NOTEB
496
tate were used as buffer components. Water de-ionized with a mixed bed exchanger and possessing a specific resistance of ca. 5 X 108 ohm/cm., and Fisher certified dioxane redistilled over sodium were used as solvent components. Allied Chemical reagent grade pyridine was fractionated over a packed column, b.p. 115". Eastman Kodak White Label acetic anhydride wm used without further purification. Measurements.-The fraction of unprotonated pyridine in the buffer solution was determined by the relationship [CsHsN]rree = .B!( - ~ H ) / ( ~ o E a=), wbere a ~a H, , and aoH are the optical densities of pyridine at 2550 A. in the buffer, in 0.1 M hydrochloric acid, and in 0.1 M sodium hydroxide solutions, respectively. The measurements were made on a Cary lModel 14 spectrophotometer in a 0.01 cm. quartz cell. The buffer solution had the same composition as that used in the kinetic determination; the concentration of total pyridine was 0.0264 M. Under these conditions pyridine was found to be 96% unprotonated. Kinetic measurements were made by following the decrease of the height of the acetic anhydride proton magnetic resonance signal which was well separated from the acetic acid-sodium acetate signal by 17 C.P.S. A Varian A-60 spectrometer was used for this purpose. The reaction was started by exliausting from a micropipet 0.10 ml. of acetic anhydride into 10.0 m!. of the desired solution, which gives 0.10 M initial concentration of acetic anhydride. Two solutions were used in the kinetic determination: (1) a blank solution consisting of 1.06 iM NaQAc and 0.53 M HOAr in 60% water-40% dioxane; (2) a solution identical with the blank solution but containing in addition 0.063 iM free pyridine. The initial concentration of pyridine under these conditions is only 2% higher than the final concentration due to the additional acid produced in the solution from the hydrolysis of acetic anhydride. This conclusion derives from the relationship
Vol. 67
was derived from the theory of Debye and Huckel.' The values of k (Table I), derived213 in 1931 from Falckenberg's measurements* of dD/dP and from molal volumes determined by Baxter and Wallace,6 could not claim high accuracy. Moreover, values of dD/dP obtained from Kyropoulos' datae led to much higher resu1t~'~~for k.
-
where Kp and K H Aare ~ the acid dissociation constants of pyridine and acetic acid, f is the fraction of unprotonated pyridine in the buffer solution, and Pt is the totaI concentration of pyridine. A plot of log ( ht - h,) us. time results in a straight line, the slope of which gives the first-order rate constants. The rate constants from solutions 1 and 2, respectively, are 5.22 X 10-8 sec.-l and 5.07 X lO-*ssec.-l.
THE MOLAL VOLUME OF ELECTROLYTES BY OTTOREDLICH Department of Chemical Engineering and InoTganic Materials Division of the Lawrence Radiation Laboratory, University of California, Berkeley, California Received June 66, 1966
Recent, obviously excellent determinations' of the dielectric constant of water between 0 and 70°, and 1 and 1000 bars, remove any doubt from a question that has been under discussion for several decades, A limiting relation, expressing the apparent molal volume of an electrolyte as a function of the valences x1 of its ions, and the dielectric constant D and compressibility p of the solvent cp =
p20+ ~
~ ~ 1 . 5 ~ 0 . ~
(1) w = 0.52; v i ~ i ' (2) k = 2N2r3(2a/1000RT)o.5D-1~6(d In D/dP - p/3) (3)
R
=
E
= 4.8029 X lQ-'O
N
=
83.1469 X lo6 erg(deg. mole)-' 6.0232 X
e.s.u. mole-'
(1) B. B. Owen, R. C. Miller, C. E. Milner, and H. L. Cogan, J. Phys. Chem., 611. 2065 (1961). See also F. E. Harris. E. W. Haycock, and B. J. Alder, ibid., S7, 978 (1953).
TABLE I COEFFICIEST k AT Temp., OC.
IC
16
1.8 i 0 . 5
25
1.7
20
2.53
25 25
1.86 i 0.02 2.517
25
1.884
OR
NEAR25'
Based on
Reference
dD/dP (Falckenberg) Molal volumes (Baxter, Wallace) dD/dP ( Kyr op oulos ) Molal volumes dD/dP (Kyropoulos) dD/dP (Owen, et al.)
(2) 1931 (3) 1931
( 7 ) 1933 (9) 1940 (8) 1949 1962
TABLE I1 COEFFICIENT k Temp., OC.
IO'* d In D / d P , dyne-' em.*
10'2 @,a dyne-1 om.2
k, cm.8 (mole/l.)-O
1
0 45 14 45.42 1.539 10 45.84 44.85 1.668 20 46.65 44.52 1.809 30 47.58 44 43 1.963 50 49.78 44.85 2.318 70 52.43 46.08 2.746 a L. B. Smith and F. G. Keyes, Proc. Am. Acad. Arts. Sci., 69, 286 (1934).
By 1940, however, accurate density determinations by Geffoken and his co-workers, and by Wirth, furnished a reliable basisg for the value k = 1.86 0.02, though the difference from the value derived from Kyropoulos' data was large. The recent measurements' are in perfect agreement with the conclusions of 1940 and eliminate any reason for using arbitrary empirical values instead of the derived values given in Table 11. Moreover, no attempt need be made to make the data fit relation 1 by introducing terms of higher orderlo a t unusually low concentrations.
*
(2) 0.Redlich and P. Rosenfeld. Z . phyeik. Chem., A M s , 6 5 (1931). (3) 0.Redlich and P. Rosenfeld, 2. Elektrochem., 3 1 , 705 (1931). (4) G. Falckenberg, Ann. Physzk, [41 61, 145 (1920). (5) G. P. Baxter and C. C. Wallace, J . A m . Chem. Soc., 38, 70 (1916). (6) 9. Kyropoulos, 2. Phyaik, 40, 507 (1926). (7) F. T. Guoker, Jr., Chem. Rea., 18, 111 (1933). (8) B. B. Owen and 6. R. Brinkley, Jr., Ann. iV. Y . Acad. Sci., 61, 753 (1949). (9) 0.Redlich, J . Phys. Chem., 44,619 (1940). (10) H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd Ed., Reinhold Publ. Corp., New York, N. Y., 1958, p. 390.
CHBRGE-TRANSFER COMPLEXES OF METHYLVIOLOGEN BYAKITSUQTJNAKAHARA~ AND JIJIH. WANQ Contribution No. 1YO3 from Sterling Chemistry Laboratory, Yale University, New Haven, Connecticut Received June 86, l g S 8
Molecular complexes with absorption spectra uncharacteristic of the components of the respective (1) On leave of absence from Institute of Chemistry, College of General Eduoation, Osaka University, Toyonaka, Osaka, Japan.
