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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: It is shown that it is possible to obtain good data for the rate constant for the decomposition of alkoxy radicals [RO] by using nitric oxide as a radical trap. Two experimental systems have been used.The first system involves the use of dialkyl peroxides [(RO)2] as thermal sources of alkoxy radicals. The peroxide concentration was ∼10-4M, nitric oxide ∼2 × 10-4M, and the extent of reaction was ∼10%. The total pressure was altered using carbon tetrafluoride as an inert gas. The mechanism is Hence R2/R3 = k2[NO]/k3. Our previous studies show that k2 lies in the range 1010.3±0.2M-1·sec-1.The second system employs alkyl nitrites [RONO] as a thermal source of alkoxy radicals. The experimental conditions are very similar, except that we chose to use an atmosphere of nitric oxide for initial experiments. If anything nitric oxide appears to be superior to carbon tetrafluoride as an energy transfer agent. The mechanism is Hence R3 = k1'k3[RONO]/(k3 + k2 + k6 [NO]).Results are given for R = t-Am, s-Bu, t-Bu, i-Pr, Et, and Me. In addition the first unequivocal evidence is given for the pressure dependence of k3 when R = t-Bu. The implications for atmospheric chemistry and combustion are also discussed.

Notes: By using isobutane (t-BuH) as a radical trapit has been possible to study the initial step in the decomposition of dimethyl peroxide (DMP) over the temperature range of 110-140°C in a static system. For low concentrations of DMP (2.5 × 10-5-10-4M) and high pressures of t-BuH (∼0.9 atm) the first-order homogeneous rate of formation of methanol (MeOH) is a direct measure of reaction (1): \documentclass{article}\pagestyle{empty}\begin{document}${\rm DMP}\mathop \to \limits^1 2{\rm Me}\mathop {\rm O}\limits^{\rm .},{\rm Me}\mathop {\rm O}\limits^{\rm .} + t{\rm - BuH}\mathop \to \limits^4 {\rm MeOH} + t{\rm -}\mathop {\rm B}\limits^{\rm .} {\rm u}$\end{document}. For complete decomposition of DMP in t-BuH, virtually all of the DMP is converted to MeOH. Thus DMP is a clean thermal source of Me\documentclass{article}\pagestyle{empty}\begin{document}$\mathop {\rm O}\limits^{\rm .}$\end{document}. In the decomposition of pure DMP complications arise due to the H-abstraction reactions of Me\documentclass{article}\pagestyle{empty}\begin{document}$\mathop {\rm O}\limits^{\rm .}$\end{document} from DMP and the product CH2O. The rate constant for reaction (1) is given by k1 = 1015.5-37.0/θ sec-1, very similar to other dialkyl peroxides. The thermochemistry leads to the result D(MeO—OMe) = 37.6 ± 0.2 kcal/mole and /H°f(Me\documentclass{article}\pagestyle{empty}\begin{document}$\mathop {\rm O}\limits^{\rm .}$\end{document}) = 3.8 ± 0.2 kcal/mole. It is concluded that D(RO—OR) and D(RO—H) are unaffected by the nature of R. From ΔS°1 and A1, k2 is calculated to be 1010.3±0.5 M-1· sec-1: \documentclass{article}\pagestyle{empty}\begin{document}$2{\rm Me}\mathop {\rm O}\limits^{\rm .} \mathop \to \limits^2 {\rm DMP}$\end{document}. For complete reaction, trace amounts of t-BuOMe lead to the result k2 ∼ 109 M-1 ·sec-1: \documentclass{article}\pagestyle{empty}\begin{document}$2t{\rm - Bu}\mathop \to \limits^5$\end{document} products. From the relationship k6 = 2(k2k5a)1/2 and with k5a = 108.4 M-1 · sec-1, we arrive at the result k6 = 109.7 M-1 · sec-1: \documentclass{article}\pagestyle{empty}\begin{document}$2t{\rm - u}\mathop {\rm B}\limits^{\rm .} \to (t{\rm - Bu)}_{\rm 2}{\rm,}t{\rm -}\mathop {\rm B}\limits^{\rm .} {\rm u} + {\rm Me}\mathop {\rm O}\limits^{\rm .} \mathop \to \limits^6 t{\rm - BuOMe}$\end{document}.

