NOTES
Feb., 1958
253
ions. The lanthanum-cerous and cerous-chromic exchanges were therefore investigated. The order of ion-resin affinity for these trivalent ions is La > Ce > Cr. The crystal radii reporjed by Paulin Equilibrium constant 4% 8% 16% for trivalent ions are: La, 1.15 A.; Cr, 0.64 DVB DVB DVB Exchange rare earths, 0.90 f 0.05 8. The selectivity data 1.40 1.71 Ca-Cu 1.26 follow the expected order, since for ions of the same Ca-Ni 1.19 1.38 1.81 valence type the ions of smaller crystal rad'ii have 0.91 1.06 1.03 Ni-Cu the larger hydrated radii. The aqueous solutions Ni-Cu (calcd.) 1.06 1.01 0.94 used for the cerous-chromic exchange were pretively. These data place beryllium ion (Table I) pared using cerous and chromic nitrates. Chromic in approximately the same position in the selectiv- nitrate also was used to convert the resin to the ity scale as manganous ion. The alkaline earth ch,rqmic form since it was found by analysis that ions are thus in the order: Mg, Be, Ca, Sr, Ba. This when chromic chloride was used, the resin was conposition is substantiated by the maximum water verted to the CrCl++ rather than the Cr+++ form. uptake data in Table 111. The 4,s and 16% berylMaximum Water Uptake Data.12--The maximum lium resinates swell more than the corresponding water uptakes (swelling) of the resins in the various calcium resinates but less than the magnesium resin- ionic forms are listed in Table I11 as grams of ates. The selectivity and swelling data for beryl- water absorbed per equivalent of resin. Although lium ion may be accounted for if beryllium exists no exchanges involving the CrC1++ ion were in the resin partially as BeC1+ ions or, because of investigated, the water uptake of the resin in this hydrolysis, as BeOH+ ions instead of as the di- form also was measured. The water uptake data valent Bef2 ion. for the trivalent ions correlate well with the The behavior of barium ion in exchanges with equilibria and crystal radii data given above. The both the divalent beryllium ion and the trivalent variation in water uptake for ions of the different cerous ion are of interest in that the affinity of the valence types as a function of the DVB content of barium ion for the resin decreases a t high barium the resins is of interest. The resins swell to such a loadings. This behavior becomes more pronounced concentration (equivalents of resin per 1000 grams as the cross-linkage of the resin increases and results of water) that the water in the resin phase is in in equilibrium quotients which are smaller for 16% equilibrium with that in the pure aqueous phase. DVB resins than for 8% DVB and 4% DVB resins This means that the product of the electrolyte, i.e., of high barium ion content. resinate, concentration and osmotic coefficient must be constant. For the lower cross-linked resins, the TABLE 111 OF DOWEX 50 RESINS IN VARIOUS trivalent resinates, being n, 3 electrolytes13 have MAXIMUM WATER UPTAKE lower osmotic coefficients than the divalent and IONIC FORMS (G./EQUIV.) univalent resinates and therefore imbibe lesser 4% 8% 16% DVB DVB DVB amounts of water. As the DVB content of the "SOH+ 356 158 92 resins increases, the swelling per equivalent of resin Mn+2 305 168 111 is almost the same for all valence types. This is Be +2 307 169 120 probably due to the greater affinity of polyvalent CrC1t2 265 168 133 ions for their water of hydration.
TABLEI1 TRIANGULAR COMPARISONS INVOLVING CUPRIC, NICKEL AND CALCIUM ION
252 206 200
Cr + 3 Ce+3 La +3
163 138 133
116 111 100
Trivalent Ions.-Two exchanges involving cerous ion, the cerous-silver and barium-cerous exchanges, were investigated to determine the position of cerous ion on the selectivity scale. The values reported in Table I for cerous ion are the averages for these two exchanges. The equilibrium constants for these two exchanges were calculated for reactions represented by the equations '/&e+++
and '/SBa++
+ AgRes = Ag+ + '/aCeRes
+ l/&eResa
=
'/&e+++
+ '/,BaResn
(4) (5)
This conforms with the conventions used in exchanges of univalent and divalent ions4 and permits the direct addition of these equations and those of other exchanges, and the consequent multiplication of equilibrium constants which is necessary for triangular comparisons. lu*iAfter determining the position of cerous ion in the selectivity scale by comparing it with previously studied ions, it was possible to use this ion as the reference ion in exchanges with other trivalent
f:
(11) L. Pauling, "The Nature of the Chemical Bond," Cornell University Press, Ithaca, N. Y., 1948, p. 350. (12) The experimental procedure for the determination of the maximum water uptake of a resin is described by 0. D. Bonnor, W. J. Argersingcr and A. W. Davidson, J . A m . Chem. Sac., 7 4 , 1044 (1952). (15) 0. D. Bonner, V. F. Holland and L. L. Smith, THISJOURNAL, 60, 1102 (1956).
SPECTRAL INVESTIGATION OF THE COMPLEXES OF ETHYL IODIDE WITH BROMINE AND WITH IODINE' BY LEONARD I. KATZIN AND ROBERT L. MCBETH Argonne National Laboralory, Lemont, Illinois Received May $9, 1967
Keefer and Andrews in 1952 reported2 that solutions of iodine or bromine in organic iodides or bromides displayed pronounced light absorption in the 300-350 mp region. This absorption,, which was lacking in the spectra of halogens in solvents such as carbon tetrachloride or heptane,. was re(1) Based on work performed under thc auspices of the U. 6. Atomic Enercy Commission. (2) R. R'I. IZeefer and L. J. Andrcws, J . A m . Chem. Suc., 74, 1891 (1962).
