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1 Dec 1947

Volume 15, Issue 12, pp. 849-886


On the Thermodynamics of Non‐Electrolyte Solutions and Its Technical Applications III. Systems with Associated Components

O. Redlich and A. T. Kister

J. Chem. Phys. 15, 849 (1947); http://dx.doi.org/10.1063/1.1746359 (7 pages) | Cited 18 times

Online Publication Date: 22 November 2004

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The influence of continuous association of one component on the free energy of a non‐electrolyte solution is derived from a few assumptions. The result is in accord with available data for hydrocarbon‐alcohol solutions. Systems of two associating components and the general interpretation of the free energy of non‐electrolyte solutions are briefly discussed.

Foreign Ion Rejection in the Growth of Sodium Chloride Single Crystals from the Melt

Raymond H. McFee

J. Chem. Phys. 15, 856 (1947); http://dx.doi.org/10.1063/1.1746360 (6 pages) | Cited 20 times

Online Publication Date: 22 November 2004

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A study is made of the process of purification that takes place when single crystals are grown from the melt. The rejection of potassium and copper ions from solid solution in sodium chloride crystals is measured quantitatively by analyzing by spectrochemical means sections of cylindrical crystals of sodium chloride to which has been added a known quantity of impurity. The theory of the distribution of impurity over the height of the crystal is derived in terms of the initial impurity concentration, height of completed crystal, and the fraction of the impurity concentration in the melt included in the crystal, which fraction is called the purification coefficient. Good agreement is obtained between theory and experiment for an initial concentration of 0.1 atomic percent potassium in sodium chloride. Purification coefficients for several crystals containing potassium and a few crystals containing copper are measured and correlated with rate of growth and temperature gradient conditions.

Theory of Burning Velocity. II. The Square Root Law for Burning Velocity

Charles Tanford and Robert N. Pease

J. Chem. Phys. 15, 861 (1947); http://dx.doi.org/10.1063/1.1746362 (5 pages) | Cited 15 times

Online Publication Date: 22 November 2004

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Previous papers by the authors having suggested that the diffusion of active particles (chiefly hydrogen atoms) from the flame front is the controlling factor in combustion in Bunsen‐type burners, an equation for burning velocity based on such a concept is derived. This equation shows that the burning velocity should be proportional to the square root of an expression in which the most important factor is the sum of the products piDi for all effective atoms or free radicals, pi being the concentration of an atom or radical at the flame front, and Di its coefficient of diffusion into cold, unburnt gas. The burning velocity equation is applied to the combustion of carbon monoxide and hydrogen, and excellent agreement between calculated and experimental values of the burning velocity is obtained. The equation is also used to account for the effect of pressure upon burning velocity.

Heat Capacities of Liquids and Vapors

Sidney W. Benson

J. Chem. Phys. 15, 866 (1947); http://dx.doi.org/10.1063/1.1746364 (2 pages) | Cited 1 time

Online Publication Date: 22 November 2004

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By the use of an equation previously derived for the energy of vaporization of a liquid: ΔE=k(Dln—Dgn) (in which k is a constant; n=5/3, and Dl and Dg are the molar densities of liquid and vapor, respectively), an expression is derived relating (∂ΔE/∂T)sat to the difference in molar heat capacities at constant pressure of the liquid and vapor (ΔCp). At low pressures this relation can be approximated by the equation ΔCp=(∂ΔE/∂T)s+ R (R is the gas constant, 1.986 cal./mole‐°A). Under these conditions (∂ΔE/∂T)s = (5/3)ΔE(∂ lnDl/∂T)s, and the resulting equation for ΔCp is ΔCp = (5/3)ΔE(∂ lnDl/∂T)s+1.99. From known properties of liquids ΔCp can be calculated using this last equation. These calculated values are found to be in good agreement with observed experimental values.

