0. D. BONNERAND DAVIDC. LUNNEY
1140
361°K9 and k d / k , = 3.55 for tritium atom-isopropyl radical reactions at 63"K, an activation energy difference of E, - E d = 0.43 kcal/mole was calculated. This calculation was done by making the assumption that the k d / k c ratio is dependent exclusively upon the temperature and is not affected by a change in the phase. A small activation energy difference of 0.42 kcal/mole with combination having the higher activation energy would also explain the difference in k d / k c values for isopropyl radical disproportionation and combination a t 373 and 77"K, 0 . W 3 and 5.5, respectively. The value of k d / k c of 0.8 for ethyl radicals
a t 63°K" is consistent with the gas phase value of 0.13 at 273"K14 for an activation energy difference of E, Ed = 0.3 kcal/mole which is essentially the same dependence that was first reported in the gas phase.14 Other small activation energy differences of the same magnitude with E, - E d have been reported for isopropyl-isopropyl and sec-butyl-sec-butyl radical disproportionation and combination.8 (13) C . A. Heller and A. S. Gordon, J. Phys. Chem., 60, 1315 (1956). (14) P.S. Dixon, A. P. Stefani, and M. Szwarc, J . Am. Chem. Soc., 85, 2551 (1963).
A Study of Some Concentration Cells with Liquid Ion-Exchanger Membranes
by 0. D. Bonner and David C. Lunneyl Department of Chemistry, University of South Carolina, Columbia, South Carolina
(Received October 11, 1965)
Concentration potentials of cells in which two aqueous electrolyte solutions are separated by a liquid ion exchanger (an oil phase with ion-exchange properties) have been measured for aqueous NaCI, NH,Cl, and CaClz solutions, using organic solutions of dinonylnaphthalenesulfonates as the liquid exchangers, for aqueous HCl solutions using General Mills Aliquat 336 and Humko Kemamine Q-19024 exchangers, and for p-toluenesulfonic acid solutions, using Humko Kemamine Q-1902-C exchanger. The behavior of the cells is analogous to that of concentration cells with solid ion-exchange membranes; their failure to reach the theoretical potential for cells with ideally permselective membranes is explained by considering the solubility of the exchanger in the aqueous phase and transport of water across the organic phase by hydrated ions. Activity coefficients were calculated from data obtained for p-toluenesulfonic acid, and they agree well with activity coefficients determined isopiestically.
Introduction The electrochemical properties of cells of the type
have been studied intensively since about 1935, but very little attention has been given to cells in which the membrane is a liquid ion exchanger (an organic solution of a water-insoluble ionogen). I n 1953 and The Journal of Physical Chemiatry
1954 Bonhoeff er, Kahlweit, and Strehlowz reported potentials of concentration cells in which two aqueous LiCl solutions were separated by an oil phase consisting of quinine hydrochloride in quinoline. The observed potentials fell far short of those obtainable with conventional ion-exchange membranes. I n 1964, Sollner and Shaen3 reported that they had obtained concentra(1) National Science Foundation Graduate Fellow.
(2) K. F. Bonhoeffer, M. Kahlweit, and H. Strehlow, Z . Elektrochem., 57, 614 (1953); Z.Physik. Chem. (Frankfurt), 1, 21 (1954).
CONCENTRATIOX CELLSWITH LIQUIDION-EXCHANGER MEMBRANES
tion potentials within several tenths a millivolt of the theoretical values for cells using benzene, xylene, and nitrobenzene as solvents for Rohm and Haas Amberlite LA-2 anion exchanger in the chloride form. The work described here was initially undertaken for the purpose of developing a cell suitable for the measurement of activity coefficients of high molecular weight disulfonic acids; conventional ion-exchange membranes were unsuitable for this application because of their for example, have small pore sizes. Sollner, et described membranes of graded porosity prepared by adsorbing polyelectrolytes on collodion films. Their least porous membranes have excellent permselectivity but will hold back molecules larger than urea, while membranes with sufficiently large pores for our application have poor permselectivity. Liquid ion exchangers, however, can admit very bulky ions, and their permselectivity is very high for electrolytes which are much more yoluble in the aqueous phase than in the organic phase. ~
1
.
