Tert-amyl methyl ether (TAME). Thermodynamic analysis of reaction equilibria in the liquid phase

Article abstract of DOI:10.1021/je990191p

Ethers obtained from C₅ olefin streams provide a good mix of octane-enhancing and CO-reducing qualities. They are considered as a replacement for the more common isobutylene-derived fuel additives, e.g., MTBE. Data on the thermodynamic reaction equilibria for the synthesis of TAME, which is produced from the etherification reaction of 2-methyl-1-butylene and 2-methyl-2-butylene with methanol, were computed by studying TAME decomposition in a recirculating batch reactor at different temperatures with activity coefficients obtained using the UNIFAC method. A comparison of experimental data and theoretical predictions showed the possibility of erroneous values reported in the literature for the standard Gibbs free energy of TAME formation. Expressions for the thermodynamic equilibrium constants as a function of temperature were developed. A corrected value for the Gibbs free energy of formation for TAME was also provided.

Full text of DOI:10.1021/je990191p

J . Chem. Eng. Data 2000, 45, 319-323
319
ter t-Am yl Meth yl Eth er (TAME). Th er m od yn a m ic An a lysis of
Rea ction Equ ilibr ia in th e Liqu id P h a se
F a isa l H. Syed ,† Ch r istin e Egleston ,‡ a n d Ra vin d r a Da tta *,†
Department of Chemical Engineering, Worcester Polytechnic Institute, Worcester, Massachusetts 01609-2180,
and Ifm efector, 805 Springdale Drive, Exton, Pennsylvania 19341
Ethers derived from C5 olefin streams provide a good mix of octane-enhancing and carbon monoxide-
reducing qualities and are, therefore, being considered as a replacement for the more common isobutylene-
derived fuel additives such as methyl tert-butyl ether (MTBE). This paper provides data on the
thermodynamic reaction equilibria for the synthesis of tert-amyl methyl ether (TAME), which is produced
from the etherification reaction of 2-methyl-1-butene and 2-methyl-2-butene with methanol. Thermo-
chemical parameters for the reaction mixture are experimentally verified, and expressions for the
thermodynamic equilibrium constants as a function of temperature are developed. A corrected value for
the Gibbs free energy of formation for TAME is provided.
In tr od u ction
Ether-based oxygenates have seen a tremendous growth
in the 1990s. Compared to alcohols, ethers have a lower
Reid vapor pressure (RVP), while providing the benefits
of increasing octane levels and reducing carbon monoxide
emissions. Methyl tert-butyl ether (MTBE), produced from
the addition reaction of methanol and isobutylene, has been
the oxygenate of choice, mainly due to the economics and
availability of the two feedstocks (Brockwell et al., 1991).
However, over the past 5 years there has been a growing
concern about contamination of groundwater by MTBE,
F igu r e 1. Simultaneous etherification and isomerization reac-
tions involved in TAME synthesis.
due to leaking storage tanks and pipelines. California,
which represents 25% of MTBE annual global demand, has
already passed legislation for its phaseout by 2002 (Morse,
1999). The major factor affecting the impact of MTBE on
drinking water stems from its high water solubility.
Therefore, refiners are seeking alternative additives with
lower water solubilities. Tert-amyl methyl ether (TAME),
produced from methanol and C5 tertiary olefins, is a
promising candidate, although other ethers such as tert-
amyl ethyl ether (TAEE) and tert-hexyl ethyl ethers
(THEE1, THEE2, and THEE3) are also being actively
investigated (Kitchaiya and Datta, 1995; Zhang and Datta,
1995). Not only are these ethers less water soluble, but they
also have lower vapor pressures, which further aid in the
reduction of air pollution for motor vehicles.
1994; Zhang and Datta, 1995a; Kitchaiya and Datta, 1995;
Zhang and Datta, 1995b). Although a number of thermo-
dynamic studies have also been reported on TAME (Ran-
driamahefa et al., 1988; Safronov et al., 1989; Rihko et al.,
1994; Hwang and Wu, 1994; Oost et al., 1995; Piccoli and
Lovisi, 1995), none of these have been able to provide
satisfactory agreement between experiments and theory.
