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Ind. Eng. Chem. Res. 2005, 44, 970-976
Measurement and Development of Solubility Correlations for Tritolylamine in Twelve Organic Solvents Touraj Manifar and Sohrab Rohani* Department of Chemical and Biochemical Engineering, The University of Western Ontario, London, Ontario N6A 5B9, Canada
Marko Saban Xerox Research Centre of Canada, 2660 Speakman Drive, Mississauga, Ontario L5K 2L1, Canada
The solubility of tritolylamine (TTA) in 12 solvents (hexane, heptane, n-octane, decane, hexadecane, toluene, benzene, 2-propanol, propanol, 2-butanone, methanol, and ethanol) was measured gravimetrically and compared with the predictions obtained from the ideal solution and UNIQUAC equations. To use the UNIQUAC equation, the specific heat and heat of fusion of TTA were measured by differential scanning calorimetry (DSC). DSC results also indicated that TTA does not have any enantiotropically related polymorphs. It is shown that the UNIQUAC predictions are close to the measured solubility values. In addition, measurement of the solubility of TTA in hexane was performed by an on-line density meter. The results compared closely with those obtained with the gravimetric method. The solvents were chosen so that their polarity indexes cover a wide range. Introduction Arylamine molecules in general and tritolylamine (TTA; for its molecular structure see Figure 1) in particular have a wide range of industrial applications such as in the xerographic, polymer,1 and pharmaceutical2,3 industries. The arylamine molecules form a stable aminium radical cation.4 In the xerographic industry, arylamine molecules are used as hole-transport materials because of their high hole drift mobilities.5-7 Typically, in organic lightemitting devices (OLEDs) and in organic photoreceptors, the arylamine molecules are used in the form of a thin solid film. Constantly increasing demand for the improvement of these devices (durability, thermodynamic stability, and higher efficiency) necessitates stricter control over the properties and characteristics of the hole-transport materials that make up such devices. Although the use of arylamine molecules in optoelectronic devices (OLEDs8-10 and organic photoreceptors11-13) has been well documented, no data on the solubilities of their constituent arylamine molecules in different solvents have been published. The solubility information is of special importance in the study of nucleation and growth kinetics of these molecules. For example, in the xerographic industry, a minor impurity in these molecules could have a serious effect on the operation of the optoelectronic devices. In OLEDs, the tendency of the hole-transport materials to crystallize on aging is thought to be one of the main causes of degradation.14-16 To prevent crystallization of these molecules, precise knowledge of the kinetics and solubility of the hole-transport materials is essential. In this study, we investigated the solubility of TTA in different solvents. In a series of subsequent papers, we will report the solubilities of other arylamine mol* To whom correspondence should be addressed: E-mail:
[email protected].
Figure 1. Molecular structure of tritolylamine molecule.
ecules that we have synthesized in our laboratory, as well as on the crystallization kinetics of these molecules. For the measurement of solubility, a gravimetric method was used. In addition, an on-line density meter was used to monitor the solubility of TTA in hexane. The two adjustable parameters of the UNIQUAC model were estimated for the prediction of the solubilities of TTA. Theory It is desirable to estimate the solubility of a solid in a solution completely from the properties of the pure components to minimize the amount of experiments required. If we assume that there is no appreciable solubility of the liquid solvent in the solid phase, then the equation showing the equilibrium between two phases becomes
f2(pure solid) ) f2(solute in liquid solution) (1) where f represents the fugacity and the subscript 2 refers to the solute. By assuming that the solution is nonideal and considering the intermolecular and intramolecular forces, the solubility of a solute in solution can be written as
x2 )
f2(pure solid) γ2f02
10.1021/ie049352v CCC: $30.25 © 2005 American Chemical Society Published on Web 01/25/2005
(2)
