Ind. Eng. Chem. Res. 1992,31, 1425-1433
1425
KINETICS AND CATALYSIS Solvent Effects on the Hydration of Cyclohexene Catalyzed by a Strong Acid Ion-Exchange Resin. 2. Effect of Sulfolane on the Reaction Kinetics H e n k - J a n P a n n e m a n and Antonie A. C . M.Beenackers* Department of Chemical Engineering, State University of Groningen, Nijenborgh 16, 9747 AG Groningen, The Netherlands
The kinetics of the hydration of cyclohexene, catalyzed by a strong acid ion-exchange resin, have been studied in a packed bed reactor at temperatures between 353 and 413 K and a pressure of 20 bar. The kinetic rate constants were measured as a function of temperature and solvent composition (0-90mol % solfolane in water). Macroporous Amberlite XE 307, a strong acid ion-exchange resin, was used as a catalyst. The influence of the solvent composition on the rate constant, which follows Arrhenius behavior, can be understood from changes in the activity coefficients of the reactant, the catalyst, and the activated complex. The decrease in rate constant between 0 and 60 mol % sulfolane is caused mainly by stabilization of cyclohexene. The observed increased rate constant above 60 mol % sulfolane is explained by an increased proton activity with increasing sulfolane content.
Introduction The production of alcohols by the catalytic hydration of the corresponding alkenes using macroporous strong acid ion-exchange resins has been known for about two decades. Deutsche Texaco (Neier and Woellner, 1973) reported the direct hydration of propene catalyzed by a strong acid ion-exchange resin in 1972. A mixture of liquid water and gaseous propene was fed to a four-stage trickle bed reactor. The temperature varied between 403 and 423 K and the pressure between 6 and 10 MPa; the annual production capacity was 100OOO tons. In 1981 Chemische Werke Hiils (Hiils, 1983) started a two-stage process for producing 50 OOO tons/yr of very pure tert-butyl alcohol (TBA). The hydration of isobutene in raffinate I (the butadiene-free C4cut from the steam cracker) is effected in the liquid phase between 313 and 373 K. By combining several reactors and various recycle streams, it was possible to achieve isobutene conversions up to 90% without the use of solubilizers. A serious drawback for a successful liquid-phase process is the low solubility of the alkene in water. The existing processes use either a trickle bed process or a high alcohol recycle stream to achieve high production rates together with a reasonable reactor capacity. The latter method is only practical for reactive alkenes with high equilibrium conversions. Chwang et al. (1977) measured the rates for the sulfuric acid catalyzed hydration of various alkenes at 298 K. They found that the rate constant for the hydration of cyclohexene is about loo00 times smaller than that for the hydration of isobutene. Therefore the existing processes are useless for the hydration of cyclohexene. A possible alternative is the use of a cosolvent to achieve one liquid phase which contains a fair amount of alkene. Different cosolvents have been suggested for the hydration of alkenes, e.g., solvents for the hydration of isoalkenes, such as 1,4-dioxaneand acetone (Chaplits, 19741, glycols and glycol mono- and diethers (Moy and Rakow, 1976), carboxylic acids (Matsuzawa et al., 1975),or lower
alcohols (Henke et al., 1962). However, a strong acid catalyst and high temperatures are necessary to obtain a satisfactory conversion of cyclohexene, a cyclic alkene. Under these circumstances none of the solvents suggested, except l,Cdioxane, are sufficiently inert in the strongly acidic environment within the resin particles. Possibly, nonbasic aprotic polar solvents like acyclic or cyclic sulfones (Van Broekhoven and Farragher, 1977; Drent, 1981; Okumura and Kaneko, 1981) can be used for the hydration of n-alkenes. In our laboratory Meuldijk et al. (1986,1989) investigated the influence of both sulfolane and 1,4dioxane on the kinetics of the hydration of 2-methylpropene in a gas-liquid system at 298 K. Prerequisites for a proper cosolvent for the specific acid-catalyzed hydration of cyclohexene are inertness under reaction conditions (ruling out solvents such as alcohols, ethers, ketones, and organic acids) and an absence of proton stabilization (ruling out solvents such as 1,4dioxane, dimethyl sulfoxide (DMSO), and other basic solvents). Only substituted sulfones satisfy both conditions and some are miscible with water in all ratios. In our study we used a cyclic sulfone, sulfolane, a relatively cheap solvent also used in extraction processes. We studied the hydration of cyclohexene catalyzed by a macroporous strong acid ion-exchangeresin. The effects of sulfolane on the equilibrium conversion of cyclohexene and on the solubility of cyclohexene in aqueous sulfolane mixtures have been described earlier (Panneman and Beenackers, 1992a,b). To our knowledge, intrinsic kinetics for the liquid-phase hydration of cyclohexene using the Amberlite XE 307 ion-exchange resin have not been published previously, nor is there information on the temperature dependence of the rate constant in different water-sulfolane mixtures. In this paper, initial reaction rate constants for the first-order reversible hydration of cyclohexene in some water-sulfolane mixtures are presented. The changes in the rate constants are explained in terms of changing ac-
0888-5885/92/2631-1425$03.00/00 1992 American Chemical Society
1426 Ind. Eng. Chem. Res., Vol. 31, No. 6,1992
tivities of the reactant cyclohexene and of the proton as a consequence of the addition of sulfolane to the reaction mixture. Experimental Section Initial hydration r a t a of cyclohexene in different solvent mixtures were measured. The catalyst preparation technique, the equipment used for measuring the kinetics, the analytical methods, and the experimental procedures applied in the determination of initial rate constants were the same as those used for the determination of the equilibrium conversion (Panneman and Beenackers, 1992b; Marsman et al., 1988). Briefly, the experiments were m, 1 = performed in a packed bed reactor (i.d. = 8 X 2 m) filled with a known amount of catalyst. The feed solution was prepared by mixing accurately weighed amounts of sulfolane and double-distilled water; the cyclohexene concentration was 50 m ~ l - m -always ~, far below its solubility. The feed was pumped through the reactor system, the preasure was kept at 2 MPa by a back-pressure valve, and the temperature was controlled within 0.2 K. Initial rate constants were calculated from the small conversions measured at the reactor outlet, when steady-state conditions were established. At least five different flow rates were employed a t each temperature. Cyclohexene and sulfolane were obtained from Janssen Chimica (purity >99% and >98%, respectively). For the determination of the Hammett acidity function, purified sulfolane (distillation under reduced pressure) was used. The catalyst used was hnberlite XE 307 from Rohm and Haas. Only narrow particle size fractions were applied in our experiments. Hammett’s method (Hammett, 1970) was used to determine the change in acidity function of a strong acid in different watel-sulfolane mixtures. The acidity function was determined by photometric detection of the dissociation of a proper indicator. Indicators with the right acidity constant and showing reversible adsorption were 3-nitroaniline and 4-nitroaniline, hereafter denoted as B. Acidity measurements with an ion-exchange resin containing immobilized sulfonic acid groups did not give reproducible results, probably due to adsorption of the indicator on the resin. Therefore these experiments were carried out with solutions of 0.1 M ptoluenesulfonic acid (pro analysi, Merck). The W absorption was measured at the wavelength of maximal absorption (k370 nm) of the neutral indicator, B. For each solvent composition four different samples were necessary. The first sample contained solvent mixture only, the second sample contained solvent with indicator (tg), and the third sample contained solvent, indicator, and p-toluenesulfonic acid (e). The final sample contained both solvent and nitrobenzene;nitrobenzene has the same absorption as fully protonated indicator (egH+). The experimentally determined Hammett acidity follows from (Braude and Stern, 1948) Ho = PKBH+- log (0 (1) where I represents the ratio of protonated and neutral indicator concentration in a solvent mixture: EB
-
I=-€ - €BH+
for all solvent mixtures, did not disturb the physical properties of the solvent mixtures. Where mass-transfer resistances could be neglected, we assumed thermodynamic equilibrium between the liquid and the resin phase. Consequently, the kinetic expressionsfor different solvent mixtures could be based on the appropriate bulk concentrations around the particle. The advantage of such an approach is that bulk concentrations can be easily measured contrary to actual concentrations in the resin. To our knowledge no open literature studies exist concerning the hydration mechanism of cyclohexene, when catalyzed by a strong acid ion-exchange resin. However, the hydration mechanism of alkenes catalyzed by strong acids in solution has been studied extensively (Nowlan and Tidwell, 1977; Chwang et al., 1977). It was found that the hydration reaction obeyed a pseudo-first-order rate law with respect to the alkene concentration as long as a large excess of water is used. For the hydration of propene and linear butenes catalyzed by strong acid ion exchangers, Petrus et al. (1984, 1986) proved that a similar reaction mechanism applies. Depending on the value of the catalytic activity of all accessible active centers, one has to use a more homogeneous or heterogeneous catalytic model. In hydrophilic solvents or solvent mixtures the catalytic activity is almost independent of the location of the active sites and the use of a homogeneous catalytic model is allowed (Buttersack, 1989). The rate-determining step of the mechanism is the protonation of carbon in the double bond, to give the corresponding carbonium ion (Chwang et al., 1977). A first-order dependence of the rate of hydration on the alkene concentration has been observed. Further, the large excess of water leads to a zero-order rate dependence for water. Also, the equilibrium conversion of cyclohexene decreases with increasing amount of sulfolane in the solvent mixture. (Panneman and Beenackers, 1992b). Therefore, it is likely that the hydration of cyclohexene can be described by a pseudo-first-order reversible rate equation. Using the symbols ENE and OL for cyclohexene and cyclohexanol, respectively, the resulting reaction rate, defined per equivalent of acid groups, can be written as -~ENE
= ~ E N -E k
- 1 ~ 0 ~
(3)
For the present purpose and under the conditions applied, the packed bed reactor can be considered as a plug flow reactor. The density of the liquid phase remains practically constant during reaction. The reactor was operated isothermally. For each experiment checks on the absence of both temperature effects and external mass-transfer limitations were made. With increasing temperatures, intraparticle mass-transfer limitations are expected to occur, as the effective intraparticle diffusion coefficients are much smaller than those in the bulk of the liquid. A number of experiments at high temperatures were done with smaller particles to avoid diffusional limitations. When internal mass-transfer limitations were present, the intrinsic rate constants were calculated with the concept of the effectiveness factor. The kinetic parameters were obtained from the experimentally observed conversions of cyclohexene using
(2)
Reaction Mechanism The reaction rate constants were determined both in water and in three wah-sulfolane mixtures containing 60, 78, and 90 mol 9%sulfolane, respectively. The very small cyclohexene concentration in the feed, which was the same
(4) and
Ind. Eng. Chem. Res., Vol. 31, No. 6, 1992 1427
0
1
2
3
4
5
Table 11. Experimentally Observed Apparent Rate Constants for the Hydration of Cyclohexene in Water with 90 mol % Sulfolane at Temperatures between 363 and 413 K and Four Different Average Catalyst Diameters 10@k,,a, m3.equiv-1-s-1, at d,, m temp, K 0.45 0.65 0.95 1.1 353.0 19.8 16.8 363.0 58.8 62.4 373.0 172 92.1 383.0 181 393.5 626 354 313 413.1 2210
10",l& (sm")
Figure 1. Typical plot of cyclohexanol conversion versus reciprocal flow rate (78 mol % sulfolane in water, T = 363 K,K, = 1.1). Table I. Experimentally Observed Apparent Rate Constants for the Hydration of Cyclohexene in Water-Sulfolane Mixtures at Temperatures between 348 K and 413 K and Average Catalyst Diameters of 0.65 X 10" and 0.95 X 1QSm i09kl,, m3.equiv-'.s-', at mol % sulfolane 78 (d, = 0.95 0 60 78 x 10-~m) temp, K 348.3 80.0 181 8.7 353.0 250 358.0 401 9.8 26.4 24.1 363.0 570 367.5 26.4 61 55 373.0 113 383.0 143 338 221 393.5 592 955 413.1
Here CAP represents the total number of acid groups present in the reactor (CAP = 0.1354 equiv). The ratio of both rate constants equals the equilibrium ratio K, multiplied by the molar fraction of water. K, is known for all solvent mixtures and at each temperature applied (Panneman and Beenackers, 1992b). A check of the first-order reversible rate equation was made by plotting the left-hand side of eq 4 versus the reciprocal volumetric flow rate, 0,. In Figure 1, typical experimental points are shown. From the slope of the straight line it follows that kl = 27 X lo* m3-equiv-'.s-'. The straight line through the origin, obtained by linear regression, confirms the validity of the proposed rate equation (eq 3).