~
~
8
~
Experimental Section Liquid Cation-Exchanger Materials. Sodium dinonylnaphthalenesulfonate (NaDNNS) was prepared by adding sodium hydroxide in ethanol to the sulfonic acid in heptane as received from King Organic Chemical Co. until the solution was strongly basic. The heptane-sulfonate solution was separated by adding a large quantity of water. The heptane layer was washed four or five times with water, and the heptane was evaporated. The product was then weighed and dissolved in o-dichlorobenzene to make a stock solution. Ammonium dinonylnaphthalenesulfonate (NH4DNNS) was similarly prepared from the sulfonic acid and alcoholic ammonia solution. Calcium dinonylnaphthalenesulfonate (Ca(DNNS)Z) was prepared by equilibrating a heptane solution of the acid with several portions of 1 il/l CaClz solution; the heptane layer way shaken with a saturated solution of Ca(OH), to ensure complete conversion. The heptane layer was then washed and evaporated by the same method used for the sodium and ammonium sulfonates. In all instances, the washing effectively removed all water-soluble materials from the organic phase. Liquid Anion-Exchanger dlaterials. Liquid anion exchangers were prepared from General Mills Aliquat 336 and from Keinamine Q-1902-C, manufactured by the Humko Chemical Division of National Dairy Products Co. Aliquat 336 was furnished as a mixture of crystals and syrupy solution in toluene and 2-propanol. This was diluted with toluene, washed five times with distilled water to remove 2-propanol, and evaporated
1141
under vacuum with a rotary evaporator in a hot water bath. The resulting syrup was diluted with o-dichlorobenzene to make a stock solution. Humko Kemamine Q-1902-C is a white, waxy solid consisting of dialkyldimethylammonium chlorides in which the alkyl groups contain from 20 to 22 carbons. It is almost insoluble in heptane, nitrobenzene, odichlorobenzene, and toluene at room temperature, but is very soluble in chloroform. To prepare the chloride exchanger the solid was dissolved in chloroform, washed twice with about 1 M HC1 to ensure complete conversion to the chloride, then washed four to five times with water to remove excess HC1. The resulting stock solution contained about 10% Kemamine chloride. Kemamine Q-1902-C was converted to the ptoluenesulfonate by dissolving it in chloroform and shaking the solution with portions of 1 M aqueous ptoluenesulfonic acid until the sulfonic acid picked up no detectable chloride. The chloroform layer was then washed with distilled water and used as a stock solution. Solvents for Exchanger Solutions. The primary purity criterion for solvents used in this work was low conductivity and freedom from water-soluble impurities. Commercial p-dichlorobenzene was purified by passing it through a column of activated alumina and storing it over alumina; this reduced its conductivity to less than 10-9 ohm-' cm-'. Nitrobenzene was purified by distilling it, storing it over activated alumina, and washing it four or five times with distilled water immediately before use, or by washing undistilled nitrobenzene four or five times with saturated aqueous sodium bicarbonate,6 four or five times with water, and passing it through a column of activated alumina. These methods produced nitrobenzene with a conductivity between 1.3 and 2 X lo-$ ohm-' cm-I (wet). Commercial chloroform was used without purification ohm-' since its conductivity was only 2.5 X cm-l. Exchanger solutions were prepared by diluting the stock solutions to the desired concentrations, washing four or five times with distilled water, and separating the emulsified organic layer. Used exchangers were recovered by repeating the washing and separation procedure. (3) K. Sollner and G. 1M. Shaen, J. Am. Chem. SOC..86, 1901 (1964). (4) C. W.Karr, R. McClintock, and K. Sollner, J . Electrochem. Soc., 109, 251 (1962). ( 5 ) M. H. Gottlieb, R. Neihof, and K. Sollner, J . Phys. Chem., 61, 154 (1957). ( 6 ) J. B. Ezell, unpublished doctoral dissertation, TSniversity of
South Carolina.
