Molality Scale Research Articles

Measurement and Correlation of the Solubility of Ammonia in Ethanol between 277.35 and 318.25 K

The solubility of ammonia in ethanol was measured by a static−analytic method. The total pressure is up to about 0.4 MPa, and temperature ranges from (277.35 to 318.25) K. The relationship between the partial molar volume of ammonia infinitely diluted in ethanol and temperature is obtained by measuring the density of (ethanol + ammonia) solution. The experimental results are used to determine the (molality scale based) Henry’s constant of ammonia in ethanol, where the activity in the liquid phase is calculated by applying the molality scale based Gibbs excess energy model of Pitzer.

Solubility, activity coefficients and acid–base properties of three naphthol derivatives in NaCl (aq) at different ionic strengths and at T = 298.15 K

Total and intrinsic solubilities of three naphthol derivatives (1-Naphthol, 1,5-Dihydroxynaphthalene and 1-Amino-2-Naphthol-4-Sulfonic acid) were determined in NaCl (aq) at different salt concentrations (0–3 mol L −1) and in pure water at T = 298.15 K. To characterize the acid–base properties of these naphthols, the protonation constants were determined potentiometrically in the same conditions. The solubility of neutral species in pure water is −1.978, −2.886, −2.594 (expressed in logarithmic scale) for 1-Naphthol, 1,5-Dihydroxynaphthalene and 1-Amino-2-Naphthol-4-Sulfonic acid, respectively. The values of Setschenow coefficients ( k m ) for these three ligands, are 0.209, 0.158 and 0.369. The dependence on the ionic strength of protonation constants K i H from potentiometric measurements was studied using three models, namely Debye–Hückel, SIT (Specific ion Ionic Interaction) and Pitzer equations. Therefore infinite dilution K i H and interaction coefficients (SIT and Pitzer) were calculated. The activity coefficients were obtained from solubility and protonation data, in the molal concentration scale. Activity coefficients in NaCl (aq) shown the same behaviour found for other O-donors ligands, previously studied.

Activity Coefficients of NaCl in Aqueous Mixtures with High Relative Permittivity Cosolvent: Ethylene Carbonate + Water at 298.15 K

The mean ionic activity coefficients of NaCl were experimentally determined in ethylene carbonate (EC) + water at 298.15 K from potential difference measurements of the following electrochemical cell containing two ion-selective electrodes (ISEs): Na-ISE|NaCl(m),EC(w),H2O(1−w)|Cl-ISE The molality (m) varied between 0.02 mol·kg−1 and almost saturation, while the mass fraction of EC in the mixture (w) was varied between 0 and 0.8 in 0.1 unit steps. Values of the standard potential difference, E0 (molal scale), were determined using a routine method of extrapolation together with the extended Debye−Hückel, Pitzer, and Scatchard equations. The results obtained produced good internal consistency, within the normal limits of experimental error encountered in these types of measurement. Once E0 was determined, the mean ionic activity coefficients of the NaCl, the Gibbs energy of transfer of the NaCl from the water to the EC + water mixture, standard solubility product of NaCl in EC + water, and the NaCl primary hydration number were estimated. The results were comparatively analyzed with those of the NaF previously obtained in similar mixtures.

Raman spectroscopic study for the determination of Cl − concentration (molarity scale) in aqueous solutions: Application to fluid inclusions

From recent studies on the structure of aqueous alkali and alkaline chloride solutions, it can be deduced that the effects of monovalent and divalent cations are small, and water–chloride interactions are primarily responsible for the solutions. Therefore, Cl − molarity can be treated as an important parameter in describing these systems. In this work, the NaCl–H 2O, KCl–H 2O, CaCl 2–H 2O, MgCl 2–H 2O and NaCl–KCl–CaCl 2–H 2O systems are synthesized, and their Raman spectra are recorded. From the Raman OH stretching bands of the solutions, the isosbestic points can be discriminated. This means that there exists the equilibrium between various structural species, the Raman OH vibration can be applied to study the salinity of these solutions. In this study, the Raman OH vibration is fitted with two Gaussian sub-bands, the equilibrium constant can be determined, then the calibration equation of Cl − concentration (molarity scale) can be determined. After the effects of the birefringence of the host mineral on the Raman OH stretching bands are discussed, the method is applied to measure the salinity of natural fluid inclusions.

