The very-low-pressure pyrolysis (VLPP) of n-propyl nitrate,tert-butyl nitrite,and methyl nitrite. Rate constants for some alkoxy radical reactions

The pyrolysis of n-propyl nitrate and tert-butyl nitrite at very low pressures (VLPP technique) is reported. For the reaction the high-pressure rate expression at 300°K, log k₁ (sec^?1) = 16.5 ? 40.0 kcal/mole/2.3 RT, is derived. The reaction was studied and the high-pressure parameters at 300°K are log k₂(sec^?1) = 15.8 ? 39.3 kcal/mole/2.3 RT. From ΔS_1,?1⁰ and ΔS_2,?2⁰ and the assumption E_?1 and E_?2 ? 0, we derive log k_?1(M^?1·sec^?1) (300°K) = 9.5 and log k_?2 (M^?1·sec^?1) (300°K) = 9.8. In contrast, the pyrolysis of methyl nitrite and methyl d₃ nitrite afford NO and HNO and DNO, respectively, in what appears to be a heterogeneous process. The values of k_?1 and k_?2 in conjunction with independent measurements imply a value at 300°K for of 3.5 × 10⁵ M^?1·sec^?1, which is two orders of magnitude greater than currently accepted values. In the high-pressure static pyrolysis of dimethyl peroxide in the presence of NO₂, the yield of methyl nitrate indicates that the combination of methoxy radicals with NO₂ is in the high-pressure limit at atmospheric pressure.     

KINETICS OF FORMATION AND STABILITY CONSTANTS OF MONO AND BIS l-TYROSINE COMPLEXES OF Cu(II), Co(II), AND Ni(II)

Abstract

The kinetics and stability constants of l-tyrosine complexation with copper(II), cobalt(II) and nickel(II) have been studied in aqueous solution at 25° and ionic strength 0.1 M. The reactions are of the type M(HL)^(3-n)+ _n-1 + HL- ? M(HL)^(2-n)+_n(k_n, forward rate constant; k_-n, reverse rate constant); where M=Cu, Co or Ni, HL^? refers to the anionic form of the ligand in which the hydroxyl group is protonated, and n=1 or 2. The stability constants (K_n=k_n/k_-n) of the mono and bis complexes of Cu²⁺, Co²⁺ and Ni²⁺ with l-tyrosine, determined by potentiometric pH titration are: Cu²⁺, log K₁=7.90 ± 0.02, log K₂=7.27 ± 0.03; Co²⁺, log K₁=4.05 ± 0.02, log K₂=3.78 ± 0.04; Ni²⁺, log K₁=5.14 ± 0.02, log K₂=4.41 ± 0.01. Kinetic measurements were made using the temperature-jump relaxation technique. The rate constants are: Cu²⁺, k₁=(1.1 ± 0.1) × 10⁹ M ^?1 sec^?1, k_-1=(14 ± 3) sec^?1, k₂=(3.1 ± 0.6) × 10⁸ M ^?1 sec^?1, k^?2=(16 ± 4) sec^?1; Co²⁺, k₁=(1.3 ± 0.2) × 10⁶ M ^?1 sec^?1, k_-1=(1.1 ± 0.2) × 10² sec^?1, k₂=(1.5 ± 0.2) × 10⁶ M ^?1 sec^?1, k_-2=(2.5 ± 0.6) × 10² sec^?1; Ni²⁺, k₁=(1.4 ± 0.2) × 10⁴ M ^?1 sec^?1, k_-1=(0.10 ± 0.02) sec^?1, k₂=(2.4 ± 0.3) × 10⁴ M ^?1 sec^?1, k_-2=(0.94 ± 0.17) sec^?1. It is concluded that l-tyrosine substitution reactions are normal. The presence of the phenyl hydroxyl group in l-tyrosine has no primary detectable influence on the forward rate constant, while its influence on the reverse rate constant is partially attributed to substituent effects on the basicity of the amine terminus.     

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.     

