Investigation of the Thermodynamic Properties of the Cationic Surfactant CTAC in EG + Water Binary Mixtures

To understand the thermodynamic characteristics of cationic surfactants in binary mixtures, the aggregation behavior of hexadecyltrimethylammonium chloride (CTAC) has been investigated in ethylene glycol (EG) + water solvent mixtures at different temperatures and EG to water ratios. The critical micelle concentration (CMC) and degree of counter ion bonding (β) were calculated from electrical conductivity measurements. An equilibrium model for micelle formation was applied to obtain the thermodynamic parameters for micellization, including the standard Gibbs energies of micellization (DG_mic^o)\Delta G_{\mathrm{mic}}^{\mathrm{o}}), standard enthalpies of micelle formation (DH_mic^o)\Delta H_{\mathrm{mic}}^{\mathrm{o}}) and standard entropies of micellization (DS_mic^o)\Delta S_{\mathrm{mic}}^{\mathrm{o}}). Our results show that DG_mic^o\Delta G_{\mathrm{mic}}^{\mathrm{o}} is always negative and slightly dependent on temperature. The process of micellization is entropy driven in pure water, whereas in EG + water mixtures the micellization is enthalpy driven.     

Stereo-dynamics study of O + HCl → OH + Cl reaction on the <Superscript>3</Superscript>A″, <Superscript>3</Superscript>A′, and <Superscript>1</Superscript>A′ states

Using three accurate potential energy surfaces of the ³A″, ³A′, and ¹A′ states constructed recently, we present a quasi-classical trajectory (QCT) calculation for O + HCl (v = 0, j = 0) → OH + Cl reaction at the collision energies (E _col) of 14.0–20.0 kcal/mol. The three angular distribution functions—P(q_r ) P(\theta_{r} ) , P(j_r ) P(\varphi_{r} ) , and P(q_r ,j_r ) P(\theta_{r} ,\varphi_{r} ) , together with the four commonly used polarization-dependent differential cross-sections, \frac2ps \fracds₀₀ dw_t , \frac2ps \fracds₂₀ dw_t , \frac2ps \fracds_{22 +} dw_t , \textand \frac2ps \fracds_{21 -} dw_t {\frac{2\pi }{\sigma }}\,{\frac{{d\sigma_{00} }}{{d\omega_{t} }}},\,{\frac{2\pi }{\sigma }}\,{\frac{{d\sigma_{20} }}{{d\omega_{t} }}},\,{\frac{2\pi }{\sigma }}\,{\frac{{d\sigma_{22 + } }}{{d\omega_{t} }}},\,{\text{and}}\,{\frac{2\pi }{\sigma }}\,{\frac{{d\sigma_{21 - } }}{{d\omega_{t} }}} are exhibited to get an insight into the alignment and the orientation of the product OH radical. There is a similar behavior of the tendency scattering direction for the two triplet electronic states (³A″ and ³A′)—backward scattering dominates, however, forward scattering prevails for the case of ¹A′ state. Also, obvious differences have been found in the stereo-dynamical information, which reveals the influences of the potential energy surface and the collision energy. The degrees of polarization and the influence of the collision energy on the stereo-dynamics characters of the title reaction are both demonstrated in the order of ³A′ > ³A″ > ¹A′.     

Specific adsorption of bromide and iodide ions from N-methyl formamide solutions with constant ionic strength on liquid ga electrode

The differential capacitance curves were measured with an ac bridge in the Ga/[N-MF + 0.1 m M KBr + 0.1 (1 − m) M KClO₄] and Ga/[N-MF + 0.1 m M KI + 0.1 (1 − m) M KClO₄] systems at the following fractions m of surface-active anions: 0, 0.01, 0.02, 0.05, 0.1, 0.2, 0.5, and 1. As compared with other solvents, N-methyl formamide (N-MF) enables one to realize the largest positive charges of Ga electrode, at which it remains ideally polarizable (up to 20 μ/cm²). The data on the specific adsorption of Br⁻ and I⁻ anions in the system can be quantitatively described by the Frumkin’s isotherm; to the first approximation, free energy of halide ion (Hal⁻) adsorption DG_{adsHal ^{- 1}} \Delta G_{adsHal^{ - 1} } is a linear function of electrode charge. It is found that, in contrast to the Hg/N-MF interface, DG_{adsHal ^{- 1}} \Delta G_{adsHal^{ - 1} } at the Ga/N-MF interface varies in the reverse order: Br^t— ∼ I⁻ < Cl⁻. From the measured results, we can conclude that the energy of metal-Hal⁻ interaction increases in series: $\Delta G_{M - Cl^ - } > \Delta G_{M - Br^ - } > \Delta G_{M - I^ - } $\Delta G_{M - Cl^ - } > \Delta G_{M - Br^ - } > \Delta G_{M - I^ - } and the difference (DG_{Ga - Hal1^-} - DG_{Ga - Hal2^-} )(\Delta G_{Ga - Hal_1^ - } - \Delta G_{Ga - Hal_2^ - } ) is larger than the difference between the solvation energies of Hal^- (DG_{S - Hal1^-} - DG_{S - Hal2^-} )Hal^ - (\Delta G_{S - Hal_1^ - } - \Delta G_{S - Hal_2^ - } ).     

