Thermochemistry of hydroxymethylphenol isomers

The standard (p° = 0.1 MPa) molar enthalpies of formation in the crystalline state of the 2-, 3- and 4-hydroxymethylphenols, $ {{\Updelta}}_{\text{f}} H_{\text{m}}^{\text{o}} ( {\text{cr)}} = \, - ( 3 7 7. 7 \pm 1. 4)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ , $ {{\Updelta}}_{\text{f}} H_{\text{m}}^{\text{o}} ( {\text{cr) }} = - (383.0 \pm 1.4) \, \,{\text{kJ}}\,{\text{mol}}^{ - 1} $ and $ {{\Updelta}}_{\text{f}} H_{\text{m}}^{\text{o}} ( {\text{cr)}} = - (382.7 \pm 1.4)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ , respectively, were derived from the standard molar energies of combustion, in oxygen, to yield CO₂(g) and H₂O(l), at T = 298.15 K, measured by static bomb combustion calorimetry. The Knudsen mass-loss effusion technique was used to measure the dependence of the vapour pressure of the solid isomers of hydroxymethylphenol with the temperature, from which the standard molar enthalpies of sublimation were derived using the Clausius–Clapeyron equation. The results were as follows: $ \Updelta_{\rm cr}^{\rm g} H_{\rm m}^{\rm o} = (99.5 \pm 1.5)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ , $ \Updelta_{\rm cr}^{\rm g} H_{\rm m}^{\rm o} = (116.0 \pm 3.7) \,{\text{kJ}}\,{\text{mol}}^{ - 1} $ and $ \Updelta_{\rm cr}^{\rm g} H_{\rm m}^{\rm o} = (129.3 \pm 4.7)\,{\text{ kJ mol}}^{ - 1} $ , for 2-, 3- and 4-hydroxymethylphenol, respectively. From these values, the standard molar enthalpies of formation of the title compounds in their gaseous phases, at T = 298.15 K, were derived and interpreted in terms of molecular structure. Moreover, using estimated values for the heat capacity differences between the gas and the crystal phases, the standard (p° = 0.1 MPa) molar enthalpies, entropies and Gibbs energies of sublimation, at T = 298.15 K, were derived for the three hydroxymethylphenols.     

Thermodynamic studies on Pr2TeO6

The standard Gibbs energy of formation of Pr₂TeO₆ $ (\Updelta_{\text{f}} G^{^\circ } \left( {{ \Pr }_{ 2} {\text{TeO}}_{ 6} ,\;{\text{s}}} \right)) $ was derived from its vapour pressure in the temperature range of 1,400–1,480 K. The vapour pressure of TeO₂ (g) was measured by employing a thermogravimetry-based transpiration method. The temperature dependence of the vapour pressure of TeO₂ over the mixture Pr₂TeO₆ (s) + Pr₂O₃ (s) generated by the incongruent vapourization reaction, Pr₂TeO₆ (s) = Pr₂O₃ (s) + TeO₂ (g) + ½ O₂ (g) could be represented as: $ { \log }\left\{ {{{p\left( {{\text{TeO}}_{ 2} ,\;{\text{g}}} \right)} \mathord{\left/ {\vphantom {{p\left( {{\text{TeO}}_{ 2} ,\;{\text{g}}} \right)} {{\text{Pa}} \pm 0.0 4}}} \right. \kern-0em} {{\text{Pa}} \pm 0.0 4}}} \right\} = 19. 12- 27132\; \left({\rm{{{\text{K}}}}/T} \right) $ . The $ \Updelta_{\text{f}} G^{^\circ } \;\left( {{ \Pr }_{ 2} {\text{TeO}}_{ 6} } \right) $ could be represented by the relation $ \left\{ {{{\Updelta_{\text{f}} G^{^\circ } \left( {{ \Pr }_{ 2} {\text{TeO}}_{ 6} ,\;{\text{s}}} \right)} \mathord{\left/ {\vphantom {{\Updelta_{\text{f}} G^{^\circ } \left( {{ \Pr }_{ 2} {\text{TeO}}_{ 6} ,\;{\text{s}}} \right)} {\left( {{\text{kJ}}\,{\text{mol}}^{ - 1} } \right)}}} \right. \kern-0em} {\left( {{\text{kJ}}\,{\text{mol}}^{ - 1} } \right)}} \pm 5.0} \right\} = - 2 4 1 5. 1+ 0. 5 7 9 3\;\left(T/{\text{K}}\right) .$ Enthalpy increments of Pr₂TeO₆ were measured by drop calorimetry in the temperature range of 573–1,273 K and heat capacity, entropy and Gibbs energy functions were derived. The $ \Updelta_{\text{f}} H_{{298\;{\text{K}}}}^{^\circ } \;\left( {{ \Pr }_{ 2} {\text{TeO}}_{ 6} } \right) $ was found to be $ {{ - 2, 40 7. 8 \pm 2.0} \mathord{\left/ {\vphantom {{ - 2, 40 7. 8 \pm 2.0} {\left( {{\text{kJ}}\,{\text{mol}}^{ - 1} } \right)}}} \right. \kern-0em} {\left( {{\text{kJ}}\,{\text{mol}}^{ - 1} } \right)}} $ .     

