J. Phys. Chem. 1980, 84,2262-2265
2262
Enthalpy Change of Hydrogen Bond Formation between Ortho-Substituted Phenols and Aliphatic Amines Gary Kogowskl, Ronald M. Scott," Department of Chemistry, Eastern Michigan University, Ypsilanti, Michigan 48 197
and Frank Flllsko Department of Chemical Engineering, University of Michigan, Ann Arbor, Michigan 48 106 (Received: October 9, 1979; In Final Form: April 30, 1980)
The enthalpy changes for the formation of hydrogen bonds between phenol, p-cresol, o-cresol, o-sec-butylphenol, 2,6-dimethylphenol, or 2,6-di-tert-butylphenol and n-butylamine, triethylamine, and tri-n-butylamine are measured in a microcalorimeter. The values obtained are interpreted in terms of steric hindrance and the freedom of movement of the amine in the adduct.
Introduction Phenols and aliphatic amines interact to form 1:l hydrogen bonded adducts. The energy involved in the formation of such bonds has been the subject of many studies.'P2 Typical values for enthalpy change, using the interactions of phenol with various aliphatic amines as examples, range from 5.9 to 9.2 kcal/mol. Bonding energies vary according to the acid strength of the phenol, the base strength of the amine, solvation, and steric factors involved in the approach of the phenolic hydroxyl to the amine nitrogen. Of special interest in this project is the interaction of amines with ortho-substituted phenols. The potential of various patterns of ortho substitution on phenol to interfere sterically with hydrogen bond formation to amines was investigated by UV spectroph~tometry.~ The verification of the enthalply changes determined in those studies as well as the extension of these studies to 2,6-disubstituted phenols, for which UV spectrophotometry is not effective, is the goal of these calorimetric studies. Materials and Methods Tributylamine, triethylamine, and n-butylamine were Aldrich Chemical Co. reagent grade. Phenol, p-cresol, o-cresol, and o-sec-butylphenol were supplied by Dow Chemical Co. 2,6-Dimethylphenol and 2;g-di-tert-butylphenol were Eastman Reagent grade. Cyclohexane was Baker Analyzed Reagent grade. All amines and cyclohexane were fractionally distilled twice before use. Phenols were sublimed twice before use, except o-sec-butylphenol which was distilled under reduced pressure. All solutions in the study were prepared by weight. Phenols were weighed in a volumetric flask and then dissolved in a weighed amount of cyclohexane. Amines were diluted in a weighed amount of cyclohexane in a volumetric flask, weighed, and then diluted again. Prepared solutions were transferred to storage bottles with ground glass stoppers. All solutions were stored under a nitrogen atmosphere in a drybox. Before use all phenol solutions were spectrophotometrically analyzed on a Beckman DK-2A to verify that no hydrogen bonding existed, either due to water vapor or intramolecular association. The experimental enthalpies of interaction were measured in a Tian-Calvet differential microcalorimeter similar to that described by Maron and F i l i ~ k o Both . ~ cells of the calorimeter were calibrated by using a calibrating resistor. 0022-3654/80/2084-2262$0 1.OO/O
A known amount of current, verified by a Keithly 179 digital multimeter, was passed through the resistor for a measured period of time. The total area under a recorded curve was then measured with a K & E compensating polar planimeter. The millicalories of heat evolved were then calculated for a series of current inputs. A bucket-type three-tube assembly was designed to measure enthalpy values of the system (Figure 1). Heats of dilution were recorded by the following method. Five milliliters of solution was pipetted into the cell and 1mL of solution into the bucket; the bucket capacity was designed to accommodate exactly 2 mL. When assembled, the solution in the cell rose to a point approximately 3/4 of the way to the top of the bucket to ensure salution (cell)-solution (bucket) thermal equilibrium. The cells were then simultaneously placed into the calorimeter until thermal equilibrium was obtained as indicated by a straight baseline of the recorder. Equilibrium time was approximately 3 h. The bucket was then lowered a calculated distance until exactly l mL of cell solution entered the bucket through holes in the top, and temperature changes were recorded. After the temperature again equilibrated the system was stirred with a paddle to ensure that the reaction was complete. This stirring process was repeated until no further