3723
OSMOTIC COEFFICIENTS OF ALKALIMETALCHLORIDE SOLUTIONS
Acknowledgment. This work was supported in part by a grant from the Pittsburgh Activated Carbon Division, Calgon Corporation. In addition, D. A. Wohle-
ber received support from an N.D.E.A. Title IV fellowship and from the Goodyear Tire and Rubber Fellowship Fund.
Osmotic Coefficients of One Molal Alkali Metal Chloride Solutions over a 300" Temperature Range1
by W. T. Lindsay, Jr.,* and Chia-tsun Liu Westinghouse Research Laboratories, Pittsburgh, Pennsylvania (Received April 6 , 1072) Publication costs assisted bg the Westinghouse Research Laboratories
Osmotic coefficients are reported for approximately 1 m solutions of LiCl and CsCl a t 25" intervals from 125 to 276 and 300". These results are summarized with previously determined osmotic coefficients for 1 m solutions of HCl, LiCl, NaCl, KC1, RbCl, and CsCl to give an overall view of individual ionic character in these solutions from the freezing points to high temperatures. Excess enthalpies and entropies are calculated for the water in 1 m HC1, LiC1, NaC1, and CsC1. Individual ionic character decreases considerably in high temperature 1 m solutions of these salts.
Introduction
Experimental Section
The individual ionic character of alkali metal halides in moderately concentrated aqueous solutions has been of long-term interest in the development of the physical chemistry of electrolytes. This subject has had various interpretations, with most recent discussions focusing on the differing effects of the ions on water structure. A principal manifestation of individual ionic character is the wide variation of activity and osmotic coefficients for alkali metal halide salt solutions at or near room temperature, yet little attention has been given to the temperature dependence of this phenomenon. In this paper we report determinations of osmotic coefficients for approximately 1 m solutions of lithium chloride and cesium chloride from 125 to 275 and 300". By many criteria, these two salts have very different effects on water structure at room temperature, LiCl being regarded as a "structure maker" and CsCl as a "structure breaker." Summarizing the experimental results of this work with other data for these and other salts gives a clear picture of the temperature dependence of 1 m osmotic coefficients for alkali metal chloride and hydrochloric acid solutions from their freezing points t o temperatures as high as 300". The comparison shows that individuality of ionic character in these solutions decreases markedly at high temperatures.
The experimental method used in this work determined the difference between the vapor pressures of solution and pure water, each contained in separate but identical pressure vessels held at the same temperature. The apparatus and techniques developed for this measurement at temperatures to 300" have been described elsewhere in some detai1.2-4 Stating the essential features in brief, the important criterion of negligible difference between the temperatures of the two vessels was achieved by a "double thermostating" method, with the two vapor pressure cells symmetrically installed in a massive aluminum block, the block being immersed in a vigorously stirred thermostat bath regulated to a constancy of +0.002" or better. A bellows arrangement confined water vapor to the isothermal environment of the cells. A differential transformer system was used to sense the position of each bellows, allowing balancing of cell internal steam pressures with externally applied helium gas pressure (1) Presented in part at the 153rd Kational Meeting of the American Chemical Society, Miami Beach, Fla., April 1967. (2) W. T. Lindsay, Jr., and C. Liu, OSW Research and Development Progress Report 347. Available from Superintendent of Documents, U. S. Government Printing Office, Washington, D. C. 20402. (3) W. T. Lindsay, Jr., and T. S. Bulischeck, Rev. Sci. Instrum., 41, 149 (1970). (4) C. Liu and W. T. Lindsay, Jr., J. Phys. Chem., 74, 341 (1970).