Cationic polymerization of N-vinyl carbazole by trityl salts—I.: Reactivity of propagating species

The cationic polymerization of N-vinyl carbazole, initiated by Ph₃C⁺ AsF₆^? and Ph₃C⁺ PF₆^? in methylene dichloride at 20 and 0°, has been studied in some detail. Reactions were very fast and rates of monomer consumption were measured using an adiabatic calorimetric technique. Initiation was relatively “slow” but complete, and termination was deduced to be insignificant during kinetic lifetimes. Values for k_p were found to vary with the initial initiator concentration; this dependence is discussed in terms of current theories regarding equilibria between ion pairs and free ions in non-aqueous solvents. k_p⁺ values estimated from two methods of extrapolation are 9.5 · 10⁵ M^?1 sec^?1 at 20° and 4.8 · 10⁵ M^?1 sec^?1 at 0°. Finally, it has been found that ion pairs are much less reactive than free ions in this system.     

The decomposition of dimethyl peroxide and the rate constant for CH3O + O2 → CH2O + HO2

The decomposition of dimethyl peroxide (DMP) was studied in the presence and absence of added NO₂ to determine rate constants k₁ and k₂ in the temperature range of 391–432°K: The results reconcile the studies by Takezaki and Takeuchi, Hanst and Calvert, and Batt and McCulloch, giving log k₁(sec^?1) = (15.7 ± 0.5) - (37.1 ± 0.9)/2.3 RT and k₂ ≈ 5 × 10⁴M^?1· sec^?1. The disproportionation/recombination ratio k_7b/k_7a = 0.30 ± 0.05 was also determined: When O₂ was added to DMP mixtures containing NO₂, relative rate constants k₁₂/k_7a were obtained over the temperature range of 396–442°K: A review of literature data produced k_7a = 10^9.8±0.5M^?1·sec^?1, giving log k₁₂(M^?1·sec^?1) = (8.5 ± 1.5) - (4.0 ± 2.8)/2.3 RT, where most of the uncertainty is due to the limited temperature range of the experiments.     

The gas-phase pyrolysis of alkyl nitrites VI. t-Amyl nitrite

The rate of decomposition of tert-amyl nitrite (t-AmONO) has been studied in the absence (120°–155°C) and presence (160°–190°C) of nitric oxide. In the absence of nitric oxide for low concentrations of tert-amyl nitrite (～10^?4M) and small extents of reaction (～1%), the first-order homogeneous rates of acetone formation are a direct measure of reaction (1) since k_3a ? k₂(NO): The rate of acetone formation is unaffected by the addition of large amounts of carbon tetrafluoride or isobutane (～1 atm) but is completely suppressed by large amounts of nitric oxide (1 atm 120°–155°C). The rate of reaction (1) is given by k₁ = 10^16.3±0.1 10^?40.3±0.1/θ sec^?1. Since (E₁ + RT) and ΔH°₁ are identical, both may be equated with D(t-AmO – NO) = 40.9 ± 0.1 kcal/mol and E₂ = 0 ± 0.1 kcal/mol. The thermochemistry leads to the result that ΔH°_f (t-AmO) = ?26.6 ± 1 kcal/mol. From ΔS°₁ and A₁, k₂ is calculated to be 10^10.5±^0.2 M^?1·sec^?1. Although the addition of nitric oxide completely suppresses acetone formation at lower temperatures, it reappears at higher temperatures. This is a result of reaction (3a) now competing with reaction (2), thus allowing k_3a to be determined. The rate constant for reaction (3a) is given by k_3a = 10^{14.7 ± 0.2} 10^{?14.3 ± 1}/θ sec^?1. There are two possible routes for the decomposition of the tert-amyloxyl radical: The dominating process is (3a). From the result at 160°C that k_3a/k_3b = 80, we arrive at the result k_3b = 10^15.0–18.7/θ sec^?1. In addition to the products accounted for by the radical split (1), methyl-2-but-1-ene and methyl-2-but-2-ene are produced as a result of the six-centre elimination of nitrous acid (5): The ratio k_5a/k_5b was 0.35. Unlike tert-butyl where the rates of the two paths were comparable [(l) and (5)], here the total rate of the elimination process was only 0.5% that of the radical split (1). The reason for this is not clear.     

