INFRARED SPECTRA OF COMPLEXES AgN03. CH&N
AND
963
AgN03.2CH3CN
Infrared Spectra of the Complexes AgN0,-CH,CN and AgNO,. 2CH,CN
and Their Solutions in Acetonitrile
by George J. Jam, Malcolm J. Tait, and Jurg Meier Department of C h b t r y , Rensselaer Polytechnic Institute, Troy, New York (Received September $0, 1966)
The complexes AgN03-CH&N and AgN03.2CHaCN have been isolated from silver nitrate solutions in acetonitrile. Their infrared spectra are reported and compared with that of solid silver nitrate and its solutions in acetonitrile of concentration up to 9.6 moles 1.-'. The results in solution can be understood in terms of the formation constants of the complexed or solvated ions Ag(NCCH&+ and Ag(NCCH3)+. Evidence is also presented that in solutions of concentration less than 3 molesl.-l there is an equilibrium between free and paired ions. The spectrum of the cation-anion pair suggests that the ions are in contact in CH&N solutions and not solvent separated as in aqueous AgN03 solutions.
Introduction Little is known about the behavior of electrolytes in concentrated solutions. Attempts to extend the dilute solution model have met with only moderate success.1-2 A thorough investigation of a suitable system by a number of experimental methods seems essential. Such an investigation is in progress in these laboratories. The solubility of silver nitrate in acetonitrile is about 9.6 moles la-' at room temperature corresponding to a mole fraction of silver nitrate of 0.43. Diffusion data are available3 for this system up to 5 moles 1.-'. Viscosity and conductance measurements* detect ion pairs at 0.1 moles I.-' and increasing concentrations of triple ions from 1 to 6 moles 1.-'. Infrared and Raman spectras can be said to confirm that ion pairs are present from 0.2 to 9 moles 1.-' and in addition they established the existence of the solvated ion Ag(NCCH3)2+in these solutions. In the present work it has proved possible to isolate from solution the complex AgNOs. 2CH3CN and also one other of composition AgN03.CH3CN. The purpose of this work has been to use the infrared spectra of these solid complexes as fingerprints in order to gain a better Of the cation-so1vent interactions in solution. It has also proved possible
nitrate and its aqueous solutions suggests that in acetonitrile the paired ions are in contact.
Experimental Section The purification of AgN03 and CH&N has been described previously! Deuterated acetonitrile, CD&N (98%), was used as received from Rferck Sharp and Dohme. Solutions of silver nitrate in acetonitrile were prepared and stored under Nz. Crystals of AgN03.2CH3CN were prepared by cooling a solution of silver nitrate in acetonitrile of mole ratio 1 :3 to - 10'. It was found necessary first to seed the solution or to initiate the crystallization by cooling rapidly to -70' for a short time. The crystals were found to melt at a temperature somewhat less than that of the room, so that it was necessary during their isolation to operate at a temperature of about 0'. The bulk of the mother liquor was removed and the crystals were washed twice with double their volume of anhydrous cold ether. The crystals were cooled to
c*
A* Krausl J. P h ~ 8them.* . 589 673 (1954). (2) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Butterworth and co, Ltd., London, 1959, (3) G . J. Janz, G. R. Lakhfinarayanan, and M. P. Klotsk, J . Phys. Chem., 70,2562 (1966).
