Rate constants for the reactions of molecular iodine with Cl,SiCl3, and SiH3 at 298 K

Rate constants k₁, k₂, and k₃ have been measured at 298 K by means of a laser photolysis-laser magnetic resonance technique and (or) by a laser photolysis-infrared chemiluminescence detection technique (LMR and IRCL, respectively). \hfill\hbox to 12em{$\rm Cl+I_2\longrightarrow ICl+I;$}\hbox to 8em{$\rm {\it k}_1=(2.5\pm 0.7)\times 10^{-10}(IRCL)$}\hfill(1)\\\hfill\hbox to 12em{}\hbox to 8em{$\rm {\it k}_1=(2.8\pm 0.8)\times 10^{-10}(LMR)$}\hfill \\\hfill\hbox to 12em{$\rm SiCl_3+I_2\longrightarrow SiCl_3I+I;$}\hbox to 8em{$\rm {\it k}_2=(5.8\pm 1.8)\times 10^{-10}(IRCL)$}\hfill (2)\\\hfill\hbox to 12em{$\rm SiH_3+I_2\longrightarrow SiIH_3+I;$}\hbox to 8em{$\rm {\it k}_3=(1.8\pm 0.46)\times 10^{-10}(LMR)$}\hfill (3)\\ As an average of the LMR and IRCL results we offer the value k₁ = (2.7 ± 0.6) × 10⁻¹⁰. Units are cm³ molecule⁻¹ s⁻¹; uncertainties are 2σ including precision and estimated systematic errors. © 1997 John Wiley & Sons, Inc. Int J Chem Kinet 29: 25–33, 1997.     

A Modified Kinetic Model for the Peroxydisulfate-Iodide Reaction

One kinetic model for the oxidation of iodide ion by peroxydisulfate ion in aqueous solution is proposed. The reaction is regarded as \documentclass{article}\pagestyle{empty}\begin{document} {\rm S}_2 {\rm O}_8^{2 -} + {\rm I}^ - {\rm IS}_2 {\rm O}_8^{3 -} \end{document}, followed by the reaction \documentclass{article}\pagestyle{empty}\begin{document} {\rm IS}_2 {\rm O}_8^{3 -} + {\rm I}l_2 + 2{\rm SO}_4^{2 -} \end{document}. If the initial rates V are obtained from the formation of the iodine molecules, the reaction rate constant k₁ and the ratio k₂/k_-1 can be estimated by plotting the values of [S₂O₈^2?][I^?]/V against that of 1/[I^?]. The extrapolated value for k₁ is 2.20×10^?2 L/mol-sec and k₂/k_-1 is calculated to be 4.25×10² mol/L at 27°C in a solution with an ionic strength of 0.420.     

Temperature dependence of europium carbonate complexation

The temperature dependencies of europium carbonate stability constants were examined at 15, 25, and 35°C in 0.68 molal Na⁺(ClO ₄ ^? , HCO ₃ ^? ) using a tributyl phosphate solvent extration technique. Our distribution data can be explained by the equilibria $$\begin{gathered} Eu^{3 + } + H_2 O + CO_2 (g)_ \leftarrow ^ \to EuCO_3^ + + 2H^ + \hfill \\ - log\beta _{12} = 9.607 + 496(t + 273.16)^{ - 1} \hfill \\ Eu^{3 + } + 2H_2 O + 2CO_2 (g)_ \leftarrow ^ \to Eu(CO_3 )_2^ - + 4H^ + \hfill \\ - log\beta _{24} = 21.951 + 670(t + 273.16)^{ - 1} \hfill \\ Eu^{3 + } + H_2 O + CO_2 (g)_ \leftarrow ^ \to EuHCO_3^{2 + } + H^ + \hfill \\ - log\beta _{11} = 1.688 + 1397(t + 273.16)^{ - 1} \hfill \\ \end{gathered}$$      

