fighting mixtures. There are many fluorocarbon surfaceactive agents able to lower the water surface tension sufficiently to permit water solutions or water-rich foams to spread spontaneously on volatile, low surface tension organic liquids such as toluene, nitromethane, and n-octane. This study has shown that thin layers of water when spread in more or less uniform films will be effective barriers to the evaporation of organic liquids. One reason for this is the low solubilities of these organic liquids in water. One would expect the thin layer of water trapped in the lamellae of foams to be barriers to the evaporation of volatile liquids. Foams may also increase the thickness of the diffusion boundary layer above the liquid surface. literature Cited
Banks, W. H., iVature 174, 365 (1954). Banks, W. H., “Proceedings of Second International Congress of Surface Activity,” Vol. 1, p. 16, Academic Press, New York, 1957
Bailom, W. D., Halper, L. A., Jarvis, N. L., Znd. Eng. Chem. Prod. Rcs. Develop. 8, 118 (1969). Bernett, hl. K., Jarvis, N. L., Zisman, W. A., J . Phys. Chem. 66, 328 (1962). Bernett, hl. K., Zisman, W. A., J . Phys. Chem. 67, 1534 (1963). Bernett, AI. K., Zisman, W. A., J . Phys. Chem. 70, 1064 (1966b). Bernett, 11.K., Zisman, W. A., “Solution Systems for the Displacing of Organic Liquids from Solid Surfaces,” NRL Rept. 6402 (JIay 25, 1966a). Blake, C. B., Ahlbrecht, A. H., Bryce, H. G., “Perfluoroalkyl
Surface-Active Agents for Hydrocarbon Systems,” Division of Petroleum Chemistry, 126th Meeting, ACS, New York, N. Y., 1954. Bowers, R. C., Jarvis, N. L., Zisman, W. A., Znd. Eng. Chem. Prod. Res. Develop. 4, 86 (1965). Brace, N. O., J . Org. Chem. 21, 1342 (1956). Davies, J. T., Rideal, E., “Interfacial Phenomena,” pp. 5, 303, Academic Press, New York, 1961. Ellison, A. H., Zisman, W. A., J . Phys. Chem. 60, 416 (1956). Frenkiel, J., “Evaporation Reduction,” Ariz. Zone Research Report XXVII, UNESCO, Paris, France, 1965. Gilby, A. R., Heymann, E., Australian J . Sci. Res. (Ser. A) 1, 197 (1948). Harkins, W. D., Jordan, H. F., J . Am. Chem. SOC.52, 1751 (1930). Jarvis, N. L., J . Colloid and Interface Sci., 29, No. 4 (April 1969). Jarvis, N. L., Zisman, W. A., J . Phys. Chem. 64, 150 (1960). Kanner, B., Reid, W. G., Petersen, I. H., Znd. Ens. Chem. Prod. Res. Develop. 6, 88 (1967). LaRfer, V. K., “Retardation of Evaporation by RIonolayers,” Academic Press, New York, 1962. Langmuir, I., Langmuir, D. B., J . Phys. Chem. 31, 1719 (1927). Lewis, W. K., Whitman, W. G., Znd. Eng. Chem. 16, 1215 (1924). Mac Iiitchie, F., Science, 163,929 (1969). hlurphy, C. hl., Zisman, W. A., Znd. Eng. Chem. 42,2415 (1950). Schwarz, E. G., Reid, W. G., Znd. Eng. Chem. 56, 26 (1964). Shafrin, E. G., Zisman, W. A., Advan. Chem. Ser., No. 43, 145 (1964). Shafrin, E. G., Zisman, W. A., J . Phys. Chem. 66, 740 (1962). Tuve, R. L., Jablonski, E. J., U. S. Patent 3,258,423 (June 28, 1966). Walker, D. C., Rev. Sci. Instr. 34, 1006 (1963). RECEIVED for review May 20, 1969 ACCEPTEDNovember 5, 1969
Solvent Effects on the Volume of a Transition State J. R. McCabe,’ R. A. Grieger, and C. A. Eckert Department of Chemistry and Chemical Engineering, University of Illinois, Urbana, Ill. 61801
The partial molal volumes in solution of the transition state of the Menschutkin reaction between pyridine and methyl iodide have been determined. Accurate partial molal volumes of both reactants have been measured and used in conjunction with existing high-pressure kinetic data to find the volumes of the complex. Such results are useful in the interpretation of the structure of the transition state and its intermolecular interactions in solution.
