NOTES
2264
their activity coefficients are equal or remain constant. Consequently, a “mixed” formation constant (terms in both activity and concentration) is obtained. The series of experiments were so designed that data could be obtained from the equilibrations of aqueous phases with benzene solutions containing TNOA and TTA which would test the relationships predicted by eq. 6. In all cases, conditions were selected so that a t equilibrium the free TTA concentration was maintained constant a t 0.1 114. Consequently, from a plot of (H+), - (€I+) us. [H+][C1-]((R3N)t (H+) - (H+),} a straight line relationship with a zero intercept should be obtained if the postulated product R3NHClHT is of significance. If such is the case, then utilizing the known formation constants for R3NHT and R3NHC1and the activity of the unassociated TTA, the formation constant, KCI,T,can be calculated from the slope of the line. The appropriate terms for the plot were calculated from the known composition of the solutions, (R3N)t and (H+)t, and from the measured aqueous hydrogen ion concentration, (H+),after equilibration. The indicated terms were calculated from the data presented in Table I. Activities were calculated for the aqueous hydrochloric acid concentrations from the known activity coefficient^.^ Utilizing the expression of King and Reas5 the activity of a 0.1 M TTA solution was calculated as 0.092. For the systems where aqueous phases contained 2 M LiCI, the activity coefficient for hydrochloric acid was taken as unity.* The resulting plot is shown in Fig. 1. I n accordance with theory, a straight line relationship is obtained with a zero intercept. From the slope of the straight line and the known formation constants, a value of (2.1 =t 0.4) X l o 5 was calculated for KCI,T. The formation constant for the organic phase reaction
+
R3NHC1
+ H T = R3NHClHT
(7)
was calculated by dividing KCI,Tby K c l . The value that was obtained is 16 5. The experiments which were performed permitted a test of the suggested reaction under a wide variety of conditions. The analytical amine concentration was varied between 0.002 and 0.04 M , the free acid concentration from O.OOO8 to 0.08 M , and the chloride concentration from 0.0008 to 2 M . The TTA concentration was maintained at 0.1 M . The formation constant was measured under conditions when each of the amine species participated as either a minor or major component of the sohtion. Consequently, it has been demonstrated that the extraction of hydrochloric acid by benecne solutions of TTA and TNOA can be completely accounted for by the presence of three species, R3NHC1,R3NHT,and R3XHC1HT. This knowledge of the nature and the extent of formation of these species has been applied for an elucidation of the synergistic effect exerted by
*
( 4 ) 13. S. Harned and B. B. Owen, “The Physical Chemistry of Elertrolytic Solutions,” Second Ed., Reinhold Publ. Corp.. New York, N. Y . , 1950. (5) E. L. King and W. H. Reas, J . Am. Chem. SOC., 73, 1806 (1951).
Vol. 66
TNOA on the solvent extraction of thorium by TTAS6 (6) L. Newman a n d P. Klots, J . Phys, Chem., in press.
THE THERRIODYKARIIICS OF AQUEOUS ELECTROLYTE MIXTURES AT ELEVATED TEMPERATURES. THE SOLUBILITY OF SILVER SULFATE IN KK03-K&304, KzS04MgS04, AND K2SO4-HZS04 MIXTURES BY M. H. LIETZHEAND R. W. STOUGHTON ChemiStTy Division, Oak Ridge National Laboratory,l Oak Ridom, Tennessee Received April 13, 1062
In a series of previous papers2 the solubility of Ag2S04in a variety of electrolyte solutions has been investigated. It was shown that single parameter empirical expressions of the type
(1) could be used to describe the variation of the solubility product of AgzS04 over a wide range of temperature and ionic strength. I n eq. 1 so is the solubility of Ag2S04 in pure water, ST is the appropriate Debye-IIuckel slope, I the ionic strength of the solution, and A the adjustable parameter. In all cases best agreement between calculated and ohserved solubilities was obtained when each single parameter A i was assumed to be temperature independent and to be either ionic strength independent or to decrease slowly with increasing ionic strength. The justification for using eq. 1 without a linear term has been discussed previously. In the present paper the treatment of the solubility of Ag2S04 in single electrolyte systems has been extended to the solubility in three electrolyte mixtures : (1) KK03-K2S04,(2) K2S04-MgS04,and (3) K2S04-H2S04. In each of the three systems the total ionic strength of the solubility medium was held constant a t two different values, while the ratio of the two components was varied. Weight concentrations were used in reporting the data, and the Debye-Huckel limiting slope at any temperature was made density dependent by multiplying its value on the molal scale by the square root of the density of water a t that temperature. Experimental All solubility measurements were performed with the same techniques described previou+y3 and were carried out in the temperature range 90 to 175
.
