Solvent structure effects on electron thermalization ranges in water–alcohol mixed solvents; preferred EAmax values for e−s in alcohols

1983 ◽  
Vol 61 (6) ◽  
pp. 1115-1119 ◽  
Author(s):  
Ah-Dong Leu ◽  
Kamal N. Jha ◽  
Gordon R. Freeman
Keyword(s):  
Energy Loss ◽  
Mixed Solvents ◽  
Solvent Structure ◽  
Solvated Electrons ◽  
Pure Solvent ◽  
Free Ion ◽  
Rotational Motions ◽  
Stiffening Effect ◽  
Water Alcohol

Addition of 2 mol% of an alcohol to water or of 1 mol% of water to an alcohol increases the value of Gfi εmax (e−s) above that in the pure solvent. The effect is attributed mainly to an increasing free ion yield Gfi. The increase of Gfi correlates with a "stiffening" of the solvent molecular rotational motions and a corresponding lower rate of energy loss by the epithermal electrons to dipole rotations. Thus the increase of Gfi is mainly attributed to a larger electron thermalization range bGP. The value of Gfi εmax passes through a maximum at 2 mol% of an alcohol in water; the maximum increases in the order methanol< ethanol< 1-propanol as additive. The solvent "stiffening" effect of these solutes increases in the same order.Some of the published values of EAmax and W1/2 for solvated electrons in alcohols appear to be too low. A set of preferred values is reported herein.

Effects of solvent structure on electron reactivity and radiolysis yields: 2-propanol/water mixed solvents

1984 ◽  
Vol 62 (7) ◽  
pp. 1265-1270 ◽  
Author(s):  
Joanna Cygler ◽  
Gordon R. Freeman
Keyword(s):  
Pure Water ◽  
Mixed Solvents ◽  
Water Structure ◽  
Trap Depth ◽  
Solvent Structure ◽  
Solvated Electrons ◽  
Absorption Energy ◽  
Diffusion Controlled ◽  
Free Ion ◽  
Different Temperatures

Reaction of solvated electrons with nitrobenzene, N, is nearly diffusion controlled in both pure solvents; kN ~ 1010 dm3/mol s. The value of kN is approximately proportional to the inverse viscosity η−1 in the pure solvents, and in the mixed solvents at different temperatures. However, on going from zero to 74 mol% water at the same temperature kN is independent of the 40% increase of η. Electron diffusion in the mixed solvents is not a simple function of fluidity.Reaction with the inefficient scavengers tryptophane (kS ~ 109 dm3/mol s) and phenol (kS ~ 107–108 dm3/mol s) correlates inversely with the electron optical absorption energy. The latter is related to the trap depth in the solvent; electrons in deeper traps have less tendency to react with molecules of low electron affinity.Addition of 3 mol% 2-PrOH to water at 296 K increases the value of Gεmax by 16%, although the value in pure 2-PrOH is three-fold smaller than that in pure water. The increase is attributed to an increase in the free ion yield, caused by an increase in the product of the electron thermalization range and the microscopic dielectric constant of the fluid between the ion and electron, averaged over the time that they exist as a correlated pair. Addition of a small amount of alcohol to water increases the orderliness of the water structure.

scholarly journals The behaviour of electrolytes in mixed solvents. Part I.—The free energies and heat contents of hydrogen chloride in water—Ethyl alcohol solutions

1929 ◽  
Vol 125 (799) ◽  
pp. 694-712 ◽  
Keyword(s):  
Hydrogen Chloride ◽  
Electric Fields ◽  
Mixed Solvents ◽  
Dielectric Constants ◽  
Free Energies ◽  
Dissolved Ions ◽  
Uniform Medium ◽  
The Mean ◽  
Solvent Molecules ◽  
Water Alcohol

Previous studies on the effect of a change of medium on tire properties of dissolved electrolytes have aimed at correlating the behaviour of the electrolyte with the mean physical properties, e. g ., dielectric constants, of the medium. While this approach may be justified in the case of solvents containing molecules of only one kind, it is not sufficient to regard a mixed solvent as a uniform medium affecting the dissolved ions solely through the effect of its dielectric constant on the electric forces between them. For the electric fields of ions exert a differential attraction on molecules possessing different degrees of polarisability and since tire more polarisable molecules must tend to congregate round the ions, the properties of the latter cannot depend solely on tire mean properties of tire medium. Studies on the behaviour of ions in such cases will throw light on the interaction between ions and solvent molecules. The present paper gives tire results of measurements of the free energies and heat contents of hydrogen chloride in water-alcohol solutions, obtained by determining the electromotive forces of cells of the type:— H 2 ( g ) | HCl ( m ), AgCl ( s ) | Ag water-alcohol

scholarly journals The Detemination of the Formation Constants of the Triiodide Ion in Water–Alcohol Mixed Solvents at Various Temperatures

1977 ◽  
Vol 50 (3) ◽  
pp. 566-569 ◽  
Author(s):  
Katumitu Hayakawa ◽  
Sumio Nakamura
Keyword(s):  
Mixed Solvents ◽  
Formation Constants ◽  
Water Alcohol

scholarly journals The behaviour of electrolytes in mixed solvents. Part. III.–The molecular refractivities and partial molar volumes of lithium chloride in water-ethyl alcohol solutions