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W. T. LINDSAY, JR.,AND CHIA-TSUN LIU
3724 with a precision usually about k 0 . 2 Torr or better. (A greater degree of variability was encountered in a few cases at some of the higher temperatures; however, the pressure differences were still measurable with acceptably small standard deviations of 0.1 to 0.4%,) The difference between balancing gas pressures was then determined by a high pressure differential manometer or by a Texas Instruments quartz spiral Bourdon gauge in a pressure capsule. This arrangement has been used successfully to measure the vapor pressure lowering of high temperature KaCl solutions as dilute as 0.1 m e 4 The measurements were carried out on analytical grade salts with a reported purity of 99.885% for LiCl (Fisher "Certified") and 99.9% for CsCl (Kawecki "High Purity"). Impurity levels many times those reported would have negligible effect on the results. The salts were dried in a vacuum oven at 100" for several days and used without further purification. For the lithium chloride runs, a dcgasscd solution at a concentration of 0,9911 m at room temperature was prepared on a weight basis from the dried salt and triply distilled water following the previously described techniques. Measurements of vapor pressure differences between this solution and pure water followed the previously established procedures at temperatures 125, 150, 175, 200, 225, 250, and 276'. At 300°, a leak developed in the solution cell, causing a slow increase in concentration of the solution; hence, no valid data were obtained at this temperature. All other information (magnitude of absolute pressures at 125O, constancy of differential pressure measurements at each temperature, absence of gas in reference cell, etc.) indicates that the results at 275" and below are accept able. For the cesium chloride runs, the solution after preparation and degassing was at a slightly higher concentration of salt, 1.0075 m at room temperature. Satisfactory measurements were obtained with this solution at all temperatures 125 to 300".
Results The results are given in Table I. The concentrations reported in the table were calculated from the original concentrations of the solutions with correction for the evaporation of water needed to raise the steam density in the vapor volume of the cell, taking expansion of the liquid into account. Vapor pressures given for pure water were calculated at each temperature from the correlation of Smith, Keyes, and Gerry.j Osmotic coefficientswere calculated from the relation + =
'Oo0 [ ~ T l nPo- vmMlR T P
where
+ is the osmotic coefficient, v is 2, m is molality,
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Table 1: Experimental Results for Xominally 1 m Solutions of LiCl and CsCl
Temp,
"C
125 150 175 200 225 230 275 300 126 150 175 200 223 2.50 275 300
Molality a t temperature, mo1/1000 g HzO
Vapor pressure lowering AP = PO P, Torr
-
Lithium Chloride Solutions 0.9934 60.46 0 9946 123.97 0,9989 233.46 1.0042 412.13 1.0121 686.48 1.0233 1101.45 1.0395 1709.89 I
...
, . ,
Cesium Chloride Solutions 1.0096 55.24 1.0116 113.88 1.0148 215.98 1.0196 385.56 1.0267 646.33 1.0369 1044.12 1.0514 1648.22 1.0720 2543.39
Osmotic coefficient + a t P = PO
0.960 0.941 0.918 0.894 0.861 0.825 0.778
... 0.862 0.849 0.835 0.823 0.798 0.771 0.740 0,695
M I is the molecular weight of water, R is the gas constant, Tis absolute temperature, p o is the vapor pressure of pure water, p is the vapor pressure of the solution, Ap is p o - p , Vlo(g) is the analytical function of pressure and temperature for the molar volume of pure water vapor given by Smith, Keyes, and Gerryj5and VI' (1) is the corresponding function for the molar volume of pure liquid water, used here as a valid approximation for the partial molal volume of water in the solution. The second term in the brackets is the correction for deviation of water vapor from perfect gas behavior, and the third term is a small correction for the purpose of reporting the osmotic coefficients at a pressure equal t o the vapor pressure of pure water at each temperature. Figure 1 shows the osmotic coefficients from this work designated by the solid circles and the solid line portions of the curves for LiCl and CsC1. Also shown by solid circles and a solid line are high tempeiature osmotic coefficients for XaC1. These XaCl data are calculated for a concentration of 1.0 m at 125 to 300" from a correlation equation that we developed earlier2 to represent our own4 and other6-s results on sodium chloride solutions at concentrations from 0.1 t o 3 m at these temperatures. The LiCl and CsCl points are those for the concen( 5 ) L. B.Smith, F. G. Keyes, and H. T. Gerry, Proc. Amer. Acad. Arts Sci., 69, 137 (1934). See also J. H. Keenan and F. G. Keyes, "Thermodynamic Properties of Steam," Wiley, Kew York, N. Y . , 1936. (6) E. R. Gardner, P . J. Jones, and H. J. de Nordwall, Trans. Faraday Soc., 59, 1994 (1963). ( 7 ) E. R.Gardner, ibid., 65, 91 (1969). (8) B. M. Fabuss and A. Korosi, Desalination, 1, 139 (1966).
