574
NOTES
sidering also the greater bulk of the phenyl group as compared with the hydrogen atom, it mould be expected that, in the approach of the attacking molecule to the nitrogen atom of the amine, oxanilic acid should encounter more steric hindrance than would oxamic acid-in other words, for the reaction in a given solvent, the AS* for the reaction should be less for oxinilic acid than for oxamic acid. Both these predictioiis are seen to be confirmed by the data in Table 111. In the studies of the decarboxylation of oxamic acid in aniline and in o-toluidine3 it was shown that the presence of a methyl group ortho to the amino group has two effects, namely, (1) a +I effect which increases the electron density on the nitrogen atom of the amine giving rise to a decrease in AH*, and ( 2 ) an ortho or steric effectBwhich hinders the approach of the oxamic acid to the nitrogen atom, resulting in a decrease in AS*. Analogy suggests that oxanilic acid should behave in this respect in a manner similar to oxamic acid. The experimental data in Table I1 confirms this expectation. It is seen that, on passing from aniline to o-toluidine, a progressive decrease in both AH* and AS* takes place for both the oxamic acid and the oxanilic acid reactions. A, further decrease in the parameters takes place in the oxanilic acid reaction on going from o-toluidine to o-ethylaniline. For each solvent it will be noted, also, that both AH* and AS* are lower for the decarboxylation of oxanilic acid than for that of oxamic acid. This is due to the fact that, in the first place, the effective positive charge on the coordinating carbonyl carbon atom of oxaiiilic acid is greater than it is on that of oxamic acid, and, secondly, the molecule of the former is the more comp1e.u. These circumstances leave little doubt but that oxanilk acid decomposes in polar solvents by the same mechanism as does oxamic acid, oxalic acid and malonic acid. I n aniline solution, the AS* for the decarboxylation of oxanilic acid is 18.5 e.u. smaller than it is for that of oxamic acid (see Table 111, lines 3 and 4). This difference in AS* is commensurate with the increascld steric hindrance which would be expected from the substitution of a phenyl group for one of the amide hydrogen atoms of oxamic acid. Since oxaiiilic acid is considerably larger than oxalic acid the A S * for the reaction in aniline would be expected to be larger for oxalic acid than for that of oxanilk acid provided the same number of molecules of each acid were iiivolved in the solvation step. We find, however, that, in aniline, the AS* for the oxalic acid reaction is actually 30 e.u./mole less than it is for that of oxanilic acid. I t is well known that the dicarboxylic acids in solution associate through hydrogen bonding to form a cluster composed of at least 3-4 m~lecules.~Wheii one of the hydroxyl groups of a dicarboxylic acid is replaced by sorne other group such as the S-phenyl amide group association evidently can no longer take place past the dimer stage. In view of the great differences in the activation entropies of (8) L. 1’. Hamrnett, ”Physical Organic Chemistry,” 3IcGraw-Hill Book Co., Inc., Ncsw York, N. Y., 1940, p . 204 (9) \V. Huchel, “Tlieoretical Principles of Organic Chemistry,” Vol. 11. Elswier Pub Co , Kew York. N Y 1958. p. 329 el seq.