1931 ◽  
Vol 131 (817) ◽  
pp. 382-390 ◽  
Keyword(s):  
Lithium Chloride ◽  
Ethyl Alcohol ◽  
Mixed Solvent ◽  
Mixed Solvents ◽  
Partial Molar Volumes ◽  
Ammonium Iodide ◽  
Dilute Aqueous Solutions ◽  
The Difference ◽  
Value Of It ◽  
Water Alcohol

It is well known that the molecular refractivity of most salts, as calculated by the Lorentz-Lorenz formula, is nearly independent of the concentration in moderately dilute aqueous solutions. Walden determined the refractivities of tetra-ethyl-ammonium iodide and other similar salts in a variety of solvents and found that, while the molecular refractivity was approximately independent of the concentration in each solvent, it varied from one solvent to another, the greatest variation from the value in water, amounting to about 2 per cent., being obtained in nitro-methane, Schreiner has recently determined the molecular refractivity of hydrogen chloride and lithium chloride in methyl and ethyl alcohols, and found that in the case of lithium chloride the value is independent of the concentration up to a concentration of about 3 M. His values for R 18 D for lithium chloride are: 8.73 in water, 8.55 in methyl alcohol and 8.38 in ethyl alcohol. The difference between the values in water and ethyl alcohol appeared to make it just possible to determine the variation of the refractivity with the composition of the solvent in mixtures of water and the alcohol. It is possible that a solvent might he found, miscible in water in all proportions, in which the value of It is further removed from that in water. Such a substance would he more suitable than alcohol for the investigation of this effect, but in order to correlate the results with the measurements recorded in Part II of the activities of alcohol and water in water-alcohol-lithium chloride solutions, it seemed desirable to investigate this case in the first instance. The variation of the refractivity of a salt with the composition of a mixed solvent may be expected to give some indication of the composition of the solvent in the immediate vicinity of the ions. For the refractivity of a salt is determined by (1) the polarisability of the ions themselves and (2) the change in the polarisability of the solvent produced by their presence. The molecular refractivity of a salt in a solution containing m grams of a salt in w grams of the solvent is taken as R = M/ m ( n 2 -1/ n 2 + 2 . w + m / d - n 2 0 -1/ n 2 0 + 2 . w / d 0 (1) where n , and d and n 0 , d 0 are the refractive index and density of the solution and of the solvent, respectively, and M the molecular weight of the salt.

CLVIII.—The solubility of picric acid in mixed solvents. Part I. Water–alcohol and water–acetone mixtures

1931 ◽  
Vol 0 (0) ◽  
pp. 1196-1201 ◽  
Author(s):  
James Cooper Duff ◽  
Edwin John Bills
Keyword(s):  
Picric Acid ◽  
Mixed Solvents ◽  
Water Alcohol

scholarly journals The role of solvent structure in the absorption spectrum of solvated electrons: Mixed quantum/classical simulations in tetrahydrofuran

2005 ◽  
Vol 122 (13) ◽  
pp. 134506 ◽  
Author(s):  
Michael J. Bedard-Hearn ◽  
Ross E. Larsen ◽  
Benjamin J. Schwartz
Keyword(s):  
Absorption Spectrum ◽  
Solvent Structure ◽  
Solvated Electrons ◽  
Role Of Solvent

Erratum: Solvent structure effects on solvated electron reactions in mixed solvents: Negative ions in 1-propanol–water and 2-propanol–water

1992 ◽  
Vol 70 (2-3) ◽  
pp. 192-192
Author(s):  
Annesley Peiris Sedigallage ◽  
Gordon R. Freeman
Keyword(s):  
Mixed Solvents ◽  
Negative Ions ◽  
Solvent Structure ◽  
Solvated Electron

scholarly journals The viscosity of dilute solutions of strong electrolytes

1931 ◽  
Vol 134 (824) ◽  
pp. 413-427 ◽  
Keyword(s):  
Empirical Equation ◽  
Relative Viscosity ◽  
Theoretical Treatment ◽  
High Dilution ◽  
Equivalent Concentration ◽  
Pure Solvent ◽  
Wide Range ◽  
Strong Electrolytes ◽  
Negative Viscosity ◽  
Stiffening Effect

Until very recently, no empirical equation had been found to represent satisfactorily the variation with concentration of the relative viscosity of electrolytes, nor had any adequate theoretical treatment of the problem been put forward. In 1929, however, Jones and Dole showed that the fluidity (or reciprocal of the relative viscosity) of a salt solution could be represented over a fairly wide range of concentration by an equation of the form ϕ = 1 + A√ c +B c , where ϕ is the fluidity, c the equivalent concentration, and A and B are empirical constants. The value of B is negative in the case of salts which increase the viscosity of water, and positive in cases of so-called “negative viscosity,” where the viscosity of the solution is less than that of the pure solvent. Jones and Dole argued further that the stiffening effect of the interionic forces would tend to make the constant A, which determines the viscosity at high dilution, always negative. A little later, Falkenhagen and Dole treated the problem theoretically from the standpoint of the ion-atmosphere theory of Debye and Huckel. They confirmed the suggestion that at high dilution the electrolyte must always increase the viscosity of the solvent, and showed that the relative viscosity of an electrolyte solution at high dilution must be represented by an equation of the form η μ / η 0 = 1 + K √ μ , where η μ is the viscosity of the solution, η 0 is the viscosity of the solvent, μ is the equivalent concentration, K is a constant.

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