Results The reaction rate constants for the hydration of cyclohexene were determined in water and in three watersulfolane mixtures with 60, 78, and 90 mol % sulfolane, respectively. At each temperature, the conversion S; was measured for five different residence times. This procedure was repeated for four or five different temperatures between 353 and 413 K. A number of different particle size fractions were used in the experiments. Most experiments were carried out with particles with sizes between 0.6 X and 0.7 X m. The experiments with 78 mol 9% sulfolane were repeated with particles between 0.9 X and 1.0X m in diameter. The experiments with 90 mol % sulfolane were done with four different particle size fractions. The apparent rate constants, kl,a,are possibly affected by intraparticle mass-transfer limitations, especially at higher temperatures, and therefore were not used to determine the intrinsic rate constants, k,. The results of experiments with water and mixtures with 60 and 78 mol % sulfolane are shown in Table I. The results of 90 mol % sulfolane are shown in Table 11. The
-I 0
I
I
0 d,=0.65. IO"m
-
L
--t -7
Y
2.3
2.5
IW/T
2.7
2.9
[K.')
Figure 2. Apparent (points) and intrinsic (line) reaction rate constants of hydration of cyclohexene in a solvent mixture with 78 mol % sulfolane.
-20
-
2.3
2.5
2.7
2.9
IOOO/T ( K - ' ] Figure 3. Apparent (points) and intrinsic (line) reaction rate constants of hydration of cyclohexene in a solvent mixture with 90 mol % sulfolane and different catalyst particle sizes.
value of the Thiele modulus 4 for a first-order reversible reaction in a porous particle is (Westerterp et al., 1984) dp
$=-
[kl(KxXH10 + 1)/%at11'2
6 DENEKJH~O and the degree of catalyst utilization follows from
(6)
4 usually is called the effectiveness factor. The experimenta with relatively large resin particles of 0.95 X and 1.1 X 10-3 m clearly show diffusion limitation at higher temperatures because here the apparent rate constanta are lower than for particles with an average size of 0.45 X and 0.65 X m. To obtain the intrinsic rate constante at the entire temperature range, the effectiveness factor 1 was determined by extrapolating the linear relationship between In (k,) and 1/T at low temperatures toward higher temperatures. See Figures 2 and 3 for 78 and 90 mol 9% sulfolane, respectively. From the effectivity factors thus obtained, a value for 4 was calculated for all experiments where the apparent reaction rate constants deviated from the regression line. All parameters in the Thiele modulus 4 are known except the value of the effective diffusion coefficient D m , which could be calculated this way. However, in most experi-
1428 Ind. Eng. Chem. Res., Vol. 31, No. 6, 1992 Table 111. Intrinsic Rate Constants and the Arrhenius Parameters for the Hydration of Cyclohexene in Water and Three SulfolaneWater Mixtures lO'kl,' m3.equiv-'d, at T,K sulfolane, 109kl,o, Ea, mol % 353 413 rn3.equiv-'d kJ.mol-* 103.2 24.75 0.277 0 0.153 0.033 108.1 60 0.0035 0.72 0.026 104.5 78 0.0089 1.54 0.014 100.4 90 0.0195 2.80 (I
-1 4
sublane ( mol%)
kl = kl,o exp(-E,/RT).
Figure 5. Intrinsic rate constanta of hydration of cyclohexene at 373 K 88 a function of solvent composition.
Table IV. Intrinsic Rate Constants for the Hydration of Cyclohexene at 373 K for Various Aqueous Sulfolane Mixtures sulfolane, 106kl,m3. sulfolane, l e k l , m3. mol % equiv-'d equiv-ld mol % 0 1.2 60 0.024 10 0.21 80 0.06 40 0.035 90 0.12 -1 0
I
)
I
~
0 MSUlf
V
-
60%sulf
-
I
(R) phase. Due to this equilibrium between both phases, eqs 8 and 9 apply for both cyclohexene and water-sulfolane
for all solvent mixtures (W/S): ai,R(W/S) = ai,L(W/S) or ~
-
0 78% SUlf
--
I
Y
t -
I
-20
2.3
I
2.5
I
2.7
IW/T
I
I
2.9
I
3.1
= Yi,Lci,L(w/s)
(9) The activity coefficients of both cyclohexene and protons are solvent dependent, as will be shown later. The number of acid groups remains the same, and also the initial cyclohexene concentration was kept the same for the different solvent mixtures. Consequently, the activity of cyclohexene and of protons changes with solvent composition. Assuming pseudothermodynamic equilibrium between the reactants and the activated complex ( a * ) ,the forward rate of the reaction is given by Yi,Rci,R(w/s)
1K.l)
Figure 4. Temperature dependence of intrinsic reaction rats constants of hydration of cyclohexene for 0,60, 78, and 90 mol % sulfolane in the feed.
