5808
JACOB SEDI,ET ARD
Agreement was very satisfactory a t d values of 1.50 or less. TABLE I1 X-RAYDATA. ALUMINUM( 111) IODIDE Sample I
d
I/Il
(est.)
d
Sample I1 I/Il (est.)
3.35 2.96
1.0 0.5
3.43 2.98
1.0 0.4
2.10 1.80 1.77 1.73
.7 .9 .2
2.13 1.82 1.80 1.74
.8
.2
.8 .2 .2
Literature's
d
I/I
3.70 3.50 3.08 2.98 2.14 1.84 1.80 1.75
0.05 1.oo 0.50 .05 .50
[CONTRIBUTION FROM
THE
1
.38 .10 .13
THOMAS
VRIES
Vol. 7 3
The differences in the highest d-values for the two samples may be due to a very slight hydrolysis, which had occurred during the time the capillaries were being loaded. This hydrolysis was evidenced by the development of a very faint yellow tinge, instead of the original pure white powder. The material giving the lower d-values had the fainter color, and would be expected to be the purer. Since Hanawalt, Rinn and Frevel do not give their method of preparation of aluminum(II1) iodide, nor the purity of the material used in their work, no conclusion from the comparison of results can be drawn. CAMBRIDGE, MASSACHUSETTS RECEIVED MAY31, 1951
DEPARTMENT OF CHEMISTRY, PURDUE UNIVERSITY]
Electrode Potentials in Liquid Ammonia. I. Activities of Potassium and Sodium Chlorides BY JACOB SEDLET~ AND THOMAS DE VRIES The activity coefficients of potassium and sodium chloride in liquid ammonia at -36 O were determined from electromotive force measurements of Fells of the type: K(Hg), KCl(m), ZnC1~6NHs(s),Zn(Hg) and Na(Hg), NaCl(m), ZnC12.6NHs(s), Zn(Hg). The closest &stance of approach of the ions and the degree of hydration was estimated according to the method of Stokes and Robinson.
In the development of the chemistry and thermo- amalgam and the ammonia could be observed. dynamics of liquid ammonia solutions the investiga- The cells were measured a t -36O, using amalgams tion of galvanic cells should prove as valuable as of constant composition and varying the salt with aqueous systems. While several cells employ- concentration from approximately 0.002 molal to ing liquid ammonia as the solvent have been mea- saturation. Since no data on the solubility of zinc chloride sured, liquid junction potentials were present in most cases, and no attempt was made to extra- could be found in the literature, this quantity was The dielectric constant of polate the data to obtain the standard potentials. determined a t -36'. Several investigators2 have studied cells of the liquid ammonia a t -36" was also measured since type: M, MA, MIA', MI, where M and M' are it was needed in the calculations. metals and MA and M'A' are soluble salts of the Experimental metals. In most cases a salt bridge was interMaterials.-Commercial anhydrous ammonia was dried posed between the two half cells. sodium and passed through a sintered glass plug t o reTwo cells free from liquid junction potentials over move solid particles before being condensed in the apparatus. have been measured a t 25'. Elliott and Yost3 Zinc amalgam was prepared by electrolyzing a zinc sulmeasured the cell: Zn(Hg), ZnC12.6NH3(s),"4fate solution between a mercury cathode and platinum anCl(m), Tl(s), Tl(Hg), and Garner, Green and Yost4 ode until a two-phase amalgam, containing about 2% zinc, was obtained. Potassium and sodium amalgams were prethe cell: Zn(Hg), ZnClz.6NH,(s), NHCl(m), Cd- pared by similarly electrolyzing a solution of their carbonC12.6NHs(s), Cd(Hg). These cells were found to ates until the desired concentration was obtained. The behave reversibly. amalgams were stored over nitrogen. All salts, except zinc chloride, were reagent grade and were In this work the zinc amalgam, zinc chloride further purified by recrystallization. Zinc chloride was electrode was combined with alkali metal amalgam prepared by electrolyzing a zinc sulfate solution with platielectrodes to form the cells: K(0.2756% amalgam), num electrodes and dissolving the metal in hydrochloric KCl(m), ZnCi2.6NHa(s), Zn (2 phase amalgam), acid. The solution was evaporated to dryness and the salt and Na(0.200070 amalgam), NaCl(m), ZnC12- dehydrated by heating in a vacuum prior to use. Apparatus and Procedure.