Ind. Eng. Chem. Res. 2007, 46, 6865-6872
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Experimental Study of the Chemical Equilibria in the Liquid-Phase Dehydration of 1-Pentanol to Di-n-pentyl Ether Roger Bringue´ , Javier Tejero,* Montserrat Iborra, Jose´ Felipe Izquierdo, Carles Fite´ , and Fidel Cunill Chemical Engineering Department. UniVersity of Barcelona. C/ Martı´ i Franque` s 1, 08028-Barcelona, Spain
The thermodynamic equilibrium of the liquid-phase bimolecular dehydration of 1-pentanol to di-n-pentyl ether (DNPE) and water was studied in the temperature range of 423-463 K over Amberlyst 70. Furthermore, the equilibrium position of two side reactions could be followed, DNPE decomposition to 1-pentanol and 1-pentene and isomerization of 1-pentene to 2-pentene. The etherification reaction proved to be slightly exothermic, with an enthalpy change of reaction at 298.15 K of -(3.8 ( 0.6) kJ mol-1. From this value, the standard formation enthalpy and molar entropy of DNPE were computed to be -(421.1 ( 1.2) kJ mol-1 and 473.71 J (K‚mol)-1, respectively. The enthalpy changes of the reaction of DNPE decomposition to 1-pentene and 1-pentanol and 1-pentene isomerization to 2-pentene were 63.4 ( 0.9 and -19.7 ( 2.1 kJ mol-1, respectively. 1. Introduction Diesel fuel specifications are becoming increasingly stringent as legislation is adopted to improve air quality by reducing emissions. It is expected that forthcoming diesel fuels will be characterized by a higher cetane number, lower density, and lower aromatic, polyaromatic, and sulfur content with respect to the present one. A feasible option to comply with these regulations might be the use of reformulated diesel fuels containing appropriate high-quality components.1 A few years ago, in a comprehensive study on the blending properties of oxygenates in diesel fuels,2 it was observed that linear ethers with g9 carbon atoms showed the best balance among blending cetane number and cold flow properties. Di-n-pentyl ether (DNPE) was selected as a good candidatefor diesel fuel reformulation because of its blending properties and the availability of potential feedstocks. DNPE has a blending cetane number of 109, and due to its density and viscosity (both a bit lower than those of commercial fuels), it behaves as a light diesel fuel. Moreover, DNPE was shown to be very effective in reducing diesel exhaust emissions such as CO, NOx, unburned hydrocarbons, particulates, and smokes.3,4 1-Butene is an appropriate feedstock for the industrial manufacturing of DNPE. The synthesis route would consist of selective hydroformylation and hydrogenation of 1-butene to 1-pentanol, followed by the bimolecular dehydration reaction of the alcohol to dialkyl ether. However, if the catalyst is not selective, the monomolecular dehydration to alkene occurs,5 which, at its turn, could react with the remaining alcohol to form undesired branched ethers of lower cetane numbers. Therefore, if high selectivity to linear ethers is desired, it is necessary to minimize the production of alkenes. In a previous work, the liquid-phase dehydration of 1-pentanol to DNPE without water removal was studied on acidic styrenedivinylbenzene (S/DVB) resins, including sulfonated and oversulfonated gel and macroporous ones.6 It was found that gel-type catalysts were suitable for the dehydration of * To whom correspondence should be addressed. E-mail: tejero@ angel.qui.ub.es.
Figure 1. Evolution of DNPE activities over time at different temperatures.
1-pentanol to DNPE, but thermal stability was an important drawback for most of them since the working temperature was limited to 423 K, which is rather low to have suitable reaction rates. Recently, thermally stable resins, such as Amberlyst 70, have been successfully obtained by halogenating the aromatic rings as well as the aliphatic polymer backbone, some of them showing losses of less than 10% of the sulfonic acid groups in hydrolytic stability tests performed at 473 K for 24 h.7 Laboratory tests on Amberlyst 70 showed that it is a good option as a catalyst for the studied reaction because its thermal stability, up to 473 K, led to high activity with great selectivity to DNPE8 at 463 K. At those activity and selectivity levels, industrial processes for obtaining DNPE from 1-pentanol appear to be feasible. However, to the best of our knowledge, thermodynamic data such as the enthalpy change of reaction and equilibrium constants have not been reported. Therefore, this work was undertaken to determine experimental values of the equilibrium constant by direct measurement of the mixture composition at several temperatures after reaching the chemical equilibrium. Then, the standard enthalpy change of DNPE synthesis was calculated and compared with estimated values. Besides DNPE synthesis, the two main side reactions, that is, the DNPE decomposition reaction to 1-pentanol and 1-pentene and the isomerization of 1-pentene to form 2-pentene (cis and trans), have been also studied at equilibrium.