Therefore, the aim of this paper is to provide a thermody-
namic analysis of the reaction network involved in TAME
synthesis and to investigate the cause(s) of the deviation
between experiments and thermodynamic analysis, as well
as some additional equilibrium data.
These tertiary ethers are industrially synthesized in the
liquid phase over acid-functionalized ion-exchange resin
catalysts. The reactions are thermodynamically limited,
with equilibrium conversion declining at higher tempera-
tures and as higher molecular weight olefin feeds are
utilized (Datta et al., 1997). Several thermodynamic studies
pertaining to MTBE, ETBE, TAEE, and ethanol-derived
C6 ethers are available in the literature (Colombo et al.,
1983; Safronov et al., 1989; Rehfinger and Hoffmann, 1990;
Izquierdo et al., 1992; Vila et al., 1993; J ensen and Datta,
Th eor etica l An a lysis
TAME can be produced from either tertiary pentene,
namely 2-methyl-1-butene (2M1B), or 2-methyl-2-butene
(2M2B) and methanol over a protonated ion-exchange resin
catalyst like Amberlyst-15. In addition to the two parallel
etherification reactions, 2M1B also isomerizes to form
2M2B, and vice-versa, on the catalyst. The resulting
reaction network is shown in Figure 1. The only significant
side reaction observed is the production of tert-amyl alcohol
if any water is present in the feed. The equilibrium analysis
of Kitchaiya and Datta (1995) and J ensen and Datta (1995)
is summarized below.
* Corresponding author. E-mail address: rdatta@wpi.edu.
† Worcester Polytechnic Institute.
‡ Ifm efector.
10.1021/je990191p CCC: $19.00 © 2000 American Chemical Society
Published on Web 02/03/2000

320 J ournal of Chemical and Engineering Data, Vol. 45, No. 2, 2000
Ta ble 1. Exp er im en ta lly Deter m in ed Mole F r a ction s a n d Id ea l Kxi a n d Non id ea l Ki Equ ilibr iu m Con sta n ts for TAME
Syn th esis a t Va r iou s Tem p er a tu r es
T/K
xMeOH,e
x2M1B,e
x2M2B,e
xTAME,e
Kx1
Kx2
Kx3
K1
K2
K3
298
308
318
328
348
0.159
0.174
0.187
0.242
0.271
0.009
0.012
0.015
0.023
0.027
0.152
0.164
0.174
0.217
0.241
0.680
0.650
0.625
0.519
0.461
451.40
311.61
229.22
95.37
28.23
22.74
19.27
9.89
8.91
9.65
10.56
13.70
15.98
119.80
85.94
65.38
29.39
19.97
7.45
6.23
5.47
3.03
2.23
16.08
13.79
11.96
9.70
62.74
7.04
8.96
For q liquid-phase reactions, the equilibrium constant
resin with ethanol, followed by 1.0 M HNO3, and then with
ethanol again to remove any excess free acid remaining
on the resin beads. After drying for 12 h in an oven under
vacuum at 105 °C, the resin was stored in a desiccator until
time of use. One gram of catalyst was loaded into the
reactor, and the reactor was evacuated using a vacuum
pump before the desired reactant mixture was pumped in.
TAME decomposition was studied by charging the reactor
with TAME (97% purity, obtained from Aldrich). The
reactor was heated to the desired reaction temperature and
the mixture agitated using a built-in stirrer. The reactor
was maintained at 1.03 MPa to ensure liquid-phase opera-
tion.
The reactor was operated in a recirculating mode, where
the charged mixture was pumped continuously via a
positive displacement pump (Gilson Medical, model 305)
through an internal liquid-sampling injector (Valco model
CL4WE) to a Perkin-Elmer AutoSystem gas chromato-
graph equipped with a Supelco capillary column (SPB-1,
0.25 mm i.d., 1.0 µm film thickness, 60 m length), before
being recycled back to the reactor. A flame ionization
detector (FID) operating at 250 °C was utilized. An oven
temperature program starting at 35 °C for 4 min and then
increased to 100 °C at a rate of 20 °C/min, where it was
held for 1 min, was sufficient to provide the necessary
separation. GC samples were taken until the successive
reactor composition remained constant over time, at which
point the temperature was set to the next desired value
and GC sampling resumed.
