3056
NOTES
Table I : Average Values of the First-Order Rate Constant ka
O C
Initial pressure, torr
No. of data points
seo-1
676 676 650 650 650 599 599 599 552 552
204 102 204 102 51 204 102 51 408 204
7 9 11 6 6 6 7 10 10 9
141 149 73.5 71.2 73.2 20.5 18.7 18.5 3.73 2.67
T,
106ka,
204 torr and various temperatures is shown in Figure 3. Since pure perfluoroisobutene was not available for calibration of the gas chromatograph, the quantitative results for this compound are based on the assumption that the thermal conductivity detector had the same sensitivity to perfluoroisobutene as the average of the perfluorobutene-2 and octafluorocyclobutane sensitivities. The rate of production of perfluoroisobutene with respect to perfluoropropene also had an order between 1 and 2. I n addition to the fluorocarbons mentioned above, the gas-phase reaction products contained carbon monoxide, carbon dioxide, and silicon tetrafluoride. Butler' also noticed side reactions with the wall when he studied the pyrolysis of octafluorocyclobutane in a Pyrex vessel in the temperature range 360460". In all cases, these three compounds were significant products. A mass balance on the carbon, including all of the gaseous carbon containing compounds, a t 650" indicated that a small carbon mass loss of 10-25% occurred. This mass loss was attributed to the small flakes of white dust that condensed in the cooler parts of the system. The present experimental results indicate that the pyrolysis of perfluoropropene, in the temperature and initial pressure ranges 550-675' and 50-410 torr, respectively, can be represented by a first-order reaction. When the pyrolysis of perfiuoropropene was carried out in a nickel vessel in the same temperature range, Atkinson and A t k i n ~ o nreported ~~ that their data could be represented by a reaction of order 1.5. In their paper Atkinson and Atkinson4bdo not specify the range of initial CaFe concentrations which were considered. However, these authors reported that the perfluoropropene half-life was approximately 60 min when the reaction temperature and initial CaFs pressure were approximately 600' and 400 torr, respectively. As discussed earlier, the half-life of the peduoropropene pyrolysis a t 599" determined in the present investigation varied between 58 and 70 min when the initial perfluoropropene pressure was varied from approximately 50 to 200 torr. A reaction of order 1.5 The Journal of Physical Chemistry
requires that the half-life be approximately doubled when the initial reactant concentration is decreased by a factor of 4. Considering these comments and the fact that the half-life as measured by Atkinson and Atkinson4b is approximately the same as determined in this investigation, it is reasonable to suggest that the pyrolysis of perfluoropropene is a first-order reaction.
Acknowledgment. The author acknowledges the help of Mr. W. C. Kelly and M r . M.J. Siemion in the experimental and data-analysis phases of the program. This research was sponsored by the Air Force Office of Scientific Research, Office of Aerospace Research, United States Air Force, under Grant No. AF-AFOSR1144-67.
Nuclear Magnetic Resonance Study of Micelle Formation in Sodium Perfluorocaprylate and -propionate by Rizwanul Haque Oregon State University, Corvallis, Oregon 97831 (Received March 26, 1968)
The various phenomena occurring in or due to surfactant materials as studied by nuclear magnetic resonance (nmr) technique include phase transition,l behavior in liquid crystalline ~ t a t e , membra ~ ' ~ ne^,^ interaction in s o l u t i ~ n ,and ~ solubilization s t u d i e ~ . ~ ~ ~ The evidence of micelle formation in-soap solutions has been obtained from the concentration dependence of proton chemical shift*and spin relaxation time of waterg and solute protonlOJ1 resonance. Changes in the proton chemical shift of quaternary ammonium soaps are small at different concentrations.12 The shielding of the fluorine nucleus is more susceptible to the environment than the shielding of the proton, and hence (1) T. J. Flautt and K. D. Lawson, Advances in Chemistry Series, No. 63,American Chemical Society, Washington, D. C., 1967, p 26. (2) A. Saupe, B.Englart, and A. Povh, ref 1,p 51. (3) M. P. McDonald, Arch. Sei. (Geneva), 12,141 (1959). (4) D. Chapman, V. B. Kanat, J. DeGier, and 5.A. Penkett, Nature, 213, 74 (1967). (5) R. M. Roseberg, H. L. Crespi and J. J. Kats, Abstracts, 155th
National Meeting of the American Chemical Society, San Francisco, Calif., April 1968. (6) J. C. Eriksson, Acta Chem. Scand., 17, 1478 (1963). (7) T.Nakagawa and K. Tori, Kolloid Z.,194 (1964). (8) J. Clifford and B. A. Pethica, Trans. Faraday Soc., 60, 1483 (1964). (9) J. CWord and B. A. Pethica, ibid., 61, 182 (1965). (10) J. Clifford, ibid., 61, 1276 (1965). (11) K. D. Lawson and T. J. Flautt, J . Phys. Chem., 69, 3204 (1905). (12) H.Inoue and T. Nakagawa, ibid., 70,1108 (1966).