The gas-phase pyrolysis of alkyl nitrites. IV. Ethyl nitrite

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}} $\end{document} products, such that k_e ～ 10¹⁰ M^?1 ～ sec^?1.Under these conditions small amounts of CH₂O are formed (3-8% of the MeOH). This is attributed to a molecular elimination of HNO from MN. The rate of MeOH formation shows a marked pressure dependence at low pressures of t-BuH. Addition of large amounts of NO completely suppresses MeOH formation. The rate constant for reaction (1) is given by k₁ = 10^{15.8°0.6-41.2°1}/· sec^?1. Since (E₁ + RT) and ΔH^Δ₁ are identical, within experimental error, both may be equated with D(MeO - NO) = 41.8 + 1 kcal/mole and E₂ = 0 ± 1 kcal/mol. From ΔS¹₁ and A₁, k₂ is calculated to be 10^10.1°0.6M^?1 · sec^?1, in good agreement with our values for other alkyl nitrites. These results reestablish NO as a good radical trap for the study of the reactions of alkoxyl radicals in particular. From an independent observation that k₆/k₂ = 0.17 independent of temperature, we conclude that \documentclass{article}\pagestyle{empty}\begin{document}$ 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} $\end{document}. From the independent observations that k₂:k_2→: k_6→ was 1:0.37:0.04, we find that k_2→ = 10^9.7M^?1 ? sec^?1 and k_6→ = 10^8.7M^?1 ? sec^?1. In addition, the thermodynamics lead to the result In the presence of NO (～0.9 atm) the products are CH₂O and N₂O (and presumably H₂O) such that the ratio N₂O/CH₂O ～ 0.5. The rate of CH₂O formation was affected by the surface-to-volume ratio s/v for different reaction vessels, but it is concluded that, in a spherical reaction vessel, the CH₂O arises as the result of an essentially homogeneous first-order, fourcenter elimination of \documentclass{article}\pagestyle{empty}\begin{document}$ {\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 = 10^{13.6°0.6-38.5-1}/? sec^?1.     

The pK 2 * for the dissociation of H2SO3 in NaCl solutions with added Ni2+, Co2+, Mn2+, and Cd2+ at 25°C

The pK ₂ ^* for the dissociation of sulfurous acid from I=0.5 to 6.0 molal at 25°C has been determined from emf measurements in NaCl solutions with added concentrations of NiCl₂, CoCl₂, McCl₂ and CdCl₂ (m=0.1). These experimental results have been treated using both the ion pairing and Pitzer's specific ion-interaction models. The Pitzer parameters for the interaction of M²⁺ with SO ₃ ^2? yielded $$\begin{gathered} \beta _{NiSO_3 }^{(0)} = - 5.5, \beta _{NiSO_3 }^{(1)} = 5.8, and \beta _{NiSO_3 }^{(2)} = - 138 \hfill \\ \beta _{CoSO_3 }^{(0)} = - 12.3, \beta _{CoSO_3 }^{(1)} = 31.6, and \beta _{CoSO_3 }^{(2)} = - 562 \hfill \\ \beta _{MnSO_3 }^{(0)} = - 8.9, \beta _{MnSO_3 }^{(1)} = 18.7, and \beta _{MnSO_3 }^{(2)} = - 353 \hfill \\ \beta _{CdSO_3 }^{(0)} = - 7.2, \beta _{CdSO_3 }^{(1)} = 13.8, and \beta _{CdSO_3 }^{(2)} = - 489 \hfill \\ \end{gathered} $$ The calculated values of pK ₂ ^* using Pitzer's equations reproduce the measured values to within ±0.01 pK units. The ion pairing model yielded $$\begin{gathered} logK_{NiSO_3 } = 2.88 and log\gamma _{NiSO_3 } = 0.111 \hfill \\ logK_{CoSO_3 } = 3.08 and log\gamma _{CoSO_3 } = 0.051 \hfill \\ logK_{MnSO_3 } = 3.00 and log\gamma _{MnSO_3 } = 0.041 \hfill \\ logK_{CdSO_3 } = 3.29 and log\gamma _{CdSO_3 } = 0.171 \hfill \\ \end{gathered} $$ for the formation of the complex MSO₃. The stability constants for the formation of MSO₃ complexes were found to correlate with the literature values for the formation of MSO₄ complexes.     

LiSO 4 − , RbSO 4 − , CsSO 4 − , and (NH4)SO 4 − ion pairs in aqueous solutions at pressures up to 2000 atm