Ind. Eng. Chem. Res., Vol. 44, No. 4, 2005 971
where x2 is the molar solubility of the solute, γ2 is the activity coefficient, and f 02 is the standard-state fugacity to which γ2 refers. The standard-state fugacity is taken as the fugacity of the pure subcooled liquid at the temperature of the solution and at some specific pressure.17 Equation 2 shows that intermolecular forces affect both the activity coefficients and the fugacity. Therefore, eq 2 can be written in terms of the ratio of the fugacities of the solid and pure subcooled liquid, assuming that the solubility of the liquid in the solid phase is negligible and defining the standard state as the pure subcooled liquid at temperature T and saturation pressure
x2γ2 )
f S2
a
(3)
f L2
The relation between the solid and liquid fugacities at equilibrium is given by17
ln
f L2 f
S 2
)
(
)
(
)
( )
∆Hfus Ttp ∆cp Ttp ∆cp Ttp -1 -1 + ln RTtp T R T R T
(4) L 2
where f is the fugacity of pure subcooled liquid and f S2 is the pure solid fugacity. Ttp is the temperature at the triple point of the solute, which can be taken as its melting point. This assumption introduces negligible error.17 ∆Hfus is the heat of fusion, and ∆cp is the difference in heat capacities of the solute in the liquid state and in the solid state at temperature T. This equation is true for all cases regardless of the ideality or nonideality of the solution. To solve this equation, the thermal properties of the pure solid are needed. Several assumptions are made for eq 4. First, ∆cp is assumed to be constant over the temperature range T-Ttp. Second, the effect of pressure on the properties of solid and subcooled liquid is assumed to be negligible. This is true unless the pressure is high. Finally, it is assumed that no solid-solid phase transition occurs. The ratio of two fugacities can be calculated using eq 4. Once γ2 is calculated, the molar solubility can easily be estimated by eq 3. For the ideal case, γ2 is assumed to be 1. For nonideal solutions, γ2 has to be determined. There are many different methods such as NRTL, Van Laar, Wilson,18,19 universal quasi-chemical theory (UNIQUAC),20,41 and UNIFAC that can be used for the calculation of the activity coefficient of a solute in a solvent. In this study, the UNIQUAC method was chosen. The UNIQUAC equation has two parts: combinatorial and residual.21,22 The combinatorial part describes the contribution due to the molecular size, the dominant entropic contribution. The residual part accounts for the contribution due to molecular interactions and intermolecular forces.20 This can be presented as
ln γ2 ) ln γC2 + ln γR2
Table 1. List of R and Q Values Used in This Study17,23
(5)
where γC2 and γR2 represent the combinatorial part and
molecule
Ra (cm3/mol)
Qb (cm2/mol)
hexane heptane n-octane decane hexadecane toluene benzene 2-propanol 1-propanol 2-butanone methanol ethanol TTA
4.499 5.174 5.8484 7.197 11.013 3.92 3.19 2.779 2.779 3.24 1.432 2.105 12.041
3.852 4.396 4.936 6.016 9.044 2.968 2.4 2.508 2.512 2.876 1.432 1.97 9.296
van der Waals volume. b van der Waals area.
the residual part for the solute, respectively. The UNIQUAC equation for a two-component system is
(
)
Φ2 z θ2 r2 + q2 ln + Φ1 l2 - l1 x2 2 Φ2 r1 τ21 τ12 q′2 ln(θ′2 + θ′1τ12) + θ′1q′2 (6) θ′2 + θ′1τ12 θ′1 + θ′2τ21
ln γ2 ) ln
(
)
where
Φ2 )
r2x2 r1x1 + r2x2
(7)
θ2 )
q2x2 q1x1 + q2x2
(8)
θ′2 )
q′2x2 q′1x1 + q′2x2
(9)
( )
τ12 ) exp -
a12 T
z l1 ) (r1 - q1) - (r1 - 1) 2
( )
τ21 ) exp -
a21 T
(10)
z l2 ) (r2 - q2) - (r2 - 1) 2 (11)
where a12 and a21 are adjustable parameters of the UNIQUAC equation and r, q, and q′ are pure-component constants that depend on the molecular size and can be calculated from van der Waals volume and area. q′ for methanol, ethanol, and C3-alcohols is equal to 0.96, 0.92, and 0.89, respectively. For fluids other than lowermolecular-weight alcohols or water, q′ ) q. x2 is the mole fraction of solute in the mixture, θ2 is the area fraction of solute, and Φ2 is the volume fraction. From the data presented in the literature17,23 (listed in Table 1), the dimensionless parameters r and q are derived as follows
ri ) qi )
Qi 15.17 Ri
2.5 × 109
(12) (13)
Note that a12 * a21, and these adjustable parameters have units of temperature in kelvin. In some cases, the adjustable parameters are not available. Using experimental thermal properties of pure solids as well as the solubility data, the experimental activity coefficients of