Electrocatalytic oxidation of nitrite at a terpyridine manganese(II) complex modified carbon past electrode

A carbon past electrode modified with [Mn(H₂O)(N₃)(NO₃)(pyterpy)], ( \textpyterpy = 4￠- ( 4 - \textpyridyl ) - 2,2￠:\text6￠,\text2￠￠- \textterpyridine ) \left( {{\text{pyterpy}} = 4\prime - \left( {4 - {\text{pyridyl}}} \right) - 2,2\prime:{\text{6}}\prime,{\text{2}}\prime\prime - {\text{terpyridine}}} \right) complex have been applied to the electrocatalytic oxidation of nitrite which reduced the overpotential by about 120 mV with obviously increasing the current response. Relative standard deviations for nitrite determination was less than 2.0%, and nitrite can be determined in the ranges of 5.00 × 10⁻⁶ to 1.55 × 10⁻² mol L⁻¹, with a detection limit of 8 × 10⁻⁷ mol L⁻¹. The treatment of the voltammetric data showed that it is a pure diffusion-controlled reaction, which involves one electron in the rate-determining step. The rate constant k′, transfer coefficient α for the catalytic reaction, and diffusion coefficient of nitrite in the solution, D, were found to be 1.4 × 10⁻², 0.56× 10⁻⁶, and 7.99 × 10⁻⁶ cm² s⁻¹, respectively. The mechanism for the interaction of nitrite with the Mn(II) complex modified carbon past electrode is proposed. This work provides a simple and easy approach to detection of nitrite ion. The modified electrode indicated reproducible behavior, anti-fouling properties, and stability during electrochemical experiments, making it particularly suitable for the analytical purposes.     

The Estimation of Standard Molar Enthalpies of Solution for VOSO<Subscript>4</Subscript>⋅<Emphasis Type="Italic">n</Emphasis>H<Subscript>2</Subscript>O(s) in Water and in Aqueous H<Subscript>2</Subscript>SO<Subscript>4</Subscript>

The molar enthalpies of solution of VOSO₄⋅3.52H₂O(s) at various molalities in water and in aqueous sulfuric acid (0.1 mol⋅kg⁻¹), Δ_sol H _m, were measured by a solution-reaction isoperibol calorimeter at 298.15±0.01 K. An improved Archer’s method to estimate the standard molar enthalpy of solution, D_solH⁰_m\Delta_{\mathrm{sol}}H^{0}_{\mathrm{m}}, was put forward. In terms of the improved method, the values of D_solH⁰_m=-24.12±0.03 kJ·mol^-1\Delta_{\mathrm{sol}}H^{0}_{\mathrm{m}}=-24.12\pm 0.03~\mbox{kJ}{\cdot}\mbox{mol}^{-1} of VOSO₄⋅3.52H₂O(s) in water and D_solH⁰_m=-15.38±0.06 kJ·mol^-1\Delta_{\mathrm{sol}}H^{0}_{\mathrm{m}}=-15.38\pm 0.06~\mbox{kJ}{\cdot}\mbox{mol}^{-1} in aqueous sulfuric acid were obtained, respectively. The data indicates that the energy state of VOSO₄ in aqueous H₂SO₄ is higher than that in pure water.     