Volunte 71, Number 4 March 1967
964
-70' to prevent evaporation of the acetonitrile component and the last traces of ether were removed in a stream of dry nitrogen. The crystals were dried to constant weight. Crystals of AgN03.CH3CN and AgNO3-CD3CN were isolated from their saturated solutions which were found to crystallize almost completely at 0'. They did not melt at room temperature. They were filtered off in a drybox, washed with ether, and dried to constant weight as described above. The composition of the crystals was determined by analysis both for silver and for acetonitrile. Both methods agreed to within 1%. Silver analysis was effected by precipitation of silver chloride in water and the acetonitrile content was gained from the weight loss after pumping under vacuum for 8 hr at room temperature. Acetonitrile evaporates readily from the complexes. Crystals of both complexes, AgN03. CH3CN and AgT\'O3+2CH3CN,are unstable at atmospheric pressure: a weight loss of about 0.01%/min is observed in an open system. They were therefore stored under Nz in a closed bottle at -10'. The melting points of the crystals were determined by differential thermal analysis. The melting point of the 1:2 complex was 8.5' and that of the 1 : l complex was 43'. However, whereas the melting point of AgX03-2CH3CYcrystals was reproducible several times with the same sample, this is not the case for AgN03. CH3CN crystals because of loss of acetonitrile on melting. Some interesting results were obtained from the data measurements on a silver nitrate solution in acetonitrile of composition 1:2. On cooling slowly, the solution crystallized at about -30'. These crystals had a well-defined melting point of 8.5' so that in fact the 1:2 complex could be prepared from its "molten state." For the infrared studies a Perkin-Elmer 221 doublebeam grating spectrometer was used. The instrument was calibrated from the spectra of liquid and solid CH3CN6z7and CD3CN.7J' I n previous work5 it had been found that AgCl windows are the most suitable to contain the solutions in the infrared cells. Cells were used which prevented any contact between the solutions and metal surfaces. For the more dilute solution studies, the cells employedg could be sealed very effectively to prevent evaporation of solvent. It was observed that solid AgN03 reacted even on simple mixing with KBr. KBr pellets were therefore not used foy any of the solid spectra. Melts were prepared with Xujol and hexachlorobutadiene. Tests showed that neither of these materials dissolved the acetonitrile out of the complexes. Because of its low melting point, the infrared specThe JOUTTU~ of Physical ChhabtTy
G. J. JANZ, M. J. TAIT,AND J. MEIER
tra of 1:2 complex had to be measured at temperatures less than 8.5'. The sample cell was mounted on a copper rod penetrating the base of a dewar-like vessel. The latter served as a cooling well which could be filled with liquid Nz. This assembly was finally placed inside a glass vessel and this outer jacket could be evacuated or filled with argon in order to prevent condensation of atmospheric moisture on the cold windows of the cell. The complete assembly has been described elsewhere in detail.l0 The temperature of the sample closely approached liquid KZtemperature (77OK). The spectrum of the 1:2 solid measured as a Nujol mull was found to be identical with that of a 1:2 solution cooled slowly to liquid Nz temperature. The latter technique, being the easier of the two, was therefore the one more often employed to determine spectral details of the 1:2 complex.
Resdts Thenitrate bands observed in the infrared spectra of silver nitrate and the solid complexes are recorded in Table I. For a11 three solids the spectral region from 1200 to 1500 cm-l was the most difficult to investigate. For the complexes, crystals of AgNOa.CD3CN and AgN03.2CD3CN had to be used because CHaCN absorbs strongly at 1376 and 1443 cm-'. Absorption by the nitrate is so intense that a reliable spectrum could not be obtained from a film of pure 1:2 complex. I n all cases, hexachlorobutadiene mulls were therefore used since hexachlorobutadiene itself contributes only a number of weak bands to this region of the spectrum. Except for the AgN03 absorption and the 1200-1400cm-l absorption by the 1 : 1 complex, the nitrate bands reported for the 1200-1500-~m-~region could be seen superimposed on the hexachlorobutadiene contour. For AgNOa it was necessary to subtract the contribution of the hexachlorobutadiene from the observed contour. The resulting nitrate band was very similar to that shown by Vratny" if the shoulder at -1400 cm-l, which may be due to an inadequate compensation for Nujol absorption in his spectrum, is neglected. The maximum absorption is a t 1370 cm-l and a shoulder is quite evident at 1250 cm-'. These frequencies are recorded in Table I, but because one band is almost sub-
(6) P. Venkateswarlu, J . Chem. Phys., 19, 293 (1951). (7)D.E.Millipan and M. E. Jacox, J. Mol. Spectry., 8, 126 (1962). (8) J. C . Evans and H. J. Bernstein, Can. J . C h a . . 33, 1746 (1955).
(9) ."Throwaway Infrared Liquid Cells," Limit Research Corp., Darien, Conn. (10) K. J. Eisentraut, Ph.D. Thesis, Rensselaer Polytechnic Inatitute, 1964. (11) F. Vratny, Appl. Spectry., 13, 59 (1959).