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the solute in different solvents can be calculated. Then, from these activity coefficients, the adjustable parameters of the activity coefficients in the UNIQUAC model can be obtained. Consequently, these parameters can be used for equilibrium solubility calculations and predictions (e.g., vapor-liquid equilibrium) of solutions consisting of different solvents or highly nonideal solutions even at high pressures. Optimization of the Adjustable Parameters of the UNIQUAC Model. The optimization procedure is based on the minimization of the errors between calculated and experimental values of activity coefficients at a number of temperatures. Therefore, we have n
min error ) a12,a21
exp cal 2 (γ2,k - γ2,k ) ∑ k)1
(14)
exp is the experimental activity coefficient of where γ2,k cal is the the solute based on the solubility data and γ2,k calculated activity coefficient of solute at the same temperature. Minimization was carried out by using the function fmincon in Matlab (Mathworks, MA). Errors arise from different sources related to the measurement of solubility and thermal parameters. In solubility measurements, errors associated with weighing were in the range of (0.01 mg. Any error associated with evaporation of solvents was minimized by employing tightly closed screw-cap vials. The main source of error in the solubility measurements was the determination of the solubility temperature. In DSC analysis, although the crucible of the sample was not thermally isolated, use of a reference crucible minimized the error. The calibration of the DSC instrument (Mettler Toledo, Greifensee, Switzerland) was performed using pure indium. The acceptable deviation in the heat of fusion was 27.85-29.05 J/g, and the deviation in melting temperature was 156.3-159.9 °C. Precise weighing of the samples ensured minimization of the errors associated with the differential scanning calorimetry.
Experimental Section Materials. Hexane was purchased from Avocado Research Chemicals Ltd. (Lancaster, PA). All other materials were purchased from Sigma-Aldrich Chemical Co., Inc. (Milwaukee, WI) and used as received. Solvents were all HPLC grade. The tritolylamine (TTA) used in the solubility studies in the present work was synthesized and purified in our laboratory. Synthesis and Purification of Tritolylamine. Tritolylamine [TTA, N,N,N-tris-(4-methylphenyl)amine] was synthesized in a glass reactor under an inert argon atmosphere. The method of copper-ligated synthesis proposed by Goodbrand and Hu24 and modified by our team25 was used. Reaction was monitored by HPLC (Varian Inc., Ontario, Canada) using an R-18 column with acetonitrile/methanol at a flow rate ratio of 1.0 mL/ min/0.2 mL/min. Samples were smaller than 1 mL, so their withdrawal from the reactor did not disturb the process. More details of TTA synthesis and purification can be found in the literature.24,25 The purified TTA material showed 100% HPLC purity. Solubility Measurements. There are several methods for the measurement of solute concentration. These include measurement of the refractive index,26,27 density,28-31 turbidity,32 or spectral characteristics.33,34 In the present work, the gravimetric method was used
in all cases. In addition, an on-line density meter was also used for solubility measurements and comparison of the results with the gravimetric method. The latter method has potential for continuous monitoring and control applications. (i) Gravimetric Method. A number of 5-mL vials with screw caps were weighed and marked. Different amounts of TTA crystals were added to each vial and weighed. Identical amounts of a specific solvent were added to each vial. Vials were immersed in a constanttemperature bath and shaken slowly for 30-45 min at each temperature increment. The temperature was increased to find the saturation temperature by visual observation. At the saturation point, no crystals could be observed in the solution. A focused light (Leica CLS 150) was used for visual monitoring. The resolution of the balance (Mettler Toledo, Greifensee, Switzerland) was (0.01 mg, and that of the temperature sensor was (0.1 °C. (ii) On-line Density Meter. A second experimental method, namely, determination of the solubility using an on-line density meter, was employed. Initially, the density meter had to be calibrated