Solubility of B-Nb2O5 and the Hydrolysis of Niobium(V) in Aqueous Solution as a Function of Temperature and Ionic Strength

B-Nb2O5 was recrystallized from commercially available oxide, and XRD analyses indicated that it is stable in contact with solutions over the pH range 0 to 9, whereas solid polyniobates such as Na8Nb6O19⋅13H2O(s) appear to predominate at pH>9. Solubilities of the crystalline B-Nb2O5 were determined in five NaClO4 solutions (0.1≤Im/mol⋅kg−1≤1.0) over a wide pH range at (25.0±0.1) °C and at 0.1 MPa. A limited number of measurements were also made at Im=6.0 mol⋅kg−1, whereas at Im=1.0 mol⋅kg−1 the full range of pH was also covered at (10, 50 and 70) °C. The pH of these solutions was fixed using either HClO4 (pH≤4) or NaOH (pH≥10) and determined by mass balance, whereas the pH on the molality scale was measured in buffer mixtures of acetic acid + acetate (4≤pH≤6), Bis-Tris (pH≈7), Tris (pH≈8) and boric acid + borate (pH≈9). Treatment of the solubility results indicated the presence of four species, \(\mathrm{Nb(OH)}_{n}^{5-n}\) (where n=4–7), so that the molal solubility quotients were determined according to: $$0\mathrm{.5Nb}_{2}\mathrm{O}_{5}\mathrm{(cr)+0}\mathrm{.5(2}n-5\mathrm{)H}_{2}\mathrm{O(l)}_{\leftarrow}^{\to}\mathrm{Nb(OH)}_{n}^{5-n}+(n-5)\mathrm{H}^{+}\quad (n=4\mbox{--}7)$$ and were fitted empirically as a function of ionic strength and temperature, including the appropriate Debye-Huckel term. A Specific Interaction Theory (SIT) approach was also attempted. The former approach yielded the following values of log 10Ksn (infinite dilution) at 25 °C: −(7.4±0.2) for n=4; −(9.1±0.1) for n=5; −(14.1±0.3) for n=6; and −(23.9±0.6) for n=7. Given the experimental uncertainties (2σ), it is interesting to note that the effect of ionic strength only exceeded the combined uncertainties significantly in the case of log 10Ks6 to Im=1.0 mol⋅kg−1, such that these values may be of use by defining their magnitudes in other media. Values of ΔfGo, ΔfHo, So and \(C_{p}^{\mathrm{o}}\) (298.15 K, 0.1 MPa) for each hydrolysis product were calculated and tabulated.

10.1007/s11175-008-1009-6

Osmotic coefficients were measured for tetra-n-butylammonium hexafluorophosphate (TBA PF6) and tetra-n-butylammonium tetrakis(pentafluorophenyl)borate (TBABARF) in tetrahydrofuran and in acetonitrile. These osmotic coefficients were fit using equations first derived by Pitzer. Mean ionic activity coefficients were determined using the integrated form of the Pitzer equations. Densities were measured for each of the solutions and the results were used to convert mean ionic activity coefficients on the molal scale to the molar scale. Remarkably, over concentrations ranging from 0.1 to 1.0 m in acetonitrile, the behavior of TBABARF is consistent with the simple Debye-Hückel equation.

Re-evaluation of the First and Second Stoichiometric Dissociation Constants of Oxalic Acid at Temperatures from 0 to 60 °C in Aqueous Oxalate Buffer Solutions with or without Sodium or Potassium Chloride

