Showing papers on "Enthalpy published in 1969"

Phase equilibria of dodecylhexaoxyethylene glycol monoether in water

Abstract: Low-angle X-ray diffraction has been used to characterize the two lyotropic mesomorphic phases (neat phase and middle phase) which are formed in the dodecylhexaoxyethylene glycol monoether + water system. The occurrence of these two anisotropic mesomorphic phases does not affect the behaviour of the thermodynamic excess functions which have been interpreted in terms of solute-water interactions. Careful measurements of water vapour pressure have been necessary to establish the limited extent of the two phase co-existence regions which separate homogeneous mesomorphic phase regions from isotropic solution in this system. Volume and heat capacity measurements as a function of temperature across mesomorphic phase boundaries have confirmed that the neat phase/isotropic solution and the middle phase/isotropic solution transformations are first-order phase transitions with very small volume and enthalpy changes.

Thermodynamic Properties of Nonideal Gases. I. Free-Energy Minimization Method

Abstract: The thermodynamic properties of nonideal or interacting gases are studied using a thermodynamically consistent method which derives a state function for the gas from a separable Hamiltonian for the many-body system. The state function - the Helmholtz free energy - is a sum of translational, configurational, and internal components, and the statistical equilibrium state is calculated on a computer by minimizing the free energy in composition space at a given volume and temperature. A series of free-energy models is constructed to investigate the effects of various interactions. The effects on the equilibrium state of including Fermi-Dirac statistics in the translational free energy is shown to be small, while Coulomb and excluded-volume configurational terms produce large thermodynamic effects. The most important interaction effect, the modification of the internal partition function, is studied using a free-atom model and a confined-atom model which introduces volume- and temperature-dependent energy eigenvalues and partition-function sum terminations. The free-energy minimization method is shown to be a versatile and efficient tool for studying interacting gases, and provides quantitative estimates of the limits of validity of the various interaction models.

Thermodynamic properties of argon from the triple point to 300 K at pressures to 1000 atmospheres

Abstract: : Tabular values of density, internal energy, enthalpy, and entropy of liquid and gaseous argon are presented for temperature from 83.8 to 300 K at pressures of 0.01 to 1000 atmospheres. Diagrams of specific heats, compressibility factor, and enthropy are included. The properties presented are calculated from an equation of state which was fitted to experimental P-p-T data from the world literature. Extensive comparisons were made between the equation of state and the experimental data, and deviation plots are presented. The second virial coefficient and Joule Thomson inversion curve were also calculated and comparisons made with values from other sources. A vapor pressure equation which covers the range from the triple point to the critical point is also given. (Author)

Enthalpies of mixing and complex formation of carbon tetrachloride with benzene, toluene, p-xylene, and mesitylene

Abstract: Measurements are presented of the enthalpy of mixing of equimolar mixtures of carbon tetrachloride with benzene, toluene, p-xylene, and mesitylene at temperatures between 281 and 362 K. The results can be understood in terms of a simple theory of complex-formation, but only if it is assumed that while the equilibrium constant for complex-formation at room temperature increases in the sequence benzene < toluene < p-xylene < mesitylene, the magnitude of the (negative) enthalpy of complex-formation decreases in the same sequence. This conclusion prompts us to ask exactly what is meant by a ‘strong’ donor or ‘strong’ complex-formation. Measurements are also presented of the enthalpy of mixing of carbon tetrachloride and toluene over a range of mole fractions at each of five temperatures between 298 and 338 K. At temperatures close to that (313 K) at which the enthalpy of mixing of equimolar mixtures changes sign, a double maximum is predicted by the theory and found by experiment, no previous example of such behaviour being known. The theory is also fairly consistent with the values of the excess Gibbs function where these have been measured.

Conversion of Existing Calorimetrically Determined Thermodynamic Properties to the Basis of the International Practical Temperature Scale of 1968.