menta the value of qj was smaller than 1so that no reliable value of DENE could be calculated. The experiments at 393 K with 90 mol 90sulfolane and using particles with mean diameters of 0.95 X and 1.1 X lV3m were the only two experiments that yielded values for the Thiele modulus that were above 1. The values obtained for the effective diffusion coefficients of cycloand 1.2 X hexene for both particle sizes are 1.4 X m2 -1 , respectively. This procedure was repeated for the other solvent mixtures, and the intrinsic rate constants obtained are shown in Table 111. This table also contains the parameters for the Arrhenius equation, which was used to correlate the temperature dependence of the rate constants. A graphical representation of the intrinsic reaction rate constant versus the reciprocal temperature for all watersulfolane mixtures is shown in Figure 4. From Figure 4, it follows that the reaction rate constant depends on the fraction of sulfolane with a minimum rate constant being found around 60 mol 9O sulfolane. To get a better description of this solvent effect, we carried out a number of experiments at 373 K with various solvent mixtures. The results are shown in Figure 5 and Table IV. Qualitatively,the solvent effect on the reaction rate constant is similar to that reported for the hydration of 2-methylpropene catalyzed by p-toluenesulfonic acid (Meuldijk et al., 1986) and catalyzed by strong acid ionexchange resins (Meuldijk et al., 1989). We will explain this solvent effect in terms of changing activities of both reactants and the activated complex. In this discussion we assume that there is always thermodynamic equilibrium between the liquid (L)and resin 0s
(8)
Since the activity of all components is the same in both liquid and resin phase, and the concentration of cyclohexene and protons is independent of the solvent mixture, the variation of the measured rate constant, kl,with solvent composition can be calculated from changes in the activity coefficients of the components in the liquid phase. Estimates of the changes in activity coefficients with solvent composition of both cyclohexene and protons were computed to show the extent to which both the initial state and the activated complex are responsible for the change in rate constant. The change in proton activity coefficients is determined in separate experiments with a water-soluble acid, and the results can be used only if the relation between the water-sulfolane ratio in the resin and that in the liquid phase is known. It is known that the distribution coefficients of small polar molecules between aqueous solutions and free water in the resin are about unity (Reichenberg and Wall, 1956). Sulfolane is a relatively small molecule with a large dipole moment, so we assume that the water-sulfolane ratio in both liquid and resin phase is the same. The solvent effect on the kinetics of the hydration of cyclohexene catalyzed by a macroporous strong acid ionexchange resin will generally be caused by a Gibbs free energy change of respectively the reactants, the catalyst, and the activated complex involved in the rate-determining reaction step (Abraham, 1974; Buncel and Wilson, 1979). To explain the differences in the kinetic parameters of the hydration of cyclohexene for different solvent mixtures, values for the transfer Gibbs free energy, AGi', of the species involved in the rate-determining step, which is the formation of the cyclohexyl carbonium ion, are required.
Ind. Eng. Chem. Res., Vol. 31, No. 6, 1992 1429 Table V. Observed Activation Parameters for the Hydration of Cyclohexene in Sulfolane-Water Mixtures and the Value of AG* at 373 K M', AG*, sulfolane, AH* mol % kJ-mol-' J.K-l-rno1-l kJ.mol-l 0 100.3 -29.0 111.1 60 104.9 -47.5 122.6 78 99.5 -54.0 119.6 90 97.3 -54.5 117.6 9
Unfortunately, it is not possible to determine directly the change in Gibbs free energy of the activated complex, so that it is impossible to give a complete theoretical description of the solvent effect on the hydration of cyclohexene. However, from changes in the Gibbs free energy of the reactants together with the kinetic data we have obtained, it is possible to calculate the change in Gibbs free energy of the activated complex. In case this change is small, the solvent effect will be caused mainly by the reactants and the catalyst. As mentioned above, an analysis of the solvent effects on the kinetics of the hydration of cyclohexene will be given by using transition-state theory (Engberta, 1979). hsuming pseudothermodynamic equilibrium between the reactants and the activated complex, the rate constant, k (d), is determined by the difference in Gibbs free energy (AG') between the activated complex and the reactants: k = kBT/h exp(-AG*/RT) = k,T/h exp(-AH*/RT) exp(AS*/R) (11)
k follows from the experimentally obtained intrinsic rate constants: k = bkatukat (12) Using eqs 11 and 12 and the experimentally obtained intrinsic rate constants, the activation parameters, AH* and AS*, can be calculated and the results are shown in Table V. The results show that the enthalpy of activation, AH*, shows the same changes with solvent composition as the energy of activation. Initially, there is an increase in enthalpy with a maximum around 60 mol ?& sulfolane; at higher sulfolane molar fractions the enthalpy decreases significantly. The entropy of activation, AS*,is decreasing continuously with increasing sulfolane percentage. The ' 9 sulfolane. greatest change occurs between 0 and 60 mol 0 The Gibbs free energy of activation, AG*, shows a course which is the inverse of the reaction rate constant's value as a function of the mole percentage sulfolane in the solvent mixture. The decrease of the rate constant between 0 and 60 mol ?& sulfolane is mainly an entropy effect, while the increase of rate constant in solvents with more than 60 mol 90sulfolane is caused exclusively by the decrease of AH*. It will be shown that both the increase of AH* and the decrease of AS*between 0 and 60 mol ?& sulfolane are mainly caused by the stabilization of cyclohexene. The decrease of AH* at higher molar fractions sulfolane is caused mainly by a destabilization of protons. For a quantitative explanation of the solvent effect, expressions for the changing Gibbs free energies of all species involved must be found. The explanation of the solvent effect for the kinetics of the hydration of cyclohexene in different solvent mixtures can be given in terms of changing activity coefficients. With the calculated activity coefficients, it is possible to obtain thermodynamic transfer functions, like the transfer Gibbs free energy, AGtr(i),which is given by AGtr(i) = R T In ( y i , s / y i , w ) (13) In eq 13 y refers to the activity coefficient of component
-
4
8
-10
B
-20 sulfobne (mol%)
Figure 6. Transfer Gibbs free energy of methylpropene and cyclohexene in aqueous sulfolane mixtures at respectively 298 and 373 K.
i in different solvents, S refers to any sulfolane-water mixture, and W refers to water. One method used to obtain activity coefficients is the measurement of solubilities. However, for sulfolane-water mixtures at higher temperatures the solubility of cyclohexene in the mixtures is quite large. Then the theory of an ideal solution has lost its validity and the determination of activity coefficients from solubility data is not straightforward. The transfer Gibbs free energy of cyclohexene was determined by using activity coefficients obtained with the UNIFAC method, according to
with YENE,S the activity coefficient of cyclohexene in any sulfolane-water mixture and YENE,W that of cyclohexene in water. The calculation of AG,(ENE) with the UNIFAC method was checked by using the same method for the calculation of the transfer Gibbs free energy of methylpropene in water-sulfolane mixtures at 298 K. The amount of methylpropene in the liquid mixture is always small (less than 2 mol W);thus, here, reliable values of AG,(MPR) can be obtained from solubility measurements. With solubility data obtained by Meuldijk et al. (1986) the experimental transfer Gibbs free energy, AGJMPR), follows from
In Figure 6, the results are shown as a function of the mole percentage sulfolane in the mixture. The points represent the transfer Gibbs free energy calculated from solubility data, and the line represents AGJMPR) calculated with activity coefficients of methylpropene obtained with UNIFAC. It can be concluded that the UNIFAC method is suitable for the calculation of activity Coefficients of methylpropene in various sulfolane-water mixtures. The excellent agreement between experimental and calculated transfer Gibbs free energies of methylpropene justify the calculation of transfer Gibbs free energies of cyclohexene in different sulfolane-water mixtures at 100 "C with UNIFAC. The results are also shown in Figure 6. The obtained negative transfer Gibbs free energies imply a decrease in the Gibbs free energy of cyclohexene and consequently a decrease in activity of cyclohexene. Addition of sulfolane leads to a stabilization of cyclohexene; it becomes considerably less reactive in solvent mixtures with sulfolane. After examination of the other possible contributions to the solvent effect on the rate constant, it will be clear whether the large changes in cyclohexene
1430 Ind. Eng. Chem. Res., Vol. 31, No. 6, 1992 (Y is the degree of ionic dissociation of acid HA, if a strong acid is used, then (Y is 1. The salt-effect ratio is usually assumed to deviate little from unity. Also, for an organic conjugate acid-base pair, the medium effect can often be - log neglected. Then, according to eq 17, the term (-Ho cHA) should represent a first approximation of log rn YH+: log, YH+ = -HO - log CHA (18)
-a
0
20
40
60
80
100
Sutfobne ( moP.6 J
Figure 7. Experimentally observed Hammett acidity function for binary mixtures of water and sulfolane at 298 K,containing 0.1 M p-toluenesulfonicacid. The value of log, yH+(W+S) at 298 K,the proton medium effect, was calculated according to eq 19.
