NOTES
548
FIGURE I P Y R R O L E - CC14 I N F R A R E D F R E E N-H 3498 C M - '
--
470 c. 33*c. 12'
c.
association were computed from the equation X+Y+-XY K = - CXY x
= y = xy = CO , =
C,o = C, = Cy =
C,, =
CdC, pyrrole pyridine or substituted pyridine pyrrole-pyridine complex initial pyrrole concentration initial pyridine concentration equilibrium pyrrole concentration equilibrium pyridine concentration equilibrium complex concentration
All of thc data are given in Table I. Because of experimental limitations, a large number of determinations were made to minimize errors. Table I1 gives thc thermodynamic data coinputed by least squares from log K us. 1/T plots of the data in Table I. Our solutions are sufficiently dilute to assume ideal behavior.
Vol. 66
solvent effect with our CCI,. Comparing with the data of Mortimer and Laidler,s which also are given in Table 11, we see that the AIIO values follow the same trend. Adding a methyl group increases the basicity of the pyridine and makes AHo, AFO, and ASo all more negative. In our study, with CC1, solvent, it is difficult to imagine any solvent clustering effects and we therefore interpret the less negative AHo and ASo of 2,G-dimethylpyridine as due entirely to a steric effect which overshadows the inductive effect of the added methyl group. This interpretation supports the ideas of Mortimer and Laidler.* In their aqueous solutions it is rioted that ASo is very much larger for 2,G-dimethylpyridim; in aqueous solution, ASo is positive since So of HsO+ is very small compared to So for BH+. Solvent clustering is greatly decreased in the dimethyl compound, thus giving it an abnormally high entropy and enthalpy. The predicted AHo for the dimethyl compound, if no steric or solvent effects were operating, would be: (a) aqueous, -2(6.95-5.70)--5.70 = -8.20 (observed is -G.13 or 75% of the predicted value) ; (b) CCL with pyrrole, -2(3.8-3.2) -3.2 = -4.4 (observed is -3.4 or 89% of the predicted value). We therefore can conclude that our values for the 2,G-dimethylpyridine reflect a steric effcct, only, and that the work of Mortimer and Laidler for aqueous pyridines depends on solvent exclusion. Acknowledgment.-The authors appreciate helpful discussions with Dr. R. G. Inskeep. Financial assistance from the Research Corporation is gratefully acknowledged. SOLUBILITIES OF SOME STRONG ELECTROLYTES I N T H E HYDROGEN PEROXIDE-WATER SYSTEM. IT. RUBIDIUM. AND CESIUM NITRATES' BY MARTINE. EVER HARD^
AKD
PAUL31,GROSS,J R . ~
CobB Chemacal Laborator?l OJ the Untverszty o i Vargznaa, Charlotlesvalle, Va., and the Department of Chemzstrv of Wake Foresl College, Winston-Salem, N . C. Received Julu 19, 1961
The solubilities of the smaller alkali inetal iiitrates in the mixed solvent hydrogen peroxidemater have been reported by Floyd and Gross.4 They found that for LiN03, NaX03, and KN03 that the solubilities in water-rich solutions deTABLE11" THERMODYKAMIC VALUESFOR PYRROLE-PYRIDINES creased with increasing cation size and that in the hydrogen peroxide rich solutions, the solubilities ALLVALUESIN KCAL.PER MOLE increased with increasing cation size. In addition Pyridincs AH2d AP298' A&d Pyridine - 3 . 2 ( - 5 70) -0 54(-7.12) - 8 g(4-4.76) LiS03 showed a discontinuity in the curve of mole fraction hydrogen peroxide in the solvent us. molal Z-Methylpyridine -3 8(--8 95) - .58(-8.13) --10.8(+3 95) solubility a t low hydrogen peroxide concentrations 2,6-Dimethylwhile the KN03 solubility curve showed a dispyridine -3 4(-6.15) - 68(-9.17) - 9 2(4-10 11) continuity in the hydrogen peroxide-rich solutions. * Values in parentheses are for aqueous solutions, I3 f H30+
BHf
+ H20,by Mortimer and Laidler.8
It is interesting to compare our values with other literature data. Happe' gives AHa = -4.3 kcal./mole and ASo = -8.0 for pyrrole-pyridine in cyclohexane solvent. We would not expect such a large difference from our values due simply to ( 8 ) C. T. Mortimer and K. J. Laidler, Trans. Faraday Soc., 66, I731 (1959).
(1) This work received support from the Ofice of Ordnance Researoh, U. s. Army. The expwlmental nork presented in this paper was carried out at the University of Virginia 8 s a part of the thesis submitted by Martln E. Everhard in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Graduate School of the University of Virginia. June, 1960. (2) Philip Francis du Pant Fellow. (3) To whom inquiries should be directed a t Wake Forest College. (4) J. D. Floyd and P. M. Gross, Jr.. J. A m . Chem. Soc., 77, 1435 (1955).