P m T I . u , hfoL.kL volume measurements of reactants in solution can be used to make high-pressure kinetic studies an even more effective tool for the investigation of the nature of the transition state of a chemical reaction in solution. Many authors have reported volumes of activation, the volume difference between reactants and transition state, and used these data to interpret mechanisms or to predict solvent effects. However, more accurate and reliable estimates can be made on the basis of the volume of the transition state itself. Although this volume cannot be measured directly, it is relatively simple to measure the partial molal volumes of the reactants and subtract them from the activation volumes. Hopefully this approach will in turn lead to a better understanding of the intermolecular interactions of the transition state in solution and help in the rational design of solvents for reactions. The activated complex theory proposed by Eyring (1935) 1
156
Present address, Chevron Research Corp., Richmond, Calif. l&EC FUNDAMENTALS
VOL. 9 NO. 1 FEBRUARY 1970
was first applied to high-pressure kinetics by Evans and Polanyi (1935). Assuming an equilibrium between the reactants and achieved complex, they showed that the pressure dependence of the reaction rate constant is given by Av*
RT where the rate constant, k , must be expressed in pressureindependent units such as mole fraction. Further justification for the assumption of equilibrium has been developed from statistical mechanics by Marcus (1967). The activation volume Av” is the difference in partial molal volumes between transition state and reactants. For a bimolecular reaction of the type, A B + products (2)
+ +
the volume of activation is rigorously AV’
=
Dli
- DA - DE
(3)
Most attempts a t interpretation have centered about Av*, which is a difference of large numbers. Av* has frequently been regarded as the sum of two terms, Avl*, a "structural" term, and Av2*, a 'kolvation" term. AY,* is related t o the intrinsic change in volume of the reacting molecules resulting from alterations of interatomic distances in the formation of the activated complex. Avz* is the volume change due to interactions between the solvent and the activated complex. Buchanan and Hamann (1953) and David and Hamann (1954) have shown that Aut* is the dominant term for some reactions in which a change in charge accompanies the formation of the transition state, due to strong electrostriction effeds. Many authors have correlated solvent effects with the AVZ* term. For example, Brower (1963) found an increase of 16 cc per gram mole in the activation volume of the reaction between n-butyl bromide and pyridine when switching from a nonpolar to a polar solvent. When the A s * term is dominant, rates for reactions of ionic reactants yielding nonionic products decrease with increasing pressure. Stewart and Weale (1965) have confirmed this effect in studying reverse llenschutkin reactions. I n this case, a negative A s " due to formation of bonds in the transition state is counterbalanced by the release of constricted solvent molecules. The electrostatic effects on the volume of polar molecules in solution can be derived from the treatment of Kirkwood (1934). He showed that the work (or Gibbs energy a t constant T and P ) required t o place a nonpolarizable spherical cavity of radius r with a point dipole moment p a t its center in a homogeneous medium of dielectric constant E is given by (4) Since Equation 4 holds for partial molal quantities, the application of the thermodynamic relationship
(E?)
=ij
T ,solvent
yields the relationship between the partial molal volume and the pressure derivative of the dielectric constant factor,
It is not necessary to deal with Av*, a difference of two large numbers. Rather DA and DB may be measured accurately to give values of DM virtually as accurate as the 'Av* values. Heydtmann et al. (1966) used such a method in a study of the llenschutkin reaction between a-picoline(X) and W bromoacetophenone(I3) in nine solvents. They measured the partial molal volumes to 50.7 cc per grain mole in each solvent with a pycnometer and used Equations 5 and 6 to calculate values for P A , P B , and P M . Using a structural sphere model to predict radii, values of PA = 1 & 1 Debye, po = 3.2 i 0.4 Debyes, and ,ur = 7.8 0.5 Debyes were calculated. Typical experimental values for PA and p B are 1.92 Debyes (Freiser and Glowacki, 1949) and 3.14 Debyes (Mohler, 1938), respectively; 7.8 Debyes appear reasonable for the activated complex. The authors conclude that the electric charges in the activated complex are less than half developed. Hartmann et al. (1965a) investigated the influence of six solvents on the pressure acceleration of the llenschutkin reaction between pyridine and methyl iodide,
*
CH3I
+ CjHjN $ CsHjN
* *
. CH3 . * I-+ CjHSNCH,@+ Ie *
where the transition state is also highly polar but nonionic. The volumes of activation reported are: Solvent
Carbon tetrachloride Benzene Acetone Chlorobenzene Methanol Xitrobenzene
Av*,
cc ~
Mole
a t 50'
C, 1 Atm
-37.5 -35.1 -33.2 -30.7 -26.6 -21.3
Data are also available for the pressure dependence of the dielectric constant for these solvents (Hartmann et aZ., 196513). I n this work we report the determination of the partial molal volumes of pyridine and methyl iodide in dilute solution in the six solvents listed above. These results are used t o determine the solvent effect on the partial molal volume of the transition state for the pyridine-methyl iodide reaction. Experimental
(5) where 8, is the partial molal volume in a solvent in which there is no electrostriction. By using Relations 1 and 3, Equation 5 can be written in terms of the activation volume.