activity are more or less responsible for the change in rate constant. The observed variation in the reaction rate constants with solvent composition also could, at least in part, be caused by a change in catalytic activity of the ion-exchange resin. The changes in the proton activity in binary sulfolane-water mixtures with p-toluenesulfonic acid will be used as a first approximation of the changes in the proton activity of the ion-exchange resin in various solvent mixtures. The proton activity is related to the Hammett acidity function (Ho),according to Ho = -log (aH+)- log (YB/YBH+) (16)
Ho is a measure of the medium's ability to protonate a given indicator B. A general assumption is that log (yB/yBH+)remains constant if the acid concentration or the solvent composition is changed (Rochester, 1970; Hammett, 1970). Ho was measured experimentally for a number of aqueous sulfolane mixtures containing 0.1 M p-toluenesulfonic acid; see Figure 7. Starting from pure water, it can be seen that the acidity of the solution, as expressed in -Ho, decreases with increasing sulfolane content. If the solvent mixture contains about 40 mol % sulfolane, a minimum is reached in the proton activity. For higher molar fractions of sulfolane, an increase in acidity is observed again. Sulfolane is a dipolar aprotic solvent (Klein et al., 1981). A common feature of these solvents is that no intermolecular hydrogen bonds exist between the solvent molecules. Although sulfolane has a moderately high dielectric constant (43.3 at 303 K), it nevertheless has very little tendency to form hydrogen bonds with other protic solutes like water. Sulfolane has neither acidic nor basic properties and breaks the structure of water (Blandamer and Burgess, 1975). At first this will decrease the acidity, because more water molecules are available for stabilizing the proton. Above 40 mol % sulfolane, the activity of the proton increases because there are fewer water molecules giving less stabilization. A similar behavior of the proton activity is shown in other water-cosolvent mixtures (Boyd, 1969). Below, a first approximation of the transfer Gibbs free energy of the proton or, in other words, of the medium effect of the proton, from water to any binary sulfolanewater mixture was made, using the Hammett acidity function. The general relationship between the medium effect of the proton, log, YH+, and the Hammett acidity function is (Popovych, 1970) log, YH+ = -HO - log CHA - log (Y YBH+
YBH+
YB
YB
log, YH++ log, -+ log, - (17) where log, are the salt-effect activity coefficient terms and
However, the true value of log, YHt will be somewhat higher, because the terms -log, YH+, -log a,and log, YBH+ all are positive. With increasing content of the organic solvent they should become increasingly more positive. This trend is observed indeed for ethanol-water mixtures + been and methanol-water mixtures, for which log, y ~has accurately measured (Popovych et al., 1972; Popovych, 1984). Unfortunately such data are not available for sulfolane-water mixtures. The Hammett acidity functionsHoof ethanol-water and methanol-water are known from literature (Boyd, 1969) and Ho of sulfolane-water mixtures from Figure 7. All three curves have a similar shape, though ethanol and methanol are more basic solvents than sulfolane, and so the minimum in the curve is lower for ethanol and to a some extent also for methanol. The difference was determined between values of log, YH+ calculated according to eq 18 and values of log, YH+ measured by Popovych (1984). The differences are similar for ethanol and methanol and thus little solvent dependent for solvents with comparable physical properties. Now it is possible to estimate a more accurate value of log, yH+for aqueous sulfolane mixtures by adding to the right-hand side of eq 18 the experimentally observed extra term as reported by Popovych. So the proton medium effect in sulfolane-water mixtures log, yH+(W-S) can be estimated as log, ')"+(W-S) = -HO - log cH++ [log, "IH+(W+M)- (-HOW- 1% cHt)l (19) where M represents a methanol-water mixture. In Figure 7 values of log, yH+(W-s), calculated according to eq 19, are shown as a function of the mole fraction of sulfolane in the mixture. The values of log, yH+(W-S) shown in Figure 7 are probably not very accurate, but they are most likely more reliable than the values calculated according to eq 18. The transfer Gibbs free energy of the proton between water and a water-sulfolane mixture, AGt,(H+,W-S), can be calculated easily from the proton medium effect according to AG,,(H+,W-S)
= RT In (10) log, yH+(W*S) (20)
The transfer Gibbs free energy of the proton calculated with eq 20 can be used as a first approximation of the transfer Gibbs free energy of a proton in an ion-exchange resin at 373 K. The values will be used for the calculation of the transfer Gibbs free energy of the initial state. We now have all the data necessary for the computation of the change in the Gibbs free energy of the activated complex, AGkT, if the solvent composition is changed. An illustration of the relationship between the transfer free energies of reactants, AGtrR,and of the activated complex, AGtIT,and the free energies of activation, AG*, for a reaction occurring in two solvent systems is shown qualitatively in Figure 8. The change in the Gibbs free energy of the activated complex, AGtIT,can be evaluated from calculated values of the transfer free energies of the reactants, AGkR, in conjunction with the measured kinetic activation parameters, bAG', according to AGtIT = AGtIR + 6AG* = AGt,(ENE) + AGt,(H+) + (AGs* - AGw') (21)
Ind. Eng. Chem. Res., Vol. 31, No. 6,1992 1431
-water - - rullolane-
water mlxtun
Surobne 1 mol%)
Figure 9. Transfer Gibbs free energies of reactants, AG,(ENE) and AG,(H+), transfer Gibba free energy of activation 6AG', and transfer Gibbs free energy of activated complex, AGt?, at 373 K.
