Study of the N-H Hydrogen Bond
150-185 "C only the t phase is formed; in the range 220-400 "C these forms are converted to the x phase; at 450 "C all these phases are transformed to stable 0 carbide, or cementite. In this study we have observed the formation of the c' carbide at 255 "C, and it appears to be resistant to relatively high temperatures (400 "C). Support interaction may be responsible for the stabilization of the t' carbide, since Le Caer et alaz1 found that silicon stabilizes the t form. Acknowledgment. This research was supported by the Donors of the Petroleum Research Fund, administered by the American Chemical Society, and by a Dow Chemical Central Research Fellowship in Catalysis.
References and Notes (1) M. A. Vannice and R. L. Garten, Paper 71e, 67th Annual Meeting of the American Institute of Chemical Englneers, Washington, D.C., Dec, 1974. (2) M. A. Vannice, J. Catal., 37, 442, 462 (1975). (3) R. L. Garten and D. F. Oliis, J. Catal., 35, 232 (1974). (4) Yu. B. Kryukov, A. N. Bashkirov, L. G. Liberov, and R. A. Fridman, Kinet. Catal., 12, 88 (1971). (5) H. H. Storch, N. Golumbic, and R. B.Anderson, "The Fischer-Tropsch and Related Synthesis", Wiley, New York, N.Y., 1951. (6) M. Araki and V. Ponec, J . Catal., 44, 439 (1976).
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(7) G. Blyholder and L. D. Neff, J. Phys. Chem., 66, 1664 (1962); J . Catal., 2, 138 (1963). (8) V. U. Vlasenko and G. E. Yuzeforich, Russ. Chem. Rev., 38, 728 (1969). (9) An excellent review of the applications of this technique in catalysis is provided by J. A. Dumesic and H. Topsoe, Adv. Catal., 28, 121 (1977). (10) J. A. Amelse, M.S. Thesis, Northwestern University, Evanston, IiI., Aug, 1977. (11) C. McIlwrick and C. Phillips, J . Phys. E , 6, 1208 (1973). (12) V. I. Goklanskii and R. L. Herber, "Chemical Applications of Mijssbauer Spectroscopy", Academic Press, New York, N.Y., 1968. (13) H. M. Gager and M. C. Hobson, Jr., Catal. Rev., 11, 117 (1975). (14) M. C. Hobson, Jr., and A. D. Campbell, J. Catal., 8 , 294 (1967). (15) W. Meisel, Proceedings of the 5th International Conference, on Mossbauer Spectroscopy, Bratislava, Chech., Sept, 1973, p 200. (16) S. M. Loktov, L. I.Makarenkova, E. V. Slivjnskii, and S. D. Entin, Kinet. Catal., 13(2), 933 (1972). (17) R. A. Arents, Yu. V. Maksinov, I. P. Suzdalev, V. K. Imshennik, and Yu. F. Krupyanskiy, Fiz. Metal. Metalloved., 36(2), 277 (1973). (18) Yu. V. Maksimov, I. R. Suzdalev, R. A. Arents, and S. M. Loktev, Kinet. Catal., 15(5), 1144 (1974). (19) E. Unmuth, private communication. (20) M. A. Vannice, J. Catal., 44, 152 (1976). (21) G. Le Caer, A. Simon, A. Lorenzo, and J. Genin, Phys. Stat. Solid, 6(A), K97 (1971). (22) H. Pichler, Adv. Catal., 4, 271 (1952). (23) H. H. Podgurski, J. T. Kummer, T. DeWitt, and P. H. Emmett, J. Am. Chem. Soc., 72, 5382 (1950). (24) W. Kundig and H. Bommei, Phys. Rev., 142(2), 327 (1966).