Ki for the ith reaction, ∑n νijAj, can be written as
j)1
n
n
n
∆G°
iT
Ki ) exp -
)
ajνeij ) ( γjνeij)( xνjeij) ) KγiKxi
∏
∏ ∏
j)1 j)1
(
)
RT
j)1
(1)
where aje ) γjexje is the activity of species j at equilibrium.
The activity coefficient γje may be calculated from the
equilibrium mole fraction xje using an appropriate model.
Columbo et al. (1983) and Izquierdo et al. (1992) have
demonstrated the adequacy of the UNIFAC group contri-
bution method for the highly nonideal MTBE reaction
mixture. Furthermore, J ensen and Datta (1995), Kitchaiya
and Datta (1995), and Zhang and Datta (1995), respec-
tively, have shown the accuracy of the UNIFAC method
for the ETBE, TAEE, and THEEx reaction systems.
Therefore, this work assumes that the UNIFAC method is
applicable to the TAME system.
The standard Gibbs free energy change and the standard
enthalpy change for reaction i can be written as a function
of temperature by relating them to the molar heat capacity
C° of species j, given as a function of temperature by the
Pj
usual form
C° ) aj + bjT + cjT2 + djT3
(2)
Pj
Thus
∆bi
∆ci
∆di
∆H° ) IiH + ∆aiT +
T2 +
T3 +
T4 (3)
iT
2
3
4
Resu lts a n d Discu ssion
Equilibrium constants were determined for the three
individual reactions using eq 1. Equilibrium mole fractions
xje determined from GC analysis were used to calculate the
species activity coefficients at equilibrium γje using the
UNIFAC method. Table 1 lists the experimental results
for the TAME decomposition reaction at different temper-
atures, along with ideal and nonideal equilibrium constants
which when used in the van’t Hoff equation provides the
standard Gibbs free energy change, or equivalently the
equilibrium constant, for reaction i
∆G°
IiH
∆ai
∆b
2R
iT
-
) ln Ki ) IiK
-
+
ln T +
iT +
RT
RT
R
∆ci
6R
∆di
T2 +
T3 (4)
K
xi and Ki, respectively. Although Kxi and Ki follow similar
12R
trends, it is clear that the equilibrium reaction mixture
significantly deviates from ideality. As the temperature
increases, the equilibrium composition of TAME for both
etherification reactions decreases, which is expected of
exothermic reactions. The composition of 2M2B relative to
2M1B also decreases with increasing temperature. Fur-
thermore, only two of the three reactions are independent,
and their reaction equilibrium constants are, thus, related
by
as a function of temperature. IiH and IiK are the respective
integration constants, evaluated from thermochemical data
at the standard temperature T° as
∆bi
∆ci
∆di
IiH ) ∆Hi°T° - ∆aiT° -
(T°)2 -
(T°)3 -
(T°)4
2
3
4
(5)
and
IiK ) ln KiT°
ln K1 ) ln K2 + ln K3
(7)
IiH
∆ai
∆b
∆ci
+
-
ln T° -
iT° -
(T°)2 -
RT°
R
2R
6R
∆di
Theoretical reaction equilibrium constants at each tem-
perature were calculated using eq 4 and the necessary
liquid-phase thermochemical data. The thermochemical
properties for the two iso-olefins were obtained from
Kitchaiya and Datta (1995), and those of methanol, from
Zhang and Datta (1995a). The standard enthalpy and
standard Gibbs free energy for TAME at T° ) 298.15 K
were obtained from the Thermodynamic Research Center
(T°)3 (6)
12R
Exp er im en ta l Section
The experimental study was carried out in a 300 mL
Parr batch reactor. Amberlyst 15 ion-exchange resin cata-
lyst (obtained from Sigma) was treated by washing the

J ournal of Chemical and Engineering Data, Vol. 45, No. 2, 2000 321
Ta ble 2. Liqu id -P h a se Th er m och em ica l Da ta Rep or ted
in th e Liter a tu r e for th e TAME Rea ction System
heat capacity coefficients/
kJ ‚mol-1‚K-1
C° ) aj + bjT + cjT2 + djT3
Pj
component
∆G°
/
∆H° /
fjT°
fjT°
j
10aj
104bj
107cj 1010dj kJ ‚mol-1 kJ ‚mol-1
methanola 0.077
1.62
2.06
5.08
7.51 -0.882
2.87 -166.64 -238.91
2M1Bb
2M2Bb
TAME
1.27 -0.609
1.69
65.00
58.40
-62.47
-69.91
1.33 -1.48
1.73c
2.29c -6.00c 20.0c
-108.20d -334.90d
Zhang and Datta (1995a). Kitchaiya and Datta (1995). c Es-
a
b
d
timated by the Missenard method (Reid et al., 1987). TRC (1992).
