0-aminophenol and 8-hydroxyquinoline as ligands of Cu(II) in 1M-Na(ClO4) at 25°C

The complex formation between Cu(II) and 8-hydroxyquinolinat (Ox) was studied with the liquid-liquid distribution method, between 1M-Na(ClO₄) and CHCl₃ at 25°C. The experimental data were explained by the equilibria: $$\begin{gathered} \operatorname{Cu} ^{2 + } + Ox \rightleftharpoons \operatorname{Cu} Ox \log \beta _1 = 12.38 \pm 0.13 \hfill \\ \operatorname{Cu} ^{2 + } + 2 Ox \rightleftharpoons \operatorname{Cu} Ox_2 \log \beta _2 = 23.80 \pm 0.10 \hfill \\ \operatorname{Cu} Ox_{2aq} \rightleftharpoons \operatorname{Cu} Ox_{2\operatorname{org} } \log \lambda = 2.06 \pm 0.08 \hfill \\ \end{gathered} $$ The equilibria between Cu(II) and o-aminophenolate (AF) were studied potentiometrically with a glass electrode at 25°C and in 1M-Na(ClO₄). The experimental data were explained by the equilibria: $$\begin{gathered} \operatorname{Cu} ^{2 + } + AF \rightleftharpoons \operatorname{Cu} AF \log \beta _1 = 8.08 \pm 0.08 \hfill \\ \operatorname{Cu} ^{2 + } + 2AF \rightleftharpoons \operatorname{Cu} AF_2 \log \beta _2 = 14.60 \pm 0.06 \hfill \\ \end{gathered} $$ The protonation constants ofAF and the distribution constants between CHCl₃?H₂O and (C₂H₅)₂O?H₂O were also determined.     

Komplexe zwischen Eisen(II)- und Tartrat-Ion

The formation of complexes between iron(II) and tartrate ion (L) has been studied at 25° C in 1m-NaClO₄, by using a glass electrode. The e.m.f. data are explained with the following equilibria: $$\begin{gathered} Fe^{2 + } + L \rightleftarrows FeL log \beta _1 = 1,43 \pm 0,05 \hfill \\ Fe^{2 + } + 2L \rightleftarrows FeL_2 log \beta _2 = 2,50 \pm 0,05 \hfill \\\end{gathered} $$ The protonation constants of the tartaric acid have been determinated: $$\begin{gathered} H^ + + L \rightleftarrows HL logk_1 = 3,84 \pm 0,03 \hfill \\ 2H^ + + L \rightleftarrows H_2 L logk_2 = 6,43 \pm 0,02 \hfill \\\end{gathered}$$ .     

Komplexe aus Kupfer(II) und Acetylaceton

The complex formation between copper(II) and acetylacetonate (L)^* has been studied by potentiometry and distribution between CHCl₃ and water. The experimental data are interpreted by postulating the following equilibria: $$\begin{gathered} Cu^{2 + } + L \rightleftharpoons CuL1g \beta _1 = 8.42 \pm 0.10 \hfill \\ Cu^{2 + } + 2 L \rightleftharpoons CuL_2 1g \beta _2 = 15.47 \pm 0.10 \hfill \\ \left( {CuL_2 } \right)_{aq} \rightleftharpoons \left( {CuL_2 } \right)_0 1g \lambda _B = 1.80 \pm 0.10 \hfill \\ \end{gathered} $$ In order to study the complex formation, the protonation constant (k) of acetylacetonate and the distribution coefficient λ _A of acetylacetone in the same experimental conditions were required. It was found: lgk=9.05±0.03; λ _A = 1.20 ± 0.02.     

Polyhydroxysäuren als Komplexbildner

The reaction of mucic acid (H₆ Mu) with Cobalt(II) and Nickel(II) ions has been studied in 1.0M-Na⁺(NO ₃ ^? ) ionic medium at 25° C using a glass electrode. The e.m.f. data in the range 8≦?log [H⁺]≦10 are explained by assuming $$\begin{gathered} Me^{2 + } + H_4 Mu^{2 - } \rightleftharpoons MeH_3 Mu^ - + H^ + \beta ''_1 \hfill \\ Me^{2 + } + H_4 Mu^{2 - } \rightleftharpoons MeH_2 Mu^{2 - } + 2 H^ + \beta ''_2 \hfill \\ \end{gathered}$$ with equilibrium constants log β′₁ = — 9.36; — 9.34; log β′₂ = — 18.11; — 18.08 for Co(II) and Ni(II) resp.     

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 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.     