Vol. 65
these two reactions it may be deduced that, in the case of oxanilic acid, probably only a single molecule coordiiiates with the solvent, whereas, in the case of oxalic acid, a supermolecule consisting of an association cluster of 3-4 single molecules coordinates. h similar interpretation has been advanced for the oxamic acid data.3 In Table I11 it will be observed that, at 140°, oxalic acid and oxanilic acid decompose in aniline a t about the same rate. This evidently is due to the circumstance that, even though the AH* for the decarboxylation of oxalic acid is nearly 10 kca1.l mole less than it is for that of oxanilic acid, the improved entropy factor in the case of oxanilic acid exactly compensates for this disadvantage. Further work on this problem is contemplated. Acknowledgment.-The support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged. HEAT OF NEUTRALIZATION OF STRONG ACIDS BY STRONG BASES IN MIXED WATER-DIOXANE SOLUTIONS’ BY HIDEHIKO KIDOAND W. CONARD FERNELIUS Department o j Chemistrv. The Pennsylvania Stale 1-niuersitg, Cniversitv Park, Pa. Received August 4, 1960
I n determining format,ion constants of complexes of organic ligands with met,al ions, it is common practice to use the mixed solvent wat’er-dioxane t,o secure the necessary h careful analysis of this mixed solvent has shown that measurements made with ordinary glass and calomel electrodes and a pH met,er can be used to calculate t’hermodynamic dissociation constants of weak electrolytes and formation constants of complexes.* This signifies that even a 75 volume dioxane-water mixt’ure (mole fraction of dioxane = 0.388) is functioning essent,ially as a “water” solvent. Indeed, the activity of water a t 25’ in a 757, dioxane solution differs by only 10% from that of pure water.9 Xevertheless, many are concerned that the state of hydration of ions in t’his (1) This investigation was carried out under contract AT(30-11-907 between The Pennsylvania State University and the U. S. Atomic Energy Commission. (2) M. Calvin and K. W. Wilson, J . Am. Ciiern. Sac., 67, 2003 (1915). (3) D. P. hIellor and L. E. hlaley, Nature, 159, 370 (1937); L. E. Jlaley and D. P. Mellor, Australian J . Sci. Research, 2 8 , 92 (1949). (4) H. Freieer, R. G. Charles and W. D. Johnston, J . Am. Chem. Sor., 74, 1383, 1386 (1952); W. D. Johnston and H. Freiser, A n d . Chem. Acta, 11, 201 (1954); T. R. Harkins and H. Freiscr, J . A m . Chem. Soc., 77, 1374 (1955); 78, 1143 (1956): 80, 1132 (1958); G . E. Cheney, 11. Freiser and Q. Fernando. ibid., 81, 2611 (1959). ( 5 ) F. Bnsolo, P. T. CIien and R. Kent Murniann, ibid., 76, 950 (1954). (6) L. G. Van Uitert. C. G. Ilaas, JV. C. Ferneliiis and B. E. Douglas, ibid.,75, 455 (1953); L. G. Van Uitert, W. C. Fernelius and B. E. Douglas, ibid., 75, 457, 2736, 2739 (1953); L. G. Van Uitert and W. C. Fernelius, ibid., 75, 3802 (1953); 76, 375 (1954); C. M. Callahan, W. C. Fernelius and B. P. Block, Anal. Chim. Acta, 16, 101 (1957); D. F. Martin and R‘. C. Fernelius, J . Am. Chem. Soc., 81, 1509 (1959); B. B. Martin and W. C. Fernelius, ibid., 81, 2342 (1959). (7) H. Irving and H. Rossotti, Acta Chem. Scand., 10, 87 (1956). (8) L. G. Van Uitert and C. G. Haas, J . Am. Chem. Soc., 75, 461 (1953); L. G. Van Uitert and W. C. Fernelius, ibid.. 76, 5887 (19.54). (9) P. Ilovorka, R. -4. Scliaefer and D. Dreisbach. J . A m . Chem. Sue., 58, 2264 (1930).
KOTES
March, 1961 solvent may be different from that in water.'O Determinations of the heat of neutralization of strong acids by strong bases in solutions of dioxane-water varying from 0 to 75 volume yo dioxane show a variation of only 2.4%. The condition of solvation of the proton and hydroxide ion must thus differ only by very little in the two solvents. Emerimental The apparatus and procedures for enthalpy titrations are essentially those described hy Jordan and Alleman." For each run, 100 ml. of 0.01 M sodiuni hydroxide in the given solvent was titrated with 1. O M hydrochloric acid in the same solvent. The results assenibkd in Table I are the avrrage of 5-9 determinations a t each mole fraction of dioxane.
TABLE I HEATSOF NEUTRALIZATIOI. I N WATER-DIOXAXE ---Dioxane--Vol C L
0 25 50 75
Mole fr
0 0 0 0
00 066 li4 388
4 H neut.
13 13 13 13
57 65 70 89
f0 f0 f0 f0
02 14 04 04
(10) Private conimunications. (11) J. Jordan and T. G. Alleman, 9naZ. Chem., 29, 9 (1957).
