ACIDOPENTAAMMINECOBALT(III) COMPLEXES
1641
Osmotic and Activity Coefficients of Acidopentaamminecobalt(111) Complexes
by L. H. Berkal and W. L. Masterton Department of C h i s t r y , University of Connecticut, Storrs, Connecticut (Received December 21, 1965)
Osmotic coefficients, measured a t 37", and activity coefficients, derived therefrom, are reported for the compounds [Co(NH3)sA]X2 where A- = acetate, propionate, or isobutyrate and X- = NOS-, I-, Br-, or C1-. Below 0.1 m,the osmotic and activity coefficients of all of these compounds fall within a narrow range of each other. At higher concentrations, 4 and Y~ appear to be independent of the organic ligand but vary with the anion in the order C1- > Br- > I- > NO3-. The osmotic and activity coefficients of these complex ion electrolytes are markedly lower than those of simple electrolytes of the same valence type. The data are interpreted in terms of association between cation and anion; calculated ion-pair dissociation constants fall between 0.05 and 0.1. The factors which are responsible for association of complex ion electrolytes are discussed. It is suggested that dispersion forces may play an important part in ion-pair formation.
Introduction Osmotic and activity coefficients have been reported previously for the 2 : 1 complex ion electrolytes [Co(NH3)5NO2lC12, [ C O ( ~ H ~ ) & ~ ] ( Cand ~ O ~ [Co(NH3)5F]C12.2 It was found that the activity coefficients of these compounds are lower than those of simple 2 :1 electrolytes and are more nearly comparable to those of simple 1 :2 electrolytes. These results were interpreted qualitatively in terms of the relatively minor role of cation hydration in 1:2 and complex ion electrolytes and extensive association in the latter compounds as indicated by the low Debye-Hiickel u0 values required to fit the data a t low concentrations. As part of a continuing study of the thermodynamic properties of aqueous solutions of complex ion electrolytes, we have chosen a series of acidopentaamminecobalt(II1) complexes, which may be represented by the general formula [ C O ( N H ~ ) ~ A ]where X ~ A- = CH,COO-, CH3CH2COO-, and (CH3)2CHCOO- and X- = NO3-, I-, Br-, and C1-. These compounds offer a systematic approach to studying the effects of ligand and anion substitution on activity coefficients in both dilute and concentrated solutions. Several of these salts have aqueous solubilities exceeding 1 m; in no case does the rate of aquation of the complex cation exceed l%/day3 at 37'. Experimental Section Cmpounds. (1) [ C O ( N H ~ ) ~ C H ~ C OXeries. O ] X ~ (a)
X - = NO-. This was prepared by the method of Basolo, et aL3 Anal. Calcd: NH3, 26.0 Found: NH3,25.9. (b) X - = I-. This was prepared according to the method of Werner.4 Anal. Calcd: NH3, 18.6; I, 55.5. Found: NH3,18.6; I, 55.4. (c) X - = Br-. [ C O ( N H ~ ) ~ C O ~ ] B Pwas - Htreated ~O~ with acetic acid; the desired product was precipitated by adding a concentrated solution of NaBr and was recrystallized from water using NaBr. Anal. Calcd: NH3, 23.5; Br, 44.0. Found: NHa, 23.4; Br, 44.3. (d) X- = CZ-. A solution of the bromide salt was passed through a column of Amberlite IRA-400 resin in the chloride form. The product was brought out of solution by freeze-drying and was recrystallized from hot water. Anal. Calcd: NH3, 31.1; Cl, 25.9. Found for sample dried over P205: NH3, 30.8; C1, 25.4. (2) [ C O ( N H ~ ) ~ C H ~ C H ~ CSeries. O O ] X ~( a ) X - = Nos-. This was prepared by the method of Basolo, et aL3 Anal. Calcd: N H , 25.0. Found: NH,, 25.0. (1) Abstracted in part from the Ph.D. thesis of L. H. Berka. (2) W. L. Masterton and J. A. Scola, J . P h y s . Chem., 68, 14 (1964). (3) Estimated from data given by F. Basolo, J. G. Bergmann, and R. G. Pearson, ibid., 56, 22 (1952). (4) A. Werner, Ber., 40, 4111 (1907). (5) A. Werner and N . Goslings, ibid., 36, 2380 (1903).