March, 1962 Subsequent data obtained as a part of this series of studies showed that these discontinuities corresponded to the changes in phase LiS03.3H20LiN03 and KN03-2KNO~.I-I~02.6This behavior indicated that cation size was an important factor in the solubility relations with varying hydrogen peroxide concentration. To examine further this possible role of the cation size, the solubilities of RbN03 and CsNOs were examined as well as the nature of the solid phases in equilibrium with these systems. Experimental The hydrogen peroxide6 used waa 95-98% by weight except the most concentrated solutions, which were obtained by distillation using the method of Gross and Taylor.? The RbNO8 was prepared from Rb2C03by a modification of the method of I'uschin and Radoicic.8 The final material was repurified to the constant melting point of 313' (cor.) of Haigh.0 The solubility i n water agreed with Jones.lo Recrystallized C.P.grsde CsNOI was used. The solubilities were determined by the method of Floyd and Gross.4 Castor and Basolo's11 procedure, as modified by Turner,&was used t o analyze the solid phases.
Results Figures 1 and 2 show the relation of the molal solubilities of RbN03 and CsN03 to the mole fraction HzOZin the solvent. The discontinuities in the RbN03 curves correspond to the changes in phase RbN03-7RbN03.3Hz02and 7RbN03.3HzOz2RbNOa-HzOzas determined by analysis of the solid phases on each side of the discontinuities. The formation of the above hydroperoxidates and the increase in solubility of the salts with larger cations in hydrogen peroxide-rich solutions indicate preferential solvation of these ions by Hz02 rather than by HzO. Conversely, the formation of hydrates and lower solubility in hydrogen peroxide-rich solutions of the smaller cation salts indicate preferential solvation of these ions by water. The deviation of the molal solubility, M ' , of the alkali nitrates in H202 from that in HzO ( M H ~ minus o ~ M13,o) a t 25' was found to follow the straight line M' = 33.5, - 39.7 (k0.03 in M ' ) , where T is the radius of the cation. CsN08, however, did not fall on the line, which probably is due to the lower charge density of the cesium ion. The caesium ion would have an effective radius of 1.53 A. if it would lie on the above line. These solubility results cannot be explained in terms of the dielectric constant, E, of these two solvents since E H ~ Ois~ about 10% lower than EH~? a t the temperature of this work. However, it is possible that the smaller ions depress EH,O~ more than the larger ions, resulting in a 1ow:r solubility of these ions as first suggested by Akerlof and Turck. Undoubtedly changes in the activity coefficients occur. These changes will be quite (5) J. W. Turner, Ph.D. Thesis, University of Virginia, Charlottesville, 1957. (6) Donated by the Becco Division of the F. M. C. Corporation. (7) P. M. Gross, Jr.. and R. C. Taylor, J. A m . Chen. Soc., 7 2 , 2075 (1950). (8) N. A. Pusohin and M. Radoioic, 2. anorg. I . allgem. Chem., 42, 233 (1937). (9) F. L. Haiah, J. A m . Chem. Soc., 34,-1144(1912). (10) E. M.Jones, J . Chem. Soc., 93,1743 (1908). (11) W. 9. Castor, Jr., and F. BaRolo, J . A m . Chem. Soc., 76, 4808 (1953). (12) G.AkerlBf and H. E. Turck. ibid., 67, 1746 (1935).
549
NOTES
SOLUBILITY AT
!547 6
.
O'C
0
15.C
0
25'C
D
o
08
MO-E
FRACTION
H&
I N SOLVENT.
Fig. 2.
significant especially with the formation of new species as occurs in the LiN03, KNOe, and RbT\'Oa systems. However, since the activity coefficient and the dielectric constant changes of these solutions have not been measured, it is difficult to assess these effects properly. To determine the extent and the nature of the solvation in these solutions, the vapor pressure, vapor composition, and liquid composition of these solutions have been measured over the full range of HzOzconcentrations. These results will be reported in a later paper. They support the contention that the smaller ions are preferentially solvated by the water molecules and the larger ions by hydrogen peroxide molecules. It is interesting to note that the solubilities of the alkali metal salts in the nitrogen analogs of H 2 0 and HzOZ,ammonia and hydrazine, show decreasing solubility with increasing ion size in both N2H, and NH3.l3-I6 Since the N-N distance in hydrazine is about the same as the 0-0 distance in Hz02, one might expect to observe the same solubility trends in the hTzH4-NH3 system as in the H~OZ-HZO system. The observed difference could be due to the weaker ion-dipole interactions of the nitrogen compounds as compared to those that occur in the oxygen solvents. 113) H. Hunt and L. Boncyk, ibid., 65, 3528 (1933). (14) H. Hunt, ibid., 64, 3509 (1932). (15) L. de Bruyn, Rec. tvau. chin., 18. 174 (1896). (16) T. W. B. Welsh and H. J. Broderson, J . A m . Chem. Soc., 37, 816 (1916).