Results and Discussion The values for the solubility of AgzS04 in each system as a function of temperature were fitted by the method of least squares to an equation (1) Operated b y Union Carbide Corporation for the United States Atomic Energy Commission. (2) M. H. Lietzke and R. W. Stoughton, J. Phys. Chem., 6 3 , 1183, 1186, 1188, 1190, 1984 (1959); 64, 133, 816 (1960). (3) M. H. Lietzke and R. W. Stoughton, J. A m . Chem. Soc., 7 8 , 3023 (1956).
NOTES
Nov., 1962
I of the
2265
TABLE I COEFFICIENTS OF EQUATION 2 FOR THE SOLUBILITY OF Ag2S04IN mixture mKNoI mKIsO4 a0 x 102
1.0
1.416
0 .1 .16 .24 .333 0 .3564 .472
1.0 0.7 .5 .28 0 1 416 0,3484 0
1,437
4.31016 3.47859 2.88124 1.90299 1.41304 5.84143 -0,202929 1 59375
ELECTROLYTE MIXTURESSTUDIED x 104 a2 x 108 Wit" x
- 1.57687
6.41420 3.57152 3.12836 4.43650 5.47167 4.86857 7.32640 5.67500
.333 .25 .167 .OS33 0 .479 ,2395
0
.0625 .125 .1875 .25 0
.I796 .3592
0 ma2so4I'
1.41304 0.642984 O.OO169903 1.42525 1.27142 1.58357 1.26804 1.22750
-
-0.875239 -1.32865 - 1.86183
5.47167 5.86084 6.99906 4.68734 5.00573 5.71667 7.94620 5.38381
-
-0.470478
- 1.49270 -1,16191
-
-0.883321 9.58466 2.63613 ,2507 .2500 1.007 13.0980 -3.06707 -3.26708 .5000 .5000 2.000 a Standard error of fit. b Formal molality based on considering HzS04 as a 1-1 electrolyte. s = a0
+ + ad2 Ult
0.16 .92 .43 1.09 0.27 0.15 2.20 0.30
- 1.06849 - 1.24572
~ K ~ S O ~
(2)
The resulting equations were solved at 25 O intervals from 25 to 200' in the case of the KN03K2S04system (IKNO~ IK~SO, = 1.0), where data were obtained over the entire temperature range, and from 100 to 175' in the other systems. The coefficients and the standard error of fit for eq. 2 describing the solubility of Ag2S04 in each system as a function of temperature are given in Table I, while a typical set of solubility curves for one of the systems ( 1 ~ ~I K 0 ~~S = O ,1.0) is shown in Fig. 1. The other families of curves were similar. Since in previous work2it was demonstrated that temperature independent values of the parameter A in eq. 1 could be calculated such that the difference between observed and calculated solubilities would be minimized for systems in which Ag2SO4 was dissolved in solutions of a simple electrolyte, it seemed of interest in the present work to calculate values of the parameter A for the mixtures studied to determine whether a simple combining relationship existed between the values of A in the simple systems and in the mixtures. In carrying out the calculations for the KNO3-KZSOh and K2S04MgS04 mixtures, values of In S and v'j in eq. 1 were computed a t 25' intervals for each system using the smoothed values of the solubility of Ag,SO4 as obtained from eq. 2 and the appropriate molalities of the components of the mixture. Then the value of A was determined in each case between 100 and 175' by selecting the value which equated the right and left sides of eq. 1. The A parameters so obtained for each system are given in Table 11. The values shown are the mean values calculated between 100 and 175'. Additional calculations mere performed in which a linear term was added to eq. 1, viz., BI. In these calculations A was set equal either to 1.0 or to 1.5 and the value of B determined such that the right
0
.
1
1
0
~
T
,
~
,r
108
2.12 1.37 2.57 1.00 0.16 0.48 1.86 0.22
-0.532662 .284457 .701309 .875239 - .697144 - 1.45714 -0.473332
mMgso4
m~ps04
1 .o
THE
1.39 0.68
,
,
,
,
+
2000 175O
\\\-
/
+
150' 125O
0.060L
0.025
0
'
0.2
0.4
0.6 If
Fig. 1.-The
0.8
1.0
K 3 o*
solubility of Ag2SO4 in KNOa-KdOd mixtures where I K N OI ~ I G S O= ~ 1.0.
+
side of eq. 1 augmented by the linear term was equal to the left side. The values of B so obtained showed a much greater dependence on temperature than did the values of A obtained without the linear term. Since the latter calculation (involving the BI term) offered no advantage over the former, no attempt was made t o correlate the B values with temperature and ionic strength fractions.
