Kinetics and mechanism of oxidation of the drug mephenesin by bis(hydrogenperiodato)argentate(III) complex anion

Mephenesin is being used as a central‐acting skeletal muscle relaxant. Oxidation of mephenesin by bis(hydrogenperiodato)argentate(III) complex anion, [Ag(HIO₆)₂]^5?, has been studied in aqueous alkaline medium. The major oxidation product of mephenesin has been identified as 3‐(2‐methylphenoxy)‐2‐ketone‐1‐propanol by mass spectrometry. An overall second‐order kinetics has been observed with first order in [Ag(III)] and [mephenesin]. The effects of [OH^?] and periodate concentration on the observed second‐order rate constants k′ have been analyzed, and accordingly an empirical expression has been deduced: k′ = (k_a + k_b[OH^?])K₁/{f([OH^?])[IO^?₄]_tot + K₁}, where [IO^?₄]_tot denotes the total concentration of periodate, k_a = (1.35 ± 0.14) × 10^?2M^?1s^?1 and k_b = 1.06 ± 0.01 M^?2s^?1 at 25.0°C, and ionic strength 0.30 M. Activation parameters associated with k_a and k_b have been calculated. A mechanism has been proposed to involve two pre‐equilibria, leading to formation of a periodato‐Ag(III)‐mephenesin complex. In the subsequent rate‐determining steps, this complex undergoes inner‐sphere electron transfer from the coordinated drug to the metal center by two paths: one path is independent of OH^? whereas the other is facilitated by a hydroxide ion. In the appendix, detailed discussion on the structure of the Ag(III) complex, reactive species, as well as pre‐equilibrium regarding the oxidant is provided. © 2007 Wiley Periodicals, Inc. Int J Chem Kinet 39: 440–446, 2007     

氢过碘酸）合银（III）配离子氧化L-脯氨酸的反应动力学及机理

L-脯氨酸独有的亚胺基使其在生物医药领域具有许多独特的功能,并广泛用作不对称有机化合物合成的有效催化剂。本文在碱性介质中研究了二（氢过碘酸）合银（III）配离子氧化 L-脯氨酸的反应。经质谱鉴定,脯氨酸氧化后的产物为脯氨酸脱羧生成的 γ-氨基丁酸盐;氧化反应对脯氨酸及Ag(III) 均为一级;二级速率常数 k′ 随 [IO₄^-] 浓度增加而减小,而与 [OHˉ] 的浓度几乎无关;推测反应机理应包括 [Ag(HIO₆)₂]^5-与 [Ag(HIO₆)(H₂O)(OH)]^2-之间的前期平衡,两种Ag（III）配离子均作为反应的活性组分,在速控步被完全去质子化的脯氨酸平行地还原,两速控步对应的活化参数为： k₁ (25 ^oC)＝1.87±0.04（mol·L^-1)^-1s^-1,∆ H₁^≠＝45±4 kJ · mol^-1, ∆ S₁^≠＝-90±13 J· K^-1·mol^-1 and k₂ (25 ^oC) ＝3.2±0.5（mol·L^-1)^-1s^-1, ∆ H₂^≠＝34±2 kJ · mol^-1, ∆ S₂^≠＝-122 ±10 J· K^-1·mol^-1。本文第一次发现 [Ag(HIO₆)₂]^5-配离子也具有氧化反应活性。     

Kinetics and mechanism of the iodine oxidation of hydroxylamine

Studies of the stoichiometry and kinetics of the reaction between hydroxylamine and iodine, previously studied in media below pH 3, have been extended to pH 5.5. The stoichiometry over the pH range 3.4–5.5 is 2NH₂OH + 2I₂ = N₂O + 4I^? + H₂O + 4H⁺. Since the reaction is first-order in [I₂] + [I₃^?], the specific rate law, k₀, is k₀ = (k₁ + k₂/[H⁺]) {[NH₃OH⁺]₀/(1 + K_p[H⁺])} {1/(1 + K_I[I^?])}, where [NH₃OH⁺]₀ is total initial hydroxylamine concentration, and k₁, k₂, K_p, and K_I are (6.5 ± 0.6) × 10⁵ M^?1 s^?1, (5.0 ± 0.5) s^?1, 1 × 10⁶ M^?1, and 725 M^?1, respectively. A mechanism taking into account unprotonated hydroxylamine (NH₂OH) and molecular iodine (I₂) as reactive species, with intermediates NH₂OI₂^?, HNO, NH₂O, and I₂^?, is proposed.     

