H. G. H A R H l S
802
ASD
D.11.HIMMELBLhU
T'ol. 67
KIKETICS OF THE REACTIONS O F ETHYLENE WITH SULFURIC ACIDREACTION OF ETHYLENE WITH SULFURIC ACID AXD ETHYL HYDROGES SULFATE BY H. G. HARRISASD D. Ill. HIMMELBLAU Department of Chemical Engineering, L'niuersity of Texas, Austin 12, Texas Received September 15, 196.9 A kinetic study has been carried out of the reactions of ethylene in diethyl sulfate-sulfuric acid-ethyl hydrogen sulfate solutions. hlthough by nature a heterogeneous system, the data were obtained from the homogeneous reactions of ethylene in a non-aqueous media. Rate data mere obtained, a kinetic model was proposed, and and rate constants determined. The rate constants proved to be functions only of the initial acid concentration. Attempts to obtain a linear correlation between rate constant gnd some type of acidity function were not successful, indicating that caution is required in the interpretation of rate data in non-aqueous solutions where the variation of activity coefficients is not well understood.
Introduction The best known reactions in the ethylene-sulfuric acid-ethyl hydrogen sulfate-diethyl sulfate system are CzH4
+ H2S04
CZHSSO~H
+ CzHbS04H Jr (CzH5)zS04 ~ C Z H S S O ~ H (Ci")zS04 + CzH,
(1) (2)
(3)
I n dilute sulfuric acid solutions a complex equilibria is established as outlined in Table I. The reactions other than (l), (a),and (3) were curbed by working in essentially non-aqueous solutions of 99.8% HzS04 added to diethyl sulfate (DES). TABLEI OVER-.4LL REACTIONS I N THE ETHYLEKE-DIETHYL SGLFATEETHYL HYDROGEN SCLFATE-SULFURIC ACID-IT'ATERSYSTEM
+ HzS04
(CzH6)zO
ti CzH6OH H2SQ4
+ +
+ CzH4 Jr CZHSHSO4 HzO
ti CZH60H CzH6HSO4
+ 112804 CZH~OH
ti HzO
+ CzHsHS04 + CzH4 I_ (CzHs)zSO4 + &SO4
ti 2CzHSHS04 While the character and mechanism of the reactions of ethylene have been investigated to sonie consider-
able extent in a qualitative way, quantitative data on the rates of reaction and kinetic constants are scarce. There have been several qualitative and semiquantitative studies of reactions 1 and 2. Plant and Sidgwickl carried out early rate measurements for these reactions by bubbling CzH4 through HZ804 and noting the weight increase of the acid solution. Dalin and Gutyrya2 measured the rate of ethylene absorption on revolving paddles wetted with H2S04. The rate of decrease in pressure of ethylene in a rotating cylinder partially filled with HzS04 was determined at various conditions by Davis and Schuler3 and Pigulevskii and Il'ina.4 Milbauer and others5 measured the reaction velocity of ethylene bubbling through H2SO4. Hellin and Jungers6 and Kerdivarenko and others' noted the decrease in the pressure n ith time of an ethylene atmosphere over HzS04 in a vibrating vessel. The niajor problem in all of the foregoing studies was that reactions 1 and 2 are heterogeneous reactions. I n each instance the authors noted that with increased turbulence (mixing) in the liquid phase or increased area available for ethylene dissolution the reaction proceeded more rapidly, indicating that the hydrodynamics of the system was obscuring the true chemical kinetics and that the rate of mass transfer of the gas to the liquid was the controlling factor in the absorption, rather than the chemical reaction rate. Other authorsSj9have attempted to determine the chemical kinetics of the reactions by examining simultaneous diffusion and reaction of ethylene in the stagnant liquid. These attempts yielded little useful information about the chemical kinetics because of the assumptions and approximations that had to be made to solve the problem, e.g., the solubility of ethylene in sulfuric acid was assumed to be the same as the solubility of ethylene in water, ethylene diffusivity in sulfuric acid was calculated from an equation developed for diffusion of colloidal particles, and a simple kinetic model assuming an irreversible reaction was used to (1)
S G. P. Plant and N
(2)
ill. A. Dalin and V. 8. Gutyrya, Khzm. Refeiat. Zh., 1, S o 8-9. 84
17.
