1074
SURJIT SINGH AND
c. N. R. RAO
Spectroscopic Studies of Self-Association Due to Hydrogen Bonding’
by Surjit Singh and C. N. R. Rao2 Department of Chemistry, Indian Institute of Technology, Kanpur, India
(Received August 29, 1066)
Thermodynamics of dimerization of alcohols, phenols, amines,’thiophenol, and haloforms have been examined by employing infrared and nmr spectroscopy. The energies of various types of hydrogen bonds have been compared.
The energy of a hydrogen bond X - H * . * Y depends on the electronegativity of X and Y as well as other environmental factors. I n recent years, self-association due to hydrogen bonding has been studied in a number of hydroxylic compounds forming 0-H . * 0 bonds employing infrared, n ~ r and , ~ electronic6 spectroscopy. The thermodynamics of self-association of alcohols has been quantitatively investigated in arid the results show that the hydrogen-bond energies in aliphatic alcohols are in the order methanol > ethanol > 2-propanol > t-butyl alcohol. Although uncertainties of the energy values are rather large, it appears likely that the trend in AH” may reflect the steric or electronegativity effects on the strength of the hydrogen bonds. Preliminary studies of trifluoroethanols indicated that trifluoroethanol was considerably less associated than ethanol. In the light of these observations, it was considered worthwhile to study quantitative aspects of the self-association of fluoro alcohols to examine the effect of acidity of the alcohols on the thermodynamic quantities. I n addition, we have also determined the thermodynamics of self-association of substituted phenols of varying acidity. Earlier e t u d i e ~indicated ~ ~ ~ ~ that sterically hindered alcohols and phenols are considerably less associated. The small magnitude of self-association of sterically hindered hydroxy compounds may be due to either the low enthalpy or the entropy contributions to the free energy. Since no quantitative thermodynamic data are available in the literature on the self-association of hindered hydroxy compounds, we have examined the self-association of a hindered alcohol and a phenol. There is little information in the literature on the enthalpies of self-association due to hydrogen bonds other than the 0-He. .O. We have now studied the The Journal of Physical Chemistry
thermodynamics of self-association of amines, mercaptans, and haloforms in order to obtain energies of N-H. - N , S-He .S, and C-H. .X (X for halogen) bonds.
-
-
Experimental Section All of the chemicals employed in the study were of analytical grade and were further purified before use. The self-association of all of the systems with the exception of thiophenol was studied by the variation of the intensity of X-H stretching and the first overtone bands. A Carl Zeiss UR-10 spectrophotometer (LiF prism) and a Cary 14-R spectrophotometer fitted with a variable-temperature cell compartment were employed for recording the infrared spectra in the X-H stretching and the overtone regions, respectively. Association of thiophenol was studied by a Varian A-60 nmr spectrometer. Optical densities of the first X-H overtone band were determined a t the absorption maxima for several concentrations and extinction coefficients calculated (1).Forms part of the Ph.D. thesis of S. Singh submitted t o the Indian Institute of Technology, Kanpur. (2) To whom all the correspondence should be addressed. (3) C. N. R. Rao, “Chemical Applications of Infrared Spectroscopy,” Academic Press Inc., New York, N. Y.,1963. (4) J. A. Pople, H. J. Bernstein, and W. G. Schneider, “High Resolution Nuclear Magnetic Resonance,” McGraw-Hill Book Co., Inc., New York, N. Y., 1959. (5) C. N. R. Rao, “Ultraviolet and Visible Spectroscopy,” Butterworth and Co. Ltd., London, 1961. (6) U. Liddel and E. D. Becker, Spectrochim. Acta, 10, 70 (1957). (7) J. C.Davis, Jr., K. S. Pitser, and C. N. R. Rao, J . Phys. Chem., 64, 1714 (1960). (8) B. D.N. Rao, P.