Equilibrium,kinetics, and mechanism of the malonic acid–iodine reaction

The equilibrium constant for the reaction CH₂(COOH)₂ + I₃^? ? CHI(COOH)₂ + 2I^? + H⁺, measured spectrophotometrically at 25°C and ionic strength 1.00M (NaClO₄), is (2.79 ± 0.48) × 10^?4M². Stopped-flow kinetic measurements at 25°C and ionic strength 1.00M with [H⁺] = (2.09-95.0) × 10^?3M and [I^?] = (1.23-26.1) × 10^?3M indicate that the rate of the forward reaction is given by (k₁[I₂] + k₃[I₃^?]) [HOOCCH₂COO^?] + (k₂[I₂] + k₄[I₃^?]) [CH(COOH)₂] + k₅[H⁺] [I₃^?] [CH₂(COOH)₂]. The values of the rate constants k₁-k₅ are (1.21 ± 0.31) × 10², (2.41 ± 0.15) × 10¹, (1.16 ± 0.33) × 10¹, (8.7 ± 4.5) × 10^?1M^?1·sec^?1, and (3.20 ± 0.56) × 10¹M^?2·sec^?1, respectively. The rate of enolization of malonic acid, measured by the bromine scavenging technique, is given by k_en[CH₂(COOH)₂], with k_en = 2.0 × 10^?3 + 1.0 × 10^?2 [CH₂(COOH)₂]. An intramolecular mechanism, featuring a six-member cyclic transition state, is postulated to account for the results on the enolization of malonic acid. The reactions of the enol, enolate ion, and protonated enol with iodine and/or triodide ion are proposed to account for the various rate terms.     

Reaction of tetramethyl-1,2-dioxetane with triphenylphosphine: Activation parameters for the formation of 2,2-dihydro-4,4,5,5-tetramethyl-2,2,2-triphenyl-1,3,2-dioxaphospholane

The reaction of tetramethyl-1,2-dioxetane ( 1 ) and triphenylphosphine ( 2 ) in benzene-d₆ produced 2,2-dihydro-4,4,5,5-tetramethyl-2,2,2-triphenyl-1,3,2-dioxaphospholane ( 3 ) in ?90% yield over the temperature range of 6–60°. Pinacolone and triphenylphosphine oxide ( 4 ) were the major side products [additionally acetone (from thermolysis of 1 ) and tetramethyloxirane ( 5 ) were noted at the higher temperatures]. Thermal decomposition of 3 produced only 4 and 5 . Kinetic studies were carried out by the chemiluminescence method. The rate of phosphorane was found to be first order with respect to each reagent. The activation parameters for the reaction of 1 and 2 were: Ea ? 9.8 ± 0.6 kcal/mole; ΔS^≠ = ?28 eu; k_30° = 1.8 m^?1sec^?1 (range = 10–60°). Preliminary results for the reaction of 1 and tris (p-chlorophenyl)phosphine were: Ea ? 11 kcal/mole, ΔS^≠ = ?24 eu, k_30° = 1.3 M^?1sec^?1 while those for the reaction of 1 and tris(p-anisyl)phosphine were: Ea ? 8.6 kcal/mole, ΔS^≠ = ?29 eu, k_30° = 4.9 M^?1 sec^?1.     

Hydrodynamic invariant of polymer molecules

The theories of hydrodynamic properties of macromolecules in solution leading to an invariant relationship between the values of the intrinsic viscosity, [η], the molecular weight, M, and the translational friction coefficient of the molecule, f, have been considered. The review of experimental data comprising as much as about 2000 fractions of various polymers suggests that for all flexible-chain and moderately rigid-chain molecules the hydrodynamic parameter A₀ = kη₀(M[η]/100)^1/3f^?1 is actually an invariant independent of the chain length and the thermodynamic strength of the solvent and for moderately polydisperse samples also independent of the degree of their polydispersity. For polymers with very rigid chains the parameter A₀ has a high value over the experimentally investigated range of M. These conclusions make it possible to recommend the use of the following average experimental values of the invariant A₀ for the determination of M of polymers from the values of [η] and f: for flexible-chain and synthetic polymers with moderately high chain rigidity (3.2 ± 0.2) · 10^?10, for polymers with high chain rigidity (3.7 ± 0.4) · 10^?10, and for cellulose derivatives and other polysaccharides with molecular dispersity of nonelectrolyte solutions (3.30 ± 0.30) · 10^?10 erg deg^?1 mol^?1/3. The fact that the experimental value of A₀ = 3.2 · 10^?10 does not coincide with the value of A_∞ = 3.8 · 10^?10 erg deg^?1 mol^?1/3 predicted by the theories of translational friction and viscosity of macromolecules implies that the theoretical values of P_∞ = 5.11 and Φ_∞ = 2.8 · 10²³ mol^?1 are mutually incompatible and these theories require further development.     

