HEPTANE ISOMERIZATION
July, 1960
where H,is the rotational part of the Hamiltonian and q i and pi are the conjugate coordinates and momenta. Since the potential energy does not depend upon the momenta, the contribution of the kinetic energy to the partition function in equation 8 just cancels in determining the free angle ratios, 6, and so we obtain =
L2
8r
f .f f e-V/RT sin 0 d0 d p dx +
(9)
849
For benzene we take ur = 0.90 (see Table I) a t 279OK. Using equation 12 me obtain VO= 90 cal./mole (benzene)
This very small value of VOindicates that on the surface there is very little interaction upon the rotational degrees of freedom indicating perhaps a loose structure on the surface. On the other hand, for chloroform, we have 6 = 0.16 a t 273OK. and equation 11 gives
where 0 and x, are Euler's angles and u is the symmetry number. We expect, as did Harkins, Davies, and Clark15 that polar molecules are oriented because of the strong electrostatic forces which act on them so that for appreciable portions of the VO= 3400 cal./mole (chloroform) solid angle, V is large, and, hence, the free angle which indicates an appreciably oriented structure ratio in equation 9 is small. For symmetrical top molecules such as benzene on the liquid surface. Since several of the free and chloroform we can estimate the amount of ori- angle ratios or condensation coefficients are smaller entation if we can write a potential energy function than for chloroform, we would expect an even for rotation. We shall first assume that rotation greater orientation for these molecules. Although about the figure axis is essentially free. For rota- the condensation coefficients indicate that there is tions a t right angles to this figure axis we shall as- orientation on the surface, we have not considered sume that we can approximate the potential energy the surface structure in sufficient detail to deterby a function of the form M V o ( l - cos ne) where n mine whether this order is long or short range. is the symmetry number about this axis. For Liquid molecules possess a cooperative structure benzene n = 2 and for chloroform n = 1. Sub- which is quite different from that of solid or gas so stituting this potential function into equation 9 that it is not surprising that a molecule whose rotation cannot pass adiabatically into the liquid we obtain structure should be rejected. This incompatibility between the structures of the liquid and other phases is reflected in the extraordinary difficulty in sin 0dOdVdx = - n e - V o ( l - c o s n 0)/2RTsinBd0 nucleating both crystallization and boiling in pure liquids as exhibited by pronounced supercooling and (10) superheating. The theory of significant liquid since u = m, where m is the symmet,ry number structure16 which has had marked quantitative about the figure axis. For the case where = 1, success in characterizing liquids pictures the averequation 10 reduces to age molecule as executing solid-like motions when it occupies the solid volume which changes over to & e - - (1 - e - V o / R T ) ( n = 1) (11) vo gas-like motions in strict proportion to the fractional increase in volume. while for n = 2, we obtain
so*'"
(15) W. D. Harkins, E. C. H. Davies and G. I,. Clark, J. A m . Chem. SOC.,39, 541 (1917).
(16)
H. Eyring, T. Rea and N. Hirai, Proc. Natl. Acad. Sci., 4 4 ,
7,683 (1958),
et seg.
HEPTANE ISOMERIZATION BY G. M. KRAMER AND A . SCHRIESHEIM Esso Research & Engineering Company, Linden, New Jersey Received November 81. 1969
The equilibrium composition of the isomeric heptanes was experimentally determined at 38.8". This study was conducted because of the scarcity of experimental data on the heptane isomer equilibrium composition, and because of the conflict between existing data and calculated equilibrium values. Equilibrium was reached among seven of the nine heptane isomers, starting with n-heptane, 3-methylhexane, 2,3-dimethylpentane and 2,4dimethylpentane. Side reactions prevented an accurate determination of the 2,2- and 3,3-dimethylpentane equilibrium values. The experimental and calculated equilibrium values of the heptane isomers were in good agreement. A discrepancy, outside the respective limits of uncertainty, was found only in the case of 2,3-dimethylpentane.
Paraffin isomerization studies1 have pointed out the existence of discrepancies between experiObtained isomer and that (1) (a) F. E. Condon in P. H. Emmett, "Catalysis," VoI. VI, Reinhold Publ. Corp., New York, N. Y.. 1958, chapter 2: (b) G. Egloff, G. Hulls and V. I. Komarewsky, "Isomerization of Pure Hydrocarbons," Reinhold Publ. Corp., New York, N. Y.,1942.
culated from thermodynamic data.2 As the number of carbon atoms is increased, the possible oc(2) (a) F. D. Rossini, E. J. Prosen and K. 5. Pitzer, J . Research Natl. Bur. Standards, 27, 529 (1941); (b) F. D. Rossini, K. S. Pitzer, R. L. .4rnett, R. M. Braun and G. C. Pimentel. "Selected Values of Physical and Thermodynamic Properties of Hydrocarbons and Related Compounds," Carnegie Press, Pittsburgh, Pa., 1953.