Raman and Infrared Spectroscopic Investigation of Speciation in BeSO<Subscript>4</Subscript>(aq)

Measurements have been made of the Raman spectra of aqueous solutions of Be(ClO₄)₂, BeCl₂, (NH₄)₂SO₄ and BeSO₄ to 50 cm⁻¹. In some cases low concentrations (0.000770 mol⋅kg⁻¹) have been used and two temperatures (23 and 40 °C) were studied. In BeSO₄(aq), the ν ₁-SO₄^2-\mathrm{SO}_{4}^{2-} mode at 980 cm⁻¹ broadens with increasing concentration and shifts to higher wavenumbers. At the same time, a band at 1014 cm⁻¹ is detectable with this mode being assigned to [BeOSO₃], an inner-sphere complex (ISC). Confirmation of this assignment is provided by the simultaneous appearance of stretching bands for the Be²⁺-OSO₃^2-\mathrm{Be}^{2+}\mbox{-}\mathrm{OSO}_{3}^{2-} bond of the complex at 240 cm⁻¹ and for the BeO₄ skeleton mode of the [(H₂O)₃BeOSO₃] unit at 498 cm⁻¹. The ISC concentration increases with higher temperatures. The similarity of the n₁-SO₄^2-\nu_{1}\mbox{-}\mathrm{SO}_{4}^{2-} Raman bands for BeSO₄ in H₂O and D₂O is further strong evidence for formation of an ISC. After subtraction of the ISC component at 1014 cm⁻¹, the n₁-SO₄^2-\nu_{1}\mbox{-}\mathrm{SO}_{4}^{2-} band in BeSO₄(aq) showed systematic differences from that in (NH₄)₂SO₄(aq). This is consistent with a n₁-SO₄^2-\nu_{1}\mbox{-}\mathrm{SO}_{4}^{2-} mode at 982.7 cm⁻¹ that can be assigned to the occurrence of an outer-sphere complex ion (OSCs). These observations are shown to be in agreement with results derived from previous relaxation measurements. Infrared spectroscopic data show features that are also consistent with a beryllium sulfato complex such as the appearance of a broad and weak n₁-SO₄^2-\nu_{1}\mbox{-}\mathrm{SO}_{4}^{2-} mode at ∼1014 cm⁻¹, normally infrared forbidden, and a broad and asymmetric n₃-SO₄^2-\nu_{3}\mbox{-}\mathrm{SO}_{4}^{2-} band contour which could be fitted with four band components (including n₃-SO₄^2-(aq)\nu_{3}\mbox{-}\mathrm{SO}_{4}^{2-}(\mathrm{aq})). The formation of ISCs in BeSO₄(aq) is much more pronounced than in the similar MgSO₄(aq) system studied recently.     

Dissolution Thermodynamics of 1,2,3-Triazole Nitrate in Water

The enthalpies of dissolution of 1,2,3-triazole nitrate in water were measured using a RD496-2000 Calvet microcalorimeter at four different temperatures under atmospheric pressure. Differential enthalpies (Δ_dif H) and molar enthalpies (Δ_diss H) of dissolution were determined. The corresponding kinetic equations that describe the dissolution rate at the four experimental temperatures are \fracdadt / s ^{- 1} = 10 ^{- 3.75}( 1 - a)^0.96\frac{d\alpha}{dt} / \mathrm{s}^{ - 1} =10^{ - 3.75}( 1 - \alpha)^{0.96} (T=298.15 K), \fracdadt /s ^{- 1} = 10 ^{- 3.73}( 1 - a)^1.00\frac{d\alpha}{dt} /\mathrm{s}^{ - 1} = 10^{ - 3.73}( 1 - \alpha)^{1.00} (T=303.15 K), \fracdadt / s ^{- 1} = 10 ^{- 3.72}( 1 - a)^0.98\frac{d\alpha}{dt} / \mathrm{s}^{ - 1} = 10^{ - 3.72}( 1 - \alpha)^{0.98} (T=308.15 K) and \fracdadt / s ^{- 1} = 10 ^{- 3.71}( 1 -a)^0.97\frac{d\alpha}{dt} / \mathrm{s}^{ - 1} = 10^{ - 3.71}( 1 -\alpha)^{0.97} (T=313.15 K). The determined values of the activation energy E and pre-exponential factor A for the dissolution process are 5.01 kJ⋅mol⁻¹ and 10^−2.87 s⁻¹, respectively.     