10.1021/ie0616646 CCC: $37.00 © 2007 American Chemical Society Published on Web 09/14/2007
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Table 1. Experimental Conditions and Calculated Equilibrium Constants for the Dehydration of 1-Pentanol to DNPE and Water, DNPE Decomposition to 1-Pentanol and 1-Pentene, and Isomerization of 1-Pentene to 2-Pentene (cis and trans) T [K]
W [g]
KDNPE x
a KDNPE γ
KDNPE a
K1-pentene ‚104 b x
K1-pentene γ
‚‚104 K1-pentene a
c K2-pentene x
K2-pentene γ
K2-pentene a
423 433 433 443 443 453 453 463
4.265 3.226 3.63 1.666 4.215 1.014 1.700 2.100
22.7 23.3 23.0 23.7 24.8 23.7 23.1 24.4
2.440 2.371 2.342 2.200 2.104 2.150 2.088 1.954
55.5 55.2 53.9 52.0 52.2 50.8 48.2 47.7
2.7 4.2 4.1 7.4 5.6 9.2 10.2 13.1
1.616 1.637 1.644 1.601 1.757 1.633 1.671 1.864
4.3 6.8 6.7 11.8 9.9 14.9 17.0 24.4
18.3 15.1 16.8 13.6 13.5 12.5 11.8 11.7
1.053 1.053 1.050 1.039 1.029 1.035 1.026 1.010
19.3 15.9 16.8 14.1 13.9 12.5 12.1 11.8
a The DNPE superscript refers to 1-pentanol dehydration to DNPE and water. b The 1-pentene superscript refers to DNPE decomposition to 1-pentanol and 1-pentene. c The 2-pentene superscript refers to 1-pentene isomerization to 2-pentene (cis and trans).
2. Experimental Section 2.1. Materials. 1-Pentanol (99% pure, 98.5% pure), and 2-pentene (cis + trans, >98.5%) were supplied by Fluka. 1,4-Dioxane (99,5%) was supplied by Panreac. Di-n-pentyl ether (DNPE, >99%) was obtained and purified in our laboratory. The thermally stable resin Amberlyst 70 (Rohm and Haas) used as the catalyst is a macroporous, halogenated, and sulfonated copolymer of styrene/divinylbenzene (S/DVB) in the H+ form. The acidity of the ion-exchange resin was measured by titration against standard base,9 resulting in 3.01 mequiv of H+ (g of dry resin)-1. The catalyst was used in its commercial form (dp ) 551 µm). 2.2. Apparatus. Experiments were carried out in a 100 mL stainless steel autoclave operated in the batch mode. A magnetic drive turbine was used for mixing, and baffles were placed inside of the reactor to improve the agitation. Temperature was controlled to within (1 K by an electric furnace. The pressure was set at 1.6 MPa by means of N2 in order to maintain the reacting mixture in the liquid phase over the whole temperature range. One of the reactor outlets was connected directly to a liquid sampling valve, which injected 0.2 µL of pressurized liquid into a gas-liquid chromatograph. More detailed information can be found elsewhere.6,8,10 2.3. Analysis. The composition of liquid mixtures was analyzed by using a split-mode operation in a HP6890A GLC apparatus (Hewlett-Packard) equipped with a TCD detector. A 50 m × 0.2 mm × 0.5 µm methyl silicone capillary column was used to separate and quantify 1,4-dioxane, 1-pentanol, DNPE, water, C5 alkenes (1-pentene, 2-pentene), and branched ethers (1-(1-methylbutoxy)pentane, 1-(2-methylbutoxy)pentane, 2-(1-methylbutoxy)pentane, and 2-(2-methylbutoxy)pentane). The column was temperature-programmed with a 6 min initial hold at 318 K, followed by a 30 K min-1 ramp up to 453 K and held for 5 min. Helium was used as the carrier gas at a total flow rate of 30 mL min-1. 2.4. Procedure. 1-Pentanol is partially soluble in water, and DNPE is not very soluble in water at 20 °C. In previous experiments, we realized that alcohol conversions higher than 80% could be obtained at equilibrium in the temperature range explored. As a consequence, to ensure that a single phase exists through equilibrium experiments and to prevent the splitting of the reaction medium into two phases, the option of adding a solvent that allowed and favored mutual solubility of water, alcohol, and ether was selected. 1,4-Dioxane was selected because it is stable (from a physical and chemical standpoint) at working conditions and also because it is easily determined in the chromatographic analysis performed. This methodology was tested by Delion et al.11 in the liquid-phase hydration of isobutene with eight different solvents. Delion et al. concluded that the presence of a solvent modified
Figure 2. Evolution of activities during an experiment at 453 K over time.