The reaction of OH with NO2 and the deactivation of O(1D) by CO

NO₂ was photolyzed with 2288 Å radiation at 300° and 423°K in the presence of H₂O, CO, and in some cases excess He. The photolysis produces O(¹D) atoms which react with H₂O to give HO radicals or are deactivated by CO to O(³P) atoms The ratio k₅/k₃ is temperature dependent, being 0.33 at 300°K and 0.60 at 423°K. From these two points, the Arrhenius expression is estimated to be k₅/k₃ = 2.6 exp(?1200/RT) where R is in cal/mole – °K. The OH radical is either removed by NO₂ or reacts with CO The ratio k₂/k^α is 0.019 at 300°K and 0.027 at 423°K, and the ratio k₂/k⁰ is 1.65 × 10^?5M at 300°K and 2.84 × 10^?5M at 423°K, with H₂O as the chaperone gas, where k^α = k₁ in the high-pressure limit and k⁰[M] = k₁ in the low-pressure limit. When combined with the value of k₂ = 4.2 × 10⁸ exp(?1100/RT) M^?1sec^?1, k^α = 6.3 × 10⁹ exp (?340/RT)M^?1sec^?1 and k⁰ = 4.0 × 10¹²M^?2sec^?1, independent of temperature for H₂O as the chaperone gas. He is about 1/8 as efficient as H₂O.     

The kinetics of the anation reaction of aquopentaamminecobalt(III) by malonate in acidic aqueous solution

The Co(NH₃)₅OH₂³⁺ ion reacts with malonate to form Co(NH₃)₅O₂CCH₂CO₂H²⁺ or Co(NH₃)₅O₂CCH₂CO₂⁺, depending on the pH of the reaction solution. The kinetics of this anation reaction have been studied as a function of [H⁺] for the acidity range 1.5 ≤ pH ≤ 6.0 in the temperature range of 60 to 80°C, the [total malonate] ≤ 0.5 M, and the ionic strength 1.0M. The anation by malonic acid follows second-order kinetics, the rate constant being 8.0 × 10^?5 M^?1·sec^?1 at 70°C, and the anations by bimalonate (Q₁, k₁) and malonate ion (Q₂, k₂) are consistent with an I_d mechanism. Typical values at 70°C for the ion pair formation constants are Q₁ = 1.3, Q₂ = 5.4M^?1; and for the interchange rate constants k₁ = 5.3 × 10^?4; k₂ = 7.3 × 10^?4 sec^?1. The activation parameters for the various rate constants are reported and the results discussed with reference to previously reported data for similar systems.     

The reaction of methoxy radicals with nitric oxide and nitrogen dioxide

By allowing dimethyl peroxide (10^?4M) to decompose in the presence of nitric oxide (4.5 × 10^?5M), nitrogen dioxide (6.5 × 10^?5M) and carbon tetrafluoride (500 Torr), it has been shown that the ratio k₂/k_2′ = 2.03 ± 0.47: CH₃O + NO → CH₃ONO (reaction 2) and CH₃O + NO₂ → CH₃ONO₂ (reaction 2′). Deviations from this value in this and previous work is ascribed to the pressure dependence of both these reactions and heterogeneity in reaction (2). In contrast no heterogeneous effects were found for reaction (2′) making it an ideal reference reaction for studying other reactions of the methoxy radical. We conclude that the ratio k₂/k_2′ is independent of temperature and from k₁ = 10^10.2±0.4M^?1 sec^?1 we calculate that k_2′ = 10^9.9±0.4M^?1 sec^?1. Both k₂ and k_2′ are pressure dependent but have reached their limiting high-pressure values in the presence of 500 Torr of carbon tetrafluoride. Preliminary results show that k₄ = 10.^9.0±0.6 10^?4.5±1.1/ΘM^?1 sec^?1 (Θ = 2.303RT kcal mole^?1) and by k₄ = 10^8.6±0.6 10^?2.4±1.1/ΘM^?1 sec^?1: CH₃O + O₂ → CH₂O + HO₂ (reaction 4) and CH₃O + t-BuH → CH₃OH + (t-Bu) (reaction 4′).     