Thermodynamic characterization of chromium tellurate

The standard Gibbs energy of formation of chromium tellurate, Cr₂TeO₆ was determined from the vapour pressure measurement of TeO₂(g) over the phase mixture Cr₂TeO₆(s) + Cr₂O₃(s) in the temperature range 1,183–1,293 K. A thermogravimetry (TG)-based transpiration technique was used for the vapour pressure measurement. This technique was validated by measuring the vapour pressure of CdCl₂(g) over CdCl₂(s). The temperature dependence of the vapour pressure of CdCl₂(g) could be represented as logp (Pa) (±0.02) = 12.06 ? 8616.3/T (K) (734 ? 823 K). A ‘third-law’ analysis of the vapour pressure data yielded a mean value of 185.1 ± 0.4 kJ mol^?1 for the enthalpy of sublimation of CdCl₂(s). The temperature dependence of vapour pressure of TeO₂(g) generated by the incongruent vapourisation reaction, $ {\text{Cr}}_{ 2} {\text{TeO}}_{ 6} (\rm s) \to {\text{Cr}}_{ 2} {\text{O}}_{ 3} (\rm s) + {\text{TeO}}_{ 2} (\rm g) + 1/2\,{\text{O}}_{ 2} (\rm g) $ could be represented as logp (Pa) (±0.04) = 18.57 – 21,199/T (K) (1,183 – 1,293 K). The temperature dependence of the Gibbs energy of formation of Cr₂TeO₆ could be expressed as $ \{ \Updelta G_{\text{f}}^{ \circ } ({\text{Cr}}_{ 2} {\text{TeO}}_{ 6} ,{\text{ s}}){\text{ (kJ}}\,{\text{mol}}^{ - 1} )\pm 4. 0 {\text{\} = }} - 1 6 2 5. 6 { \,+\, 0} . 5 3 3 6\,T({\text{K}}) \, (1{,}183 - 1{,}293\,{\text{K}}). $ A drop calorimeter was used for measuring the enthalpy increments of Cr₂TeO₆ in the temperature range 373–973 K. Thermodynamic functions viz., heat capacity, entropy and Gibbs energy functions of Cr₂TeO₆ were derived from the experimentally measured enthalpy increment values. $ \Updelta H_{{{\text{f}},298\,{\text{K}}}}^{ \circ } ({\text{Cr}}_{ 2} {\text{TeO}}_{ 6} ) $ was found to be ?1636.9 ± 0.8 kJ mol^?1.     

Thermal investigations of SmFeTeO6 and SmCrTeO6 compounds

SmFeTeO₆ and SmCrTeO₆ were synthesized by heating the respective oxides in molar quantities and characterized by X-ray technique. Thermogravimetric studies suggested that SmFeTeO₆ and SmCrTeO₆ vapourize incongruently according to the reactions: $$ \begin{aligned} {\text{SmFeTeO}}_{ 6}{({\text{s}})} & \to {\text{SmFeO}}_{ 3} {( {\text{s}})} + {\text{TeO}}_{ 2} {( {\text{g}})} + \left( { 1/ 2} \right){\text{O}}_{ 2}{( {\text{g}})} \\ {\text{SmCrTeO}}_{ 6} {( {\text{s}})} & \to {\text{SmCrO}}_{ 3} {( {\text{s}})} + {\text{TeO}}_{ 2}{( {\text{g}})} + \left( { 1/ 2} \right){\text{O}}_{ 2}{( {\text{g}})}. \\ \end{aligned} $$ X-ray diffraction data of both the compounds have been indexed on the hexagonal system. Partial pressures of TeO₂(g) were measured over SmFeO₃(s) and SmCrO₃(s) by employing the Knudsen effusion mass loss technique. The standard Gibbs free energy of formation of (Δ_f G°) SmFeTeO₆(s) and SmCrTeO₆(s) were obtained from partial pressures and represented by the following relations: $$\Updelta_{\text{f}} G^{\circ} \left( {{\text{SmFeTeO}}_{6}{( {{\text{s}},\,T})}} \right) \pm 2 5\,{\text{kJ}}\,{\text{mol}}^{ - 1} = - 1 5 1. 6 5+ 0. 1 5\left(T \right)\quad \left( 1 ,0 90{-} 1,1 80\,{\text{K}} \right) \\ \Updelta_{\text{f}} G^{\circ } \left( {{\text{SmCrTeO}}_{ 6} {( {{\text{s}},\,T})}} \right) \pm 2 5\,{\text{kJ}}\,{\text{mole}}^{ - 1} = - 2 5 2. 8 6+ 0. 1 2(T)\quad \left( { 1,100 {-} 1 , 1 7 5\,{\text{K}}} \right).$$      

