Heat of Dilution of Cesium Chloride A. S. LEVINE and R. H. WOOD Department of Chemistry, University of Delaware, Newark, Del.
19711
The heat of dilution of CsCl at 2 5 O C. was determined a t concentrations of 1.5,
2.0, 2.5, and 3.0 m. The excess entropy,?E, was calculated from the experimental heat of dilution and compared to the entropy expected for a "normal" electrolyte. The experimental and normal values differ greatly, confirming the conclusion from other evidence that CsCl is ion-paired to an appreciable extent.
W H E N these measurements were started, the heat of dilution of CsCl had not been reported above 1.1 molal. Recently, Stakhanova et al. ( 8 ) reported values of +I, from 2.0 to 5.0m, which differed appreciably from t h e present results.
where q is t h e heat given out and n is the moles of CsCl. The equation for the thermistor is
EXPERIMENTAL
The calorimeter, experimental procedure, and CsCl solution have been described in detail elsewhere ( 4 , 7, I O ) . Briefly, the calorimeter is a 60-ml. Dewar flask stirred by a magnetic stirrer. The temperature is sensed by a thermistor and electrical calibration is used. Tests (7, 1 0 ) show that for high heats the calorimeter is accurate to about 2 1 % , and for low heats the accuracy is about +- 0.01 cal. The initial temperature was adjusted to 25.00 -C 0.02" C. and the pipette was opened. Electrical calibration followed completion of the measurement of the heat of dilution. TREATMENT OF DATA
The values of A+I. were calculated from the data using t h e equation
where Ri and R, a r e the initial and final thermistor resistances, and Tiand T, a r e the initial and final temperatures. The average heat capacity is (3)
and therefore
The relative apparent molal heat content a t the initial concentration, $r,(mi),was calculated from the value of A+r, obtained above and t h e relative apparent molal heat content a t the final concentration, +,(mf), by the equation
Table I. Heat of Dilution of CsCl at 25" C. &L,
Cal./Mole 1.5 1.5 1.5
2.0 2.0 2.0
2.5 2.5 2.5
3.0
3.0 3.0
0.2123 0.2040 0.2168
0.2837 0.2684 0.2881
0.2330 0.2423 0.2398
0.2850 0.2725 0.2787
3.206 3.083 3.296
5.640 5.326 5.748
4.234 4.406 4.357
5.969 5.583 5.871
35.6 37.2 34.7
20 6 23.8 19.7
31.3 29.3 29.9
20.3
23.0 21.7
+L (mi), Cal. /Mole
271.7 273.1 271.5
-236.1 -235.9 -236.8
Average
-236.3
357.4 359.4 357.0
-336.8 -335.6 -337.3
Average
-336.6
443.9 440.1 441.5
-412.6 -410.8 -411.6
Average
-411.7
510.2 51 1.7 512.7
-489.9 -488.7 -491.0
Average
-489.9
+Lh(mi),
Cal. /Mole
- 347.0
-423.6
-494.6
From Parker's low-concentration data, NSRDS-NBS 2, April 1965. Calculated from Stakhanova et al (a), using Parker's value at m = 1.0.
