JOHNH.
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SAYLOR -4ND
RUBINBATTINO
Vol. 62
THE SOLUBILITIES OF THE RARE .GASES IN SOME SIMPLE BENZENE DERIVATIVES BY JOHNH. SAYLOR AND RUBINBATTINO Contribution from the Department of Chemistry, Duke University, Durham, North Carolina Received June 83. 1968
The solubility of helium, neon, argon, krypton and xenon have been determined in toluene, A uorobenzene, chlorobenzene, bromobenzene, benzene and nitrobenzene a t four temperatures and a pressure of one atmosphere. Heats and entropies of solution have been calculated. As previously found for the solubility of the rare gases in some hydrocarbons, there is B linear relationship between the logarithm of the Ostwald coefficient, and the surface tension of the solvent.
This investigation is a continuation of studies on the solubilities of the rare gases in selected nonelectrolyte solvent^.^-^ Toluene, fluorobenzene, chlorobenzene, bromobenzene, iodobenzene and nitrobenzene were chosen as solvents because of their increasing polarity. The measurements were made a t four temperatures over a 40" range. Experimental The solubility apparatus and procedures have been previously described.2 Excepting xenon, which was furnished by the Linde Air Products Co., the gases were the same as used previously. Reagent grade Mallinckrodt toluene was shaken with successive portions of sulfuric acid until the acid layer showed no color, was washed, dried over Drierite and fractionated in a 75 cm. vacuum jacketed glass column packed with 1 / 8 ' 1 glass helices. Table I gives the boiling ranges for the six purified solvents. Eastman Kodak white label fluorobenxene, chlorobenzene and bromobenzene were
TABLE I BOILINGRANGESOF SOLVENTS Solvent
Boiling range, "C.
Lit.ovalue," C.
84.28-84.68 84.75 Fluorobenzene 131.67-131.71 131.70 Chlorobenzene 155.86-155.90 155.79 Bromobenzene 77.40-77.60 (20 mm.) 78.38 (20 mm.) Iodobenzene 110.40-110.60 110.625 Toluene 81.0-81.2 (10 mm.) 8 3 . 9 (10 mm.) Nitrobenzene 4 R. R. Dreisbach, "Physical Properties of Chemical Compounds," Advances in Chemistry Series No. 15, Amer. Chem. Soc., 1955.
dried for several days over phosphoric anhydride before being fractionated. Eastman Kodak white label iodobenzene was shaken with successive portions of dilute sodium thiosulphate, washed, and dried for several days over phosphoric anhydride before being fractionated under reduced pressure (20 mm.). Nitrobenzene, Eastman Kodak white label, was distilled under reduced pressure (10 mm.) from phosphoric anhydride. The middle fraction was refractionated and the second middle fraction was used. The refractive indices of the freshly purified solvents and the solvents recovered a t the end of the measurements, and the boiling points agreed satisfactorily with the literature values .0 Results obtained with the present apparatus and procedure for some solubilities of hydrogen and argon compare (1) (a) Presented before the Division of Physical and Inorganic Chemistry, 131st National Meeting of the American Chemical Society, Miami, Florida, April, 1957. (b) Part of a thesis submitted by Rubin Battino to the Graduate School of Duke University in partial fulfillment of the requirements for the degree of Doctor of Philosophy, June, 1957. (2) H. L. Clever, R. Battino, J. H. Saylor and P. M . Gross, T H r s JOURNAL, 61, 1078 (1957). (3) H.L. Clever, ibid., 61, 1082 (1957). (4) H.L. Clever, J . H. Saylor and P. A t . Gross, ibid., 62,89 (1958). (5) H.L. Clever, ibid., 62, 375 (1958). (6) R. R. Dreisbach, footnote a of Table I. (7) M . W.Cook, U. 9. Atomic Energy Comm. Publ. No. UCRL2459 (1954).
satisfactorily with values obtained by Cook7 and Reeves and Hildebrands (see Table 11).
