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

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.     

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.     

Cross-sections of 14 MeV neutron reactions on phosphorus and calcium

Cross-section values for 14.7 MeV neutrons are measured for the following reactions: $$\begin{gathered} ^{31} P(n,\alpha )^{28} Al,(132 \pm 10)mb;^{42} Ca(n,p)^{42} K,(173 \pm 19)mb; \hfill \\ ^{43} Ca(n,p)^{43} K,(111 \pm 9)mb;^{44} Ca(n,p)^{44} K,(42 \pm 2)mb; \hfill \\ ^{44} Ca(n,\alpha )^{41} Ar,(27 \pm 2)mb;^{48} Ca(n,2n)^{47} Ca,(616 \pm 54)mb. \hfill \\ \end{gathered}$$ An analysis of available cross-section data for these reactions is performed and preferred mean values for each reaction are given.     

Effects of washing conditions on the thermal decomposition behaviour of gelled UO3 microspheres

DTA, TG and DTG curves obtained in various atmospheres using different heating rates were used together with X-ray examinations to study the thermal decomposition mechanisms of two types of gelled UO₃ microspheres: ammonia-washed (UN) and hot water-washed (UH) microspheres. The kinetics of the thermal decompositions were studied. The specific reaction rate constantk _r for the decomposition of UO₃ to U₃O₈ could be expressed in terms of the activation energy and the pre-exponential factor by the expressions: $$\begin{gathered} K_r (s^{ - 1} ) = 1.277 \times 10^{18} \exp \frac{{ - 295.4}}{{RT}}for the UN spheres, \hfill \\ K_r (s^{ - 1} ) = 8.406 \times 10^{19} \exp \frac{{ - 263.2}}{{RT}}for the UH spheres. \hfill \\ \end{gathered} $$      

Determination of standard pH values for sodium hydrogen diglycolate buffer (0.05m) from 5–65°C

Values of pa _H ^o for 0.05 mole-kg^?1 aqueous solutions of sodium hydrogen diglycolate in the temperature range 5–65°C have been obtained from cells without transport, and can be fitted to the equation $$\begin{gathered} pa^\circ _H = 3.5098 + 2.222 \times 10^{ - 3} ({T \mathord{\left/ {\vphantom {T {K - 298.15}}} \right. \kern-\nulldelimiterspace} {K - 298.15}}) \hfill \\ + 2.628 \times 10^{ - 5} ({T \mathord{\left/ {\vphantom {T {K - 298.15}}} \right. \kern-\nulldelimiterspace} {K - 298.15}})^2 \hfill \\ \end{gathered} $$ The analysis has been carried out by a multilinear regression procedure using a form of the Clarke and Glew equation. This buffer standard may be a useful alternative to the saturated potassium hydrogen tartrate buffer.     

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.     

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.     

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 $$\begin{gathered} Eu^{3 + } + H_2 O + CO_2 (g)_ \leftarrow ^ \to EuCO_3^ + + 2H^ + \hfill \\ - log\beta _{12} = 9.607 + 496(t + 273.16)^{ - 1} \hfill \\ Eu^{3 + } + 2H_2 O + 2CO_2 (g)_ \leftarrow ^ \to Eu(CO_3 )_2^ - + 4H^ + \hfill \\ - log\beta _{24} = 21.951 + 670(t + 273.16)^{ - 1} \hfill \\ Eu^{3 + } + H_2 O + CO_2 (g)_ \leftarrow ^ \to EuHCO_3^{2 + } + H^ + \hfill \\ - log\beta _{11} = 1.688 + 1397(t + 273.16)^{ - 1} \hfill \\ \end{gathered}$$      

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.     

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

Copenhangen Standard Sea Water: Conductivity and salinity

In the present work, the conductivity of Standard Sea Water was studied over the temperature range 12–40°C using high-precision conductance techniques. One of the purposes of this investigation was to extend the temperature range of the equation that describes the conductivities tabulated by Cox. The conductivities reported in this work differ from the Cox equation by no more than 0.1% in the range of temperature overlap (to 30°C). The combined data base leads to the expression $$\begin{gathered} \kappa \left( {mmho - cm^{ - 1} } \right) = 29.05128 + 0.88082t - 1.98312 \times 10^{ - 4} t^2 + 3.33663 \hfill \\ \times 10^{ - 4} t^3 - 1.0776 \times 10^{ - 5} t^4 + 1.12518 \times 10^{ - 7} t^5 \hfill \\ \end{gathered} $$ for the electrical conductance of Copenhagen Standard Sea Water, with an overall uncertainty estimate (absolute accuracy) of ±0.1% for the temperature range of 0–40°C.     

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: $$\begin{gathered} H^ + + L \rightleftarrows HL logk_1 = 3,84 \pm 0,03 \hfill \\ 2H^ + + L \rightleftarrows H_2 L logk_2 = 6,43 \pm 0,02 \hfill \\\end{gathered}$$ .     

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.     

Kinetics of the initiated and inhibited oxidation of methyl oleate in homogeneous and aqueous emulsion media

The kinetics of the 2,2′-azobisisobutyronitrile-initiated oxidation of methyl oleate (MO) in the medium of the oxidized substrate itself (homogeneous system) and in an aqueous solution of cetyltrimethylammonium bromide (aqueous emulsion system, AES) was studied. The oxidation was found to occur as a nonbranched chain reaction with quadratic-law chain termination in both neat methyl oleate ([MO] ≈ 2.6 mol/l) and AES. The temperature dependence of the oxidizability parameter for methyl oleate in the temperature range from 303 to 333 K was described by the following expressions:

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.     

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.     

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.     

Potentiometric investigation of complexes between lead(II) and ethylenedithiodiacetic acid

Complex formation between lead(II) and ethylenedithio diacetic acid (H₂ L) has been studied at 25°C in aqueous 0.5M sodium perchlorate medium. Measurements have been carried out with a glass electrode and with a lead amalgam electrode. In acidic medium and in the investigated concentration range experimental data can be explained by assuming the following equilibria: $$\begin{gathered} Pb^{2 + } + L^{2 - } \rightleftharpoons PbL log\beta _{101} = 3.62 \pm 0.03 \hfill \\ Pb^{2 + } + H^ + + L^{2 - } \rightleftharpoons PbHL^ - log\beta _{111} = 6.30 \pm 0.07 \hfill \\ \end{gathered} $$