NOTES
2266
Vol. 66
The observed solubility of AgzS04in each system as a function of temperature was compared with values calculated using the average A values given in Table 11. I n all cases the agreement was very good-within a few tenths of a per cent in most cases and within 201, in the remainder. Hence the behavior of the A values in the electrolyte mixtures was very similar to the behavior of the A values in the individual components.
value for the mixture lay between values computed from the A values of the simple systems by eq. 3 and 4. Since a simple general combining relationship based on the A values could not be found it seemed of interest to try a direct correlation of the solubility of Ag,S04 in the electrolyte mixtures with the solubilities in the pure component solutions using eq. 5. However, the coefficient s12was not con-
TABLEI1 A PARAMETERS FOR EQUATION 1
slfl Szf2 SlZflfZ (5) stant but varied quadratically with either fi or f z and quadratically with temperature. Thus a direct correlation of the solubilities is neither simpler nor more complicated than a correlation through an activity coefficient function. Acknowledgment.-We wish to thank Stanley Anderson and Sarah Ledford for carrying out the experimental solubility measurements and Dr. W. D. Larson for interesting discussions during the early phasea of this work.
Total I
1.o
1.416
mKNOa
1.0 0.7 .5 -28 0 1.416 0.3484 0 mKZs04
1. o
1.437
0.333 .25 .167 ,0833 0 0.479 .2395 0
A
Std. dev. from the mean
0,991 .994 .956 .815 .656 0.997 ,811 .630
0.004 ,018 ,030 .014 .002 0.012 .021 .000
0.656 .794 ,834 ,892 .940 0.634 ,815 .885
0,002 .006 .006 .003 ,006 0.000 .043 ,004
mKzso4
0 0.1 .16 .24 * 333 0 0.3564 0.472
SM
m~~so4
0 0,0625 .125 .1875 .25 0 0.1796 ,3592
=
+ +
THE ACTIVATION ENERGY FOR THE DISPROPORTIONBTIOlJ OF THE HOP RADICAL IN ACID SOLUTIOXSl BY BENON H. J. BIELSKIAND EIICHISAITO Chemistry Department, Brookhaven Natzonal Laboratory, Upton, L. I., N e w Yorlc Recezoed Aprzl 16, 1961
The e.p.r. spectrum of the HOz radical has been reported independently by several research groups.2-b The spectra show good agreement with respect to line shape and line width, although each investigation produced the radical by a different method. I n the case of the solubility of Ag2S04in K2SOr By means of a previously described flow techH2S04 mixtures the calculations were performed n i q ~ eHOz , ~ was generated by reaction of hydrogen essentially as described previously2 for the solu- peroxide with ceric sulfate. This method has the bility of AgzSOd in HzS04 solutions except that the advantage of producing the radical in a relatively ionic strength term included the concentration of high concentration in the liquid phase. K2S04. The denominator parameter in the term Experiment a1 correcting the bisulfate acid constant for the ionic Chemicals.-The solutions were prepared from triply strength of the solution was taken as 0.4, consistent distilled water and reagent grade chemicals. All solutions with previous calculations.2 At all temperatures were in 0.8 A7 sulfuric acid. The concentrations of the ceric from 100 t o 175’ the observed and caIculated solu- sulfate solutions were determined spectrophotometrically. bilities of AgzS04agreed to within 1% with a mean The hydrogen peroxide concentrations were determined by A value for I = 1.0 of 0.922 and for I = 2.0 of the ceric sulfate and iodide method.6 Apparatus.-The experimental equipment consisted of a 0.832. Varian e.p.r. spectrometer, a flow system, and a thermostat. An attempt was made to relate the A values The flow system had a double-jet mixing chamber and a shown in Table I1 for each system. In no system flow tube with an inner diameter of 0.1 cm. The diameter of tube was checked for uniformity with a manganous was it possible to relate the A value calculated for a the flowsolution by sliding it up and down in the microwave mixture with the A values calculated for the corre- sulfate cavity of the spectrometer. The maximum deviation from sponding pure systems by taking a linear combina- tion of the latter values weighted according t o the (1) Research performed under the auspices of the U. S. Atomic ionic strength fractions fi of the appropriate elec- Energy Commission. (2) 3. Kroh, B. C. Green, a n d J. W. T. Spinks, J . A m . Chem. Sac., trolytes; i.e. 88, 2201 (1961). (3) L. I. Awramenko and R. V. Kolesnikowa, DokZ. A k a d . Nauk A m i x t u r e = Alfi Adz (3) S S S R , 140, No. 5 (1961). (4) E. Saito a n d B. H. J. Bielski, J . Am. Chem. Soc., 88, 4467 did not hold. However, in the KKO3-K2SO4 (1961 1. system ( I = 1.0) the relationship ( 5 ) J. Kroh, B. C. Green, and J . W. T. Spinks, Can. J . Chem., 40, m ~ m 4
mH*SOaU
1.007 0,2500 0.2507 0.922 0.036 0,5000 0.832 0.5000 0.013 2.000 Formal molality based on considering Hi304 as a 1-1 electrolyte. Q
+
+
+
Alfi Adz -41~Az~fifz(4) held very well. I n the remaining systems the A Amixture =
413 (1962). (6) I. M. Kolthoff and E. B. Sandell. “Textbook of Quantitative Inorganic Analysis,” 1st Ed., The Maomillan Co., New York, N. Y., 1943.