Sidgaick, J Soc Chem. Ind
40, 14T 11921).
(1938). (3) H. S. DaTis and R Sohulei, J . Am. Chem. Soc., 52, 721 (1930). (4) V V. Pigulevskii and S I. Il'ina. Doklady Akad. Nauk S S S R , 45, 352 (1944). (5) J. Milbauer, J. Kurka and J Mikolasek, Chem. Obzor, 14, 213 (IQJB). (6) M. Hellin a n d J C. .Tungem, Bull soc chzm. Prance, 380 (1987) (7) >I. 4. Kerdixarenko. P. K. lIixa1, a n d RI. Kh. Xishinebskli, Zh. Priklad Khim., 28, 459 (19%). (8) A. I. Gel'buhtein and a1 I. Temhin, Zh. rzz. Khzm , 31, 2697 (1957) (9) hf. S.Nemtsov, Khzm. Prom, No. 8, 16 (1960).
April, 1963
REBCTIONS OF
ETHYLENE WITH
effect several mathematical simplifications. This lack of pure chemical reaction rate data makes the determination of such data of interest and importance. It became apparent with preliminary trials that the study of a homogeneous reaction would be much preferable to trying t o remove the influence of mass transfer from the over-all reaction rate in a heterogeneous reaction. It also became apparent that reactions 1 and 2 are in some respects similar to reaction 3: at high sulfuric acid concentration the reactions are very rapid, but with a decrease in acid concentration the rate of reaction is decreased significantly. The several investigations of (1) and ( 2 ) that had been attempted had been carried out a t high acid concentration. It proved feasible to collect data on reactions 1 and 2 as homogeneous reactions for mixtures of Hi304 and DES in solutions of relatively high DES concentration.
SULFURIC
Results and Analysis Preliminary investigations showed that the rate of the ethylene reaction was much more rapid than reaction 3; it was found that the ethylene reaction proceeded to completion with virtually no change in the concentrations of (CzH&S04, HzS04,and CzH5S04Hby reaction 3, even though some of the mixtures were far from equilibrium with respect to reaction 3. I n addition, the amount of ethylene dissolved was so small that its reaction to coinpletion did not noticeably af( 10) A. Weissberger, Ed., "Physical Methods of Organic Chemistry," 2nd Ed., Vol. I, Part I, Intrrsoience, New York, E.Y., 1949, p. 297. (11) A. M. Truchard, €1. G. Harris and D. A I . Himmelblau, J . Phys. Chem., 66, 575 (1961).
803
feet the concentrationq of the other components. Typical results are given in Table 11. TABLE I1 ETKYLENE REACTIOX WITH ( C ~ H ~ ) Z S O ~ - H Z S O ~ - C ~ H ~ S ~ ~ H MIXTURE6b
Sample
1 2 3 4 5 6 7
8
Ethylene reaction with mixture 6b Time reacted CCiHa (moles/l. 1 (min.)