deflection of recorder from baseline was observed. The stirring shaft for the paddle was connected to a synchronous motor (120 rpm) and rotated for approximately 10 s. Using this procedure, we determined heats of cyclohexane-cyclohexane mixing, phenol-cyclohexane dilutions, amine-cyclohexane dilutions, and phenol-amine interactions. In each case the area under the recorded curve was measured with a K & E compensating polar planimeter, and the enthalpies of interaction were calculated. To verify that hydrogen bonding did occur, random triethylaminephenol samples were taken and the spectra were recorded. Reproducibility was estimated by measuring the o-cresol-triethylamine system three times. Enthalpy values of amine dilutions were substracted from the values for enthalpy of formation of the hydrogen-bonded complex. Error analysis was based on the standard deviation of the calibrated cells, the error via use of the compensating planimeter, and standard deviation from reproducibility data. Results The calibration curves for the cells (Figures 2 and 3) establish the linearity of response of the system. The 0 1980 American Chemical Society
The Journal of Physical Chemistry, Vol. 84, No. 18, 1980 2263
Phenol-Amine Hydrogen Bond Formation
n
1
ll!r
CELLTWO - CALIBRATION
towering shaft
Support shaft
Teflon cap
Stainless steel cell
Agitation shaft Phenol ( 1 mi I
L--l
I
- Amine ( 5 ml
0
0
1
Flgure 1. Calorimeter cell. A three-tube bucket-type assembly was designed. After allowing the bucket to reach thermal equilibrium it is lowered by use of the middle shaft, causing mixing of phenol and amine. Then after thermal equilibrium is reestablished, completeness of reaction is tested by stirring the solution by use of the inner shaft. CELL ONE - CALIBRATION
‘1
/
Figure 2. Cell otie-calibration.
amount of heat generated by the mixing process was determined by mixing cyclohexane with cyclohexane and was found to be insignificantly small. Phenols are capable of hydrogen bonding to themselves. A portion of such hydrogen bonds would be broken on dilution of the sample during the mixing process. It is important to know whether such phenol-phenol hydorgen bonding is extensive, since it is significant whether we are measuring the initial formation of a hydrogen bond at the
8
12
AREA ( c m 2 )
Figure 3. Cell two-calibration. Calibration curves for the cells of the microcalorimeter demonstrate the linearity of response of the instrument to heat generated by a calibrating resistor.
phenol, or are replacing an already existant hydrogen bond. M were diluted 1:l with Phenol solutions of 3-4 X cyclohexane resulting in exothermic changes of less than 1mcal/mol. This was considered adequate evidence that the phenol samples were virtually free of hydrogen bonding before mixing. Amine stock solutions were concentrated, consequently enthalpy changes on dilution of these solutions were measured, and values for phenol-amine interaction were appropriately adjusted. Mixing amine stock with an equal volume of cyclohexane in the calorimeter produced the following values: triethylamine (initially 3.52 M), AH = +9.34 f 0.48 cal/mol; n-butylamine (initially 2.49 M), AH +99.23 f 5.1 cal/mol; tri-n-butylamine (initially 4.02 M), AH = -29.35 f 1.5 cal/mol. Previous experience titrating these systems for spectrophotometric analysis and the spectrum of the samples taken after calorimetric analysis assure that the hydrogen bonding between phenol and amine was virtually complete at the concentration of amine utilized. Unsubstituted phenol was included in this study because numerous thermodynamic studies have already been performed with phenol and aliphatic amines. It was felt that comparison ,of our method with others previously employed would be useful. By comparing o- and p-cresol we are able to isolate the effect of a single ortho substitution on the enthalpy change of hydrogen bond formation. Such variables as amine base strength, solvation, and phenol and amine concentration are held constant by the design of the experiment, and the acid strengths of the two phenols are nearly identical (pK, 10.2). Assaying the effect of 2,6 disubstitution on hydrogen bond formation is possible by comparing the data for p-cresol and 2,6-dimethylphenolaIt must be assumed that the effect of any difference brought about by the addition of another alkyl group to the ring is very small when compared to the steric factors introduced. Finally, the
-
2264
Kogowski et
The Journal of Physical Chemistry, Vol. 84, No. 18, 1980
al.