Volume 70,Number 4
A p r i l 1966
1142
Aqueous Solutions. The most concentrated solution of each chloride was prepared by weight from the carefully dried salt, then used as a stock solution for preparation of the more dilute solutions. All of the chlorides were Baker and Adamson reagent grade and were subjected to no purification except drying. Acid solutions were prepared by weight from stock solutions whose concentrations had been determined by titrating three weighed aliquots with freshly standardized NaOH solution. The most concentrated solution of each acid was then used as a stock solution to prepare the more dilute solutions, except for lo-* m solutions, which were made from the 0.01 m solutions. Baker and Adamson CP hydrochloric acid was used without purification. Matheson Coleman and Bell p-toluenesulfonic acid was found to be sulfate free and was also used without purification. Electrodes. Silver-silver chloride electrodes were made by a method similar to that of Bonner and Smith.' Platinum wires sealed in soft-glass tubing were washed with concentrated nitric acid and rinsed well with distilled water, then plated for 6-12 hr in 0.2 iM SaAg(CS).?solution at a current density of about 10 ma/cm*. The plated wires were washed with dilute ammonia and soaked in distilled water, then anodized in 1 1l.l HC1 at about 5 ma/cm2 for 30-45 min. The completed electrodes were stored in distilled water or dilute HC1 and were usually reliable for 6 weeks or more. Two matched Beckman glass electrodes with mercury internals were used for the HC1 and p-toluenesulfonic acid measurements. They were calibrated against a silver-silver chloride electrode in HC1 solutions from loW3to 1.0 m; the corrections obtained were applied to the data reported herein for HC1 and p-toluenesulfonic acid. The correction for lo-* m HC1 was obtained by extrapolating the calibration curve. Electrical Measurements. Cell potentials were measured with a Beckman Research pH meter, which is reproducible to =tO.O5 niv and has an absolute accuracy of k0.37 mv over its l500-mv span. The cell was contained in a shielded box maintained at 25.0" by an air thermistor and control amplifier. In initial experiments. attempts were made to use static cells, but these exhibited extremely poor performance because of the surface-active nature of the materials. The cell described below in which all surfaces were continually renewed was necessary to give reproducible results. The Cell. The cell used in this work is shown in Figure 1. The liquid ion exchanger flowed from 1.5mm capillaries C into the aqueous solutions contained in two polyethylene beakers B. The fine glass frit The Journal of Physical Chemistry
0. D. BONNERAND DAVIDC. LUNNEY
Figure 1. Flowing cell for liquid ion exchangers.
F across the lower compartment allowed the drop rates to be independently adjusted with the two needle valves N, while maintaining a low-resistance electrical path through the cell. The reservoirs were ordinary pear-shaped separatory funnels attached by the ball joints J. The beakers rested on a laboratory jack which allowed them to be lowered to clear the capillaries; the electrodes E were held in clamps attached to the jack platform. When both beakers contained identical solutions, the cell developed an asymmetry potential (about 1 mv), but it was usually reproducible. The effect of the asymmetry was canceled by repeating each emf reading several times and interchanging the two beakers after each reading until two pairs of readings were reproduced within *0.2 mv. When glass electrodes were used, this procedure also canceled the emf arising from their differing asymmetry potentials (about 5 mv). It was found that the following conditions must be met to obtain a sharp interface (no emulsion) and thus proper operation of the cell. (1) The capillaries must be rendered hydrophobic to prevent water from wetting their inner walls. They were treated with Siliclad silicone and coated with Dow-Corning silicone grease dissolved in methylene chloride. If the inner walls of the capillaries were (7) 0.D. Bonner and L. L. Smith, J . Phys. Chem., 64, 261 (1960).
CONCENTRATION CELLSWITH LIQUIDION-EXCHANGER MEMBRANES
allowed to become wet with water, a large unreproducible asymmetry potential developed, and the cell emf became too noisy to be readable. (2) The capillary tips must be ground a t an angle as shown in Figure 1 and rendered hydrophobic to encourage the formation of well defined drops. This was not critical for the anion exchangers but was indispensable for cation exchangers; failure to form clean drops of cation exchanger produced noisy, unreproducible potentials. The cell worked well with drop rates ranging from the 2 drops/sec to 1 drop/lO sec. The emf usually oscillated as drops separated from the capillary tips, but the amplitude of the oscillation was below 0.1 mv for cation exchangers when the cell was operating properly. Anion exchangers were in general noisier because the drops did not separate cleanly from the capillaries but wet the external conical tips; this caused a sharp pulse of several tenths of a millivolt as each drop fell, after which the emf would return t o its former value. The input of the Beckman Research pH meter is shunted by a 0.001-pf capacitor which charged during the pulses and discharged slowly through the high resistance of the cell, so a very slow drop rate (about 1 drop/lO sec) was used to ensure that the emf could return to its equilibrium value between drops. Anion exchangers were also slower to reach constant potentials than cation exchangers. With cation exchangers the pot entia1 usually reached its equilibrium value within about 5 min, while anion exchangers required as long as 30 min to reach a constant potential.