by obtaining the relationship between C (solute concentration), F (solution density), and t (solution temperature). The experiments were carried out in a 1000-mL double-jacketed glass vessel (Bellco, Vineland, NJ) equipped with a stirrer (AC Tech, Minnesota City, MN). The solution temperature was controlled using a water bath system (RTE 220, NESLAB Instruments Inc., Portsmouth, NH). Lab View (National Instruments, Austin, TX) hardware/ software system was used for data acquisition. Temperature and density data were displayed and stored using an on-line density meter (MPDS 2000, Anton Paar, Graz, Austria). To pump a solid-free liquid sample through the tube of the Anton Paar density meter, an inverted cylindrical glass tube with a diameter of 1 in. reduced to 1/4 in. at the top, was placed in the vessel (to create a quiescent zone). The 1/4-in. tube was connected to a peristaltic pump via a flexible tube. Glass wool was placed at the entrance of the latter tube to stop the flow of particles into the density meter. The flow rate of the sample was kept low (approximately 4 mL/min) to ensure withdrawal of a solid-free sample. Solutions with different known concentrations of TTA in hexane (0.26, 1.322, 2.03, 3.35, 4.73, 5.99, 7.44, and 9.45 g per 100 g of hexane) were prepared. The temperature of the solution was increased, and the solution density was measured at each temperature increment and concentration. The data were used to develop a relationship between density, concentration, and temperature (eq 15). Results and Discussion The solubility of tritolylamine was measured in different solvents using the gravimetric method. These solvents covered a wide range of polarities. The solvents employed were hexane, heptane, n-octane, decane, hexadecane, toluene, benzene, 2-propanol, propanol, 2-butanone, methanol, and ethanol. Table 2 lists the polarity indexes of the different solvents used in this study.23,37 The solvents studied were divided into two main categories: polar and nonpolar. The nonpolar solvents included both cyclic and chain hydrocarbons, whereas polar solvents covered alcohols and ketones. Figures 2 and 3 show the solubilities of TTA in nonpolar and polar
Ind. Eng. Chem. Res., Vol. 44, No. 4, 2005 973 Table 3. Solubility Equations for TTA in Different Solventsa
Figure 2. Logarithmic graph of solubility of TTA in nonpolar solvents. C* has the units of g of TTA/100 g of solvent. These solvents include cyclic and chain hydrocarbons. Solid lines show predictions by the UNIQUAC model, and dashed lines show predictions by the ideal law equation.
Figure 3. Logarithmic graph of solubility of TTA in polar solvents. CC* has the units of g of TTA/100 g of solvent. Predictions obtained using the UNIQUAC equation and the ideal law equation are shown in solid and dashed lines, respectively. Table 2. List of Solvents Used in This Study solvent
polarity index
solvent
A
B
D
R2
hexane heptane n-octane decane hexadecane toluene benzene 2-propanol 1-propanol 2-butanone methanol ethanol
0.0242 0.0074 0.0002 0.014 0.0018 0.0227 0.0439 0.0037 0.0024 0.0725 0.0005 0.0013
-0.9823 0.0724 0.4414 -0.7822 0.0977 0.9425 -0.0535 -0.2009 -0.0271 -3.7274 -0.0017 -0.0114
20.967 2.4735 -1.9951 21.002 0.3729 14.305 57.642 3.9727 0.1819 72.6 0.2274 0.5881
0.998 0.999 0.999 1 0.999 0.996 1 0.992 0.993 0.99 0.94 0.999
hexane heptane n-octane decane hexadecane toluene benzene
14.98 15.208 15.36 15.538 15.938 18.346 18.706
2-propanol 1-propanol 2-butanone methanol ethanol
Polar Solvents 3.9 4 4.7 5.1 5.2
23.575 24.557 18.796 29.523 26.421
solvents, respectively. Because TTA has a symmetrical molecular structure, it is expected that its solubility in nonpolar solvents should be high. As is shown in Figure 2, over the temperature range studied, benzene and toluene have the highest solubilities. Such high solubilities facilitate the cooling crystallization process.28,29 The range of validity of the TTA solubility correlation is reported in Table 3. The solubility of TTA decreases with increasing polarity of the solvent. Therefore, methanol and ethanol should have the lowest solubilities (see Figure 3). However, 2-butanone, a polar solvent, shows a high
5-50 20-60 10-50 40-66 23-50 9-50 18.5-45 25-79 24.5-54 22-50 22-45 12-50
a Equations are in the form C ) At2 + Bt + D in which C is in grams of solute per 100 grams of solvent and t is in degrees Celsius.