TABLE I1 SOMEGAS SOLUBILITIES IN TOLUENE A N D METHYLCYCLOHEXANE: COMPARISON WITH PREVIOUS MEASUREMENTS Solvent
Gas
Toluene Toluene Methylcyclohexane
HI A A
Temp., OC.
25 25 25
Ostwald coefficient Present investigaLit. tion value
0.0728 .249 ,354
0.07260' .2498 .35398
Results The solubilities of helium, neon, argon and krypton were determined at a total pressure of one atmosphere and the temperatures 15, 25, 40 and 55" in fluorobenzene, chlorobenzene, bromobenzene, iodobenzene, nitrobenzene and toluene. The xenon solubilities were measured at 15 and 5 5 O , and in fluorobenzene and bromobenzene a t 25'. The solubilities (Table 111)are presented in terms of both the Ostwald coefficient L (defined as the ratio of the concentration of the gas in solution to the concentration of the gas in the gas phase) and the mole fraction solubility x2. Over the temperature range studied a plot of log L against 1/T is essentially linear. The average precision of the measurements was 4% for helium and neon, 2y0 for argon and 0.5% for krypton and xenon. Discussion A. Heats, Free Energies and Entropies of Solution.-The standard heats, free energies and entropies a t 25" have been calculated for the transfer of one mole of the gas from the gas phase a t one atmosphere pressure t o the hypothetical unit mole fraction solution. There is a well known empirical linear relationship between heats and entropies of condensation of pure non-associated liquids and heats and entropies of dilute solutions. The solubilities of the rare gases in hydrocarbons2 have been shown to satisfy 0.00124AH.9 the equation AX = -12.75 The present data are plotted in Fig. 1. The deviations from the equation are much larger than in the case of the hydrocarbons and increase from xenon to helium. The heats and entropies of conden-
+
(8) L. W. Reeves and J. H. Hildebrand, J . A m . Chem. Soc., 79,1313 (1957). (9) H. 8. Frank, J . Chem. Phys., 13, 493 (1945).
Oct., 1958
1335
SOLUBILITIES OF RAREGASESIN BENZENE DER,IVATIVES TABLE I11 SOLUBILITIES OF HELIUM,NEON,ARGON, KRYPTON A N D XENON I N SOMEBENZENE DERIVATIVES t,
Solvent
oc.
Fluorobenzene
15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0
Chlorobenzene
Bromobenzene
Iodobenzene
Toluene
Nitrobenzene
----Helium-OatMole wald fraction
x 102 2.57 3.00 3.60 4.19 1.39 1.66 2.11 2.53 0.997 1.27 1.68 1.94 0.634 0.840 1.14 1.39 1.89 2.24 2.77 3.52 0.613 0.897 1.22 1.38
x 104 1.01 1.16 1.35 1.52 0.597. .696 .853 .990 .441 ,550 ,701 .782 .298 .385 ,504 .592 , ,846 .981
1.17 1.45 0.265 ,377 .494 ,540
--Neon-Ostwald
x io*
3.69 3.95 4.91 5.67 1.99 2.36 2.90 3.57 1.60 1.79 2.24 2.66 0.960 1.18 1.41 1.84 2.81 3.21 3.84 4.67 0.736 1.22 1.42 1.73
Mole fraction
x 104 1.46 1.52 1.84 2.07 0.853 0.986 1.17 1.40 0.706 .771 ,932 1.07 0.452 ,539 .621 .787 1.26 1.40 1.62 1.91 0.317 .509 .575 ,676
sation of the solvents used are found to agree very well with the equation. B. Solubility and Surface Tension.-A linear relationship is found for the present data between the logarithm of the Ostwald coefficient L and the surface tension CT of the solvent as shown in Figs. 2 and 3. This relationship was first pointed out by UhliglO who gave a theoretical explanation involving the radius of the cavity in the solvent formed by the gas molecule and a free energy of interaction between solute and solvent. When Uhl3g's equation was used with the present data to calculate cavity radii, it was found that the calculated radii decrease in going from helium to xenon while the atomic radii increase. Table IV gives values for the slopes and intercepts calculated by least squares for the linear relationship between the logarithm of the Ostwald coefficient and the solvent surface tension. Solubilities calculated with these values reproduce the experimental data within approximately 5%. The wide applicability of the surface tension relationship is illustrated in Fig. 3 which represents the solubility data for argon in 25 varied solvents at 25". The average percentage deviation for solubilities calculated from the least squares line 0.025) was 8%, and the (log L = -0.02210 greatest deviations were for iodobenzene (- 29%), water (21%) and carbon tetrachloride (18%). This would appear to be very good considering the diversity of the solvents. C. Hildebrand Equation.-A number of investigators have used equations derived from regular solution theory to correlate gas solubilities with
+
(10) H. H. Uhlig, THISJOURNAL, 41, 1215 (1937).
