Die Kristallstruktur von Cr3GeN

The crystal structure of Cr₃GeN has been determined by single crystal methods: $$\begin{gathered} Space group P \bar 42_1 m - D_{2d}^3 ; \hfill \\ a = 5.375{\AA}; c = 4.012 {\AA}; c/a = 0.7465 \hfill \\ \end{gathered} $$ The close relationship of Cr₃GeN with the phases having filled U₃Si-type and the perowskite carbides and-nitrides will be discussed.     

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.     

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

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

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.     

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.     

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: $$\begin{gathered} Pb^{2 + } + L^{2 - } \rightleftharpoons PbL log\beta _{101} = 3.62 \pm 0.03 \hfill \\ Pb^{2 + } + H^ + + L^{2 - } \rightleftharpoons PbHL^ - log\beta _{111} = 6.30 \pm 0.07 \hfill \\ \end{gathered} $$      

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.     

Beiträge zur Chemie des Sulfamids, 1. Mitt.: Zur Konstitution des Disilber-sulfamids

According to spectroscopic (IR, broadline proton NMR) and chemical (alkylation) investigations of disilver sulphamide the following molecular structure is assumed: $$\begin{gathered} O \hfill \\ || \hfill \\ H_2 N\_\_S\_\_NAg| \hfill \\ OAg \hfill \\ \end{gathered}$$ From the IR and NMR data deduction concerning the nature of the chemical bonds in this compound is possible. The instability of the still unknown mono-and trisilver sulphamide is discussed with regard to the structure of disilver sulphamide.     

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.     

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 $$\begin{gathered} 2 MCl_3 AN + 6AN \rightleftharpoons [MCl_2 (AN)_4 ]^ + [MCl_4 ]^ - + 4 AN \rightleftharpoons \hfill \\ 2 [MCl_2 (AN)_4 ]^ + Cl - \hfill \\ \end{gathered} $$      

Thermodynamic excess properties of ternary alcohol-hydrocarbon systems. Simplified method for predicting enthalpies from binary data

Two general equations for estimation of excess enthalpies of ternary systems consisting of an alcohol and two hydrocarbons from observed excess properties of the various binary combinations have been developed. The first expression is based on the Kretschmer-Wiebe association model and takes the form $$\Delta \overline H _{ABC}^{ex} = h_A x_A K_A (\phi _{A1} - \phi _{A1}^o ) + Q_{ABC}$$ where $$\begin{gathered} Q_{ABC} = (x_A + x_B )(\phi _A + \phi _B )(\Delta \overline H _{AB}^{ex} )_{phys}^ \bullet + (x_A + x_C )(\phi _A + \phi _C )(\Delta \overline H _{AC}^{ex} )_{phys}^ \bullet \hfill \\ + (x_B + x_C )(\phi _B + \phi _C )(\Delta \overline H _{BC}^{ex} )_{phys}^ \bullet \hfill \\ \end{gathered}$$ \((\Delta \overline H _{ij}^{ex} )_{phys}^ \bullet\) represents the physical interactions in each of the individual binary systems, and the term involving φ _A1 ^o represents the chemical contributions (caused by self-association) to the excess enthalpies of mixing. The second predictive expression is based on the Mecke-Kempter association model and is given by $$\Delta \overline H _{ABC}^{ex} = - h_A x_A [In(1 + K_A \phi _A )/K_A \phi _A - In(1 + K_A )/K_A ] + Q_{ABC}$$ where the first term (contained within brackets) represetns the chemical contributions to the enthalpies of mixing. The predictions of both expressions are compared with experimental data for the excess enthalpies of six ternary systems.     

Thermal properties of Na2TeO4(s) and TiTe3O8(s)

Molar heat capacity measurement on Na₂TeO₄(s) and TiTe₃O₈(s) were carried out using differential scanning calorimeter. The molar heat capacity values were least squares analyzed and the dependence of molar heat capacity with temperature for Na₂TeO₄(s) and TiTe₃O₈(s) can be given as, $$ \begin{gathered} {\text{C}}^{\text{o}}_{{{\text{p}},{\text{m}}}} \left\{ {{\text{Na}}_{ 2} {\text{TeO}}_{ 4} \left( {\text{s}} \right)} \right\} \,={159}.17 { } + 1.2\,\times\,10^{-4}T-{55}.34\,\times\,10^{5}/T^{2};\hfill \\ C^{\text{o}}_{{{\text{p}},{\text{m}}}} \left\{ {{\text{TiTe}}_{ 3} {\text{O}}_{ 8} \left( {\text{s}} \right)} \right\}\,=\,{ 275}.22{ }+{4}.0\,\times\, 10^{-5}T-{58}.28\,\times\,10^{5}/T^{2};\hfill \\ \end{gathered} $$ From this data, other thermodynamic functions were evaluated.     