Results and Discussion Concentration potentials obtained with the cell in Figure 1 for NaC1, NH4CI, and CaClz solutions are given in Figure 2. In every case silver-silver chloride electrodes were used, and the cation exchanger was the corresponding dinonylnaphthalenesulfonate in 50% nitrobenzene-jO$& o-dichlorobenzene by volume; the exchanger concentration was 0.02 M . Activity coefficients for KC1 tabulated by Pitzer and Brewer* were used for NH4Cl up to 0.1 m because the activity coefficients of KC1 and NH&l are almost identical up to 2.0 171.9 Activity coefficients for 0.003 and 0.03 m CaClzwere obtained by interpolation. Data obtained for HC1 with glass electrodes and anion exchangers are shown in Figure 3. The theoretical potentials shown were calculated from activity coefficients tabulated by Robinson and Stokes.’O Solvents used for the anion exchangers were o-dichlorobenzene for Aliquat 336 and 50% chloroform-50~o nitrobenzene by volume for Kemamine Q-1902-C.
1143
100
98 A
.-+
2a
96
7.-
OCaCI,
U
e
94
8
ANH~CI
0, I
oNaCI
;j 92
z
881
I ’ 0
Figure 2.
I
I
-1 Log mi.
-2
Concentration potentials obtained with
dinonylnaphthalenesulfonate cation exchangers.
100
98
4
i
96
s
a 0
0 ALIQUAT 336 0
*alC
;
0-1902-C
90
EL.10 mi
PI
88
.
I,
86
0
-1
Log
-2
-
mp.
Figure 3. Concentration potentials obtained for HCl solutions with anion exchangers.
These solvents were chosen because of their relatively high dielectric constants and low solubility in water (8) G. N. Lewis and M. Randall, “Thermodynamics,” revised by K. S. Pitzer and L. Brewer, 2nd ed, McGraw-Hill Book Co., Ino., New York, N. Y., 1961. (9) B. F. Wishaw and R. H. Stokes, Trans. Faraday SOC.,49, 27 (1953). (10) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Academic Press Inc., New York, N. Y., 1959.
Volume 70,Number 4 April 1966
1144
0. D. BONNER AND DAVID C. LUNNEY
except in the case of chloroform, a choice necessitated by the virtual insolubility of Kemamine Q-1902-C in our usual solvents. In discussing our results we shall use the quasithermodynamical approach of Scatchard. l 1 For any cell with a transference across a membrane he writes
EF - _ G0,F_ _ Eo,F_ -
RT
RT
RT
-
Cui,
In ai,
Cui, In ai, -
r t i d In ai (1) J a
where the subscripts a and w refer to the left and right compartments, respectively, and ai is the activity of species i; vi, and vi, are the number of moles of species i produced at the left and right electrodes and t i is the number of moles of species i transported from left to right when 1 faraday of negative charge flows reversibly from left to right in the external circuit. G o , and So, are the standard electrode potentials of the left and right electrodes, & is the cell potential, F is the faraday, and R and T are the gas constant and the temperature, respectively. It can be shown that for a cationexchanger membrane with electrodes reversible with respect to the anion equation (1) is equivalent to
where m, and m, are the molalities of the electrolyte MAX, in the left and right compartments; ZM and zx are the ionic charges of the cation and anion; nM is the hydration number of the cation in the organic phase and the y's and 4's are mean activity coefficients and osmotic coefficients, respectively. The first term in eq 2 is the theoretical emf for an ideally permselective membrane, the second term is a correction for water transport, and the third term is a correction for leakage of coion. We would expect the error due to water transport t o be large at high concentrations and to go in the order Ca2+ > Na+ > NH4+ = C1-; examination of Figures 2 and 3 shows that this is indeed the case. The error due to leakage is probably negligible in the data reported here. Sollner and Shaen2 have reported that the ionic selectivities of the liquid anion exchanger Amberlite LA-2 range from 100,000 to 400,000 for aqueous solutions of alkali metal salts; this indicates that the effect of the integral in eq 2 should be negligibly small compared to other errors, even if our exchanger is a great deal less selective than Sollner's. For electrolytes which are appreciably soluble in the organic phase, however, the The Journal of Physical Chemistry
liquid exchanger membrane would leak, and the integral would then be significant. The sources of error thus far discussed are common to both solid and liquid ion-exchanger membranes. There are two additional errors present in liquid exchangers: the effects of solubility of the exchanger and those of the organic solvent in the aqueous phases. The effect of the organic solvent on the activities of aqueous electrolytes is not predictable from any simple theoretical treatment; the error can be minimized, however, by using solvents with very low solubility in water. The effect of exchanger solubility can be treated more rigorously. Let us consider the cell
A g l A ~M~X. (\mi ~R)organic ( m , * MX(mw) ~ s ~ ~ ~ AgX,Ag ~~R~m,*) We assume that both aqueous solutions are in equilibrium with the organic phase, so that the chemical potential of the organic salt MR is the same in all three phases. Ignoring transport of water and Xions yields EF
= -(px,
- px,) - JwtRd&R a - smUtMdpM
-
- PX,)
-(IZX,
- (PM,
- PM,) -
Since the chemical potential of MR is the same in all three phases, the integral is zero. We now assume that the presence of a low concentration of MR in the aqueous solutions does not change the activity coefficients of the ions M+ and x-. If MR is regarded as being completely ionized in aqueous solutions, we obtain UM,
= ydm,
+ ma*)
UM,
= yM,(m,
+ mu*)
where y, and y, are mean activity coefficients of MX. The cell potential now becomes
The first term is the theoretical potential, and the (11) G.Scatchard, J. Am. Chem. SOC.,7 5 , 2853 (1953).