Volume 70,N u m b e r 6 M a y 1966
L. H. BERKAAND W. L. MASTERTON
1642
(b) X- = I-. A solution of K I was added to a solution of the nitrate salt. The product was recrystallized from hot water. Anal. Calcd: "3, 18.1; I,53.9. Found: NH3, 18.1; 1,533. ( c ) X - = Br-. A solution of the iodide salt was shaken with solid AgBr. After filtration, the product was brought out of solution by freeze-drying. Anal. Calcd: WH:, 22.6; Br, 42.4. Found: "3, 22.6; Br, 42.2. ( d ) X - = C1-. This was prepared by shaking a solution of the iodide salt with solid AgC1. The product was brought out of solution by freeze-drying. Anal. Calcd: NH,, 29.6; C1, 24.6. Found: NH3, 29.6; C1, 24.6. (3) [CO(A'H~)~(CH~)~CHCOO]X~ Series. The method given by Basolo3 for the preparation of the nitrate salt gave a product which analyzed correctly for NH3 but which was found to be contaminated by a relatively insoluble compound. Recrystallization gave a very poor yield owing to the high solubility (>2.5 m) of the nitrate. For these reasons, the procedure was modified first to prepare the iodide salt, which was then readily converted to the nitrate. ( a ) X - .= I - . Ten grams of [Co(NH3)&03]N03.0.5Hz0 was suspended in 17 ml of water and 37 ml of isobutyric acid was slowly added. The mixture was concentrated on the steam bath for 90 min, then cooled in ice. The solid product was removed by filtration and tQe iodide salt was precipitated from the filtrate using KI. The product was recrystallized twice from hot water. Anal. Calcd for [Co(NH3)6(CHJ&HCOO]Iz*0.5HzO: "3, 17.2; I, 51.4; C, 9.7; H,4.7. Found: "3, 17.2; I, 49.4; C, 9.9; H, 4.8. (b) X- = NO3-. Prepared by titrating a solution of the iodide to the equivalence point with AgN03. After filtration, the product was brought out of solution by freeze-drying Anal. Calcd: NH3, 24.0; C, 13.5; H, 6.2. Found: NH3, 24.0; C, 14.1; H, 6.0. (c) X - = Rr-. This preparation was analogous to that of [CO(X~H~)~CH~CHZCOO]B~~. Anal. Calcd for [Co(NH3)j(CH3)2CHC00]Br2~H~O: "3, 20.8; Br, 39.1; C, 11.7; H, 5.9. Found: NH,, 20.7; Br, 38.0; C, 11.9; H, 5.9. (d) X- = Cl-. This preparation was analogous to O ] C ~ ~Calcd . that of [ C C ~ ( N H ~ ) ~ C H ~ C H ~ C OAnal. O O ] C ~: ~NH3, 27.4; for [ C O ( N H ~ ) ~ ( C H ~ ) ~ C H C*0.5H20 C1, 22.7; C, 15.4; H, 7.5. Found: NH3, 27.3; C1, 22.0; C, 15.8; H, 7.3. Apparatus. Osmotic coefficients were measured a t 37" with a Nechrolab vapor pressure osmometer. The method described previously2 was modified as follows. Potassium chloride solutions rather than pure water The Journal of Physical Chemistry
were used for both the reference bead and the solvent cup. A resistance reading, AR, was taken on a solution of complex ion electrolyte nearly isopiestic with the KCl solution used as a reference. The osmotic coefficient, $, of the complex ion electrolyte is then given by $ =
( ~ ~ $ ) K c I AR 3m + k g
(1)
The constant k in eq 1 was calculated from resistance readings taken with two KCl solutions of different concentrations. Instead of using a AR reading to calculate $ directly, as was done previously, the AR reading, by this method, is used only to obtain a small correction to be added to the accurately known first term on the right of eq 1. Consequently, the error in $ introduced by an error in k is minimized. For each compound, measurements were taken a t approximately 0.01 m intervals in the range from 0.010.10 m. With the more soluble compounds, additional readings were taken a t approximately 0.1 m intervals up to 1.0 m. Typical data are given in Table I. Treatment of Data and Results