The Experimental Determination of the Intensities of Infra‐Red Absorption Bands. III. Carbon Dioxide, Methane, and Ethane

A. M. Thorndike

J. Chem. Phys. 15, 868 (1947); http://dx.doi.org/10.1063/1.1746366 (7 pages) | Cited 62 times

Online Publication Date: 22 November 2004

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The absolute intensities of the chief infra‐red absorption bands of carbon dioxide, methane, and ethane have been determined experimentally to be (in cycles per centimeter at N.T.P.):
math
These results are in satisfactory agreement with available data on infra‐red dispersion and atomic polarization. They may be interpreted in terms of dipole moments, μ′, of bonds between vibrating atoms, and their rates of change with internuclear distance, ∂μ′/∂r. When this is done for the CH bond, however, unexpected variations in these quantities are found for the different vibrations and molecules studied. The average values are about μ′=0.4×10−18, ∂μ′/∂r=±0.6×10−10. For the CO bond the values are considerably larger.

A Note on the Relation between Entropy and Enthalpy of Solution

O. K. Rice

J. Chem. Phys. 15, 875 (1947); http://dx.doi.org/10.1063/1.1746367 (5 pages) | Cited 2 times

Online Publication Date: 22 November 2004

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The relation between entropy and enthalpy of solution for a series of non‐polar solutes in a given non‐polar solvent is discussed. It is considered that solution of a gaseous solute, without changing its concentration on going from gas phase to liquid solution phase, does not change its own entropy, all change of entropy being referred to the solvent. The entropy of the solvent changes because of surface effects around the solute molecules and a kind of long‐range order introduced by the solute molecules. Two extreme cases are considered, (1) the case of an ideal solution, and (2) the case of a solute of hard attractionless spheres. The difference in entropy of solution between these extreme cases can be estimated. It can also be estimated by extrapolation from the experimental data on the entropy and enthalpy of solution, and these two estimates agree in order of magnitude. The fact that the relation between entropy and enthalpy of solution is linear is also shown to be a reasonable expectation, and effect of changing solvent as well as solute is considered. The groundwork is thus laid for a qualitative understanding of this relation between entropy and enthalpy of solution.
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Halogen‐Halogen Distances in Halogen‐Polymer Complexes

R. E. Rundle

J. Chem. Phys. 15, 880 (1947); http://dx.doi.org/10.1063/1.1746369 (1 page) | Cited 8 times

Online Publication Date: 22 November 2004

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Abstract Unavailable

Infra‐Red Dichroism Involving Hydrogen Bonds

Leonard Glatt and Joseph W. Ellis

J. Chem. Phys. 15, 880 (1947); http://dx.doi.org/10.1063/1.1746371 (2 pages) | Cited 5 times

Online Publication Date: 22 November 2004

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Abstract Unavailable

A Spectroscopic Evidence for Activation of Fluorescence by High Valency Manganese Ions

G. Szigeti, E. Nagy, and E. Makai

J. Chem. Phys. 15, 881 (1947); http://dx.doi.org/10.1063/1.1746373 (2 pages) | Cited 2 times

Online Publication Date: 22 November 2004

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Abstract Unavailable

Raman Effect of Dichlorohexafluorocyclobutane

Walter F. Edgell and Francis E. Kite

J. Chem. Phys. 15, 882 (1947); http://dx.doi.org/10.1063/1.1746412 (1 page) | Cited 10 times

Online Publication Date: 22 November 2004

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Abstract Unavailable

Phases of Fourier Coefficients Directly from Crystal Diffraction Data

D. Harker and J. S. Kasper

J. Chem. Phys. 15, 882 (1947); http://dx.doi.org/10.1063/1.1746413 (1 page) | Cited 2 times

Online Publication Date: 22 November 2004

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Infra‐Red Dichroism in Aligned Polythene and ``Parowax''

Leonard Glatt and Joseph W. Ellis

J. Chem. Phys. 15, 884 (1947); http://dx.doi.org/10.1063/1.1746415 (1 page) | Cited 8 times

Online Publication Date: 22 November 2004

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Abstract Unavailable

On the Frequency Factor in the Viscosity‐Temperature Relationship of Liquids

M. S. Telang

J. Chem. Phys. 15, 885 (1947); http://dx.doi.org/10.1063/1.1746417 (1 page) | Cited 5 times

Online Publication Date: 22 November 2004

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Prism Spectrometry from 24 to 37 Microns

Earle K. Plyler

J. Chem. Phys. 15, 885 (1947); http://dx.doi.org/10.1063/1.1746419 (2 pages) | Cited 2 times

Online Publication Date: 22 November 2004

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Abstract Unavailable
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