Spectroscopic Studies on the Thermodynamics of the Complexation of Trivalent Curium with Propionate in the Temperature Range from 20 to 90 °C

The thermodynamics of the stepwise complexation reaction of Cm(III) with propionate was studied by time resolved laser fluorescence spectroscopy (TRLFS) and UV/Vis absorption spectroscopy as a function of the ligand concentration, the ionic strength and temperature (20–90 °C). The molar fractions of the 1:1 and 1:2 complexes were quantified by peak deconvolution of the emission spectra at each temperature, yielding the log₁₀ $ K_{n}^{\prime } $ values. Using the specific ion interaction theory (SIT), the thermodynamic stability constants log₁₀ $ K_{n}^{0} (T) $ were determined. The log₁₀ $ K_{n}^{0} (T) $ values show a distinct increase by 0.15 (n = 1) and 1.0 (n = 2) orders of magnitude in the studied temperature range, respectively. The temperature dependency of the log₁₀ $ K_{n}^{0} (T) $ values is well described by the integrated van’t Hoff equation, assuming a constant enthalpy of reaction and $ \Updelta_{\text{r}} C^\circ_{{p,{\text{m}}}} = 0, $ yielding the thermodynamic standard state $ \left( {\Updelta_{\text{r}} H^\circ_{\text{m}} ,\Updelta_{\text{r}} S^\circ_{\text{m}} ,\Updelta_{\text{r}} G^\circ_{\text{m}} } \right) $ values for the formation of the $ {\text{Cm(Prop)}}_{n}^{3 - n} $ , n = (1, 2) species.     

Synergistic extraction of some univalent cations into nitrobenzene by using sodium dicarbollylcobaltate and nonactin

From extraction experiments and γ-activity measurements, the exchange extraction constants corresponding to the general equilibrium $ {\text{M}}^{ + } \left( {\text{aq}} \right) \, + {\mathbf{1}}\cdot{\text{Na}}^{ + } \left( {\text{nb}} \right) \Leftrightarrow {\mathbf{1}}\cdot{\text{M}}^{ + } \left( {\text{nb}} \right) \, + {\text{Na}}^{ + } \left( {\text{aq}} \right) $ taking place in the two-phase water–nitrobenzene system $ \begin{gathered} ({\text{M}}^{ + } = {\text{ Li}}^{ + } ,{\text{ K}}^{ + } ,{\text{ Rb}}^{ + } ,{\text{ Cs}}^{ + } ,{\text{ H}}_{ 3} {\text{O}}^{ + } ,{\text{NH}}_{4}^{ + }, {\text{ Ag}}^{ + } ,{\text{ Tl}}^{ + } ;{\mathbf{1}} \\ = {\text{ nonactin}};{\text{ aq }} = {\text{ aqueous phase}},{\text{ nb }} = {\text{nitrobenzene phase}}) \\ \end{gathered} $ were determined. Moreover, the stability constants of the 1·M⁺ complexes in water-saturated nitrobenzene were calculated; they were found to increase in the series of $ {\text{Cs}}^{ + } < {\text{ Rb}}^{ + } < {\text{ H}}_{ 3} {\text{O}}^{ + } ,{\text{ Ag}}^{ + } < {\text{ Tl}}^{ + } < {\text{ Li}}^{ + } < {\text{ K}}^{ + } < {\text{NH}}_{4}^{ + } $ .     