3057
NOTES
0
1
1
1
I
I
*
CAPRYLATE PROPIONATE
*
I
I
I
CONCENTRATION OF SOAP (MOLAL)-I Figure 1. Plot of 19F chemical shift against the inverse concentration of soap: (a) -CF2 peak; (b) -CFa peak. figure, the ordinate represents the absolute chemical shift (to avoid the confusion of sign of chemical shift of -CFa and -CF2 peak).
salts containing fluorine show stronger concentration dependence of l9F chemical shift.l3-l6 Thus fluorinated colloidal electrolytes show greater changes in the l9F chemical shift during the micelle formation. Recently, Muller and Birkhahn17p18 demonstrated the concentration dependence of the l9F chemical shift of partially fluorinated salts of some long-chain carboxylic acid. In this note we report the '9F chemical shift of two completely fluorinated salts of relatively short-chain carboxylic acids, namely sodium perfluoropropionate and -caprylate. The critical mi-
I n this
celle concentration (cmc) of these two less studied soaps has been calculated from the results.
(13) J. N.Shoolery and B. J. Alder, J. Chem. Phys., 23,806 (1966). (14) R. Haque and L. W. Reeves, J. Phys. Chem., 70,2763 (1966). (16) R. Haque and L. W. Reeves, Can. J . Chem., 44, 2769 (1966). (16) R. Haque and L. W. Reeves, J. Amer. Chem. SOC.,89, 260 (1967). (17) N.Muller and R. H. Birkhahn, J. Phys. Chem., 71, 967 (1967). (18) N.Muller and R. H. Birkhahn, ibid., 72,683 (1968). Volume 7& Number 8 August 1968
NOTES
3058 Experimental Section Sodium salts of perfluorocaprylate (contains 8 carbon atoms, C,) and -propionate (contains 3 carbon atoms, C3) were supplied by K and K Chemical Co. Aqueous solutions of each surfactant ranging from 0 t o 3 m were prepared in double-distilled water. The nmr spectra were recorded on a Varian HA 100 nmr spectrometer operating a t 94.1 MHz.lg Chemical shifts were measured with a side-band technique using trifluoroacetic acid as an external standard and have an accuracy of *0.1 Hz. No bulk susceptibility corrections were made, as they were small." All the experiments were performed at the probe temperature of -30". The 19Fnmr spectra of propionate soap gave peaks due to -CF3 and -CF2 groups while the caprylate soap gave peaks due to CF3, a-CF2,and the remaining -CF, groups.
Table I : Fluorine Chemical Shift in Perfluoropropionate and -caprylate as a Function of Soap Concentration Propionate Chemical shift," ppm -CFa -CFz
Concn, m
2.93 2.38 1.515 1.215 0.570 0.195
-0.356 -0.262 -0.159 -0.0956 -0.085 -0.069
Caprylate Chemical shift,a ppm -CFs -CFa
Concn, m
0.76 0.66 0.605 0.515 0.546 0.467
1.48 1.365 0.940 0.690 0.479 0.366 0.1065
-0.436 -0.415 -0.234 -0.160 -0.138 -0.117 -0.085
0.790 0.780 0.635 0.566 0.54 0.535 0.446
a To obtained chemical shift values from trifluoroacetic acid, a value of -0.425 and $41.5 ppm should be added to the chemical shift's values of -CFa and -CF2 peaks, respectively.