OSMOTICCOEFFICIENTS OF ALKALIMETALCHLORIDE SOLUTIONS
1.05
1.00
.95
-.a
a c,
.90
.-
V
5
.E
.85
4-
8 .80
.75
.70
I
0
I
50
I
I
1m 1% Temperature.
I
I
200
I
I
250
I
O C
Figure 1. Temperature dependence of osmotic coefficients of 1 m alkali metal chlorides and HCl. This work and ref 2 (a), ref 18, 24, and 25 (0),ref 20 (m), ref 19 (U), ref 17 (A), ref 13, 14, and 15 (A), ref 16 (V),ref 12 (V), ref 10 (a),ref 11 (a), ref 21 (8).
trations of Table I. We have no way of correcting these t o exactly 1.0 m without additional data on concentration dependence of the osmotic coefficients. However, the high temperature osmotic coefficients for NaCl are only slightly dependent on concentration in the vicinity of 1 m because of the minima in the curves of osmotic coefficient vs. concentration. Considering (1) the effects of temperature on the shapes of these curves for NaCl and on the positions of their minima4 and (2) the relative positions of the minima for the curves for LiC1, NaC1, and CsCl at room temperature, it appears probable that osmotic coefficients for LiCl and CsCl at high temperatures are also only slightly dependent on concentration in the neighborhood of 1 m.
Discussion On Figure 1 we have also shown points summarizing 1 m osmotic coefficient data for HCI, LiCl, NaC1, KC1, RbC1, and CsCl at room temperature and a t as many other temperatures as possible. Some of these data were obtained from freezing point depressionsJ10-12 some from boiling some by isopiestic cornpari~on,~~-~7 point elevations,18some by measurement of vapor pres-
3725
sure l o ~ e r i n g , ~ ~ ~and J ~ Jsome O were calculated from activity coefficients determined by emf methods.21 Although the data are of varying degrees of reliability, they fall into an overall pattern that is consistent with the high temperature results we report. The complete set of curves shows a regular and systematic variation from HCl to CsCl, with no exceptions or anomalies. We believe Figure 1t o give a much clearer picture of the temperature dependence of 1 m osmotic coefficients for the alkali metal chlorides (and, by inference, of the effects of temperature on individual ionic character) than has heretofore existed. Some observations may be noted from the information of Figure l. First, the relative order of l m osmotic coefficients, HCl > LiCl > KCl > RbCl > CsCl, does not change from the freezing points to the highest temperatures for which there are data. Second, the spread between the osmotic coefficients decreases very greatly as the temperature increases. Third, above lOO", the temperature coefficients are all of the same sign and the curves are quite similar, in contrast to the situation at room temperature. Fourth, the temperatures of the maxima in the curves progress regularly from about 80" for CsCl to 60" for NaCI, with the maximum for LiCl about at the freezing point and the apparent maximum for HCl below the freezing point. All of these observations are consistent with the notion that individual ionic character at room temperature is due primarily to structure in the solvent, which gives possibilities for both structure-making and structurebreaking effects of the ions according to the two-zone co-sphere model elaborated by Frank,22Gurney,23and (9) G. Scatchard in "The Structure of Electrolytic Solutions," W. J. Hamer, Ed., Wiley, New York, N. Y., 1959, Chapter 2. (10) G. Scatchard and S. S. Prentiss, J . Amer. Chem. Soc., 55, 4355 (1933). (11) H. M. Chadwell, ibid., 49, 2795 (1927). Interpolated for 1.0 m. (12) G. Karagunis, A. Hawkinson, and G. Damkohler, Z. Phys. Chem. (Leipzig), 151, 433 (1930). Interpolated for 1.0 m as required. (13) C. S. Patterson, L. 0. Gilpatrick, and B. A. Soldano, J . Chem. Soc., 2730 (1960), using 4 ~ = ~0.935 ~ at 199.6'. (14) B. A . Soldano and C. S. Patterson, J . Chem. SOC.,937 (1962), using @NaCl = 0.924 at 121.1O. (15) B. A. Soldano and M. Meek, ibid., 4424 (1963), using+Naci = 0,911 at 140.3'. (16) K. L. Hellams, C. S. Patterson, B. H. Prentice, 111, and M . J. Taylor, J. Chem. Eng. Datu, 10, 323 (1965), using +xaci = 0.942 a t 45". (17) L. L. Makarov, Y. G. Vlasov, and V. I. Isotov, Russ. J . Phys. Chem., 38, 1297 (1964). Isopiestic ratio at 1.3 m NaCl used with 4 ~ = ~0.942 ~ a t 145' and 1.0 m. (18) R. P . Smith, J . Amer. Chem. SOC.,61, 500 (1939). (19) H. F. Gibbard, Jr., Ph.D. Thesis, Department of Chemistry, Massachusetts Institute of Technology, 1966. Data are for concentrations 1.046 to 1.049 m. (20) H. F. Gibbard, Jr., personal communication, July 7, 1970. Data are for concentrations 1.0208 to 1.0238 m.. (21) R. S. Greeley, W. T. Smith, Jr., M. H. Lietake, and R. W. Stoughton, J . Phys. Chem., 64, 1445 (1960). (22) H. 8 . Frank and M. W. Evans, J . Chem. Phys., 13, 507 (1945). (23) R. W. Gurney, "Ionic Processes in Solution," Dover Publications, New York, N. Y., 1953.