Figure 3. Evolution of the calculated equilibrium constant of an experiment at 453 K over time. Table 2. Mean Experimental Values of the Equilibrium Constants, with the Associated Standard Error, of the Three Chemical Reactions with Temperature T [K]
KDNPE a
K1-pentene‚104
K2-pentene a
423 433 443 453 463
55.5 ( 0.5 54.6 ( 0.9 52.1 ( 0.1 49.5 ( 1.9 47.7 ( 1.2
4.3 ( 0.3 6.8 ( 0.1 10.9 ( 1.3 15.9 ( 1.4 24.4 ( 2.5
19.3 ( 0.6 16.8 ( 0.6 14.0 ( 0.2 12.3 ( 0.3 11.8 ( 0.1
the equilibrium shift of the reaction since the activities of all the compounds were also modified. However, the equilibrium constants computed from the equilibrium activities of all of the systems were in agreement with the thermodynamic constant, that is, the methodology proved to be suitable for equilibrium constant measurements. Catalyst particle diameter measurements were performed in 1,4-dioxane to check whether it has any influence on the
Ind. Eng. Chem. Res., Vol. 46, No. 21, 2007 6867
catalyst structure. A Beckman Coulter LS particle size analyzer coupled to a universal liquid module was used for this purpose. Results showed that 1,4-dioxane did not swell the catalyst particles. Blank experiments performed with the same amount of catalyst used in subsequent equilibrium experiments showed that 1,4-dioxane did not react under the experimental conditions. A volume of 70 mL of different mixtures of 1,4-dioxane, 1-pentanol, di-n-pentyl ether, water, 1-pentene, and 2-pentene (cis and trans) near the equilibrium composition were charged into the reactor together with the dried catalyst and, after checking for leakages, heated to the desired temperature. To monitor the concentration variation of chemicals with time, liquid samples were taken out of the reactor periodically and analyzed as mentioned above. Experiments were performed in the temperature range of 423-463 K. The stirring speed was fixed at 350 rpm, a lower value that ensured a good agitation without external masstransfer effects,8 to prevent catalyst particles attrition. The resin was dried at 390 K under vacuum for 2h. The catalyst load varied from 1 to 4.3 g, depending on the temperature, in order to shorten experiment landing at lower temperatures. Experiments finished when equilibrium conditions were reached, which was shown when the experimental thermodynamic equilibrium constants had the same value within the limits of the experimental error (typically after 48 h of running).
S
Ka )
S
(a′i)νe ) ∏ (γi)νe (xi)νe ) ∏ i)1 i)1 i
i
S
∏ i)1
S
(γi)νe i
(xi)νe ) Kγ‚Kx ∏ i)1 i
(1)
Kγ values were calculated by
) KDNPE γ
K1-pentene ) γ
γDNPE‚γwater γ1-pentanol2
(2)
γ1-pentene‚γ1-pentanol γDNPE
(3)
γ2-pentene γ1-pentene
(4)
) K2-pentene γ
where superscripts DNPE, 1-pentene, and 2-pentene refer to 1-pentanol dehydration to DNPE and water, DNPE decomposition to 1-pentanol and 1-pentene, and 1-pentene isomerization to 2-pentene, respectively. Kx was calculated in a similar way, by changing activity coefficients by molar fractions
) KDNPE x
3. Results and Discussion Besides the etherification reaction products, 1-pentene and especially 2-pentene (cis and trans) were detected to some extent in preliminary experiments at high temperatures and long reaction times. 1-Pentene could be produced by means of the monomolecular dehydration of 1-pentanol to 1-pentene and water and/or by decomposition of DNPE to 1-pentanol and 1-pentene. Since the 1-pentanol concentration was very low throughout the experiments, the most likely reaction for 1-pentene production was DNPE decomposition. 2-Pentene was produced through isomerization of 1-pentene. Some branched ethers, for example, 1-(1-methylbutoxy)pentane, 1-(2-methylbutoxy)pentane, 2-(1-methylbutoxy)pentane, and 2-(2-methylbutoxy)pentane, were detected in very small amounts. Since 2-pentanol was not detected, these ethers were formed by direct reaction of 1-pentanol and 2-methyl-1-butanol with olefins. A more detailed description of byproduct formation can be found elsewhere.6 The formation of branched ethers shifted the main reaction to the decomposition of the ether; therefore, the system reached a quasi-equilibrium state wherein the molar fractions of DNPE and 1-pentanol had a very slight