dissolving power for TTA similar to those of aliphatic hydrocarbons such as hexane. The reason might be the carbonyl group of 2-butanone that has the electron resonance capability, similarly to toluene or benzene. Among the polar solvents that were under investigation, ethanol and 2-propanol have almost the same dissolution capabilities. Increasing the number of carbons in aliphatic hydrocarbons decreases the solubility of TTA (see Figure 2). This is an important result, given that, for the synthesis of arylamine molecules, aliphatic hydrocarbons are used as the reaction medium.24,38-40 Correlation between Density, Concentration, and Temperature. For the on-line density meter measuring method, hexane was selected as the test solvent due to its high dissolving capability and its potential application in the synthesis of arylamine molecules. A series of experiments was conducted to find a correlation between temperature, concentration, and density of the TTA/hexane system. The correlation that best fitted the data is
solubility parameter [(J/cm3)0.5]
Nonpolar Solvents 0.1 0.2 0.4 0.4 2.4 2.7
temperature validity range (°C)
F ) (R1t + R2)(R3C2 + R4C + R5)
(15)
where
R1 ) -0.000078 R2 ) 0.058 R3 ) -0.00016 R4 ) 0.047 R5 ) 11.925 Figure 4 shows a good agreement between the solubility curves obtained from the gravimetric and density meter methods over a narrow temperature range. However, the R2 value is equal to 0.99, which shows a good fit for these data. Although parameters R1 and R3 appear to be statistically insignificant, dropping these two terms from eq 15 resulted in a poor fit with R2 equal to 0.70. Therefore, we retained all of the parameters in eq 15. Table 4 presents the percent relative difference in the measured solubilities obtained by the two different methods.
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Figure 4. Solubility of TTA in hexane based on the on-line density meter and gravimetric methods. Table 4. Comparison of the Solubility (C*) of TTA in Hexane Obtained by the Two Methods solubility (g of TTA/100 g of hexane) temperature (°C)
gravimetric method
density meter
relative difference (%)
3.15 5.21 7.43 10 15 20 25 30 33
4.52 4.92 5.39 6.00 7.37 9.05 11.13 13.67 15.47
4.54 4.93 5.38 5.97 7.29 8.90 10.88 13.29 14.98
-0.35 -0.11 0.15 0.46 1.06 1.67 2.28 2.90 3.27
a
/ / Percent relative difference ) [(Cgravimetry - Cdensity meter)/ × 100.
/ ] Cgravimetry
Figure 5. DSC result for tritolylamine.
Prediction of TTA Solubility Using the Ideal Solution Theory and the UNIQUAC Model. To determine the thermal properties of TTA that are needed for the UNIQUAC parameters, a differential scanning calorimeter (DSC) was used. The thermogram of TTA is shown in Figure 5. Because polymorphism or molecular packaging results in variations in different physical properties such as stability, crystal shape, solubility, dissolution rate, density, melting point, and optical or electrical properties,35,36 it was necessary to determine whether TTA exhibits polymorphic behavior.
The results from DSC analysis confirmed that there is no solid-solid enantiotropic transition over the temperature range studied (see Figure 5). The specific heat of TTA, cp, calculated based on DSC results and presented in Table 5. By substitution of the TTA melting point, 115 °C (388 K), as the triple point of TTA and ∆Hfus ) -79.88 J/g as the heat of fusion and calculating ∆cp in eq 3, the ratio of fugacities can be calculated. Using eq 4 and assuming that γ2 is equal to unity, the ideal solubility can be calculated. As shown in Figure 2, the measured solubilities of TTA in benzene and toluene are higher than the ideal solubilities, whereas the other nonpolar solvents have lower measured solubilities than the ideal predictions. Usually, whenever there is a significant difference in the nature and size of the solute and solvent molecules, we expect γ2 to deviate from unity. In solutions where only dispersion forces are important, γ2 is generally larger than unity, and therefore, solubility is lower than that corresponding to the ideal behavior. All solvents show the same behavior except toluene and benzene, which have activity coefficients almost equal to and less than unity, respectively. Toluene has partially the same structure as TTA; therefore, the solution of TTA in toluene exhibits less deviation from the ideal behavior. According to the data, the activity coefficients of TTA in alkane solutions are in the range of 1-10; activity coefficients for TTA solution in toluene and benzene are in the ranges of 0.85-0.90 and 0.57-0.78, respectively. However, for polar solvents, these values are completely different. The range changes mostly between 10 and 100, but in the case of methanol, it reaches 240. As stated earlier, the only exception in this group is the methyl ethyl ketone (MEK) 2-butanone, which has a activity coefficient in the range of 1-2. Significant deviations also exist between the measured solubility and the solubility predicted with the ideal equation as is shown in Figures 2 and 3. Therefore, to improve the solubility predictions, the UNIQUAC equation is recommended. The adjustable parameters of the UNIQUAC equation were estimated and are listed in Table 6 along with the R2. Using these parameters, solubilities were calculated. Figures 2 and 3 clearly demonstrate that the solubilities of TTA in various solvents predicted with the UNIQUAC model are close to the experimental observations. The predictive capability of the UNIQUAC model in a single solvent can be further utilized to estimate the solubility of TTA in mixed-solvent systems even at high pressure and nonideal conditions. Conclusion The solubility of TTA in seven polar and five nonpolar solvents was measured experimentally. The UNIQUAC predictions of the solubility of TTA in 12 selected solvents were shown to agree with the experimental results. The specific heat and heat of fusion of TTA were measured by DSC. As shown in Figure 2, toluene and benzene had the highest solubilities of TTA. Among the
Table 5. Heat Capacities (cp) of TTA in Both the Liquid and Solid Phases validity range (K) 298-373 393-413 Temperature in K.
