Ethyl
I
3
L Y
r-BJtyl
I
3
4
5
6
7
0
5
7
8
9
'0
11
'2
Time
m nutes
'
1
2
3
4
Figure 1. Separations of the n-alkyl ether derivatives of the methyl phenols on column IVb at 122 "C
(21). For example, the separation factor for meta- and para-methoxyphenyl alkyl ethers decreases as the temperature increases when using column IVb. The mesomorphic transition temperature, from Table I, for column IVb is 118°C.
1
NOTES
In the case of these phenols and their ether derivatives, all of which have high vapor pressures, even far below their boiling points, the temperature a t which the column is used should be 75 to 100 "C below the boiling point of the most strongly retained isomer. Figure 1 illustrates the separation obtained for several of the n-alkyl derivatives of the isomeric methyl phenols on column IVb. The figure shows the degree of separation obtained :or the various sets of isomers, the increase in retention time with increasing alkyl chain length, and the symmetrical peak shapes obtained for the ether derivatives. I n summary, the identification of various phenols is possible on liquid crystal phases. The separation of isomers is best accomplished by modifying the structure of the solutes so as to emphasize any differences in the shapes of the molecules and by eliminating groups causing extreme molecular interactions, such as inter- and intramolecular hydrogen bonding. Received for review April 18, 1973. Accepted August 1, 1973. This work was supported in part by a Faculty Research Grant administered by the University of North Dakota.
I
Calorimetric Investigation of Adsorption of Aromatic Compounds by Linde Molecular Sieve 13X Delbert J. Eatough, Sedigheh Salim, Reed M. Izatt, James J. Christensen, and Lee D. Hansen Departments of Chemistry and Chemical Engineering, Brigham Young University, Provo, Utah 84602
Titration calorimetric data have been used to calculate the thermodynamic quantities log K , AH", and AS" for many different systems, e.g., proton ionization and metal ligand complex formation (1-4). Titration calorimetry can be used to determine log K values under conditions where other methods are sometimes difficult to apply with accuracy, e.g., the study of reactions in nonaqueous, highly acidic, or highly basic solutions, or the study of the formation of weakly associated complexes in which the concentrations of participants are not easily measured or in which the ligand concentrations have no p H dependence. In the present study, a n additional application of titration calorimetry is presented. Log K, AH", AS", and adsorption capacity values have been determined for the adsorption of aromatic compounds by Linde Molecular Sieve 13X (LMS 13X). The calorimetric log K value and the capacity for the adsorption of aniline by LMS 13X were (1) J. J. Christensen, J. Ruckman, D. J. Eatough, and R. M. izatt, Thermochimica Acta, 3, 203 (1972). (2) D. J. Eatough, J. J. Christensen, and R . M. Izatt, Thermochimica Acta, 3, 219 (1972). (3) D. J. Eatough, R. M. Izatt, and J. J. Christensen, Thermochimica Acta, 3, 233 (1972). (4) D. P. Wrathall and W. Gardner, "Temperature, its Measurement and Control in Science and Industry," H. Plumb, Ed., Vol. 4 , part 3, 1973, p 2223.
126
checked by an independent method of analysis based on the potentiometric titration of aniline in the supernatant solution.
EXPERIMENTAL Reagents and Solutions. Solutions of aniline (vacuum distilled); toluene (Wasatch Chemical, analytical); nitrobenzene (Fisher, certified); and LMS 13X (Linde Division of Union Carbide, powder) in n-hexane (Baker, "Analyzed," dried over molecular sieves) were prepared by dissolving or suspending weighed amounts of the compounds in n-hexane. A solution of perchloric acid in glacial acetic acid was prepared by dissolving HClOl (SO7270, ACS grade) in glacial acetic acid (ACS grade). The resulting solution was standardized against potassium hydrogen phthalate (National Bureau of Standards. acidimetric standard.) The molecular sieve was used directly as received from Linde. All compounds were prepared and stored in a dry nitrogen atmosphere. Procedure. The isothermal titration calorimeter and operational procedures have been described (2, ,5, 6). In the calorimetric experiments. 0.4.44 solutions of each aromatic compound in n-hexane were titrated into the reaction vessel which contained a suspension of 0.5 g LMS 13X in 50 ml of n-hexane. The temperature of all runs was 25.00 " C . Three or more duplicate runs were made for each system studied. Heats of dilution of reactants were (5) J. J. Christensen, H. D. Johnston, and R. M. Izatt. Rev. Sci. lnstrum., 39, 1356 (1968). (6) D. J. Eatough, J. W. Gardner. J. J. Christensen. and R. M . Izatt, unpublished data, 1973.