Venkateswarlu, A. S. N. blurthy, and C. N. R. Rao, Can. J . Chem., 40, 387 (1962). (9) F. A. Smith and E. C. Creits, J . Res. Natl. Bur. Std., 46, 145 (1951). (10) B. G. Somers and H. 9. Gutowsky, J. A m . Chem. Soc., 85, 3065 (1963).
SPECTROSCOPIC STUDIES OF SELF-ASSOCIATION DUE TO HYDROGEN BONDING
1075
from Beer's lam. The concentration was 0.01-0.30 M in the case of alcohols and phenols, 0.02-0.8 M in the case of haloforms, and 0.05-1.0 M in the case of amines. Equilibrium constants, K , of dimerization were determined employing the method of Liddel and Becker.'j The equilibrium constant is related to the slope of the plot of the apparent molar extinction coefficient,e, against the concentration, c
where em0 is the molar extinction coefficient of the monomer X-H stretching or overtone band extrapolated to infinite dilution (c = 0) and K , is the equilibrium constant of dimerization in liters per mole. The use of the limiting-slope method in the evaluation of K , is illustrated in Figure 1. Similar curves are found with both CC14 and benzene solvents, although the K , value in benzene is lower than in cc14. Equation 1 holds good only for cyclic dimers. For open dimers, if one assumes that the absorption coefficient of the open dimer is r X emo, the equilibrium constant is given by n
Kopen =
L
-Kcyclic
2-r
For the particulm case where r = 1, Kopen = 2Kcyclio. It can be seen that the KO,,, and Kcyclicare proportional to each other. Even if the assumption regarding the structure of the dimer is wrong, the value of the enthalpy of dimerization, A H " , will not be affected6 since this is obtained from a linear plot of log K against 1/T (Figure 2). The energy per hydrogen bond will, however, be different, depending on whether the dimer is linear or cyclic:. The equilibrium constants were calculated assuming closed dimers in all the cases a t several temperatures in the range of 5-50' and the enthalpy of hydrogen bonding was calculated (see Figure 2 for typical plots). The values of equilibrium constants are correct to f10%. Only two values of K are mentioned in Table I as well as in the text for purpose of brevity. The uncertainty in AH" is f0.5 kcal mole-' and the uncertainty in AvoH is f5 cm -1.
Results and Discussion Alcohols and Phenols. Thermodynamic data for the dimerization of two fluoro alcohols are compared with the data on the corresponding unfluorinated alcohols in Table I. The enthalpy values are considerably lower for the more acidic fluoro alcohols. However, the AH" values for other aliphatic alcoholssJ are not in line with the data for the fluoro alcohols when cor-
CONCENTRATION
Figure 1. Plot of the extinction coefficient of the monomer 0-H stretching first overtone band of 2-propanol us. concentration (solvent CCL) in moles per liter: open circles, 40'; filled circles, 10'.
Table I : Thermodynamics of Dimerization of Alcohols and Phenols (ROH)
ROH
pKs
16.0 12.4 17.0 12.7 9.95 9.38 10.20
*.. ...
KiP, K4Oo, 1. mole-1 1. mole-1
0.39 0.19 1.25 0.22 0.74 2.20 3.00 0.53 1.00
0.26 0.14 1.00 0.20 0.49 1.70 2.20 0.27 0.43
-AHa, kC81 mole-1
5.0 3.8 2.6 1.2 5.1 3.2
3.8 8.4 10.0
AVOH,
om-1
125 230 125 195 120 170 150 260 127
I n benzene solvent; benzene has been used as the solvent owing to the poor solubility of the fluoro alcohols in C c 4 . The K value as well as the -AH" of ethanol has been found to be lower in benzene solvent than in Cc14.' It may not be unreasonable to compare the K and AH' values for the self-association of a series of alcohols in benzene solvent although there may be some O-H...?r bonding. CCla solvent. M. M. Maguire and R. West, Spectrochim. Acta, 17, 369 (1961).
