Vol. 75
MARTINKILPATRICK, R. DEANEANESAND J. G. NOKSE
588 [CONTRIBUTION FROM
THE
THE UNIVERSITY OF PENNSYLVANIA AND THE CHEMISTRY OF ILLINOIS INSTITUTE OF TECHNOLOGY]
DEPARTMENT OF CHENISTRY OF
DEPARTMENT OF
The Dissociation Constants of Acids in Salt Solutions. N. Cyclohexanecarboxylic Acid BY MARTINKILPATRICK, R. D. EANESAND J. G. MORSE' RECEIVED AUGUST22, 1952 By measurement of cells with liquid junction the dissociation constant of cyclohexanecarboxylicacid has been determined in solutions of alkali chlorides in hydroxylic solvents and the results extrapolated to infinite dilution to give the thermodynamic dissociation constant.
The equilibrium constant for the isoelectric reaction
Potassium chloride 0.695 1.80 1.44 .695 1.96 1.57 HCy BoA0 4- e y (1) .703 2.03 1.62 where HC, represents cyclohexanecarboxylic acid .698 2.09 1.67 and A0 benzoic or acetic acid, has been determined .698 2.13 1.70 in salt solutisns as given earlier2:' and the results ,698 2.13 1.70 are summarized in columns two and four of Table I. .703 2.14 1.71 There is very little change in the ratio of the disso.703 2.06 1.65 .80 .697 2.01 1.61 TABLE I .90 .709 2.02 1.62 DISSOCIATION CONSTANT OF CYCLOHEXANECARBOXYLIC ACID 1.oo .717 1.97 1.58 AT 25' 1.50 .723 1.77 1.42 Electrolyte 2.00 .728 1.53 1.22 moles/ K O X 10' litK A = ~ K e X 10' K A . ~ Ka X 10' Hydro2.50 ,736 1.27 1.02 Ref. acid Benzoic Acetic cbforic K. 3.00 .742 1.06 0.85 Lithium chloride Based on K O values from paper I.e b Based on K, 1.80 1.46" values from paper 11.: K , from column 5, K, = 1.25 X 0.05 0.693 1.82' 0.195 1.79" 10-6. 2.12 1.76 .695 2.20 10 .I95 2.16 2.33 1.82 .693 2.27 .20 .205 2.35 ciation constants with equivalent salt concentration 2.46 1.88 .693 2.35 .30 ,208 2.47 up to 0.5 molar. This result appears surprising 2.61 2.02 ,690 2.53 .40 .200 2.59 when one considers the difference in structure of 2.52 1.98 .682 2.48 .50 .195 2.53 the acids. The above observation applies to 25' 1.97 .682 2.46 .60 and would not be expected to hold at other temper. . 1.97 atures. .684 2.46 ... .70 . . 1.93 .. .672 2.41 .80 ... From a knowledge of the dissociation constants .. .665 2.41 *. 1.93 .90 *.. of benzoic acid and acetic acid in the same salt solu2.42 1.90 .657 2.38 1.00 0.191 2.46 tions the values of the dissociation constant of cy2.24 1.85 .652 2.31 .179 2.26 1.50 clohexanecarboxylic acid are calculated and pre2.09 1.64 .645 2.05 2.00 .188 2.12 sented in columns three and five of Table I. These 1.69 1.37 .635 1.71 2.50 .179 1.69 results are in good agreement with the direct meas.. 1.27 .. ,615 1.59 3.00 ...
+
0.05 .IO .20 .30 .40 .50 .60 .70
a
...
0.05
...
.20 .30 -40
.198 .198 ,198 .198
.50
.198
.10
.60 .70 .80 .90 1.00 1.50 2.00 2.50 3.00
-
.. ..
Sodium chloride .. 0,699 1.78 1.99 .702 1.96 2.14 .706 2.10 .699 2.19 2.25 .702 2.27 2.36 .702 2.30 2.36 .702 2.27 .699 2.24 .696 2.17 .696 2.13 .696 2.08 .702 1.92 .710 1.64 .722 1.34 .724 1.14
..