-
reactlon
ordlnate
Figure 8. Relationship between transfer Gibbs free energies of reactants, activated complex, and products and Gibbs free energies of activation, in two solvent systems. Table VI. Values of ACtrB, bAG', and AG,,T for the Hydration of Cyclohexene in Various Sulfolane-Water Solvent Mixtures at 373 K sulfolane, AGtXR, 6AG', mol 90 kJ-mol-' kJ-mol-' kJ.mo1-' 0 0 0 0 60 -16.6 11.5 -5.1 78 -12.4 8.5 -3.9 90 -8.0 6.5 -1.5
The transfer Gibbs free energy of cyclohexene and that of the proton follow from eqs 14 and 20, respectively. The last term of eq 21, the difference in Gibbs free energy of activation between solvent system S (a sulfolane-water mixture) and W (water), directly follows from Table V. In Table VI, the transfer Gibbs free energies of the initial state, the transfer Gibbs free energies of activation, and the transfer Gibbs free energies of the activated complex in mixtures of sulfolane and water are presented for various molar fractions of the cosolvent. A qualitative graphical representation of the transfer Gibbs free energies as a function of the molar percentage of sulfolane in the solvent mixture at 373 K is given in Figure 9. From the variation of the transfer Gibbs free energies of the initial and activated complex as a function of the molar fraction of sulfolane, it is clear that the large stabilization of cyclohexene is almost entirely responsible for the decrease of the rate constant with increasing sulfolane content of the reaction mixture until 60 mol % sulfolane. At the lower molar fractions of sulfolane, the proton is stabilized relative to pure water; the same is true for the activated complex. The change of both ionic species is of comparable size. At higher molar fractions of sulfolane the proton activity increases with increasing sulfolane content, while the corresponding decrease in Gibbs free energy of cyclohexene is relatively small. This results in an increasing value (less negative) of the transfer Gibbs free energy of the initial state with increasing sulfolane content. Above 60 mol % sulfolane the transfer Gibbs free energy of the activated complex has a rather small negative value over the entire solvent range, indicating a small stabilization of the activated complex or an extension of the number of realizations, leading to an increasing entropy of the activated complex. The ambiguity of the activated complex-it contains a apolar organic part and an ionic part-is responsible for a small sensitivity toward the
solvent mixture composition. The solvent effect on the kinetics visualized by the transfer Gibbs free energy of activation, 6AG', is mainly caused by changes in the Gibbs free energy of the initial state. For low sulfolane contents a stabilization of both cyclohexene and protons gives considerably lower rate constants relative to pure water. Above 60 mol % sulfolane, the change in Gibbs free energy of cyclohexene is small and protons become destabilized so that the rate constant increases again with increasing sulfolane content. In general the reaction rate for the hydration of an alkene remains nearly the same if cosolvent is added as long as the solution remains saturated with the alkene and Y ~ + does not change. This is true because the decrease in rate constant is compensated by the increase of solubility of the alkene: POL
= ~YH+CH+(YENECENE- YOLCOL/K) (22)
Whether a cosolvent is practically useful depends on the behavior of the proton activity coefficient and of the change in Gibbs free energy of the activated complex. Solvents, like 1,4-dioxane and DMSO,which stabilize protons will never give an increased rate. Only aprotic polar solvents like sulfolane are suitable. At high mole percentages cosolvent yH+will increase considerably,while the activated complex is still stabilized, because sulfolane has a high dipole moment.
Conclusions The liquid-phase hydration of cyclohexene catalyzed by the macroporous strong acid ion-exchange resin Amberlite XE 307 proceeds according to a reversible rate equation, which for 353 K < T < 413 K is first order in cyclohexene and zero order in water, provided water is in sufficient excess. The rate of hydration of cyclohexene shows a strong dependence on the composition of the sulfolane-water solvent mixture. For mole fractions of sulfolane increasing from 0 to 0.6, the rate constant decreases continuouslyand considerably. The rate constant shows a minimum value, which is 40 times smaller than without sulfolane, at about 60 mol % sulfolane. Above 60 mol % sulfolane the rate constant sharply increases again with sulfolane content. The intrinsic reaction rate constants show an Arrhenius type dependency on temperature. The preexponential factor and EA depend on the amount of sulfolane in the solvent mixture. The rate constants are experimentally determined between 353 and 413 K, while the sulfolane fraction varied between 0 and 0.9. With Amberlite XE 307, there exist intraparticle diffusional limitations in solvent mixtures with 78 mol % sulfolane and particles of 0.95 X and 0.65 X m at temperatures above 393 and 413 K, respectively. In
1432 Ind. Eng. Chem. Res., Vol. 31, No. 6, 1992 solvent mixtures with 90 mol % sulfolane, diffusional limitations occurs at 393 K with particles of 0.95 X m. The kinetic results can be understood mainly from changes in the Gibbs free energy of the initial state of both
cyclohexene and the proton, while the activated complex shows only minor changes in activity. Between 0 and 60 mol 70sulfolane, there is a large stabilization of cyclohexene, which is responsible for the decrease of the rate constant in that solvent composition region. To a lesser extent also the protons and the activated complex are stabilized here. Above 60 mol % sulfolane, the activity of protons increases with increasing sulfolane content while the activities of cyclohexene and the activated complex show only minor variation in that region. Taken together, this explains the observed increase of the rate constants above 60 mol % sulfolane.