Equations were developed for the calculation of the first stoichiometric (molality scale) dissociation constant (Km1) of oxalic acid in buffer solutions containing oxalic acid, potassium hydrogen oxalate, and potassium chloride from the determined thermodynamic values of this dissociation constant (Ka1) and the molalities of the components in the solutions. Similar equations were also developed for the second stoichiometric dissociation constant (Km2) of this acid in buffer solutions containing sodium or potassium hydrogen oxalate, oxalate and chloride. These equations apply at temperatures from 0 to 60 °C up to ionic strengths of 1.0 mol⋅kg−1 and they have been based on single-ion activity coefficient equations of the Huckel type. For the equations for Km1, the activity parameters of oxalate species and the Ka1 values were determined at various temperatures from the Harned cell data of a recent tetroxalate buffer paper (Juusola et al., J. Chem. Eng. Data 52:973–976, 2007). By using the resulting equations for Km1, the activity parameters of oxalate species for Km2 and the Ka2 values were then determined from the new Harned cell data and from those of Pinching and Bates (J. Res. Natl. Bur. Stand. (U.S.) 40:405–416, 1948) for solutions of sodium or potassium oxalates with NaCl or KCl. The resulting simple equations for calculation of Km1 and Km2 for oxalic acid were tested with all important thermodynamic data available in the literature for this purpose. The equations for ln (Ka1) and ln (Ka2) are of the form ln (Ka)=a+b(t/°C)+c(t/°C)2. The coefficients for ln (Ka1) are the following: a=−2.8737, b=0.000159, and c=−0.00009. The corresponding coefficients for ln (Ka2) are −9.6563, −0.003059, and −0.000125, respectively. The new activity coefficient equations were used to evaluate the pH values of the tetroxalate buffer solution (i.e., of the 0.05 mol⋅kg−1 KH3C4O8 solution) for comparison with the pH values recommended by IUPAC at temperatures from 0 to 60 °C and to develop a new two-component oxalate pH buffer of 0.01 mol⋅kg−1 KHC2O4+0.05 mol⋅kg−1 Na2C2O4 for which pH values are given from 0 to 60 °C. Values of p(mH) calculated from these equations are tabulated for these buffers as well as for buffer solutions with KCl and KH3C4O8 as the major component and minor component, respectively. Tables of p(mH) are also presented for 0.001 mol⋅kg−1 KHC2O4+0.005 mol⋅kg−1 Na2C2O4 solutions in which KCl is the supporting electrolyte.

Solubility of the Single Gases Carbon Dioxide and Hydrogen in the Ionic Liquid [bmpy][Tf2N

Experimental data for the solubility of the two single gases carbon dioxide and hydrogen in the ionic liquid 1-n-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide ([bmpy][Tf 2 N]) are reported for temperatures between (293.1 and 413.2) K. The maximum pressure (the maximum gas molality in the solvent) was 10.8 MPa (4.43 mol·kg -1 ) for carbon dioxide and 8.9 MPa (0.14 mol·kg -1 ) for hydrogen. The experiments were performed using a high-pressure view-cell technique operating on the synthetic method. The solubility of carbon dioxide in [bmpy][Tf 2 N] is more than 1 order of magnitude higher than the solubility of hydrogen. Carbon dioxide becomes less soluble in [bmpy][Tf 2 N] with increasing temperature, whereas hydrogen shows the opposite effect. An extension of Henry's law is employed to correlate the solubility pressures. The final results for the Henry's constant (at zero pressure) of carbon dioxide and hydrogen in [bmpy][Tf 2 N] (on the molality scale) are represented within the experimental uncertainty (about 0.8 % for carbon dioxide and 1.3 % for hydrogen) by In(k (0) H,CO2 /MPa) = 8.5128 - 2057/(T/K) - 0.004274·(T/K) and In(k (0) H, H2 /MPa) = 3.5746 + 343/(T/K) - 0.0008·(T/K), respectively.

Solubility of the Single Gases Carbon Monoxide and Oxygen in the Ionic Liquid [hmim][Tf2N

Experimental data for the solubility of the two single gases carbon monoxide and oxygen in the ionic liquid 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ([hmim][Tf2N]) are reported for temperatures between (293.25 and 413.2) K. The maximum pressure (the maximum gas molality) was 9.8 MPa (0.27 mol·kg−1) for carbon monoxide and 9.1 MPa (0.31 mol·kg−1) for oxygen. The experiments were performed using a high-pressure view-cell technique operating on the synthetic method. Oxygen shows a slightly higher solubility than carbon monoxide under all conditions investigated, but the solubility generally remains very low. Both gases become less soluble in [hmim][Tf2N] with increasing temperature; however, the effect is not very pronounced and becomes ambiguous for carbon monoxide above 373 K. An extension of Henry’s law is employed in correlating the solubility pressures. The final results for the Henry’s constant (at zero pressure) of carbon monoxide and oxygen in [hmim][Tf2N] (on the molality scale)...