INFRARED SPECTRA OF COMPLEXES AgNOz.CH&N
AND
965
AgN03.2CH3CN
Table I: Infrared Frequencies (cm-1) of the Nitrate Ion for Crystals of AgNOs, Ag(CHaCN)NOa, and Ag(CHaCN)zNOa P6(Al)
AgNOs AgNOs*CHsCN AgNOs CD3CN AgN03 * 2CHsCN AgNOs. 2CD3CN
708 710
vs(Bz)
vdBi)
724 718 720 721 721
801 809 809 816 815
vdAd -1015 sh -1010 sh 1023 sh 1014
vi(Ai)
-
V4(BZ)
2 vs
12501”
1370?”
1200
1400
1460 w
1369
1444 w
1214 w, 1224 w
The “?” indicates that the broad band in the 1200-1400 em-’ region is not clearly resolved and that the above values are estimates.
merged in the other, they cannot represent the true vmsx. A similar calculabion for the 1 : l complex revealed absorption in the 1200-1400-~m-~region. This band may consist of unresolved components. Comparison of the spectra of the complexes with that of pure acetonitrile shows that the major changes occur in the CN and CC stretching frequencies. These are recorded in Table 11. Figure 1 illustrates the concentration dependence of the CN stretching frequencies in concentrated acetonitrile solutions.
\
(6)\
V!
\
Table 11: CN and CC Stretching Frequencies (cm-1) of the Complexes
CDsCN AgNOa. CDsC?J CHICN AgN03 CHsCN AgNOs.CHaCY CHsCN AgNOs * 2CH3CN AgNOa * 2CDsCN CDsCN
Temp
VCN
vcc
20°C 20°C 20°C 20OC 77°K 77 “K 77°K 77°K 77°K
2263 2274 2254 2264 2261 2252 2269 2278 2260
833 917 924 917 929
2290 2 2 8 0
2270
2260
2 2 5 0 cm-‘
Figure 1. The concentration dependence of the CN stretching frequencies for AgN03-CHsCN solutions. The mole fractions of AgNOa are (a) 0.00, (b) 0.33, and (e) 0.43.
832
In the previous studyS it had been found that the asymmetric stretching vibration of the nitrate ion was no longer degenerate in solutions of silver nitrate in acetonitrile. In solutions of concentration from 3 to 9 moles l.-l, absorpt~ionmaxima were found at 1300 and 1425 cm-l. However it appeared that at -0.2 mole 1.-’ the low-frequency band shifted to 1350 cm-l. This concentration dependence has now been more thoroughly investigated and the results are shown in Figure 2. Discussion Solid Spectra-VibrationaE Assignments. It can be seen from Table I that four frequencies are observed in the infrared spectrum of silver nitrate. If the point group of the nitrate ion is D3hl only three
infrared frequencies (vz, Az”; vtt E’; v4, E’) should be observed. It is apparent the symmetry of the nitrate ion is lower than Dah in silver nitrate since one of the doubly degenerate modes (v3, E’) has split into two bands, 1250 and 1370 cm-l, respectively. Similarly in the spectra of the 1 : l and 1:2 complexes of silver nitrate with acetonitrile, there is evidence that the symmetry of the nitrate ion is lower than D3h. The symmetrical stretching vibration at -1016 cm-I is infrared active in both complexes, a definite splitting of a degenerate made is detected in the 700-cm-I region for the 1 : 1 solid, and, as in the spectrum of silver nitrate, the absorption band of the 1:2 solid at -1300 cm-l is split. The symmetry of the nitrate-cation pair in all three of the solids in Table I may therefore be either Czvor C,. For each of these point groups all six fundamental vibrations should be infrared active. The classification of the fundamental vibrations of the Volume 71,Number 4
March 1967
G. J. JANZ, M. J. TAIT,AND J. MEIER
966
4
'
I
I
I
1500 1400 1300 1200
cm-1
Figure 2. The concentration dependence of t h e NO2 bands between 1250 and 1450 cm-1 for AgN03-CD&N solutions. T h e concentrations of AgNOa are (a) 0.02, (b) 0.2, (c) 1.3, and (d) 2.7 moles 1.+.