trend to decrease, whereas those of water rose very slowly. In Figure 1, the evolution of DNPE activities over time at different temperatures is shown. This fact was observed at higher temperatures when attaining a constant composition was rather difficult. In this case, the assessment of the equilibrium state was done by checking the constancy of the calculated equilibrium constant, within the limits of the experimental error, instead of checking the composition constancy. In order to consider the nonideality of the mixture, activity coefficients of compounds, γ, were estimated by the UNIFACDORTMUND predictive method.12 The thermodynamic equilibrium constant for a liquid-phase reaction of a nonideal system is given by
i
) K1-pentene x
xDNPE‚xwater x1-pentanol2
(5)
x1-pentene‚x1-pentanol xDNPE
(6)
x2-pentene x1-pentene
(7)
K2-pentene ) x
Figure 2 shows the evolution of activities of 1,4-dioxane, DNPE, 1-pentanol, water, 1-pentene, and 2-pentene of a model experiment at 453 K. Activities increased or decreased to reach equilibrium or quasi-equilibrium conditions in ∼48 h. As the system was diluted with 1,4-dioxane, its activity was clearly higher than that of the others (it should be noted that the y axis is broken from 0.4 to 0.7). Water activities were also quite high compared to those of the others mainly because of its high activity coefficient (higher than 2). On the other hand, 1-pentene activities were always very small (around 0.004) due to its low molar fraction. This fact supposed a drawback for the equilibrium characterization since 1-pentene concentrations were near the detector threshold of the chromatograph, especially at lower temperatures. In Figure 3, the evolution of calculated Ka’s over time is shown for the same experiment as that in Figure 2, which states that pseudo-equilibrium conditions were achieved. Table 1 shows the experimental conditions of the performed experiments and the calculated equilibrium constants for the dehydration of 1-pentanol to DNPE and water, the DNPE decomposition reaction to 1-pentanol and 1-pentene, and the isomerization of 1-pentene to 2-pentene. Duplicate runs were made at some temperatures, and the reproducibility of experiments was found to be reliable. Equilibrium constant values presented in Table 1 are the average of those calculated at quasiequilibrium during each experiment. As it can be seen, the catalyst mass used had no effect on the measured Ka, as expected. On the other hand, values of Kγ were significantly and K1-pentene , which different from unity, especially KDNPE γ γ proves the nonideality of the mixture.
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Table 3. Molar Volumes of 1-Pentanol, DNPE, Water, 1-Pentene, and 2-Pentene and KΓ Correction Factors for the Three Reactions T [K]
VPeOHa [L mol-1]
VDNPE [L mol-1]
Vwater [L mol-1]
V1-pentene [L mol-1]
V2-pentene [L mol-1]
KDNPE Γ
K1-pentene Γ
K2-pentene Γ
423 433 443 453 463
0.125 0.127 0.129 0.132 0.134
0.205 0.208 0.212 0.215 0.219
0.019 0.019 0.020 0.020 0.020
0.170 0.180 0.193 0.207 0.225
0.144 0.153 0.164 0.178 0.195
0.989 0.989 0.989 0.989 0.989
1.039 1.042 1.046 1.051 1.056
0.989 0.988 0.988 0.988 0.988
a
Calculated by the HBT method.14
Table 4. Thermochemical Data of 1-Pentanol, DNPE, Water, 1-Pentene, and 2-Pentene units ciT2 +
Cpi ) ai + biT + ai bi ci di ∆fH°(l) (298.15 K) S° (298.15 K)
diT3
mol-1
1-pentanol
DNPE
water
1-pentene
2-pentene
177.73a 0.1872 -3.456‚10-4 7.892‚10-7 (-351.62 ( 0.28)e 258.9h
-126.85b 2.3799 -3.240‚10-3 1.550‚10-6 (-435.2 ( 3.0)f 394.44i
106.61c -0.2062 3.777‚10-4 -1.226‚10-7 (-285.830 ( 0.040)c (69.95 ( 0.03)c
46.12d 0.8055 -2.694‚10-3 4.236‚10-6 (-46.94 ( 0.42)g 262.6j
56.75d 0.7201 -2.563‚10-3 4.096‚10-6 (-58.24 ( 0.42)g 256.6k
K-1
J J mol-1 K-1 J mol-1 K-2 J mol-1 K-3 J mol-1 K-4 kJ mol-1 J mol-1 K-1
a The a , b , c , and d are estimated by the Rowlinson-Bondi method and fitted to a third-order equation.14 b Obtained by calorimetry and fitted to a i i i i third-order equation. c Calculated from Shomate equation and fitted to a third-order equation.15 d Estimated by the Lyman-Dannen method and fitted to a third-order equation.14 e Mosselman et al.16 f Murrin et al.17 g Wilberg et al.18 h Counsell et al.19 i Estimated by a modified Benson method.20 j Messerly et al.21 k Chao et al.22