Solubility Measurements of Crystalline NiO in Aqueous Solution as a Function of Temperature and pH

Results of solubility experiments involving crystalline nickel oxide (bunsenite) in aqueous solutions are reported as functions of temperature (0 to 350 °C) and pH at pressures slightly exceeding (with one exception) saturation vapor pressure. These experiments were carried out in either flow-through reactors or a hydrogen-electrode concentration cell for mildly acidic to near neutral pH solutions. The results were treated successfully with a thermodynamic model incorporating only the unhydrolyzed aqueous nickel species (viz., Ni²⁺) and the neutrally charged hydrolyzed species (viz., Ni(OH)₂⁰)\mathrm{Ni(OH)}_{2}^{0}). The thermodynamic quantities obtained at 25 °C and infinite dilution are, with 2σ uncertainties: log₁₀K_s0^o = (12.40 ±0.29),\varDelta_rG_m^o = -(70. 8 ±1.7)\log_{10}K_{\mathrm{s0}}^{\mathrm{o}} = (12.40 \pm 0.29),\varDelta_{\mathrm{r}}G_{m}^{\mathrm{o}} = -(70. 8 \pm 1.7) kJ⋅mol⁻¹; \varDelta_rH_m^o = -(105.6 ±1.3)\varDelta_{\mathrm{r}}H_{m}^{\mathrm{o}} = -(105.6 \pm 1.3) kJ⋅mol⁻¹; \varDelta_rS_m^o = -(116.6 ±3.2)\varDelta_{\mathrm{r}}S_{m}^{\mathrm{o}} =-(116.6 \pm 3.2) J⋅K⁻¹⋅mol⁻¹; \varDelta_rC_p,m^o = (0 ±13)\varDelta_{\mathrm{r}}C_{p,m}^{\mathrm{o}} = (0 \pm 13) J⋅K⁻¹⋅mol⁻¹; and log₁₀K_s2^o = -(8.76 ±0.15)\log_{10}K_{\mathrm{s2}}^{\mathrm{o}} = -(8.76 \pm 0.15); \varDelta_rG_m^o = (50.0 ±1.7)\varDelta_{\mathrm{r}}G_{m}^{\mathrm{o}} = (50.0 \pm 1.7) kJ⋅mol⁻¹; \varDelta_rH_m^o = (17.7 ±1.7)\varDelta_{\mathrm{r}}H_{m}^{\mathrm{o}} = (17.7 \pm 1.7) kJ⋅mol⁻¹; \varDelta_rS_m^o = -(108±7)\varDelta_{\mathrm{r}}S_{m}^{\mathrm{o}} = -(108\pm 7) J⋅K⁻¹⋅mol⁻¹; \varDelta_rC_p,m^o = -(108 ±3)\varDelta_{\mathrm{r}}C_{p,m}^{\mathrm{o}} = -(108 \pm 3) J⋅K⁻¹⋅mol⁻¹. These results are internally consistent, but the latter set differs from those gleaned from previous studies recorded in the literature. The corresponding thermodynamic quantities for the formation of Ni²⁺ and Ni(OH)₂⁰\mathrm{Ni(OH)}_{2}^{0} are also estimated. Moreover, the Ni(OH)₃ ^-\mathrm{Ni(OH)}_{3}^{ -} anion was never observed, even in relatively strong basic solutions (m_{OH ^-} = 0.1m_{\mathrm{OH}^{ -}} = 0.1 mol⋅kg⁻¹), contrary to the conclusions drawn from all but one previous study.     

Evidence of Different Stoichiometries for the Limiting Carbonate Complexes across the Lanthanide(III) Series

The stoichiometries of limiting carbonate complexes of lanthanide(III) ions were investigated by solubility measurements of hydrated NaLn(CO3)2 solid compounds (Ln = La, Nd, Eu and Dy) at room temperature in aqueous solutions of high ionic strength (3.5 mol⋅kg−1 NaClO4) and high CO32-\mathrm{CO_{3}^{2-}} concentrations (0.1 to 1.5 mol⋅kg−1). The results were interpreted by considering the stability of carbonate complexes, with limiting species found to be La(CO3)45-\mathrm{La(CO_{3})_{4}^{5-}}, Nd(CO3)45-\mathrm{Nd(CO_{3})_{4}^{5-}}, Eu(CO3)33-\mathrm{Eu(CO_{3})_{3}^{3-}} and Dy(CO3)33-\mathrm{Dy(CO_{3})_{3}^{3-}}. TRLFS measurements on the Eu and Dy solutions confirmed the predominance of a single aqueous complex in all the samples. Equilibrium constants were determined for the reaction Ln(CO3)33-+CO32-\mathrm{Ln(CO_{3})_{3}^{3-}}+\mathrm{CO_{3}^{2-}} ⇌ Ln(CO3)45-\mathrm{Ln(CO_{3})_{4}^{5-}}: log10K3.5m NaClO44,La=0.7±0.3\log_{10}K\mathrm{^{3.5m\:NaClO_{4}}_{4,La}=0.7\pm0.3}, log10K3.5m NaClO44,Nd=1.3±0.3\log_{10}K\mathrm{^{3.5m\:NaClO_{4}}_{4,Nd}=1.3\pm0.3}, and for Ln = Eu and Dy, log10K3.5m NaClO44,Ln ￡ -0.4\log_{10}K\mathrm{^{3.5m\:NaClO_{4}}_{4,Ln}\leq-0.4}. These results suggest that tetracarbonato complexes are stable only for the light lanthanide ions in up to 1.5 molal CO32-\mathrm{CO_{3}^{2-}} aqueous solutions, in agreement with our recent capillary electrophoresis study. Comparison with literature results indicates that analogies between actinide(III) and lanthanide(III) ions of similar ionic radii do not hold in concentrated carbonate solutions. Am(CO3)33-\mathrm{Am(CO_{3})_{3}^{3-}} was previously evidenced by solubility measurements, whereas we have observed that Nd(CO3)45-\mathrm{Nd(CO_{3})_{4}^{5-}} predominates in similar conditions. We may speculate that small chemical differences between Ln(III) and An(III) could result in macroscopic differences when their coordination sphere is complete.     