F igu r e 3. Experimental (9, this study; 0, Rihko et al., 1994) and
theoretical reaction equilibrium constant as a function of temper-
ature for TAME synthesized from methanol and 2M2B using
literature (dotted line) and regressed (solid line) ∆H° and ∆G°
iT°
iT°
values from Table 3.
F igu r e 2. Experimental (9, this study; 0, Rihko et al., 1994) and
theoretical reaction equilibrium constant as a function of temper-
ature for TAME synthesized from methanol and 2M1B using
literature (dotted line) and regressed (solid line) ∆H° and
iT°
∆Gi°T° values from Table 3.
(TRC), while the heat capacity information was estimated
using the Missenard method described by Reid et al. (1987).
This method is a group contribution method providing
estimates with <5% error. The thus obtained thermo-
chemical data for all components are summarized in Table
2.
F igu r e 4. Experimental (9, this study; 0, Rihko et al., 1994) and
theoretical reaction equilibrium constant as a function of temper-
ature for the isomerization of 2M1B to 2M2B using literature
(dotted line) and regressed (solid line) ∆H° and ∆G° values
iT°
iT°
from Table 3.
The experimental and theoretical reaction equilibrium
constants for the TAME reaction system as a function of
temperature are plotted in Figures 2-4, along with the
data of Rihko et al. (1994), who studied TAME synthesis
as well as decomposition using different feed mixtures. The
experimental results of this study are similar to the results
obtained by Rihko et al. (1994). However, experimental
results for both of the etherification reactions are ap-
proximately an order of magnitude higher than the results
predicted using the literature thermochemical data listed
in Table 2. However, the equilibrium constant for the
isomerization reaction agrees well with predictions. The
disagreement between theory and experiments raises ques-
tions about the accuracy of the thermochemical data listed
in Table 2. It is known that slight changes in the enthalpy
and the Gibbs free energy of formation values can cause
significant changes in the equilibrium constant. Rihko et
al. (1994) found that an increase of 1% in the enthalpy of
formation of TAME caused the value of the equilibrium
constant to increase by as much as 3-fold. Furthermore,
Kitchaiya and Datta (1995) found that the equilibrium
constant is fairly insensitive to errors in the heat capacity
data, but a 1 kJ /mol difference in the standard Gibbs free
energy can result in a 50% difference in the thermodynamic
equilibrium constant.