Abstract: Formulas are derived for converting the relative enthalpy, heat capacity, entropy, and Gibbs energy from the basis of one practical temperature scale to the basis of another, when these properties on either scale have been derived from calorimetric measurements of enthalpy as though that scale were the thermodynamic one. These formulas are directly applicable for converting certain other properties as well. The conversion relates the values of the property at the same numerical temperature on both scales. The formulas, given as exact infinite series, are applicable to widely differing scales, one of which may vary linearly with a temperature-measuring quantity such as electrical resistance. However, great simplification is well within most calorimetric accuracy when the conversion is from the International Practical Temperature Scale of 1948 to the corresponding scale of 1968, which has recently replaced it, provided the heat capacity is not changing abnormally rapidly, as in a transition region. For convenient application to conversion between these two scales, relatively simple numerical equations are derived giving the differences between the two scales at temperatures from 90 K to 10,000 K. The problem of avoiding the introduction of discontinuities with temperature in converted tables, arising from the existing discontinuities in the temperature derivative of the differences between the two scales, is discussed.

Concise Chemical Thermodynamics

Abstract: 1 Energy.- 1.1 The Realm of Thermodynamics.- 1.1.1 Energy Bookkeeping.- 1.1.2 Nature's Driving Forces.- 1.2 Setting the Scene Basic Ideas.- 1.2.1 System and Surroundings.- 1.2.2 Functions of State.- 1.2.3 Mechanical Work and Expanding Gases.- 1.2.4 The Absolute Temperature Scale.- 1.3 Forms of Energy and Their Interconversion.- 2 The First Law Of Thermodynamics.- 2.1 Statement of the First Law.- 2.1.1 Reversible Expansion of an Ideal Gas.- 2.1.2 Constant Volume Processes.- 2.1.3 Constant Pressure Processes.- 2.2 A New Function-Enthalpy.- 2.2.1 Relationship between ?H and ?U.- 2.3 Uses and Conventions of ?H.- 2.3.1 Enthalpy Change of Reaction.- 2.3.2 Standard Enthalpies of Formation.- 3 Thermochemistry.- 3.1 Calorimetry.- 3.1.1 Bomb Calorimeters.- 3.2 Concepts of Heat Capacity.- 3.2.1 Combustion and Flame Temperatures.- 3.2.2 Variation of Reaction Enthalpies with Temperature.- 3.3 Bond Energies.- 3.3.1 Average Bond Dissociation Energies.- 4 Spontaneous Changes.- 4.1 Everyday Processes.- 4.2 Exothermicity, a Possible Criterion.- 4.2.1 Spontaneous Exothermic Processes.- 4.2.2 Spontaneous Processes Involving no Heat Change.- 4.2.3 Endothermic Processes.- 4.3 The Second Driving Force.- 5 Entropy.- 5.1 Measurement of Entropy.- 5.1.1 The Second Law of Thermodynamics.- 5.1.2 Reversibility and Entropy.- 5.1.3 Changes of Entropy with Temperature.- 5.1.4 An Adiabatic Compression.- 5.2 Absolute Entropies.- 5.2.1 The Third Law of Thermodynamics.- 5.2.2 ?S for Phase Changes.- 5.3 The Direction of Time.- 5.4 A Statistical Approach to Entropy.- 5.4.1 The Boltzmann Entropy Equation.- 5.4.2 Spectroscopic Entropies.- 6 Free Energy, The Arbiter.- 6.1 Processes in Isolated Systems.- 6.2 Gibbs Free Energy, G.- 6.3 Gibbs Free Energy and Maximum Work.- 6.4 Some Processes in Terms of Gibbs Free Energy.- 6.4.1 Adsorption Processes.- 6.4.2 Evaporation Phenomena.- 6.4.3 Endothermic Chemical Processes.- 6.5 Standard Free Energy Changes.- 7 Chemical Equilibrium.- 7.1 Preamble.- 7.2 Variation of G with Gas Pressure.- 7.2.1 Chemical Potential, ?.- 7.2.2 Pressure and Chemical Potential for Ideal Gases.- 7.2.3 Chemical Potential for Real Gases.- 7.2.4 Activity.- 7.3 The Active Mass of Pure Liquids and Solids.- 7.4 Activity of Materials in Solution.- 7.4.1 Solvents.- 7.4.2 Solutes or Minor Components.- 7.5 A Summing Up Activity as a Unifying Concept.- 7.6 Practical Aspects of Activity.- 7.7 Equilibrium and the Reaction Isotherm.- 8 Equilibrium Experiments and Their Interpretation.- 8.1 The Reaction Isochore Equation.- 8.1.1 Le Chatelier up to Date.- 8.2 Applications of the Isochore Equation.- 8.2.1 Vaporization Processes.- 8.2.2 The Decomposition of the Compound UIn3.- 8.2.3 The High Temperature Dissociation of Water Vapour.- 8.3 The Clapeyron Equation.- 9 Electrochemistry.- 9.1 Electrochemical Cells.- 9.2 Cell Energetics.- 9.3 Standard Electrode Potentials.- 9.4 Variation of Cell e.m.f. with Activity.- 9.4.1 Ionic Activities.- 9.4.2 Analysis of e.m.f. Data to Find E0.- 9.5 Variation of e.m.f. with Temperature.- 10 Free Energy And Industrial Processes.- 10.1 Free Energies as a Function of Temperature.- 10.1.1 The Gibbs-Helmholtz Equation.- 10.1.2 The Integrated Form of the Gibbs-Helmholtz Equation.- 10.1.3 Tabulated Forms of Free Energy.- 10.2 The Synthesis of Ethanol.- 10.2.1 Equilibrium Calculations.- 10.2.2 Use of Activity Coefficients.- 10.3 Ellingham Diagrams.- 10.3.1 Corrosion Prevention.- 10.3.2 Electrolysis of Alumina.- 10.3.3 Thermal Reduction of Magnesia.- 10.3.4 Titanium and the Kroll Process.- Suggested Further Reading.- Appendix I: The Twenty Most Useful Equations.- Appendix II: Fundamental Constants and Conversion Factors.- Answers to Problems.

