. The rate constants predicted using the derived rate laws, as [OH-] varies, are in excellent agreement with the observed rate constants, thus justifying these rate laws and hence the proposed mechanistic schemes.

Notes: Rates of oxidation of XCOO- (X = H, D) by Br2 in acid aqueous media were measured between 274 and 332 K. The derived Arrhenius parameters for both reactions \documentclass{article}\pagestyle{empty}\begin{document}$$\log k_{\rm H} (M^{- 1} {\rm s}^{{\rm - 1}}) = (11.18 \pm 0.10) - (14.33 \pm 0.13)/\theta $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$\log k_{\rm D} (M^{- 1} {\rm s}^{{\rm - 1}}) = (13.77 \pm 0.13) - (17.62 \pm 0.04)/\theta $$\end{document} where θ = 4.575T × 10-3 kcal/mol, with (kH/kD)298K = 2.85, reveal a primary isotope effect, but the difference (ED - EH) = 3.29 kcal/mol and the ratio AD/AH = 91 fall beyond the limits imposed by semiclassical transition-state theory, suggesting tunneling or a multiple-stage mechanism. However, it can be shown that either tunneling in a single step or a three-step, internal return mechanism can be ruled out as alternative models, since both require unreasonable kinetic parameters to fit the data. The simplest scheme accounting for the present observations involves tunneling in the decomposition of a charge transfer complex in equilibrium with the reactants.

Notes: Kinetic investigations on the reaction between U(IV) and H2O2 have been carried out at different acidities in chloride medium at an ionic strength of 2M. The observed bimolecular rate constant has been found to be dependant on [H+]-1.3. The activation energy of the overall reaction has been found to vary from 13.4 ± 0.7 to 18.0 ± 0.8 kcal/mol in the range of acidity from 0.3 to 1.5M. The results have been explained on the basis of three parallel rate-controlling reactions involving unhydrolyzed species of U(IV) and hydrolyzed species UCl(OH)2+ and UO2+. The values of the rate constants for these three reaction paths have been found to be of the order of 3.95, 5.59 × 103, and 1.49 × 105M-1 min-1, respectively.

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

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

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

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

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

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