--Argon--
I
__ Krypton--
--Xenon
7
Mo!e fraction
Ostwald
Mole fraction X 103
Ostwald
0.291 .298 ,310 .313 .202 .204 .214 ,225 .153 .157 .165 .176 .lo4 ,109 ,118 ,131 .240 .249 .262 .275 ,100 .lo5 ,117 .123
1.15 1.15 1.16 1.14 0.864 .852 ,864 ,882 .675 .676 .687 .710 ,486 ,497 .521 .560 1.07 1.09 1.11 1.13 0.431 .439 ,477 ,483
0.880 .871 ,838 .814 .664 ,659 .634 .625 ,538 .532 ,514 .508 ,368 .370 .369 .371 .785 .775 ,741 .711 ,318 .338 ,338 ,348
t, OC.
X 103
3.47 3.36 3.13 2.96 2.84 2.75 2.56 2.44 2.38 2.29 2.14 2.04 1.73 1.70 1.63 1.58 3.50 3.37 3.13 2.91 1.37 1.42 1.37 1.36
'
Ost-
p
-
7
Mole fraction
15.2 25.0
3.798 3.350
x 102 1.479 1.298
55.0 15.2
2.732 3.288
0,985 1,390
55.0 15.2 25.0
2.304 2.798 2.478
0.894 1.222 1.057
55.0 15.5
2.062 2.084
0.824 ,968
55.0 15.2
1,608 3.722
,682 1.637
55.0 15.1
2.655 1.465
1.078 0.627
55.0
1.240
0.483
wald
ASs=OOO124AHg-1275
@/
@Om
0 HELIUM
a NEON
m
ARGON
0 KRYPTON
0 XENON
i
0
I
I -2
I -1
,
I 0
I AHs
I
I
2
3
(KCALIYOLE),
Fig. 1.-Relation between heats and entropies of solution for the gases in the six benzene derivatives.
the properties of the pure components. For polar solvents Gjaldbaek and Andersen'l have included a term to account for the dipole contribution to the energy of vaporization. Various attempts to apply these equations to the present solubility data have resulted in indifferent success. This is not surprising for solubilities of gases in polar solvents a t temperatures above their critical temperatures. D. Ideal Solubilities:-An ideal solubility xi may be evaluated from Raoult's law. At a pressure of one atmosphere pz the ideal solubility is defined as x2i =
l/p,O
(1)
where pzo is the saturation vapor pressure of the pure solute. In the present work, the ideal solu(11) J. Chr. Gjaldbaek and E. K. Andersen, Acta Chem. Scand., 8 , 1398 (1954).
JOHN H. SAYLOR AND RUBINBATTINO
1336
Vol. 62
c I
0
I
e
3
4
gc,
SURFACE TENSION
Fig. 2.-Logarithm
CDYNEICMI,
of Ostwald coefficient us. solvent surface tension.
Fig. 4.-Logarithm of mole fraction solubility of the rare gases us. gas phase polarizability.
