Rm Mol Research Articles

Kinetics and Equilibria for Complex Formation between Palladium(II) and Chloroacetate

Kinetics and equilibria for the formation of a 1:1 complex between palladium(II) and chloroacetate were studied by spectrophotometric measurements in 1.00 mol\(\cdot {\rm dm}^{-3}\) HClO4 at 298.2 K. The equilibrium constant, K, of the reaction $$ {\rm Pd}^{2 + } + {\rm HL}^{K} {\rm PdL}^+ +{\rm H}^+ $$ was determined from multi-wavelength absorbance measurements of equilibrated solutions at variable temperatures as log\(_{10} \, K (298.2\,K) = 0.492 \pm \) 0.006 with \(\Delta H^{\circ}=-15.{6}\pm 1.{9} \,{\rm kJ} \cdot {\rm mol}^{-1}\) and \(\Delta^{\circ}=-43.{0}\pm 6.{3} {\rm J} \cdot {\rm K}^{-1} \cdot {\rm mol}^{-1}\), and spectra of individual species were calculated. Variable-temperature kinetic measurements gave rate constants for the forward and backward reactions at 298.2 K and ionic strength 1.00 mol\( \cdot\,{\rm dm}^{-3}\) as \(k_{1}=6.46\pm 0.08 {\rm dm}^{3} \cdot {\rm mol}^{-1} \cdot {\rm s}^{-1}\) and \(k_{-1}=2.91\pm 0.03 {\rm s}^{-1}\), with activation parameters \(\smash{\Delta H_1^{\ne} = 56.55 \pm 0.75 } {\rm kJ} \cdot {\rm mol}^{-1}, \smash{\Delta S_1^{\ne} = - 40.0 \pm 2.5} {\rm J}\cdot {\rm K}^{-1} \cdot {\rm mol}^{-1}, \smash{\Delta H_{ - 1}^{\ne} = 64.42 \pm 0.50 } {\rm kJ} \cdot {\rm mol}^{-1}\) and \(\smash{\Delta S_{ - 1}^{\ne} = - 20.3 \pm 1.7} {\rm J} \cdot {\rm K}^{-1} \cdot {\rm mol}^{-1}\), respectively. From the kinetics of the forward and reverse processes, \(\log _{10} \,K = 0.346 \pm 0.007, \Delta H^{\circ}=-7.9\pm 0.9 {\rm kJ} \cdot {\rm mol}^{-1}\) and \(\Delta^{\circ}=-19.{8}\pm 3.{0} {\rm J}\cdot {\rm K}^{-1} \cdot {\rm mol}^{-1}\) were derived in good agreement with the results of the equilibrium measurements. Specific Ion Interaction Theory was employed for determination of thermodynamic equilibrium constants for the protonation of chloroacetate (\(\log _{10} \,K^\circ _ {\rm p} = 2.857 \pm 0.016 \)) and formation of the PdL+ complex (\(\log _{10} K^\circ = 1.002 \pm 0.046, \log _{10} \beta^\circ = 3.859 \pm 0.048\)). Specific ion interaction coefficients \(\varepsilon({\rm L}^{-},{\rm Na}^{+})=0.04{2}\pm 0.02{2} {\rm kg}\cdot {\rm mol}^{-1}, \varepsilon({\rm L}^{-},{\rm K}^{+})=0.08{1}\pm 0.02{1} {\rm kg}\cdot {\rm mol}^{-1}, \varepsilon ({\rm PdL}^+ ,{\rm ClO}_4^-) = 0.92 \pm 0.21 {\rm kg}\cdot {\rm mol}^{-1}\) were derived.