Smoothed Osmotic and Activity Coeficients. These quantities, obtained as described previously,2 are given in Tables II-IV. The experimental error is estimated to be f1%for $ and =k3% for -yh in the concentration range 0.03-0.10 m, and *0.5% for $ and f 3 % for 7 % above 0.1 m. The values of a0 obtained were virtually identical for all of the compounds studied, averaging 2.5 f 0.2A. . Extent of Ion-Pair Formation. The fact that the osmotic and activity coefficients of these electrolytes are significantly lower than those of "normal" 2: 1 electrolytes implies a considerable degree of association between cation and anion. This is confirmed by the small magnitude of the Debye-Huckel a0 values required to fit the data a t low concentrations. I n these respects, the data reported here resemble those for other complex ion electrolytes studied previously.2 For the process MA+ PI2+ A- the ion-pair dissociation constant, K , is given by the expression
+
K = where
+
41 ff)m(7*'2d3 (1 - 4 ( Y * ' 1 J 2
(2)
CY is the degree of dissociation of the ion pair, and and 1 ~ ~ ' 1 are 1 the mean activity coefficients of completely dissociated 2 : 1 and 1: 1 electrolytes, respectively. Values of K for the 12 compounds reported here were calculated by the approach described in a previous article,$making use of the equation ~
~
'
2
ACIDOPENTAAMMINECOBALT(III) COMPLEXES
1643
~
Table I : Measured Osmotic Coefficients of [Co(NH3)&H&H&OO]C12 m
9
m
9
m
9
m
9
0.0103 0.0104 0.0202 0.0302 0.0400 0.0503
0.863 0.878 0.844 0.826 0.817 0.805
0.0600 0.0706 0.0802 0.100 0.198 0.295
0.793 0.788 0.769 0.766 0.739 0.715
0.390 0.396 0.485 0.570 0.653 0.732
0.701 0.700 0.691 0.687 0.681 0.677
0.852 0.999 1.50 1.81 2.10 2.40
0.672 0.670 0.669 0.671 0.676 0.681
Table I1 : Smoothed Osmotic and Activity Coefficients of [Co(NH3)&H3COO]XZ -----No8---
--
--I--
m
4
Yi
9
Yi
0.03 0.04 0.05 0.07 0.10 0.20 0.30 0.40 0.50 0.60
0.829 0,812 0.800
0.563 0.526 0.498
0.833 0.817 0.805 0.788 0.776
0.570 0.533 0.505 0.464 0.424
p a - -
Br--
9
Yi
9
Y t
0.833 0.816 0.803 0.783 0.763 0.722 0.697 0.678 0.662 0.650
0.568 0.531 0.502 0.459 0.415 0.333 0.289 0.259 0.237 0.220
0.820 0.805 0.794 0.777 0.760 0.725 0.699 0.678 0.664 0.661
0.552 0.515 0.487 0.446 0.404 0.326 0.283 0.253 0.232 0.217
Table 111: Smoothed Osmotic and Activity Coefficients of [Co(NH3)&H3CH2COO]X2 -------NOa---
-
1-
-c1---
-Br------
m
4
7.42
9
Yi
9
Yf
9
Yi
0.03 0.04 0.05 0.07 0.10 0.20 0.30 0.40 0.50 0.60 0.80 1.0 1.2 1.6 2.0 2.4
0.829 0.812 0.798 0.777 0.754 0.703 0.667 0.641
0.563 0.525 0.496 0.452 0.407 0.320 0.272 0.240
0.829 0.813 0.800 0.780 0.760 0.719 0.692 0.667 0.642
0.565 0.528 0.499 0.456 0.412 0.330 0.285 0.254 0.229
0.829 0.813 0,800 0.781 0.761 0.724 0.701 0.681 0.664 0.648 0.625 0.618 0.609
0.565 0.528 0.499 0.456 0.412 0.332 0.289 0.259 0.237 0.219 0.193 0.176 0.163
0.830 0.814 0.802 0.784 0.766 0.735 0.718 0.704 0.692 0.682 0.672 0.670 0.670 0.671 0.674 0.681
0.565 0.529 0.500 0.458 0.415 0.339 0.298 0.270 0.250 0.234 0.210 0.196 0.186 0.168 0.156 0.148
.
b 2--1
=
adJ’21 .~
+
2/30
- a>4‘11
(3)
where Cpzl is the measured osmotic coefficient while 4/21 and 4’11 are calculated osmotic coefficients for completely dissociated 2 : 1 and 1 : 1 e1ectrolYt% respectively. The values of 4/21, 4/11)Y+’ZI, y+’u used in eq 2 and 3 to solve for K are given in Table V. These parameters are slightly smaller than those used in a previous papery6
which was concerned with the calculation of ion-pair dissociation constants of “simple” 2 :1 electrolytes. Presumably, the osmotic coefficient of a completely dissociated complex ion electrolyte should be somewhat smaller than that of a simDle electrolvte such as c a c l , because of hydration effects. The entry of solvent (6) w. L. Maaterton and L. H. Berks, J . Phus. chm.,in press.