Volumetric Properties of Amino Acids in Aqueous N-methylacetamide Solutions at 298.15 K

The apparent molar volumes (V _φ ) of glycine, L-alanine and L-serine in aqueous 0 to 4 mol⋅kg⁻¹ N-methylacetamide (NMA) solutions have been obtained by density measurement at 298.15 K. The standard partial molar volumes (V_f⁰)V_{\phi}^{0}) and standard partial molar volumes of transfer (D_trV_f⁰)\Delta_{\mathrm{tr}}V_{\phi}^{0}) have been determined for these amino acids. It has been show that hydrophilic-hydrophilic interactions between the charged groups of the amino acids and the –CONH– group of NMA predominate for glycine and L-serine, but for L-alanine the interactions between its side group (–CH₃) and NMA predominate. The –CH₃ group of L-alanine has much more influence on the value of D_trV_f⁰\Delta_{\mathrm{tr}}V_{\phi}^{0} than that of the –OH group of L-serine. The results have been interpreted in terms of a co-sphere overlap model.     

Ion-Solvation Behavior of Pyridinium Dichromate in Water–N,N-Dimethyl Formamide Solvent Mixtures

The electrical conductances of pyridinium dichromate have been measured in N,N-dimethyl formamide–water mixtures of different compositions in the temperature range 283–313 K. The limiting molar conductance, Λ₀, association constant of the ion pair, K _A, and dissociation constant K _C have been calculated using the Shedlovsky and Kraus–Bray equations. The effective ionic radii (r _i ) of C₅H₅NH⁺ and Cr₂O₇ ^-\mathrm{Cr}_{2}\mathrm{O}_{7}^{ -} have been determined from the L_i⁰\Lambda_{i}^{0} values using Gill’s modification of Stokes’ law. The influence of the mixed solvent composition on the solvation of ions is discussed with the help of the ‘R’-factor ( R = \frachL _±⁰(solvent)hL _±⁰(water)R = \frac{\eta \Lambda_{ \pm}^{0}(\mathrm{solvent})}{\eta\Lambda_{ \pm}^{0}(\mathrm{water})}). Thermodynamic parameters are evaluated and reported. The results of this study are interpreted in terms of ion–solvent interactions and solvent properties.     