Protonierungskonstanten der 8-Hydroxychinolinat-Ionen bei 25°C und in 1m-NaClO4

The protonation of the 8-hydroxyquinolinate ion (Ox ^?) has been studied at 25°C in 1m-NaClO₄ by the potentiometric method and the distribution between CHCl₃ and H₂O. The experimental data are explained by the following equilibria: $$\begin{array}{*{20}c} {H^ + + Ox^ - \rightleftharpoons HOx} \\ {H^ + + Ox \rightleftharpoons H_2 Ox^ + } \\ {HOx_w \rightleftharpoons HOx_{org} } \\ \end{array} \begin{array}{*{20}c} {\log k_1 = 9.42 \pm 0.08} \\ {\log k_2 = 5.46 \pm 0.10} \\ {\log \lambda = 2.40 \pm 0.10} \\ \end{array} $$      

Die Ionisation von Triphenylchlormethan durch Metallchloride in Acetonitril

In acetonitrile (AN) solutions the gross constants are determined for the reactions $$Ph_3 CCl + MCl_n ANPh_3 C^ + MCl_{n + 1}^ - + AN$$ (MCl _n =SbCl₅, GaCl₃, InCl₃, and FeCl₃). The relaxation spectra are interpreted for the reactions of metal(III) chlorides according to the equilibria $$\begin{gathered} 2 MCl_3 AN + 6AN \rightleftharpoons [MCl_2 (AN)_4 ]^ + [MCl_4 ]^ - + 4 AN \rightleftharpoons \hfill \\ 2 [MCl_2 (AN)_4 ]^ + Cl - \hfill \\ \end{gathered} $$      

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: $$\begin{gathered} KSO_4^ - : - log K_m = - 1.03 + 1.04 \times 10^4 P \pm 0.020 \Delta \bar V^o = - 5.9 \hfill \\ NaSO_4^ - : - log K_m = - 1.00 + 1.30 \times 10^4 P \pm 0.019 \Delta \bar V^o = - 7.3 \hfill \\ MgCl^ + : - log K_m = - 0.75 + 0.71 \times 10^4 P \pm 0.028 \Delta \bar V^o = - 4.0 \hfill \\ \end{gathered} $$      

Die Kristallstruktur des Chromarsenids Cr4As3

The crystal structure of Cr₄As₃ has been determined by single crystal photographs: $$\begin{gathered} space group Cm - C_s ^3 \hfill \\ \alpha = 13.16_8 {\AA} \hfill \\ b = 3.54_2 {\AA} \hfill \\ c = 9.30_2 {\AA} \hfill \\ \beta = 102.1_9 \circ \hfill \\ \end{gathered}$$ Cr₄As₃ crystallizes with a novel structure type, which can be derived from the MnP-structure type.     

The ionization of water in NaCl media to 300°C

The apparent ionization quotient for water has been measured potentiometrically near the saturation pressure from 25 to 295°C in 1 and 3m NaCl using a previously described hydrogen-electrode concentration cell. The results are presented in terms of a modification of the Brönsted-Guggenheim treatment for activity coefficients of ions. The mathematical form of the temperature dependence for the interaction coefficients was indicated by the more extensive data on \(\gamma _{{\rm H}^ + } \gamma _{{\rm O}{\rm H}^ - } \) in KCl media. From a least-squares analysis of these data in NaCl along with the very precise data from the literature for NaCl media from 0 to 50°C, the following expression for the effect of salt concentration and temperature on logQ _w ^′ is obtained $$\begin{gathered} log Q'_W = log K_W + 2.0AI^{1/2} /(1 + I^{1/2} ) - [p_1 + p_2 /T + p_3 T^2 + p_4 F(I)]I \hfill \\ - 0.0157\phi m_{NaCl} \hfill \\ F(I) = [1 - (1 + 2I^{1/2} - 2I) exp( - 2I^{1/2} )]/4I \hfill \\ \end{gathered} $$      

Solvent extraction of rare earth metal ions with 1-(2-pyridylazo)-2-naphthol (PAN)