To obtain more accurate values of the standard enthalpy
change ∆H° and the standard Gibbs free energy change
iT°
∆G° for reaction i at T° ) 298.15 K, eqs 5 and 6 for the
iT°
constants IiH and IiK are substituted into eq 4 to obtain
φi
∆H°
iT° 1
1
T
ln Ki - ) KiT°
+
-
(8)
(
)
R
R
T°
where
∆bi
2
∆ci
∆di
φ ≡ ∆a T° +
(T°)2 +
(T°)3 +
(T°)4
×
(
)
i
i
3
4
∆b
1
T
1
T
T°
-
+ ∆a ln
+
i(T - T°) +
∆di
(
)
i
T°
2
∆ci
(T2 - (T°)2) +
(T3 - (T°)3) (9)
6
12
A plot of the LHS of eq 8, using experimentally determined
values for Ki, versus (1/T° - 1/T) results in a straight line,
as shown in Figure 5, with a slope of ∆H° /R and an
intercept equal to KiT°. The resulting standard enthalpy
∆Hi°T° and standard Gibbs free energy ∆Gi°T° of reaction for
iT°

322 J ournal of Chemical and Engineering Data, Vol. 45, No. 2, 2000
Ta ble 3. Sta n d a r d En th a lp y a n d Gibbs F r ee En er gy of Rea ction i a t T° ) 298.15 K
regressed from eq 8
determined from literature values
∆G° /kJ ‚mol-1 ∆H° /kJ ‚mol-1
reaction i
∆Gi°T°/kJ ‚mol-1
∆H° /kJ ‚mol-1
iT°
iT°
iT°
1
2
3
-12.00
-5.50
-6.76
-32.96
-23.50
-9.46
-6.56
-33.52
-26.08
-7.44
0.04
-6.60
Ta ble 4. Ca lcu la ted Sta n d a r d Gibbs F r ee En er gy a n d Sta n d a r d En th a lp y of F or m a tion for TAME a t T° ) 298.15 K
regressed from eq 8
lit. values (TRC)
exp at 298.15 K
reaction i
∆Gf°iT°/kJ ‚mol-1
∆Hf°iT°/kJ ‚mol-1
∆Gf°iT°/kJ ‚mol-1
-108.2
∆Hf°iT°/kJ ‚mol-1
-334.9
∆Gf°iT°/kJ ‚mol-1
1
2
-113.9
-113.7
-334.32
-332.34
-113.5
-113.2
comparison to the literature values, the ∆H° value is
fiT°
similar, but the ∆G° value differs significantly. Thus, it
fiT°
appears that a more accurate value for ∆G°fiT° for TAME is
-113.8 kJ /mol, rather than the -108 kJ /mol value reported
in the literature (TRC, 1992).
Con clu sion s
The thermodynamic equilibrium constants for the si-
multaneous etherification and isomerization reactions
involved in the liquid-phase synthesis of TAME are ana-
lyzed. Experimental equilibrium constants were calculated
by studying TAME decomposition in a recirculating batch
reactor at different temperatures with activity coefficients
determined using the UNIFAC method. Comparison of
experimental data and theoretical predictions revealed the
possibility of erroneous values reported in the literature
for the standard Gibbs free energy of formation for TAME.
Experiments conducted at T° ) 298.15 K, coupled with
the linear regression of data at other temperatures, re-
F igu r e 5. Estimation of standard enthalpy change and standard
Gibbs free energy change of reaction at T° ) 298.15 K using eq 8
for reaction 1 (b), reaction 2 (9), and reaction 3 (2).
vealed that the actual value of ∆G° for TAME is closer
fiT°
to -113.8 kJ /mol, rather than to -108 kJ /mol reported in
the literature (TRC) and used by all studies on TAME. The
use of these values in the thermodynamic analysis results
in a much better agreement between experiments and
theory, allowing for the determination of more accurate
expressions for the equilibrium constants for the TAME
reaction network.