Solubility of Hydrogen in the Cyclic Alkylene Ester 1,2-Butylene Carbonate

The solubility of hydrogen in the cyclic alkylene ester 1,2-butylene carbonate was determined at (283.3, 298.2, and 323.1) K up to a pressure of 9.3 MPa with a high-pressure view-cell technique based on the synthetic method. The solubility of hydrogen in 1,2-butylene carbonate increases with increasing temperature. An extension of Henry’s law was employed to correlate the solubility pressures, and the final results for Henry’s constant (at zero pressure) of hydrogen in 1,2-butylene carbonate (on the molality scale) are correlated within the experimental uncertainty (about ± 1.4 %) by ln(kH,H2(0)/MPa) = 2.123 + 650.96/(T/K).

Activity coefficients of NaCl in aqueous mixtures with ɛ-increasing co-solvent: Formamide–water mixtures at 298.15 K

The electromotive force of the cell containing two ion-selective electrodes (ISE), Na-ISE|NaCl( m), formamide ( Y), H 2O(100 − Y)|Cl-ISE has been measured at a temperature of 298.15 K as a function of the weight percentage Y of formamide in a mixed solvent. Y was varied between 0 and 100% in ten-unit steps and the molality of the electrolyte ( m) was between ca. 0.01 and saturation. The values of the standard electromotive force, E 0 (molal scale), were determined using routine methods of extrapolation together with Debye-Hückel, Pitzer and Scatchard equations. The results obtained produced good internal consistency, within the normal limits of experimental error encountered in these types of measurement. Once E 0 was determined, the mean ionic activity coefficients for NaCl, the Gibbs energy of transfer from the water to the formamide–water mixture, the standard solubility product and the primary NaCl hydration number were calculated. The variation of these magnitudes with the composition of this mixture with ɛ-increasing co-solvent is discussed in comparison with those previously obtained with ɛ-decreasing methanol–water and ethanol–water systems in terms of the ion–solvent and ion–ion interactions and their changes with the properties of the medium.

Solubility and activity coefficients of 2,2′-bipyridyl, 1,10-phenanthroline and 2,2′,6′,2″-terpyridine in NaCl (aq) at different ionic strengths and T = 298.15 K

The solubility of 2,2′-bipyridyl (Bipy), 1,10-phenanthroline (Phen) and 2,2′,6′,2″-terpyridine (Terpy) was determined in NaCl (aq) at different salt concentrations ( c NaCl = 0–5 mol L −1) at 298.15 K. Some solubility data were also obtained for 1,10-phenanthroline in (CH 3) 4NCl (aq) ( c ( C H 3 ) 4 NCl = 0 – 3 mol L − 1 ). To better characterize the acid–base properties of these azines, the protonation constants were determined potentiometrically in the same supporting electrolyte concentrations. The solubility of neutral species is 43.1, 18.8 and 5.4 mmol L −1 for Bipy, Phen and Terpy, respectively, in pure water. In NaCl (aq) solubility decreases by ∼1.2 orders of magnitude. The opposite trend is shown by the solubility of Phen in (CH 3) 4NCl (aq), with a small increasing effect. By the dependence on salt concentration Setschenow coefficients were also determined. All the considered azines (Bipy, Phen and Terpy) form a monoprotonated species HA + and, in addition, Phen forms the diligand species HA 2 +, and Terpy the diprotonated species H 2A 2+. The dependence on ionic strength of protonation constants was modelled by the SIT (Specific ion Interaction Theory) and Pitzer approaches. The activity coefficients of different species were obtained by using solubility and protonation data, in the molal concentration scale. Activity coefficients, for the neutral species, in (CH 3) 4NCl (aq) (only for Phen) show an opposite trend with respect to that in NaCl (aq), and this was interpreted in terms of weak complex formation.