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.     

Mechanistic study of the oxidation of N-phenyldiethanolamine by <Emphasis Type="Italic">bis</Emphasis>(hydrogen periodato)argentate(III) complex anion

The kinetics of oxidation of phenyldiethanolamine (PEA) by a silver(III) complex anion, [Ag(HIO₆)₂]⁵⁻, has been studied in an aqueous alkaline medium by conventional spectrophotometry. The main oxidation product of PEA has been identified as formaldehyde. In the temperature range 20.0–40.0 °C , through analyzing influences of [OH⁻] and [IO ₄ ⁻ ]_tot on the reaction, it is pseudo-first-order in Ag(III) disappearance with a rate expression: k _obsd = (k ₁ + k ₂[OH⁻]) K ₁ K ₂[PEA]/{f([OH⁻])[IO ₄ ⁻ ]_tot + K ₁ + K ₁ K ₂ [PEA]}, where k ₁ = (0.61 ± 0.02) × 10⁻² s⁻¹, k₂ = (0.049 ± 0.002) M⁻¹ s⁻¹ at 25.0 °C and ionic strength of 0.30 M. Activation parameters associated with k ₁ and k ₂ have also been derived. A reaction mechanism is proposed involving two pre-equilibria, leading to formation of an Ag(III)-periodato-PEA ternary complex. The ternary complex undergoes a two-electron transfer from the coordination PEA to the metal center via two parallel pathways: one pathway is spontaneous and the other is assisted by a hydroxide ion.     

Kinetics of cerium (IV) oxidation of mandelic acid. Ionic strength and specific cation effects

The kinetics of the redox reaction between mandelic acid (MA) and ceric sulfate have been studied in aqueous sulfuric acid solutions and in H₂SO₄? MClO₄ (M⁺ = H⁺, Li⁺, Na⁺) and H₂SO₄? MHSO₄ (M⁺ = Li⁺, Na⁺, K⁺) mixtures under various experimental conditions of total electrolyte concentration (that is, ionic strength) and temperature. The oxidation reaction has been found to occur via two paths according to the following rate law: rate = k[MA] [Ce(IV)], where k = k₁ + k₂/(1 + a)²[HSO₄^?]² = k₁ + k₂/(1 + 1/a)²[SO₄^2?]², a being a constant. The cations considered exhibit negative specific effects upon the overall oxidation rate following the order H⁺ ? Li⁺ < Na⁺ < K⁺. The observed negative cation effects on the rate constant k₁ are in the order Na⁺ < Li⁺ < H⁺, whereas the order is in reverse for k₂, namely, H⁺ ? Li⁺ < Na⁺. Lithium and hydrogen ions exhibit similar medium effects only when relatively small amounts of electrolytes are replaced. The type of the cation used does not affect significantly the activation parameters.     

Persulfate-initiated polymerization of acrylamide

A dilatometric technique was used to obtain conversion–time data for the polymerization of acrylamide initiated by potassium persulfate in water. The results are summarized by the empirical rate expression, ?d[M₁]/dt = R_p = k_1.25[K₂S₂O₈]^0.5[M₁]^1.25, and k_1.25 = 1.70 × 10¹¹ exp {?16,900/RT} 1.^0.75/mole^?0.75-min. Persulfate was varied over the range 9.5 × 10^?4 to 5.2 × 10^×2 mole/l., and initial monomer concentration [M₁] was varied from 0.05 to 0.4 mole/l. The temperature range was 30?50°C. Results of analysis of the kinetics and energetics of the polymerization favor a cage-effect theory rather than a complex-formation theory to explain the order with respect to monomer.     

