NOTES
July, 1958 THE ANION-EXCHANGE RESIN ADSORPTION OF ZINC(I1) FROM AMINECHLORIDE SOLUTIONS1 BY R. A. HORNE* Contribution o f the Department of Chemistry and Laboratory for Nuclear Science o f the Maesachusetts Institute o f Technology, Cambridge, Masaathwetts Received January 80, 1968
873
ethanol-water solution. The retention at high chloride concentrations is greatest for the smallest cation, CH3NH3+,and least for the largest, (CH& Nf. The similarity of the CH3NH3+curve with the NH4+curve reported earlier*suggests some orientation preference in the ion-pair formation process. The results of lowering the dielectric constant of the external solution by ethanol-water media is comparable for (CH3)4N+with those reported earlier6 for Lis, H+, K+ and Csf. The present observations are compatible with the conclusion that ion-pair formation of the type described previouslya and above is important in the resin phase only when the dielectric constant of the external solution is high. Presumably, by using a supporting electrolyte of sufficiently large cation, it should be possible to obtain an "ideal" adsorption curve for which the only important equilibrium at high chloride concentration is
A theory of the dependence of the adsorption of metal complexes from aqueous solution onto anionexchange resins on the nature of the cation of the supporting electrolyte has been described in previous papers.a-s This theory attributes the observed effect to ion-pair formation in the low dielectric constant medium of the resin phase or in both resin and external solution if the dielectric constants of these phases are comparable. The extent of ion-pairing between anionic metal complexes and the cation of the supporting electrolyte 2RC1 ZnCl," RzZnC14 2C1(1) depends, among other factors, on the size of the latter. Unfortunately the degree of hydration of and which lends itself im'mediately to analysis of the simple, relatively small cations, i.e., alkali and the type proposed by Coryell and Marcus.6 alkaline earth metal cations, results in ambiguities (6) C. D. Coryell and Y. Marcus, Bull. Res. Council Israel, 4, 7 in their relative size assignments. This difficulty (1954); Y. Marcus and C. D..Coryell, ibid., in press (1958). can be circumvented in part by the use of aminechlorides as supporting electrolytes. INFRARED SPECTRUM AND BARRIER TO Experimental INTERNAL ROTATION IN ETHYL The experimental procedurea involved in batch anionFLUORIDE1 exchange resin studies of the present type, using tracer radiozinc, have been described previously.8 The chloride BYEDWARD CATALANO AND KENNETH S. PITZER
+
+
concentrations were maintained by the addition of CHaNH2. HCl, (CH&NH.HCl, (CH&N.HCl, or (CHs)4NC1 as supporting electrolyte. Chloride was determined by rapid Volhard analyses.
Results and Discussion Figure 1 shows the adsorption of tracer radiozinc-
t
I"""'
30
A
Department of Chemistry, University of California, Berkeley, California Received February 14, 1968
The infrared spectrum of ethyl fluoride has been measured in the 200-300 cm.-l region in a PerkinElmer 12-C spectrometer with CsI prism. The observed spectrum, Fig. 1, shows distinct bands a t 243.5 and 227 cm.-l together with less distinct absorption in the region of the low frequency cut off the prism. 40
I
2
I
I
I
I
C H 3 C H 2 F lOcm CsI CELL
0
C 30-
10.. 0
\
J
0 LCH.), (CHs)~N NCI 0 (CH&NCI
In
50Vol %CiHsOH
0 200
210
220
230
240
250
260
2/ c 6 ' Fig. 1.-The
absorption spectrum of ethyl fluoride.
Herschbach2 obtained a value 3300 cal./mole for the torsional barrier in ethyl fluoride from the splitting of the first excited state lines in the microwave spectrum. He assumed the usual threefold cosine barrier function V = 1/zV3(1 - cos 39). (1) This researoh was a part of the program of Research Project 50 of the Amerioan Petroleum Institute. (2) D. R. Herschbach, J . Chem. Phys., 86, 358 (1956).
NOTES
874
Kraitchman and Daileys give the reduced moment of inertia as 4.359 X l O F 4 0 g. cm.2. These data and assumptions yield 243 and 226 cm.-l for the 0-1 and 1-2 torsional level separations, respectively. The agreement is most satisfactory and confirms the values 3300 d= 30 cal./mole for the barrier and 243 cm.-l for the energy of the first excited torsional state. The value 278 cm.-l had been proposed for the latter quantity on the basis of a very weak line in the Raman spectrum4 but the possibility that this line arose from an impurity also was mentioned. Kraitchman and DaileyS gave 300 j = 40 cm.-l for the energy of the first excited state from the temperature coefficient of intensity of satellite lines in the microwave spectrum, but it now appears t o be necessary to increase somewhat their estimated uncertainty. The calculated location of the 2-3 band is 207 cni. -I. The observed spectrum is consistent with this but is not reliable because of heavy absorption by the prism at this frequency. The thermodynamic properties were calculated by Smith, et u Z . , ~ and presented in itemized form so that it is easy to correct for the new barrier and the new moments of inertia. Table I gives the revised values .