Dynamics of Hydration of Alkylsulfonate Anions in Aqueous Solutions

The ¹⁷O-NMR spin-lattice relaxation times (T ₁) of water molecules in aqueous solutions of n-alkylsulfonate (C₁ to C₆) and arylsulfonic anions were determined as a function of concentration at 298 K. Values of the dynamic hydration number, (S^-) = n_h ^- (t_c^- /t_c⁰ - 1)(\mathrm{S}^{-}) = n_{\mathrm{h}}^{ -} (\tau_{\mathrm{c}}^{-} /\tau_{\mathrm{c}}^{0} - 1), were determined from the concentration dependence of T ₁. The ratios (t_c ^-/t_c⁰\tau_{\mathrm{c}}^{ -}/\tau_{\mathrm{c}}^{0}) of the rotational correlation times (t_c ^-\tau_{\mathrm{c}}^{ -} ) of the water molecules around each sulfonate anion in the aqueous solutions to the rotational correlation time of pure water (t_c⁰\tau_{\mathrm{c}}^{0}) were obtained from the n _DHN(S⁻) and the hydration number (n_h ^-n_{\mathrm{h}}^{ -} ) results, which was calculated from the water accessible surface area (ASA) of the solute molecule. The t_c ^-/t_c⁰\tau_{\mathrm{c}}^{ -}/\tau_{\mathrm{c}}^{0} values for alkylsulfonate anions increase with increasing ASA in the homologous-series range of C₁ to C₄, but then become approximately constant. This result shows that the water structures of hydrophobic hydration near large size alkyl groups are less ordered. The rotational motions of water molecules around an aromatic group are faster than those around an n-alkyl group with the same ASA. That is, the number of water–water hydrogen bonds in the hydration water of aromatic groups is smaller in comparison with the hydration water of an n-alkyl group having the same ASA. Hydrophobic hydration is strongly disturbed by a sulfonate group, which acts as a water structure breaker. The disturbance effect decreases in the following order: $\mbox{--} \mathrm{SO}_{3}^{-} > \mbox{--} \mathrm{NH}_{3}^{ +} > \mathrm{OH}> \mathrm{NH}_{2}$\mbox{--} \mathrm{SO}_{3}^{-} > \mbox{--} \mathrm{NH}_{3}^{ +} > \mathrm{OH}> \mathrm{NH}_{2}. The partial molar volumes and viscosity B _V coefficients for alkylsulfonate anions are linearly dependent on their n _DHN(S⁻) values.     

The kinetics of the reduction of tetraoxoiodate(VII) by <Emphasis Type="Italic">n</Emphasis>-(2-hydroxylethyl)ethylenediaminetriacetatocobaltate(II) ion in aqueous perchloric acid

The stoichiometries, kinetics and mechanism of the reduction of tetraoxoiodate(VII) ion, IO₄ ⁻ to the corresponding trioxoiodate(V) ion, IO₃ ⁻ by n-(2-hydroxylethyl)ethylenediaminetriacetatocobaltate(II) ion, [CoHEDTAOH₂]⁻ have been studied in aqueous media at 28 °C, I = 0.50 mol dm⁻³ (NaClO₄) and [H⁺] = 7.0 × 10⁻³ mol dm⁻³. The reaction is first order in [Oxidant] and [Reductant], and the rate is inversely dependent on H⁺ concentration in the range 5.00 × 10⁻³ ≤ H⁺≤ 20.00 × 10⁻³ mol dm⁻³ studied. A plot of acid rate constant versus [H⁺]⁻¹ was linear with intercept. The rate law for the reaction is:

Dissolution properties of hexanitrohexaazaisowurtzitane (CL-20) in ethyl acetate and acetone

The enthalpies of dissolution in ethyl acetate and acetone of hexanitrohexaazaisowurtzitane (CL-20) were measured by means of a RD496-2000 Calvet microcalorimeter at 298.15 K, respectively. Empirical formulae for the calculation of the enthalpy of dissolution (Δ_diss H), relative partial molar enthalpy (Δ_diss H _partial), relative apparent molar enthalpy (Δ_diss H _apparent), and the enthalpy of dilution (Δ_dil H _1,2) of each process were obtained from the experimental data of the enthalpy of dissolution of CL-20. The corresponding kinetic equations describing the two dissolution processes were \frac\textda\textdt = 1.60 ×10 ^{- 2} (1 - a)^0.84 {\frac{{{\text{d}}\alpha }}{{{\text{d}}t}}} = 1.60 \times 10^{ - 2} (1 - \alpha )^{0.84} for dissolution process of CL-20 in ethyl acetate, and \frac\textda\textdt = 2.15 ×10 ^{- 2} (1 - a)^0.89 {\frac{{{\text{d}}\alpha }}{{{\text{d}}t}}} = 2.15 \times 10^{ - 2} (1 - \alpha )^{0.89} for dissolution process of CL-20 in acetone.     