Kinetic Study of the Oxidation of N,N-Dimethylethanolamine by Bis(Hydrogenperiodato)Argentate(III) in Aqueous Solution

The oxidation of N,N-dimethylethanolamine (DMEA) by bis(hydrogenperiodato) argentate(III) ([Ag(HIO₆)₂]⁵⁻) was studied in aqueous alkaline medium. Formaldehyde and dimethylamine were identified as the major oxidation products after the oxidation of DMEA. The oxidation kinetics was followed spectrophotometrically in the temperature range of 25.0 °C–40.0 °C. It was found that the reaction was first order in [Ag(III)]; the oberved first-order rate constants k _obsd as functions of [DMEA], [OH⁻] and total concentration of periodate ([IO₄^-]_tot[\mathrm{IO}_{4}^{-}]_{\mathrm{tot}}) were analyzed and were revealed to follow a rate expression: k_obsd = (k₁ +k₂[OH^-])K₁K₂[DMEA]/{f([OH^-])[IO₄^-]_tot+ K₁ + K₁K₂[DMEA]}k_{\mathrm{obsd}} = (k_{1} +k_{2}[\mathrm{OH}^{-}])K_{1}K_{2}[\mathrm{DMEA}]/\{f([\mathrm{OH}^{-}])[\mathrm{IO}_{4}^{-}]_{\mathrm{tot}}+ K_{1} + K_{1}K_{2}[\mathrm{DMEA}]\}. Rate constants k ₁ and k ₂ and equilibrium constant K ₂ were derived; activation parameters corresponding to k ₁ and k ₂ were computed. In the proposed reaction mechanism, a peridato-Ag(III)-DMEA ternary complex is formed indirectly through a reactive intermediate species [Ag(HIO₆)(OH)(H₂O)]²⁻. In subsequent rate-determining steps as described by k ₁ and k ₂, the ternary complex decays to Ag(I) through two reaction pathways: one of which is spontaneous and the other is prompted by an OH⁻.     

Temperature and pressure effects on the kinetics of the Bromate ion-iodide ion-L-ascorbic acid clock reaction

The kinetics of the bromate ion-iodide ion-L-ascorbic acid clock reaction was investigated as a function of temperature and pressure using stopped-flow techniques. Kinetic results were obtained for the uncatalyzed as well as for the Mo(VI) and V(V) catalyzed reactions. While molybdenum catalyzes the BrO-I^? reaction, vanadium catalyzes the direct oxidation of ascorbic acid by bromate ion. The corresponding rate laws and kinetic parameters are as follows. Uncatalyzed reaction: r₂ = k₂[BrO] [I^?][H⁺]², k₂ = 38.6 ± 2.0 dm⁹ mol^?3 s^?1, ΔH? = 41.3 ± 4.2 kJmol^?1, ΔS? = ?75.9 ± 11.4 Jmol^?1 K^?1, ΔV? = ?14.2 ± 2.9 cm³ mol^?1. Molybdenum-catalyzed reaction: r′₂ = k₂[BrO] [I^?] [H⁺]² + k_Mo[BrO] [I^?] [ H⁺]²[M₀(VI)], k_Mo = (2.9 ± 0.3)10⁶ dm¹² mol^?4 s^?1, ΔH? = 27.2 ± 2.5 kJmol^?1, ΔS? = ?30.1 ± 4.5 Jmol^?1K^?1, ΔV? = 14.2 ± 2.1 cm³ mol^?1. Vanadium-catalyzed reaction: r′₁ = k_V[BrO] [V(V)], k_V = 9.1 ± 0.6 dm³ mol^?1 s^?1, ΔH? = 61.4 ± 5.4 kJmol^?1, ΔS? = ?20.7 ± 3.1 Jmol^?1K^?1, ΔV? = 5.2 ± 1.5 cm³ mol^?1. On the basis of the results, mechanistic details of the BrO-I^? reaction and the catalytic oxidation of ascorbic acid by BrO are elaborated. © 1995 John Wiley & Sons, Inc.     

Reactions of [NiIII(cyclam)] with aqueous sulfur dioxide: Exhibition of a deuterium isotope effect and catalysis by chloride and bromide