Notes: The rate of decomposition of s-butyl nitrite (SBN) has been studied in the absence (130-160°C) and presence (160-200°C) of NO. Under the former conditions, for low concentrations of SBN (6 × 10-5 - 10-4M) and small extents of reaction (∼1.5%), the first-order homogeneous rates of acetaldehyde (AcH) formation are a direct measure of reaction (1) since k3c » k2(NO): . Unlike t-butyl nitrite (TBN), d(AcH)/dt is independent of added CF4 (∼0.9 atm). Thus k3c is always » k2 (NO) over this pressure range. Large amounts of NO (∼0.9 atm) (130-160°C) completely suppress AcH formation. k1 = 1016.2-40.9/θ sec-1. Since (E1 + RT) and ΔH°1 are identical, within experimental error, both may be equated with D(s-BuO-NO) = 41.5 ± 0.8 kcal/mol and E2 = 0 ± 0.8 kcal/mol. The thermochemistry leads to the result ΔH°f (s-\documentclass{article}\pagestyle{empty}\begin{document}${\rm Bu}\mathop {\rm O}\limits^{\rm .}$\end{document}) = - 16.6 ± 0.8 kcal/mol. From ΔS°1 and A1, k2 is calculated to be 1010.4 M-1 · sec-1, identical to that for TBN. From an independent observation that k6/k2 = 0.26 ± 0.01 independent of temperature, \documentclass{article}\pagestyle{empty}\begin{document}${\rm s - Bu}\mathop {\rm O}\limits^{\rm .} + {\rm NO}\mathop \to \limits^{\rm 6} {\rm MEK} + {\rm HNO}$\end{document}, we find E6 = 0 ± 1 kcal/mol and k6 = 109.8M-1 · sec-1. Under the conditions first cited, methyl ethyl ketone (MEK) is also a product of the reaction, the rate of which becomes measurable at extents of conversion 〉2%. However, this rate is ∼0.1 that of AcH formation. Although MEK formation is affected by the ratio S/V for different reaction vessels, in a spherical reaction vessel, this MEK arises as the result of an essentially homogeneous first-order 4-centre elimination of HNO. \documentclass{article}\pagestyle{empty}\begin{document}${\rm SBN}\mathop \to \limits^{\rm 5} {\rm MEK} + {\rm HNO}$\end{document}; k5 = 1012.8-35.8/θ sec-1. Sec-butyl alcohol (SBA), formed at a rate comparable to MEK, is thought to arise via the hydrolysis of SBN, the water being formed from HNO. The rate of disappearance of SBN, that is, d(MEK + SBA + AcH)/dt, is given by kglobal = 1015.7-39.6/θ sec-1. In NO (∼1 atm) the rate of formation of MEK was about twice that in the absence of NO, whereas the SBA was greatly reduced. This reaction was also affected by the ratio S/V of different reaction vessels. It was again concluded that in a spherical reaction vessel, the rate of MEK formation was essentially homogeneous and first order. This rate is given by kobs = 1012.9-35.4/θ sec-1, very similar to k5. However, although it is clear that the rate of formation of MEK is doubled in the presence of NO, the value for kobs makes it difficult to associate this extra MEK with the disproportionation of s-\documentclass{article}\pagestyle{empty}\begin{document}${\rm Bu}\mathop {\rm O}\limits^{\rm .}$\end{document} and NO: s-\documentclass{article}\pagestyle{empty}\begin{document}$s{\rm - Bu}\mathop {\rm O}\limits^{\rm .} + {\rm NO}\mathop \to \limits^{\rm 6} {\rm MEK} + {\rm HNO}$\end{document}. NO at temperatures of 130-160°C completely suppresses AcH formation. AcH reappears at higher temperatures (165-200°C), enabling k3c to be determined. Ignoring reaction (6), d(AcH)/dt = k1k3 (SBN)/[k3c + k2(NO)]; k3c = 1014.8-15.3/θ sec-1. Inclusion of reaction (6) into the mechanism makes very little difference to the result. Reaction (3c) is expected to be a pressure-dependent process.