Table 5. Temperature Dependence Parameters of K, ∆H°(l), ∆S°(l), and ∆G°(l) for the Three Chemical Reactions units IK IH a b c d
DNPE
1-pentene
2-pentene
J/mol 47749.1 11844.8 -19969.58 adim 253.004 -227.536 -8.65609 J/(mol K) -375.689 350.693 10.6329 J/(mol K2) 1.79924 -1.38719 -8.5368 × 10-2 J/(mol K3) -2.171 × 10-3 2.007 × 10-4 1.307 × 10-4 J/(mol K4) -1.514 × 10-7 3.476 × 10-6 -1.401 × 10-7
Mean values of the equilibrium constant at each temperature are shown in Table 2. As can be observed, dehydration of 1-pentanol to DNPE and isomerization of 1-pentene to 2-pentene (cis and trans) are exothermic since their equilibrium constants decrease with temperature. KDNPE values are high enough to state that the main reaction is clearly shifted to the products at equilibrium, which assures good conversion levels of 1-pentanol to ether in industrial etherification processes. Moreover, it hardly changes with temperature, which points out that conversion is quite promising to produce the ether in all of the experimental temperature range. On the other hand, the DNPE decomposition reaction to 1-pentanol and 1-pentene is endothermic. As a consequence, this reaction could be a drawback in an industrial reactor operated at conversions of 1-pentanol close to equilibrium, especially at high temperatures. Fortunately, the equilibrium constant of this side reaction is very low in the temperature values suggest that the range explored. Finally, the K2-pentene a isomerization reaction is clearly shifted to 2-pentene (cis and trans). This agrees with the general rule of alkene stability, which states that the most substituted alkene should predominate in the thermodynamic equilibrium state. Deviation in Ka values due to the difference between the working pressure and the pressure at the standard state was evaluated by means of the Poynting correction factor KΓ. It can be estimated by the following expression13
KΓ ) exp
[
P-1 RT
S
]
νiVi ∑ i)1
(8)
where Vi is the molar volume of compound i. Molar volumes14 and KΓ correction factors for the three reactions are shown in Table 3. It can be seen that neglecting
KΓ introduced an error in the calculation of Ka that is lower than the experimental error. Therefore, it can be assumed that the equilibrium constant is only a function of temperature. The thermodynamic equilibrium constant can be related to thermodynamic variables of the reaction system by
Ka ) exp
(
)
-∆rG°(l) RT
(9)
The standard free-energy change for the liquid-phase reaction can be computed from the standard enthalpy and entropy changes, as follows
∆rG°(l) ) ∆rH°(l) - T‚∆rS°(l)
(10)
The temperature dependence of the equilibrium constant can be expressed by using eqs 9 and 10
ln Ka )
∆rS°(l) R
-
∆rH°(l) RT
(11)
If the enthalpy change of reaction is assumed to be constant over the temperature range, by fitting eq 11 to experimental values of the equilibrium constant (Figure 4), the standard molar enthalpy change of reaction, ∆rH°(l), can be obtained from the slope and the standard molar entropy change of reaction, ∆r S°(l), from the intercept. The experimental temperature dependence of Ka for each reaction was found to be
) ln KDNPE a
(783.42 ( 75.52) + (2.18 ( 0.17) T
(12)
(-8464.8 ( 106.2) + (12.26 ( 0.24) (13) T (2549.3 ( 249.0) ) - (3.08 ( 0.53) (14) ln K2-pentene a T
) ln K1-pentene a
Ind. Eng. Chem. Res., Vol. 46, No. 21, 2007 6869 Table 6. Energy, Enthalpy, and Entropy Changes of DNPE Synthesis, DNPE Decomposition, and 2-Pentene Formation in the Liquid Phase at 298 K DNPE
∆rH°(l) constant ∆rH°(l) as f(T) theoreticala a
1-pentene
2-pentene
∆rH°(l),298 (kJ mol-1)
∆rS°(l),298 (J mol -1 K-1)
∆rG°(l),298 (kJ mol-1)
∆rH°(l),298 (kJ mol-1)
∆rS°(l),298 (J mol -1 K-1)
∆rG°(l),298 (kJ mol-1)
∆rH°(l),298 (kJ mol-1)
∆rS°(l),298 (J mol -1 K-1)
∆rG°(l),298 (kJ mol-1)
-6.5 ( 0.6 -3.8 ( 0.6 -17.8
18.1 ( 1.4 25.7 ( 3.1 -48.4
-11.9 ( 1.1 -11.5 ( 0.3 -3.4
70.4 ( 0.9 63.4 ( 0.9 36.7
101.9 ( 2.0 83.1 ( 4.8 122.1
40.0 ( 2.0 38.6 ( 0.5 0.3
-21.2 ( 2.1 -19.7 ( 2.1 -11.3
-25.6 ( 4.7 -21.6 ( 11.0 -6.0
-13.6 ( 3.5 -13.3 ( 1.2 -9.5
Computed from data given in Table 4.