Electrical conductance data at 25°C for Li₂SO₄, Rb₂SO₄, Cs₂SO₄, and (NH₄)₂SO₄ aqueous solutions are reported at concentrations up to 0.01 eq.-liter^?1 and as a function of pressure up to 2000 atm. The molal dissociation constants are as follows: $$\begin{gathered} LiSO_4^ - : - log K_m = - 1.02 + 1.03 \times 10^4 P \pm 0.019 \Delta \bar V^o = - 5.8 \hfill \\ RbSO_4^ - : - log K_m = - 1.12 + 0.58 \times 10^4 P \pm 0.020 \Delta \bar V^o = - 3.3 \hfill \\ CsSO_4^ - : - log K_m = - 1.08 + 1.10 \times 10^4 P \pm 0.014 \Delta \bar V^o = - 6.2 \hfill \\ \left( {NH4} \right)SO_4^ - : - log K_m = - 1.12 + 0.58 \times 10^4 P \pm 0.020 \Delta \bar V^o = - 3.3 \hfill \\ \end{gathered} $$ whereP is in atmospheres and \(\Delta \bar V^o \) is in cm³-mole^?1. These values were obtained by using the Davies-Otter-Prue conductance equation and Bjerrum distance parameters. A simultaneous Λ°,K _m search was used to determine the equilibrium constantK _m, a different procedure than used earlier for KSO ₄ ^? , NaSO ₄ ^? , and MgCl⁺. Recalculated values for these salts are as follows: $$\begin{gathered} KSO_4^ - : - log K_m = - 1.03 + 1.04 \times 10^4 P \pm 0.020 \Delta \bar V^o = - 5.9 \hfill \\ NaSO_4^ - : - log K_m = - 1.00 + 1.30 \times 10^4 P \pm 0.019 \Delta \bar V^o = - 7.3 \hfill \\ MgCl^ + : - log K_m = - 0.75 + 0.71 \times 10^4 P \pm 0.028 \Delta \bar V^o = - 4.0 \hfill \\ \end{gathered} $$      

The kinetics of the chlorate-bisulfite and chlorate-sulfite/bisulfite reaction systems

The kinetics of the reactions of ClO₃ ⁻ with HSO₃ ⁻ and H₂SO₃ was studied by measuring the concentration of [Cl⁻] and [H⁺] both in chlorate-bisulfite and chlorate-sulfite/bisulfite solutions. A reaction mechanism was applied for simulation of the experimental observations. Rate constants k₁ = (1±0.5)·10⁻⁴ M⁻¹ s⁻¹ and k₂ = (0.23±0.01) M⁻¹ s⁻¹ were determined for the following reactions:

Effect of the excluded volume on specific rate of combination reactions between polymer radicals

The specific rate k_D for reaction between polymer radicals is formulated when the potential of average force on the basis of the excluded volume affects the motion of the polymer radicals. This rate is given by \documentclass{article}\pagestyle{empty}\begin{document}$ k_D = Fk_S \left( {{\rm with}\ {F} = \sum\limits_{s = 0}^\infty {{{[ ‐ 2(\alpha ^2 ‐ 1)]} \mathord{\left/ {\vphantom {{[ ‐ 2(\alpha ^2 ‐ 1)]} {(s + 1)^{{3 \mathord{\left/ {\vphantom {3 2}} \right. \kern‐\nulldelimiterspace} 2}} }}} \right. \kern‐\nulldelimiterspace} {(s + 1)^{{3 \mathord{\left/ {\vphantom {3 2}} \right. \kern‐\nulldelimiterspace} 2}} }}} } \right) $\end{document} where k_S is specific rate of reaction between radical chain ends and α is the average expansion of the polymer arising from the long-range effects. The effect of the excluded volume reduces k_D. F depends on the degree of polymerization of the polymer radical when α ≠ 1. These results are discussed in terms of the experimental data for very low polymer concentrations.     

Iodine Hydrolysis Equilibrium

The equilibrium constant for the hydrolytic disproportionation of I₂

Tautomerism and Atropisomerism in Free‐Base (meso)‐Strapped Porphyrins: Static and Dynamic Aspects

We report herein some outstanding examples of atropisomerism and tautomerism in five (meso‐)strapped porphyrins. Porphyrins S0 – S4 have been synthesised, characterised and studied in detail by spectroscopic and spectrometric techniques, and their isomeric purity verified by HPLC analysis. In particular, they exhibit perfectly well‐defined NMR spectra that display distinct patterns depending on their average symmetry at room temperature: C_2v, D_2d, C_2h, C_2v, and D_2h for S0 – S4 , respectively. NH tautomerism was evidenced by variable‐low‐temperature ¹H NMR experiments in [D₂]dichloromethane performed on S0 (Δ${G{{{\ne}\hfill \atop {\rm 298K}\hfill}}}$ =48±1 kJ mol^?1) and S1 (Δ${G{{{\ne}\hfill \atop {\rm 298K}\hfill}}}$ =55±3 kJ mol^?1), which has led to an understanding of the average spectra observed for the five porphyrins at room temperature. On the other hand, S2 and S3 are stable atropisomers at room temperature, easily separated and characterised, as a result of restricted rotation of their strapped bridges due to their high rotational barrier energies. Upon heating to 82 °C, they slowly equilibrate to a thermodynamic ratio of 64:36 in favour of the more stable S2 isomer. This atropisomerisation process was evidenced by ¹H NMR spectroscopy and monitored by HPLC, from which high rotational energy barriers of 115.2 (Δ${G{{{\ne}\hfill \atop {\rm S2}\rightarrow {\rm S3}\hfill}}}$ ) and 116.9 kJ mol^?1 (Δ${G{{{\ne}\hfill \atop {\rm S2}\rightarrow {\rm S3}\hfill}}}$ ) were deduced.     