Volume 70, Number 6 May 1986
L. H. BERKA AND W. L. MASTERTON
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Table IV : Smoothed Osmotic and Activity Coefficients of [GO(NHa)5(CHa)zCHCOO]Xz m
1-
Yi
9
Y*
9
7i
9
Y*
0.823 0.805 0.791 0.769 0.745 0.697 0.664 0.636 0.611 0.588 0.552 0.535 0.484 0.450 0.427
0.557 0.519 0.489 0.444 0.398 0.313 0.266 0.234 0.210 0.190 0.162 0.144 0.113 0.094 0.081
0.830 0.814 0.801 0.781 0.759 0.715 0.683 0.657 0.633 0.614 0.593
0.565 0.528 0.499 0.456 0.411 0.328 0.282 0.249 0.225 0.206 0.180
0.822 0.806 0.793 0.774 0.754 0.716 0.691 0.671 0.654 0.640 0.630
0.555 0.517 0.488 0.445 0.401 0.321 0.278 0.249 0.227 0.210 0.187
0.827 0.811 0.799 0.781 0.763 0.732 0.715 0.701 0.690 0.680 0.669 0.671 0.671 0.671 0,672
0.560 0.523 0.495 0.453 0.410 0.334 0.293 0.266 0.246 0.230 0.207 0.194 0.169 0.152 0.141
0.03 0.04 0.05 0.07 0.10 0.20 0.30 0.40 0.50 0.60 0.80 1.0 1.5 2.0 2.5
7
P
water molecules into the first coordination sphere of an ion such as Ca2+will have the effect of increasing the apparent values of t$ and T* for CaC12. The magnitude of this correction was estimated from the equal ions of Stokes and Robinson.' I n the concentration range covered here (m I O.l), the effect is very small; the maximum correction in 421, for example, amounts to only -0.009. Table V : Osmotic and Activity Coefficients Calculated for Completely Dissociated Complex Ion Electrolytes Ionic strength
0.010 0.015 0.020 0.030 0.040 0 050 0.060 0.080 0.100 0.120 0.150 0.180 0.210 0.240 0.270 0.300 I
r-
-Nos---------. 9
9'21
0.938 0.928 0.920 0.909 0.901 0.894 0 888 0.882 0.877 0.872 0.867 0.864 0.862 0.861 0.860 0.861 1
+'I1
0.969 0.964 0.959 0.952 0.948 0.944 0.942 0.938 0.934 0.931 0.928 0.925 0.924 0.924 0.923 0.923
Y*'"
0.818 0.789 0.765 0.732 0.706 0.686 0.669 0.643 0.622 0.605 0.585 0.570 0.557 0.546 0.537 0.529
Yf'"
0.903 0.886 0.873 0.851 0.836 0.822 0.812 0.794 0.780 0.768 0.753 0.741 0.732 0.725 0.717 0.710
Equations 2 and 3 were used to calculate K for each complex ion electrolyte at a series of molalities ranging from 0.03 to 0.10. The excellent agreement between calculated K values at different molalities (Table VI) is particularly striking. The Journal of Phyaieal Chemistry
B
------Cl--
Table VI: Calculated Values of the Dissociation Constant of [GO(NHa)sCH3CHzCOO]2+, C1m
991
a
K
0.03 0.04 0.05 0.06 0.07 0.08 0.09 0.10
0.830 0.814 0.802 0.792 0.784 0.777 0.771 0.766
0.793 0.752 0.722 0.695 0.668 0.648 0,630 0.608
0.088 0.085 0.084 0.083 0.081 0.082 0.081 0.080
The ion-pair dissociation constants given in Table
VI1 for the 12 electrolytes studied were calculated a t m = 0.1, except €or [Co(NHa)&H3COO](N03)2, where, because of solubility limitations, the calculation was made a t m = 0.05. Values of K were also calculated by the more involved thermodynamic approach described in ref 6 with virtually identical results.
