NOTES
Downloaded by FLORIDA INTL UNIV on August 25, 2015 | http://pubs.acs.org Publication Date: December 1, 1958 | doi: 10.1021/j150570a041
Dee., 1958 value in 1 M cesium chloride is about 50% greater than that in 1 M lithium chloride. The trend is consistent with the positions of the alkali metals in the periodic table. The vapor pressure studies of Lannung4 indicate that the activity of water is nearly the same in the various salt solutions used, ranging from 0.964 in 1 M lithium chloride to 0.971 in 1 M cesium chloride. Therefore, although it appears in the above expression for Q, it need not be regarded as a changing factor. If the principal orange form is RH2Cl2,equation 3 shows that the mean activity coefficient of hydrochloric acid is involved. Since it has been shown6 that this quantity is constant in 1 M mixtures of lithium chloride and hydrochloric acid, acceptance of equation 3 would imply that the changing value of Q in the lithium series is due solely to specific effects upon the rhodamine B activity coefficients. A few measurements of Q in solutions of 0.1 M hydrochloric acid with varying concentrations of alkaline earth metal chlorides (from 0.2 to 0.8 M ) showed the same type of specific effect, ie., Q increased with ionic weight of the cation. For solutions containing 0.8M salt the values of Q were 0.090, 0.098, 0.098 and 0.100 for magnesium, calcium, strontium and barium, respectively. It is not possible to compare this series directly with the 1:l salts because of the impossibility of maintaining both ionic strength and chloride concentration constant simultaneously, but it may be noted that the variation is not as marked as in the 1:1 series. The abnormal values of Q when pyridinium chloride is present are probably due to a different sort of interaction between the cation and rhodamine B. The molar absorptivity of RHCl is the same as with the other cations, but the position of the absorption maximum is shifted to 560 m/J. Acknowledgment.-We are grateful to the Research Corporation and t o the du Pont Corporation for the support of this work. (4) A. Lannung, Z . physik. Chem., Al70, 134 (1934). (5) See H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolyte Solutions,” third edition, John Wiley and Sons, Inc.. New York, N. Y., 1958.p. 698.
T H E DISSOCIATION CONSTANT FOR THE CYCLOHEXAAMYLOSE-IODINE COMPLEX BY JOHN A. THOMA AND DEXTER FRENCH Department of Chemistry, Iowa State CoEEeoe, Ames, Iowa Received J u l y 68,1968
The reactions of cyclohexaamylose (a) with iodine-iadide solutions are of considerable interest as models for the more complicated starch-iodine reaction. It has been reported that a in I2 solutions can form an a12 complex in the absence of I- and an a13- complex in the presence of I-. The dissociation constant for the a11- complex (1) J. A. Thoma and D. Frenoh, J . A m . Chem. SOC., 80,8142(1958).
(2) H. Dube, unpublished Ames, Iowa, 1947.
Ph.D. thesis, Iowa State College
1603
0 10S
z
B
L O O C O N C OF T O T A L d IN M I L
Fig. 1.-Variation
I
I
e
of the absorbancy of a-12 system a t 480 mp at differing 01 levels.
was measured electrometrically by Dube, who also attempted to determine the dissociation constant for the a12 complex by partitioning IZbetween benzene and an aqueous a solution. Because the Schardinger dextrins are notorious complexing reagents,8 it seems likely that Dube measured an “average” dissociation constant (Le., for the dissociation of aIzand a benzene. Moreover, he did not take into consideration the interaction of I-, formed by the hydrolysis of I%,with the other chemical species in solution. A spectrophotometric study was therefore undertaken to measure the dissociation constant for the a 1 2 complex at room temperature in which HIOa was added to prevent competing equilibria involving I-. Consider the equilibrium CVIZ -
+ Iz [a!1[I~1/[01IzI= K
Then when IZis titrated with a until [ a L ]= [IPadded/2]
(1) [I2
free] =
log K = log [011/2 - 1121
(2)
- 1 2 added/Z] (3) where aI/nis the concentration of a added to make [12free] = [aIz]. or K =
The requirements necessary for a spectrophotometric determination of K are that the complex and the reagent held constant are the only absorbing species and obey Beer’s law, the law of mass action is operative, and there are no competing equilibria. Materials.-The preparation of 01 has been described elswhere. The per cent. hydration was measured polarimetrically .a Iodine and iodic acid were of the best grade commercially available and used without further purification. This spectral investigation was conducted in a Beckman DU Spectrophotomep in conjunction with capped one cm. silica cells a t 24 f 2
.
Experimental The optical densitiea of a series of a-12 solutions containing a constant concentration of Iz(2.45 X lo4 and 0.2 M HI03 were measured at 480 mp and plotted against the log of the concentration of total a! in Fig. 1. The blank was 0.2 M HIOi.
y)
Result The value of a l / z was determined graphically from Fig. 1 and the dissociation constant for aIz, utilizing equation 3, was calculated to be 1.07 f 0.10 x 10-4.
Library, ( 3 ) D. French, Adv. i n Carbohydrate Chem., 12, IS9 (1957).