cp (J/g‚K)a
state solid liquid
10-9T4
10-6T3
cp ) 3 × -4× + - 0.441T + 38.927 cp ) -1 × 10-7T4 + 0.0002T3 - 0.123T2 + 31.967T - 3106.1 0.002T2
R2 0.98 0.97‡
Ind. Eng. Chem. Res., Vol. 44, No. 4, 2005 975 Table 6. Adjustable Parameters of the UNIQUAC Equation for Mixtures of TTA with Various Solvents solvent
a12 (K)
a21 (K)
R2 a
hexane 204.89 -90.83 0.99 heptane 177.13 -79.4 0.99 n-octane -233.74 441.16 0.99 decane 25.83 24.28 1 hexadecane 23.03 21.24 0.99 toluene -126.21 201.65 0.99 benzene 420.31 -241.71 0.99 isopropyl alcohol 487.04 -70.24 0.99 propanol -23.41 212.87 0.97 2-butanone 35.92 36.10 0.97 methanol 653.86 23.09 0.99 ethanol 470.86 -33.45 0.99
cumulative relative differenceb(%) 0.91 3.87 5.0 8.60 3.3 2.43 5.96 9.1 0.62 2.13 4.5 2.6
a R2, relative difference calculated for the estimated data and experimental solubility data. b Cumulative relative difference (%) / )/C/exp] × 100. ) ∑[(C/exp - Cest
nonpolar solvents studied, hexane had the highest solubility. In the nonpolar aliphatic hydrocarbon solvents studied, increasing the number of carbon atoms decreased the solubility of TTA. Acknowledgment The authors thank the University of Western Ontario, Xerox Research Centre (XRCC) in Mississauga and the Natural Sciences and Engineering Council of Canada (NSERC) for their financial support of this project. In addition, we extend our sincere thanks to Material and Manufacturing Ontario, ORDCF, and CFI/OIT for their support. Symbols a ) adjustable parameter C ) concentration, g of TTA/100 g of solvent cp ) heat capacity, J/mol‚K f ) fugacity, bar Hfus ) enthalpy of fusion, J/mol l ) adjustable parameter Q ) van der Waals area, cm2/mol q ) dimensionless area q′ ) pure constant; equal to q for fluids other than water or alcohol R ) universal gas constant, 8.314 J/mol‚K (eq 4) Ri ) van der Waals volume, cm3/mol r ) dimensionless volume constant T ) temperature, K (eqs 4 and 10) Ttp ) triple-point temperature, K t ) temperature, °C x ) molar solubility, mol of solute/mol of solution z ) coordination number, 10 Greek Letters Φ ) segment fraction θ ) area fraction θ′ ) area fraction γ ) activity coefficient τ ) adjustable parameter F ) density, g/cm3 ∆ ) difference Superscripts/Subscripts C ) combinatorial cal ) calculated est ) estimated exp ) experimental L ) liquid phase
R ) residual S ) solid phase 1 ) solvent component 2 ) solute component
Literature Cited (1) Singer, R. A.; Sadighi, J. P.; Buchwald, S. L. A General Synthesis of End Functionalized Oligoanilines via PaladiumCatalyzed Amination. J. Am. Chem. Soc. 1998, 120, 213-214. (2) Hong, Y.; Senanayake, C. H.; Xiang, T.; Vandenbossche, C. P.; Tanoury, G. J.; Bakale, R. P.; Wald, S. A. Remarkably Selective Palladium-Catalyzed Amination Process: Rapid Assembly of Multiamino Based Structure. Tetrahedron Lett. 1998, 39, 31213124. (3) Negwer, M. Organic Drugs and Their Synonyms; Akademie Verlag GmbH: Berlin, 1994. (4) Louie, J.; Hartwig, J. F. Discrete High Molecular Weight Triarylamine Dendrimers Prepared by Palladium Catalyzed Amination. J. Am. Chem. Soc. 1997, 119, 11695-11696. (5) Brantly, T. B.; Contois, L. E.; Fox, C. J. Photoconductive Elements Containing Substituted Triarylamine Photoconductors. U.S. Patent 3,567,450, 1971. (6) Brantly, T. B.; Contois, L. E.; Fox, C. J. Photoconductive Elements Containing as Photoconductors Triarylamines Substituted by Active Hydrogen Containg Groups. U.S. Patent 3,658,520, 1972. (7) Klupfel, K.