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 $$\begin{gathered} log Q'_W = log K_W + 2.0AI^{1/2} /(1 + I^{1/2} ) - [p_1 + p_2 /T + p_3 T^2 + p_4 F(I)]I \hfill \\ - 0.0157\phi m_{NaCl} \hfill \\ F(I) = [1 - (1 + 2I^{1/2} - 2I) exp( - 2I^{1/2} )]/4I \hfill \\ \end{gathered} $$      

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.     

Vacuum sublimation kinetics of urea nitrate

Urea nitrate completely sublimes in a continuously pumped vacuum at a rate dependent upon the extent of the surface area. The fraction sublimed (α) vs. time (t) curve is sigmoidal in shape with an inflection point fluctuating between 5.3 and 8.1 % weight loss in the temperature range 56 to 97°. The experimental data fit the grain burning model $$\begin{gathered} 1 - (1 - \beta )^{{1 \mathord{\left/ {\vphantom {1 3}} \right. \kern-\nulldelimiterspace} 3}} = k_s (t - t_0 ) \hfill \\ \beta = \alpha - \beta _0 ,\beta = 0 at inf{\text{lection}} {\text{point}} \hfill \\ \end{gathered} $$ where k_s is a sublimation rate constant, and the constantsβ ₀ andt ₀ are the respective values ofα andt at the inflection point. This equation yielded an activation enthalpy of sublimation of 79±5.4kJ/mole.     

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.     

Spektrophotometrische Bestimmung der Dissoziationskonstanten von Chloranilsäure

The ionization constants of chloranilic acid are determined at 25° C by the spectrophotometric method. The following values are obtained: $$\begin{gathered} pk_1 pk_2 ionic strength [m] \hfill \\ 0.762.580.50 \hfill \\ 0.972.552.00 \hfill \\ \end{gathered}$$      

Bond dissociation energies and bond orders for diatomic alkali halides

The bond dissociation energies for Alkali halides have been estimated based on the derived relations: $$\begin{gathered} D_{AB} = \bar D_{AB} + 31.973{\text{ e}}^{0.363\Delta x} {\text{ and}} \hfill \\ D_{AB} = \bar D_{AB} (1 - 0.2075\Delta xr_e ) + 52.29\Delta x, \hfill \\ \end{gathered} $$ where \(\bar D_{AB} = (D_{AA} \cdot D_{BB} )^{{1 \mathord{\left/ {\vphantom {1 2}} \right. \kern-0em} 2}} \) , Δx represents Pauling electronegativity differences x(_A ?x_B) and r _e is the internuclear distance. A simplified formula relating bond orders, q, to spectroscopic constants is suggested. The formula has the form q = 1.5783 × 10^?3 (ω _e ² r_e/ B_e)^1/2. The ambiguity arising from the Parr and Borkman relation is discussed. The present study supports the view of Politzer that q/(0.5r _e)² is the correct definition of bond order. The estimated bond energies and bond orders are in reasonably good agreement with the literature values. The bond energies estimated with the relations we suggested, for alkali halides give an error of 4.5% and 5.3%, respectively. The corresponding error associated with Pauling's equation is 40.2%.     

Paradigms and Paradoxes: Diazomethane and Ethyl Diazoacetate: The Role of Substituent Effects on Stability

Diazomethane and ethyl diazoacetate are highly reactive and highly versatile synthetic reagents that undergo numerous related reactions. However, while the former is highly dangerous because of its toxicity and explosive behavior; the latter is much more benign. This is usually ascribed to resonance stabilization in ethyl diazoacetate involving an extra carbonyl group that is absent in diazomethane, cf. $$\begin{gathered} {\text{EtOOC}}---{\text{CH}} = {\rm N}^ + = {\rm N}^ - \leftrightarrow {\rm E}{\text{tOOC}}---{\text{CH}}^ - ---{\text{N}}^{\text{ + }} \equiv {\text{N}} \leftrightarrow {\text{EtOC(O}}^ - {\text{)}} = {\text{CH}}---{\text{N}}^{\text{ + }} \equiv {\rm N} \hfill \\ {\text{CH}}_{\text{2}} = {\rm N}^ + = {\rm N}^ - \leftrightarrow {\text{CH}}_{\text{2}}^ - ---{\text{N}}^{\text{ + }} \equiv {\rm N} \hfill \\ \end{gathered}$$ The additional resonance stabilization is derived using a recent literature measurement of the enthalpy of an ethyl diazoacetate/aldehyde reaction, key enthalpies of formation, also from the literature, and some simplifying assumptions. The resonance stabilization is deduced to be but 16 kJ/mol, merely 4 kcal/mol. But, oh how grateful we are for this!