One unit of S(IV) (SO₂ or SHO₃^?) is oxidized per 2 units of [Ni^III(cyclam)] species to obtain sulfate. Kinetic analyses have been done by varying the acidities (0.013 ? [H⁺] ? 1.0 M) and halide concentrations (0.000 ? [X^?] ? 0.012 M; X=Cl and Br) at constant ionic strength (μ = 1.0 M). The rate law that incorporates the [X^?] and [H⁺] dependence is ?d[Ni^III]_T/dt=2k[Ni^III]_T[S(IV)]_T where 2k={k_a[H⁺] + k_bK + kK′_X[H⁺] [X^?] + kK′_XK[X^?]} {[H⁺] + K}^?1 {1 + K′_X[X^?]}^?1, here k_a=87 ± 7 M^?1 s^?1, k_b=(2.5 ± 0.5)×10³ M^?1 s^?1 and pK = 1.8 ± 0.2. Rate constants k_a and k_b are attributed to the reactions of [Ni^III(cyclam) (H₂O)₂]³⁺ with SO₂ and SHO₃^?, respectively. Monohalo species apparent equilibrium constants K′_Cl=(1600 ± 400) M^?1 and K′_Br=(190 ± 20) M^?1 and rate constants k=80 ± 8 M^?1 s^?1 and k = 140 ± 15 M^?1 s^?1 are ascribed to the protonated pathway, involving the [Ni^III(cyclam) (H₂O)X]²⁺ and SO_2(aq) reaction pairs. The other two rate constants of k=(5 ± 1)×10³ M^?1 s^?1 and k=(3.1 ± 0.5)×10⁴ M^?1 s^?1, refer to the deprotonated pathway and are assigned to the [Ni^III(cyclam) (H₂O)X]²⁺ /SHO₃^? redox couple. A deuterium H₂O / D₂O isotope effect of 2.1–2.8 can be attributed partially to an equilibrium isotope effect at low acidity though a small kinetic isotope (2.5 ± 0.5) effect is evident for the dihydrogen sulfito pathway, k_a. The kinetic isotope effect and the absence of sulfite radical scavenging effects are explained by a mechanism entailing migration of a hydride from sulfur to the Ni^III center to produce a Ni^III—H species, which rapidly comproportionates, and S(VI). © 1993 John Wiley & Sons, Inc.     

Kinetic Study of the Bromine-Catalyzed Isomerization of Maleic Acid in Aqueous Ce(IV)-Br−-H2SO4 Medium

The presence of ceric and bromide ions catalyzes the isomerization of maleic acid (MA) to fumaric acid (FA) in aqueous sulfuric acid. A kinetic study of this bromine-catalyzed reaction was carried out. The reaction between ceric ion and maleic acid is first order with respect to Ce(IV). For [Ce(IV)]₀=5.0×10^?4 M, [H₂SO₄]₀=1.2 M, μ=2.0 M (adjusted by NaClO₄), and [MA]₀=(0.5–1.0)M, the observed pseudo-first-order rate constant (k₀₃) at 25° is k₀₃=7.622×10^?5 [MA]₀/(1+0.205[MA]₀). The reaction between ceric and bromide ions is first order with respect to Ce(IV). For [Ce(IV)]₀=5.0×10^?4 M, [H₂SO₄]₀=1.2 M, μ=2.0 M, and [Br^?]₀=(0.025–0.150)M, the pseudo-first-order rate constant (k₀₂) at 25° is k₀₂= (4.313±0.095)x10^?2[Br^?]²+(2.060±0.119)x10^?3[Br^?]. The reaction of Ce(IV) with maleic acid and bromide ion is also first order with respect to Ce(IV). For [Ce(IV)]₀=5.0×10^?4 M, [MA]₀=0.75 M, [H₂SO₄]₀=1.2 M, μ=2.0 M, and [Br^?]₀= (0.025–0.150)M, the pseudo-first-order rate constant (k₀₃) at 25° is k₀₃= (5.286±0.045)x10^?2[Br^?]²+(3.568±0.056)x10^?3[Br^?]. For [Ce(IV)]₀=5.0 × 10^?4 M, [Br^?]₀=0.050 M, [H₂SO₄]₀=1.2 M, μ=2.0 M, and [MA]₀=(0.15–1.0)M at 25°, k₀₃=(2.108×10^?4+2.127×10^?4[MA]₀)/(1+0.205[MA]₀). A mechanism is proposed to rationalize the results. The effect of temperature on the reaction rate was also studied. The energy barrier of Ce(IV)—Br^? reaction is much less than that of Ce(IV)—MA reaction. Maleic and fumaric acids have very different mass spectra. The mass spectrum of fumaric acid exhibits a strong metastable peak at m/e 66.5.     