Acknowledgment This work was part of the research program of the Dutch Foundation for Chemical Research (SON)and was made possible by financial support from the Dutch Organisation for the Advancement of Pure Research (NWO). We thank DSM for additional financial support. Nomenclature uH+ = proton actvity in solution ci = concentration of species i in the liquid phase, k m ~ l - m - ~ CAP = number of acid groups in reactor system, equiv Di = internal diffusion coefficient of i in a porous particle, m2 *S-1
d, = particle diameter, m EA = energy of activation, J-mol-l ENE = cyclohexene AC&) = transfer chemical potential of species i between two solvent mixtures (W-S), J-mol-'
AG*= Gibbs free energy difference between activated complex and initial state, J-mol-' ACtrT = transfer Gibbs free energy of activated complex
(W-S),J.mo1-' AG,? = transfer Gibbs free energy of the reactants (W-S), J-mol-'
6AG* = transfer Gibbs free energy of activation (W-S), J. mol-'
h = Planck constant, J.s
Ho= Hammett acidity function HA = acid in solvent mixture AH* = enthalpy difference between activated complex and initial state, J-mol-' Z = ratio of protonated and neutral indicator B IS = initial state kB = Boltzmann constant, J-K-' k = reaction rate constant of t h e forward reaction, s-' kl = reaction rate constant per equivalent of acid, m3equiv-l-s-' kl,a = apparent reaction rate constant with diffusion, m3. equiv-l-s-' kl,o = preexponential factor of Arrhenius equation, s-l K, = mole fraction equilibrium ratio log, yH+ = medium effect on proton activity coefficient MPR = 2-methylpropene OL = cyclohexanol p = pressure, Pa pKBH+ = acidity constant of BH+ R = gas constant, J.mol-'.K-' ri = molar rate of production of species i, k m ~ l . m - ~ d Si = solubility of species i, k m ~ l - m - ~ AS* = entropy difference between activated complex and initial state, J.mol-'.K-' T = temperature, K
x L = liquid mole fraction of species i Greek Symbols a = degree of dissociation of acid HA yi = activity coefficient of i 5; = conversion of cyclohexene for a specific flow rate feq = equilibrium conversion of cyclohexene q = effectiveness factor of reaction in porous particle &.at = density of resin, kgIE.mc3 u = specific capacity of the solvent-swollen resin, equiv.kgE-' I$ = Thiele modulus 9, = volumetric flow rate, m3-s-* Registry No. Amberlite XE 307, 52038-28-5; cyclohexene, 110-83-8;sulfolane, 126-33-0;hydrogen ion, 12408-02-5.
Literature Cited Abraham, M. H. Solvent Effects on Transition States and Reaction Rates. Progr. Phys. Org. Chem. 1974,11, 1. Blandamer, M. J.; Burgess, J. Kinetics of Reactions in Aqueous Mixtures. Chem. SOC. Rev. 1975, 4, 55. Boyd, R. H. Acidity Functions. In Solute-Solvent Interactions; Coetzee, J. F., Ritchie, C., Eds.; Marcel Dekker: New York, 1969. Braude, E. A.; Stern, E. S. Acidity Functions. Part 11. The Nature of Hydrogen Ion in Some Aqueous and Non-aqueous Solvents. The Exceptional Solvating Properties of Water. J. Chem. SOC. 1948,1976. Buncel, E.; Wilson, H. Initial-State and Transition-State Solvent Effects of Reaction Rates and the Use of Thermodynamic Transfer Functions. Acc. Chem. Res. 1979, 12, 42. Buttersack, C. Accessibility and Catalytic Activity of Sulfonic Acid Ion-ExchangeResins in Different Solvents. React. Polym. 1989, 10, 143. Chaplits, D. N. Recovery of isobutylene, US. Patent 4 012 456,1974. Chwang, V. K.; Nowlan, V. J.; Tidwell, T. T. Reactivity of Cyclic and Acyclic Olefinic Hydrocarbons in Acid-Catalyzed Hydration. J. Am. Chem. SOC. 1977,99, 7233. Drent, E. A Process for the Preparation of an Olefiiically Unsaturated Ether. (Shell Internationale Research Maatschappij B.V.) European Patent 0 025 240, 1981. Engberts,J. B. F. N. Mixed Aqueous Solvent Effecta on Kinetica and Mechanisms of Organic Reactions. In Water (a Comprehensive Treatise); Franks, F., Ed.; Plenum Press: New York, London, 1979; Vol. 6. Hammett, L. P. Physical Organic Chemistry, 2nd ed., McGraw-Hik New York, 1970. Henke, A. M.; Odioso, R. C.; Schmid, B. K. Hydration of olefins in the present of a solvent. (Gulf Research & Development Co.) Br. Patent 973 832, 1962. H a s , Chemische Werke. Production of tert-butanol and isobutylene. Chem. Eng. 1983, Dec 12,60. Klein, S.D.; Pawlak, Z.;Fernandez-Prini, R.; Bates, R. G. Conductance of HCl in Water-Sulfolane Solvents at 25,30, and 40 O C ; A Comparison of Conductance Equations. J. Solution Chem. 1981, 10, 333. Marsman, J. H.; Panneman, H. J.; Bennackers, A. A. C. M. Automated On-Line Gas ChromatographicInjection of Samples Completely Liquified by Pressure. Chromatographia 1988, 26, 383. Matsuzawa, H.; Ikeda, M.; Sugimoto, Y.; Uchida, S. Tertiary Butyl Alcohol. (Mitsubishi Rayon Co., Ltd.) Ger. Offen. 2 430 470,1975. Meuldijk, J.; Joosten, G. E. H.; Stamhuis, E. J. Effect of Solvent on the Hydration of 2-Methylpropene in Solutions of p-toluenesulfonic acid and poly(styrenesulfonic acid). J. Mol. Catal. 1986, 37,75. Meuldijk, J.; Joosten, G. E. H.; Stamhuis, E. J.; Beenackers, A. A. C. M. Effect of Solvent on the Kinetics of the Hydration of 2methylpropenecatalyzed by Strong Acid Ion Exchangers. J.Mol. Catal. 1989,54, 183. Moy, D.; Rakow, M. S. High Purity Isobutylene Recovery. (Cities Service Co.) U.S. Patent 4 096 194, 1976. Neier, W.; Woellner, J. Isopropyl Alcohol by Direct Hydration. Chemtech 1973,3,95. Nowlan, V. J.; Tidwell, T. T. Structural Effects on the Acid-Catalyzed Hydration of Alkenes. Acc. Chem. Res. 1977, 10, 252. Okumura, Y.; Kaneko, K. Tert-Butyl Alcohol. (Toa Nenryo Kogyo K.K.) US.Patent 4 270 011,1981. Panneman, H. J.; Beenackers, A. A. C. M. Solvent Effects on the Hydration of Cyclohexene Catalyzed by a Strong Acid Ion Exchange Resin. 1. Solubility of Cyclohexene in Aqueous Sulfolane