all three reactions are shown in Table 3, along with values
obtained using the thermochemical data in Table 2. A much
better fit between theory and experiments is obtained if
the ∆H° and ∆G° values thus obtained from eq 8 are
used in eqs 4-6, as shown in Figures 2-4 (corrected). The
resulting equilibrium constants for the TAME reaction
network can thus be written in their final form:
iT°
iT°
Nom en cla tu r e
5018.61
ln K1 ) -39.065 +
+ 4.6866 ln T +
T
aj, bj, cj, dj ) coefficients of molar heat capacity expres-
sion
0.00773T - (2.635 × 10-5)T2 + (1.547 × 10-8)T3 (10)
aj ) activity of species j, ≡γjxj
3918.02
C° ) molar heat capacity of species j, J /mol‚K
IiHPj ) constant of integration
ln K2 ) -34.798 +
+ 3.9168 ln T +
T
0.01293T - (3.121 × 10-5)T2 + (1.805 × 10-8)T3 (11)
IiK ) constant of integration
Ki ) thermodynamic equilibrium constant of reaction i
R ) gas constant, 8.3143 J /mol‚K
T ) temperature, K
1100.69
ln K3 ) -4.159 +
+ 0.7698 ln T - 0.00521T +
T
(4.865 × 10-6)T2 - (2.580 × 10-9)T3 (12)
T° ) standard temperature, 298.15 K
xj ) mole fraction of species j
On average, eqs 10-12 provide values with errors of 9.6%,
9.2%, and 1.8%, respectively, within the studied temper-
ature range.
Greek Letters
γj ) activity coefficient of species j
∆ai ≡ ∑n νijaj
Since studies by Zhang and Datta (1995) and Kitchaiya
and Datta (1995) have demonstrated the accuracy of the
thermochemical parameters for methanol and the C5 iso-
olefins, respectively, those for TAME are likely the cause
of the discrepancy between theory and experiments. Thus,
the reaction was conducted at T° ) 298.15 K to directly
determine the standard Gibbs free energy of formation of
TAME from the equilibrium composition. This is compared
to the value obtained from the regression of eq 8 in Table
4, and the two are seen to be in excellent agreement. In
j)1
∆bi ≡ ∑n νijbj
j)1
∆ci ≡ ∑n νijcj
j)1
∆d ≡ ∑nj)1νijdj
∆C° ) liquid-phase heat capacity difference between
iP
products and reactants of reaction i, ∆C° ≡ ∑nj)1νijC° ,
iP
Pj
J /(mol‚K)
∆G° ) standard Gibbs energy of formation of species j
fjT
in the liquid phase at temperature T, kJ /mol

J ournal of Chemical and Engineering Data, Vol. 45, No. 2, 2000 323
Daubert, T. E.; Danner, R. P. Physical and Thermodynamic Properties
∆G° ) standard Gibbs energy change for reaction i in
iT
of Pure Chemicals, Data Compilation; Hemisphere Publishing
Corp.: Washington, DC, 1992.
the liquid phase at temperature T, ∆G° ≡ ∑nj)1νij∆G° ,
iT
fjT
kJ /mol
Hwang, W. S.; Wu, J . C. Kinetics and Thermodynamics of Synthesis
of Tertiary-Amyl Methyl Ether Catalyzed by Ion-Exchange Resin.
J . Chin. Chem. Soc. 1994, 41, 181-186.
∆H° ) standard enthalpy of formation of species j in
fjT
the liquid phase at temperature T, kJ /mol
Ignatius, J .; J a¨rvelin, H.; Lindqvist, P. Use TAME and Heavier Ethers
to Improve Gasoline Properties. Hydrocarbon Process. 1995, 74 (2),
51-53.
∆H° ) standard enthalpy change for reaction i in the
iT
liquid phase at temperature T, ∆H° ≡ ∑nj)1νij∆H°
iT
Izquierdo, J . F.; Cunill, F.; Vila, M.; Tejero, J .; Iborra, M. Equilibrium
Constants for Methyl tert-Butyl Ether Liquid Phase Synthesis. J .
Chem. Eng. Data 1992, 37, 339-343.
n
∑
j)1νij∆H° , kJ /mol
νij ) stoichifoiTmetric coefficient of species j in reaction i
J ensen, K. L.; Datta, R. Ethers from Ethanol. 1. Equilibrium Ther-
modynamic Analysis of the Liquid-Phase Ethyl tert-Butyl Ether
Reaction. Ind. Eng. Chem. Res. 1995, 34, 392-399.
Subscripts
Kitchaiya, P.; Datta, R. Ethers from Ethanol. 2. Reaction Equilibria
of Simultaneous tert-Amyl Ethyl Ether Synthesis and Isoamylene
Isomerization. Ind. Eng. Chem. Res. 1995, 34, 1092-1101.