0.0 30.3 68.0 120.5 201 .o 300.0 435.0 1091.5
4.79 x 10-2 3 . 3 9 x 10-2 2.30 X 1.42 X 0.716 X ,429 X l o d 2 ,254 X .2.50 x 10-2 ( C C ~final) H~
Concentration of mixture 6b
Experimental Ethylene (C.P. grade 1 was physically dissolved by violent agitation in mixtures of purified (CgH6)2SO4,H2S04(99.8 wt.%), and C2H6S04H having relatively high initial diethyl sulfate concentrations. The mixtures were then isolated from the source of ethylene, and the component concentrations monitored as reactions 1 and 2 proceeded at 30.00 i 0.02". First, the entire ethylene reaction system was evacuated for more t h m 30 minutes to remove from the liquid any dissolved inert gases m d any light component (such as ether) which might have been formed in trace amounts by some very slow side reaction wliile reaction 3 was taking place during preparation of the liquid phase. After evacuation, ethylene was added until the pressure in the system was about 200 mm. greater than atmospheric. Upon agitation of the liquid, the system pressure decreased rapidly by several mm., due to physical dissolution of the ethylene, and then the premure began to decrease more slowly, corresponding to reaction of ethylene with the solution. The mitxure wm allowed to react in this manner for about 15 minutes. Then the Rtirrer was turned off, and the reaction vessel filled with mercury until the reaction mixture level had risen into a capillryy a t the top of the system. Portions of solution were periodically withdrawn through a rubber serum cap with a calibrated hypodermic syringe. During withdrawal of the samples, mercury was released into the reaction vessel to maintain the reaction mixture level in the capillary. There was thus virtually no area for gas transfer into or out of the solution. The ethylene in each sample was measured volumetrically in a fashion quite similar to that employed in determining gas solubi1ity.lo To check the accuracy and reproducibility of this volumetric method of ethylene determination, ethylene was dissolved in diethyl sulfate in the ethylene reaction apparatus and samples analyzed for ethylene content. The average of five different determinations in this manner for ethylene solubility in diethyl sulfate gave a Henry's law constant H = 98.3 atm./mole fr., with a standard deviation of 0.6 atm./mole fr. Although the reproducibility was good, this value for H was 6.4% lower than that reported by Truchard, et aZ.11 No suitable explanation of this discrepancy was apparent.
ACID: KINETICS
c(C2Hs)zSO4
CHzS04
Sample
(moles/l.)
(moles/l.)
7" 2"
5.67 5.76 5.53 5.52
2.39 2.39 2.25 2.24
3b
4b
C (moles/l.) C~HKSO~H
1.33 1.23 1.59 1.62
Initial concentration of mixture 6b
CK~H~),J =O6.34 ~ moles/l. C H ~ S O ~ = 3.04 moles/l. Sampled 143 min. before ethylene reaction began. 133 min. after ethylene reaction began.
Sampled
I n order to be able to separate the effect of initial acid concentration from component concentration a t the time of the ethylene reaction, two mixtures were prepared at each of the different initial acid concentrations. The ethylene reaction mas then carried out for different degrees of completion of reaction 3 for each of the two mixtures until the ethylene concentration reached a final, seemingly unchanging value, This was not necessarily the value denoted by CC,II~~. which would be noted for C C ~ aHt ~system equilibrium, ie., equilibrium with respect to each of reactions 1, 2, and 3, so the concentration was not referred to as the equilibrium concexitration. The final value of CC,H~ for a fast reaction mixture was estimated with fairly good accuracy from a C C ~ us. H ~time plot. The rate of reaction of ethylene for some mixtures was so slow that it was not possible to follow them to find this terminal value; C C ~ for H ~ these ~ mixtures was assumed to be 0.3 X mole/l. for purposes of data analysis as discussed later. As will be shown, this assumed value was not critical. Ethylene could not be reacted with mixtures with a higher initial acid concentration than 37 mole % because the sampling of ethylene would have been necessary at periods of less than 25 minutes, the minimum time for one ethylene determination. Reactions a t lower acid concentration than 22 mole yo were not run because of the inordinate length of time that would have been required to get useful data. For most reaction mixtures plots of C C ~ H us.~time resulted in smooth curves. For two mixtures, however, the experimental points on the Cc2H4-time plot had a small amount of scatter, and the data did not seem as accurate. To check the reproducibility of the data, one complete run was repeated; these results were practically identical with the first trial.