The solvent extraction of Yb(III) and Ho(III) by 1-(2-pyridylazo)-2-naphthol (PAN or HL) in carbon tetrachloride from aqueous-methanol phase has been studied as a function ofpH ^× and the concentration ofPAN or methanol (MeOH) in the organic phase. When the aqueous phase contains above ～25%v/v of methanol the synergistic effect was increased. The equation for the extraction reaction has been suggested as: $$\begin{gathered} Ln(H_2 0)_{m(p)}^{3 + } + 3 HL_{(o)} + t MeOH_{(o)} \mathop \rightleftharpoons \limits^{K_{ex} } \hfill \\ LnL_3 (MeOH)_{t(o)} + 3 H_{(p)}^ + + m H_2 0 \hfill \\ \end{gathered} $$ where:Ln ³⁺=Yb, Ho; $$\begin{gathered} t = 3 for C_{MeOH in.} \varepsilon \left( { \sim 25 - 50} \right)\% {\upsilon \mathord{\left/ {\vphantom {\upsilon \upsilon }} \right. \kern-\nulldelimiterspace} \upsilon }; \hfill \\ t = 0 for C_{MeOH in.} \varepsilon \left( { \sim 5 - 25} \right)\% {\upsilon \mathord{\left/ {\vphantom {\upsilon \upsilon }} \right. \kern-\nulldelimiterspace} \upsilon } \hfill \\ \end{gathered} $$ . The extraction equilibrium constants (K _ex ) and the two-phase stability constants (β ₃ ^× ) for theLnL ₃(MeOH)₃ complexes have been evaluated.     

Lattice parameters and thermal expansion of δ-VN1−x from 298–1000 K

The thermal expansion of VN_1?x was determined from measurements of the lattice parameters in the temperature range of 298–1000 K and in the composition range of VN_0.707–VN_0.996. Within the accuracy of the results the expansion of the lattice parameter with temperature is not dependent on the composition. The lattice parameter as a function of composition ([N]/[V]=0.707?0.996) and temperature (298–1000 K) is given by $$\begin{gathered} a([N]/[V],T) = 0.38872 + 0.02488([N]/[V]) - \hfill \\ - (1.083 \pm 0.021) \cdot 10^{ - 4} T^{1/2} + (6.2 \pm 0.1) \cdot 10^{ - 6} T. \hfill \\ \end{gathered} $$ . The coefficient of linear thermal expansion as a function of temperature (in the same range) is given by $$\alpha (T) = a([N]/[V],T)^{ - 1} [( - 5.04 \pm 0.01) \cdot 10^5 T^{ - 1/2} + (6.2 \pm 0.1) \cdot 10^{ - 6} ].$$ . The average linear thermal expansion coefficient is $$\alpha _{av} = 9.70 \pm 0.15 \cdot 10^{ - 6} K^{ - 1} (298 - 1 000K).$$ . The data are compared with those of several fcc transition metal nitrides collected and evaluated from the literature.     

Vanadyl complexes with ethylenedithiodiacetic acid

The complex formation between vanadyl ions and ethylenedithiodiacetic acid (H₂ L) has been studied at 25 °C in 0.5M-NaClO₄ as inert medium, by measuring the hydrogen ion concentration with a glass electrode. In acidic medium and in the investigated concentration ranges (2.07mM≤C _M≤7.50mM, C _L up to 12.5mM) the emf data can be explained assuming the equilibrium: $$VO^{2 + } + L^{2 - } \rightleftharpoons VOL log\beta _{101} = 2.68 \pm 0.03$$      

Dichlorosilylene: Rate constant for reaction with carbon monoxide and nitrous oxide

The absolute rate constanss for the gas-phase reactions of 1,1-dichlorosilylene with carbon monoxide and nitrous oxide have been determined using the flash photolysts-kinetic absorpiton spectroscopy technique. The bimolecular rate constant values at 25° C are: $$\begin{gathered} k\left( {Cl_2 Si + CO} \right) = \left( {6.3 \pm 0.7} \right) \times 10^8 M^{ - 1} s^{ - 1} \hfill \\ k\left( {Cl_2 Si + N_2 O} \right) = \left( {5.7 \pm 0.3} \right) \times 10^8 M^{ - 1} s^{ - 1} \hfill \\ \end{gathered} $$      

Calibration of conductance cells at various temperatures

Precise conductance measurements on aqueous potassium chloride solutions at 0, 10, 18, and 25°C have been made under various conditions over a concentration range 10^?4?2 mole-dm^?3, yielding the conductance equations $$\begin{gathered} 25^\circ C:\Lambda = 149.873 - 95.01\sqrt c + 38.48c log c + 183.1c - 176.4c^{3/2} \hfill \\ 18^\circ C:\Lambda = 129.497 - 80.38\sqrt c + 32.87c log c + 154.3c - 143.0c^{3/2} \hfill \\ 10^\circ C:\Lambda = 107.359 - 64.98\sqrt c + 27.07c log c + 125.4c - 110.3c^{3/2} \hfill \\ 0^\circ C:\Lambda = 81.700 - 47.80\sqrt c + 20.60c log c + 93.8c - 79.3c^{3/2} \hfill \\ \end{gathered} $$ which are proposed for calibration of conductance cells.     