Determination of the E-pO2− stability diagram of plutonium in the molten LiCl–KCl eutectic at 450 °C

Plutonium trichloride solution in the molten LiCl–KCl eutectic was prepared by carbochlorination of plutonium oxide. Kinetics of this reaction was compared in different conditions in the range of 443–550°C. Using this molten salt solution, the redox potential of the Pu(III)/Pu couple at inert tungsten electrode was measured at 450°C by electromotive force measurement and was found to be E′○=−2.76V vs. the Cl2(g)(1atm)/Cl− reference electrode (molar fraction scale). Reaction between plutonium trichloride and oxide ions was studied by potentiometric titration, using yttria stabilized electrodes. In our experimental conditions, the titration curves indicate the precipitation of the sesquioxide Pu2O3. The solubility product cologarithm calculated from these curves is found to be pKs(Pu2O3)=22.8±1.1 (molality scale). Using the experimentally obtained values for E′○, activity coefficient and pKs joined to the published thermodynamic data, the stability phase diagram of the Pu–O species was then drawn.

Solubility of Tetrafluoromethane in the Ionic Liquid [hmim][Tf2N

Experimental results for the solubility of tetrafluoromethane (CF4, R14) in the ionic liquid 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ([hmim][Tf2N]) are presented for temperatures between 293.3 and 413.3 K, at pressures (gas molalities) up to 9.6 MPa (0.22 mol kg-1). The experimental results were determined with a high-pressure view-cell technique operating on the synthetic method. The experimental data were used to determine Henry's constant of tetrafluoromethane in [hmim][Tf2N]. The results for the Henry's constant (at zero pressure) are correlated (on the molality scale) within the experimental uncertainty (i.e., about 1.1%) by ln(k(0)(H,CF4)/MPa) = 7.537 - 893.8/(T/K) - 0.003977(T/K). Henry's law was also extended to describe the gas solubility at higher pressures. Furthermore, a cubic equation of state was used to correlate the gas solubility over the entire range of experimentally investigated temperature and pressure. Both methods proved suited for a reliable correlation of the new experimental data.

Assessable Consequences of Through‐Bond Donor–Acceptor Interactions in β‐Aminoketones

Reported are the syntheses of ester-functionalized (6-8) and alkyl-substituted (9) 1-aza-adamantanones; the easy handling of the compounds provides an opportunity to comprehensively study the fundamental changes in structure and reactivity that can accompany the donor-acceptor arrangement in rigid beta-aminoketones. X-ray structural analysis of trione 6 and dione 7 reveals bond length and angle variations consistent with through-bond (hyperconjugative) donor-acceptor interactions. Observed is a shortening of the C-N bond, elongation of the central C-C bond (to approximately 1.6 A), and a significant pyramidalization of the carbonyl carbon within the donor-sigma-acceptor pathway. UV/Vis spectra of 6-9 show a new absorption maximum (lambda(max)=260-275 nm in three solvents), the so-called "sigma-coupled transition"; the molar absorptivity scales with the number of carbonyl groups (for trione 6, epsilon approximately 3000, for dione 7, epsilon approximately 2000) and the band reversibly disappears upon addition of acid. IR and (13)C NMR spectroscopic data show trends consistent with through-bond donation to the carbonyl acceptor groups and commensurate weakening of the carbonyl pi bond. High yielding acid-mediated fragmentations are used to illustrate the effects of the donor-acceptor arrangement on the reactivity of the molecules. Given that donor-sigma-acceptor molecules have recently been found to show self-assembly behavior and macromolecular properties linked to their unusual structure, the current analysis encourages further consideration of the systems in advanced materials applications.

Solubility and Acid–Base Properties of Ethylenediaminetetraacetic Acid in Aqueous NaCl Solution at 0 ≤ I ≤ 6 mol·kg−1 and T = 298.15 K

Solubility and protonation constants of ethylenediaminetetraacetic acid (EDTA) were determined in aqueous NaCl solutions at different ionic strengths (0 ≤ I ≤ 6.0 mol·kg−1), at T = 298.15 K. Dependence on the ionic strength of protonation constants was modeled by the modified SIT approach on the molal concentration scale. From total solubility, we calculated the solubility of the neutral species S0, and from these values, as a function of NaCl concentration, the Setschenow (km) coefficient was obtained, together with the values of KS0.