Molecular Freedom of the Ammonium Ion. Heat Capacity and Thermodynamic Properties of Ammonium Perchlorate from 5°–350°K

Abstract: The heat capacity of NH4ClO4 has been determined by adiabatic calorimetry from 5°–350°K and found to be of simple sigmate character without thermal anomalies. The heat capacity (Cp), entropy (S°), enthalpy function (H°−H°0) / T, and Gibbs energy function (G°−G0°) / T evaluated at 298.15°K from these data are 30.61, 44.02, 20.24, and −23.78 cal/(gfm °K). Combination of these values with aqueous NH4ClO4 thermochemical data suggests the absence of zero‐point entropy. Comparison with the heat capacity of isostructural KClO4 permits resolution of the molecular dynamics of the ammonium ions and leads to the conclusion that these ions are restricted rotators, prevented from freely rotating by comparatively low‐energy barriers.

Diffraction Pattern and Structure of the System Tetra‐n‐butylammonium Fluoride–Water at 25°C

Abstract: The diffraction pattern of a concentrated solution of tetra‐n‐butylammonium fluoride in water (Bu4NF·41H2O) has been measured and analyzed at 25°C. The radial distribution function for the solution is surprisingly similar to that of pure water at the same temperature. Each water molecule has ∼ 3.8 neighbors at an average distance of 2.80 A, compared to ∼ 4.4 neighbors at 2.85 A in pure water. Intensity and radial distribution functions have been calculated for two different models. The respective models assume modified ice‐I and gas hydrate structures, surrounded by a random distance distribution, as an adequate description for the average short‐range order in the solution. In both models the F− ions and the N+ atoms of the cations are part of a hydrogen‐bonded network, while the butyl chains of the cations are located inside the cavities which, although of different size and shape, are typical of both structures. The ice‐I model describes the x‐ray data of the solution and of pure water, and its parameters are related to bulk thermodynamic properties in a physically reasonable way. The gas hydrate model is incompatible with the experimental radial distribution function of liquid water, and the parameters of this model for the solution do not explain the large enthalpy and entropy of fusion of the solid hydrate.