Czv point group is given in Table I according to the assignments developed by Topping12 from a study of metal-unidentate nitrato complexes. This anion-cation pair model is unquestionably an oversimplification. It has been shownla to be preferable to interpret the spectra of solid alkali metal nitrates in terms of their crystal structure. This approach has been used successfully for fused silver nitrate14 t o support the hypothesis that local order prevails in the melt. However in the absence of crystallographic data for the silver nitrate complexes with acetonitrile, no detailed treatment can be attempted. It is interesting that for silver nitrate no bands were recorded at -1450 or -1015 cm-1 which could be identified with the 2v5 or v2 vibrations. However Katzinl6 has reported an infrared band at 1471 cm-' for this salt. The presence of the 1015-cm-l absorption in only the spectra of the silver nitrate complexes with acetonitrile is analogous to the infrared resultslB for Co(K03):! and Co(SOp)~.2H20where only the dihydrate absorbs at this frequency. I n the spectra of the complexes, the number of fundamental frequencies observed for the acetonitrile component was the same as in the spectrum of pure acetonitrile. CH3CN must therefore retain the point group CaVsymmetry in the complexed state. X-Ray analysis" of CH3CN.BF3has shown that the CNB angle is 180°. The donor-acceptor bond in this'* and otherlg acetonitrile Complexes is thought to be formed by bhe lone The Journal of Physical Chemistry
pair of electrons in an sp orbital on the nitrogen atom. Consequently it would be expected that the complex ion [AgNCCHa]+ also has a linear Ag-N-C-C skeletal structure. However the NAgN angle in the 1:2 complex ion [CH3CNAgNCCH3]+ need not be 180' to preserve the CaVsymmetry of the acetonitrile molecules. From Table I1 it can be seen that the CN and CC bands of acetonitrile are shifted to higher frequencies in the complexes. Thus AVCN = 18 cm-l and AVCC = 12 cm-' for the 1:2 complex, whereas in the 1:1 complex AWN = 11 cm-l and Avcc = 7 cm-'. The fact that the CN frequency increases in acetonitrile complexes has been shownm by normal coordinate analysis to be a result of the increase in the force constant of the CN bond and can be explained18 by a change in hybridization of the nitrogen orbital contributing to the u b o d of the CN group. I n the 1: 1 complex the C N frequency is lower than in the 1:2. With SnCL, the same order of C N frequencies is observedlg with propionitrile complexes but the order is reversed in the acetonitrile complexes. It was found20 in the normal coordinate analysis treatment of acetonitrile complexes that the lower the CN frequency, the weaker is the donor-acceptor bond. The weaker Ag-N bond in the [AgNCCH3]+ ion may be due to the more extensive interaction in the anion-cation pair of the 1: 1 solid. Some evidence of this is recorded in Table I where the nitrate band in the 700-cm-' region of the spectrum, which, for a NO3- ion of D3h symmetry arises from a degenerate vibration, is split for the 1: 1solid but not for the 1:2. Solution Spectra-Vibrational Assignments. The spectra of concentrated solutions in the region of the CN absorption is given in Figure 1. For the solution of mole fraction 0.33, three bands are observed at 2254, 2265, and 2271 cm-l. The first two of these frequencies correspond exactly t o those of the C N bands of pure acetonitrile and AgN03 .CH3CNsolid at room temperature and the third frequency is 2 cm-' higher than that ~
~~
~~
(12) G. Topping, Spectrochim. Acta, 21, 1743 (1965). (13) (a) K. Buijs and C. J. H. Schutte, ibid., 18, 307 (1962); (b) C. J. H. Schutte, 2. Physik. Chem. (Frankfurt), 39, 241 (1963). (14) J. P. Devlin, K. Williamson, and G. Austin, J . Chem. Phys., in press. (15) L. I. Katzin, J. Imrg. Nucl. Chem., 24, 245 (1962). (16) J. R. Ferraro and A. Walker, J . Chem. Phys., 42, 1278 (1965). (17) J. L. Hoard, T. B. Owen, A. Buzzell, and 0. N. Salmon, Acta Cryst., 3 , 130 (1950). (18) V. N. Filimonov and D. S. Bystrov, O p t . Spectry., 12, 31 (1962). (19) V. H. A. Brune and W. Zeil, 2. Naturforsch., 16a, 1251 (1961). (20) K. F. Purcell and R. S. Drago, J . Am. Chem. SOC.,88, 919 (1966).