-The thermostat consisted of 8NH3(s), Zn (2 phase amalgam). The composition a one gallon dewar vessel filled with acetone and cooled by of solid zinc chloride under the experimental condi- alcohol circulated by a centrifugal pump through Dry Ice tions was determined by Biltz and Me~serknecht.~and then through the bath. A bimetallic regulator conI n the present investigation a stationary pool of trolled the operation of the pump as the bath temperature Temperatures were measured with a five junction amalgam was used, and no reaction between the varied. copper-constantan thermel calibrated a t the sublimation (1) Abstracted from the Ph.D. thesis of J. Sedlet. whose present address is Argonne National Laboratory, Chicago, Ill. (2) F. M. G. Johnson and N. T. M. Wilsmore, Trans. Faraday Sac., 3, 77 (1907); G. I. Costeanu, Compf. rend., 198, 778 (1932). 19'7, 1113 (1933); V. A. Pleskov and A. M. Monossohn, Acta Physrorhim. (U.R. S. SS.),1, 871 (1935); 3, 615 (1935); 6, l(1537); 1% e59 (1940); ai, 235 (1940). (3) N. Elliott and D M Yost, Tzus JOURNAL, 66, 1057 (1934). (4) C.S.Garner, E. W. Green and D. M. Yost,rbid.. 61,2055 (1935). ( 5 ) 'A'. Biltz and C Messerkriecht, 2. anorg. Chem , 129, 161 (1523).
point of carbon dioxide. All measurements were made a t -36 i 0 . 0 5 O . The electromotive force cell was of the conventional type with platinum wires sealed into the bottom of the H-cell which was closed at the top with ball and socket joints. A sintered glass plug in the cross arm prevented solid zinc chloride from coming over into the sodium amalgam compartment. The amalgams and zinc chloride were put into the cell with a stream of dry and oxygen-free nitrogen passing through the apparatus. Solutions of salts in liquid
ELECTRODE POTENTIALS IN LIQUID AMMONIA
Dec., 1951
ammonia were prepared in a separate side arm fitted with a magnetically operated stirrer. The solutions were forced over into the H-cell by a slight excess of ammonia gas under pressure. The potentials reached a constant value in from two to five hours, depending on the concentration, and were observed for about four hours after becoming constant. The readings varied from f O . l mv. for the most concentrated solutions t o f 0 . 3 mv. for the most dilute. Samples of the alkali amalgams were analyzed each time they were used. The concentration of the potassium amalgam remained constant, but not the sodium amalgam. Hence cells were also measured in which the sodium chloride concentration was kept constant, and the composition of the amalgam varied. From these results the potentials of the cells in which the sodium chloride concentration varied were corrected to a sodium concentration of 0.2000 per cent. The concentration of the potassium amalgam was 0.2756 per cent. The solubility of zinc chloride in liquid ammonia at -36' was determined to be (1.70 f 0.05) X lo-' molal. The apparatus consisted of a closed glass vessel fitted with a magnetically operated glass stirrer so that about 50 ml. of liquid ammonia could be saturated with zinc chloride. After stirring for five hours, the solution was forced over through a sintered glass plug into a graduated vessel. After the ammonia was evaporated, the salt remaining was dissolved in dilute hydrochloric acid and analyzed for zinc with a polarograph. The dielectric constant was measured with a General Radio Company Type 916-A Radio Frequency Bridge at a frequency of 4.110 Mc./sec. The liquid ammonia had a specific conductance of mho. Three determinations of the dielectric constant gave values of 23.6, 23.8 and 23.8. These values plot favorably with data given by other workers at various temperatures.