The Enthalpy of Solution of the Alanine-Based Ionic Liquid [C<Subscript>4</Subscript>mim][Ala]

The molar enthalpies of solution of an alanine-based ionic liquid (IL) [C₄mim][Ala], 1-butyl-3-methylimidazolium alanine, containing various amount of water and various molalities Δ_sol H _m(wc), were measured with a solution-reaction isoperibol calorimeter at (298.15±0.01) K, where wc denotes water content. According to Archer’s method, the standard molar enthalpies of solution of [C₄mim][Ala] containing known amounts of water, D_solH_m^o(wc)\Delta_{\mathrm{sol}}H_{\mathrm{m}}^{\mathrm{o}}(\mathrm{wc}) , were obtained. In order to eliminate the effect of the small amount of residual water in the source [C₄mim][Ala], a linear fitting of D_solH_m^o(wc)\Delta_{\mathrm{sol}}H_{\mathrm{m}}^{\mathrm{o}}(\mathrm{wc}) against water content was carried out, yielding a good straight line where the intercept is the standard molar enthalpy of solution of anhydrous [C₄mim][Ala], D_solH_m^o(pure IL)=-(61.42±0.08)\Delta_{\mathrm{sol}}H_{\mathrm{m}}^{\mathrm{o}}(\mathrm{pure}\ \mathrm{IL})=-(61.42\pm 0.08) kJ⋅mol⁻¹. The hydration enthalpy of the alanine anion [Ala]⁻ was estimated using Glasser’s lattice energy theory.     

Solubility of LIDOCAINE in Ionic, Nonionic and Zwitterionic Surfactants

The solubility of an anesthetic drug, LIDOCAINE, in water was investigated in the presence of ionic, nonionic and zwitterionic surfactants at 25 °C, and the solubility was found to increase linearly with the surfactant concentration. The molar solubilization ratio, R _m,s, and Gibbs free energy, DG_s^o\Delta G_{\mathrm{s}}^{\mathrm{o}} values for nonionic surfactants fall in the order DDAO > Brij 35 > Brij 30, whereas for ionic and zwitterionic surfactants the order is DDAPS > DTAB > SDS. The high negative values of the Gibbs energies in the cases of DDAO and DDAPS prove them to be better surfactants for solubilizing this drug as compared to the other surfactants.     

Thermochemistry of 2-Aminopyridine (C<Subscript>5</Subscript>H<Subscript>6</Subscript>N<Subscript>2</Subscript>)(s)

The molar enthalpies of solution of 2-aminopyridine at various molalities were measured at T=298.15 K in double-distilled water by means of an isoperibol solution-reaction calorimeter. According to Pitzer’s theory, the molar enthalpy of solution of the title compound at infinite dilution was calculated to be D_solH_m^￥ = 14.34 kJ·mol^-1\Delta_{\mathrm{sol}}H_{\mathrm{m}}^{\infty} = 14.34~\mbox{kJ}\cdot\mbox{mol}^{-1}, and Pitzer’s ion interaction parameters b_MX^(0)L, b_MX^(1)L\beta_{\mathrm{MX}}^{(0)L}, \beta_{\mathrm{MX}}^{(1)L}, and C_MX^fLC_{\mathrm{MX}}^{\phi L} were obtained. Values of the relative apparent molar enthalpies ( ^φ L) and relative partial molar enthalpies of the compound ([`(L)]₂)\bar{L}_{2}) were derived from the experimental enthalpies of solution of the compound. The standard molar enthalpy of formation of the cation C₅H₇N₂ ⁺\mathrm{C}_{5}\mathrm{H}_{7}\mathrm{N}_{2}^{ +} in aqueous solution was calculated to be D_fH_m^o(C₅H₇N₂⁺,aq)=-(2.096±0.801) kJ·mol^-1\Delta_{\mathrm{f}}H_{\mathrm{m}}^{\mathrm{o}}(\mathrm{C}_{5}\mathrm{H}_{7}\mathrm{N}_{2}^{+},\mbox{aq})=-(2.096\pm 0.801)~\mbox{kJ}\cdot\mbox{mol}^{-1}.     

Solute–Solvent Interactions of Amino Acids in Aqueous 1-Propyl-3-Methylimidazolium Bromide Ionic Liquid Solutions at 298.15 K

Densities, viscosities, and refractive indices of three amino acids (glycine, L-alanine, and L-valine) in aqueous solutions of an ionic liquid, 1-propyl-3-methylimidazolium bromide, have been measured at 298.15 K. These data have been used to calculate apparent molar volumes (V _φ ), viscosity B-coefficients, and molar refractions of these mixtures. The standard partial molar volumes (V_f⁰V_{\phi}^{0}) and standard partial molar volumes of transfer (D_trV_f⁰\Delta_{\mathrm{tr}}V_{\phi}^{0}) have been determined for these amino acid solutions from these density data. The resulting values of V_f⁰V_{\phi}^{0} and D_trV_f⁰\Delta_{\mathrm{tr}}V_{\phi}^{0} for transfer of amino acids from water to aqueous ionic liquid solutions have been interpreted in terms of solute + solvent interactions. These data also indicate that hydrophobic interactions predominate in L-alanine and L-valine solutions. Linear correlations were found for both V_f⁰V_{\phi}^{0} and the viscosity B-coefficient with the number of carbon atoms in the alkyl chain of the amino acids, and have been used to estimate the contribution of the charged end groups (NH₃⁺\mathrm{NH}_{3}^{+}, COO⁻), the CH₂ group, and other alkyl chains of the amino acids. The viscosity and molar refractivity results have been used to confirm the conclusions obtained from volumetric properties.     