As a consequence, the standard enthalpy and entropy changes of reaction for every reaction were, respectively
∆rH°(l) ) -6.5 ( 0.6 kJ/mol
DNPE:
∆rS°(l) ) 18.1 ( 1.4 J/(K‚mol) 1-Pentene: ∆rH°(l) ) 70.4 ( 0.9 kJ/mol ∆rS°(l) ) 101.9 ( 2.0 J/(K‚mol) 2-Pentene: ∆rH°(l) ) -21.2 ( 2.1kJ/mol ∆rS°(l) ) -25.6 ( 4.7 J/(K‚mol) On the other hand, considering that the standard enthalpy of reaction changes significantly over the temperature range, its dependence on temperature can be computed by the Kirchoff equation
d∆rH°(l) dT
S
)
νi‚Cp(l)i ∑ i)1
(15)
where Cp(l) are the molar heat capacities in the liquid phase of the compounds that take part in the reaction and are usually expressed in the polynomial form13,14
Cp(l)i ) ai + biT + ciT + diT 2
3
(16)
By integrating eq 15, combined with eq 16, the following expression can be obtained
b c d ∆rH°(l) ) IK + aT + T2 + T3 + T4 2 3 4
(17)
where S
a)
νiai ∑ i)1
S
b)
νibi ∑ i)1
S
c)
νici ∑ i)1
S
d)
νidi ∑ i)1
(18)
The dependence of the equilibrium constant on temperature is obtained from the van’t Hoff equation
d ln Ka ∆rH°(l) ) dT RT2
(19)
Combined with eq 17, the integration of 19 leads to
ln Ka ) IH -
IK d 3 a b c 2 + ln T + T + T + T RT R 2R 6R 12R
(20)
In Table 4, thermochemical data of the compounds involved in the three reactions are shown.13-21 The integration constants IK and IH can be calculated from the temperature dependence relationship for the experimental equilibrium constant of each reaction. By fitting eq 20 to the experimental values of equilibrium constants at different tem-
Table 7. Standard Molar Enthalpy and Entropy of the Liquid-Phase Synthesis of Some Alkyl Ethers Computed from Molar Formation Enthalpies and Entropies ether
∆rH°(l) [kJ/mol]
∆rS°(l) [J/(K‚mol)]
DME DEE DNPrE DNBE DNPE
-10.2 ( -11.1 ( 2.6c -9.4 ( 1.4e -9.8 ( 1.5e (-17.8 ( 4.1g) -6.5 ( 0.6h -3.8 ( 0.6i
-37.86b 3.73d 8.25f N/Aj (-48.41g) 18.10h 25.86i
0.9a
a Pilcher et al.,23 Majer et al.,24 and Chao et al.25 b Carlson et al.26 and Kennedy et al.27 c Pedley et al.28 and Murrin et al.17 d Haida et al.29 and Counsell et al.30 e Mosselman et al.16 and Colomina et al.31 f Counsell et al.19 and Andon et al.32 g Theoretical from data of Table 4. h Experimental with ∆rH°(l) constant. i Experimental with ∆rH°(l) variable with temperature. j N/A: the molar entropy of DNBE is not reported yet in data banks.
peratures (Figure 5), IK can be obtained from the slope and IH from the intercept. By means of eqs 10, 11, and 17, the standard molar changes of reaction ∆S°(l) and ∆G°(l) as a function of temperature for each reaction can be obtained as follows
c d ∆rS°(l) ) R‚IH + a + a ln T + bT + T2 + T3 (21) 2 3 b c d ∆rG°(l) ) IK - R ‚IHT - aT ln T - T2 - T3 - T4 (22) 2 6 12 In Table 5, values of IK, IH, and a, b, c, and d for each reaction studied are shown. Values of the standard molar enthalpy, entropy, and free energy changes of the three reactions at 298.15 K determined for ∆rH°(l) assumed to be constant and for the ∆r H°(l) variable with temperature are gathered in Table 6. As can be seen, ∆rH°(l) for the DNPE synthesis reaction at 298.15 K was not very high, even if ∆rH°(l) was considered to be constant or a function of temperature, which was in agreement with the low sensitivity to temperature of KDNPE . a Both values were lower than that obtained from the molar formation enthalpies of 1-pentanol, DNPE, and water shown in Table 4. However, they are in the trend showed by values of other di-n-alkyl ethers computed from experimental enthalpy changes of formation, ∆fH°(l), found in data banks, as can be seen in Table 7.23-32. The ∆fH°(l) for 1-pentanol given in Table 4 seems to be a reliable value, as a similar one was found in another experimental work, -(352.57 ( 0.72).33 On the other hand, only Murrin et al.17 found a value for DNPE computed from combustion enthalpies. From ∆rH°(l) values in Table 6 and ∆fH°(l) values of 1-pentanol and water in Table 4, ∆fH°(l) values for DNPE at 298.15 K could be obtained: -(423.9 ( 1.2) kJ mol-1 if ∆rH°(l) was considered to be constant or -(421.1 ( 1.2) kJ mol-1 if it was variable with temperature. These values were lower by 3% than that proposed by Murrin et al.,17 -(435.2 ( 3.0), and by 1.5% than that estimated by an improved Benson group-additive method,20 (-430.1 kJ mol-1). In Table 7, the standard entropy of reaction of some alkyl ethers are also shown. By comparing ∆rS°(l) for these reactions,
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Figure 4. ln K versus 1/T. Comparison between values obtained experimentally and those predicted from eqs 12, 13, and 14 (solid line). (A) DNPE synthesis reaction, (B) DNPE decomposition reaction, and (C) 1-pentene isomerization reaction.