Kinetic and mechanistic studies on the oxidation of pyrrolidine by bis(hydrogenperiodato)argentate(III) complex anion

The kinetics of oxidation of pyrrolidine by bis(hydrogenperiodato)argentate(III) complex anion ([Ag(HIO₆)₂]^5?) was studied in alkaline medium, with reaction temperatures in the range of 15.0–30.0 °C. The experiments indicated that the oxidation follows an overall second-order reaction, being first-order in both Ag(III) and pyrrolidine. The observed second-order rate constants, k′, decreased with increasing [IO₄ ^?] but increased slightly with increasing [OH^?]. The influence of ionic strength on the reaction rate was also investigated. The oxidation resulted in oxidative deamination of pyrrolidine, giving 4-hydroxybutyrate as the product. A reaction mechanism is proposed which includes an equilibrium between [Ag(HIO₆)₂]^5? and [Ag(HIO₆)₂(OH)(H₂O)]^2?; these two Ag(III) species are reduced by pyrrolidine in parallel rate-determining steps. The rate equation derived from the proposed mechanism can explain the experimental observations. The rate constants of the rate-determining steps, together with the associated activation parameters, were calculated accordingly.     

Kinetics and mechanism of the aquation of pentaaqua (trichloroacetato-O)-chromium(2+) ion

The kinetics of the aquation of (H₂O)₅Cr(O₂CCCl₃)²⁺ have been examined at 35–55°C and 1.00M ionic strength with [H⁺] = 0.01?1.00M. The reaction follows the rate equation -d ln [Cr_total]/dt = (a[H⁺]^?1 + b + c[H⁺])/(1 + d[H⁺]), where [Cr_total] is the stoichiometric concentration of the complex. At 45°C a = (1.41 ± 0.03) × 10^?7M/s, b = (1.66 ± 0.02) × 10^?5 s^?1, c = (7.0 ± 0.8) × 10^?5M^?1·S^?1 and d = 2.3 ± 0.3M^?1. Two mechanisms consistent with this rate law are discussed, with evidence being presented in favor of an ester hydrolysis mechanism involving steady-state intermediates. Equilibrium and activation parameters were determined.     

Kinetics and Mechanism of the Cerium(IV) Oxidation of Arnold's Base and the Protonation and Hydroxylation of Michler's Hydrol Blue

In aqueous H₂SO₄, Ce(IV) ion oxidizes rapidly Arnold's base((p-Me₂NC₆H₄)₂CH₂, Ar₂CH₂) to the protonated species of Michler's hydrol((p-Me₂NC₆H₄)₂CHOH, Ar₂CHOH) and Michler's hydrol blue((p-Me₂NC₆H₄)₂CH⁺, Ar₂CH⁺). With Ar₂CH₂ in excess, the rate law of the Ce(IV)-Ar₂CH₂ reaction in 0.100 M H₂SO₄ is expressed -d[Ce(IV)]/dt = k_app[Ar₂CH₂]₀[Ce(IV)] with k_app = 199 ± 8M^?1s^?1 at25°C. When the consumption of Ce(IV) ion is nearly complete, the characteristic blue color of Ar₂CH⁺ ion starts to appear; later it fades relatively slowly. The electron transfer of this reaction takes place on the nitrogen atom rather than on the methylene carbon atom. The dissociation of the binuclear complex [Ce(III)ArCHAr-Ce(III)] is responsible for the appearance of the Ar₂CH⁺ dye whereas the protonation reaction causes the dye to fade. In highly acidic solution, the rate law of the protonation reaction of Michler's hydrol blue is -d[Ar₂CH⁺]/dt = k_obs[Ar₂CH⁺] where K_obs = ((ac + 1)[H*] + bc[H⁺]²)/(a + b[H⁺]) (in HClO₄) and k_obs= ((ac + 1 + e[HSO₄^?])[H⁺] + bc[H⁺]² + d[HSO₄^?] + q[HSO₄^?]²/[H⁺])/(a + b[H⁺] + f[HSO₄^?] + g[HSO₄^?]/[H⁺]) (in H₂SO₄), and at 25°C and μ = 0.1 M, a = 0.0870 M s, b = 0.655 s, c = 0.202 M^?1s^?1, d = 0.110, e = 0.0070 M^?1, f = 0.156 s, g = 0.156 s, and q = 0.124. In highly basic solution, the rate law of the hydroxylation reaction of Michler's hydrol blue is -d[Ar₂CH⁺]/dt = k_OH[OH^?]₀[Ar₂CH⁺] with k_OH = 174 ± 1 M^?1s^?1 at 25°C and μ = 0.1 M. The protonation reaction of Michler's hydrol blue takes place predominantly via hydrolysis whereas its hydroxylation occurs predominantly via the path of direct OH attack.     