Table VII: Dissociation Constants for [Co(NH&,A]2+, X -
CHaCOOCHaCHzCOO(CHa)zCHC00-
Nos-
1-
Br -
c1-
0.080 0.064 0.054
0.097 0.071 0.071
0.075 0.072 0.064
0.071 0.080 0.075
Discussion Up to a molality of 0.1, the osmotic and activity coefficients of the compounds listed in Tables 11-IV are (7) R. H. Stokes and R. A. Robinson, (1948).
J. Am. Chem. Soc., 70, 1870
ACIDOPENTAAMMINECOBALT(III) COMPLEXES
1645
very close to one another, except for an indication of The only simple 2 : l salts which resemble those lower values for the nitrates. This is, of course, also studied here in the variation of 4 and y h with contrue of the ion-pair dissociation constants listed in centration are ones which are extensively ion paired, Table VII. notably the alkaline earth nitrates. Even with these Above 0.1 m, I$ and yi. appear to change very little compounds, the extent of ion pairing is much smaller. when one acid ligand is replaced by another. The For example, the ion-pair dissociation constant of agreement between the osmotic coefficients of the Ca(NO& is estimated by Daviese to be about 0.5; those for the compounds listed in Table VI1 fall between compounds [Co(NH3)~CHaCH2COO]Clzand [Co0.05 and0.1. (NH3)6(CH3)2CHCOO]Clnover the entire concentration range from 0.03 to 2.00 m is particularly striking. A The question arises as to why these complex ion similar agreement is observed with [Co(NH3)&H3- electrolytes ion pair to a much greater extent than COO]Brz and [ C O ( N H ~ ) ~ C H ~ C H ~ Cfrom O O ]0.03 B ~ ~to simple salts of the same valence type. The answer 0.60 m. The equivalence of propionate and isobutyrate cannot lie in the coulombic factor frequently invoked to ligands is indicated in the first instance and that of explain relative extents of ion pairing. Complex acetate and propionate ligands, in the second. cations of the type [ C O ( N H ~ ) ~ Amust ] ~ + be a t least as On the other hand, a t concentrations above 0.1 m, large as hydrated alkaline earth cat,ions (e.g., Cathere is an appreciable change in I$ or y i when one (H20),j2+,et c.) . anion is replaced by another. Almost without exPerhaps the simplest explanation of the effects ception, the osmotic coefficients decrease in the order observed here is that offered by Stengle and Langford’O C1- > Br- > I--> Nos-. At very high concentrations, to rationalize nmr studies which indicate that salts of the spread becomes particularly large. Compare, for Cr(en)g3+ are extensively associated while those of example, the two compounds [ C O ( N H ~ ) ~ ( C H ~ )Cr(H20)E3+ ~are not. They suggest that hydrated CHCOOIClz and [CO(NH~>~(CH~)~CHCOO](NO~)~, cations can fit readily into the surrounding network of whose osmotic coefficients at 2.0 m are, respectively, water molecules and consequently a,re unlikely to come into close contact with anions. Complex cations such 0.671 and 0.450. If one interprets the data at concentrations above as Cr(en)a3+or, in this case, [ C O ( N H ~ ) ~ Ashould ~ + , not 0.1 m solely in terms of varying extents of association bond as readily with water molecules in the second between anion and cation, two conclusions can be coordination sphere and hence may be able to associate drawn. more readily with anions. Carrying this idea one step further, it might be (1) Association is little affected by a change in ligand involving the replacement of one organic anion postulated that large ions such as [CO(NH&A]~+, by another. which cannot fit readily into the surrounding water Association appears to increase in the order structure, may create a cavity in the solvent large (2) C1- < Br- < I- < NOa-. This order is the reverse of enough to accommodate simultaneously an anion. This what one would expect if the association were strictly type of “structure-enforced” ion pairing has been coulombic, in which case, its extent should be inversely proposed by Diamond“ to explain the abnormally low related to anionic size. osmotic and activity coefficientsof tetraalkylammonium The most striking feature of these data is the fact iodides. Quite possibly, it may be characteristic of that the osmotic and activity coefficients are considerwater solutions of complex ion electrolytes as well. ably lower than those of most simple 2: 1 electrolytes Both of these approaches treat ion pairing in terms of the effect of a large cation upon the surrounding (e.g., CaClZ, etc.). This becomes obvious at concentrawater structure. We should like to suggest an altertions as low as 0.1 m where T~ is about 0.41 f 0.02 as native explanation which emphasizes instead the compared to values of 0.50 or greater commonly observed with simple 2 : 1 salts.