Notes: The rate of decomposition of isopropyl nitrite (IPN) has been studied in a static system over the temperature range of 130-160°C. For low concentrations of IPN (1-5 × 10-5M), but with a high total pressure of CF4 (∼0.9 atm) and small extents of reaction (∼1%), the first-order rates of acetaldehyde (AcH) formation are a direct measure of reaction (1), since k3 » k2(NO): \documentclass{article}\usepackage{amssymb}\pagestyle{empty}\begin{document}$ {\rm IPN}\begin{array}{rcl} 1 \\ {\rightleftarrows} \\ 2 \\ \end{array}i - \Pr \mathop {\rm O}\limits^. + {\rm NO},i - \Pr \mathop {\rm O}\limits^. \stackrel{3}{\longrightarrow} {\rm AcH} + {\rm Me}. $\end{document} Addition of large amounts of NO (∼0.9 atm) in place of CF4 almost completely suppressed AcH formation. Addition of large amounts of isobutane - t-BuH - (∼0.9 atm) in place of CF4 at 160°C resulted in decreasing the AcH by 25%. Thus 25% of \documentclass{article}\pagestyle{empty}\begin{document}$ i - \Pr \mathop {\rm O}\limits^{\rm .} $\end{document} were trapped by the t-BuH (4): \documentclass{article}\pagestyle{empty}\begin{document}$ i - \Pr \mathop {\rm O}\limits^. + t - {\rm BuH} \stackrel{4}{\longrightarrow} i - \Pr {\rm OH} + (t - {\rm Bu}). $\end{document} The result of adding either NO or t-BuH shows that reaction (1) is the only route for the production of AcH. The rate constant for reaction (1) is given by k1 = 1016.2±0.4-41.0±0.8/θ sec-1.Since (E1 + RT) and ΔH°1 are identical, within experimental error, both may be equated with D(i-PrO-NO) = 41.6 ± 0.8 kcal/mol and E2 = 0 ± 0.8 kcal/mol. The thermochemistry leads to the result that \documentclass{article}\pagestyle{empty}\begin{document}$ \Delta H_f^\circ (i - {\rm Pr}\mathop {\rm O}\limits^{\rm .} ) = - 11.9 \pm 0.8{\rm kcal}/{\rm mol}. $\end{document} From ΔS°1 and A1, k2 is calculated to be 1010.5±0.4M-1·sec-1. From an independent observation that k6/k2 = 0.19 ± 0.03 independent of temperature we find E6 = 0 ± 1 kcal/mol and k6 = 109.8+0.4M-;1·sec-1: \documentclass{article}\pagestyle{empty}\begin{document}$ i - \Pr \mathop {\rm O}\limits^. + {\rm NO} \stackrel{6}{\longrightarrow} {\rm M}_2 {\rm K} + {\rm HNO}. $\end{document}In addition to AcH, acetone (M2K) and isopropyl alcohol (IPA) are produced in approximately equal amounts. The rate of M2K formation is markedly affected by the ratio S/V of different reaction vessels. It is concluded that the M2K arises as the result of a heterogeneous elimination of HNO from IPN. In a spherical reaction vessel the first-order rate of M2K formation is given by k5 = 109.4-27.0/θ sec-1: \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm IPN} \stackrel{5}{\longrightarrow} {\rm M}_2 {\rm K} + {\rm HNO}. $\end{document} IPA is thought to arise via the hydrolysis of IPN, the water being formed from HNO. This elimination process explains previous erroneous results for IPN.

Notes: The rate of decomposition of methyl nitrite (MN) has been studied in the presence of isobutane-t-BuH-(167-200°C) and NO (170-200°C). In the presence of t-BuH (∼0.9 atm), for low concentrations of MN (∼10-4M) and small extents of reaction (4-10%), the first-order homogeneous rates of methanol (MeOH) formation are a direct measure of reaction (1) since k4(t-BuH) »k2(NO): . The results indicate that the termination process involves only \documentclass{article}\pagestyle{empty}\begin{document}$ t - {\rm Bu\, and\, NO:\,\,}t - {\rm Bu} + {\rm NO\stackrel{e}{\longrightarrow}} $\end{document} products, such that ke ∼ 1010 M-1 ∼ sec-1.Under these conditions small amounts of CH2O are formed (3-8% of the MeOH). This is attributed to a molecular elimination of HNO from MN. The rate of MeOH formation shows a marked pressure dependence at low pressures of t-BuH. Addition of large amounts of NO completely suppresses MeOH formation.The rate constant for reaction (1) is given by k1 = 1015.8°0.6-41.2°1/· sec-1. Since (E1 + RT) and ΔHΔ1 are identical, within experimental error, both may be equated with D(MeO - NO) = 41.8 + 1 kcal/mole and E2 = 0 ± 1 kcal/mol. From ΔS11 and A1, k2 is calculated to be 1010.1°0.6M-1 · sec-1, in good agreement with our values for other alkyl nitrites. These results reestablish NO as a good radical trap for the study of the reactions of alkoxyl radicals in particular. From an independent observation that k6/k2 = 0.17 independent of temperature, we conclude that \documentclass{article}\pagestyle{empty}\begin{document}$ E_6 = 0 \pm 1{\rm kcal}/{\rm mol\, and\,}\,k_6 = 10^{9.3} M^{- 1} \cdot {\rm sec}^{- 1} :{\rm MeO} + {\rm NO}\stackrel{6}{\longrightarrow}{\rm CH}_2 {\rm O} + {\rm HNO} $\end{document}. From the independent observations that k2:k2→: k6→ was 1:0.37:0.04, we find that k2→ = 109.7M-1 ċ sec-1 and k6→ = 108.7M-1 ċ sec-1. In addition, the thermodynamics lead to the result In the presence of NO (∼0.9 atm) the products are CH2O and N2O (and presumably H2O) such that the ratio N2O/CH2O ∼ 0.5. The rate of CH2O formation was affected by the surface-to-volume ratio s/v for different reaction vessels, but it is concluded that, in a spherical reaction vessel, the CH2O arises as the result of an essentially homogeneous first-order, fourcenter elimination of \documentclass{article}\pagestyle{empty}\begin{document}$ {\rm HNO}:{\rm MN\stackrel{5}{\longrightarrow}CH}_{\rm 2} {\rm O} + {\rm HNO} $\end{document}. The rate of CH2O formation is given by k5 = 1013.6°0.6-38.5-1/ċ sec-1.