A kinetic investigation of the oxidation of <Emphasis Type="SmallCaps">dl</Emphasis>-pipecolinate by bis(hydrogenperiodato)argentate(III) complex anion

Kinetics of oxidation of dl-pipecolinate by bis(hydrogenperiodato)argentate(III) complex anion, [Ag(HIO₆)₂]⁵⁻, has been studied in aqueous alkaline medium in the temperature range of 25–40 °C. The oxidation kinetics is first order in the silver(III) and pipecolinate concentrations. The observed second-order rate constant, decreasing with increasing [periodate] is virtually independent of [OH⁻]. α-Aminoadipate as the major oxidation product of pipecolinate has been identified by chromatographic analysis. A reaction mechanism is proposed that involves a pre-equilibrium between [Ag(HIO₆)₂]⁵⁻ and [Ag(HIO₆)(H₂O)(OH)]²⁻, a mono-periodate coordinated silver(III) complex. Both Ag(III) complexes are reduced in parallel by pipecolinate in rate-determining steps (described by k ₁ for the former Ag(III) species and k ₂ for the latter). The determined rate constants and their associated activation parameters are k ₁ (25 °C) = 0.40 ± 0.02 M⁻¹ s⁻¹, ∆H ₁^≠ = 53 ± 2 kJ mol⁻¹, ∆S ₁^≠ = −74 ± 5 J K⁻¹ mol⁻¹ and k ₂ (25 °C) = 0.64 ± 0.02 M⁻¹ s⁻¹, ∆H ₂^≠ = 41 ± 2 kJ mol⁻¹, ∆S ₂^≠ = −110 ± 5 J K⁻¹ mol⁻¹. The time-resolved spectra, a positive dependence of the rate constants on ionic strength of the reaction medium, and the consistency of pre-equilibrium constants derived from different reaction systems support the proposed reaction mechanism.     

The oxidation of thiosulphate ions by octacyanotungstate(V) in weak acidic medium

The kinetics of the oxidation of thiosulphate ions by octacyanotungstate(V) ions has been studied in the pH range 3.9–5.0. The reaction showed zero-order kinetics with respect to [W(CN)₈^3?] and is consistent with the rate law R = k[H⁺][S₂O₃^2?]². A reaction mechanism is proposed for the reaction with a third-order rate constant of 0.26 M^?2 s^?1 at 25°C.     

Kinetics and mechanism of pyrogallol autoxidation: Calibration of the dynamic response of an oxygen electrode

The kinetics of the reaction between 1,2,3-trihydroxybenzene (pyrogallol) and O₂ (autoxidation) have been determined by monitoring the concentration of dissolved dioxygen with a polarographic oxygen electrode. The reaction is carried out in pseudo-first-order excess pyrogallol, 25°C, 0.08 M NaCl, and 0.04 M phosphate buffer in the pH range 6.9–10.5. Data collection precedes reaction initiation, but only the data recorded after the estimated 3.2 s dead time are used in kinetics calculations. Observed rate constants are corrected for incomplete mixing, which is treated as a first-order process that has an experimentally determined mixing rate constant of 4.0 s^?1. The rate law for the reaction is ?d[O₂]/dt=k_app[PYR]_tot[O₂], in which [PYR]_tot is the total stoichiometric pyrogallol concentration. A mechanism is presented which explains the increase in rate with increasing [OH^?] by postulating that H₂PYR^? (k₂) has greater reactivity with dissolved dioxygen than does H₃PYR (k₁). The data best fit the equation k_app=(k₁ + k₂K_H[OH^?])/(1 + K_H[OH^?]) when the value of the hydrolysis constant K_H (the quotient of the pyrogallol acid dissociation and water autoprotolysis constants) for this medium equals 3.1×10⁴ M^?1. The resulting values of k₁ and k₂, respectively, equal (0.13 + 0.01) M^?1 s^?1 and (3.5 plusmn; 0.1) M^?1 s^?1. This reaction is recommended as a test reaction for calibrating the dynamic response of an O₂-electrode. © 1993 John Wiley & Sons, Inc.     

