R. CARUSO, F. J. LOPREST,AND A. LUM
1716
Heat of Formation of Hydrazinium Diperchlorate
by R. Caruso, F. J. Loprest, and A. Lum Thiokol Chemical Corporation, Reaction Motors Division, Denville, New Jersey
(Received December $3, 1964)
A value of -70.1 f 1.0 kcal./mole has been obtained for the heat of formation (298.15”K.) of solid hydrazinium diperchlorate (NzH4.2HC104) by solution calorimetry.
Introduction Samples of pure hydrazinium diperchlorate were recently prepared a t this laboratory. I n view of the current interest in perchlorates, it was desirable to determine the heat of formation of this compound. The methed employed was similar to that used recently for determining the heat of formation of nitronium perchlorate. 1 It involves a measurement of the heat of reaction of the solid with aqueous sodium hydroxide. In addition, it was necessary to determine the heat of solution of hydrazine in aqueous sodium perchlorate.
Experimental Materials. The hydrazinium perchlorate was prepared as a white crystalline powder. Analysis indicated that the only impurities present were HC104, N2H4.HC1O4,and HzO. The purity of each sample is given in Table 11. The hydrazine employed in the heat of solution measurements was an anhydrous grade obtained from Olin Mathieson Chemical Corp. Its purity was greater than 99.5% and it was used as received. Apparatus and Procedure. The calorimeter consisted of a 400-cc. silvered dewar vessel with a watertight silicone rubber cap. Three Pyrex tubes inserted through the cap permitted a glass stirrer, a glass rod, and a thermometer to extend into the solution contained in the calorimeter. The entire assembly was placed in a water bath thermostatically controlled at 25”. The thermometer employed for measurement of the calorimeter temperature was a Parr calorimeter thermometer which could easily be read to f0.003”. A thin-walled glass bulb containing the solid or liquid was attached to the submerged end of the glass rod. When thermal equilibrium had been attained, the The Journal of Physical Chemistry
bulb was broken by crushing it against the bottom of the dewar. Calibration. The calorimeter constant was determined chemically from the heat of neutralization of sodium hydroxide with hydrochloric acid. A carbonate-free aqueous solution of NaOH was prepared and its concentration was determined by titration with a standardized solution of HC1. Samples were mixed in the calorimeter with 300 cc. of a 0.121 m HC1 solution. An excess of HC1 was present in all runs. During a measurement, temperature readings were taken from 10 min. before to 10 min. after the reaction period. A method applicable to rapid reactions was employed for obtaining the corrected temperature increment in all cases.2 The results of the calibration are presented in Table I. Using appropriate thermodynamic datal3a value of -14.041 kcal./mole was calculated for the AH of the calorimeter reaction. A value of 333.9 cal./deg. was obtained for the calorimeter constant. The calorimeter constant for each measurement was obtained from the calibration value by correcting for the slight difference in heat capacity of the solution and the calibration solution. All heats of reaction refer to a standard temperature of 298.15”K. The calorie unit used is the thermochemical ~ a l o r i e . All ~ atomic weights were taken from the 1957 “International Table of Atomic Weights.”
Results and Discussion The calorimeter reaction involved the mixing of 1g. samples of hydrazinium diperchlorate with 300 cc. (1) A. A. Gilliand, J . Res. Nail. Bur. Std., 66A, 447 (1962). (2) J. M. Sturtevant, “Physical Methods of Organic Chemistry,” Vol. I, Part I, A. Weissberger, Ed., Interscience Publishers, Inc., New York, N. Y., 1949, Chapter 14. (3) E’. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine, and I. J d e , “Selected Values of Chemical Thermodynamic Properties,” National Bureau Standards Circular 500,U. S. Government Printing Office, Washington, D. C., 1952.
1717
HEATOF FORMATION OF HYDRAZINIUM DIPERCHLORATE
Table 11: Data for Heat of Reaction of NzH4.2HC104 and Aqueous NaOH
Table I : Calibration of the Calorimeter W t . of NaOH soln., g.
3.6390 2.6543 3.7639 3.1356
Concn. of NsOH s o h , mole/g.
0.006360 0.006360 0.006360 0.006360
AT, 'C.
Calorimeter constant, cal./deg.
0.970 0.711 1.017 0.837
335.0 333.4 332.7 334.5
Mean = 333.9 Std. dev. = 0.5
of 0.0802 N NaOH solution. The experimental values of the heat of reaction were corrected to account for perchloric acid and hydrazinium monoperchlorate impurities. Corrections for these impurities were made by assuming they react with aqueous NaOH in accordance with the following equations. HC1O4(1)
+ NaOH(aq) +NaC104(aq) + H20(1) AH = -33.8 kcal./mole
+
N2H4.HC104(c) NaOH(aq) + X2H4(aq)
+ NaC104(aq) + H2O(l) AH = 5.3 kcal./mole
Mass of sample, g.
NzH4.2HC104, %
1.3550 1.1459 0.7776 1.0654 1.0249 0.6522
99.58 99.52 99.79 99.79 99.64 99.64
HClO4, NzHcHCIO4, % %
...
0.04
0.22
*..
... ...
0.15 0.15 0.15 0.13
... ...
AHi, A T , "C.
kcal.
0.218 0.172 0.125 0.162 0.158
-12.3 -11.5 -12.2 -11.5 -11.7 -11.6
0.100
Mean = -11.8 Std. dev. = f 0 . 2
In order to calculate the heat of formation by use of eq. 1, it was also necessary to determine the heat of solution of hydrazine in aqueous NaC104. The heat of solution of N2H4 was measured by mixing a sample in 300 ml. of an aqueous solution containing NaC104 in about the same concentration as it was present in the final state of the heat of solution measurements for NzH4.2HC104. The measured heat of solution was corrected to the AH of reaction 2.