Spectrophotornetrlc and Calorimetric Study of the Nitrogen-Hydrogen Bond J. N. Spencer,' Jeffrey E. Gleim, M. Loulse Hackman, Charles H. Blevins, and Robert C. Garrett Department of Chemistty, Lebanon Valley College, Annville, Pennsylvania 17003 (Received September 16, 1977) Publication costs assisted by the Petroleum Research Fund
Thermodynamic functions for the hydrogen bonded complexes N-methylaniline-pyridineand pyrrole-Me2S0 have been determined by the pure base and high dilution calorimetric methods and by IR spectroscopy. The three methods agree for the enthalpy of formation of the pyrrole-Me2S0 complex but do not agree for the N-methylaniline-pyridine complex. A possible pyridine-CC14 interaction may be responsible for the different enthalpies obtained by the three methods. Pure base calorimetric methods have been used to determine the enthalpy of formation of the hydrogen bonded complexes of aniline and N-methylaniline with ethyl acetate, pyridine, and dimethylformamide. These results indicate that two hydrogen bonds are formed by aniline with each base. An estimate of the dimerization enthalpy of adenine has been made by using pure base enthalpies.
Introduction In 1960, Pimentel and McClellanl wrote "it is hardly an exaggeration to say that in the chemistry of living systems the H bond is as important as the carbon-carbon bond". At this time the a helix, pleated-sheet structures, and the DNA double helix were known mostly as models.2 Since that time studies on many other systems in which the hydrogen bond has been shown to be a vital part of the structure have been carried Despite the interest in the hydrogen bond due to its importance in biological processes, thermodynamic parameters for hydrogen bond formation by molecules of biological interest are lacking. One reason for this is that these compounds generally involve weak N-H.-0 or N-H-N linkages. This coupled with the small formation constants for these hydrogen bonded adducts makes accurate measurement of the thermodynamics difficult. Most molecules which participate in biological processes are complex. Thus the thermodynamics of hydrogen bond 0022-385417812082-0563$01.00/0
formation in model compounds containing carbonyl and amide groups has been studied in connection with the stability of protein structures4 and model compounds which are closely related to constituents of nucleic acid have provided a means to investigate molecular features which contribute to the structural stability of nucleic acid! However, the use of model compounds does not alleviate the difficulty in determining thermodynamic properties of weakly hydrogen bonded systems. The standard spectrophotometric analysis is pushed to its limits for the determination of formation constants and enthalpy changes as small as those encountered in these model compounds. Even high dilution calorimetric methods may produce large errors under these condition^.^^' The pure base calorimetric method was proposed by Arnett et ala8in 1970 to deal with those cases where the formation constant for hydrogen bonded adducts was small or difficult to obtain. The pure base method has since been shown to give enthalpies in good agreement with reliable 0 1978 American Chemical Society
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TABLE I: Spectroscopic Determination of Thermodynamic Parameters for N-H.-.N and N-Ho.*O Hydrogen Bonds in CCI, Solvent Proton Aula Proton donor acceptor kcal mol-' K Z g 8 cm-I N-Methylaniline Pyridine 1.9 i 0.6b 1.4 109 Pyrrole Me,SO 4.2* 0.5c 14.7 175
Spencer et al.
Results and Discussion Table I gives the equilibrium constants and enthalpies determined by spectroscopic methods for the formation of the pyrrole-Me2S0 and N-methylaniline-pyridine complexes in CCl& The reported error in the enthalpy is large in both cases and reflects the difficulty encountered with the spectroscopic measurements. King14 has demonstrated that uncertainties in enthalpy changes detera Frequency shift measured relative to the free acid. mined from the equilibrium constant as a function of Equilibrium constants determined at 15 and 30 C. Error estimate from propagation of errors analysis. Error temperature are usually underestimated. For systems with estimate obtained from the slope of a least-squares plot of small enthalpy changes the errors are magnified. In K vs. T 1 . The calorimetric pure base approach to the determination of enthalpy changes is independent of the equiresults obtained by other investigators using different librium constant provided that the equilibrium constant methods of analysis.6 The advantage of the pure base is large enough to ensure complete complexation.s Acmethod is that larger and more easily measured expericording to the pure base method, if a small quantity of mental heats of solution are used for the calculation of the hydrogen bonding acid is injected into the base as a enthalpy of formation. solvent, two contributions to the heat observed are inPure base analysis has shown that bulk medium effects volved; the heat due to hydrogen bonding and that heat on the enthalpy of hydrogen bond formation may be small term which might