Emmerich Wilhelm

Estimation of Lennard‐Jones (6,12) Pair Potential Parameters from Gas Solubility Data

The Lennard‐Jones (6,12) pair potential parameters for solvents can be calculated from gas solubility data by using the scaled particle theory to calculate the work of cavity formation. Values of the hard sphere diameter σ1 and the energy parameter e1/k at 298.15°K and 1 atm pressure are reported for 39 liquids, polar as well as nonpolar. These parameters are discussed in relation to values determined by other methods.

On the temperature dependence of the effective hard sphere diameter

For a number of solvents, the effective hard sphere diameter at several temperatures is determined from gas solubility data using the expression of the scaled particle theory of the work of cavity formation. From the results, the linear coefficients of expansion l1 of the effective hard core diameter are calculated and discussed in relation to values determined by other methods.

Solubilities of gases in liquids II. The solubilities of He, Ne, Ar, Kr, O2, N2, CO, CO2, CH4, CF4, and SF6 in n-octane 1-octanol, n-decane, and 1-decanol

Abstract The solubilities of 11 gases in n-octane, n-decane, 1-octanol, and 1-decanol have been determined at atmospheric pressure in the range 293 to 313 K. From the temperature variation of the experimental solubilities, partial molar enthalpies and entropies of solution for 1 atm partial pressure of gas and 298.15 K have been derived. Comparison with results obtained through application of scaled-particle theory showed satisfactory agreement.

Excess molar heat capacities of (1,4-dioxane+ an n-alkane): an unusual composition dependence

Abstract Excess molar heat capacities C p ,m E at constant pressure were measured, as a function of mole fraction x at 298.15 K and atmospheric pressure, for { x 1,4-C 4 H 8 O 2 + (1− x )C n H 2 n +2 } for n = 7, 10, and 14. The instrument used was a Picker flow calorimeter. The composition dependence of C p ,m E of these three mixtures is strikingly unusual in that two minima are observed in each case, the more prominent being situated at small values of x : x ′ min = 0.185 for n = 7; x ′ min = 0.251 for n = 10; and x ′ min = 0.289 for n = 14. With increasing chain length the corresponding excess molar heat capacities become more negative: C p ,m E ( x ′ min )/(J·K −1 ·mol −1 ) is, respectively, −0.95, −1.53, and −2.53. The less prominent minima are located around x ″ min = 0.9, with C p ,m E ( x ″ min )/(J·K −1 ·mol −1 ) being −0.40 for n = 7, −0.54 for n = 10, and −0.63 for n = 14. The mole fractions x max of the maxima of the three curves and the corresponding values of C p ,m E ( x max )/J·K −1 ·mol −1 ) are 0.605, 0.03; 0.699, 0.03; 0.780, 0.22.

Thermodynamics of solutions: selected aspects

Abstract This article reviews recent advances in the thermodynamics of (dilute) non-electrolyte solutions. The focus is on high-precision experimental techniques. Some of the problems encountered in data reduction and data correlation are discussed. “Der einzige sichere Fuhrer auf dem Weg der weiteren Entwicklung bleibt stets die Messung...” Max Planck (1926), Physikalische Gesetzlichkeit .

Excess volumes and excess heat capacities of oxane + cyclohexane and 1,4-dioxane + cyclohexane

Abstract Molar excess volumes VE at 298.15 K have been determined as a function of mole fraction x for the two binary liquid systems oxane (tetrahydropyran, C5H10O) + cyclohexane(c-C6H12) and 1,4-dioxane(1,4-C4H8O2) + cyclohexane by vibrating-tube densimetry. In addition, a Picker flow calorimeter was used to obtain molar excess heat capacities CPE at constant pressure at the same temperature. VE is positive for both systems and rather symmetric, with VE(x1 = 0.5) amounting to 0.317 cm3 mol−1 for &[;x1(C5H10O)+x2(c-C6H12)&];, and to 0.961 cm3 mol−1 for &[;x1(1,4-C4H8O2)+x2(c-C6H12)&];. For the former system, CPE is negative and noticeably skewed toward the cyclohexane side. The composition dependence of CPE for the latter system is rather unusual in that two minima are observed, the more pronounced being situated at x′1,min = 0.173, with CPE(x′1,min) = −0.97 J K−1 mol−1. The second minimum is rather shallow and is located at x″1,min = 0.761, with CPE(x″1,min) = −0.56 J K−1 mol−1. The maximum of the curve is at x1,max = 0.539, with CPE(x1,max) = −0.47 J K−1 mol−1.