Yol. 64 currence of discrepancies increases and reliable experimental values are not usually available. Thus, only a limited amount of experimental data on the isomerization of heptane has been r e p ~ r t e d . ~ This paper discusses the results of an investigation of the equilibrium isomerization composition of the heptanes. Two acid catalysts were used and five of the nine heptane isomers were employed as starting materials. Apparatus and Experimental Phillips Research Grade heptane isomers were used, except for 2,2-dimethylpentane which was kindly supplied by D r . SI.R. Fenske of the Petroleum Refining Laboratory of the Pennsylvania State University. The isomers were passed over a zeolitic adsorbent to remove water and olefins. Isobutane (Matheson) was used u-ithout further treatment. Benzene containing 0.0049& sulfur was used, as supplied by the Baker Company. The aluminum halide was purified by repeated distillations in a vacuum system and handled in a dry box.. Phosphoric acid, (87.4%), was used as a cocatalyst,4 with a metallic oxide as a catalyst support.6 The oxide x a s calrined a t 590". One catalyst consisting of an aluminum halide and phosphoric acid was inhibited with 10 vol. 70benzene (on the heptane), and reactions were carried out in stirred glass flasks at 36.8". A gas bag attached t o the end of a reflux condenser eollected any light material. Samples were periodically ivithdrawn from the reaction flask with a hypodermic syringe for gas chromatographic analysis. A squalene column o n firebrick a t 0" was used to separate and identify all the C? isomers. A second system consisting of aluminum halide supported on a metallic oxide, with isobutane as an inhibitorla was studied in a stainless steel reactor. Samples were periodically removed for analysis through a product withdrawal line.
Results In order to secure accurate equilibrium data, t\Yo conditions must be met. One is the achievement of equilibrium and the second is the knowledge that side reactions have not perturbed the experimental equilibrium. In these studies, both catalytic systems produced the same isomer distribution from n-heptane, 3-methylhexane, 2,3-dimethylpentane and 2.4-dimethylpentane with less than 5% conversion to side products. However, 2,2-dimethylpentane could not be isomerized to equilibrium over these catalysts because of interfering side reactions. Interaction of the hydrocarbons with the catalyst leading to adsorption of the heptanes could not be prevented. It is unlikely that selective adsorption was important since both catalytic systems yielded similar products. Table I shows the isomer distribution reached a t 36.8'. The experimental values are averages from the four starting isomers. The reproducibility of the individual values is in the neighborhood of h 1 5 7 , of the numbers shown. The aluminum halide-metal oxide catalyst caused side reactions and the indicated product distribution was obtained by extrapolating the data to zero degradation. Good agreement was obtained over both catalysts. However, there are apparent differences b e h e e n the experimental values and those calculated from thermodynamic data. These differences are due mainly to the fact that equilibrium (3) J. J. H 1-m Eigk van Voorthuijsen, Ree. traa. chzm., 66, 323 (1947). (4) V. N. Ipatieff and L Schmerllng, U. S. Patent 2,358,011, U. S. Patent 2,402,051. ( 5 ) J J Onen and E E. Stably, U. S. Patent 2,349,458.
TABLE I Isomer, mole %
AlXa-HaPO4 Vapor
Liquid
AlXsOxide Liquid
-Calculated aLiquid Vapor
22 DMP 5.0 7 4.4 23.0 29.4 24 DMP 17.0 22 19.7 8.0 9.6 30.1 26.3 23 D M P 9.0 8 8.2 11.4 33 DMP 9.5 10 8.2 11.2 4.4 8.0 10 5.4 223 TMB 5.5 11.9 26.0 22 30.0 2 MH 9.9 7.6 18.5 16 17.2 6.0 3 MH 2.7 6.0 4 5.9 1.5 n C7 0.9 1.0 1 1.0 3 EP 0.6 * Vapor phase calculations are made using vapor pressure data from API Research Project 44, revised December 31, 1952,2band Raoult's law. The thermodynamic values are calculated from equilibrium data in source.2B
has been reached only among seven of the nine isomers. This point is explained more fully in the following paragraphs. In comparing experimental equilibrium data with the values calculated from thermodynamics, an experimental or calculated error in one compound will introduce errors into the composition values of the other isomers. In order to avoid this, a standard technique is to compare equilibrium values of pairs of isomers. 2,4-Dimethylpentane 7;vas picked as the reference compound (Table 111). This compound was chosen because of its high equilihrium concentration, which in turn, provides confidence in its relative value. The ratios of n-heptane, 2-methylhexane, 3-methylhexane, 3-ethylpentane and 2,2,3-trimethylbutane to this common base are in good agreement with values calculated from thermodynamic data. TABLE I1 RELATIVE RATIOSOF C7 ISOMERS (36.8') Ratio
n-C,
2,4-DnIP 3-b4H 2,4-DMP 2-MH 2,4-DMP 2,2,3-T1113 ____2,4-Dh.IP 3-EP ___ 2,4-DMP
Exptl.
Thermod ynainiczm
0.18
0.16
0.73
0.62
1.OO
1.03
0.46
0.57
0.05
0.06
The remaining isomer ratios differ considerably from the thermtdynamic ratios, as shown in Table 111. TABLE I11 REL.4TIVE RATIOSO F Cs. ISOMERS (36.8") Ratio
2,2-D1IP
Exptl.