Notes: The reaction of methyl radicals (Me) with hexafluoroacetone (HFA), generated from ditertiary butyl peroxide (dtBP), was studied over the temperature range of 402-433 K and the pressure range of 38-111 torr. The reaction resulted in the following displacement process taking place: where TFA refers to trifluoroacetone. The trifluoromethyl radicals that were generated abstract a hydrogen atom from the peroxide: such that k6a is given by: \documentclass{article}\pagestyle{empty}\begin{document}$$ {\rm log}\,\,k_{6{\rm a}} \left( {M^{ - 1} \cdot {\rm s}^{ - 1} } \right)\,\, = \,\,8.2\, \pm \,0.7\, - \,8.9\, \pm \,{{1.3} \mathord{\left/ {\vphantom {{1.3} \theta }} \right. \kern-\nulldelimiterspace} \theta } $$\end{document} where θ = 2.303RT kcal/mol. The interaction of methyl and trifluoromethyl radicals results in the following steps: Product analysis shows that k17/k151/2k161/2 = 2.0 ± 0.2 such that k17 = 1010.4±0.5M-1 · s-1. The rate constant k5 is given by: \documentclass{article}\pagestyle{empty}\begin{document}$$ {{{\rm log}\,\,k_5 } \mathord{\left/ {\vphantom {{{\rm log}\,\,k_5 } {\left( {M^{ - 1} \cdot {\rm s}^{ - 1} } \right)}}} \right. \kern-\nulldelimiterspace} {\left( {M^{ - 1} \cdot {\rm s}^{ - 1} } \right)}}\,\, = \,\,8.0\, \pm \,1.3\, - \,5.1\, \pm \,{{0.7} \mathord{\left/ {\vphantom {{0.7} \theta }} \right. \kern-\nulldelimiterspace} \theta } $$\end{document} It is concluded that the preexponential factor for the addition of methyl radicals to ketones is lower than that for the addition of methyl radicals to olefins.

Notes: By allowing the t-butoxy radical to decompose in the presence of nitric oxide, it has been possible to determine rate constants for decomposition by the measurements of the relative rates (2) and (3) Process (3) is clearly pressure dependent. The value of k3(∞) has been determined in the presence of several inert gases (CF4, SF6, N2, and Ar) and a value of k3 interpolated for atmospheric conditions. The results may be compared with those for other relevant alkoxy radicals at room temperature. Extrapolated values for k3 in the presence of CF4 lead to the result \documentclass{article}\pagestyle{empty}\begin{document}$$ k_3 (\infty)/s^{ - 1} = 10^{14.6 \pm 0.6} \exp (- 8052 \pm 604/T) $$\end{document}

Notes: The near U-V photolysis of t-butyl nitrite has been studied over the temperature range 303-393 K. Under these conditions t-butyl nitrite was shown to be a very clean photochemical source of t-butoxy radicals. This allows a study of the decomposition of the t-butoxy radical to be made over this temperature range (3). Extrapolation of the rate constants k3 to high pressure and combination with our previous thermal data give the results: \documentclass{article}\pagestyle{empty}\begin{document}$$k_3 (\infty)/{\rm s}^{{\rm - 1}} = 10^{14.04 \pm 0.37} \exp (- 7519 \pm 70.5/T)$$\end{document}