occur if there were no hydrogen bonding. or at least small in comparison with experimental e r r ~ r . ~ If ~ ~a ~ ~ proper model compound is chosen, the hydrogen This would suggest that different thermodynamic pabonding heat term can be isolated. An inert reference rameters obtained for a given system in different solvents solvent is chosen to correct for heats of solution of the acid may be primarily due to an interaction of the acid or base and model compound. Then the enthalpy of formation with the solvent. If these interactions predominate over may be calculated from bulk solvent effects, measurements of hydrogen bond enthalpies by the pure base method may give better esaH= ( A T s A - aHsM)base - (AHsA - nHsM)ref(1) timates of the actual hydrogen bond strength formed by where the superscripts A and M refer to acid and model, interpeptide linkages or nucleic acid pairs than those respectively. enthalpies obtained by other methods in solvating media. Bertrand and Duer15 have presented evidence derived Experimental Section from measurements on phenol-base systems that both the choice of model compound and reference solvent can Purification of most of the solvents used in this work influence the enthalpy calculated from eq 1. Arnett et a1.16 has been previously describeda1@13Fisher certified ACS have also shown that high estimates of the enthalpy obSpectranalyzed ethyl acetate was dried over dri-Na and tained from adducts of p-fluorophenol by the pure base distilled under nitrogen. Aldrich Analyzed pyrrole, Fisher method relative to the enthalpy obtained in CC14solvent certified ACS aniline, Fisher reagent grade N,N-dimay be related to the polarity of the medium. Thus for methylaniline, and Eastman practical grade N-methylsolvents such as dimethylformamide or MezSO, pure base aniline were dried over NaOH and distilled under nitrogen. measurements may lead to higher calculated enthalpies Aldrich Analyzed N-methylpyrrole was distilled over BaO than those found in CCl,. under nitrogen. Only freshly distilled samples of all amines The high dilution calorimetric method8 assumes that if were used. a small quantity of acid or base is injected into a medium Spectra were recorded on the Beckman Acta M-IV which contains a species capable of hydrogen bonding to recording spectrophotometer. The methods of calculation the acid or base, the enthalpy of hydrogen bond formation have been previously described.'O For the N-methylmay be found from aniline-pyridine and the pyrrole-MezSO systems in C C 4 solvent, the free acid and complex absorption bands A H ' = AHCcV (2) overlap. Concentrations were chosen so as to minimize the where AH' is the heat observed corrected for the heat of overlap and when this was not possible, the overlapping solution of the injected species, C, is the concentration of peaks were resolved by assuming a Gaussian band shape the hydrogen bonded complex, and V is the total volume and symmetrically reflecting the bands about their centers. of the system. If the equilibrium constant, usually dePyrrole concentrations of about 0.007 M and MezSO termined from spectroscopic measurements, is known C, concentrations of 0.02 M were used. N-Methylaniline may be calculated and AH directly solved for by eq 2. concentrations were 0.01 M and pyridine concentrations Even though eq 2 is not totally reliable for use with small were 0.1 M. Enthalpy changes were calculated from the equilibrium constants? eq 2 still is superior to the plotting temperature dependence of the equilibrium constant as of spectroscopically derived equilibrium constants vs. outlined previou~1y.l~ temperature to obtain enthalpy values for such systems. The calorimeter was the same as that previously Table I1 lists the heat of solution at infinite dilution for described13except that a Houston Instrument Omniscribe the acids and bases of this work. Table I11 gives the strip chart recorder replaced the computer for recording enthalpy of hydrogen bond formation calculated by the data. Water from a temperature bath maintained at 25 pure base method for the two reference solvents cyclo"C was pumped through the Dewar flask containing the hexane and CC14. The model compounds chosen were solvent. The sample to be injected was brought to the N,N-dimethylaniline for aniline and N-methylpyrrole for same temperature as the solvent by equilibration in the pyrrole. The enthalpy values obtained for the two reftemperature bath. Samples (1-10 mmol) were injected into erence solvents are not markedly different within error 200 mL of solvent. Injections were made into a minimum limits. However, a clear trend is evident. The calculated of three different freshly prepared volumes of solvent. All enthalpy values are in all cases higher when cyclohexane samples for spectroscopic or calorimetric analysis were is the reference solvent than for CCl& These results are prepared in a drybox.
The Journal of Physical Chemistry, Vol. 82, No. 5, 1978 565
Study of the N-H Hydrogen Bond
TABLE 11: Heat of Solution at Infinite Dilution (kcal mol-' )" Solvent
cc1, Aniline N-Methylaniline N,N-Dimethylaniline Pyrrole N-Methylpyrrole a Temperature is 298 K.