Quantitative description of the chloride effect on chlorine dioxide generation from the ClO2‐HOCl reaction

AbstractThe generation of chlorine dioxide from the reaction between hypochlorous acid and chlorite with or without an initial chloride addition has been studied under slightly acidic conditions. Chloride (Cl‐), one of the products from the reaction, not only changes the reaction stoichiometry, but also alters the rate law. It was found that the formation of chlorine dioxide from the HOCI‐ClO2‐ system consists of two distinct parts, one is promoted by chloride, the other is independent of chloride. The overall kinetics of the chlorine dioxide generation from the reaction is: This model can very well predict the reaction under the following conditions: \documentclass{article}\pagestyle{empty}\begin{document}$ 273 - 303{\rm K},\left[ {ClO_2^ - } \right] = 0.001 - 0.00267\;{\rm mol}/{\rm L},\left[ {ClO_2^ - } \right]/\left[ {HOCl} \right] = 2,\;{\rm and}\;{\rm pH}\;3.86 - 4.91 $\end{document}

Acid Rain and Ozone Effects on Gas Exchange of Pinus ponderosa: A Comparison between Trees and Seedlings

We studied the effects of increased rain acidity and ozone on net photosynthesis (A), internal CO2 concentration $(C_{i})$, stomatal conductance (g1) and instantaneous water use efficiency (A/E) of tree clones and their half-sib seedlings of a ponderosa pine (Pinus ponderosa Dougl. ex Laws.) genotype in the field. Branches of trees and whole seedlings were exposed to rain (pH 5.1 and 3.0) and ozone (charcoal filtered, ambient, and twice-ambient) treatments using open-bottom, 1.5-m-long × 0.7-m-diameter chambers. The rain treatment was applied weekly from January 1992 to April 1992, to foliage only. The ozone treatment was applied from September 1991 to November 1992. Midday charcoal filtered, ambient, and twice-ambient ozone concentrations were 0.025, 0.038, and 0.075 μmol mol-1 in February, peaked at 0.028, 0.058, and $0.115 \mu {\rm mol} {\rm mol}^{-1}$ in July, and decreased to the lowest values at 0.01, 0.02, and $0.044 \mu {\rm mol} {\rm mol}^{-1}$ in November, respectively. Monthly midday gas-ex...

Leaf and Plant Water use Efficiency of C4Species Grown at Glacial to Elevated CO2Concentrations

Leaf gas exchange was measured on C4 plants grown from near glacial to current CO2 concentrations $(200-350 \mu {\rm mol} {\rm mol}^{-1})$ and from the current concentration to possible future levels (near 700 and $1000 \mu {\rm mol} {\rm mol}^{-1}$) to test the prediction that intrinsic water use efficiency ( CO2 assimilation [A]/stomatal conductance to water [g]) would rise by a similar relative amount as CO2 concentration. Studied were species differing in growth form or life history, the perennial grass Schizachyrium scoparium (little bluestem), perennial shrub Atriplex canescens (four-wing saltbush), and annual grass Zea mays (maize). Contrary to our prediction, leaf A/g of the C4 species examined was stimulated proportionally more by a given relative increase in CO2 over subambient than by elevated concentrations. The ratio of the relative increase in A/g to that in CO2 exceeded unity in S. scoparium and, in 1 of 2 yr, in Z. mays as CO2 rose from 200 to $350 \mu {\rm mol} {\rm mol}^{-1}$, but ...

Dynamics of ternary complex formation in the reaction of diaquoanthranilato-N, N-diacetatonickelate(II) with 2,2?-bipyridine and 1,10-phenanthroline

