NOTES
Oct., 1956
so no appreciable error was involved in this procedure. The process of solution of a single 0.8264-g. sample of nitramide in 500 ml. of 0.001 N sulfuric acid a t 24’ was found to be endothermic by 3700 ca1.l mole. For calculation of the heat of formation of nitramide a t 24” from the elements, the following values were used. Enthalpy for the process of solution of nitrous oxide in water -5303 cal./mole (AC, between 0 and 40” +35.9 cal./OC.) as derived from the solubility data of Markham and Kobe8; heat of formation of nitrous oxide 19490 cal./mole and of liquid water -68310 as given by Rossini, et aL6 The value for the heat of decomposition of nitramide (c) to form nitrous oxide gas and liquid water was calculated to be exothermic by 28200 f 85 cal./mole. The uncertainty in this value includes that of sample purity, and that of calorimetric measurement. The heat of formation of nitramide (e) from the elements was calculated to be 20630 i 85 cal./mole (exothermic) at 24”, that for aqueous nitramide, - 16930 cal./mole. Comparison of this latter value with the heat of formation of hyponitrous acid in aqueous solution from the element,s as determined by L a t h e r and Z i m r n e r m a ~ and ~ , ~ recalculated by Rossini, et u Z . , ~ - 13700 cal./mole shows that aqueous nitramide is more stable by 3230 cal./mole than hyponitrous acid.
+
( 8 ) A. E. Markham and K. A. Kobe, J. A m . Chem. SOC.,63, 449 (1941).
THE VAPOR PRESSURE OF HC1 ABOVE NON-AQUEOUS SOLUTIONS BY J. J. FRITZ Department of chemistry, The Pennsylvania State University, Univerm’ly Park, P a . Received M a y 7 , I065
The writer and Fugetl have recently published tables of the vapor pressures of HC1 above aqueous solutions, as calculated from available cell measurements. The same methods are applicable to non-aqueous solutions. The purpose of this communication is to make available data for facilitating the calculation of vapor pressures of HC1 above non-aqueous media. In general, the vapor pressure of HCl is given by F InfHol = - (Ego - E ) RT
(1)
where E is the voltage of the cell Hz, Pt/HCl (solvent)/AgCl, Ag, (corrected to 1 atm. pressure of hydrogen); Egois the standard potential of the cell for unit fugacity of HCl. Values of Egofrom 0 t o 40’ are given by Aston and Gittler,2 and can be calculated readily for other temperatures. Some typical vapor pressures of HC1 a t 25”, so calculated, are: in water (1 m), 2.2 X 10-4 mm. Hg; 50 mole % ’ ethanol-water (1 m),1.2 X mm. Hg; ethanol (1 m ) , 0.46 mm. Hg; methanol (0.56 m),8 X mm. Hg. (1) J. J. Fritz and C. R. Fuget, Ind. Eng. Chem., to be published. (2) J. G. Aston and F. L. Gittler, J . Am. Chem. SOC., 77, 3173 (1955).
1461
By use of the standard potential for hypothetical 1 molal solution (EO)in equation 1, the standard fugacity of HC1 is obtained for each system. Table I gives some representative values of the standard fugacity a t 25”. FUQACITY OF HCl
IN
Solvent
TABLEI HYPOTHETICAL 1 MOLALSTANDARD STATEAT 25’
f” HCI, atm.
Source of f” HCl, mm. Hg cell data
4.92 x 10-7 3.74 x 10-4 Water 4.19 X 3.18 Methanol 5.06 X 38.5 Ethanol Dioxane-water“ N2 = 0.0187 1.045 X 10-6 7.93 X Nz = 0.1433 4.86 X 10-8 3.69 X l o - * N Z 0,3231 2.34 X 0.178 Na = 0.4823 1.41 X 10.7 0 Nz is the mole fraction of dioxane in the solvent. =i
3 4 4
5 5 5 5
The vapor pressure of HC1 is then given by fi
= a&
= may2&fi”
(2)
where 7~ is the mean in activity coefficient as ordinarily tabulated. The fugacity is a measure of the absolute activity of HC1 in the solution, as contrasted to the relative activity usually tabulated. As demonstrated in our analysis of the aqueous system,’ the vapor pressures obtained from cell measurements are frequently more accurate than the results of direct measurements. Where measurements are available to sufficient dilution, the vapor pressure of the solvent can, of course, be calculated by means of the Duhem equation. (3) H.8.Harned and R. W. Ehlers, ibid., 64, 1350 (1932). (4) H.8. Harned and B. B. Owen, “Physical Chemistry of Electrolytio Solutions,” Reinhold Publ. Corp., New York, N. Y.,1950, p. 336. (5) H. S. Harned, J. 0. Morrison, F. Walker, J. G. Donelson and C. Calman, J . A m . Chem. SOC.,61,49 (1939).
NUCLEATION OF COPPER METAL FROM AQUEOUS SOLUTION BY WELBYG. COURTNEY~ Chemical Construction Corporation, 6.86 West 43 Street, New York, N . Y. Received M a y 7 , 1066
Theories of nueleation can be divided into the classical “supersaturation” viewpoint2 that the size of the critical nucleus varies significantly with temperature and supersaturation (i.e., reagent activities), and the less widely accepted “constant number” viewpoint3 that the critical nucleus is essentially independent of temperature and activities. I n view of the vanishingly small solubility of metals in water, supersaturation would appear to be an uncertain concept when applied to the nucleation of metal particles by reduction of the metal ion from aqueous solution. However, the results of the following exploratory study of the nucleation of Cu metal from aqueous solution can (1) Experiment Incorporated, Richmond, Virginia. (2) M. Volmer,“Kinetik der Phasenbildung,”T. Steinkopff, Dresden
1939. (3) J. A. Christiansen and A. Nielssn, Acta Chem. Scand., 6 , 103 (1949): F. R. Duke, R. J. Bever and H. Diehl, Iowa Slate. College J. Sci., 83, 297 (1949).