Results and Discussion The fluorine chemical shifts of -CFI and a-CFz resonance at different concentrations of salts are given in Table I. The equilibrium states occurring in the surfactant solutions are represented as
S+ + A-
nS+
+ nA-
SA
(1)
z(SA), KE
A,-"
(2)
+ nSf
(3)
Ka
A,-"
z(SA),
Perfluoropropionate Perfluorocaprylate u-Fluorocaparate a-Fluorolaurate
(4)
S+ and A- represent the cation and the anion, respectively; (SA), the ion pair; (SA),, the polymeric species; A,-", the micelle; and m and n are constants. K I , K 2 ,K E , and K 3 are appropriate equilibrium constants. It is assumed that the concentration of the ion pair is very small as compared to that of the micelle, and hence only the KE is important in these solutions. Following the assumptions made by Rhller and Birkhahn,l7 the observed chemical shift 6o is expressed in terms of , the chemical cmc, chemical shift of the micelle 6 ~ and shift of the free anion 6,. (5)
where A is the concentration of the surfactant. A plot of (A)-I against So gives two straight lines intersecting at the crnc (Figure 1). The value of cmc obtained in this way is given in Table 11. An examination of Figure 1 indicates that a t infinite dilution [(l/A) = a ] , both the surfactants have nearly the same chemical shift; however, the extrapolated value of chemical shift a t [(l/A) = 01 (which is 6M) for the two surfactants is not identical. Muller and Birkhahn17 report 6 M and 6, to be independent of the number of carbon atoms. Our results indicate The Journal of Physical Chemistry
(6M
Soap
K1
Kz
4SA)
Table I1 : Critical Micelle Concentration of Various Fluorinated Soaps
u-Fluorotridecaneate
No. of carbons
-
of -CFa, ppm
Cmc
Ref
3
-1.33 M
0.535
This work
8
-0.69 M
0.63
This work
10
-0,167 M
1.28
12
-0.051 M
1.28
13
-0.024
M
1.28
Muller and Birkhahn Muller and B irk h ah n Muller and Birkhahn
a The value of (&,I - 8,) has been obtained by subtracting 8, (chemical shift a t (A-1) = a) from 8~ (chemical shift a t (A-l) = 0).
Sa to be independent of the number of carbon atoms while 6~ is slightly dependent upon the number of carbon atoms. The independent nature of 6, on the number of carbon atoms is reasonable because the ions are completely hydrated at infinite dilution. The magnitude of the change in chemical shift appears smaller for caprylate and propionate surfactants when compared t o clO-c13 soaps. The difference in chemical shift magnitude may result from different values of n for the various surfactants. If this is the case, ( 6 ~ 6,) should give a measure of n or degree of micellization. The (6M - 6,) values are 1.28 ppm for CIOCI3 soaps, 0.53 ppm for CB,and 0.63 ppm for Cg soap (Table 11). Estimated cmc values for the n m r technique are summarized in Table 11. It appears from (19) Preliminary measurements were made during the summer of 1966 at the University of British Columbia, Vancouver, Canada.
3059
NOTES this study that the cmc values of surfactants obtained by other methods20321 are considerably lower than values obtained by nmr technique. (20) H. B.Klevins and M. Raison, J . Chim. Phys., 51,l (1954). (21) C. H. Arrington and G. D. Paterson, J . Phys. Chem., 57, 247 (1953).