+ (2NaC104+ 6000H20)(soh.) + (NzH4+ 2NaC104 + 6000H20)(soln.)
NzH4(1)
A value of -43 kcal./mole for the AHr of N2H4. HC104 was estimated from data in the literature for the heat of solution of NzH4.~ ~ 1 and 0 from ~ 4 values for the heat of formation of N2H6+(aq)and C104-(aq) avaihble in tables of thermochemical data.3 The experimental heat of reaction was corrected to the heat of the reaction shown in eq. 1 by standard methods. Thermochemical data required in the cal-
The results are tabulated in Table 111. A value of -4.3 kcal./mole Was found for the AH of reaction 2. This value is comparable to the literature value of -3.89 kcal. for the heat of solution of N2H4 in 300 molesofH~0.~
+ 6000H20) (l\SZH4 + 2NaC104 + 6000H20)(soh.) + 2H20(1)
Table IJI : Data for Heat of Solution of N2H4in
N2H4*2HC104(c) f (2NaoH
-
(1)
culations were taken from NBS table^.^ Since the final state is basic, some consideration was given to the enthalpy change due to a slight shift in the equilibrium N2Hj+
Dilute NaC101 Masa of sample, g.
NZH4, %
AT, "C.
AHe, kcal,
0.2214 0.1653 0.5567
99.5 99.5 99.5
0.093 0.068 0.220
-4.4 -4.4 -4.2
+ OH- If NzHsOH
and to the possible decomposition or oxidation of the hydrazine solution. Using a value of 8.5 X lo-' for the equilibrium constant15the heat effect amounted to 0.1 cal. in all runs. From kinetic data available in the literature, it was estimated that the decomposition of the hydrazine solution during the run time would be negligible.6 The results of t,hese measurements are presented in Table 11. A d u e Of -11.8 kcal*/mole was found for the heat of reaction 1.
(2)
Mean = - 4 . 3 Std. dev. = f O . l
B~ substituting the following thermodynamic quanin the equation that results from combining (lland (2)
tities3
(4) E. C. Gilbert and A. W. Cobb, J. Am. Chem. SOC., 57,39 (1935). (5) L.F.Audrieth and B. A. Ackerson, "The Chemistry of Hydrasine," John Wiley and Sons, Inc., New York, N.Y., 1951. (6) E. C. Gilbert, J. Am. Chem. Soc., 51, 2744 (1929).
Volume 60,Number 6 M a y 1966
C. W. GARLAND, S. TONG, AND W. H. STOCKMAYER
1718
A.Hf(NaOH.3000H20) = -112.170
the heat of formation of N2H4*2HC1O4is computed to be -70.1 kcal./mole. Taking into consideration errors in temperature measurement, calibration, and weighing, an estimated absolute error of * l . O kcal. is placed on the reported value.
AHf(H20(1)) = -68.317 AHf(NaC104.3000H20)= -88.38 AHf(NzHd(1)) = 12.05
Diffusion in the System Cadmium Iodidewater
by C. W. Garland, S. Tong, Department of Chemistry, Massachuaetts Institute of Technology, Cambridge, Massachusetts
02139
and W. H. Stockmayer Department of Chemistry, Dartmouth College, Hanover, New Hampshire 03766 (Received December WY, 1964)
Mutual diffusion coefficients at 25’ for aqueous solutions of cadmium iodide have been measured at concentrations from 0.01 to about 1 M by the porous-frit technique. The apparent mobility of the salt ‘is a rapidly varying function of concentration, from which estimates of the diffusivities and stabilities of the species CdIf and CdI2 have been made, in satisfactory agreement with other existing data. Diffusion coefficients for aqueous solutions of mercuric chloride are also reported, leading to a limiting value of D = 10.35 X sec.-l for the HgClz molecule in water.
Introduction The analysis of isothermal diffusion in weak-electrolyte systems presents several challenging problems. I n any two-component system, the experimental mutual diffusion coefficient is given by
D
=
(Q/c)RTv(l
+ b In y+/b In c)
(1)
where v is the number of moles of ions per formula weight of completely dissociated electrolyte and the thermodynamic factor involves the stoichiometric mean ionic activity coefficient y, on the volume-formal concentration scale c. 1,2 However, diffusion in weakelectrolyte solutions differs from that in strong-electrolyte solutions for two reasons: the factor (1 b In y*/b In c) must change with concentration in a different manner than for completely ionized solutes, and the mobility factor W/c can be expected to change with the
+
The Journal of Physical Chemistry
degree of dissociation since the various solute species will have different mobilities. A general relation between the observable quantity Q and the mobility matrix elements Ois has been givena for weak electrolytes of arbitrary valence type with any number of dissociation or polymerization equilibria consistent with a single thermodynamic component in addition to the solvent. This relation reduces to special formulas which were obtained previously for ~ni-univalent~ and uni-bivalent5 weak electrolytes and (1) H. 5. Harned and B. B. Owen, “Physical Chemistry of Electrolytic Solutions,” 2nd Ed., Reinhold Publishing Corp., New York, N. Y.,1950,pp. 86-90. (2) L. Onsager and R. M. Fuoss, J . Phys. Chem., 36, 2689 (1932). (3) W.H. Stockmayer, J . Chem. Phys., 33, 1291 (1960). (4) H. S. Harned and R. M. Hudson, J . Am. Chem. SOC.,7 3 , 3781, 5880 (1951). (5) J. M.Creeth and R. H. Stokes, J . Phys. Chem., 64, 946 (1960).