TABLE IV SLOPES AND INTERCEPTS CALCULATED FOR log L = mu b
+
EQUATION
THE
Temp., Gas
Helium
Neon
Argon CULBON TFlRACHLMlIDE
.Krypton
20
40
SURFACE
60
TENSION.
Fig. 3.-Logarithm of Ostwald coefficient for argon vs. surface tension of diverse solvents.
bility was calculated using the Clausius-Clapeyron equationfor xenon12 at 1 5 as ~ 1/57.55 or o.01738. The mole fraction solubilities of xenon in a group of hydrocarbons and polar solvents a t 15' are cornpared with the ideal solubility in Table V. All of the hydrocarbons except benzene show negative deviations from Raoultls law and hence greater solubilities. As the dipole moment of the solvent (12) "International Critical Tables," Vol. Co., New York, N. Y., 192G, p. 204.
a,
Xenon
OC.
15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 25.0 40.0 55.0 15.0 55.0
m
b
-0.0372 - .0310 - .0310 - .0350 - .0403 - ,0318 - .0356 - .0364 - ,0282 - ,0277 - .0283 - .0288 - ,0265 - .0252 - ,0257 - ,0258 - ,0274 - .0210
-0.603 - .730 - .714 - .586 - .337 ,591 - .452 .415 ,232 .I93 .180 .161 .689 .612 .566 .509 1.281 0.975
-
increases the deviations from Raoult's law grow more and more positive and the solubility rapidly decreases. The solubility of xenon in highly polar water, which also has an extremely high internal pressure, is very E. Empirical Correlations.-As suggested by Hildebrand13a plot of the logarithm of the solubility against the solvent solubility ,parameter yields a relatively good linear relationship for the gases and solvents used in the present work.
McGraw-Hill Book (la) J. H. Hildebrand, THISJOURNAL,68, 671 (1954).
THEIONIZATION CONSTANTS OF ISOCITRIC ACID
Oct., 1958
1337
Friedman14 plotted the logarithm of the mole fraction solubility against an arbitrary scale for the abscissa chosen so as to make the solubility curve THE for one of his solvents linear. A plot of this type % Solvent xa Dev." was made for the present data with the abscissa 0.0001~ 99 Water chosen to mako the bromobenxene curve linear. Nitrobenzene .0063 58 The points for all the other solvents also fell on .0097 44 Iodobenzene straight lines. .0122 30 Bromobenzene This indicates that the solubility is a function of .0134d 23 Benzene some physical property of the rare gases. Several .0139 20 Chlorobenzene physical properties of the rare gases were plotted Fluorobenzenc .0148 15 against this arbitrary abscissa. The property .0164 + G Toluene which yielded the best fit was the polarizability. Perfluoromethyl.0188" - 8 Figure 4 shows a plot of log 5 2 against the rare gas cyclohexane polarizability. Figure 4 shows a plot of log z2 Cyclohexane .0237d -36 against the rare gas polarizability for the present Methylcyclohexane .0252' -45 data a t 25". .0303d -74 n-Hexane Acknowledgment.-We express our gratitude to .031fjd -83 Isooctane Dr. H. L. Clever for helpful discussions, and to the n-Dodecane .0352d - 102 Tennessee Eastman Company, Division of EastT. J. man Kodak Company for financial support in the a Signs indicate deviations from Raoult's law. Morrison and M. B. Johnstone, J . Chem. Soc., 3441 (1954). form of a fellowship for R. B. H. H. L. Clever, et al., THISJOURNAL, 62, 89 (1958).
COMPARISON OF
TABLE V SOLUBILITIES OF XENONAT 15' IDEAL SOLUBILITY x2' = 0.0174 THE
WITH
+ + + + + + +
L . Clcvor, et al., ibid., 61, 1078 (1957).
(14) H.L. Friedman, J . A m . Chem. Soc., 76, 8294 (1954).