Journal of Chemical and Engineering Data, Vol. 15, No. 1, 1970
33
A+L = h ( m i ) - mi)
(5)
The value for CL(mf) was obtained from Parker’s lowconcentration data ( 5 ) . RESULTS AND DISCUSSION
Table I shows the heat of dilution of CsCl a t 25” C. The last column in this table shows the values of +r, calculated from the data of Stakhanova et al. ( 8 ) . together with Parker’s ( 5 ) value a t l m . The present results a r e 5 to 12 cal. per mole more positive. This could be due to a n error of this magnitude in Parker’s value at l m . Wood (9) showed t h a t the partial molal entropies of normal strong electrolytes followed a single family of curves and could be correlated by a single parameter for each ion. The low-concentration data then available indicated that CsCl did not follow this correlation and was a n “abnormal” electrolyte. The present results confirm this conclusion. The values of TgE were obtained from the equation ( 9 ) T%E
=
-
@
=
+L
+ m (ddL/dm) - vRT In y r
and -659 cal. per mole for m = 1.5, 2.0, and 3.0, respectively, a s compared to the values predicted f o r a normal electrolyte (9) of 190, 45, and -179 cal. per mole for m = 1.5, 2.0, and 3.0, respectively. The most likely reason for abnormal behavior is ion-Dairinn. and there is other evidence to support the hypothesi’s’ that CsCl is ion-paired (1, 2, 3 ) . LITERATURE CITED Davies, C. W., “Ion Association,” pp. 30, 169, Butterworths, London, 1962. Davies, C. W., Trans. Faraday SOC.23, 351 (1927). Hindman, J., J . Chem. Phys. 36, 1000 (1962). Levine, A. S., Bhatt, N., Ghamkhar, M., Wood, R. H., J. CHEM.ENG.DATA15, 34 (1970). Parker, V. V., “Thermal Properties of Aqueous Univalent Electrolytes,” p. 61, U. S. National Bureau of Standards NSRDSNBS 2, Washington, D. C., 1965. Robinson, R. A,, Stokes, R. H., “Electrolyte Solutions,” p. 495, Butterworths, London, 1959. Rooney, R. A., M.S. thesis, University of Delaware, Newark, Del., 1968. Stakhanova, M. S., Vlasenko, K. K., Karapet’yants, M. Kh., Bazlova, I. V., Russ. J . Phys. Chem. 42, 274 (1968). Wood, R. H., J . Phys. Chem. 63,1347 (1959). Wood, R. H., Rooney, R. A,, Braddock, J. N., Ibid., 73,1673 (1969).
(6)
using the values of y3 from Robinson and Stokes ( 6 ) together with t h e dr, from these measurements. The values of d + , / d m a r e from the slope of a large-scale m, using the present results and those graph of +I, of Parker. The experimental values for TTE are -385, -431,
RECEIVED for review February 27, 1969. Accepted November 6, 1969. This study was aided by a grant from the Office of Saline Water, U. S. Department of Interior, Washington, D. C.
Heats of Mixing of Some Aqueous Alkali Metal Chloride Solutions A. S. LEVINE, NALlNl BHATT, MEHDI GHAMKHAR, and R. H. WOOD Department of Chemistry, University of Delaware, Newark, Del. 1971 1 The heats of mixing CsCl with LiCI, CsCl with KCI, and KCI with LiCl were determined a t various ionic strengths. The heat of mixing CsCl with KCI was very small and the coefficient RTh, changed rapidly with concentration, indicating that the likecharged pair interactions are smaller in this mixing and about as important as triplet interactions. For the CsCl with LiCl and KCI with LiCl mixings, the coefficient
RTh, changed fairly slowly with concentration, indicating that like-charged pair interactions contribute more to RTh, than triplet interactions.
I N CONNECTION with a study of the heat of mixing alkaline earth chlorides with alkali metal chlorides, the heats of mixing CsCl with LiCI, CsCl with KCl, and KCl with LiCl were needed a t ionic strengths of 1, 2, and 3. W u et al. reported results a t unit ionic strength ( I O ) . Stakhanova e t al. (.?I measured the heats of mixing LiCl with KCl and LiCl with CsC1, but their results were not accurate enough for our purposes.
Table
I. Results of the H e a t of Mixing Experiments at 25’ C.
Mixture CsC1-LiCI
CsC1-KC1
34
Journal of Chemical and Engineering Data, Val. 15,
No. 1, 1970
6
1 2
4 5 6
1 2
3 LiCI-KC1
2
3 a
D
w
0.5 3
EXPERIMENTAL
The lithium chloride and potassium chloride stock solutions were prepared from reagent grade salts. All impurities reported for these salts were less than 0.01%. The stock solutions were analyzed for chloride in a gravimetric silver chloride precipitation. The pH of both solutions was between 5 and 6.5.
I
RThoh
RThib
-172 i 11 -192 i 6 -215 i 4 -229 f 3
-13 f 11
7
8.0 f 0.9
6 6
2.8
8 15
z!z
- 7 f 11 -17 f 7 -22 f 5
0.0 f 1
0.2 -0.3 i 0.4
-1.7 & 0.4
-0.7 & 0.7
-80.4 f 0.5
-3.8 i 0.7
-93 f 1
-5
i 2
N is the number of experiments. Units are cal. per kg. of solvent per ionalz (Equation 1).