CONCENTRATION CELLSWITH LIQUIDION-EXCHANGER MEMBRANES
second term is a correction for solubility of the exchanger. We would expect this correction to increase rapidly in dilute solutions. This is seen in the HC1 data obtained with Aliquat 336 (Figure 3). For 10-3/10-2 m HC1 the error is only 0.1%; for 10-4/10-3 m HC1 it increases to more than 12%. Furthermore, the observed potential for 10-4/10-3 m HC1 decreased with time as the aqueous solutions became saturated with the exchanger. The relatively high solubility of Aliquat 336 in water is also indicated by the fact that the conductivity of the o-dichlorobenzene-exchanger solution decreased by about 25% when it was washed four times with water. Studies of the Kemamine anion exchanger and HC1 were quite promising, however, and encouraged an attempt to measure activity coefficients for a sulfonic acid. p-Toluenesulfonic acid was chosen for the first attempt because its activity coefficients have been accurately determined isopiestically down to 0.1 m.12 The results obtained with Kemamine Q-1902-C are shown in Table I. The exchanger solvent was 50% chloroform-50% nitrobenzene by volume. ~~~~~
~
Table I : Concentration Potentials for p-Toluenesulfonic Acid Obtained with Kemamine Q-1902-C in 50% C e , H ~ N 0 ~ - 5CHCI," 0~~ mdm
Obsd emf, mv
0.0001/o ,001 0.001/0.01 0.003/0.03 0.01/0.10 0.03/0.30 O . O O l / O 003 0.003/0.01 0.01/0.03 0.03/0 10 0.10/0 30
126.91 115.56 112.73 110.00 106.33 56.06 59.26 53.38 56.69 49.82
Exchanger conductivity, 9.3 X loT5ohm-' cm-l.
The activity coefficients reported in Table I1 were calculated from the data in Table I by choosing 0.003 m as a reference concentration, calculating the potentials for all other solutions against the reference (using every possible sum of experimental emf values and averaging the results), and plotting the functions E - .-
118.3
m log 7= log m
y
- log y'
where m' and y' are the molality and mean activity coefficient for the 0.003 m reference solution, respectively.
1145
Extrapolation of this function to m = 0 gives the value of y'. The extrapolation was taken along the DebyeHuckel limiting slope, which is equivalent to assuming that p-toluenesulfonic acid obeys the limiting law up to 0.003 m. This procedure is justified by the fact
Table I1 : Activity Coefficients of p-Toluenesulfonic Acid -Log
-Emf
-Isopiestic
data--
m
Y
Y
d
Y
0.003 0.01 0.03 0.10 0.30
0.0279 0.0495 0.0754 0.1195 0.1483
0.938 0.892 0.841 0.759 0.667
... ... ...