A N A L Y T I C A L C H E M I S T R Y , V O L . 46, N O . 1, J A N U A R Y 1974
a
:
0
:
Nitrobenzene, ( b )
0
:
Aniline, ( c )
Toluene, (a)
0 1.0
0.5
1.5
m o l e s titrant added
Figure 1. Typical calorimetric titration curves for the interaction of ( a ) toluene, ( b ) nitrobenzene, and (c) aniline with 0.5 gram of the LMS 13X in hexane
Table I . Log K , AH", AS",and Adsorption Capacity Values for the Adsorption of Aromatic Compounds by LMS 13X in n-Hexane at 25 " C
Compound
a
Dipole moment, debyes
Toluene Aniline
0.36 1.53
Nitrobenzene
4.22
AH",
log K 1.17 f 0.10 3.08 f 0.14 3.05 f 0.03b 3.21 f 0.04
kcal/mole -3.3 f 0 . 4 -6.21 f 0.10
-3.50
*
0.12
AS", gibbs/mole -5.7 f 1.3 -6.7 f 0.5 2.9 f 0.4
Adsorption capacity, mmole/g LMS 13X 1.64a 1.64 f 0.10 1.67 f O . O l b 1.37 f 0.01
Value assumed to be the same as that for aniline, see text. Determined by potentiometric titration.
obtained by addition of titrant to hexane in the absence of LMS 13X. A nitrogen atmosphere was maintained over the reacting solutions at all times to avoid contamination of the system by H20. In the determination of the log K value and adsorption capacity for aniline by the analytical method. 0.5 g LMS 13X and 50 ml n-hexane were placed in a titration flask under a nitrogen atmosphere in a constant temperature water bath a t 25.00 "C. One-half milliliter portions of a 0.4M aniline solution were added at 6- to i minute intervals. Two-milliliter portions of the equilibrated supernatant solution in the flask were extracted 5 minutes after each addition. Care was taken to avoid removal of the zeolite. Each 2-ml portion was then added to glacial acetic acid and titrated with perchloric acid in glacial acetic acid using previously described techniques (7, 8). It was desirable to keep the total time of the additions and extractions as short as possible since the decomposition of LMS 13X became appreciable after about 30 minutes in the presence of aniline. Calculations. The calculation of log K , AH", and 15" values from calorimetric titration data has been described (1-4). The variable metric minimization method (3, 9) was modified to allow for the least squares analysis of the data to determine the total number of moles of adsorption sites as well as the log K and AH" values ( 4 ) . In the system studied here, K refers to the reaction: A + Z ZA, and may be expressed as K = [ZA]/[Z][A], where [A] = solute concentration; [ZA] = initial solute concentration, [A]; and [Z] = (total number of adsorption sites)/(liter of solution). [ZA]. Since the terms, ZA and Z , representing the solid phase in the above reaction, appear in both the numerator and denominator in the expression for K , the units used for these terms cancel each other. For convenience in calculations, units of moles/l. were used for each term and the system under study was treated mathematically as a homogeneous system.
-
(7) R. T. Keen and J. S. Fritz, A n a / . Chem., 24, 564 (1952) (8) P. C. Markunas and J. A. Riddick, Anal. Chem. 2 3 , 337 (1951). (9) R . M . Izatt, D. J. Eatough, J. J. Christensen, and R. L. Snow, J.
Phys. Chem.. 72, 1208 (1968).
Adsorption by toluene was relatively weak (log K = 1.2). As a result the log K and adsorption capacity values were highly correlated and both could not be reliably determined from the calorimetric data. The capacity for toluene was assumed to be identical to that for aniline because: (1) the capacity for aniline was the same when determined by two different methods. ( 2 ) the calculated log K and 1fP values for toluene were relatively insensitive to the value of the capacity used. e . g . , use of 1.64 or 1.37 (Table I) gave nearly the same results. Therefore, using 1.64 as the value for the adsorption capacity of toluene should result in reasonable values for other thermodynamic constants, e . g . , log K , AH", and
AS". The log K value and the adsorption capacity for aniline were determined independently by analysis for the amount of aniline adsorbed by the zeolite as a function of total aniline added. The adsorption capacity was found from a plot of aniline adsorbed us. total aniline added. This value combined with the data on the amount of adsorbed and unadsorbed aniline for each addition of aniline, allowed the calculation of log K .
RESULTS The adsorption capacity, log K , AH", and A S o values for adsorption of toluene, aniline, and nitrobenzene by
LMS 13X along with the dipole moments of the aromatic molecules (10) are given in Table I. The uncertainties are expressed as the standard deviations from the mean. Typical calorimetric results for each system are given in Figure 1. The data from which these values were calculated have been deposited as NAPS document No. 02266 (6 pages of supplementary material). Order from ASIS/NAPS, C / O (10) "Handbook of Physics and Chemistry," S l s t ed., 1970-71, Chemical Rubber Co., Cleveland, Ohio, p E71.