Some authors (Laidler and Eyring, 1940) add a term to the right-hand side of Equations 5 and 6 to account for all nonelectrostatic interactions, such as hydrogen bonds and London forces. Corrections for factors such as nonspherical molecules, polarizable molecules, and noncentral dipoles can be made (Boettcher, 1952), but they change the functional dependence only slightly in most cases. A particularly suitable type of reaction to study for comparison with the Kirkwood approach is the Alenschutkin reaction between tertiary amines and organic halides. The rate of reaction increases with pressure and depends strongly on the nature of the solvent. Electrostrictive effects are important in the determination of Av*, since there is a large difference in polarity between the reactants and transition state.
The partial molal volumes were determined by the dilatometric method of Hepler et al. (1965) with some modifications.
A 500-cc Erlenmeyer flask was fitted with a side arm of 2-mm precision bore capillary tubing and a modified ground glass stopper as shown in Figure 1. The solute was weighed into the capsule (stopper) and loaded into the flask containing the solvent. The flask and capsule were carefully sealed with stopcock grease to prevent leaks and the formation of air bubbles. This entire assembly was weighed and immersed in a constant temperature water bath. The bath was controlled t o within *0.002' C with a Hallikainen Thermotrol and a Beckniann thermometer, a t a temperature of 25' =t 0.02' C. The data were taken at 25' rather than 50" C because of the high volatility of methyl iodide. Accurate temperature control was necessary, since a change of 0.002O C will cause a change in volume of approximately 10-3 CC, depending on the solvent. After thermal equilibrium was attained, mixing was accomplished by pulling away the magnet holding the capsule lid with a large horseshoe magnet and stirring with a magnetic stirrer. The volume change on mixing was measured with a Gaertner cathetometer to +0.001 cm, which corresponds to better than 10-4 cc. Duplicate runs of four points each were made in each solvent. The capsule was refilled with pure solute and mixed with the solvent-solute mixture in the flask for each determination. The volume changes were plotted against the solute concenVOL 9 NO. 1 FEBRUARY 1970
I&EC FUNDAMENTALS
157
2 m m PRECISION BORE CAPILLARY TUBING
24/40 I JOINT MODIFIED 24/40 T GROUND GLASS llcc STOPPER CAPACITY
-
TEFLON COATED MAGNETS STAINLESS STEEL
GROUND GLASS SURFACE
TEFLON COATED
FLAT GLASS L I D
AGNETIC STIRRER
Figure 1.
81-
-
Dilatometer for measuring volume changes on mixing
NITROBENZENE
s
CHLOROBENZENE
-
4
BENZENE
80-
cc
MOLE
-
"0
,ACETONE
fi
r
6
4
ACARBON TETRACHLORIDE
"tI---+
I
CARBON TETRACHLOR
METHANOL
78
I 0.5
OO
NITROBENZENE
,r
1
I
1.0
CONCENTRATION,
1.5
62 O0
-
Figure 2. Partial molal volumes of pyridine in various solvents a t 25" C
05
CONCENTRATION,
IO
1.5
*i
Figure 3. Partial molal volumes of methyl iodide in various solvents a t 25" C
tration (expressed as moles of solute per mole of solvent)
Results and Discussion
and the points fitted by a least squares method to the form
The partial molal volumes of methyl iodide and pyridine at 25" C in six solvents are shown in Figures 2 and 3. The data presented are the average of the two best runs; the maximum difference in partial molal volume calculated from duplicate runs was 0.06 cc per gram mole. The values of the partial molal volumes at infinite dilution are listed in Table I. No partial molal volume data for these systems were found in the literature for comparison. However, one may compare the data of Heydtmann et al. (1966) for the partial molal volume of a-picoline (1-methylpyridine) in acetone, chlorobenzene, and nitrobenzene with the results reported here for pyridine in those solvents. As would be expected, the sign and magnitude of the volume changes on mixing are in good agreement. These results may be used in Equation 3 (A = pyridine, B = methyl iodide) to separate the volume of the transition state from Av*. Since the kinetic data were taken a t reactant concentrations of 0.2M, partial molal volumes a t that concentration mere used. The data from this work were
Av = a
+ bn + cnz
(7) Equation 7 was differentiated with respect to the solute concentration and the partial molal volumes were calculated from the expression b Av 02
=
VZ+
(T-)
=
VZ+
b+2m
where v2 is the molar volume of the solute. The accuracy of this method was checked by comparing partial molal volumes obtained for the system KaOH-water with literature values. Hepler et al. (1965) obtained an apparent molal volume of -4.98 cc per gram mole a t a concentration of 0.77:11, which compares with -4.95 cc per gram mole determined using the technique described here. Reagent grade materials were found to be of sufficient purity for the purposes of this work and were used as received from the supplier. Further experimental details are described by hIcCabe (1969). 156
lhEC FUNDAMENTALS
VOL 9 NO. 1 FEBRUARY 1970
Table 1.