Pyrolysis of dimethyl peroxide

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 = 10^15.5?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(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 and A₁, k₂ is calculated to be 10^10.3±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(k₂k_5a)^1/2 and with k_5a = 10^8.4 M^?1 · sec^?1, we arrive at the result k₆ = 10^9.7 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}.     

Critical evaluation and experimental determination of the nuclear activation and decay parameters for the reactions:50Cr(n, ψ)51Cr,58Fe(n, ψ)59Fe,109Ag(n, ψ)110mAg

Isotopic abundance values for⁵⁰Cr,⁵⁸Fe and¹⁰⁹Ag and the absolute gamma-intensities for⁵¹Cr,⁵⁹Fe and^110mAg were evaluated. These evaluated data, together with experimental k₀-determinations (i.e. from the “activation method”), made it possible to calculate the following 2200 m.s^?1 cross-sections, which considerably deviate from the hitherto generally published ones [between brackets]: $$\begin{gathered} {}^{5 0}Cr(n,\gamma )^{5 1} Cr; \sigma _0 = (15.2 \pm 0.2) barn [cf.:15.8 - 16.0] \hfill \\ {}^{5 8}Fe(n,\gamma )^{5 9} Fe; \sigma _0 = (1.31 \pm 0.03) barn [cf.:1.14 - 1.16] \hfill \\ {}^{1 0 9}Ag(n,\gamma )^{1 1 0 m} Ag;\sigma _0 = (3.89 \pm 0.05) barn [cf.:4.4 - 5.0] \hfill \\ \end{gathered} $$      

Thermodynamic quantities for the interaction of SO 4 2− with H+ and Na+ in aqueous solution from 150 to 320°C

The aqueous reactions,

Kinetic and equilibrium studies of the acid dissociation of the copper(II) and nickel(II) complexes of 5,7-dioxo-1,4,8,11-tetra-azacyclotetradecane

The acid catalysed dissociation of the copper(II) and nickel(II) complexes of 5,7-dioxo-1,4,8,11-tetra-azacyclo-tetradecane (dioxocyclam = LH₂) has been studied using nitric acid solutions over a range of temperatures at I = 1.0 mol dm si-³. The kinetic data for the copper complex can be fitted to the rate expression k _obs = k K ₂[H⁺]/(1 + K ₂ [H⁺] with K = 24.7s⁻¹ and K ₂ = 65dm³mol⁻¹ at 25° C. The analogous constants for the nickel(II) complex are K = ³.³s⁻¹ and K ₂ = 45dm³mol⁻¹. The acid dissociation can be rationalized in terms of the kinetic scheme

Dynamics of magnesium-bicarbonate interactions

The complexation kinetics of Mg²⁺ with CO ₃ ⁼ and HCO ₃ ^? has been studied in methanol and water by means of the stopped-flow and temperature-jump methods. Kinetic parameters were obtained in methanol by coupling the magnesium-carbonato reactions with the metal-ion indicator Murexide. Relatively high stability constants were found in methanol (K=1.0×10⁵ liters-mole^?1 for Mg²⁺-Murexide,K=7.0×10⁴ liters-mole^?1 for Mg²⁺?HCO ₃ ^? , andK=2.0×10⁵ for Mg²⁺?CO ₃ ⁼ liters-mole^?1). The corresponding, observed formation rate constants were determined to be $$\begin{gathered} k_f = 4.0 \times 10^6 M^{ - 1} - sec^{ - 1} (Mg^{2 + } - Murexide) \hfill \\ k_f = 5.0 \times 10^5 M^{ - 1} - sec^{ - 1} (Mg^{2 + } - HCO_3^ - ) \hfill \\ k_f = 6.8 \times 10^5 M^{ - 1} - sec^{ - 1} (Mg^{2 + } - CO_3^ = ) \hfill \\ \end{gathered} $$ The relaxation times were found to be much shorter (τ≈5–20 μsec) in aqueous solutions, primarily due to the relatively high dissociation rate constants. The data could be interpreted on the basis of a coupled reaction scheme in which the protolytic equilibria are established relatively rapidly, followed by a single relaxation process due to the formation of MgHCO ₃ ⁺ and MgCO₃ between pH 8.7 and 9.3. The observed formation rate constants were determined to be $$\begin{gathered} k_f = 5.0 \times 10^5 M^{ - 1} - sec^{ - 1} (Mg^{2 + } - HCO_3^ - ) \hfill \\ k_f = 1.5 \times 10^6 M^{ - 1} - sec^{ - 1} (Mg^{2 + } - CO_3^ = ) \hfill \\ \end{gathered} $$ These results, in conjunction with NMR solvent exchange rate constants, are analyzed in terms of a dissociative (S _N1) mechanism for the rate of complex formation. The significance of these kinetic parameters in understanding the excess sound absorption in seawater is discussed.     