Vaporization of Uranium Mononitride and Heat of Sublimation of Uranium

Abstract: The vaporization of uranium mononitride UN has been investigated by the Knudsen effusion technique in combination with a mass spectrometer. The vaporization occurs incongruently by preferential loss of nitrogen forming the two‐phase system nitrogen‐saturated liquid uranium–uranium‐saturated nonstoichiometric uranium mononitride. The vapor pressures of U and N2 have been obtained by silver calibration over the two‐phase system UN0.4–UN0.9 in the temperature range 1910–2230°K. The enthalpies ΔH°298 for the simplified reaction (1) UN(s)=U(g)+0.5 N2 (g) and its partial reactions (2) UN(s)=U(l)+0.5 N2(g) and (3) U(l)=U(g) have been obtained from second‐ and third‐law treatments. The experimental third‐law value of Partial Reaction (3) of 130.7 ± 2.2 kcal mole−1 may be considered an upper value for the heat of vaporization of uranium metal. The third‐law enthalpy of Partial Reaction (2) of 68.9 ± 0.4 kcal mole−1 is in general agreement with literature data for this reaction. The selected value for the enthalpy ...

Particle size effects and thermal hysteresis in crystal structure transformations

Abstract: Particle size affects the temperature and enthalpy of phase transformations as well as the magnitude of thermal hysteresis ΔT, which is shown to be approximately equal to (1/Cpβ)ΔV2/V. The variation of ΔT with particle size (0–150 µm) is interpreted in terms of Turnbull's theory of heterogeneous nucleation. The dependence of the enthalpy of transformation on particle size seems to result from the variation of surface area and surface energy with particle size.

Thermodynamic properties of ideally crystalline polyethylene

Abstract: Results are presented of parallel calorimetric and dilatometric experiments on several meltcrystallized polyethylenes, with various thermal histories, over the range –35°→ liquidus. Enthalpic and weight fraction degrees of crystallinity are in good agreement up to 60°C, where direct comparison with n-alkane data is possible. Hence the interfacial surface enthalpy of the crystallites is small and there is evidence to show such behaviour continues to the melting point. The derived heat of fusion of perfectly crystalline polyethylene exhibits considerable curvature at low temperatures; at the melting point it is 307 J g–1, some 5 % above the highest previous estimate. The temperature dependence of the free energy of fusion indicates an extended chain melting point of 141 ± 1°C.

Ultrasonic Relaxation and the Volume Change during Internal Rotation in 1,1,2‐Trichloroethane

Abstract: Ultrasonic absorption measurements have been carried out in solutions of 1,1,2‐trichloroethane in various solvents. The observed relaxation has been attributed to the perturbation of the equilibrium between the trans and gauche rotational isomers. The concentration dependence of the relaxation parameters have served as an indirect check on the assumptions used in developing the ultrasonic theory. These results were used with those from a recent NMR study to calculate the volume change during isomerization, ΔV / V. This quantity had values in the range 1%–3% showing that the assumption used in deriving the enthalpy difference ΔH from ultrasonic data (ΔV / V)(Cp / θ) ≪ ΔH is not justified. (Cp is the specific heat at constant pressure and θ the coefficient of thermal expansion.) The temperature dependence of the relaxation time was found to be approximately the same as the pure liquid.

Uranium Diboride: Preparation, Enthalpy of Formation at 298.15°K, Heat Capacity from 1° to 350°K, and Some Derived Thermodynamic Properties

Abstract: A sample of uranium diboride was prepared and characterized as UB1.979±0.006 with 0.06 ± 0.03 wt % of identified impurities. The standard enthalpy of combustion in fluorine was determined to be − 1021.2 ± 1.1 kcal mole−1. The heat capacity was measured from 0.84° to 350°K. At 298.15°K the heat capacity CP°, entropy S°, and enthalpy increment H° − H°0 are 13.23 ± 0.03 cal K−1·mole−1, 13.17 ± 0.03 cal °K−1·mole−, and 2108 ± 4 cal mole−1, respectively. The following values were obtained for the standard enthalpy, entropy, and Gibbs energy of formation of UB2 at 298.15°K: ΔHf° = − 39.3 ± 4.0 kcal mole−1, ΔSf° = − 1.54 ± 0.05 cal °K−1·mole−, and ΔGf° = − 38.8 ± 4.0 kcal mole−1. These agree within experimental error with values calculated from high‐temperature effusion measurements. The heat‐capacity results below 4.2°K follow the equation CP = (9.40 ± 0.01)T + (3.18 ± 0.14) × 10−2T3mJ °K−1·mole−1. The relatively high value for the coefficient of the linear term indicates that uranium diboride is a good electri...