related with the pK. values. This is probably because in halogen-substituted alcohols, the halogen atoms may competitively form hydrogen bonds. The same problem would arise in any alcohol where the acidity is changed by substitution with electron-withdrawing groups. It is also possible that there will be no simple relation between the acidity of alcohols Volume 71, Number .I March 1067
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SURJIT SINGH AND C.N. R. RAO
0.2 0 .o
Y 0
5
-0.4
-0.8 3.2
3.4
( I / T ) x 10'
Figure 2. Plob of log K us. 1/T for various alcohols (ROH): 1, CaH,; 2, csH6; 3, C2H6; 4, CFZHCFZCHZ;5, CFsCHe.
and their self-aasociation equilibria since the same oxygen atom serves as the donor as well as the acceptor site in self-association. The more acidic alcohols, however, form appreciably strong hydrogen bonds with other electron donors."J2 The AVOHvalues are greater in the fluoro alcohols even though the enthalpy values are lower (Table I). Similar absence of proportionality between AVOHand AH" has been noticed in the earlier studies of selfassociation of aliphatic alcohols by Liddel and Becker.6 There appears to be no reason why AVO= should be linear with respect to AH", particularly when different alcohols are involved. The magnitude of AH" is determined by the O - . . O distance and the linearity between AVOHand AH" should be expected only in a comparable range of Ro ...o.13 Recent studies of hydrogen bonding in various donor-acceptor systems in this laboratory" have also indicated that the linear AvoH-AH" relation is found for the interaction of an acceptor with donors of comparable basicity. The linearity particularly fails in the interaction of sterically hindered hydroxy compounds with donors. '4 Data on the dimerization of a few substituted phenols of varying acidity (Table I) show that the K and AH" of hydrogen bonding do not vary systematically with the pK, values of phenols just as in the case of aliphatic alcohols. The thermodynamics of hydrogen bonding of substituted phenols with other donor molecules have, however, been found to vary linearly with their acidity." Sterically Hindered Alcohols and Phenols. The highly hindered 2,6-di-tbutylphenol and 2,2,4,Ptetramethyl-3-isopropyl-3-pentanol show no evidence of selfThe Journal of Physicol Chemttry
association and remain monomeric even in the pure state. I n these hydroxylic compounds, the &,E in the nmr and the V O H in the infrared spectra were found to be concentration independent. The slightly less hindered 2,&diisopropylphenol and 2,4-dimethyl-3-ethyl3-pentanol show some dimerization and the thermodynamic data on the self-association values of these two hydroxy compounds are given in Table I. The enthalpy values for the self-association of these hindered hydroxy compounds are comparable to those of simple alcohols and phenols. These results clearly show that the low values of equilibrium constants are mainly due to entropy contributions to the free energy (AF" = AH" - T A S " ) . The AVOEvalue for the association of 2,&diisopropylphenol is much greater than that found for phenol (120 cm-') and the AVOHfor 2,4dimethyl-3-ethyl-3-pentanol is comparable to that for t-butyl alcohol (124 cm-') . In the evaluation of the equilibrium constants for the dimerization of alcohols and phenols, the cyclic structure for the dimer has been assumed. In the hindered alcohols, however, there is a good possibility that the dimers are linear. If this is true, the enthalpy data from the present study indicate that a linear hydrogen bond in the open dimer is as strong as two bent hydrogen bonds in the cyclic dimer. Further, the enthalpy values for dimerization reported here for the hindered alcohol and phenol are more reliable since higher polymers were totally absent in these cases unlike in simple alcohol^.^ Amines. The only reliable value of the enthalpy of dimerization for amines is that of aniline (-1.6 kcal mole-') determined by Lady and Whet~e1.l~From the equilibrium constant data of Lady and WhetseP on N-methylaniline, an enthalpy value of -2.1 kcal mole-' was calculated for its dimerization. The self-association of di-n-butylamine has now been examined and the K values at 25 and 40" were found to be 0.13 and 0.12 1. mole-', respectively, in carbon tetrachloride solution. Triplicate determinations of the equilibrium constants at three temperatures showed that the enthalpy value was of the order of 1.0 kcal mole-'. It is noteworthy that the enthalpy values for the self-association of amines are considerably (11) S. Singh, A. S. N. Murthy, and C. N. R. Rao, Trans. Faraday Soc., 62, 1056 (1966). (12) A. Balasubramanian and C. N. R. Rao, Spectrochim. Acta, 18, 1337 (1962). (13) E. R. Lippincott and R. Schroeder, J . Phys. Chem., 61, 921 (1957). (14) 8. Singh and C. N. R. Rao, J . Am. Chent. Sac., 88, 2142 (1966). (15) J. H.Lady and K. B. Whetsel, J . Phys. Chem., 68, 1001 (1964). (16) J. H.Lady and K. B. Whetsel, ibid., 69, 1596 (1965).
