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 a t 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 A4 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.
NOTES
Feb., 1963 complexes are often called charge-transfer complexes.2 We wish to report here a new group of charge-transfer complexes formed between N,N’-dimethyl-4,4’-dipyridinium ion, hereafter referred to as methylviologen after Michaeli~,~ and various electron-donor anions. Because they are good electron acceptors, all viologens have strong tendency to form charge-transfer complexes with donor anions. An aqueous solution of methylviologen chloride is colorless, that of the bromide slightly yellow, and that of the iodide is yellow. In organic solvents the colors become more intense. Fine crystals of methylviologen chloride appear almost white, with absorption maxima a t 260 and 377 mp; those of the bromide light yellow, absorption maxima a t 265 and 387 mp; and those of the iodide red, absorption maxima a t 277, 285, and 475 mp. Since neither the viologen ion itself nor t,he halide ions in ordinary ionic crystals absorb in the visible region, the absorption bands a t the longer waxe lengths must be due to charge-transfer. Qualitatively, the longest wave length of the charge-transfer band for the iodide crystal is also consistent to the strongest electron-donor property of the iodide ion. The methylviologen ion, (CH3. NCaH4.CsH4N. CH3)+2, is known to be a fairly good electron acceptor because of the stability of the reduced free-radical ion (C&. NCsH4. C6H4N CH3)+ discovered by M i ~ h a e l i s .When ~ aqueous solutions of methylviologen chloride and potassium ferrocyanide are mixed, a deep purple color develops instantly. The absorption spectrum of the mixture is distinctly different from those of either components or that of the intensely colored free-radical ion (CI13.NC6H4.C6H4N* CH,) +, and hence must also be of the charge-transfer type. Crystals of methylviologen ferrocyanide have indigo blue color with a broad absorption maximum near 600 mp. Aqueous solutions of methylviologen ferricyanide do not show charge-transfer spectrum. The composition of the charge-transfer complex formed between methylviologen and ferrocyanide ion in aqueous solution was determined from the contjnuous variation data shown in Fig. 1. Using 0.36 and -0.45 v. as the standard reduction potentials of ferricyanide and methylviologen,a respectively, it can be shown readily that the concentration of the intensely colored free-radical ion (CHI. NC6H4.C6% N.CH3)+ is negligible in all of the solutions used for these studies. Consequently, it may be concluded fram the data in Fig. 1 that the ratio of methylviologen to ferrocyanide in the charge-transfer complex in these solutions is 1:1. The absorption spectra of a number of mixtures of methylviologen chloride and potassium ferrocyanide in aqueous sodium chloride solutions of constant ionic strength. are given in Fig. 2 . Figure 2 shows that if the stoichiometric ratio of methylviologen to ferrocyanide is equal to or less than unity, the chargetransfer absorption maximum remains constant a t 530 mp, presumably due to a one-to-one complex. But, as the methylviologen to ferrocyanide concentration increases above unity, the absorption maximum slightly shifts to longer wave length. This observa-
-
(2) E . M. Kosower, “Charge-Transfer Complexes.” Vol. 111. “The Enzymes” (edited by Boyer, Lardy, and Myrbaok), Academic Press, New York, N. Y..1960, Chapter 13. (3) L. Miohaelia, Chen. Rev., 16, 243 (1935).
497
15
j,
+.’
2
IC
3 3
1
j ’ 0
I
0 0
450
550
500
650
600
Wave length (ma).
Fig. 1.-Continuous variation data on the formation of chargetransfer complex between methylviologen and ferrocyanide ion in aqueous solution a t 23”. Total stoichiometric concentration of methylviologen ahloride and potassium ferrocyanide = v f = 0.05 F . Total ionic strength wa8 maintained a t the constant value of 0.93 M by the addition of sodium chloride. Dois the optical density of the components in the absence of complex formation.
+
4
I
(
0
0.I
0.2 0.3
I
I
I
I
I
94 0.5 0.6 0.7’ 0.8 0.9 V/(P
+ f).
Fig. 2.-Charge-transfer spectra of methylviologen-ferrocyanide complexes in aqueous solutions a t 23’. 2) -1f = 0.05 F , total ionic strength := 0.93 M . As v/f increases above unity, the absorption maximum shifts to longer wave length, indicated by the broken straight line in the figure, due to the formation of the ternary complex ( CHa.;PVIC6HI.CsHaN.CHa)zFe(CN)p.
NOTES
498
tion suggests the existence of a new minority molecular species in the equilibrium mixture, which was later shown to be a two-to-one complex. The molecular formulas and formation constants of these complexes are determined by equilibrium studies described in the Experimental part of this note. The formation constants of (CH3.SC6H4. C5H4rUI.CH3)Fe(CN)6-2and (CHs.NC5H4. CbH4N-CH3)zFe(CS)6determined at 23' respectively. are 52 i 5 and 0.21 h 0.03 The first value is about ten times larger than the previously reported formation constants of chargetransfer complexes for the simple pyridinium ion.2 Experimental Determination of Equilibrium Constants.-For simplicity let V represent methylviologen and F represent ferrocyanide ion. The association equilibria may be written as
+ I' VF, K1 VF + V J_ VzF, Kz V
where K1 and KZare the apparent stability constants, i.e., the true equilibrium constants multiplied, respectively, by the appropriate activity coefficient ratios. Let v and f represent the total stoichiometric concentrations of V and F, z1 and zzthe equilibrium concentrations of VF and VzF, respectively. We have K1 =
21
(v - 21 - 222)(f
Kz
- 21 - x2)
x2
= x1(v
- x1 - 222)
(1)
(2)
The optical density, D, of the solution is
D
=
+
Z(Z~C~
XZC~)
(3)
where 1 is the length of optical path and €1, EZ are the molar extinction coefficients of VF and V*F, respectively. The molar extinction coefficients of V and F are negligible a t the wave lengths selected for the computation. In the presence of a large excess of V, 21