Ind. Eng. Chem. Res. 1992,31, 1433-1440 Mixtures. Ind. Eng. Chem. Res. 1992a,31, 1227. Panneman, H. J.; Beenackers, A. A. C. M. Solvent Effects on the Hydration of Cyclohexene Catalyzed by a Strong Acid Ion-Exchange Resin. 3. Effect of Sulfolane on the Equilibrium Conversion. I d . Eng. Chem.Res. 1992b,followingpaper in this iesue. P e w , L.; de Roo,R. W.; Stamhuis, E. J.; Joosten, G. E. H. Kinetics and Equilibria of the Hydration of Propene over a Strong Acid Ion-Exchange Resin aa Catalyst. Chem. Eng. Sci. 1984,39,433. P e t m , L.;de Roo, R. W.; Stamhuis, E. J.; Joosten, G. E. H. Kinetics and Equilibria of the Hydration of Linear Butenes over a Strong Acid Ion-ExchangeResin aa Catalyst. Chem. Erg. Sci. 1986,41, 217. Popovych, 0. Estimation of Medium Effects for Single Ions in Non-Aqueous Solvents. Crit. Rev. Anal. Chem. 1970,1,73. Popovych, 0. Transfer Activity Coefficients of Ions in Methanol-
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Water Solvents Based on the Tetraphenylborate Assumption. J. Phys. Chem. 1984,88,4167. Popovych, 0.;Gibofsky,A.; Berne, D. H. Medium Effects for Single Ions in Acetonitrile and Ethanol-Water Solvents Based on Reference-Electrolyte Assumptions. Anal. Chem. 1972,44,811. Reichenberg,D.; Wall, W. F. The adsorption of uncharged molecules by ion exchange resins. J. Chem. SOC.1956,3364. Rochester, C. H. Acidity Functions;Academic Press: London, 1970. Van Broekhoven, J. A. M.; Farragher, A. L. Alcohols (Shell Internationale Research MaatachappijB.V.) Ger. Pat., 2,721,206,1977. Westertarp, K. R.;van Swaaij, W. P. M.; Beenackers, A. A. C. M. Chemical Reactor Design and Operation; Wiley: Chicester, 1984.
Received for review June 25, 1991 Accepted January 15, 1992
Solvent Effects on the Hydration of Cyclohexene Catalyzed by a Strong Acid Ion-Exchange Resin. 3. Effect of Sulfolane on the Equilibrium Conversion Henk-Jan Panneman and Antonie A. C. M. Beenackers* Department of Chemical Engineering, State University of Groningen, Nijenborgh 16,9747 AG Groningen, The Netherlands
The liquid-phase hydration of cyclohexene, a pseudo-first-order reversible reaction catalyzed by a strong acid ion-exchange resin, was investigated in solvent mixtures of water and sulfolane. Macroporous Amberlite XE 307 was used because of its superior catalytic activity. Chemical equilibrium conversions were measured as a function of temperature, 353 K C T < 423 K, and solvent composition (between 0 and 90 mol % sulfolane in water). A decrease by a factor of 3 and 6 is observed in the experimentally measured equilibrium conversion for solvent mixtures with 60 and 90 mol % sulfolane, respectively. This effect is explained in terms of activity coefficients of the reaction species involved. Classical methods of thermodynamics allowed the computation of the equilibrium constant K,,as exp(30236/RT). From use of the predictive UNIFAC method to describe the K , = 2.37 X liquid-phase nonideality, it was possible to estimate the equilibrium conversion of cyclohexene for every solvent composition and temperature. The predicted equilibrium conversions agree well with experimental results.
Introduction The industrially important hydration of propene and butenes (Neier and Wollner, 1973) to produce the corresponding alcohols is catalyzed by a strong acid. There are two different processes to produce the alcohols on a commercial scale. The first and oldest process uses a 5 0 4 % aqueous solution of sulfuric acid. Some serious drawbacks of this process are corrosion and environmental problems. The second process uses a strong acid ion-exchange resin as a catalyst. The hydration of alkenes catalyzed by strong acid ion-exchange resins was studied by Petrus (1982). The very low mutual solubility of alkenes and water gives small production rates of the alcohols per unit volume of ionexchange resin. A possible alternative for obtaining higher production rates is the use of a cosolvent, which results in a single liquid phase and a greatly improved solubility of alkenes in the solvent mixture. In the patent literature various solvents are mentioned for the hydration of alkenes, such as 1,4dioxane, acetone, and glycol ether (Schmerling, 1971) and sulfolane and related sulfones (Van Broekhoven and Farragher, 1977). In our laboratory Meuldijk et al. (1986, 1989) investigated the influence of both sulfolane and 1,4-dioxaneon the kinetics of the hydration of 2-methylpropene. In the present study the hydration of cyclohexene, catalyzed by a macroporous strong acid ion-exchange resin, OSSS-5SS5/92/2631-1433$03.00/0
was investigated. Investigating the hydration of cyclohexene has some experimental advantages relative to the hydration of propene and butenes. Firstly, no complicafhg isomerization or side reactions occur in cyclohexene hydration. Secondly, butenes and propene are gases at room temperature, giving practical problems in pumping the feed and in the analysis of the product stream. According to our findings, sulfolane (tetrahydrothiophene 1,l-dioxide)is the only suitable cosolvent, which satisfies all conditions (Panneman and Beenackers, 1992a). The cosolvent has to be completely miscible with both water and all the organic components. The solvent must behave really inert under reaction conditions, which are high temperatures and strong acidities, and the solvent is not allowed to have any proton affinity, which would decrease the acidity of the resin. Addition of sulfolane will not only increase the solubility of cyclohexene in the reaction mixture, but also change the equilibrium conversion, the reaction rate, and the activity of the ion-exchange resin. Also the costs of product separation and coeolvent recycle have to be taken into account, for determining the optimal economical process conditions. Recently, two Japanese companies have published results about the liquid-phase hydration of cyclohexene. A patent of Fukuhara et al. (1990) describes the preparation of cyclohexanol by hydration of cyclohexene in the presence of a strong acid ion-exchange resin, immediately 0 1992 American Chemical Society