Calculations and Results
5809
TABLE I DATAFOR CELL: K(0.2756% amalgam), KCl(m), ZnClz. GNHa(s), Zn(2 phase amalgam) m E Eo" Y Y' 0.1045 0.001492 0.3834 0.740 0.740 .001923 .3750 .lo47 .705 .705 ,003204 .3575 ,1037 .650 ,650 .lo35 .591 ,591 .004696 .3457 .lo03 .425 .425 ,01398 .3147 .426 .427 ,01426 ,3137 .0997 .0978 ,356 .357 ,02151 .3042 .0972 .351 .352 .02266 .3027 By extrapolation, Eo' = 0.1052
TABLE 11 DATAFOR CELL: Na(0.2000% amalgam), NaCl(m), ZnCln. 6NHa(s), Zn(2 phase amalgam) m
0.002148 .002763 .004968 ,008551 .01518 .01719 .03014 .03528 .04427 .07497 ,09220 .lo58 ,1399 .2666 .2779 .4316 .5283
E
EO"
Y
0.4835 0.2198 0.700 .2196 .666 .4762 .591 .4559 .2186 .2179 .500 .4421 .2150 .416 .4262 .2142 .398 .4235 .313 .4106 .2098 ,288 .4073 .2083 .257 .4014 .2053 .1974 .185 .3952 .1909 .169 .3892 .I56 .3865 .E368 .135 .1762 .3820 .1482 .0826 .3752 .3739 .1447 .0822 .0569 .3702 .1157 .3672 .0958 .0520 By extrapolation, &?' = 0.2208
7'
0.700 .666 .591 ,500 ,417 .399 .314 .289 .258 .186 .170 .157 .136 .0834 .OB1 .0578 ,0531
To obtain the standard potentials of the cells, the electromotive force data (Tables I and 11) were extrapolated according to the method of Hitchcock.6 For the cell with sodium chloride as electrolyte, the reaction for the passage of one 1/2ZnCI2.6NH3faraday of current is: Na(Hg) (s) = NaCl(m) 1/22n(Hg) 3NHa. The e.m.f. of the cell is given by According to the above equation, a plot of the left member, designated as EO", against m should be a straight line with the ordinate equal to Eo' at m = 0. For the more concentrated solutions this Since the concentration of amalgam was maintained constant, the alkali metal activity is not included curve showed some curvature, but approached in equation (1) and the standard potential desig- linearity in dilute solutions such that reasonable nated as EO'. The molal activity coefficient y of a extrapolations were possible for both cells. The activity coefficients of potassium and uni-univalent electrolyte is given by the Debyesodium chloride were calculated by means of equaHiickel theory. tion (1). The coefficients which were calculated assuming the activity of the ammonia to be unity log y = - -___ 1.824 (DT)'h -Iare given in the columns headed y in Tables I BC - log ( 1 0.034 m ) (2) and 11. To obtain the true activity coefficients In this equation D is the dielectric constant and the it is necessary to know the activity of ammonia. concentration is given in moles per liter. Equa- A reasonable approximation for this quantity may tion (1) can be rewritten with the aid of equation be obtained by using the mole fraction of ammonia, ( 2 ) to give an expression suitable for extrapolation : assuming the solute to be 50% dissociated. The E 0.09409 m - 0 . 3 3 7 2 6 - 0.09409 log (1 coefficients so calculated are given in the column 0.034 m) = Eo" = Eo' - B'm. Two approxi- headed y'. An examination of the data showed mations were made in obtaining the latter equation. that the coefficients approach those predicted by The concentration of electrolyte in volume units the Debye-Hiickel limiting law in dilute solutions. was approximated from the molality by the relation The deviations, however, increased more rapidly c = dm, with d = 0.6851 g./ml., the density of with increasing concentration in ammonia than in ammonia. The activity of ammonia in these solu- water a t equal concentrations. Stokes and Robinson' proposed a modification tions is not known, and was assumed to be unity. These approximations will not affect the extrapola- of the Debye-Hiickel equation, introducing a tion since they become valid in dilute solutions. term to correct for the solvation of the ions and a
+
+
+
+
+
+
(6)
D. I. Hitchcock, THISJOURNAL, 60, 2076 (1928).
(7) R. H.Stokes and R. A. Robinson, ibid., 70, 1870 (1948).
EUGENE E. KETCHEN AND W. E. WALLACE
5810
Vol. 73
solutions and, if so, to compare solvation in water and ammonia. The parameters were varied to obtain the best fit to the data for sodium chloride. The results are given in Table 111. Good agreement was obtained for a = 3.53 f 0.02 A., and n = 2.5 =t0.2 with an average difference between the calculated and observed coefficients of 4 in the TABLE 111 third significant figure. Stokes and Robinson COMPARISON O F EXPERIMENTAL AcrrvITY COEFFICIEXTS O F found for *sodium chloride in water solution that SODIUMCHLORIDE WITH THOSE CALCULATED BY E Q U A T I O ~a = 3.97 A. and n = 3.5. This difference between OF STOKES A N D ROBIXSON the two solvents is reasonable in view of their diParameters: a = 3.53; n = 2.6 electric constants. llt Ycnptl. The hydration parameter could not be deter(0.780) 0.778 0 * 001 mined for potassium chloride since the terms in the .706 .002 .i 0 5 equation containing n are negligible a t the highest ,655 ,658 .003 concentration used. The parameter, a, was deter.588 .005 ,589 mined by fitting the data to the Debye-Hiickel * 480 ,484 .01 equation for uni-univalent electrolytes, log y = .372 ,377 .Q2 -AV“&/(l aBdG), where A and B are con.242 ,248 .05 stants equal to 3.5835 an4 0.3922, respectively. A ,164 . 10 . i6o value for a of 3.3 f 0.1 A. reproduces the data to ,104 ,103 .20 f0.003 in y. X value of 3.63 was found by Stokes .30 .Q767 ,0775 and Robinson for potassium chloride. .40 . 0014 ,0629
term to convert from rational to molal activity coefficients. They were able to reproduce] quantitatively, the experimental activity coefficients in aqueous solution of a large number of electrolytes up to four and five molal. It is of interest to determine if the equation is valid for liquid ammonia
YCdCii.
