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).
DAVID I. HITCHCOCK
1338 3.35,
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3 25
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1
15
20
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IO IONIC
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STRENGTH.
upper curve, Fig. 1.-Extyapolation for pK1 at 25': p k i 4- l . O W d I / ( l