Measurement of mean ionic activity coefficients for tetra-n-butyl ammonium salts of PF 6 − and [B(C6F5)4]− in tetrahydrofuran and acetonitrile

Osmotic coefficients were measured for tetra-n-butylammonium hexafluorophosphate (TBA PF6) and tetra-n-butylammonium tetrakis(pentafluorophenyl)borate (TBABARF) in tetrahydrofuran and in acetonitrile. These osmotic coefficients were fit using equations first derived by Pitzer. Mean ionic activity coefficients were determined using the integrated form of the Pitzer equations. Densities were measured for each of the solutions and the results were used to convert mean ionic activity coefficients on the molal scale to the molar scale. Remarkably, over concentrations ranging from 0.1 to 1.0 m in acetonitrile, the behavior of TBABARF is consistent with the simple Debye-Huckel equation.

Solubility of the Single Gases Methane and Xenon in the Ionic Liquid [hmim][Tf2N

This work reports new experimental results for the solubility of two single gases (methane and xenon) in the ionic liquid 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ([hmim][Tf2N]). The solubility was measured with a high-pressure view-cell technique operating on the synthetic method. The temperature ranged from 293.3 to 413.3 K, and the maximum pressure (the maximum gas molality) was 9.3 MPa (0.51 mol·kg-1) for methane and 9.6 MPa (2.08 mol·kg-1) for xenon. Xenon shows a significantly higher solubility than methane at all conditions investigated. Both gases become less soluble in [hmim][Tf2N] with rising temperature. An extension of Henry's law is employed to correlate the solubility pressures. The final results for the Henry's constants (at zero pressure) of methane and xenon in [hmim][Tf2N] (on the molality scale) are represented within the experimental uncertainty (about 1.2% for methane and 1% for xenon) by ln( /MPa) = 6.4929 − 835.9/(T/K) − 0.003471(T/K) and ln( /MPa) = 6.3357 − 13...

Solubility of the Single Gases Methane and Xenon in the Ionic Liquid [bmim][CH3SO4

Solubility measurements of methane and of xenon in the ionic liquid 1-n-butyl-3-methylimidazolium methyl sulfate ([bmim][CH3SO4]) were performed with a high-pressure view-cell technique based on the synthetic method. The temperature ranged from 293.1 K to 413.2 K, and the maximum pressure (the maximum gas molality) was 8.9 MPa (0.19 mol·kg-1) for methane and 11.3 MPa (0.65 mol·kg-1) for xenon. Both gases become less soluble in [bmim][CH3SO4] with rising temperature, but xenon shows a significantly higher solubility at all conditions investigated. An extension of Henry's law is employed to correlate the solubility pressures in both cases. The final results for the Henry's constant (at zero pressure) of methane and xenon in [bmim][CH3SO4] (on the molality scale) are correlated within the experimental uncertainty (about ± 1.3 %) by ln( /MPa) = 6.216 − 564.5/(T/K) − 0.002566(T/K) and ln( /MPa) = 5.5906 − 928/(T/K) − 0.00141(T/K), respectively.

Determination of the Dissociation Constants of Urocanic Acid Isomers in Aqueous Solutions

Summary. (E)- and (Z)-Urocanic acids are endogenous chemicals in the normal mammalian skin. The first and the second thermodynamic dissociation constants (pKa1 and pKa2) of urocanic acid isomers were determined using UV spectrophotometry in aqueous solutions. The values with standard deviation (pKa1 ¼ 3.43 � 0.12 and pKa2 ¼ 5.80 � 0.04) and (pKa1 ¼ 2.7 � 0.3 and pKa2 ¼ 6.65 � 0.04) were obtained to (E)- and (Z)-urocanic acids, respectively. The second dissociations were studied also by potentiometric titration in aqueous sodium chloride solutions up to the isotonic salt concentration (0.154 mol dm � 3 ), and the second thermodynamic dissociation constants as well as activity parameters for both isomers were determined at temperature 25 � Ca nd for (E)urocanic acid also at 37 � C. The obtained pKa2 values were close to those found by UV spectrophotometry. The equations for the calculation of the second stoichiometric dissociation constants of urocanic acid isomers on molality and molarity scale in aqueous sodium chloride solutions were derived. The obtained pKa1 and pKa2 values for (Z)-urocanic acid appear to be essentially lower than some previously reported values in literature.