.. 1.96 2.17 2.28 2.32 2.35
1.42 1.57 1.68 1.75 1.82 1.84 1.82 1.79 1.74 1.70 1.66 1.,54 1.31 1.07
0.91
(1) This paper waa abstracted from the dissertation presented by Jerome Gilbert Morse to the Graduate School of Illinois Institute of Technolonv in Dartid fulfillment of the reauirements for the degree of Doctor olhhilosophy, June, 1952. (2) M. Kilpatrick, THIS JOURNAL,76, 584 (1953). (3) M. Kilpatrick and R. D.Eanes, ibid., 71, 586 (1953).
urements against hydrochloric acid shown in column six. The small dependence of the equilibrium constant of equation (1) on ionic strength makes possible a more accurate extrapolation to infinite dilution in accordance with the equation KA.B~
E
[KA,B,]' f BC
(2)
The intercept [KA,B,]~ and the slopes B are given
in Table 11. TABLE I1 Range Reference electrcacid lyte
0.05-0.5 M
Intercept
Benzoic LiCl 0.199 NaC1 .198 Acetic LiCl .700 NaCl .695 KCl .696
Slope
0 . 0 5 - 3 . 0 hf Intercept Slope
0.00471 0.201 ,000 --
...
,0339 0,702 .00247 .696 .00564 .690
0.0091 ..,...
-0.0329 .0084
.0206
The ratios multiplied by the dissociation constant of benzoic acid 6.32 X lov6and acetic acid 1.754 X 10-6 yield values for the dissociation constant of ( 4 ) D. H. Everett and B. R. W. Pinsett, 1.Chcm. Soc., 1029 (1948).
Feb. 5, 1953
PROPERTIES OF ION-EXCHANGE RESINSIN RELATION TO STRUCTURE
589
HC,of 1.26 and 1.22 X 10-6 which are to be com- salt lithium chloride are summarized in Table 111. pared with literature values from conductance data of 1.26,61.346and 1.32 X 10-6.’ Similar measurements in the solvents methyl and ethyl alcohol and ethylene glycol, for the solvent
In these cases there are no literature values for comparison and the thermodynamic dissociation constant is based on the values for benzoic acid. Justification for the assumption that the ratio of the dissociation constants is independent of the lithium chloride concentration is given in Table IV.
TABLE I11 THEDISSOCIATION CONSTANT OF CYCLOHBXANECARBOXYLIC ACIDIN ALCOHOLS T. 25’: solvent salt. LiCl fi‘
&BC 0 0.10
0.923 0.220 8.32
9.0
0.196
0.170 7 . 8 1 44.8 0.213
( 6 ) J. s. Lumsden, J . Chrm. Soc., 81, 90 (1905). (6) N . Zelinsky and N. Izgatuisher. J. RUSJ.Phrs. Cham. jot., 40, 1379 (1908). (7) C. € Spiers I. and J. F. Thorpe, J . Cham. Soc., 117, 638 (1926).
THE
0.103 .193
0.196 .I95
0.279 .193
0.382 ,195
0.475 ,195
Mainly LiCl. 94.0
Reference acid, benzoic acid.
[CONTRIBUTION FROM
TABLE IV KA.B, for HCy/HB in the solvent ethanol
F.prther evidence for the independence of the ratio of dissociation constants on electrolyte concentration in alcohols will be presented in another paper. CHICAGO, ILLINOIS
CHEMICAL RESEARCH LABORATORY, DEPARTMENT OF SCIENTIFIC AND INDUSTRIAL RESEARCH 1
Properties of Ion-Exchange Resins in Relation to their Structure. III. Kinetics of Exchange BY D. REICHENBERG RECEIVED MAY28,1952
A study has been made of the kinetics of sodium-hydrogen exchange on sulfonated cross-linked polystyrenes in bead form. Resins of 5 and 17% (nominal) divinylbenzene content were fractionated by wet elutriation and matched according to swollen particle diameter. The rates of exchange were measured by an :‘indicator” method. For both resins, the kinetica a t high sodium ion concentrations in solution (>1 N) are clearly distinct from those a t low concentrations (