Thermochemical properties of the coordination complex of 8-hydroxyquinolinato-bis-(salicylato) praseodymium (III)

The product, [Pr(C₇H₅O₃)₂(C₉H₆NO)], which was formed by praseodymium nitrate hexahydrate, salicylic acid (C₇H₆O₃), and 8-hydroxyquinoline (C₉H₇NO), was synthesized and characterized by elemental analysis, UV spectra, IR spectra, molar conductance, and thermogravimetric analysis. In an optimalizing calorimetric solvent, the dissolution enthalpies of [Pr(NO₃)₃·6H₂O(s)], [2 C₇H₆O₃(s) + C₉H₇NO(s)], [Pr(C₇H₅O₃)₂(C₉H₆NO)(s)], and [solution D (aq)] were measured to be, by means of a solution-reaction isoperibol microcalorimeter, $ \begin{gathered}\Updelta_{\text{s}} H_{\text{m}}^{\theta}\left[ {{ \Pr }\left( {{\text{NO}}_{ 3} } \right)_{ 3} \cdot 6{\text{H}}_{ 2} {\text{O}}\left( {\text{s}} \right), 2 9 8. 1 5{\text{ K}}} \right] \, = - ( 20. 6 6 { } \pm \, 0. 29)\,{\text{kJ}}\,{\text{mol}}^{ - 1} , \\\Updelta_{\text{s}} H_{\text{m}}^{\theta } \left[ { 2 {\text{C}}_{7} {\text{H}}_{ 6} {\text{O}}_{ 3} \left( {\text{s}} \right) +{\text{ C}}_{ 9} {\text{H}}_{ 7} {\text{NO}}\left( {\text{s}}\right),{ 298}. 1 5 {\text{ K}}} \right] \, = \, ( 4 2. 2 7 { }\pm \, 0. 3 1)\,{\text{kJ}}\,{\text{mol}}^{ - 1} , \\\Updelta_{\text{s}} H_{\text{m}}^{\theta } \left[ {{\text{solutionD }}\left( {\text{aq}} \right), 2 9 8. 1 5 {\text{ K}}} \right] \,= - \left( { 8 9. 1 5 { } \pm \, 0. 4 3}\right)\,{\text{kJ}}\,{\text{mol}}^{ - 1} , \\\end{gathered} $ Δ s H m θ [ Pr ( NO 3 ) 3 · 6 H 2 O ( s ) , 2 9 8.1 5 K ] = ? ( 20.6 6 ± 0.2 9 ) kJ mol ? 1 , Δ s H m θ [ 2 C 7 H 6 O 3 ( s ) + C 9 H 7 NO ( s ) , 298.1 5 K ] = ( 4 2.2 7 ± 0.3 1 ) kJ mol ? 1 , Δ s H m θ [ solution D ( aq ) , 2 9 8.1 5 K ] = ? ( 8 9.1 5 ± 0.4 3 ) kJ mol ? 1 , and $ \Updelta_{\text{s}} H_{\text{m}}^{\theta } \left\{ {\left[ {{\Pr }\left( {{\text{C}}_{ 7} {\text{H}}_{ 5} {\text{O}}_{ 3} }\right)_{ 2} \left( {{\text{C}}_{ 9} {\text{H}}_{ 6} {\text{NO}}}\right)} \right]\left( {\text{s}} \right),{ 298}. 1 5 {\text{ K}}}\right\} \, = - \left( { 4 1.0 4 { } \pm \, 0. 3 3}\right)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ Δ s H m θ { [ Pr ( C 7 H 5 O 3 ) 2 ( C 9 H 6 NO ) ] ( s ) , 298.1 5 K } = ? ( 4 1.0 4 ± 0.3 3 ) kJ mol ? 1 , respectively. Through an improved thermochemical cycle, the enthalpy change of the designed coordination reaction was calculated to be $\Updelta_{\text{r}} H_{\text{m}}^{\theta} = \, ( 2 1 3. 1 8\pm0. 6 9)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ Δ r H m θ = ( 2 1 3.1 8 ± 0.6 9 ) kJ mol ? 1 , the standard molar enthalpy of the formation was determined as $ \Updelta_{\text{f}} H_{\text{m}}^{\theta} \left\{ {\left[ {{\Pr }\left( {{\text{C}}_{ 7} {\text{H}}_{ 5} {\text{O}}_{ 3} }\right)_{ 2} \left( {{\text{C}}_{ 9} {\text{H}}_{ 6} {\text{NO}}}\right)} \right]\left( {\text{s}} \right), 2 9 8. 1 5 {\text{K}}}\right\} \, = \, - \, ( 1 8 7 5. 4\pm 3.1)\,{\text{kJ}}\,{\text{mol}}^{ - 1} $ Δ f H m θ { [ Pr ( C 7 H 5 O 3 ) 2 ( C 9 H 6 NO ) ] ( s ) , 2 9 8.1 5 K } = ? ( 1 8 7 5.4 ± 3.1 ) kJ mol ? 1 .     