1077
SPECTROSCOPIC STUDIES OF SELF-ASSOCIATION DUE TO HYDROGEN BONDING
lower than those for the corresponding hydroxylic compounds. Thiophenol. Although it has been pointed out that thiophenol undergoes weak dimeri~ation,'~ no thermodynamic data :ire available for its self-association. The &,H of thiophenol in the nmr spectra varies linearly with mole fraction in the entire 0-1 mole fraction range and the association shifts are 0.29 and 0.26 ppm, respectively, a t 21 and 35" in carbon tetrachloride solution. Making use of these association shifts, the -AH" for the dimerization of thiophenol is estimated to be about 1.2 kcal mole-l. This enthalpy value is close to the value of 1.0 kcal mole-' estimated for the dimerization of phosphinodithionic acid.18 Haloforms. Becker19 pointed out that chloroform forms dimers by self-association. Jumper and coworkers20t21 have found the dimerization constant for chloroform by rimr to be -0.01 1. mole-' in carbon tetrachloride a t -25". Calculations based on the infrared data of Becker on the C-H stretching band of chloroform give an equilibrium constant of dimerization of -0.3 I. mole-' if the method of Liddel and Beekel.6 is employed for the calculations. Since no bonded C-H peak is seen due to the association of chloroform, the equilibrium constant was also evaluated by the method of Spurr and Byers,22where the contribution from dimers to the intensity of the band is also taken into consideration. The K value thus calculated for the association of chloroform with Becker's data was found to be 0.01 1. mole-'. I n the present study, ;he K values for the dimerization of chloroform and bromoform were determined6 by a study of the con centration dependence of the intensity of the first overtone of the C-H stretching band. The data in CC1, solvent are given in Table 11.
Table II -AHo,
1. mole-1
Kkoo , 1. mole-'
mole -1
0.13 0.62
0.10 0.57
3.2 1.0
Kw,
Chloroform Bromoform
k08l
While the individual values of the equilibrium constants may vary with the method of evaluation, the enthalpy values for the dimerization will remain nearly the same. The AH" of dimerization found for chloroform is considerably greater than for bromoform, although the equilibrium constant is higher in the latter case. The nmr association shift of bromoform (0.95
ppm) is also greater than that of chloroform (0.28 ppm) in cyclohexane solvent.
Concluding Remarks The energy ranges for various types of hydrogen bonds caused by self-association, after taking into consideration all the data available in the l i t e r a t ~ r e , ~ ~ are shown in Figure 3. Even though there is no simple trend in the energy ranges, a comparison of the hydrogen-bond energies found in systems of related structures shows (Figure 3) the energy order of 0H - * S O > N-H- . N > S-H. * .S,a trend which has been found by Singh, Murthy, and Rao" in donoracceptor hydrogen bonding. This trend is similar to what one would expect from electronegativity considerat i o n ~ . The ~ ~ energies found for the CH. * .X bonds in the self-association of haloforms do not fall in line. (17) B. D.N. Rao, P. Venkateswarlu, A. S. N. Murthy, and C. N. R. Rao, Can. J. Chem., 40, 963 (1962). (18) J. Allen and R. 0. Colclough, J. Chem. SOC.,3912 (1957). (19) E. D.Becker, Spectrochim. Acta, 743 (1959). (20) B. B. Howard, C. F. Jumper, and M. T. Emerson, J. Mol. Spectry., 10, 117 (1963). (21) C. F. Jumper, M. T. Emerson, and B. B. Howard, J. Chem. Phys., 35, 1911 (1961). (22) R. A. Spurr and H. F. Byers, J. Phys. C h a . , 62, 425 (1952). (23) G. C. Pimentel and A. L. McClellan, "The Hydrogen Bond," W. H. Freeman and Co., San Francisco, Calif., 1960.