+
.50
,0535
I0540
[CO~TRIBUTIONS O . 804 FROM
THE
RECEIVED JUXE18, 1951
LAFAYETTE, ISD.
DEPARTMENT O F CHEMISTRY, UNIVERSITY
OF PITTSBURGH]
Thermal Properties of the Alkali Metals. I. The Heats of Reaction of Sodium and Potassium with Water at 2 5 O 1 l 2 BY EUGENE E.KETCHEN AND W. E. WALLACE The heats of reaction of potassium and sodium with an infinite quaiitity of watcr a t 28’ were measured aiid found to be -47,190 zk SO and -44,350 i 70 calories/gram atom, respectively. These results are i n reasonably good agreement with other recent determinations except for Roth and Kaule’s work on the sodium-water reaction. with the forniation of ionic solid solutions.6 Constructional Introduction details and operating characteristics have been reported in While determining the heats of formation of connection with that work. The only modification needed sodium-potassium alloys reported in the following in the present study was the replacement of the crushing used to introduce the salts into the calorimeter by a paper, it became necessary to have reliable values deviceschamber for carrying out the metal-water reaction. for the heats of reaction of the two pure metals with special The reaction chamber consisted essentially of a monel rod water. Since several earlier determinations of these drilled out to receive the reactants and nickel plated on the heat effects led to somewhat divergent results13~*interior to improve its corrosion resistance to the alkaline particularly for potassium, it seemed desirable to solution produced. Its operation will be made clear in the which follows. carry out a redetermination of these important discussion In a typical experiment approximately 0.002 t o 0.006 thermochemical quantities in the hope that more gram atom of alkali metal contained in a crushable glass reliable values might emerge. The measured heat bulb was allowed to react with about 20 g. of water. From effects are of interest not only for their use in 90 to 270 calories of heat were developed, being communifor the most part to water (about 1 kg.) surrounding obtaining heats of alloying but also because much cated the reaction chamber. of the thermochemical information pertaining to There are several sources of difficulty customarily assosodium and potassium compounds depends on the ciated with experiments of this type. These have presumdeterminations of the heat produced when the ably been eliminated in the reaction chamber used. Much of the trouble stems from the violence with which the reacmetals react with water. tion takes place. In such reactions the reactants are fre-
Experimental Methods Calorimeter .-The calorimeter employed in this study was the instrument constructed for studying the heats associated (1) From a thesis submitted to the University of Pittsburgh by Eugene E. Ketchen in partial fulfillment of the requirements for the Doctor of Philosophy degree, August, 1950. (2) This work was assisted by the Office of Naval Research and the Atomic Energy Commission. (3) P. R. Bichowsky and F. D. Rossini, “The Thermochemistry of the Chemical Substances,” Reinhold Publishing Corn, New York, N. Y., 1936. (4) W. A. Roth and H. 1, Katile, 2. onorg Chcm., 163, 352 (1047)
quently blown apart before the process has gone to completion. To minimize this possibility a series of baffles was provided. Evidently these baffles were very effective in confining the reactants for the relatively short reaction time for if any alkali metal had been thrown up and lodged above the water level, reaction would have been continued with water vapor and this mishap would have been recognized by the thermal response of the calorimeter toward the conclusion of an experiment. With the chamber employed no such behavior was ever observed. In preliminary experiments using a chamber of a somewhat different design a slow
__
(51 31. .I. Ijiticluau and W. li. Wallace, Tms ( 19.18).
JUURNAI,,
TO, 4162