* Furthermore, these relatively strong noncoulombic forces which must complex ion electrolytes show no evidence of a minimum exist between ions of the type dealt with here. It was reported earlier12 that the neutral complex [Co(NH3)ain activity coefficient at concentrations as high as 2.5 m; T & values for simple 2 : l salts ordinarily pass through a minimum somewhere between 0.1 and 1.0 m and then increase rapidly with increasing concentration. (8) R. H. Stokes, Trans. Faraday SOC.,44, 295 (1948). (9) C.W. Davies, “Ion Association,” Butterworth Inc., Washington, The activity coefficient of CaCl2, for example, after D . C., 1962. passing through a minimum of 0.448 a t about 0.5 m, (10) Thomas R. Sten& and C. H. Langford, J . Phys. C h m . , 69, rises to 1.063 a t 2.5 m; y i for [CO(NH~)~(CH&- 3299 (1965). CHCOOlCl2 at 2.5 m is only 0.141. (11) R. M. Diamond, ibtd., 67, 2513 (1963). Volume 70,Number 6 May 1966
RICHARD L. HANSEN
1646
(NO&] is strongly salted-in by alkali halides. The salt effects in these systems were explained largely in terms of dispersion forces between the complex molecule on the one hand and the ions of the salt on the other. Extending this idea to complex ion electrolytes, one might expect dispersion forces to make a significant contribution to the formation of ion pairs involving a large, highly polarizable cation such as [CO(NH&A]~+. This would explain, among other things, why the extent of association appears to increase in the order C1- < Br- < I- since this is the order of increasing polarizability of the anions. I n this connection, Rosseinsky13has pointed out that, contrary to popular belief, ion-pair formation in salts of
the alkali or alkaline earth metals usually increases with the size of the cation. This lends support to our belief that, while coulombic forces and solvent structure effects are undoubtedly important in ion pairing, dispersion forces may play a much more significant role than has been generally recognized.
Acknowledgment. This work was supported in part by the Research Foundation of the University of Connecticut and by an NSF graduate fellowship awarded to L. H. B. ~
(12) W. L. Masterton and Robert N. Schwartz, J . Chem. Phys., 69, 1546 (1965).
(13) D. R. Rosseinsky, J. Chem. SOC.,785 (1962).
Nitro-p-terphenyls. I. Dual Charge-Transfer Properties and Spectral Correlations
by Richard L. Hansen Contribution No. $60 jrom the Central Research Laboratories, Minnesota Mining and Manujacturing Company, St. Paul, Minnesota 66119 (Receiued November 10, 1966)
Seven nitro-p-terphenyls have been found to form charge-transfer complexes with both electron donors and acceptors. The results of quantitative st,udies of complexes with tetracyanoethylene and with N,N-dimethyl-p-toluidine are reported. A partial interpretation of the absorption spectra of the nitroterphenyls has been made on the basis of their charge-transfer spectra and elementary molecular orbital theory.
Introduction Organic charge-transfer complexes have been extensively investigated for a number of years. Typically, a given organic molecule acts as either an electron donor or an electron acceptor but seldom as both. In addition to intramolecular charge-transfer interactions in certain molecules, a limited number of compounds such as iodine can play a dual role and form self-complexes in which one molecule behaves as an electron donor toward a second molecule as acceptor.’ Examples of materials possessing more general dual The Journal o j Physical Chemistry
charge-transfer properties are quite rare. The “chargetransfer’’ or “a-complex” theories developed by Mulliken,2 by Dewar and L e ~ l e y and , ~ enunciated by Briegleb4 intimate that organic molecules should be capable of functioning as both charge-t,ransfer donors and ac(1) P. A. D.deMaine, J . Chem. Phys., 24, 1091 (1956). (2) R. S. Mulliken, J . A m . Chem. Soc., 72, 600 (1950); 74, 811 (1952); J . Phys. Chem., 56, 801 (1952). (3) M. J. S. Dewar and A. R. Lepley, J. A m . Chem. SOC.,83, 4560 (1961). (4) G. Briegleb, Angew. Chem., 76, 326 (1964).