Synthetic thermally reversible gel systems. IV

The polymerization kinetics in water of acrylylglycinamide (AG) initiated by K₂S₂O₈ was studied over the temperature range 40.0 to 60.0°C. Monomer concentration was varied from 7.8 × 10^?3 to 31.2 × 10^?3M and catalyst from 1.85 × to 11.10 × 10^?5M. The rate expression is ?d[M]/dt = R_p, = k_1.22[K₂S₂O₈]^0.5[M]^1.22, and the overall empirical rate constant, k_1.22 = 1.14 × 10¹¹e^?15,800/RT 1.^0.72 mole^?0.72 min^?1. To explain the dependence on monomer, a kinetic scheme which includes a bimolecular reaction (k₂) between monomer and initiator is suggested. The simplified expression which describes the initial rate of polymerization is: ?d[M]/dt = R_p, = k₄(2[I]/k₅)^1/2[M](k₁ + k₂[M])^1/2, where k₁, k₂, k₄ and k₅ are rate constants for S₂O₈ = decomposition, a bimolecular reaction between monomer and initiator, propagation, and termination, respectively. Individual bimolecular rate constants are expressed in liter/mole-min. The equation predicts a dependence on monomer concentration between 1.0 and 1.5 with 1.5 being approached a t high monomer concentrations. Plots of R_P²/[M]² versus [M] are linear, as predicted by the postulated reaction route and values for k₂ and k₄/k₅^1/2 were obtained from the slopes and intercepts of these plots. The temperature dependence of the bimolecular monomer-initiator reaction is k₂ = 5.19 × 10²¹e^?36,000/RT. Instead of the usual behavior, the k₄/k₅^1/2 ratio was found to decrease with temperature and the difference of activation energies, (E₄ ? E₅/2), is ?1.50 kcal. The temperature dependence of the propagation to square root of the termination rate constant ratio is k₄/k₅^1/2 = 6.16e^1500/RT. These rather unusual results may be related to the ability of AG polymers in water to form thermally reversible gels; even above the gel melting points, the polymers are considerably aggregated in solution. This would tend to make the bimolecular termination reaction more temperature dependent and also account for the high values (59–69) for the k₄/k₅^1/2 ratios. For similar temperatures, the overall rate constants for AG are approximately four times those for acrylamide.     

Kinetics and mechanism of silver(I)-catalysed oxidation of malonic acid by peroxodiphosphate in acetate buffers

Summary The kinetics of the silver(I)-catalysed oxidation of malonic acid by peroxodiphosphate (pdp) was studied in acetate buffers. The rate law as represented by-d[pdp]/dt = {(k ₁ K _inf2 ^sup-1 [H⁺]² + k ₂[H⁺] + k ₃ K ₃)/ ([H⁺]²/K ₂ + [H⁺] + K ₃)}[pdp][Ag(I)] conforms to the proposed mechanism. The rate is independent of malonic acid concentrations. Acetate ions do not affect the rate; however, the rate decreases as the ionic strength increases. A probable portrait of reaction events is suggested. A comparative analysis of the reactivity pattern of malonic acid towards peroxodiphosphate and peroxodisulphate in presence of silver(I) has been made.     

A Thermodynamic Study of Chromotropico-Titanium(III) Complex from Spectrophotometric Measurements

The formation of 1 : 2 titanium(III) complex with chromotropic acid (4, 5-dihydroxy-2, 7-naphthalene-disulfonic acid) was observed by spectrophotometric measurements at various ionic strengths. An expression, [Ti(III)]/D=1/Δ? + α_H²+/KΔ?[H₂R^2?]², was derived for the determination of the formation constant, K=7.2×10² liter² mol^?2 for the Ti(III).(HR)₂ ion in the pH range of 1.3–1.8 at constant ionic strength, I=0.2 M, at 25°C. The thermodynamic data for the reaction, Ti(III)+2H₃R^2?=Ti(III) (HR)₂+2H⁺, were calculated to be ΔG° = ?16 kJ mol^?1 ΔH° = 18 kJ mol^?1, ΔS° = 110 JK^?1 mol^?1, at 25°C.