Vol. 62
vent cooled to -78". This cell could be closed to prevent escape of glyoxal and the absorption of atmospheric moisture. All solvents were dried and purified by conventional methods. To minimize reaction of the glyoxal with the solvents and polymerization, spectra were recorded a t 0' in a Beckman DU spectrophotometer. The spectra in mixed solvents were determined by first preparing a heptane solution and recording the extinctions at the maxima. The cell was opened, a measured volume of the second solvent added, and the cell closed again. This procedure was repeated for the successive dilutions. The recorded extinctions were then corrected for the volume changes. It was determined that t h e loss of glyoxal vapor during this procedure was negligible. The spectrum of glyoxal vapor was determined in a 10 em. cell. Solutions of glyoxal in heptane, chloroform, dioxane and ethyl ether showed little change in spectral intensity in the time required for a determination (cu. 30 minutes). The change of intensity was appreciable for nitromethane, tetrahydrofuran and acetone solutions, but it was possible to determine the shape of the spectra by working rapidly. Solution of glyoxal in n-butyl alcohol, n-octyl alcohol, anisole, pyridine and acetonitrile resulted in an immediate reaction in which a white s o h appeared (probably glyoxal polymer). Even in the stable solutions a slight turbidity often was observed. A correction for this turbidity was made by subtracting the extinction a t 560 mfi where glyoxal absorption is negligible. Incorporation of the X4 dependence for this small scattering correction would only serve to exaggerate the effects and would leave the over-all conclusions unchanged.
Results.-Solvents
containing oxygen atoms re-
TABLE I move the structure in the glyoxal spectrum; Fig. 1 CALCULATED THERMODYNAMIC PROPERTIES FOR ETHYL shows dioxane spectrum. This may be compared FLUORIDE (CAL./DEG. MOLE)
with the heptane spectrum (Fig. 2, curve 1). The spectra of nitromethane, tetrahydrofuran and ethyl ether solutions are very similar t o that in 235.5 12.16 9.52 50.73 60.25 dioxane. Absorbance values are relative and n o 298.15 14.10 10.26 53.04 63.30 quantitative comparisons between the spectra are 400 17.60 11.70 56.26 67.96 possible. Chloroform does not remove the structure 14.67 600 23.45 61.57 76.24 completely. These effects are similar to those observed in biacetyl solutions.2 However, ether is (3) J. Kraitchman and B. P. Dailey, ibid., 28, 184 (1955). (4) D. C. Smith, R. A. Saunders, J. R. Nielsen and E. E. Ferguson, much more effective in removing glyoxal structure ibid., 20, 847 (1952). than biacetyl structure. This is consistent with the interpretation that interaction between solvent THE SPECTRA OF GLYOXAL SOLUTIONS1 oxygen and carbonyl oxygen atoms is responsible for the loss of structure. The steric interference to BY CLARENCE L. CARPENTER, JR., AND LESLIES. FORSTER close oxygen-oxygen approach is less in glyoxalether than in biacetyl-ether systems. Department o,+ Chemistrv, Universily o j Ariaono, Tucson, Brimma Received Februarv 8.4, 1968 In heptane the spectrum is shifted 300 cm.-l to the red compared t o the gas phase spectrum. This It has been found that solvents containing oxygen and nitrogen markedly diminish the intensity is the same polarization red shift shown by biof portions of the visible absorption spectrum of a~ety1.~&.4 The mixed solvent spectra are shown in Fig. 2-4. biacetyl.2 This study was initiated to determine For the mixed solvents heptane-dioxane and hepwhether analogous solvent effects exist in solutions of the simplest dicarbonyl, glyoxal. The gas phase tane-chloroform, the integrated intensity (21,000spectrum is much more structured than that of 27,000 cm.-1) is reduced progressively as the mole biacetyl and much of this structure remains in fraction of heptane decreases. The dilution technique for these spectra makes quantitative comhydrocarbon solvents. parison meaningful. No correlation between the Experimental integrated intensity and the refractive index exists. Glyoxal was prepared by the oxidation of ethylene with Discussion.-Solvent effects in biacetyl and glyselenium dioxideJ and was stored in uucuo a t -78". Samples oxal solutions are closely making it were prepared by warming the storage bulb and freezing the glyoxal into an absorption cell containing degassed sol- unlikely that hyperconjugation enhanced by solvent interaction is the cause of the biacetyl spectral (1) (a) This research was supported by a grant from the National changes, since hyperconjugation is not possible in Science Foundation. Reported in part at the 33rd annual meeting glyoxal. The small blue shift is comparable to of the Southwestern and Rocky Mountain Division of the A.A.A.S,, Tucrson, Arizona, hlay 1, 1957; (b) taken in part from the M.S. thesis that observed for pyrazine in propionitrileof Clarence L. Carpenter, Jr., May, 1957. hexane mixture.6 12) (a) L. 8. Forster, J . A m . Chem. Soc., 77, 1417 (1955); (b) J. Chem. Phvs., 26, 1761 (1957). (4) N . 8. Bayliss, J. Chem. Phya.. 18, 292 (1950). HQ
T ,"K.
CPO
- HoQ T
FO
- Hog T
a
(3) J. C. Calvert and G. Lane, J . Awa. Chem. Soc., 76, 856 (1953).
(5) G. Brealey and M. Kaeha,
J, Am. Chem. SOC.,77, 4462 (1955).