Volume 71, Number 4
March 1967
1078
ARPADELO,JR.,AND PORTER CLEMENTS
It is quite likely that the carbon atom in haloforms is much more electronegative owing to the presence of the highly electron-withdrawing CX1 groups.
Acknowledgments. The authors’ thanks are due to the Council of Scientific and Industrial Research, India, for a research grant.
Hydrothermal Aging of Silica-Alumina Cracking Catalysts
by Arpad Elo, Jr., and Porter Clements Catalyst Division, Nalco Chemical Company, Chicago, Illinois 60629 (Received September 6, 1966)
A logical consequence of the steam deactivation mechanism proposed by Schlaff er, Morgan, and Wilson is that stability to steam should be inversely related t o virgin surface area. When several catalysts were steamed for times between 0.18 and 199 hr at temperatures between 568 and 746” and pressures between 1.0 and 7.7 atm, the expected inverse relationship was found. At the most severe steaming conditions, an unexpected dramatic reduction was found in the pore volume as measured with carbon tetrachloride; however, the pore volume as measured with water remained high, and skeletal density measurements show that the phenomenon reflects the formation of trapped voids rather than catastrophic structural collapse.
Introduction The physical structure of silica-alumina catalysts is generally thought to consist of loose, irregular aggregates of spheroidal ultimate particles. To account for the BET surface areas of fresh catalysts, the ultimate particles would need to have diameters of about 50 A, and particles of this size have been observed with the electron microscope.lg2 The pore volume of the catalyst must then consist of the interstices in the primary aggregates of these ultimate particles. If lower orders of aggregation exist within the 50-A particles, the interstices are of dimensions too small to admit the hydrocarbon molecules whose reaction is to be catalyzed or the adsorbate molecules used for measuring the surface properties. In this sense, it is useful to regard the 50-A particles as the ultimate particles. Numerous investigators have studied the hydrothermal aging of silica-alumina catalysts. Inasmuch as it is generally acknowledged that, for any given conventional silica-alumina composition, catalytic activity is directly proportional to surface area, the term The Journal of Physical Chemistry
“deactivation” has frequently been used as synonymous with loss of surface area. Ries3 showed that in steam deactivation large decreases in surface area are not accompanied by correspondingly large decreases in pore volume, whereas in thermal deactivation in the absence of steam the pore radius (strictly, the ratio of pore volume to surface area) remains relatively constant. Ashley and Innes12and later Adams and Voge,’ presented electron micrographs to support the view that the loss of surface area results from the coalescence of numbers of ultimate particles t o form larger spheroidal bodies. Schlaffer, Morgan, and Wilson4 studied the kinetics of thermal and steam deactivation, reaching the conclusion that the mechanism is the same for both processes, the presence of steam serving merely to hasten the process of thermal deactivation. They (1) C. R. Adams and H. H. Voge, J. Phys. Chem., 61, 722 (1957). (2) K.D.Ashley and W. B. Innes, Ind. Eng. Chem., 44,2857 (1952). (3) H.E.Ries, Jr., Advan. Catalysis, 4,87 (1952). (4) W.G.Schlaffer, C. 2.Morgan, and J. N. Wilson, J . Phys. Chern., 61, 714 (1957).