NOTES
254
garded as evidence for formation of a 1:1 molecular complex, and equilibrium constants for such complexes were estimated from spectrophotometric data. The workers did not find it possible to locate the peak of the absorption, but made their measurements on the rising slope of the curve. By physicochemical methods Jepson and Rowlinson3 have given clear evidence that molecular complexes are formed strongly between, for example, ethyl iodide and iodine. We have investigated the ethyl iodide-halogen complexes spectrophotometrically, and through the device of using short path lengths, have determined the location of the peaks for the complexes. The maxima are rather broad, but for the bromineethyl iodide complex in carbon tetrachloride the peak comes at 270 mp, and for the iodine-ethyl iodide complex the peak is at 275 mp in bdth carbon tetrachloride and n-heptane. Keefer and Andrews2 had inferred that the peak would be found in the 300-283 mp region. We have also made quantitative measurements on the association constantextinction coefficient product at the wave length of the maximum absorption for the ethyl iodideiodine complex which indicate the absorption to be very intense (extinction coefficient of the order of 45,000). Experimental Stock solutions of iodine, ethyl iodide and bromine were made in n-heptane or CC14 (CClr alone for bromine), and aliquot portions were taken for ex erimental solutions, which were then made up to volume. &he stocks were made in sufficiently small portions that repeats of experiments could be made with fresh stocks, to eliminate stock compositions as a source of systematic error. For an individual spectrophotometric measurement, the desired solution wa8 made in a 10-ml. volumetric flask. Observation cells used were 1-cm. square quartz cells, with a quartz spacer to give a 0.1 cm. path length. For the uantitative measurements with ethyl iodide and iodine, io&ne concentrations ranged from 0.0015-0.017 M , and ethyl iodide from 0.012-0.066 M. Carbon tetrachloride was reagent grade. The heptane used was Eastman white-label, further treated with sulfuric acid and fractionally dist,illed, to remove impurities absorbing in the ultraviolet. When tested, iodine in n-heptane showed a slightly higher absorption than in carbon tetrachloride, in the region 240-300 mp, but there was no suggestion of any absorption peak in this region. I n both solvents there appears a curvature of the absorption concentration relation in the 275 mp region (most recently discussed by Keefer and Allen'). The spectrophotometric blanks (Beckman DU spectrophotometer) consisted of the solvent with a concentration of ethyl iodide equal to the concentration in the experimental cell. The absorption due to complex formed is given then by the apparent absorption of the iodide plus halogen minus the absorption due to halogen alone, read from the calibration curve above.
Results and Discussion For weak complexes, such that the fraction of the reactants consumed to form the complex is nfgligible, the exact form of the equilibrium equation K = r/(A - X ) (B
-X)
(1)
reduces to the approximate expression K
=
x/AB
(2)
Measuring x spectrophotometrically with unit path length, ex = D , the optical density at the wave length used, or (3) W.B. Jepson and J. 9. Rowlinson, J . Chem. SOC.,1278 (1956). R.M. Keefer and T.L. Allen, J . Chsm. Phye., 26,1059 (1956).
(4)
Vol. 62 D/AB = K E
(3)
If both sides of this equation are inverted, it becomes A B / D = l/Ke, I n either form it is clear that from this equation one can obtain only the Ke product. Another or different relation is needed to obtain the equilibrium constant and the extinction coefficient of the complex explicitly. Keefer and Andrews,2 starting from the exact equilibrium equation, derived and used in their experiments an equation which, converted to our notation, may be written ( A B I D ) [1/(A B - $11 = ( l / e ) [ l / K ( Af B - X ) 11 = ( 1 / K e ) M A B - % ) I f I/€ It may readily be seen that when on the right-hand side 1/e becomes negligible with respect to the other term, the equation reduces to the one we have given above. This in turn requires that 1/K. (A B - z) be large with respect t o unity. From the values for the equilibrium constant given by Keefer and Andrews, working at ( A B ) values of 0.1-1, a change in the quantity plotted of the order of 20-25%, between the extremes, could have been found (their data are unfortunately not available). At our values of (A B) of not more than 0.06, the maximum observable difference from infinite dilution was 2%) which is within the limits of the experimental error. The use of eq. 3 is therefore justified. Thus, based on 41 determinations, over an ( A B) range of 0.015 to 0.06, and a to 4.6 X range of AB values from 0.45 X the K E product for ethyl iodide-iodine in heptane was (1.66 f 0.02) X lo4, and in carbon tetrachloride, with about the same error limits, 1.33 X lo4. These values are for the peak wave length, which was 275 mp in both cases. The ethyl iodidebromine complex in carbon tetrachloride showed a similar absorption peak at 270 mp, but no quantitative equilibrium measurements were made. Using the Keefer and Andrews equilibrium constants, the extinction a t the peak is essentially the same for the two solvents, 45,000-46,000. However, as may be seen in a full spectrum between the peak and say 350 mp, a t longer wave lengths than the peak, the extinction in the carbon tetrachloride solution is greater than in heptane-almost by a factor of 1.5 at 340-345 mp, according to the values of Keefer and Andrews. Since the apparent extinction is so high, such a large difference off the peak seems anomalous. This might mean that there is in actuality a small difference in the peak wave lengths in the two solvents, masked by the fact that measurements are being made against a heavy ethyl iodide background absorption, with a peak at about 258 mp. The visible peak of iodine itself conies a t about 523 mp in heptane, and 516 mp in carbon tetrachloride. Addition of 0.037 M ethyl iodide to 0.007 M iodine in both of the solutions shifts the peak perhaps 2-3 millimicrons. Insofar as the difference in equilibrium constants is concerned, the apparent lesser complex stability in carbon tetrachloride may reflect competition of the halogenated solvent with the ethyl iodide for the iodine. The value of the absorption peak wave length
+
+
+
+
+
+
+