KSO 4 − , NaSO 4 − , and MgCl+ ion pairs in aqueous solutions up to 2000 atm

Electrical conductance data at 25°C for K₂SO₄, Na₂SO₄, and MgCl₂ 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: $$\begin{gathered} {\text{ }}KSO_4^ - :log K_m = ( - 1.02{\text{ }} + 1.6 \times 10^{ - 4} P - {\text{ }}2.5 \times 10^{ - 8p2} ) \pm 0.03 \hfill \\ NaSO_4^ - :log K_m = ( - 1.02 + 9.6 \times 10^{ - 5} P - {\text{ }}4.3 \times 10^{ - 9p2} ) \pm 0.03 \hfill \\ MgCl^ + :log K_m = ( - 0.64 + 1.1 \times 10^{ - 4} P - {\text{ }}1.7 \times 10^{ - 8p2} ) \pm 0.04 \hfill \\ \end{gathered} $$ withP in atmospheres. These values cannot be chosen solely on the basis of minimizing errors in fitting conductance data to theoretical equations. For the values cited above, the Bjerrum distances for 1–2 (or 2-1) and 1-1 salts were used. However, the conductance fits for KSO ₄ ^? and NaSO ₄ ^? were equally good for half-Bjerrum distances and resulted in higher dissociation constants. Ultrasonic data are used to argue in favor of the lower dissociation values derived by using Bjerrum distances. Our results for MgCl⁺ disagree with those of Havel and Högfeldt.     

Ion association of dilute aqueous sodium hydroxide solutions to 600°C and 300 MPa by conductance measurements

The limiting molar conductances Λ₀ and ion association constants of dilute aqueous NaOH solutions (<0.01 mol-kg^?1) were determined by electrical conductance measurements at temperatures from 100 to 600°C and pressures up to 300 MPa. The limiting molar conductances of NaOH_(aq) were found to increase with increasing temperature up to 300°C and with decreasing water density ρ_w. At temperatures ≥400°C, and densities between 0.6 to 0.8 g-cm^?3, Λ₀ is nearly temperature-independent but increases linearly with decreasing density, and then decreases at densities <0.6 g-cm^?3. This phenomenon is largely due to the breakdown of the hydrogen-bonded, structure of water. The molal association constants K _Am for NaOH( _aq ) increase with increasing temperature and decreasing density. The logarithm of the molal association constant can be represented as a function of temperature (Kelvin) and the logarithm of the density of water by $$\begin{gathered} log K_{Am} = 2.477 - 951.53/T - (9.307 \hfill \\ - 3482.8/T)log \rho _{w } (25 - 600^\circ C) \hfill \\ \end{gathered} $$ which includes selected data taken from the literature, or by $$\begin{gathered} log K_{Am} = 1.648 - 370.31/T - (13.215 \hfill \\ - 6300.5/T)log \rho _{w } (400 - 600^\circ C) \hfill \\ \end{gathered} $$ which is based solely on results from the present study over this temperature range (and to 300 MPa) where the measurements are most precise.     

14 MeV neutron cross-sections on Mg,Cl and Sr

Cross-section values for 14.7 MeV neutrons have been measured for the following reactions: $$\begin{gathered} ^{24} Mg(n,p)^{24} Na - (187 \pm 7)mb;^{25} Mg(n,p)^{25} Na - (74 \pm 9)mb; \hfill \\ ^{26} Mg(n,\alpha )^{23} Ne - (55 \pm 6)mb;^{35} Cl(n,p)^{37} S - (22 \pm 3)mb; \hfill \\ ^{35} Cl(n,2n)^{34m} Cl - (9.3 \pm 1.5)mb;^{86} Sr(n,p)^{86m} Rb - (14 \pm 2)mb; \hfill \\ ^{88} Sr(n,p)^{88} Rb - (19 \pm 3)mb;^{86} Sr(n,2n)^{85m} Sr - (244 \pm 32)mb; \hfill \\ ^{88} Sr(n,2n)^{87m} Sr - (289 \pm 33)mb. \hfill \\ \end{gathered}$$ An analysis of available cross-section data for these reactions has been performed and preferred mean values for each reaction are given.     

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]: $$\begin{gathered} {}^{5 0}Cr(n,\gamma )^{5 1} Cr; \sigma _0 = (15.2 \pm 0.2) barn [cf.:15.8 - 16.0] \hfill \\ {}^{5 8}Fe(n,\gamma )^{5 9} Fe; \sigma _0 = (1.31 \pm 0.03) barn [cf.:1.14 - 1.16] \hfill \\ {}^{1 0 9}Ag(n,\gamma )^{1 1 0 m} Ag;\sigma _0 = (3.89 \pm 0.05) barn [cf.:4.4 - 5.0] \hfill \\ \end{gathered} $$