it seems that the entropy change increases as the number of carbon atoms of the molecule does. The value found in this work agreed with this trend, but the expected one from the data of Table 4 did not. The molar entropy of DNPE is still not reported in data banks, and the modified Benson method seems to underestimate it. An S°(l) of 473.71 J (K‚mol)-1 (465.95 if ∆f H°(l) is considered constant with temperature) for DNPE could be computed from ∆rS°(l) and from data in Table 4, which was nearly 20% higher than the one estimated by the modified Benson method (394.4 J (K‚mol)-1). Deviations of the improved Benson method in the estimation of the standard entropy, although not so important, were also observed for di-n-propyl
Figure 5. ln K + f(T) versus 1/T. Comparison between values obtained experimentally and those predicted from eq 20 with data of Table 5 (solid line). (A) DNPE synthesis reaction, (B) DNPE decomposition reaction, and (C) 1-pentene isomerization reaction.
ether, DNPrE, the symmetrical ether derived from propanol. The experimental value,31 323.9 J (K‚mol)-1, is 5% higher than the one estimated by the modified Benson method (308.6 J/(K‚ mol)). To the best of our knowledge, no experimental value of S°(l) for di-n-butyl ether, DNBE, has been published yet. As a general conclusion, it seems that deviations on the estimation of S°(l) of lineal symmetrical ethers by the Benson method increase with the number of carbons in the molecule. Experimental values of a ∆rH°(l)1-pentene differed considerably from the theoretical ones (see Table 6). Activities of 1-pentene and 1-pentanol were very low in all experiments, especially for 1-pentene at 423 and 433 K. Thus, equilibrium constant values of the DNPE decomposition reaction at these
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temperatures probably had an important inaccuracy associated. Although less important, the same occurred when computing ∆rH°(l)2-pentene; experimental values were higher than those obtained theoretically from thermochemical data, which are very similar to gas-phase values found in the literature, that is, ∆rH°(g)2-pentene ) -10.9 ( 0.8 kJ/mol.34 Taking into account the vaporization enthalpies of 1-pentene and 2-pentene, ∆rH°(l)2-pentene ) -12.2 ( 0.8 kJ/mol is obtained, relatively close to the value estimated by us within the limits of the experimental error (see Table 6) 4. Conclusions The thermodynamic equilibrium of the bimolecular dehydration of 1-pentanol to DNPE and water, the DNPE decomposition reaction to 1-pentanol and 1-pentene, and the isomerization of 1-pentene to 2-pentene were studied in the temperature range of 423-463 K. The etherification reaction proved to be slightly exothermic, with an enthalpy change of reaction of -(3.8 ( 0.6) kJ mol-1 at 298.15 K (-6.5 ( 0.6 kJ mol-1 if ∆rH°(l) was assumed to be constant over the temperature range). These values agreed fairly well with the trend of values of other di-n-alkyl ether reactions from the corresponding alcohol and allowed computation of a ∆fH°(l) of -(421.1 ( 1.2) kJ mol-1 (-423.9 ( 1.2 kJ mol-1 if considered constant) and a S°(l) of 473.71 J (K‚mol)-1 (465.95 J (K‚mol)-1 if ∆fH°(l) was considered constant with temperature) for DNPE formation. The DNPE decomposition reaction to 1-pentanol and 1-pentene proved to be endothermic, with a ∆rH°(l) of 63.4 ( 0.9 kJ mol-1 (70.4 ( 0.9 kJ mol-1 if considered constant), and the 1-pentene isomerization reaction was exothermic, with a ∆rH°(l) of -19.7 ( 2.1 kJ mol-1 (-21.2 ( 2.1 kJ mol-1 if considered constant). Acknowledgment The authors are thankful for financial support from State Education, Universities, Research & Development Office of Spain (Project PPQ2000-0467-P4-02). The authors are also grateful to Rohm and Haas for providing the ion-exchange resin used in this work. Nomenclature ai ) activity of compound i ai, bi, ci, di ) polynomial form coefficients of heat capacity expression of compound i a, b, c, d ) temperature dependence coefficients Cp(l)i ) molar heat capacity of compound i in the liquid phase (J mol-1 K-1) dp ) particle diameter (µm) DME ) dimethyl ether DEE ) diethyl ether DNPrE ) di-n-propyl ether DNBE ) di-n-butyl ether ∆rG°(l) ) standard free energy change of reaction in the liquid phase (kJ mol-1) ∆fH°(l) ) liquid-phase standard molar enthalpy change of formation (kJ mol-1) ∆rH°(l) ) standard molar enthalpy change of reaction in the liquid phase (kJ mol-1) IH ) van’t Hoff integration constant IK ) Kirchoff equation integration constant Kja ) thermodynamic equilibrium constant of reaction j Kjγ ) activity coefficients ratio of reaction j
