within the experimental uncertainty.13 Under the colored substance and chromium were extracted most favorable conditions employed (the last expt. in the same proportion, all of the chromium must in Table I), the molar absorbancy index at 546 mp have been complexed. The close similarity of the is 3.6 X lo4 based on chromium(I1). This is magenta product of the chromium(I1)-diphenylsubstantially larger than the value of 2.5 X lo4 carbazone reaction, both in the form of its absorp(at 540 mp) obtained by Pflaum and H o ~ i c k . ~tion spectrum and in its molar absorbancy index, It seems likely that when the reaction is carried out suggests that it is the same substance, and that under 0ptimu.m conditions the molar absorbancy here too all of the chromium is complexed. The absorption spectrum of the chromium comindex will be the same as that of the magenta product of the chromium(VI)-1,5-diphenylrarbo- plex in aqueous solution is rather similar t o that of hydrazide reaotion, which is (4.17 0.04) x l o 4 diphenylcarbazone in glacial acetic acid (or, to a lesser extent, in certain other solvents). It ic a t 546 mp bmed on chromium(VI).ll Extra.ction studies of the latter sitbstance have markedly different from that of diphenjrlcarbazone shown that i t is a chromium ~ o m p l e x . ' ~As the in dilute acetic acid and other aqueous soliitions. Previous investigators45 have obtained evidencc (13) IT" are indehtrd t o Dr. .I. I. Poi)ov for pointing o u t this that the complex contains diphenylcarbazone and method. chromium(I1I). Because of the similarity ill (14) I. 11. Lichti3netein and T. I,. Allen, J . Am. C h ~ r n S. o c . , 81, 1040 spectra, it is reasonable to suppose that diphenyl(1959). I n this article tlie complex was inadvert,ently called a n organometallic cornpleu. Altliorigh t h e s u r d organometallic commonly carbazone has a similar structure in the complex indicates that the inptal is bonded directly t o a carbon a t o m , such a n and in glacial acetic acid; complexing with chroimplication \vas not intcnded, inaerniicii as t!ie complex is prohably a mium(II1) makes this structure stable with respect chelate in which chroniiiiin is bonded t o oxygen and/or nitrogen to the form normally present in aqueoiis snliition. atoms.
THER&IODI'KA&TICSOF SOLUTION OF HIPPURIC ACID IN WATER AND I N VARIOUS SODIUM CHLORIDE SOLUTIOKS',2 BY RICCIJ. LARESEAND WILLIAMJ. CANADY West Virginia University Medical Center, Department of Biochemistry, Morgantown, West Virginia Received February 10, 1961
I n view of the fact that hippuric acid is a compound of considerable biological importance, and since crystallization is used as a step in the quantitative determination of hippuric acid, the solubility of hippuric acid as a function of temperature and ionic strengcth has been studied. The solvents varied from pure water to 3.0 m sodium chloride solution. Variatioiis of the free energies, heats and entropies of solution with ionic strength are presented. The free energy of solution a t 25" is linear with ionic strength, while the heat of solution varies very little with ionic strength. The variation of the change in heat capacitv for the solution process with salt concentration may be explained in a rough qualitative manner in terms of the iceberg theory of Frank and Evans. The variation of the entropy change with salt concentration is very slight at 25" but when extrapolated to lower temperatures it tends to increase with an increase in ionic strength as would be expected from iceberg theory.
The solubility of hippuric acid in water and in salt solutions has been of practical interest for a number of years, since crystallization is used as a step in the quantitative determination of hippuric acid.3-) Very little work to date has had to do with the temperature dependence of the solubility of hippuric acid or the effect of added neutral salts upon this dependence. Equilibrium mas studied from both undersaturation and oyersaturation. Solvents ranged from pure water to three molal salt soliition; seven temperatures were investigated, ranging from npprosimntcly I5 to approximately 45'. Experimental Equipment.--Constant trmperaturr was maintained hy means of a Sarqrnt thermistor controllrd water-bath. Th(A temperaturr riav be maintained within f0.01'. The (1) This investigation wns supported by a grant from t h e National Science Foundation. (2) A h , s t r a c t ~ from d a thesis by R. J . Larese in partial fulfillment of t h e requirements for the 11,s.degree, \Vest Virginia University. ( 3 ) A . J. Quick, A m . J . M e d . Sei., 186, 030 (1933). (4) 4. .I. Quick, A m . J . Clin. Path., 10, 222 (1940). ( 5 ) T. E. Weictiselbaum and J. G. Probstein, J . Lab. Clin. Med.. 24, 636 (:Y39).