H. G. HARRISAND D. 1LI. HIMMELBLAU
804 10,000
-
I
-
thus be readily integrated, and the rate constants determined. The following kinetic model based 011 the over-all stoichiometric equations 1and 2 as proposed
I
REACTION MIXTURE 6 0 . SLDPE.60X10'a(m~n)'' m -14~10'~(rninl" INITIAL GONCENTRATWS
C(C2H&W40,.62e mobs/l cHISO, - 3 1 4 mdos/l
-
T'ol. 67
d(CcZH4) ~-
- -liF~(cCzHa)
dt
(CHzSO,)
+ +
kF,(CCz€I,) (CCzIl,SOaH)
= m(CCzH,)
+
-
lcRl(CC~H6S04H)
kRz(C(C2H6)zS04)
(4)
where
Integrating (4) between the limits
gave 1
10
1
x
10-1
I
I
Thus the data for reactions 1 and 2 plotted according to equation 5 on semilogarithmic graph paper should give straight lines with slopes of m/2.303 if the proper kinetic model is that given in equation 4. As can be seen from a typical plot, Fig. 1, the data are well represented by equation 4. The rate of reaction of ethylene also proved to be dependent only upon the initial acid concentration and not a function of the DES, sulfuric acid, or ethyl hydrogen sulfate concentrations. This suggests one of two possibilities. The first is that kF, and JLIFp in equation 4 are identical and thus the reaction rate is constant for a fixed sum of CH~SO, and CC~H,SO,H, and is subsequently constant for all ( C2H&S04-C2H,S04H mixtures having the same initial concentration. Another possibility is that the actual mechanism involves not H2S04and CzHsSOdH but some property of the system that is constant for a fixed initial concentration. h substantial case has been establishedle-16 indicating that the mechanism for the hydration of olefins dissolved in aqueous acid is through reversibly formed r-complex precursors (or perhaps the u-type) with the later formation of a carbonium ion. The proposed scheme16is
3 1
0
I
I
1 2 3 Initial &SO4 concn., moles/l.
4
TEMPERATURE
30-c
I
Fig. 2.-Rate constants for C2H4 reaction with HzS04 and C2H6S04Has functions of initial HzS04 concentration.
Kinetic Models Because of the previously mentioned large difference between the rates of reaction 3 and the ethylene reaction, the concentrations of the components other than ethylene were virtually constant during each of the ethylene reactions. This greatly simplified the data analysis. The rate for any kinetic model could be expressed in terms of only one variable, CGH~,could
\
/
/
\
CH-C+
I + HS04- +\ CH-C-SOdH /
(8)
I
(12) E. R. Alexander, "Principles of Ionic Organic Reactions," John Wiley a n d Sons, Inc., New York. N. Y., 1850. (13) B. T. Brooks, "The Chemistry of the Non-benzoid Hydrocarbons," Reinhold Publ. Corp., New York, N. Y., 1950. (14) E. Kohn, Doctoral Diasertation, The University of Texas, 185G. (15) R. J. Gillespie and J . A. Leisten. Quad. Rcu. (London), 8, 40 (1854). (16) R. W. T a f t , E. L. Purlee, P. Ries. and C. A. Farico, J . Am. Chem. Soc., 77, 1582 (1985).
S. M.R. SPECTRA OF SILICOX HYDRIDE DERIVATIVES
April, 1963
with the limiting stage the transition of the a-complex into the carbonium ion, equation 7. Since reaction 6 is presumed t o be in equilibrium, it is easy to show
(9) where r represents the net product of all the activity coefficients(non-idealities) in the system. According to the Hammett hypothesis, if the transition state is formed simply by the addition of a proton to the substrate, then the reaction velocity will be proportional to the acidity function ho,but if the transition state also contains a water molecule, then the reaction velocity will increase less rapidly than ho and perhaps may be approximately proportional to the hydrogen ion concentration. The use of the Hammett acidity function to linearly correlate J C ~ ~data , ~ in the absorption of ethylene in solutions of S0-9570 ?&SO4 (and water) has been demonstrated by Korovina, Entelis and Chirkovl’ and Lumbroso, Hellin, and Coussemant. l 8 On the other hand Gel’bshtein arid Ternkin*were unable to find such a linear correlation from their data. I n any case, the pertinency of these results t o this investigation is not clear, since the diethyl sulfate and ethyl hydrogen sulfate concentrations in their solutions were small, and the water concentrations were substantial-in contrast to the present work which was essentially carried out in non-aqueous solutions. Furthermore, Melander and Myhrelg point out that the direct proportionality between the acidity function ho and the reaction rate is pot proof of the mechanism proposed by Taft and (17) G.V. Korovina, S. 0.Entelis, and N. M. Chirkov, J . A p p l . Chem. (English Transl.), 81,597 (1958). (18) D . Lumbroso, M. Hellin, and F. Coussemant, Compt. rend. Congr. Intern. Chim. 81’ Liege, 1068, 1, b24 (1959). (19) L. Melander and P. 0.Myhre, Arlcivfdr Kemi, 13, 507 (1959).