Morse, P. M. Producers Brace for MTBE Phaseout. Chem. Eng. News
1999.
Piccoli, R. L.; Lovisi, H. Kinetics and Thermodynamic Study of the
Liquid-Phase Etherification of Isoamylenes with Methanol. Ind.
Eng. Chem. Res. 1995, 34, 510-515.
e ) at equilibrium
f ) of formation
i ) of reaction i
j ) of species j
T ) at temperature T
T° ) at standard temperature
Rehfinger, A.; Hoffmann, U. Kinetics of Methyl Tertiary Butyl Ether
Liquid Phase Synthesis Catalyzed by Ion Exchange Resin-I. Intrin-
sic Rate Expression in Liquid-Phase Activities. Chem. Eng. Sci.
1990, 45, 1605-1617.
Abbreviations
2M1B ) 2-methyl-1-butene
2M2B ) 2-methyl-2-butene
RVP ) Reid vapor pressure
ETBE ) ethyl tert-butyl ether
MTBE ) methyl tert-butyl ether
TAEE ) tert-amyl ethyl ether
TAME ) tert-amyl methyl ether
THEE ) tert-hexyl ethyl ethers
Reid, R. C.; Prausnitz, J . M.; Poling, B. E. The Properties of Gases and
Liquids, 4th ed.; McGraw-Hill: New York, 1987.
Rihko, L.; Linnekoski, J . A.; Krause, O. I. Reaction Equilibria in the
Synthesis of 2-Methoxy-2-methylbutane and 2-Ethoxy-2-methylbu-
tane in the Liquid Phase. J . Chem. Eng. Data 1994, 39, 700-704.
Safronov, V. V.; Sharanov, K. G.; Rozhnov, A. M.; Alenin, V. J .; Sidorov,
S. A. Thermodynamics of Sythesis of tert-Amyl Methyl Ether. Zh.
Prikl. Khim. (Leningrad) 1989, 62, 824-828.
TRC Thermodynamic Tables, Hydrocarbons and Nonhydrocarbons;
Thermodynamic Research Center: Texas A&M University, College
Station, TX, 1992; Vol. 5.
Vila, M.; Cunill, F.; Izquierdo, J . F.; Tejero, J .; Iborra, M. Equilibrium
Constants For Ethyl tert-Butyl Ether Liquid-Phase Synthesis.
Chem. Eng. Commun. 1993, 124, 223-232.
Zhang, T.; Datta, R. Integral Analysis of Methyl tert-Butyl Ether
Synthesis Kinetics. Ind. Eng. Chem. Res. 1995a , 34, 730-740.
Zhang, T.; Datta, R. Ethers from Ethanol. 3. Equilibrium Conversion
and Selectivity Limitations in the Liquid-Phase Synthesis of Two
tert-Hexyl Ethyl Ethers. Ind. Eng. Chem. Res. 1995b, 34, 2237-
2246.
Liter a tu r e Cited
Brockwell, H. L.; Sarathy, P. R.; Trotta, R. Synthesize Ethers.
Hydrocarbon Process. 1991, 70 (9), 133-141.
Colombo, F.; Cori, L.; Dallora, L.; Delogu, P. Equilibrium Constant
for the Methyl tert-Butyl Ether Liquid-Phase Synthesis by use of
UNIFAC. Ind. Eng. Chem. Fundam. 1983, 22, 219-223.
CRC Handbook of Chemistry and Physics, 73rd ed.; Lide, D., Ed.;
CRC: Boca Raton, FL, 1992.
Datta, R.; J ensen, K.; Kitchaiya, P.; Zhang, T. Thermodynamic
Transition-State Theory and Extrathermodynamic Correlations for
the Liquid-Phase Kinetics of Ethanol Derived Ethers. In Dynamics
of Surfaces and Reaction Kinetics in Heterogeneous Catalysis;
Froment, G. F., Waugh, K. C., Eds.; Elsevier: Amsterdam, 1997;
pp 559-564.
Received for review J uly 12, 1999. Accepted December 8, 1999.
J E990191P