CHARGED CARBOXYLATE BASESIN DILUTEACIDSOLUTIOXS
.4pril, 1963
779
TABLE IV LEASTSQUARES DATA Temp., OC.
Efficiency of fit, %
25 40
96 90
Sbandard dev.
x
Slope
Intercept
- 0.0248
- 0.5734 - 0.3384
104
1.78 1.42
- 0.0232
solute-solvent interaction energy. For the 25 and 40’ data, respectively, the water radii calculated were 1.37 and 1.36 A.,while the interaction energies were -782 and -485 cal./mole. Although literature values for the radius of the water molecule range from 1 to 3 A., depending on the assumptions made, some of these are very close to the values determined in this work: 1.25,13 1.32,13 1.39,14 and 1.44.13 The negative interaction energies can be interpreted, according to Uhlig,12 as an absorption of energy by the system as water dissolves, which is indicative of a low solubility. The negative interaction energies account in part, according to Le Chatelier’s rule, for the increase in solubility of water with increasing temperature.
SOLVENT SURFACE TENSION, d y n e r l c m
Fig. 3.-Alkane
Conclusions Results of water solubilities in normal alkanes from C7 to CI6show a gradual decrease with increasing molec(13) N. E. Dorsey, “Properties of Ordinary Water Substance,” ACS Monograph, Reinhold Publ. Corp.. New York, N. Y.. 1940. (14) C. ,J. F. Bbttcher, Rac. trau. chim. (Rotterdam), 66, 14 (1946).
water solubilities and surface tensions.
ular weight, when the solubilities are expressed on a weight basis; when expressed on a mole fraction basis, the reverse trend occurs. The experimental water solubilities agree well with two different theoretical treatments.
ACID-BASE EQUILIBRIA I N CONCENTRATED SALT SOLUTIOKS. 11. CHARGED CARBOXYLATE BASES I N DILUTE ACID SOLUTIONS1 BY JAMES S. DWYER AND DONALD ROSEXTHAL~ Department of Chemistry, The University of Chicago, Chicago, Illinois Received September 6, 1968 ~HG.E - pMH is shown to be a constant for a given salt solution in slightly acidified 1 to 8 &r LiCI. (pHa E . is the pH measured using a cell with a glass indicator electrode and a saturated calomel reference electrode. pMH = -lag total strong acid concentration.) Values of p H c . ~ . pMH are reported. In dilute acid solutions of acetate or formate ion KBH= ( [B-][H+]/[BH])QBH where KBHis the thermodynamic acid dissociation constant of the acid (BH), and QBR is a constant for a particular base, B-, and a particular salt solution. The values of ~ H G . Ecalculated . using this equation are in good agreement with the eyperimental values. Values of log &BE are reported for acetic acid in 1-8 M LiCl. It is shown that calculated values of log QBHusing the equation log QBH = - ( d Z / l A dZ) BM do not differ significantly from the experimental values. Calculated and experimental QBH values are compared with published results for dilute LiCl solutions.
+
+
Introduction I n a previous study3PHG.E. - pMH was shown to be a constant for dilute acid solutions in 4 and 8 M LiC1, 6 M NaC104,6 M Na,:l‘Oa, and 4M CaCL (PHG.E.is the pH measured using a cell with a glass indicator electrode and a saturated calomel reference electrode and pMH is -logarithm of the total molar concentration of strong acid.) Further, it was found that the quantitative aspects of equilibrium between a weak uncharged base, B, and its conjugate acid, BH+, in such solutions can be accounted for using the equation KBHc =
([B][total strong acid c ~ n c n . ] / [ B H + ] ) & ~ ~ (1)
(1) Taken in p a r t from the Ph.D. research of James 9. Dwyer. This work s u p p o ~ t e dby research grants from the U. S. Public Health SerTiee
\vas
(RG-9583 and RG-8069). (2) To n h o m inquiries should be direrted a t Clarkson College of Technolony, Potsdam. New York. (31 D. Rosenthal and J. S. Dwyer .I. Phrs. Chem., 66. 2687 (1962).
where K B H + is the thermodynamic molar acid dissociation constant and QBH+ depends upon the nature and concentration of salt and the nature of the base, B. Apparently, the various factors4-” which are important in concentrated salt solutions can be satisfactorily incorporated into the QBH+term. (4) H. S.Harned and B. B. Owen, “The Physical Chemistry 01 Elert,rolytic Solutions,” Reinhold Publ. Corp., New York, N. Y., 1958, pp. 509-547. ( 5 ) H . 9. Frank. et at., J . Chem. Phus., 13, 507 (1945); Ann. Res. P h y s . Chem., 6, 43 (1954). R. W. Gurney, “Ionic Processes in Solution.” 410Craw- Hill Book Co.. Inc., New York, N. Y., pp. 248-260. ( 6 ) J. B. Hasted, et al., J . Chem. Phys., 16, 1 (1948); 29, 17 (1959). (7) (a) R . 4. Robinson and R . H . Stokes, “Electrolyte Solutions.” Butterworths Publications Ltd.. London, 1959, p. 62; (b) E. Glueckauf in “The Structure of Electrolyte Solutions,” M7. J. Hamer (editor), John Wiley and Sons, Inc., New York, N. Y . , 1959, pp. 97-112. (8) J. Beck, Physik. Z., 40, 474 (1939); G. W. Brady, J. Chem. rhus., 28, 464 (1958); H. S.Frank and P. T. Thompson in ref. 7b, pp. 113-134. (9) Ref. 4, pp. 614, 634-643. (10) E. Grunwald, G. Baughman, and G. Kohnstam, J . A m . Chem. Soc.. 62, 5801 (1960). (11) J. K. BrSnsted, Trans. Faraday Soe., 2S, 430 (1927); G. Scatchard in rof. 7b, pp. 9-18.
JAMES 8. DWYERAND
780
T’ol. 67
DONALD 13OSESTHtlL
TABLE I COMPARISON OF CALCULATED AND EXPERIMENTAL VALUES OF p H o . ~OBTAISED . IN % Neutralized
10 20 50 90 100 110 120 140 pMH - ~ H G . E . log Q=*cafb
S . D * E
7-4
JV LiCl--
Exptl.
Calcd.
4.76 4.76 4.42 4.41 3.80 3.81 2.88 2.88 2.22 2.21 1.57 1.57 1.28 1.29 0.99 0.99 0.93 - ,015 ..007
THE
SEUTRALIZATION OF 0.03 SaAc JI NaNOa-
7--8 Jf LiCl-Exptl. Calcd.
-6 .li KaCIOdExptl. Calcd.
Exptl.
4.01 3.69 3.07 2.13 1.28 0.23