Kinetics of the reactions of 2,4,6-tri-t-butylphenoxyl with cumene hydroperoxide,cumylperoxyl radicals,and molecular oxygen

Reactions of 2,4,6-tri-t-butylphenoxyl (TBP) with cumene hydroperoxide (ROOH), cumylperoxyl radicals (RO₂), and molecular oxygen in benzene solution have been investigated kinetically by the ESR method. The rate constant of the reaction TBP + ROOH has been estimated in the temperature range 27°-75°C: log₁₀(k_?7/M^?1sec^?1) = (7.1 ± 0.4) - (10.9 ± 0.6 kcal mole^?1)/θ The ratio of the rate constants of reactions TBPH + RO₂ products has been determined from the experimental dependence of the rate constant of reaction TBP with ROOH on [TBPH]₀/[TBP]₀. Putting k₇ = 4.0 × 10³M^?1sec^?1, we obtain k₈ = (2.0 ± 0.2) × 10⁸M^?1sec^?1 at 30°C. The reaction of TBP with O₂ obeys the kinetic law ?d[TBP]/dt = k′[O₂][TBP]². This is in accordance with scheme TBP + O₂ ← TBP ?O₂ [I]; TBP ?O₂ + TBP · products, log₁₀ (k′/M^?2sec^?1) = (?14.5 ± 0.9) + (27.2 ± 1.4)/θ at 66°?78°C, where ° = 2.303RT.     

The gas-phase pyrolysis of alkyl nitrites. II. s-butyl nitrite

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 k_3c » k₂(NO): . Unlike t-butyl nitrite (TBN), d(AcH)/dt is independent of added CF₄ (～0.9 atm). Thus k_3c is always » k₂ (NO) over this pressure range. Large amounts of NO (～0.9 atm) (130–160°C) completely suppress AcH formation. k₁ = 10^16.2–40.9/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 10^10.4 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 = 10^9.8M^?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 = 10^12.8–35.8/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 k_global = 10^15.7–39.6/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 k_obs = 10^12.9–35.4/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 k_obs 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 k_3c to be determined. Ignoring reaction (6), d(AcH)/dt = k₁k₃ (SBN )/[k_3c + k₂(NO)]; k_3c = 10^14.8–15.3/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.     

Kinetics and rate constants for the reaction of tri-n-butylgermanium hydride with methyl iodide and carbon tetrachloride. The combination of methyl and trichloromethyl radicals in solution.

The kinetics of the photoinitiated reductions of methyl iodide and carbon tetrachloride by tri-n-butylgermanium hydride in cyclohexane at 25°C have been studied and absolute rate constants have been measured. Rate constants for the combination of CH₃? and CCl₃? radicals are equal within experimental error and are also equal to the values found for the self-reactions of most non-polymeric radicals in low viscosity solvents, i.e. ～1–3 × 10⁹ M^?1 sec^?1. Rate constants for hydrogen atom abstraction by CH₃? and CCl₃? radicals are both ～1?2 × 10⁵ M^?1 sec^?1. Tri-n-butyltin hydride is about 10–20 times as good a hydrogen donor to alkyl radicals as is tri-n-butylgermanium hydride. The strength of the germanium–hydrogen bond, D(n-Bu₃Ge–H) is estimated to be approximately 84 kcal/mole.     