+2.11 t +1.34 k +0.20 t t2.22 * t0.38 t
0.04 0.03 0.04 0.06 0.02
Ethylacetate
Cyclohexane
-1.76 i 0.02 -1.03 t 0.02 +0.03 ?I 0.01
-0.76 * 0.03 -0.26 t 0.04 t 0 . 2 4 0.02
t3.72 k +2.77 t t1.54 i t3.74 k t1.85 t
-AH," kcal mol-' (reference solvent) CCl, 3.70 i 0.11 2.91 t 0.13 Aniline-N,N-dimethylformamide4.56 * 0.11 2.20 i 0.10 N-Methylaniline-pyridine 1.64 * 0.13 N-Methylaniline-ethylacetate N-Methylaniline-N,N-dimethyl-2.55 * 0.10
Aniline-pyridine Aniline-ethylacetate
C6H'Z 3.97 t 3.18 t 4.83 i 2.29 i 1.73 t 2.64 i
0.10 0.12 0.10 0.10 0.13 0.10
formamide 4.23 t 0.12 4.28 i 0.24 Pyrrole-Me,SO a kcal per mol of proton donor. TABLE IV: Comparison of Pure Base, High Dilution, and Spectroscopic Enthalpies of Hydrogen Bond Formation -AH, kcal mol-' ~~
High
Pure base
N-Methylaniline- 2.20 t pyridine 2.29 i Pyrrole-Me,SO 4.23 t 4.28 t
0.04 0.04 0.03 0.13 0.06
Me,SO
-2.29 * 0.01 +0.10 t 0.03
formamide
-2.51 i 0.02 -1.27 i 0.02 +0.14 i 0.01
The error reported is the standard deviation of a single measurement.
TABLE 111: Enthalpy of Formation of N-H...N and N-H-.O Bonds Calculated by the Pure Base Method
Adduct
N, N-Dimethyl-
Pyridine
dilution
Spectrosc
1.3 * 0.2c
1.9 i 0.6 O.lOb 0.12a 4.2 t O . l d 4.2 * 0.5 0.24b 4.08 t 0 . 1 V O.lOa
Reference solvent is a Reference solvent is CCl, . C,H, * . Calculated from eq 2 using k given in Table I. M. S. Nozari and R. S. Drago, J. Am. Chem. SOC.,92, 7086 (1970). Determined calorimetrically through use of the Drago-Bolles equation. consistent with those of Bertrand and Duer,15who found that for three out of four hydrogen bonded adducts with phenol, the enthalpy calculated using cyclohexane as the reference solvent was 0.2 kcal m4-l higher than that calculated when CCll was the reference solvent. Bertrand and Duer attributed the difference in calculated enthalpy for the two reference solvents to an interaction of phenol with CC,. In Table IV, a comparison of the enthalpies obtained for the formation of the N-methylaniline-pyridine and pyrrole-Me2S0 complexes by the pure base, high dilution, and spectroscopic method is given. The good agreement between all three methods for the enthalpy of the pyrrole-Me2S0 adduct is surprising, especially when it is considered that Me2S0 is a highly polar solvent. The agreement between the three methods for the N methylaniline-pyridine complex is not good. The spectroscopic value for the enthalpy has a large uncertainty but the high dilution value, although suspect for reasons previously stated, should be somewhat reliable. A possible reason for the discrepancy between the pure base and high dilution results may be that an interaction of the Nmethylaniline or pyridine with the solvent occurs. Because pyrrole would be expected to undergo a stronger interaction with CC14 than N-methylaniline and because the pure base and high dilution methods agree well for the
pyrrole-Me2S0 enthalpy, the most likely interaction is one between pyridine and CCl,. Drago et al.16J7have long argued for a pyridine-CC1, interaction of 0.9 kcal mol-l. If this interaction is accepted, the pure base and high dilution results are brought into agreement. A further implication is that enthalpies determined in CCll solvent for the important N-H-mN bond would underestimate the actual hydrogen bond strength in the apolar interior of certain molecules by 0.9 kcal molv1 of hydrogen bonds. The data of Table I11 show that the enthalpy of formation of the aniline complexes with pyridine, ethyl acetate, and N,N-dimethylformamide are larger than for the corresponding complexes of,N-methylaniline. This would be expected if hydrogen bonding to both hydrogens of the aniline amino group occurs. It would also be expected that the enthalpy change for aniline complexes would not be twice as great as that