... ...
a
0.923" 0.894"
... 0.759 0.660
data1Z4
... ... ... 0.922 0.887
Calculated by graphical integration of 1 - ( l / m )f m d l n y
-+=
that the curve between 0.003 and 0.01 m deviates only minutely from the limiting slope. Points for 0.001 and 0.0001 m solutions were ignored in making the extrapolation because their inclusion would have led to grotesque deviations from the limiting law. The high potentials observed for solutions below 0.003 nz are probably attributable to the effect of the organic solvent of the exchanger on the activity of p-toluenesulfonic acid in the aqueous phases since all the other sources of error lead to potentials lower than the theoretical. This is predicated on the assumption that we have written the cell reaction correctly, i.e., that current is carried only by the p-toluenesulfonate ion. If an appreciable fraction of current were carried by a dimer with a single negative charge, the observed emf would be higher than the value calculated from our assumed cell reaction. Attempts to apply the method described here to measurement of activity coefficients of disulfonic acids may not be feasible because of inability to specify a unique cell reaction, i e . , if the acid HS03RS03H dissociates incompletely, so that some current is carried by the ion HS03RS03-. In summary, we believe that we have described a method of some potential usefulness in measuring activity coefficients of bulky electrolytes for which cells with solid ion-exchange membranes would be unsuitable. At the present stage, however, great care must be exercised in the choice of the exchanger and solvent, and data are reliable only over the approximate range 3 x lo-'> m > 3 X (12) 0. D. Bonner, G. D. Easterling, D. L. West, and V. F. Holland, J . Am. Chem. Soc., 7 7 , 242 (1955).
Volume 70, Number 4 April 1955
RYOHEINAKANEAND TOSHIYUKI OYAMA
1146
Acknowledgments. The authors gratefully acknowledge the assistance of Charles F. Jumper and William J. Sutton, who did the preliminary work on this prob-
lem. The project was partially supported by the Atomic Energy Commission under Contract No. AT(40-1)-1437.
Boron Isotope Exchange between Boron Fluoride and Its Alkyl Halide Complexes, 11.1 Infrared Spectrum of Boron FluorideMethyl Fluoride Complex
by Ryohei Nakane and Toshiyuki Oyama The Institute of Physical and Chemical Research, Bunkyo-ku, Tokyo, Japan
(Received October 15, 1965’)
The infrared spectra of liquid B10F3,B1*F3,CH3F, B10F3.CH3F complex, and Bl1F3.CH3F complex are observed in the region from 400 to 4000 cm-l. The isotopic data are used to calculate the theoretical equilibrium constant for boron isotope exchange between gaseous boron fluoride and its methyl fluoride complex. The calculated values are found to agree fairly well with observed values.
Previously, one of the authors found that, for boron isotope exchange, the equilibrium constant is much smaller between gaseous boron fluoride and weak boron fluoride complexes in the liquid form than between gaseous boron fluoride and strong boron fluoride complexes in the liquid form. Namely, with boron fluoride-alkyl halide complexes1 or boron fluoride-alkyl halide-alkylbenzene 1: 1: 1 addition oriented 7~ complexes,2 which exist only a t low temperatures, the constant is much smaller than with boron fluoride-ether complexes which are stable even at room temperature. The known infrared spectra of gaseous boron fluoridea and liquid boron fluoride-ether complexes4 were used to calculate the theoretical equilibrium constant for boron isotope exchange. I n the present work, the infrared spectrum of liquid boron fluoride-methyl fluoride complex was observed a t low temperatures and the equilibrium constants, observed and calculated from isotopic data of infrared spectrum, were compared. The Journal of Physical Chemistry
Experimental Section For the measurements of infrared spectra at low temperatures a cell as shown in Figure 1 was made. The specimen to be studied was introduced into the space between two KRS-5 plates (1 cm diameter by 5 mm thick) A and A’, between which spacer B of Teflon was inserted to give a space of about 0.05 mm thickness. The t,wo KRS-5 plates were firmly held in brass holder C with two Teflon gaskets D and D‘ for vacuum tightness. The cryostat was a stainless steel dewar with two KRS-6 plate windows (2 cm diameter by 3 mm thick), E and E’. The thickness, (1) Part I : R. Nakane, 0. Kurihara, and -4. Natsubori, J . Phys. Chem., 68, 2876 (1964). (2) R. Nakane, A. Natsubori, and 0. Kurihara, J . Am. Chem. Soc., 87, 3597 (1965). (3) J. Vanderryn, J . Chem. Phys., 30, 331 (1959). (4) A. A. Palko, G. M. Begun, and L. Landau, ibid., 37, 552 (1962); G. M. Begun and A. A. Palko, ibid., 38, 2112 (1963); G.M. Begun, W. H. Fletcher, and A. A. Palko, Spectrochim. Acta, IS, 655 (1962).