Crystal structure and thermochemical properties of bis(1-octylammonium) tetrachlorocuprate

Bis(1-octylammonium) tetrachlorocuprate (1-C₈H₁₇NH₃)₂CuCl₄(s) was synthesized by the method of liquid phase reaction. The crystal structure of the compound has been determined by X-ray crystallography. The lattice potential energy was obtained from the crystallographic data. Molar enthalpies of dissolution of (1-C₈H₁₇NH₃)₂CuCl₄(s) at various molalities were measured at 298.15?K in the double-distilled water by means of an isoperibol solution-reaction calorimeter, respectively. In terms of Pitzer??s electrolyte solution theory, the molar enthalpy of dissolution of (1-C₈H₁₇NH₃)₂CuCl₄(s) at infinite dilution was determined to be $ \Updelta_{\rm s} H_{\text{m}}^{\infty } = \, - 5. 9 7 2\,{\text{kJ}}\,{\text{mol}}^{ - 1} , $ and the sums of Pitzer??s parameters $ (4\beta_{{{\text{C}}_{ 8} {\text{H}}_{ 1 7} {\text{NH}}_{ 3} , {\text{Cl}}}}^{ ( 0 )L} + 2\beta_{\text{Cu,Cl}}^{ ( 0 )L} + \theta_{{{\text{C}}_{ 8} {\text{H}}_{ 1 7} {\text{NH}}_{ 3} , {\text{Cu}}}}^{L} ) $ and $ (2\beta_{{{\text{C}}_{ 8} {\text{H}}_{ 1 7} {\text{NH}}_{ 3} , {\text{Cl}}}}^{ ( 1 )L} + \beta_{\text{Cu,Cl}}^{ ( 1 )L} ) $ were obtained.     

Enthalpy increments and thermodynamic functions of rare earth samarium titanates RESmTi<Subscript>2</Subscript>O<Subscript>7</Subscript>(s) (RE = Gd,Dy)

Enthalpy measurements have been taken on GdSmTi₂O₇ and DySmTi₂O₇ by using a high-temperature differential calorimeter at temperature between 800 and 1655 K. Thermodynamic function, such as heat capacity, entropy and Gibbs energy functions of GdSmTi₂O₇ and DySmTi₂O₇, was derived using the data obtained in this study. The results are presented and compared with the data available in the literature. The polynomial expression of enthalpy increments obtained for GdSmTi₂O₇(s) and DySmTi₂O₇(s) in the temperature range 298–1700 K is given as: \(\begin{aligned} H_{\text{T}}^{0} - H_{298}^{0} / {\text{J}}\,{\text{mol}}^{ - 1} & = 252.961\,T \, + 1.596 \times 10^{ - 2} \,T^{2} + 3.705 \times 10^{6} \,T^{ - 1} - 89{,}265\quad ({\text{GdSmTi}}_{2} {\text{O}}_{7} ) \\ H_{\text{T}}^{0} - H_{298}^{0} / {\text{J}}\,{\text{mol}}^{ - 1} & = 256.504\,T \, + 1.576 \times 10^{ - 2} \,T^{2} + 3.531 \times 10^{6} \,T^{ - 1} - 89{,}721\quad \left( {{\text{DySmTi}}_{2} {\text{O}}_{7} } \right). \\ \end{aligned}\)     

Imidazolium-Based Ionic Liquids Containing the Trifluoroacetate Anion: Thermodynamic Study

New experimental vapor pressures and vaporization enthalpies of the ionic liquids \( [ {\text{C}}_{2} {\text{mim][CF}}_{3} {\text{CO}}_{2} ] \) and \( [ {\text{C}}_{4} {\text{mim][CF}}_{3} {\text{CO}}_{2} ] \) have been measured by the QCM method. The solution enthalpies of these ionic liquids were measured by using high-precision solution calorimetry and were used for calculation the aqueous enthalpy of formation \( \Delta_{\text{f}} H_{\text{m}}^{ \circ } ({\text{CF}}_{ 3} {\text{CO}}_{2}^{ - } ,_{{}} {\text{aq}}) \) of the anion for combination with quantum-chemical results. The solubility parameters of the ILs under study have been derived from experimental \( \Delta_{\text{l}}^{\text{g}} H_{\text{m}}^{ \circ } \)(298.15 K) values and were used for estimation of miscibility of some common solutes with \( [ {\text{C}}_{n} {\text{mim][CF}}_{3} {\text{CO}}_{2} ] \).     