temperature was measured with two EXAX solid point thermometers having temperature ranges of ,O to 30' and 20 t o 50°, respectively, in increments of 0.1 . These thermometers were calibrated against a Leeds and Northrup platinum resistance thermometer over the entire experimental range. The platinum resistance thermomrter had been previously calibrated by the U.S. Bureau of Standards. A Beckman DIiV Spectrophotometer v a s used for t h r photometric determinations. Materials .-The hippuric acid was ohtained from the Fisher Corporation and was the "highest purity" grade. It was recrystallized twice from water and its melting point remained unchanged a t 187.2'. The water was doubly distilled through all glass apparatus. The sodium chloridc used was J. T. Baker C.P. grade. Procedure.-For equilibrium approached from under saturation, one gram of hippuric acid was placed in a 100 ml. screw cap Pyrex erlenmeyer flask along with 80-90 ml. of solvent. Water-proofed corks were placed around the necks of the flasks and this whole assembly was wired to lead sinkers. The lengths of wire linking them m t h the lead sinkers a t the bottom of the bath were arranged YO that the flasks would float with a gentle bobbing, rotating motion due to the brisk circulatory flow of water in the bath, but were not allowed t o move about freelv. For equilibrium approached from oversaturation, thr procedure was similar to that described above except that the flasks containing hippuric acid crystals and solvent were maintained a t approximately 55" for six hours before being placed in the bath. Samples were removed a t various times from both the
.July, 1Nil
T I I C l t \ 2 0 D Y S b \ l I C S O F SOLUTIOTS O F
under and oversaturated flasks. For the studies with pure water 2.5 ml. was removed; for the studies with salt solutions the samples consisted for 10 ml. In each case care was eyercised to avoid taking up any solid. The aliquot was transferred t o a voliimetric flask which was held in the same bath. After a 15-minute equilibration, the volume WRS adjusted evactlv i o the mark. This made it possible to be certain thnit no appreciable volume change took place from the di&t warming or cooling effect introduced by removal of the sample. The flasks then were stoppered and wei,ghed. The contents of each volumetric flask were tliluted in such a w:ty as to produce a solution ranging from 2-5 X lo-. J1 ic hippuric acid for spectrophotometric determination at 230 nip Srrid dilutions mere found t o follow Beer’s law. All readings mere compared to n stitndard curve dwived by least squares from known solutions of hippuric n t d .
Results The r(wlts ohtaind from undersaturation did iiot differ significantly from those obtained from oversaturn ticn. Approximately half of the experiments nTerc clone in each way. The average deviation from the mean a t each temperature for 281 determinations was 1.13%. The logarithm of the solubility, s, expressed as molality, was plotted against the reciprocal of the absolute temperature on large scale graph paper. A smooth cwve mas drawn and the values read a t mrious values of 1/T. These smoothed values are listed in Table I. The sinoothed data were fittrd to an equation of the form AH I n s = --,! 121
+ ACp In T + c R
where AH = A H 0
-
+ ACpT
It is assumed that the conventional activity coefficient is independent, of concentration.
IIIPPURIP hCID
1241
I N WATER
for pure water. Since these terms vary considerably with small experimental errors, it is doubtful that this is significant. The constants appearing in eq. 1 are given in Table 111 so that the solubility may be calculated for any temperature within the experimental range. ,A linear relationship is found when AFO is plotted against ionic strength. Since the logarithms of the solubility of many non-electrolytes produce a linear result when plotted against ionic strength, the linearity of the free energy plot is not surprising. .4 plot of relative activity coefficient, y, against ACp, the change in heat capacity, results in a fairly good straight line. TABLE I1 NaCl concn.
AF, cal.
A H . oal.
A& e.u.
0.0 .2 .5 1.0 2.0 3.0
2293 2342 2401 2511 2723 2949
6264 6421 6485 6479 6561 6813
13.3 13.7 13.7 13.3 12 9 13.0
ACP, cal./’C.
Y
103.7 104.G 100.1 91.0 54.4 21.1
1.00 1.09 1.20 1.44 2.07 3.03
TABLE 111 CONSTAXTS O F EQUATION
NaC (.onen.
AHo
0.0 .2 .5 1.o 2 0 9 0
-24647.60293 -26112.56126 -23371 88995 -20651 03184 - 9672 34056 513 62820
+
_ACP
1
R
-C
52.18857 54.92649 50.41238 45.80447 27.40730 10 63547
342.84110 360.Y9648 330.74994 300.08691 177.08581 64.70971
Discussion 11 relatively small variation of the heat of solution a t 25” with ionic strength was observed. If 1,OGARITHMS O F THE SOLCBILITJES O F HIPPURICA C I D I N the heat of solution is calculated for higher tempera\VA4TErL AWD hTaCl : h L T T I O N S AT lrARIoUS TEMPERATURES e.g., 45O, the effect becomes even smaller, tures, Concentrations are expressed as 1molalities while a t lower temperatures considerable varia- Logs T,or