Thernrodynamir:
2.4-DMP
0.32
3,3-DMP 2.4-DMP 2,3-DMP a,snnlP
.46
1.66
.36
2.74
1
2,2-Dimethylpentane could not be isomerized without degrading. Therefore, it is not known whether the differences between experimental and calculated isomer ratios are due to errors in thermodynamic data or to interfering side reactions.
July, 1960
SORPTION OF WATERVAPORB Y XATIVE: AND DENATURED EGGALBUMIN
Since the isomerization of 3,3-dimethylpentane was not attemptmedlit is also not known whether this isomer was a t equilibrium. However, 2,3-dimethylpentane was isomerized without side reactions, and its concentration was found to check that produced from the other isomers under conditions of little degradation. Thus, it is felt that the differences in the ratios for 2,3-dimethylpentane are due to errors in the thermodynamic values. It is important to determine what causes the differences between the experimental and calculated equilibrium values. The calculated values are derived from the equation AFQ = -RT In Kes where AFO is obtained from thermodynamic measurements. Values of AFo/T may be experimentally determined using equilibrium constants based on 2,4-dimethylpentane as the reference isomer. These experimental AFo/T values are compared with the calculated values in Table IV. There is excellent agreement in all cases except 2,3dimethylpentane, where a discrepancy outside the respective limits of uncertainty was found. The cause of this discrepancy probably is an error in
851
the thermodynamic data and the subject bears further investigat,ion. TABLE IT' VAPOR PH.4SE
EQUILIBRIUM OF
Isonrer
-Mole Obsd.
SEVEX
V0-Ca1cd.a
2,3-Dimethylpentane 9 . 6 44.3 2,4-Dimethylpentane 26.5 16.2 3-Methylhexane 19.3 10.0 2-Methylhexane 26.5 16.7 2,2,3-Trimethylbutane 12.1 9.3 3-Ethylpentane 1.2 1.0 n-Hept'ane 4.8 2.5 "Ref. 2a. "robable experimental cal./deg. mole. "he uncertainty is mole.
HEPTASES,36.8" --AFB/T-Obsd. b Calcd.08c
2.02 0 0.63 0 1.57 6.17 3.39 error is =t 1.3
- 2.02 0 0.94 -0.06 1.11 5.54 3.70 i 0.46 cal./deg.
Acknowledgment.-The aut,hors wish to thank Esso Research and Engineering Company for permission to publish this work. They also wish to thank Mr. D. L. Baeder and Dr. P. J . Luchesi for their helpful suggestions, and Mr. J. L. Carter who participated in much of the experimental program.
THE SOILYTIOS OF WATEK VAPOR BY XATII~E~ ALBUMIN
s DEKATUKEI) n
EGG
BY ROBERT L. ALTMAN~ AND SIDX'EY W. BEWON Departmeid of Chemistry, University of Southern California, Los Angeles 7, California Received December 8. 1969
The sorption-desorption cycles for water on native and on steam, heat and alcohol-denatured egg albumins have been studied from 25 to 70" for the first and up to 100' for the latter three. The amount of water sorbed is a very weak function of the temperature at high relative humidity and only slightly more sensitive a t low humidities. The size of the hysteresis loop decreases with increasing temperature and for denatured albumin a t 100" it has almost disappeared. Between 70 and 100' denatured albumin shows no hysteresis above a relative humidity of 0.7. The sorption isotherms of the three denatured egg albumins were significantly different both a t 25 and 40' indicating that denatured albumin is not a uniquely defined material.
Introduction The sorption of water vapor on solid proteins is different from the sorption of non-polar gases on these solids. For non-polar gases the adsorption and desorptioii paths of non-polar gases coincide and there is no hysteresis I O O P . ~ ~ ~ ( I ) This work has been supported by a grant (G-8541) from the
U. S. Public Health Service of the National Institutes of Health. (2) The material in this paper has been included in a dissertation hubmitted by R . L. Altman t o the Giaduate School of the University of Southern California (1958) in partial fulfillment of the rprjuirementR for the degree of Doctor of Philosophy. (3) S. Vi'. Benson and I). -4. Ellis, J . .4m. Chem. Soc., 7 0 , :35ti:3 (1948). (4) S. W. Benson and D. A. Ellis, ibid., 73, 2085 (1950).
Experimental data on the sorption of water vapor by egg albumin are relatively complete only in the room temperature region. Barker5 provides sorption and desorption isotherms a t 20" for both native and hea,t denatured egg albumin, and Mellon, Korn and Hoover6 have done the same a t 30". Further observations on mater sorption in this t*emperaturc ~~* range are report'ed by Benson and E l l i ~ ,Shaw,' Bull,*and Benson and Richardson.9 ( 5 ) H. A. Barker, J. Cen. Physiol., 17, 2 1 (1833). (6) E. F. Mellon, A. H. Korn and S.R. Hoover, J . A m . P h e m . S o c . , 71, 2761 (1948). (7) 1'. M. SLaw, .I. Chem. Phys., 12, 391 (1944). (8) H. B. Biill, J . A m . Chem. Soc.. 66, 1499 11944). (91 S. I\-. Benson and R. I,. Ricliardson, i b i d . , 7 7 , 2585 (19.55).