805
co-workers. TaftZ0postulated that the first step in the hydration of olefins has a rapid proton transfer to the a-bond, and that the rate determining step came later in the migration of the proton to one of the carbon atoms forming a carbonium ion. Melander and Myhre’s argument indicates that the kinetic evidence alone cannot establish the existence or non-existence of the intermediate a-complex. How well do the rate constants determined in this work correlate, and what is their significance in connection with the mechanism indicated by equation 91 If log IC is plotted against H o or -log CH+, neither plot gives a straight line with slope of unity; neither do they give straight lines with other slopes. Attempts to obtain a linear correlation with other proton indices were also unsuccessful. One of the basic problems is how to evaluate realistically both the proton concentration and the activity coefficients in diethyl sulfate solutions. One might adjust the activity coefficients so the calculated data fit the proposed mechanism, or modify the mechanism assuming the activity coefficient product is essentially constant (or do both a t the same time). Figure 2 shows the highly nonlinear variation of the rate constant; adjustment of the abscissa scale by one scheme or another to straighten out the curve does not appear to be fruitful. The “true” mechanism must be considerably more complex than that indicated by Taft, perhaps involving orders of reaction greater than unity. One must conclude that considerable caution is needed in interpreting kinetic data in essentially non-aqueous solutions based on niodels proposed for aqueous solutions. Acknowledgment.-This research was supported by a grant from The Petroleum Research Fund administered by the American Chemical Society. Grateful acknowledgment is hereby made to the donors of said fund. (20) R. W. Taft, Jr., etal., J . Am. Chem. Soc., 74, 5372 (1952); 77, 1584 (1955); 78, 5807, 5Sll (1956).
N.M.R. SPECTRA aF SILICON HYDRIDE DERIVATIVES: 11. CHEMICAL SHIFTS IN SOME SIMPLE DERIVATIVES’ BY E. A. V. EBSWORTH ASD J. J. TURNER University Chemical Laboratory, LensJield Road, Cambridge, England Received September 15, 1962 The proton and fluorine chemical shifts in a number of simple compounds containing SiH bonds are presented; the SiH chemical shifts are much less sensitive to the nature of the substituents than are the CH chemical shifts in analogous carbon compounds. The SiH resonance shifts to high field from fluorosilane to trifluorosilane. The 19Fchemigal shifts in fluorosilanes are consistent with a recently developed interpretation of lgF chemical shifts in fluorocarbon derivatives.
We have recently recorded the chemical shifts for protons and fluorine nuclei in a number of simple compounds containing SIH bonds; the results are presented in Tables 1-111. The values of the H-H, 2gSSiH, H-F, and 29Si-F coupling constants, which were also measured, are discussed elsewhere. Experimental The compounds mere prepared by standard methods, and were studied in the liquid phwe in concentrated (ca. 95%) and dilute (1) P a i t I, J . Chem. Phge., 86,2028 (1962).
solutions, using cyclohexane or tetramethylsilane (for proton resonances) or trichlorofluoromethane (for fluorine resonances) as solvent and internal standard. T h e samples were held in Pyrex tubing of 5 mm. 0.d.; the spectra were recorded using a Varian Associates V4300B spectrometer and 12 in. electromagnet, operating a t 40 Mc./sec., with flux stabilization and sample spinning; chemical shifts were measured using side-bands generated by a Muirhead-Wigan D695A decade oscillator. Errors quoted for proton resonances are derived from extrapolations, taking a t least six meaaurements a t each concentration, unless otherwise stated.