THE IONIZATION CONSTANTS OF ISOCITRIC ACID BY DAVIDI. HITCHCOCK From the Department of Physiology, Yale University School of Medicine, New Haven, Connecticut Received June $8, 1.968
Buffer solutions of isocitric acid, partly neutralized by KOH in amounts corresponding t o the mid-points of the three stages of ionization, were adjusted to known ionic strengths by the addition of KCl and by dilution. Hydrogen ion concentrations, obtained by the use of the glass electrode, were based on comparison with HCl-KCI solutions. From these measurements the three ionization constants in terms of concentration, were evaluated for each of 4 ionic strengths. Extrapolation gave the negative logarithms of the true ionization constants a t 25" as 3.287, 4.714 and 6.396. The p K values for 38" were lower by 0.02 t o 0.04.
Although isocitric acid is a substance of considerable biochemical importance, it was available only in small amounts before Vickery and Wilson' discovered a method for the isolation of its monopotassium salt from certain leaves. Several grams of this salt, recrystallized and dried a t the Connecticut Agricultural Experiment Station, were made available for the determination of the ionization constants of the acid. Experimental Samples of the primary salt weighing about 0.65 g. lost only 0.4 mg. in weight during 1 day at 110". Quantitative conversion of these samples to KzS04 indicated that 1 equivalent of K was contained in 229.8 g. of the dry sa.lt. Titration to pH 9.2 with a standard KOH solution indicated that 1 equivalent of acid hydrogen was contained in 115.8 g. of the dry salt. As the formula weight of the salt is 230.21, both analyses imply a slight contamination with the dipotassium salt. Since the KOH solution was used in preparing the buffers to be investigated, the titration figure was used in the calculation of their composition. Solutions were prepared from 3.0720 g. of the primary salt, dissolved in distilled water and diluted to 100 ml. a t 23'. Three portions of this, 30.00 ml. each, were pipetted into 100-ml. flasks for the preparation of 3 stock solutions, A, B and C. Solution A contained also, in 100 ml., 20.00 ml. of 0.1000 N HC1 and 1.193 g. of KC1. Solution B contained also 6.62 ml. of 0.3100 N KOH and 0.894 g. of KCl. Solution C contained also 19.49 ml. of the KOH solution (1) H. B. Vickery and D. G. Wilson, J . B i d . Chem., 233, 14 (1958).
and 0.149 g. of KCl. The KCl had been purified by recrystallization from water. The KOH had been freed from carbonate by addition of CaO to a 1 M solution, according to Kolthoff ,2 followed by settling, decantation and dilution. It was standardized by titration against 0.1 N HCl which had been standardized against Na&Oa and against borax. Solut,ions A, B and C, which contained approximately 0.04 M isocitrate at 0.2 ionic strength, were diluted to give ionic strengths of about 0.15, 0.10 and 0.05. The hydrogen ion concentrations of these 12 buffer solutions were obtained from measurements with a glass electrode (Beckman No. 1190-80) supported in a waterjacketed vessel a t 25" (or 38') rt0.05". The other half cell was a saturated KC1 calomel electrode, also in a waterjacketed vessel, and a liquid junction with saturated KCI was formed a t the 2 mm. bore of a stopcock. This electrode system gave results, with solutions of pH 2 to 7, which ran parallel, within 0.2 millivolt, with those obtained by Hitchcock and Taylora who used a hydrogen electrode and made the liquid junction in a bulb about 10 mm. wide. A different standard solution was em loyed for each ionic strength; these solutions were 0.01 N k C l in 0.19, 0.14 and 0.09 N KC1, and 0.005 N HCl in 0.045 N KCI. Since the difference, E , between the electromotive force observed with the standard and the unknown was equivalent to the e.m.f. of a hydrogen electrode concentration cell, the Nernst equation E = k log (CHOIICH) was used to evaluate the hydrogen ion concentration Cn in (2) I. M. Kolthoff, 2. a n d . Chem., 61, 48 (1922).
(3) D.I. Hitohcock and A. C. Taylor, J . Am. Chem. Soc., 69, 1812 (1937).