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W. T. LINDSAY, JR.,AND CHIA-TSUN LIU
3726
others. The net effect of these opposing influences can be widely variable in sign and magnitude for ions of various sizes. On the other hand, only structure enhancement is possible in the thermally disrupted solvent a t high temperatures, where the similarities of the curves of Figure 1 are more striking than the differences. Particularly, the similar slopes at high temperatures suggest that CsCl, RbC1, KC1, and NaCl become more like LiCl, which is regarded as a structure maker even at room temperature. The similarities and differences in the data for these solutions can be examined in another way by displaying correlations between excess enthalpy and excess entropy for the solvent. For this purpose, the data for HCl, LiC1, NaC1, and CsCl were fitted by a leastsquares procedure to a cubic polynomial of the form
4
= a0
+ alT + azTZ+ a2T3
(2)
where T is temperature in degrees Kelvin and ao, al,a2, and a3 are constants. Double weight was given to the identical 25” values tabulated by Robinson and Stokes24 and by Wu and Hamer,26because of the large number of high quality measurements that these authors have correlated. All others, including our own, were given single weight. The excess partial specific enthalpy H l e x was calculated by differentiating the resulting expressions26according to the relation
Hlex = vmRT2b4/bT
(3) and the excess partial specific entropy SleXwas calculated from the relations Flex = vmRT(1 - 4 )
(4)
SieX= (Hiex - F1PX)/T
(5)
and where Flex is the excess partial specific Gibbs free energy (per kilogram of H20).27 The results are shown by Figure 2, where Hlex is plotted against TSlex. The magnitude of Flexcan be determined for any point on this plot by the vertical or horizontal distances from the line marked Flex = 0, for which Hlex = TSlex. It should be recognized that the exact magnitudes of the values calculated for SleXand H l e x at each temperature depend t o some extent on the analytical form of the equations assumed t o represent the data. On the other hand, the correlation of XleX with HLex for each salt is relatively insensitive to the choice of equation; Le., qualitatively equivalent results are obtained whether quadratic, cubic, or quartic polynomials are used. The cubic equations fit the data with an accuracy sufficient t o illustrate the points discussed below. Table I1 gives the values of the parameters of eq 2 for HC1, LiC1, NaC1, and CsCl. The solid lines of Figure 2 may be viewed as a correlation between solvent excess enthalpy and solvent excess entropy for solutions of a single salt in a series of solThe Journal of Physical Chemistry, Vol. 76, N o . 24, 1972
H Y , calories / kg H20
Figure 2. Correlations of excess enthalpy with excess entropy for the water in 1 m CsCl, NaC1, LiC1, and HC1 a t temperatures 0 to 300”. Points are shown a t 25” intervals beginning a t 0” for each solution.