Dynamics of ternary complex formation in the reaction of diaquoanthranilato-N, N-diacetatonickelate(II) with 2,2′-bipyridine and 1,10-phenanthroline. $\rm Ni(ada)(H_2O)_2^{-}$ $+$ $L\rightleftharpoons Ni(ada)(L)^{-}$ $+$ $2 H_20;$ $- {{d[Ni(ada)^{-}]}\over{dt}}$ $=$ $k_f[Ni(ada)^{-}][L]+k_d\ [Ni(ada)(L)];$ $\ ada^{3-}=$anthranilate-N, N-diacetate; and L=bipy or phen. The kinetics of formation of ternary complexes by diaquoanthranilato-N, N-diacetatonickelate(II). [Ni(ada)(H2O)]− with 2,2′-bipyridine (bipy) and 1,10-phenanthroline (phen) have been studied under pseudo-first-order conditions containing excess bipy or phen by stopped-flow spectrophotometry in the pH range 7.1–7.8 at 25°C and λ = 0.1 mol dm−3. In each case, the reaction is first-order with respect to both Ni(ada)− and the entering ligand (ie., bipy, phen). The reactions are reversible. The forward rate constants are: $k^{\rm Ni(ada)}_{\rm Ni(ada)(bipy)}=0.87\times10^3{\rm dm}^3 {\rm mol}^{-1}{\rm s}^{-1}$,. $k^{\rm Ni(ada)}_{\rm Ni(ada)(phen)}=1.87\times10^3{\rm dm}^3 {\rm mol}^{-1}{\rm s}^{-1}$; and the reverse rate constants are: $k^{\rm Ni(ada)(bipy)}_{\rm Ni(ada)}=1.0{\rm s}^{-1}$ and $k^{\rm Ni(ada)(phen)}_{\rm Ni(ada)}=2.0{\rm s}^{-1}$. The corresponding stability constants of ternary complex formation are: and , . The observed rate constants and huge drops in stability constants in ternary complex formation agree well with the mechanism in which dissociation of an acetate arm of the coordinated ada3− prior to chelation by the aromatic ligand occurs. The observations have been compared with the kinetics of ternary complex formation in the reaction Ni(ada)− - glycine in which the kinetics involves a singly bonded intermediate, N(ada)((SINGLE BOND)O(SINGLE BOND)N)2− in rapid equilibrium with the reactants followed by a sluggish ring closure step. The reaction with the aromatic ligands conforms to a steady-state mechanism, while for glycine it gets shifted to an equilibrium mechanism. The cause of this difference in mechanistic pathways has been explained. © 1996 John Wiley & Sons, Inc.

Photosynthetic and Water Relations Responses to Elevated CO<sub>2</sub> in the C<sub>4</sub> Grass Andropogon gerardii

Undisturbed tallgrass prairie, dominated by the C₄ grass Andropogon gerardii, was exposed to ambient and elevated (double ambient) levels of atmospheric CO₂ in large open-top chambers throughout the 1991 and 1992 growing seasons. Responses in leaf xylem pressure potential (ψ), net photosynthesis (A), and stomatal conductance (g) were measured in both years for A. gerardii grown within chambers and from adjacent field plots. In 1992, maximum photosynthetic capacity (A_max), apparent quantum requirement (Q_r), the photosynthetic light compensation point (LCP), and dark respiration (<tex-math>$R_{d}$</tex-math>) were also measured. Midday ψ was significantly higher in plants grown at elevated CO₂ in both years, and seasonally averaged ψ was 0.48-0.70 MPa lower in 1991 (a dry year) than 1992 (a wet year). In 1991, A and g were significantly higher (regardless of measurement CO₂ level) in plants grown at elevated vs. ambient CO₂. These increases were measured in well-watered plants insuring that these plants differed only in CO₂ growth conditions and previous exposure to low ψ. Increased A at elevated CO₂ occurred (as much as <tex-math>$7.1\ \mu {\rm mol}\ {\rm m}^{-2}\ {\rm s}^{-1}$</tex-math>) over a broad range of temperatures (17-35 C), but the temperature optimum for A was similar at both 350 and <tex-math>$700\ \mu {\rm L}\ {\rm L}^{-1}\ {\rm CO}_{2}$</tex-math>. In 1992, no differences in A, A_max, Q_r, LCP, or <tex-math>$R_{d}$</tex-math> were detected when ambient and elevated CO₂ plants were compared. In plants collected from field plots, <tex-math>$R_{d}$</tex-math>, LCP, and leaf N were significantly higher than in plants within the chambers indicating that a chamber effect exists for these parameters. In both years, g was significantly reduced (21%-51%) when measured at 700 vs. <tex-math>$350\ \mu {\rm L}\ {\rm L}^{-1}\ {\rm CO}_{2}$</tex-math>. Peak aboveground biomass was increased at elevated CO₂ in 1991 but not in 1992. These data indicate that for C₄ grasses, effects of elevated CO₂ may only be detectable in years with significant water stress, a common occurrence in the central North American tallgrass prairies.