Specific heat capacity and thermodynamic properties of CuTeO3 and HgTeO3

Tellurites of CuTeO₃ and HgTeO₃ are synthesized and their specific molar heat capacities are experimentally determined for the first time. The tellurites discussed in the present paper are used for preparation of optical glasses with special properties for optoelectronics, nuclear and power industries. The tellurites synthesized are prepared for chemical analysis, differential thermal analysis and X-ray analysis. The use of the tellurites studied is related to knowing their thermodynamic properties like specific molar heat capacity (C _p,m), enthalpy \( \left( {\Delta_{{{\text {T}}^{\prime}}}^{\text{T}} H_{\text{m}}^{0} } \right), \) entropy \( \left( {\Delta_{{{\text {T}}^{\prime}}}^{\text{T}} S_{\text{m}}^{0} } \right) \) and Gibbs energy \( \left( { - \Delta_{{{\text {T}}^{\prime}}}^{\text{T}} G_{\text{m}}^{0} } \right) \) . The temperature dependences of their molar heat capacities are determined using the least squares method. The thermodynamic properties are calculated: entropy, enthalpy and Gibbs function.     

A Potentiometric Study of the Association of Copper(II) and Sulfate Ions in Aqueous Solution at 25 °C

A study of the association between copper(II) and sulfate ions in aqueous solution has been made using copper ion-selective electrode potentiometry at constant ionic strengths (I) of 0.05, 0.1, 0.25, 0.5, 1.0, 3.0 and 5.0 mol·L^?1 in NaClO₄ media at 25 °C. Only one complex was detected, corresponding to the equilibrium: \( {\text{Cu}}^{ 2+ } ({\text{aq}}) + {\text{SO}}_{4}^{2 - } ({\text{aq}}) \rightleftarrows {\text{CuSO}}_{4}^{0} ({\text{aq}}). \) No higher order complexes were detected even at sulfate/copper(II) concentration ratios of up to 1,000. The present potentiometric values of log₁₀ K ₁(I) are shown to be consistently higher than those obtained by UV–Vis spectrophotometry because of the failure of the latter technique to detect all of the solvent-separated ion pairs present. Extrapolation of log₁₀ K ₁(I) to infinite dilution using an extended Guggenheim equation yielded a standard state value of log₁₀ \( K_{1} \{ {\text{CuSO}}_{4}^{0} ({\text{aq}})\} = 2.32 \pm 0.09 \) , which is in excellent agreement with a recent IUPAC-recommended value.     

Isosaccharinate Complexes of Fe(III)

Prior to this study there were no thermodynamic data for isosaccharinate (ISA) complexes of Fe(III) in the environmental range of pH (>~4.5). This study was undertaken to obtain such data in order to predict Fe(III) behavior in the presence of ISA. The solubility of Fe(OH)₃(2-line ferrihydrite), referred to as Fe(OH)₃(s), was studied at 22?±?2?°C in: (1) very acidic (0.01?mol·dm^?3 H⁺) to highly alkaline conditions (3?mol·dm^?3 NaOH) as a function of time (11?C421?days), and fixed concentrations of 0.01 or 0.001?mol·dm^?3 NaISA; and (2) as a function of NaISA concentrations ranging from approximately 0.0001 to 0.256?mol·dm^?3 and at fixed pH values of approximately 4.5 and 11.6 to determine the ISA complexes of Fe(III). The data were interpreted using the SIT model that included previously reported stability constants for $ {{\text{Fe(ISA}})_{n}}^{3 - n} $ (with n varying from 1 to 4) and Fe(III)?COH complexes, and the solubility product for Fe(OH)₃(s) along with the values for two additional complexes (Fe(OH)₂(ISA)(aq) and $ {\text{Fe(OH)}}_{ 3} ( {{\text{ISA}})_{2}}^{2 - } $ ) determined in this study. These extensive data provided a log₁₀ K ⁰ value of 1.55?±?0.38 for the reaction $ ({\text{Fe}}^{ 3+ } + {\text{ISA}}^{-} + 2 {\text{H}}_{ 2} {\text{O}} \rightleftarrows {\text{Fe(OH}})_{ 2} {\text{ISA(aq}}) + 2 {\text{H}}^{ + } ) $ and a value of ?3.27?±?0.32 for the reaction $ ({\text{Fe}}^{ 3+ } + 2 {\text{ISA}}^{-} + 3 {\text{H}}_{ 2} {\text{O}} \rightleftarrows {\text{Fe(OH)}}_{ 3} ( {\text{ISA}})_{2}^{2 - } + 3 {\text{H}}^{ + } ) $ and show that ISA forms strong complexes with Fe(III) which significantly increase the Fe(OH)₃(s) solubility at pH?<~12. Thermodynamic calculations show that competition of Fe(III) with tetravalent ions for ISA does not significantly affect the solubilities of tetravalent hydrous oxides (e.g., Th and Np(IV)) in ISA solutions.     