The free-radical reaction of tri-n-butyltin hydride with peroxides

Rate constants for the tri-n-butyltin radical ( Sn · ) induced decomposition of a number of peroxides have been measured in benzene at 10°C. The values range from ～100 M^?1 sec^?1 for di-t-butyl peroxide to 2.6 × 10⁷ M^?1 sec^?1 for di-t-butyl diperoxyisophthalate. The majority of the peroxides, including diethyl peroxide, diacetyl peroxide, and t-butyl peracetate, have rate constants of ～10⁵ M^?1 sec^?1. It is shown that di-n-alkyl disulfides are ten times as reactive toward Sn · as di-n-alkyl peroxides, although the exothermicities of these reactions are ～15 and ～39 kcal/mole, respectively. The enhanced reactivity of the disulfides is attributed to the easier formation of an intermediate or transition state with 9 electrons around sulfur, compared with an analogous species with 9 electrons around oxygen. The following bond strengths (kcal/mole) have been estimated: D[ Sn ? OR] = 77; D[ Sn ? H] = 82; D[ Sn ? SR] = 83; and D[ Sn ? OC(O)R] = 86, where R = alkyl. Rate constants for reaction of Sn · with some benzyl esters have also been measured. It has been found that t-butoxy radicals can add to benzene and abstract hydrogen from benzene at ambient temperatures.     

Rate constants for radical recombination. III. The trifluoromethyl radical

Products of the radical reactions arising from t-Bu₂O₂, CF₃I, and CH₃I at 146°C in the vapor phase have been measured over a 33-fold range of CH₃I/CH₃I ratios and shown to be governed by the rapidly established equilibrium Together with K estimated by thermochemical methods, the results yield, for the rate of recombination for CF₃· radicals, k_r = 10^{9.7 ± 0.5} M^?1 sec^?1.     

On the reactivity of diethyl hydroxyl amine toward free radicals

Diethyl hydroxyl amine is an efficient trap for alkyl, alkoxy, and peroxy radicals. The specific rate constant for the reaction of ethyl radicals (gas phase, 25°C), tert-butoxy radicals (benzene solution, 115°C), and poly (peroxystyryl) peroxy radicals (styrene solution, 50°C) were evaluated as 7.2 × 10⁵, 7.7 × 10⁷, and 2.9 × 10⁵ M^?1·sec^?1, respectively. Several possible secondary reactions of the nitroxide radicals are discussed.     

The triplet state of 4-nitronaphthylamine

Nanosecond spectroscopic and kinetic studies of 4-nitronaphthylamine (4-NO₂NA) in aerated and deaerated nonpolar solvents at room temperature show a transient species with absorption maxima at 470 and 665 nm. The rate constant for the decay of this species in deaerated benzene is 6.7 × 10⁵ sec^?1, while in aerated benzene solutions the species is quenched by oxygen with arate constant k = 2.0 × 10⁹M^?1·sec^?1. The transient absorption at 470and 665 nm is assigned to the lowest triplet excited state of 4-NO₂NA. In polar solvents, however, electronic excitation of 4-NO₂NA does not lead to any detectable transient absorption between 400 and 800 nm for the temperature range of 25 to ?150°C. This is attributed to lack of intersystem crossing of 4-NO₂NA in polar solvents.     

The Kinetics of the Anation Reaction of Aquopentaamminerhodium(III) lon by Oxalate in Aqueous Acidic Solution

The anation reaction of aquopentaamminerhodium(III) by oxalate has been studied in the temperature range 51–69°C and acidity range 0 ≤ pH ≤ 4.5 for oxalate concentrations up to 0.25 M and at ionic strength 1.0 M. The kinetic results provide evidence for the formation of an ion-pair between the complex ion and HC₂O(Q₁) and C₂O₄^2?(Q₂), where Q₁ = 2.3 M^?1 and Q₂ = 8.1 M^?1 at 60°C, but no evidence for an ion-pair with H₂C₂O₄ exists. The values of the rate constants at 60°C for anation by H₂C₂O₄, HC₂O and C₂O₄^2? are k₀ = 1.5 · 10^?4 M^?1 sec^?1, k₁ = 1.4 · 10^?4 sec^?1 and k₂ = 1.2 · 10^?4 sec^?1. The corresponding values for ΔH≠ and ΔS≠ are reported and the results discussed with reference to analogous reactions of Rh(III) and Co(III).