for the N-methylaniline complexes because the ability of the second amino proton to hydrogen bond should be less after the first has bonded.18 It is the ultimate goal of the study of model compounds to be able to apply the results to more complex molecules. Although the model compounds of this work cannot be said to be truly representative of the more complex nucleic acid bases, an estimate of the hydrogen bond strength in the adenine dimer may be attempted. There are three possible structures for the cyclic dimer of eth~1adenine.l~ The two adenine residues may associate by two hydrogen bonds from the amino group to the imidazole or the pyrimidine nitrogen of adenine. Previous work has shown that hydrogen bonding in one part of a molecule does not significantly affect hydrogen bonding in another part of the molecule.20 Thus if other structural features are ignored, the enthalpy of dimerization of adenine may be approximated by twice the hydrogen bond enthalpy of the Nmethylaniline-pyridine adduct, giving a dimerization enthalpy for adenine of -4.40 f 0.20 kcal mol-l. Kyogoku, Lord, and Rich19 using IR spectroscopic methods report the enthalpy of association for the adenine dimer to be -4.0 f 0.8 kcal mol-l.
Acknowledgment. Acknowledgment is made to the Donors of The Petroleum Research Fund, administered by the American Chemical Society, for support of this research. The authors also acknowledge Sherry Etter and J. Christopher Peiper for most of the preliminary calorimetric work and Sherri Gettle, a Project SEED student, for her assistance in many phases of this work. Supplementary Material Available: Spectroscopic concentrations and absorbances for the calculation of the thermodynamic parameters in Table I (3 pages). Ordering information is available on any current masthead page. References and Notes (1) G. C. Pimental and A. L. McClelland, "The Hydrogen Bond", W. H. Freeman, San Francisco, Calif., 1960 p 9. (2) G. C. Pimentel and A. L. McClelland, Annu. Rev. Phys. Chem., 22, 347 (1971).
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(3) M. D. Joesten and L. J. Schaad, “Hydrogen Bonding”, Marcel Dekker, New York, N.Y., 1974. (4) S. J. Gill and L. Noll, J . Pbys. Chem., 76, 3065 (1972). (5) S. J. Gill, M. Downing, and G. F. Sheats, Biochemistry, 6, 272 (1967). (6) E. M. Arnett, E. J. Mitchell, and T. S. S.R. Murty, J. Am. Cbem. Soc., 96, 3875 (1974). (7) S.Cabanl and P. Glanni, J . Chem. SOC.A , 547 (1968). (8) E. M. Arnett, L. Joris, E. Mitchell, T. S. S. R. Murty, T. M. Gorrie, and P. v. R. Schleyer, J. Am. Cbem. Soc., 92, 2365 (1970). (9) E. M. Arnett, T. S. S. R. Murty, P. v. R. Schleyer, and L. Joris, J . Am. Cbem. Soc., 89, 5955 (1967). (10) J. N. Spencer, R. A. Heckman, R. S. Harner, S. L. Shoop, and K. S. Robertson, J . Pbys. Cbem., 77, 3103 (1973). (11) J. N. Spencer, R. S. Harner, and C.D. Penturelli, J . Pbys. Cbem., 79, 2488 (1975).
Stephen E. Stein (12) J. N. Spencer, J. R. Seigart, M. E. Brown, R. L. Bensing, T. L. Hassinger, W. Kelly, D. L. Housel, and G. W. Reisinger, J . Pbys. Cbem., 80, 811 (1976). (13) J. N. Spencer, J. R. Sweigart, M. E. Brown, R. L. Bensing, T. L. Hassinger, W. Kelly, D. L. Housel, G. W. Reisinger, D. S. Reifsnyder, J. E. Gleim, and J. C. Peiper, J . Pbys. Cbem., in press. (14) E. J. King, “Acid-Base Equilibria”, Pergamon Press, Oxford, 1965. (15) W. C.Duer and G. L. Bertrand, J. Am. Chem. SOC.,92,2587 (1970). (16) M. S. Nozari and R. S.Drago, J. Am. Cbem. Soc., 94, 6877 (1972). (17) R. M. Guidry and R. S. Drago, J . Pbys. Cbem., 78, 454 (1974). (18) P. A. Kollman and L. C. Allen, Chem. Rev., 72, 283 (1972). (19) Y. Kyogoku, R. C.Lord, and A. Rich, J . Am. Cbem. Soc., 89,496 (1967). (20) J. N. Spencer, R. S. Harner, L. I. Freed, and C.D. Penturelll, J. Pbys. Cbem., 79, 332 (1975).