INFRARED SPECTRA OF COMPLEXES AgN03. CH&N
AND
of AgN03.2CH3CN solid a t liquid nitrogen temperature (see Table 11). Freezing acetonitrile t o liquid nitrogen temperature decreases the CN absorption by 2 cm-l. One can therefore with confidence interpret the origin of the bands observed in solution of mole fraction 0.33 as the CN vibrations of the species CHICN, [AgNCCH3]+,and [Ag(NCCH&]+. The relative intensity of these bands is concentration dependent. It can be seen from Figure 1 that at-a mole fraction of silver nitrate of 0.43 there is no “free” solvent absorption band. I n very dilute solutions, mole fraction less than 0.05 ( i e . , c < 1 mole L-’), the species [AgNCCD3]+ is no longer detectable; neither is the concentration of the [Ag(NCCD&]+ ion sufficient for observation at mole fractions less (ie., c < 2 X than mole I.-l), in which solutions only the “free” solvent can be detected. These results are explicable in terms of the equilibria Ag+ 1- CH3CN IJ[AgNCCH3]+ [AgNCCHp]+
+ CH3CN
[Ag(NCCH3)2]+
If a solution of mole fraction 0.33 in which all three species are detectable at room temperature is cooled, then the concentration of “free” acetonitrile decreases. Both equilibria must therefore be shifted to the right. To obtain the spectrum of the solid AgN03.2CH3CN complex it was therefore necessary to cool this solution slowly t o liquid nitrogen temperature. Some very accurate measurements of the proton magnetic resonance spectra of silver nitrate solutions in acetonitrile have recently been publishedS2l They are in adequate agreement with our previous work.22 From their results, Schneider and Strehlow calculate solvation numbers of 4 for the Ag+ ion and 2 for the NOB- ion. I n view of our infrared results it would seem possible to rationalize the concentration dependence of the chemical shift in terms of the formation constants of [AgNCCH3]+and [Ag(NCCH&]+. For the unperturbed nitrate ion with D3h point group symmetry, the assymetric stretching vibration ( v 3 , E’) occurs a t 1384 cm-I. This degeneracy is in aqueous Ca(N01)2 where the cation-anion pair has CWsymmetry. Two bands are found at 1355 and 1435 cm-’. For silver nitrate solutions in acetonitrile of concentration between 3 and 0.2 moles I.-l there are three absorption bands in this region of the spectrum. The concentration dependence of these bands is shown in Figure 2. The relative intensity of the high- and low-frequency bands decreases as the concentration falls until a t 0.02 mole I.-‘ only the central band is detectable. We therefore assign the bands at 1306 and 1411 cm-l to cation-anion pairs of symmetry CZv
AgN03.2CH3CN
967
point group and the central band at 1353 cm-l to NO3- ions (D3a)in equilibrium with ion pairs. The structure of these ion pairs is of interest. I n solutions of very high ~oncentration,~ 3-9 moles L-I, only the ion-pair absorptions a t 1300 10 and 1425 zt 3 cm-I are observed. In the infrared24and Raman s p e ~ t r a ~of ~fused s ~ ~ silver nitrate, bands are detected a t 1310, 1395 and 1280, 1410 cm-’, respectively. An interpretation of the fused-salt spectrum of molten silver nitrate based on normal coordinate vibrational analysis has been a d v a n ~ e d ~in~which p ~ ~ this observed splitting correlates with a contact ion-pair model. In aqueous AgN03 soIutions the frequencies of the equivalent Raman lines (at 1340 and 1420 cm-’ in a saturated solution,n -8 moles I.-’) show marked concentration dependence;28 this has been attributed to the existence of solvent-separated ion pairs. Irish and Walrafen23present further evidence that ion pairs (in Ca(N03)2solutions) giving rise to the concentrationdependent Raman frequencies at 1358 and 1450 cm-l are solvent-separated ion pairs. I n molten AgN03, a contact ion-pair model appears adequate for a theoretical analysis of the observed spectral properties. It is of interest to note that the anion-cation pair frequencies in acetonitrile solutions, at -1300 and -1415 cm-’, correspond closely to values for AgN03 in the molten state. A further observation is that the above ion-pair frequencies remain almost invariant over the entire concentration range studied (0.2-9 moles l.-l, Figure 2) ; this is in marked contrast to the spectra of aqueous AgN03 solutions. The spectrum of the cation-anion pair thus suggests that the ions are in contact in CH3CN solutions and not solvent separated as in aqueous AgK03 solutions. This feature may be a significant factor to be considered in any complete analysis of the physicochemical properties of such nonaqueous electrolytes.