Growth and Gas Exchange of Oats (Avena sativa) and Wild Mustard (Brassica kaber) at Subambient CO&lt;sub&gt;2&lt;/sub&gt; Concentrations

A repeated sequence of monocultures and mixtures of oats (Avena sativa L.) and wild mustard (Brassica kaber (DC.) Wheeler) was grown along a daytime gradient of CO 2 concentrations ([ CO 2 ]) from near 330 to a minimum of $150 \mu {\rm mol} {\rm mol}^{-1}$ . The objectives were to determine effects of subambient [ CO 2 ] on leaf gas exchange, biomass production, and competitive interactions of these C 3 species. A decrease in stomatal conductance did not prevent a nearly linear increase in leaf internal [ CO 2 ] and net assimilation of oat leaves as [ CO 2 ] increased. Net assimilation of oats and wild mustard increased from 5.0 and $2.5 \mu {\rm mol} {\rm m}^{-2} {\rm s}^{-1}$ at $150 \mu {\rm mol} {\rm mol}^{-1}$ , respectively, to 16.1 and $15.9 \mu {\rm mol} {\rm m}^{-2} {\rm s}^{-1}$ at $330 \mu {\rm mol} {\rm mol}^{-1}$ CO 2</s...

The pyrolysis kinetics of 4‐chloro‐2‐butanone in the gas phase

AbstractThe rates of pyrolysis of 4‐chloro‐2‐butanone in the gas phase have been determined in a static system seasoned with the products of decomposition of allyl bromide. The reaction is catalyzed by hydrogen chloride. Under maximum catalysis of HCl, the kinetics were found to be of order 1.5 in the substrate suggesting that a complex elimination is involved. The reaction, when maximally inhibited with propene, appears to undergo a unimolecular elimination and follows a first‐order law kinetics. The products are methylvinyl ketone and hydrogen chloride. The kinetics have been measured over the temperature range of 402.0–424.4°C.The rate coefficients are given by the Arrhenius equation \documentclass{article}\pagestyle{empty}\begin{document}$ \log k_1 (\sec ^{ - 1}) = (13.67 \pm 0.69) - (225.2 \pm 8.6)\,{\rm kj}/{\rm mol}/2.303RT\angle $\end{document}. Thepyrolysis of 4‐chloro‐2‐butanone is 31 times greater in rate than that of ethyl chloride at 440°C. This large difference in rate may be attributed to the ‐M effect of the acetyl substituent in the pyrolysis of the former halo compound.

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

AbstractThe 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.

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

AbstractThe 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.

Oxygen vibrational and dissociation relaxation behind regular reflected shocks

The oxygen vibrational and dissociation relaxation behind regular reflected shocks has been calculated and measured. Numerical calculations using published rate coefficients supplied the relaxation-zone data needed to estimate the range of most useful experimental conditions. Then photographs of the shock reflexion were taken using a complementary double-exposure interferometer. The density profiles in the relaxation zones behind the reflected shocks were measured by means of a multibeam laser-differential interferometer. The results of these experiments confirmed the theoretical model adopted for the calculations within a certain range of experimental conditions, but clearly revealed the need for revising the rate coefficients. New calculations with different vibrational relaxation times and dissociation rate coefficients then had the result that the best fit of calculated to measured profiles was obtained when the following values were inserted.Vibration\begin{eqnarray*} &amp; p\tau_v = A_v\exp(B_vT^{-\frac{1}{3}}),\\ &amp; A_v = (2.1\pm 0.2)\times 10^{-5}\,{\rm kg/ms},\quad B_v = 129\,{}^{\circ}{\rm K}^{\frac{1}{3}}. \end{eqnarray*}Dissociation: O2+ O_2[rlarr ] 2O + O_2\begin{eqnarray*} &amp; {\mathop {k_1}\limits^{\rightharpoonup}} = A_1T^{-2.5}\exp (-\theta_D/T),\\ &amp; A_1 = (6.2 \pm 0.5)\times 10^{18}\,{\rm m}^3\,{}^{\circ}{\rm K}^{2.5}/{\rm mol}\,{\rm s},\quad\theta_D = 59\,136\,{}^{\circ}{\rm K}. \end{eqnarray*}Dissociation: O2+ O[rlarr ]3O\begin{eqnarray*} &amp; {\mathop {k_1}\limits^{\rightharpoonup}} = A_2T^{-1.0}\exp (-\theta_D/T),\\ &amp; A_2 = (4.0 \mp 0.5)\times 10^{-13}\,{\rm m}^3\,{}^{\circ}{\rm K}/{\rm mol}\,{\rm s}. \end{eqnarray*}