Contact Charge-Transfer Spectra of Iodine in Some Hydrocarbon Solvents
by Larry M. Julien and Willis B. Person’ Department of Chemistry, University of Iowa, Iowa city, Iowa 62240, and Department of Chemistru, University of Florida, Gainewille, Florida 88601 (Received April 8,1068)
The phenomenon of contact charge-transfer absorption on the long wavelength side of the very strong 182-mp band of Iz solutions in hydrocarbons has been known for some time.2-6 It is illustrated in Figure 1 for I2 in some typical paraffin solvents. There is a more or less normal broadening and solvent shift of the strong Iz band to the red, together with some new additional absorption near 220-250 mp. The latter has been attributed to the “contact” charge transfer abn-heptane), s ~ r p t i o n , ~between -~ contact pairs (Iz differing from an ordinary charge transfer band’ only in that the molecular pair is not stable in the ground state, so that AHfO = 0 and Kr = 0. We wish to report here a, few studies which confirm the earlier work2-6 and extend it to 180 mp using differential double-beam techniques. The spectra were obtained with a Beckman DK-2A far-ultraviolet spectrometer, flushed continuously with dry N2 to remove atmospheric 0 2 absorption. For solution studies we used short path-length cells (UV-0-2) with far-ultraviolet silica windows (from Limit Research Corp.). Path lengths were measured by interference fringes. For the vapor phase studies, we dsed a 1-cm Beckman cell with excess 1 2 crystals. The concentration of 1 2 was computed from the known vapor pressure a t the temperature of the cell.* I n all experiments we used the Beckman temperature-regulated cell holder, circulating coolant from a Haake circulator and Sargent water bath cooler. The hydrocarbon solvents (Phillips chemically pure grade) were further purified by passing through silica gel column^.^ The iodine was resublimed from a Mallinckrodt Analytical Reagent grade sample. Stock solutions were prepared by weighing Izinto volumetric flasks and filling with the appropriate solvent. These solutions were then used shortly after preparation; they were diluted to reach the appropriate concentrations for study.
+
Figure 1. The absorption spectrum near 200 mp of 1 2 vapor (-), in n-heptane ), and in cyclohexane (- - -). The dotted line ( , , ) gives the spectrum estimated for the Iz “182-mp” band. Absorptivities are based on the total IZ concentration in the solutions. (--.--e
. .
The spectra were obtained in double-beam experiments with pure solvent (or mixed solvent) in the reference beam. For example, solutions of n-heptane and cyclohexane were prepared and a base line was measured with this solution in each cell. Then 1 2 was added to the sample cell and the absorbance was measured again. The difference between this spectrum and the spectrum of a solution of the same concentration of 1 2 in pure n-heptane was then analyzed to obtain the contact absorption due to the cyclohexane-12 pair. In other experiments a solution of 1 2 in pure n-heptane was used instead in the reference beam and the differential absorbance (ranging from -0.30 to +0.70) was recorded directly. Both methods gave the same results. Using the vapor spectrum as a guide, we estimated the absorbance in solutions (such as in Figure 1) due to the strong 182-mp Izband. This estimate is shown with a dotted line in Figure 1. Subtracting this absorbance from the total absorbances in the solution gave the resulting contact bands shown in Figure 2. These bands are of about the same appearance as ordinary charge-transfer bands. For example, the half(1) T o whom reprintrequests should be directed at the University of Florida. (2) J. 8. Ham, J. R. Platt, and H. McConnell, J. Chem. Phys., 19, 1301 (1951). (3) D.F. Evans, ibid., 23, 1424,1426(1955). (4) D.F. Evans, J.Chem. SOC.,4229 (1967). (5) R. 8. Mulliken, REC.Trav. Chim.,75,845 (1956). (6) L. E. Orgel and R. 8. Mulliken, J. Amer. Chem. SOC.,79, 4839 (1967). (7) R. 8.Mulliken, ibid., 74,811 (1952). (8) National Research Council, “International Critical Tables of Numerical Data,” Vol. 111, McGraw-Hill Book Co., Inc., New York, N. Y.,p 201. (9) W.J. Potts, J.Chem. Phys., 20,809 (1952). Volume 72, Number 8 Auguat 1068