*
Acknowledgments. We thank Professor Dr. H. Strehlow (Gottingen) for a copy of the Ph.D. Thesis of H. Schneider and Drs. J. P. Delvin (Oklahoma State University), G. E. Walrafen (Bell Telephone Laboratories), and D. E. Irish (University of Waterloo) for (21) H. Schneider and H. Strehlow, Z. Physik. Chem. (Frankfurt), 49, 44 (1966).
(22) Erroneous magnetic susceptibility corrections were applied t o our results in ref 5. (23) D. E. Irish and G. E. Walrafen, J . C h a . Phys., in press. (24) J. K. Wilmshurst and S. Senderoff, ibid., 35, 1078 (1961). (25) 5. C . Wait and A. T. Ward, ibid., 44, 448 (1966). (26) 9. C . Wait, A. T. Ward, and G . J. Janz, ibid., 45, 133 (1966). (27) R. E. Hester and R. A. Plane, J. Am. Chem. Soc., 3 , 769 (1964). (28) H. Lee and J. K. Wilmshurst, Australian J . Chem., 17, 943 (1964).
Volume 71,Number 4 March 1967
T. F. LIN, S. D. CHRISTIAN, AND H. E. AFFSPRUNG
968
providing prepublication copies of their article^.^^.^* This work was made possible in large part by financial
support from the U. S. Atomic Energy Commission, Division of Chemistry, Washington, D. C.
Self-Association and Hydration of Ketones in Carbon Tetrachloride1
by T. F. Lin, S. D. Christian, and H. E. Affsprung Department of Chemistry, University of Oklahoma, Nmnaan, Oklahoma (Received October 6, 1966)
The self-association and hydration of acetone, 2,3-butanedione, and acetylacetone in carbon tetrachloride solution has been investigated by partition, solubility, and isopiestic techniques. The results indicate that the complex species ketone monomer monohydrate, ketone dimer, and ketone dimer monohydrate are needed to obtain a satisfactory explanation of all the data. A study of acetone in hexadecane was made to substantiate the observation of dimerization of ketones in nonpolar solution. The dimers are assumed to be formed by dipole-dipole interactions. The basic strengths of the ketones were found to be in the order acetone > acetylacetone > 2,3-butanedionea Equilibrium constants are reported at 15 and 25" and Henry's law constants for acetone are given through the range 20-60".
Introduction
Experimental Section and Results
There are few published reports of the relative basicities of ketones involved in hydrogen-bonding reactions.293 Iiecently4 .we reported the results of an investigation of the hydration of acetone in 1,2-dichloroethane using partition, solubility, nmr, and dielectric constant techniques.6 We concluded that the acetone monomer monohydrate is the major associated species in the ternary system acetone-l,2-dichloroethane-water at concentrations of acetone less than 0.55 mole/l. It was determined that the basic strength of the acetone carbonyl group is comparable to that of oxygen in the hydroxyl group of water and aliphatic alcohols. We have now completed a similar study of three additional ternary systems of the type ketone-water-CCL The ketones employed were acetone,. 2,3-butanedione, and acetylacetone. Since the nmr and dielectric constant techrliques are limited by the low solubility of water in CCL, only the partition and water solubility techniques were used in the present investigation.
Acetone, Z13-butanedione, acetylacetone, and CC14 were purified by fractional distillation; the middle portion of distillate was collected. Hexadecane was purified by vacuum distillation. The partition, sohbility, and isopiestic techniques used are similar to methods described ~ r e v i o u s l y . ~Concentrations ~~ of ketones in the 0 3 4 and aqueous phases were deter-
The Journal
of
Physical Chemistry
~-
(1) Abstracted in part from the Ph.D. Dissertation of T. F. Lin. University of Oklahoma, Norman, Okla., 1966. (2) J. M.Wisdom, R. J. Philippe, and M. E. Hobbs, J . Am. Chem. Soc., 79, 1383 (1957). (3) H.Fritzche and H. Dunken, Acta Chim. A d . Sci. Hung., '10, 37 (1964). (4) T.F. Lin, S. D. Christian, and H. E. Affsprung, J . Phya. Chem., 69, 2980 (1965). (5) T.F. Lin, 5.D. Christian, and H. E. Affsprung, submitted. (6) S. D. Christian, H. E. Affsprung, and J. R. Johnson, J . C h m . SOC.,1896 (1963). (7) s. D. Christian, E. Affspwg, J. R. Johnson, and J. Worley, J . Chem. Educ., 4 0 , 419 (1962).