ANALYTICAL CHEMISTRY, VOL. 46, NO. 1, JANUARY 1974
127
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DISCUSSION As indicated by the d a t a in Table I, there is good agreement between t h e two log K values and also between the adsorption capacities calculated for aniline by the two different methods. It is therefore concluded t h a t titration calorimetry can provide reliable thermodynamic values for reactions involving adsorption on solids. The speed and wide applicability of titration calorimetry makes it preferable to other methods for the study of solid adsorption reactions especially when decomposition of products or reactants can occur with time. T h e calorimetric titration curves given in Figure 1 clearly show the adsorption capacity regions for aniline and nitrobenzene but not for toluene. The gradual change in slope of Q us moles of added toluene (Figure 1) results in highly correlated adsorption capacity and K values. Thus the two parameters could not be independently measured as previously mentioned. The calculated adsorption capacities for aniline and nitrobenzene are similar as would be expected for these similar size molecules. Whereas log K values for nitrobenzene and aniline are significantly larger than t h e corresponding value for toluene, this is not observed in the case of the AH" values. Rather,
the AW value for adsorption of aniline is significantly more negative than t h a t for either nitrobenzene or toluene. The change in AH" values does not parallel the change in dipole moments in going from toluene with the smallest dipole moment to nitrobenzene with the largest. T h e larger negative A H " value for aniline may be attributed to the presence of a free electron pair on the nitrogen in aniline which allows Lewis acid-base interactions with the zeolite. This argument is strengthened by the results of Abramov et. al. (11) who observed an unusually large change in the infrared spectrum associated wiLh the adsorption of aniline by the zeolite, which they attributed to the effect of t h e free electron pairs in aniline. However, in the case of nitrobenzene, the change in the spectrum was less pronounced and was attributed to the presence of the benzene ring.
ACKNOWLEDGMENT Appreciation is expressed to Linde Inc. for donation of the zeolite used in the study and to E. A. Butler and John Lamb for helpful discussions. Received for review May 3, 1973. Accepted ,July 5 , 1973. Contribution No. 40 from t h e Center for Thermochemical Studies, Brigham Young University. (11) V. N. Abramov, A. V. Kiselev, and V. I . Lygin, Zh. Fiz. Khim., 38, 1044 (1964); Chem. Abstr., 61,38209 (1964).
New Method of Purity Determination by Means of Calorimetric Differential Thermal Analysis Heiner Staub and Werner Perron Mettler lnstrurnente A G , CH-8606 Greifensee, Switzerland
Purity determinations with adiabatic calorimetric measurements based on Van't Hoff's law have been successfully performed for quite some time [see, for example, the work by Glasgow et al. ( I ) ] . Gunn (2) has described a method which allows measurements also with dynamic melting tests. Various authors ( 3 , 4 ) have modified this method for use with DSC or calorimetric DTA. Both the DSC and DTA methods require a correction factor for the calculation of the molar impurity which has no physical basis. Its magnitude is determined by trial and error or according to a formula by Sondack ( 5 ) .Relative errors of u p t o more than 1070of the determined impurity concentration do occur. The present article proposes the use of a new "step heating programming technique" which does not require the above mentioned correction factor and which yields results t h a t are more accurate than those obtained by the dynamic method. (1) A. R . Glasgowetal.. A n a / . Chim. Acta. 17, 54 (1957). (2) S.R. Gunn. Ana/. Chem.. 34, 1262 (1962). (3) G. L. Driscoll. I . N. Duling, and F. Magnotta, "Analytical Calorimetry," Plenum Press, New York, N.Y., 1968, p 271 (4) "Mettler TA2000 Information," No. 1 , 1972, Mettler Instrument Corp.. Princeton, N.J. (5) D.L. Sondack, Ana/. Chem.. 44, 888 (1972).
128
THEORETICAL The method is based on the following equation which is derived from Van't Hoff's law:
T.
=
RT,?x,* 1 T,, - A H , F ~
whereby x-* = x2F = mole fraction of the impurity in the liquid phase = mole fraction of the impurity in the original substance F = fraction melted T , = equilibrium melting temperatuie [K] T o = equilibrium melting point of the pure substance [K] R = gal constant = 1.987 cal/mol K AHi = melting enthalpy of the pure substance {cal/moll x2
x2*
The validity of Equation 1 depends on the applicability of Van't Hoff's law and, for this reason, is subject to the following limitations. Only eutectic systems can be investigated. The impurity may, therefore, not form a solid solution and must, in it5 liquid phase. be ideally soluble in the main component. Moreover, the impurity concentration must be small, and the solid and the liquid phase must be in thermodynamic equilibrium.
A N A L Y T I C A L C H E M I S T R Y , VOL. 46, N O . 1, J A N U A R Y 1974