On the High Temperature Chemical Equilibria of Polycyclic Aromatic Hydrocarbons Stephen E. Stein Department of Chemistry, West Virginia Universiv, Morgantown, West Virginia 26506 (Received September 26, 1977) Publlcation costs assisted by the Department of Energy, Morgantown Energy Research Center
Concentrations of selected benzenoid polycyclic aromatic hydrocarbons in equilibrium with acetylene and hydrogen over the range 1400-3000 K are estimated by extension and application of existing predictive methods. Implications of these results concerning chemical equilibria and reaction kinetics of high temperature hydrocarbon systems are explored. Of special interest is the finding that in idealized homogeneous equilibrium systems above 1700 K, there exists a “critical” highly condensed polycyclic aromatic species whose concentration is at a minimum relative to other highly condensed aromatics. Equilibrium concentration and size of this critical species depend on temperature and partial pressures of light gases (chiefly acetylene and hydrogen) and this species may constitute a natural thermodynamic barrier for homogeneous high temperature carbon particle formation. Also, concentrations of polyaromatic members of the most thermodynamically stable “benzene polymerization” pathway are examined in selective equilibrium with benzene and hydrogen. For this series, over a wide range of conditions, o-terphenyl is found to have the lowest equilibrium concentration.
Introduction Incomplete combustion of hydrocarbon fuels not only wastes energy but also creates serious environmental problems resulting from the release of soot and carcinogenic polycylic aromatic hydrocarbons (PCAHs) into the atmosphere. These problems may become more acute as coal-derived fuels with higher carbon-to-hydrogen ratios than conventional petroleum derived fuels increase in use, since the former substances are more prone to incomplete combustion. Bituminous coal itself, when burned, can produce copious quantities of smoke, since the highly aromatic tars evolved upon heating rapidly degrade to soot. Soot and PCAHs are also formed in many industrial processes involving high temperature hydrocarbon reactions, as in carbon black manufacture, pure pyrolysis of hydrocarbons, and coal pyrolysis. The ubiquitous nature and importance of such hydrocarbon “polymerization” processes is reflected by the enormous number of studies reported in this area.l However, due to the intrinsic complexity and experimental difficulties inherent to this area of study few details concerning chemical mechanisms in even rather wellcontrolled chemical systems are known with certainty. Studies of species concentration in premixed hydrocarbon-oxygen flames,2 and pyrolysis studies in shock tubes3 and flow reactor^,^ have provided fundamental information regarding soot and PCAH formation. Flame studies are exemplified by the elegant work of Homann 0022-3654/78/2082-0566$01 .OO/O
and Wagner2 who studied concentration profiles of a number of species in soot formation zones of premixed flames, and deduced that under their conditions polyacetylenes are the primary soot precursors. However, these experiments could not answer mechanistic questions with certainty. Moreover, other soot and PCAH forming systems may involve different mechanisms in formation, growth, and agglomeration of soot particles and it is not straightforward to apply results of one system to another. Shock tube experiments have clarified conditions and rates of soot formation under isothermal homogeneous conditions, however, little unambiguous chemical mechanistic insight has resulted from such studies. High temperature flow studies have been severly hindered by undesired heterogeneous reactions on reactor walls. The chemical steps leading to a nascent soot particle are particularly difficult to study experimentally since precursor species need be present in only trace concentrations. Observation of major species profiles in soot forming zones cannot, in themselves, reveal these chemical steps. Clearly, to elucidate soot formation, thermochemistry and kinetics of all reactive species present in soot forming regions must be well understood. A fundamental area of uncertainty in these systems concerns the thermodynamic and kinetic behavior of PCAHs. The influence of temperature, size, degree of ring condensation, etc. on individual PCAH stability has never been quantitatively analyzed primarily due to the lack of 0 1978 American Chemical Society