-W.; Sus, O.; Behmenburg, H.; Neugebauer, W. Material for Electrophotographic Purposes. U.S. Patent 3,180,730, 1965. (8) D’lorio, M. Molecular Materials for Microelectronics. Can. J. Phys. 2000, 78, 231-241. (9) Kaliowski, J. Electroluminescence in Organics. J. Phys. D: Appl. Phys. 1999, 32, R-179-R-250. (10) Mitschke, U.; Ba¨uerle, P. The Electroluminescence of Organic Materials. J. Mater. Chem. 2000, 10, 1471-1507. (11) Borsenberger, P. M.; Wiess, D. S. Organic Photoreceptors for Xerography; Marcel Dekker Inc.: New York, 1998. (12) Law, K.-Y. Organic Photoconductive Materials: Recent Trends and Developments. Chem. Rev. 1993, 93, 449-486. (13) Loutfy, R. O.; Hor, A. M.; Hsiao, C. K.; Baranyi, G.; Kazmaier, P. Organic Photoconductive Materials. Pure Appl. Chem. 1988, 60, 1047-1054. (14) Dodabalapur, A.; Jordan, R.; Strakelj, M. Blue Emitting Materials and Electroluminescent Devices Containing These Materials. U.S. Patent 5,904,994, 1999. (15) Forest, S. R.; Burrows, P.; Garbuzov, D. Z. Light Emitting Devices Having High Brightness. U.S. Patent 6,125,226, 2000. (16) Hosokawa, C. Highly Efficient Blue Electroluminescence from a Distyrylarylene Emitting Layer With a New Dopant. Appl. Phys. Lett. 1995, 67, 3853-3855. (17) Prausnitz, J. M.; Lichtenthaler, R. N.; Azevedo, E. G. Molecular Thermodynamics of Fluid-Phase Equilibria, 3rd ed.; Prentice Hall International: Upper Saddle River, NJ, 1999; pp 638-670. (18) Scatchard, G.; Wilson, G. M.; Satkiewicz, F. G. VaporLiquid Equilibrium. X. An Apparatus for Static Equilibrium Measurements. J. Am. Chem. Soc. 1964, 86, 125-127. (19) Wilson, G. M. Vapor-Liquid Equilibrium. XI. A New expression for the Excess Free Energy of Mixing. J. Am. Chem. Soc. 1964, 86, 127-130. (20) Abrams, D. S.; Prausnitz, J. M. Statistical Thermodynamics of Liquid Mixture. A New Expression for the Excess Gibbs Energy of Partly and Completely Miscible Systems. AIChE J. 1975, 21, 116-128. (21) Fredenslund, A.; Jones, R. L.; Prausnitz, J. M. Group Contribution Estimation of Activity Coefficients in Nonideal Liquid Mixtures. AIChE J. 1975, 21, 1086-1099. (22) Derr, E. L.; Deal, C. H. Predicted Compositions during Mixed Solvent Evaporation from Resin Solutions Using the Analytical Solution of Groups Methodo. Adv. Chem. Ser. 1973, 124. (23) Yaws, C. L.; Wang, M. A.; Satyro, M. A. Chemical Properties Hand Book: Physical, Thermodynamic, Environmental, Transport, Safety, and Health Related Properties for Organic and Inorganic Chemicals; McGraw-Hill: New York, 1999; pp 340-363. (24) Goodbrand, H. B.; Nan-Xing, H. Ligand Accelerated Catalysis of the Ullmann Condensation: Application to Hole Conducting Triarylamines. J. Org. Chem. 1999, 64, 670-674.
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Received for review July 23, 2004 Revised manuscript received November 22, 2004 Accepted December 6, 2004 IE049352V