Electrochemical Studies on <Emphasis Type="Italic">cis</Emphasis>-[Cr<Superscript>III</Superscript>(bipy)<Subscript>2</Subscript>(SCN)<Subscript>2</Subscript>]I<Subscript>3</Subscript> (Where bipy Denotes 2,2′-Bipyridine) in Acetonitrile

The molar conductivities (Λ) of solutions of bis(2,2′-bipyridine)bis(thiocyanate)chromium(III) triiodide [Cr^III(bipy)₂(SCN)₂]I₃ (where bipy denotes 2,2′-bipyridine, C₁₀H₈N₂), [ _3^-\mathrm{A}^{+}\mathrm{I}_{3}^{-} ], were measured in acetonitrile (ACN) at the temperatures 294.15, 299.15, and 305.15 K. In addition, cyclic voltammograms (CVs) of [ A⁺I₃^-\mathrm{A}^{+}\mathrm{I}_{3}^{-} ] were recorded on platinum, gold, and glassy carbon working electrodes in ACN, using n-tetrabutylammonium hexafluorophosphate (NBu₄PF₆) as the supporting electrolyte, at scan rates (v) ranging from 0.05 to 0.12 V⋅s⁻¹. Furthermore, electrochemical impedance spectroscopic (EIS) measurements were carried out in the frequency range 50 Hz<f<50 kHz using these three working electrodes. The measured molar conductivities (Λ) demonstrate that [ A⁺I₃^-\mathrm{A}^{+}\mathrm{I}_{3}^{-} ] behaves as uni-univalent electrolyte in ACN over the investigated temperature range. The Λ values were analyzed by means of the Lee-Wheaton conductivity equation in order to estimate the limiting molar conductivities (Λ _o), as well as the thermodynamic association constants (K _A), at each experimental temperature for formation of [A⁺ I₃^-\mathrm{I}_{3}^{-} ] ion-pairs. The limiting ionic conductivities ( l_±^o\lambda_{\pm}^{\mathrm{o}} ), the diffusion coefficients at infinite dilution (D _±), as well as the Stokes’ radii (r _St) were determined for both A⁺ and I₃^-\mathrm{I}_{3}^{-} ions. The thermodynamic parameters for the ionic association process, i.e. the Gibbs energy ( DG_A^o\Delta G_{\mathrm{A}}^{\mathrm{o}} ), enthalpy ( DH_A^o\Delta H_{\mathrm{A}}^{\mathrm{o}} ), and entropy ( DS_A^o\Delta S_{\mathrm{A}}^{\mathrm{o}} ), were also determined. The mobility and diffusivity of the A⁺ ion increase linearly with increasing temperature because the solvent medium becomes less viscous as the temperature increases. The K _A values indicate that significant ion association occurs that is not influenced by temperature changes. The ion-pair formation process is exothermic ( DH_A^o < 0\Delta H_{\mathrm{A}}^{\mathrm{o}}<0 ), leading to the generation of additional entropy ( $\Delta S_{\mathrm{A}}^{\mathrm{o}}>0$\Delta S_{\mathrm{A}}^{\mathrm{o}}>0 ). As a result, the Gibbs energy DG_A^o\Delta G_{\mathrm{A}}^{\mathrm{o}} is negative ( DG_A^o < 0\Delta G_{\mathrm{A}}^{\mathrm{o}}<0 ) and the formation of [A⁺I₃^-][\mathrm{A}^{+}\mathrm{I}_{3}^{-}] becomes favorable. CV studies on [A⁺I₃^-][\mathrm{A}^{+}\mathrm{I}_{3}^{-}] solutions indicated that the redox pair Cr^3+/2+ appears to be quasi-reversible on a glassy carbon electrode but is completely irreversible on platinum and gold electrodes. EIS experiments confirm that, among these three electrodes, the glassy carbon working electrode has the smallest resistance to electron transfer.