Solubility constants and free enthalpies of metal sulphides,part 6: A new solubility cell

A solubility cell which can be operated continuously over the temperature range 5–95 °C has been developed. The solubility of Fe_0.88S (monoclinic pyrrhotite) in solutions $$S_0 = ([H^ + ]) = H{\text{ }}m,{\text{ }}[Na^ + ] = (1.00---H) m,{\text{ }}[ClO_{4^ - } ] = 1.00 m)$$ at fixed partial pressures of H₂S has been investigated at 50.7 °C. The hydrogen ion concentration and the total concentration of iron(II) ion in equilibrium with the solid phase was determined by e.m.f. and analytical methods respectively. The data were consistent with $$\log ^* K_{pso} = \log \frac{{[Fe^{2 + } ]pH_2 S}}{{[H^ + ]^2 }} = 3.80 \pm {\text{ }}0.10{\text{ }}[50.7^\circ C,{\text{ }}1 m(Na)ClO_4 ]$$ according to the overall reaction $$1.14{\text{ }}Fe_{0.88} S_{(s)} {\text{ }} + {\text{ }}2H_{(I = 1m)}^ + {\text{ }} \rightleftharpoons {\text{ }}Fe_{(I = 1m)}^{2 + } {\text{ }} + {\text{ H}}_{\text{2}} S_{(g)} {\text{ }} + {\text{ }}0.14{\text{ }}S_{(s)} $$      

Extraction kinetics of Uranium(VI) and Thorium(IV) with Tri-iso-amyl phosphate from nitric acid using a Lewis Cell

The extraction kinetics of uranium(VI) and thorium(IV) with Tri-iso-amyl phosphate (TiAP) from nitric acid medium has been investigated using a Lewis Cell. Especially, dependences of the extraction rate on stirring speed, temperature, interfacial area were firstly measured to elucidate the extraction kinetics regimes. The experimental results demonstrated that extraction kinetic of U(VI) is governed by chemical reactions at interface with an activation energy, E_a, of 43.41 kJ/mol, while the rate of Th(IV) extraction is proved to be intermediate controlled, of which the E_a is 23.20 kJ/mol. Reaction orders with respect to the influencing parameters of the extraction rate are determined, and the rate equations of U(VI) and Th(IV) at 293 K have been proposed as $$ {\text{r}} = - {\text{dcUO}}_{ 2} \left( {{\text{NO}}_{ 3} } \right)_{ 2} /{\text{dt}} = 1. 80 \times 10^{ - 3} \left[ {{\text{UO}}_{ 2} \left( {{\text{NO}}_{ 3} } \right)_{ 2} } \right]^{ 1.0 1} \left[ {\text{TiAP}} \right]^{0. 5 5} , $$ $$ {\text{r}} = - {\text{dcTh }}\left( {{\text{NO}}_{ 3} } \right)_{ 4} /{\text{dt}} = 1. 8 8\times 10^{ - 3} \left[ {{\text{Th }}\left( {{\text{NO}}_{ 3} } \right)_{ 4} } \right]^{ 1.0 4} \left[ {\text{TiAP}} \right]^{ 1. 7 7} \left[ {{\text{HNO}}_{ 3} } \right]^{0. 3 8} , $$ respectively.     

The dissolution properties of N-guanylurea-dinitramide (FOX-12) in dimethyl sulfoxide (DMSO)

The enthalpy of dissolution of FOX-12 in dimethyl sulfoxide (DMSO) was measured by means of a RD496-III Calvet microcalorimeter at 298.15 K. Empirical formulae for the calculation of the enthalpy of dissolution ( $ \Updelta_{\text{diss}} H $ ), relative partial molar enthalpy ( $ \Updelta_{\text{diss}} H_{\text{partial}} $ ), and relative apparent molar enthalpy ( $ \Updelta_{\text{diss}} H_{\text{apparent}} $ ) were obtained from the experimental data of the enthalpies of dissolution of FOX-12 in DMSO. The kinetic equation that describes the dissolution process of FOX-12 in DMSO at 298.15 K is determined as $ \frac{{{\text{d}}\alpha }}{{{\text{d}}t}} = 8.5 \times 10^{ - 3} (1 - \alpha )^{0.59} $ .     