KΓ ) Poynting correction factor Kjx ) thermodynamic equilibrium constant of reaction j based on molar fractions P ) pressure R ) gas constant (J mol-1 K-1) S° ) liquid-phase molar entropy (J mol-1 K-1) ∆rS°(l) ) standard molar entropy change of reaction in the liquid phase (J mol-1 K-1) T ) temperature (K) Vi ) molar volume of compound i W ) weight of dry catalyst (g) xi ) molar fraction of i Greek Letters γi ) activity coefficient of i νi ) stoichiometric coefficient of compound i Superscripts DNPE ) 1-pentanol dehydration to DNPE and water 1-pentene ) the DNPE decomposition reaction to 1-pentanol and 1-pentene 2-pentene ) the 1-pentene isomerization reaction to 2-pentene Literature Cited (1) Douaud, A. Tomorrow Engines and Fuels. Hydrocarbon Process. 1995, 74, 55. (2) Pecci, G. C.; Clerici, M. G.; Giavazzi, M. G.; Ancillotti, F.; Marchionna, M.; Patrini, R. Oxigenated Diesel Fuels. Part 1 Structures and Properties Correlation. Proceedings of the Ninth International Symposium on Alcohol Fuels Vol. 1, Florence, Italy, 1991; p 321. (3) Giavazzi, F.; Terna, D.; Patrini, D.; Ancillotti, F.; Pecci, G. C.; Trere`, R.; Benelli, M. Oxigenated Diesel Fuels. Part 2. Practical Aspects of Their Use. Proceedings of the Ninth International Symposium on Alcohol Fuels Vol. 1, Florence, Italy, 1991; p 327. (4) Van Heerden, J.; Botha, J. J.; Roets, P. N. J. Improvement of Diesel Performance with the Addition of Linear Ethers to Diesel Fuels. Proceedings of the Twelfth International Symposium on Alcohol Fuels Vol. 1, Beijing, China, 1998; p 188. (5) Patrini, R.; Marchionna, M. Process for the Production of Ethers from Alcohols. U.K. Patent Application GB 2.323.844 A, November 7, 1998. (6) Tejero, J.; Cunill, F.; Iborra, M.; Izquierdo, J. F.; Fite´, C. Dehydration of 1-Pentanol to Di-n-pentyl Ether over Ion-Exchange Resin Catalysts. J. Mol. Catal. A: Chem. 2002, 183, 541. (7) Collin, J. R.; Ramprasad, D. Methods, Systems and Catalysts for the Hydration of Olefins. European Patent Application EP 1 479 665, 2004. (8) Bringue´, R.; Iborra, M.; Tejero, J.; Izquierdo, J. F.; Cruz, V. J.; Cunill, F.; Fite´, C. Thermally Stable Ion-Exchange Resins as Catalysts for the Liquid-Phase Dehydration of 1-Pentanol to Di-n-pentyl Ether (DNPE). J. Catal. 2006, 244, 33. (9) Fisher, S.; Kunin, R. Routine Exchange Capacity Determinations of Ion Exchange Resins. Anal. Chem., 1955, 27, 1191. (10) Bringue´, R. M.S. Thesis, University of Barcelona, Spain, 2004. (11) Delion, A.; Torck, B.; Hellin, M. Equilibrium Constant for the Liquid-Phase Hydration of Isobutylene over Ion-Exchange Resin. Ind. Eng. Chem. Process. Des. DeV. 1986, 25, 889. (12) Wittig, R.; Lohmann, J.; Vapor-Liquid Equilibria by UNIFAC Group Contribution. 6. Revision and Extension. Ind. Eng. Chem. Res. 2003, 42, 183. (13) Smith, J. M.; Van Ness, H. C. Introduction to Chemical Engineering Thermodynamics, 4th ed.; McGraw-Hill: Singapore, 1987; Chapter 15. (14) Reid, R. C.; Prausnitz, J. M.; Poiling, B. E. The Properties of Gases and Liquids, 4th ed; McGraw-Hill: New York, 1987; Chapter 5. (15) Chase, M. W., Jr. NIST-JANAF Thermochemical Tables, 4th Ed. J. Phys. Chem. Ref. Data, Monogr 9 1998, 1. (16) Mosselman, C.; Dekker, H. Enthalpies of Formation of n-Alkan1-ols. J. Chem. Soc., Faraday Trans. 1 1975, 417. (17) Murrin, J. W.; Goldenhagen, S. Determination of the C-O Bond Energy from the Heats of Combustion of Four Aliphatic Ethers; NAVORD Report No. 5491; U.S. Naval Powder Factory Research and Development Department: MD, 1957; p 1.
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ReceiVed for reView December 22, 2006 ReVised manuscript receiVed July 20, 2007 Accepted July 23, 2007 IE0616646
