ALBERT

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

Notes: 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 heats of formation of C3 and C4 alkyl nitrites (RONO) have been determined via their heats of combustion by bomb calorimetry, thereby providing a complete set of values of ΔHºf for C1-C4 alkyl nitrites. The experimental values are in excellent agreement with values derived from group additivity rules. For branched compounds these calculations involve corrections for gauche interactions. In these cases, the gauche interactions are reflected in the activation energies E1 determined by recent kinetic studies, required for breaking the RO-NO bond. The heats of formation of the alkoxy radicals involved together with ΔHºf(NO) = 21.6 kcal/mole leads to the result D(RO-NO) = 41.5 ± 1 kcal/mole. The concordance between D(thermochemical) and D(kinetic), unlike previous kinetic studies, implies that E2 = 0 ± 1 kcal/mole.

Notes: The rate of decomposition of t-butyl nitrite (TBN) has been studied in a static system over the temperature range of 120-160°C. For low concentrations of TBN (10-5- 10-4M), but with a high total pressure of CF4 (∼0.9 atm) and small extents of reaction (∼1%), the first-order homogeneous rates of acetone (M2K) formation are a direct measure of reaction (1), since k3» k2 (NO): TBN . Addition of large amounts of NO in place of CF4 almost completely suppresses M2K formation. This shows that reaction (1) is the only route for this product. The rate of reaction (1) is given by k1 = 1016.3-40.3/θ s-1. Since (E1 + RT) and ΔH°1 are identical, both may be equated with D(RO-NO) = 40.9 ± 0.8 kcal/mole and E2 = O ± 1 kcal/mole. From ΔS°1 and A1, k2 is calculated to be 1010.4M-1 ·s-1, implying that combination of t—BuO and NO occurs once every ten collisions. From an independent observation that k2/k2′ = 1.7 ± 0.25 independent of temperature, it is concluded that k2′ = 1010.2M-1 · s-1 and k1′ = 1015.9-40.2/θ s-1; . This study shows that MeNO arises solely as a result of the combination of Me and NO. Since NO is such an excellent radical trap for t-Bu\documentclass{article}\pagestyle{empty}\begin{document}${\rm Me\dot O}$\end{document}, reaction (2) may be used in a competitive study of the decomposition of t—Bu\documentclass{article}\pagestyle{empty}\begin{document}${\rm Me\dot O}$\end{document} in order to obtain the first absolute value for k3. Preliminary results show that k3 (∞) = 1015.7-17.0/θ s-1. The pressure dependence of k3 is demonstrated over the range of 10-2-1 atm (160°C). The thermochemistry for reaction (3) implies that the Hg 6(3P1) sensitised decomposition of t-BuOH occurs via reaction (m): In addition to the products accounted for by the TBN radical split, isobutene is formed as a result of the 6-centre elimination of HONO: TBN \documentclass{article}\pagestyle{empty}\begin{document}$\mathop \to \limits^7 $\end{document} isobutene + HONO. The rate of formation of isobutene is given by k7 = 1012.9-33.6/θ s-1. t-BuOH, formed at a rate comparable to that of isobutene-at least in the initial stages-is thought to arise as a result of secondary reactions between TBN and HONO. The apparent discrepancy between this and previous studies is reconciled in terms of the above parallel reactions (1) and (7), such that k + 2k7 = 1014.7-36.2/θ s-1.

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.