Synergistic extraction of europium and americium into nitrobenzene by using hydrogen dicarbollylcobaltate and 1,3-bis(diphenylphosphino)propane dioxide

Extraction of microamounts of europium and americium by a nitrobenzene solution of hydrogen dicarbollylcobaltate (H⁺B^?) in the presence of 1,3-bis(diphenylphosphino)propane dioxide (DPPPrDO, L) has been investigated. The equilibrium data have been explained assuming that the species $ {\text{HL}}^{ + } ,\,{\text{HL}}_{2}^{ + } ,\,{\text{ML}}^{3 + } \,{\text{and}}\,{\text{ML}}_{3}^{3 + } $ HL + , HL 2 + , ML 3 + and ML 3 3 + (M³⁺ = Eu³⁺, Am³⁺) are extracted into the organic phase. The values of extraction and stability constants of the complex species in nitrobenzene saturated with water have been determined. It was found that the stability constants of the corresponding complexes $ {\text{EuL}}_{n}^{ 3+ } \,{\text{and}}\,{\text{AmL}}_{n}^{ 3+ } , $ EuL n 3 + and AmL n 3 + , where n = 1, 3 and L is DPPPrDO, in water-saturated nitrobenzene are comparable.     

Experimental and DFT study on complexation of Eu3+ with a macrocyclic lactam receptor

From extraction experiments and $ \gamma $ -activity measurements, the extraction constant corresponding to the equilibrium $ {\text{Eu}}^{ 3+ } \left( {\text{aq}} \right) + 3 {\text{A}}^{ - } \left( {\text{aq}} \right) + {\mathbf{1}}\left( {\text{nb}} \right) \Leftrightarrow {\mathbf{1}} \cdot {\text{Eu}}^{ 3+ } \left( {\text{nb}} \right) + 3 {\text{A}}^{ - } \left( {\text{nb}} \right) $ taking place in the two-phase water–nitrobenzene system ( $ {\text{A}}^{ - } = \text {CF}_{3} \text{SO}_{3}^{ - } $ ; 1 = macrocyclic lactam receptor—see Scheme 1; aq = aqueous phase, nb = nitrobenzene phase) was evaluated as $ { \log } K_{{{\text{ex}} }} ({\mathbf{1}} \cdot {\text{Eu}}^{ 3+ } ,{\text{ 3A}}^{ - } )\; = \; - 4. 9 \pm 0. 1 $ . Further, the stability constant of the 1·Eu³⁺ cationic complex in nitrobenzene saturated with water was calculated for a temperature of 25 °C: $ { \log } \beta_{{{\text{nb}} }} ({\mathbf{1}} \cdot {\text{Eu}}^{ 3+ } ) \; = \; 8. 2 \pm 0. 1 $ . Finally, using DFT calculations, the most probable structure of the cationic complex species 1·Eu³⁺ was derived. In the resulting 1·Eu³⁺ complex, the “central” cation Eu³⁺ is bound by five bond interactions to two ethereal oxygen atoms and two carbonyl oxygens, as well as to one carbon atom of the corresponding benzene ring of the parent macrocyclic lactam receptor 1 via cation-π interaction.

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!     

Standard molar formation enthalpies of flower-like NH4CdPO4·H2O

Flower-like ammonium cadmium phosphate monohydrate was synthesized by solid-state reaction at low temperature and characterized by X-ray diffraction, Fourier transform infrared spectroscopy, scanning electron microscope, and elemental analysis. The product NH₄CdPO₄·H₂O was obtained with flower-like morphology by the addition of fatty alcohol-polyoxyethylene ether surface-active agent. Based on Hess??s law, thermochemical cycle was designed to determine the dissolution enthalpies of reactants and products using a solution-reaction isoperibol calorimeter at 298.15?K, and the molar reaction enthalpy was calculated on the basis of above dissolution enthalpies. With the aid of other auxiliary thermodynamic data, the standard molar formation enthalpy of the title compounds was concluded: $ \Updelta_{\text{f}} H_{\text{m}}^{-\!\!\!\!\circ} [{\text{NH}}_{4} {\text{CdPO}}_{4} \cdot {\text{H}}_{2} {\text{O}}] = ( - 1749.82 \pm 0.76)\; {\text{kJ}}\; {\text{mol}}^{ - 1} . $