TEMPERATURE DEPENDENCE OF AQUEOUSELECTROLYTES
Heats of Mixing.
1455
I. Temperature Dependence of Aqueous
Electrolytes with a Common Anion1 by Henry L. Anderson and Linda A. Petree Department of Chemistry, University of North Carolina at Greensboro, Geensboro, North Carolina !3'41l? (Received September 4, 1969)
The heats of mixing of LiC1-(CHs)4NC1-H20, KC1-(CH&NC1-H2O1 LiCl-KCl-H20, LiC1-NaC1-H20J HC1NaC1-H20, and NaCl-KCl-H20 at constant total ionic strength have been measured at 40, 60, and 80'. The first three mixtures had heats of mixing independent of the temperature whereas the latter three had heats of mixing dependent on the temperature. The results have been interpreted in terms of the solutesolvent structural properties.
Introduction Aqueous electrolyte solutions have received considerable attention in recent years with much of this attention being stimulated by the need to develop an efficient and relatively inexpensive method of converting sea water to potable water. Over the years, a wealth of experimental and theoretical information has been amassed concerning electrolyte solutions and from an a priori point of view the Debye-Huckel limiting law has been firmly established. Unfortunately, efforts to make a priori calculations of thermodynamic properties of electrolyte solutions in real concentration ranges often fall short of the desired level of understanding. A large contribution to this lack of success can be attributed to an incomplete understanding of the specific role water plays in influencing ionic interactions in aqueous solutions. Recently13it has been shown that determination of the excess thermodynamic properties of aqueous mixed electrolyte solutions is an excellent way of studying specific ion interactjons. For example, mixing aqueous salt solutions at constant ionic strength, I , according to the scheme RIX(1)
+ NX(I)
+
mixture(1)
has the advantage of canceling effects due to the ionic atmosphere. Further, in the presence of a common ion (e.g., the common "X" anion) the effects due t o oppositely charged pair interactions cancel, and one observes phenomena due only to like-charged pair and higher order interactions. Young, Wu, and Krawetz4observed that for common ion mixtures the sign of the heat of mixing could be predicted by classifying the ions into two groups. When ions of like classification are mixed, the heat of mixing is endothermic and when ions of unlike classification are mixed, the heat of mixing is exothermic. Wood and Anderson5rs have shown that this classification should be made on the basis of the ion-water struc-
tural relationship. They used the classification of Frank, Evans, and Wen7v8where an ion is classified as either a structure maker or a structure breaker. In order to gain more specific information concerning the nature of the solute-solvent interaction, the present work was undertaken. Since water structure is strongly temperature dependent, the temperature dependence of the heats of mixing should shed considerable light on the role water plays in determining the nature of the specific ion interactions upon mixing. There has been considerable discussion in the literature concerning the nature of specific ion interactions. On the one hand, the Brgnsted specific ion interaction principle has been defended by Scatchard'" and on the other hand it has been refuted by Friedman1' and Wood and coworlters.s$12 Stern and c o ~ o r k e r s ~ 3measured ~~4 the heat of mixing of HC1-NaCl-H20 and HC1-KC1HzO from 0 to 40". Both mixtures showed a decrease in absolute magnitude in the heat of mixing as tempera(1) This study was aided by a grant from the Office of Saline Water, U . S. Department of the Interior. (2) See, for example, J. L. Kavanau, "Water and Solute-Water Interactions," Holden-Day, Inc., San Francisco, Calif., 1964. (3) R. H . Wood and R. W. Smith, J. Phys. Chem., 69, 2974 (1965). (4) T. F . Young, Y . C. Wu, and A. A. Krawete, Discuss. Faraday Soc., 24, 37 (1957). (5) R. H . Wood and H. L. Anderson, J. Phys. Chem., 71, 1869 (1967). (6) R. H. Wood and H. L. Anderson, {bid., 71, 1871 (1967). (7) H.S. Frank and M. W. Evans, J. Chem. Phys., 13, 507 (1945). (8) H . S. Frank and W. Y. Wen, Discuss. Faraday Soc., 24, 133 (1967). (9) Seef for example: (a) G. E. Walrafen, J. Chem. Phys., 36, 1035 (1962); (b) 0 . D. Bonner and G. B. Woolsey, J . Phys. Chem., 72, 899 (1968). (10) G. Scatchard, J. Amer. Chem. SOC.,83, 2636 (1961). (11) H. L. Friedman, J. Chem. Phys., 32, 1134 (1960). (12) R. H. Wood and H. L. Anderson, J. Phus. Chem., 70, 992 (1966). (13) J. H. Stern and A. A. Passchier, ibid., 67, 2420 (1963). (14) J . H. Stern, C. W. Anderson, and A. A. Passchier, ibid., 69, 207 (1965). Volume 74, Number 7
April 8, 1970
1456
HENRY L. ANDERSON AND LINDAA. PETREE
ture increased, although the former did so at a much larger rate than the latter. Stern and coworkers discussed their heats of mixing data in terms of deviations from the Brgnsted specific ion interaction principle. They interpreted the decrease in the heats of mixing as temperature increased as a reduction in the deviation from the specific ion interaction principle approaching the Brgnsted prediction of a zero heat of mixing. They recognized that 40” was not a high enough temperature to make a prediction as to whether or not the heat of mixing mould actually go to zero. The present work was designed to go to higher temperatures to see if in fact the heat of mixing does decrease to zero as temperature is increased. Experimental Section
Calorimeter. Common ion heats of mixing are quite small (10-40 cal/kg of solvent for the mixtures in this study) and require a solution calorimeter having a sensitivity in the microdegree range. Because of the required sensitivity it was decided to set an upper limit of 80” in this initial investigation. A schematic diagram of the calorimeter constructed for this research is given in Figure 1. The “double calorimeter”16 construction has an advantage in that twice as many experiments can be performed during a given run. This is especially advantageous when experiments are carried out at temperatures greatly different from ambient temperatures, because considerable time is spent in bringing the calorimeter to the desired experimental temperature. Although the calorimeter incorporates many of the ideas presented e l ~ e w h e r eits , ~configuration ~~~~ from the point of view of high-temperature utility and simplicity warrants some comment. The calorimeter consisted of a 240-ml dewar flask1* (I) cemented with silicon rubber cement to a brass collar (D). The collar was bolted to a brass lid (C) and sealed with a rubber “0 ring” (I