Laserflash-photolysis of the p-Chloranil/naphthalene system: Characterization of the naphthalene radical cation in a fluid medium,

Excitation of p-Chloranil ( CA ) in propylcyanide (PrCN) at room temperature leads to rapid production of ³ CA * which decays predominantly to CA H· with k = 1.6 · 10⁵ s^?1. Observation of a photoinduced current suggests simultaneous production of CA ^? formed by electron transfer quenching of ³ CA * by the medium. Added naphthalene ( NP ) quenches ³ CA * with k_q = 7.0 · 10⁹M ^?1S ^?1; NP ⁺ is unambigously identified as product (besides CA ^?) of the electron transfer process. Dissociation of the ion pair occurs with essentially unit probability. Higher concentrations of NP lead to the formation of ( NP )⁺₂. Pertinent spectroscopic parameters established for NP ⁺ under the conditions used are λ_max = 685 nm (? = 2970) using the known parameters of CA \documentclass{article}\pagestyle{empty}\begin{document}$ 1^{+ \atop \dot{}} $\end{document} as reference. NP ⁺ and CA \documentclass{article}\pagestyle{empty}\begin{document}$ 1^{+ \atop \dot{}} $\end{document} decay by charge annihilation with k_r = 4.5 · 10⁹ M ^?1S ^?1. The deviation from the diffusion controlled rate constant expected for ionic species, is discussed in view of the spin characteristics of the process. Comparison with two other ion recombination reactions leads to the conclusion that ‘inverted behaviour’ as expected from Marcus' theory does also not show up for backward e^?-transfer between two ions (produced by forward e^?-transfer between two neutrals). Residual absorptions in the system are ascribed to CA H·, tentatively proposed to originate from H⁺-abstraction by CA \documentclass{article}\pagestyle{empty}\begin{document}$ 1^{+ \atop \dot{}} $\end{document} from the solvent. NP ⁺ appears to be a rather stable species with respect to the medium if the latter is meticulously purified.     

The gas-phase pyrolysis of alkyl nitrites. III. Isopropyl nitrite

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}). $\end{document} The result of adding either NO or t-BuH shows that reaction (1) is the only route for the production of AcH. The rate constant for reaction (1) is given by k₁ = 10^{16.2±0.4–41.0±0.8}/θ sec^?1. Since (E₁ + RT) and ΔH°₁ are identical, within experimental error, both may be equated with D(i-PrO-NO) = 41.6 ± 0.8 kcal/mol and E₂ = 0 ± 0.8 kcal/mol. The thermochemistry leads to the result that \documentclass{article}\pagestyle{empty}\begin{document}$ \Delta H_f^\circ (i - {\rm Pr}\mathop {\rm O}\limits^{\rm .} ) = - 11.9 \pm 0.8{\rm kcal}/{\rm mol}. $\end{document} From ΔS°₁ and A₁, k₂ is calculated to be 10^10.5±0.4M^?1·sec^?1. From an independent observation that k₆/k₂ = 0.19 ± 0.03 independent of temperature we find E₆ = 0 ± 1 kcal/mol and k₆ = 10^9.8+0.4M^?;1·sec^?1: \documentclass{article}\pagestyle{empty}\begin{document}$ i - \Pr \mathop {\rm O}\limits^. + {\rm NO} \stackrel{6}{\longrightarrow} {\rm M}_2 {\rm K} + {\rm HNO}. $\end{document} In addition to AcH, acetone (M₂K) and isopropyl alcohol (IPA) are produced in approximately equal amounts. The rate of M₂K formation is markedly affected by the ratio S/V of different reaction vessels. It is concluded that the M₂K arises as the result of a heterogeneous elimination of HNO from IPN. In a spherical reaction vessel the first-order rate of M₂K formation is given by k₅ = 10^9.4–27.0/θ sec^?1: \documentclass{article}\pagestyle{empty}\begin{document}$ {\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.