I
Experimental Results at Infinite Dilution
5"
V"
Pyridine,
25" C, Cc per Solvent
Gram Mole
Carbon tetrachloride Benzene Acetone Chlorobenzene Xlethanol n'itrobenzene
79 80 79 80 78 81
Methyl Iodide, 25' C, Cc per Gram Mole
45 08 59 91 25 08
62 63 64 62 64 62
72 74 95 84 25 78
Carbon tetrachloride Reiizeiie Acetone Chloiobenzene Methanol Sitrobcnzene
Av
-VA
*
-V B
-"M
120-
cc MOLE
\
A
CHLOROBENZENE
ACETONE
-V 1 1
OO
CARBON
05
15
10
d
I., Weale, K. E., J . Chem. SOC.1965, 2849, 2854. RECEIVED for review February 10, 1969 ACCEPTED July 16, 1969
activated complex solute
literature Cited
Boet tcher, C. J. F., “Theory of Electric Polarisation,” Elsevier, Amsterdam, 1952. Brower, K. R., J . Am. Chem. SOC. 85, 1401 (1963). Uuchanan, J., Hamann, S. D., Trans. Faraday SOC.49, 1425 (1953). David, 1% G., Hamann, S. D., Trans. Faraday SOC.50, 1188 (1954). Eckert, C. A., Ind. Eng. Chem. 59, No. 9, 20 (1967). Evans, RI. G., Polanyi, RI., Trans. Faraday SOC.31, 875 (1935). Eyring, H., J . Chem. Phys. 3, 107 (1935).
,
Spatial Distribution of Electron Density and Electric Field Strength in a High-Frequency Discharge Criteria for Similarity Alexis T. Bell Department of Chemical Engineering, University of California, Berkeley, Calif. ~ 4 ~ 2 0
The development of electric discharge processes of direct industrial application has generated an interest in the relationship between the rates at which activated species are formed and the electrical properties of the discharge. Since the rates of discharge processes are related to the electron density and the electric field strength, there is an interest in their spatial variation. These two distributions are examined for a helium discharge sustained between parallel plates of a radio-frequency discharge. The electric field and density distributions are used to calculate the voltage across the discharge and the volume-averaged power density. By maintaining PA, and ( P ) constant, the rates of excitation in two reactors of different size can b e made identical.
1s THE: last few years a number of process applications have been developed for free radicals created in high-frequency electric t1isch:irges. The majority have dealt with the niodificat,ion of the surface properties of solids and the preparation of thin films. Treatment of polymer films with ionic and metastable species of a noble gas have been used to produce a highly crosslinked surface layer to which adhesive bonding is greatly improved (Hansen and Schonhorn, 1966). Printing on polymers has been made possible by exposure of the surface to a stream containing atomic oxygen (Mantell and Orm:intl, 1964). Sevcral of the newest developments relate t o the semiconduct,or industry. One of the first steps in the manufacture of traasistors or integrat,ed circuit,s is the deposition of a thin laycr of silicon oxide or nitride on the surface of a silicon wafer. This film passivates the surface and acts as a diffusion mask 160
I&EC FUNDAMENTALS
VOL. 9 NO. 1 FEBRUARY 1970
in subsequent processing. A number of investigators have shown that such films can be prepared either by direct oxidation of the silicon surface (Ligenza, 1965) or via the decomposition of a silane with atomic oxygen or nitrogen (Kuwano, 1968; Secrist and Xackenzie, 1966). Atomic oxygen has also been used to remove light-sensitive polymer films (photoresist) used for photoengraving purposes (Irving, 1968). Finally, recent work has shown that uniform pinhole-free polymer films, useful for encapsulation of microcircuits, can be formed by the interaction of excited argon with a variety of gaseous hydrocarbons (Smolinsky and Heiss, 1968). The development of discharge processes of direct industrial importance has generated an interest in exploring the relationship between the rates a t which activated species are formed and the electrical properties of the discharge. One question of immediate interest concerns the spatial variation