On Solvophobic Interaction

An extended version of the scaled particle theory is used to calculate various thermodynamic functions pertaining to the solubility of gases in liquids. In particular, the Gibbs free energy of cavity formation Ḡc is determined for several solvent—gas systems. The usefulness of this quantity as a practical and relatively simple guide for comparing solvophobic interaction in various media is discussed.

Solubility of Gases in Liquids. XVI. Henry's Law Coefficients for Nitrogen in Water at 5 to 50°C

The solubility of nitrogen in pure liquid water was measured in the pressure range 45 to 115 kPa and in the temperature range 5 to 50°C. These data are used to obtain Henry coefficients H2,1 (T,Ps,1) at the vapor pressure Ps,1 of water. The temperature dependence of H2,1 (T,Ps,1) is accounted for by both a Clarke-Glew (CG) type fitting equation, and a power series in T−1, as suggested by Benson and Krause (BK). The imprecision of our measurements is characterized by an average deviation of ±0.038% from a four-term CG equation, and by an average deviation of ±0.042% from a three-term BK equation. From the temperature variation of H2,1 (T,Ps,1) partial molar quantities referring to the solution process, such as enthalpies and heat capacities of solution, are obtained. They are given in tabular form, together with H2,1 (T,Ps,1) and derived Ostwald coefficients L∞, at rounded temperatures. Finally, experimental results are compared with values calculated via scaled particle theory.

The solubility of gases in liquids I. The solubility of a series of fluorine-containing gases in several non-polar solvents

The solubilities of CF3Cl and CF2Cl2 in n-heptane, n-octane, cyclohexane, benzene, and carbon tetrachloride, and additionally CF4 and SF6 in n-octane were determined in the temperature range 24 to 35 °C. The solubility, Gibbs free energy, enthalpy, and entropy of solution were calculated from the experimental results for 1 atm partial pressure of gas and 25 °C. These values were compared with results calculated using the scaled particle theory. The agreement was good. The scaled particle theory was also used to determine the Lennard-Jones (6–12) pair potential parameters for the solvents, yielding results which are in excellent agreement with literature values.

Thermodynamics of (a halogenated ethane or ethene + an n-alkane). VE and CPE of mixtures containing either 1,1,2,2-tetrachloroethane or tetrachloroethene

Abstract Excess molar volumes V E and excess molar heat capacities C P E at constant pressure have been determined at 298.15 K as a function of mole fraction x 1 for mixtures belonging to series I: {x 1 1,1,2,2-C 2 H 2 Cl 4 + x 2 n-C n H 2n+2 }, and series II: {x 1 C 2 Cl 4 + x 2 n-C n H 2n+2 }, n = 7 and 14. While 1,1,2,2-tetrachloroethane (1,1,2,2-TCE) exhibits trans-gauche rotational isomerism, tetrachloroethene (TCEe) is a rigid molecule without permanent electric dipole moment. The instruments used were a vibrating-tube densimeter and a Picker flow calorimeter. For series I, V E (x 1 =0.5)/(cm 3 .mol −1 ) = 0.153 for n = 7, and 1.029 for n = 14, as compared to −0.192 for n = 7, and 0.313 for n = 14 in series II. The highly asymmetric shape of V E vs. x 1 of (1,1,2,2-TCE + n-C 7 H 16 ) is noted. For series I, the composition dependence of C P E as well as its dependence on n are similar to those for the series (1,2-dichloroethane + an n-alkane) in that for n = 7 the minimum (−2.18 J.K −1 .mol −1 ) is at x 1,min = 0.364 and a shoulder extends to, roughly, x 1 ≈ 0.75. For n = 14, C P E (x 1,min ) = −4.77 J.K −1 .mol −1 at x 1,min = 0.423, and no shoulder is discernible. The curves C P E vs. x 1 for series II are more or less parabolic, with C P E (x 1,min )/(J.K −1 .mol −1 ) = −0.14 at x 1,min = 0.512 for n = 7, and −1.66 at x 1,min = 0.502 for n = 14.