SOLUBILITY AND THERMODYNAMIC FUNCTIONS OF ETHYLENE IK DIETHYL SULFATE BY A. M. TRUCHARD, H. G. HARRIS ANJ) D. M. H~MMELBLALI Department of Chemical Engineering, The Unzverszty of Texas, A u s t t n
=
log
575 3.04666 - o.132845 5.9302
x
10-4
was calculated by the least squares technique to give an excellent fit (maximum deviation from 0 to 80" was 1.51Yo). With the aid of this equation, the partial molal heats of solution were calculated as
and the partial molal entropies of solution were calculated as
The calculated thermodynamic values are listed in Table I. Womenclature and discussion related to equat,ions2 and 3 is in reference 1. TABLE I HENRY'SLAWCONST.4NTG AND THERMODYNAMIC DATA FOR ETHYLENE - (EL - Bo), tal./ - Bo), %(""-_) - (RLO 1.
"C.
mole f r I n EtnSOh
cal./g. molea
In EtL304 In H20
( g . mole)('K.) b I n EtsSOa In H20
2430 5260 17.9 36.4 0 68.71 2300 4390 16.8 33.3 20 92.24 2270 4170 16.7 32.6 25 98.73' 30 105.1 2240 3880 16.6 31.7 2180 3340 16.4 29.9 40 117.3 2080 2100 16.1 26.0 60 144.2 1995 850 15.9 22.3 80 172.8 "infinite dilution." * Standard State: a Standard State: mole fraction = 1.0. Calculated.
12, Texas
Reeelired October 84, 1960
As part of a study of the reaction of olefins with sulfuric acid, a thorough investigation of the solubility of ethylene in diethyl sulfate has been made. From the solubility data the partial molal heats and entropies of solution have been calculated. (a) Apparatus.-The experimental apparatus consisted of two calibrated volumetric bulbs which contained approximately equal volumes of gas after 200 ml. of diethyl sulfate had been added to the larger bulb. The flasks were immersed in a thermostat (good to f0.05O) and attached to a mercury manometer. The system was evacuated by prolonged pumping, and ethylene at about 2 atm. pressure was added to the empty smaller flask. Pressures in the system were measured, and then the ethylene was admitted to the bulb coiitaining the diethyl sulfate. The system was allowed to reach equilibrium while being stirred. For each temperature successive additions of ethylene were brought into contact with the diethyl sulfate a t pressures ranging from 50 to 1300 mm. Plots of corrected partial pressure of ethylene us. mole fraction ethylene dissolved in diethyl su!fate could be fit with excellent precision by straight linrs. T h r diethyl sulfate wap Eastm:tn practical grade purified by washing with Xa,CO, solution and drying with CaC11. A t temperatures much higher than 100" diethyl sulfate begins to degrade so the maximum temperature reported is 80".
(bj Results.--Tahle I shows the Henry's law constants from 0 to 80" computed as X = p (atm.)/x (mole fraction). Since the logarithms of the Henry's law constants were not exactly a lirieur function of temperature, a curve to fit the values of x of the form
Discussion The reliability of t,he Henry's law constants was excellent since the standard deviation for X for any temperature was less than 0.01. The accuracy of the apparatus was tested by measuring the solubility of ethylene in water at 0'. X was determined t'o be 5.28 X lo3, which is 4.5yoless than the value of Winkler.2 Winkler's value may be high, however, since the data of Davis and McKetta3 and Bradbury, McKult8y,Savage and McS ~ e e n e y ,when ~ extrapolated to lower tempera tures, yields a lower value (less than 5 X lo3) for
x.
The partial molal heats and entropies of solution of ethylene in diet,hyl sulfate do not change as much mit,h temperature as do corresponding values for ethylene dissolved in water, which have been calculated from the data in reference 5, although the trend with temperature is the same. The values of (& - SG)in diethyl sulfate are onehalf as negative as those in water and correspond closely t'o the values shown in Table I1 for other non-polar compounds. The values in Table I1 have been computed from the data of Horiuti6 (1) D. hl. Hirnmelblau, J . Phya. Chem., 63, 1863 (1959). ( 2 ) L. W. Winkler, Z. phusik. Chem.. 55, 350 (1906). (3) J. E. Davis and J. J. McKetta, J . Chem. Eng. D z t a , 6 , 37-1 (1960). (4) E. J. Bradbury, D. RlcNulty, R . L. Savage and E. E. McSweeney, I n d . Eng. Chem., 44, 211 (1952). (5) D. M. Hirnmelblau and E. Arends, Chem. Iny. Tech., 31, 791 (1959).