vents with progressively varying dielectric constant and extent of structure. These lines are quite accurately linear for both HC1 and LiCl over the complete temperature range of the data. The lines for KaC1 and CsCl are both curved at the lower temperatures, but they approach linearity at the higher temperatures. These observations indicate no major qualitative change in the nature of the interactions in solution for HCl and LiCl as the temperature increases, They indicate also that the interactions in solution for NaCl and CsCl at the lower temperatures are qualitatively different from those of these same salts a t higher temperatures, and from HCl and LiCl at any temperature, conclusions that are consistent with those that can be inferred from Figure 1. (24) R . A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Academic Press, New York, N. Y.,1959. (25) Y X . Wu and W. J. Hamer, U. S. National Bureau of Standards Report 9908, “Electrochemical Data. Part XI. Osmotic Coefficients and Mean Activity Coefficients of Aqueous Solutions of HC1, HBr, H I and,:he Chlorides, Bromides and Iodides of the Alkali Metals at 25OC, Sept 6, 1968. (26) Direct differentiation of the data with respect to temperature is not strictly at constant pressure, since the high temperature osmotic coefficients are reported a t pressures equal to the vapor pressure of water a t each temperature. We have estimated the correction for NaCl solutions that is required to compensate for this procedure, by using partial molal volumes derived from the density correlation of J. L. Haas, Jr., Amer. J. Sei., 269, 490 (1970). The corrections are negligible for 1 m solutions of NaC1. They are assumed negligible also for 1 m solutions of LiCl and CsC1. (27) See H. L. Friedman, “Ionic Solution Theory,” Interscience, New York, N. Y . , 1962, for a discussion of these excess partial functions.
OSMOTIC COEFFICIENTS OF ALKALIMETALCHLORIDE SOLUTIONS
Table 11: Parameters for Representation of 1 m Osmotic Coefficient Data
Temp range, "C Number of data points Parameters for eq 2 a0
lOzal 106az 109as Standard error of estimate for osmotic coefficient
a6
3727
a Function of Temperature
HCl
LiCl
NECI
CSCl
-3.93t0200
-3.80 to 275 20
-3.39 to 300 19
-2.14 to 300 14
1,31819
0.279727 0.431668 - 0.809964 3.31285 0,003640
0.00625499 0.551187 - 1.05079 5.25157 0.007537
7 1.22202
- 0.130086 0.408428
- 5 * 75394
0.006793
Significant new information from Figure 2 is shown by the dashed lines. These indicate linear correlations, a t each temperature from 300" down to about 125", between solvent excess enthalpy and solvent excess entropy for the different solutes. Both excess enthalpy and excess entropy for the water are more negative in solutions of the chlorides with the smaller cations. We believe these correlations and their progressively varying slopes are as one would expect for the effect of ion size on degree of solvation in a series of unstructured polar solvents with differing dielectric constants. It would be very desirable to make a quantitative test of this contention by comparison of the data with theoretical calculations based on a suitable model. This task may be more readily achievable for unstructured aqueous solutions at high temperatures than accomplishment of the more complex theoretical development needed at lower temperatures. The above-mentioned linear correlations break down completely at temperatures below loo", as shown in clearer detail in the lower part of Figure 2. We believe this is evidence strongly suggestive of a new factor coming into play at the lower temperatures, namely, the effect of water structure in amplifying the influence of ion
- 0.220256
0.608859
- 7,04722
0.003585
size on the thermodynamic properties of electrolyte solutions. Osmotic coefficients are only one of the many kinds of solution properties that show the influence of solute individuality despite similarity of ionic charge and electronic structure. The variation of 1 m osmotic coefficients for the alkali metal chlorides and hydrochloric acid over a very wide temperature range is consistent with the following hypotheses concerning the more general phenomenon of individual ionic character. (1) Individual ionic character for these solutions diminishes greatly a t high temperatures. (2) Individual ionic character for these solutions a t room temperature is due primarily to the presence of hydrogen-bonded structures in the solvent. (3) Residual individual ionic character a t high temperatures is due mostly or entirely t o differences in extent of hydration. The comparisons presented in this paper illustrate the value that measurements over a very wide temperature range have for interpretation of phenomena in aqueous solutions.
Acknowledgment. This work was performed under contract from the Office of Saline Water. We wish also to acknowledge Mr. T. S. Bulischeck for his valuable assistance with the experimental work.
The Journal of Phgsical Chem$stry, Vol. 76,No. a4, 1971