TABLE I: A Summary of Enthalpy Values for the Phenol-Amine Hydrogen Bonding of the Systems Under Study data from this study phenola phenol
p-cresol o-cresol o-sec-butylphenol 2,6-dimethylphenol
- A H , C kcal/mol
method used
TEA
8.91 * 0.42
UVd calorimetry IRd near-IRd
heptane cyclohexane CCl CCl,
9.2 9.08 7.2, 8.35 7.8, 9.2
BA TEA BA TBA TEA BA TBA BA TEA
9.87 f 7.61 i: 8.61 i: 7.63 f 7.35 f 5.20 i 6.05 f 5.52 f 2.71 f 0.05 f 2.41 ?: 0.33 f 0.37 f 0.14 f
5 6 7, 8 9,lO
UVd UVd
cyclohexane cyclohexane
7.24 8.09
3 3
UVd
cyclohexane cyclohexane
6.59 8.49
3 3
BA 2,6-di-tert- butylphenol
data from previous studies solvent . - A H , kcal/mol
amineb
TBA TEA BA TBA
0.46 0.36 0.41 0.36 0.35 0.24 0.28 0.25 0.13 0.002 0.11 0.02 0.02 0.007
uvd
ref
-
a Final phenol concentrations were as follows: phenol, 1.458 X M;p-cresol, 1.785 X M;o-cresol, 1.739 X lo-, M ; o-sec-butylphenol, 2.463 X lo-, M ; 2,6-dimethylphenol, 1.49 X M ; 2,6-di-tert-butylphenol, 1.935 X lo-, M. Final amine concentrations were as follows: triethylamine (TEA), 1.76 M ; n-butylamine (BA), 1.25 M ; tri-n-butylamine Error analysis is wed on the standard deviation of the cell calibration (0.03140) and of the o-cresol(TBA), 2.01 M. triethylamine reaction (0.01570). Spectrophotometry.
’
effect of the size of the ortho substitution is studied by comparing o-cresol with o-sec-butylphenol and 2,6-dimethylphenol with 2,6-di-tert-butylphenoL The values obtained in the experiments and values obtained by previous researchers are summarized on Table
I. Discussion If we first look at the phenol-triethylamine values it is noteworthy that previous values in the literature vary from 7.2 to 9.2 kcal/mol. Calorimetry is generally conceded to be the most accurate of the methods for measuring enthalpy change. The value previously determined with the calorimeter, 9.08 kcal/mol, and our value of 8.91 kcal/mol show good correspondence. The data for p-cresol indicate that the tertiary amines are essentially identical in enthalpy change while the primary amine has a higher enthalpy change. Two of these values are similar to values obtained by UV spectrophotometry3 and follow the same pattern as the equilibrium constants for the same interactions. The corresponding values for the enthalpy change of the primary and tertiary amines forming adducts with phenol are higher but show the same pattern, the primary amine value being approximately 1 kcal/mol higher. Previous studies have interpreted this relative reluctance of the tertiary amine to interact with the phenol as being steric interference due to the geometry of the tertiary amine.syll In comparison with values obtained for p-cresol and tertiary amines, the o-cresol values are lower. Triethylamine is slightly lower, but tributylamine is sharply lower, indicating that the degree of interference to the formation of an o-cresol-tertiary amine hydrogen bond is greater due to the larger alkyl groups of tributylamine. A single ortho substitution should restrict the approach of the amine to a place adjacent to the other (unsubstituted) ortho position. Within limits the size of the ortho substitution should have little effect on the formation of the hydrogen bond, particularly with the primary amine. Replacing the methyl group with a sec-butyl group in the ortho position had little effect on the formation of a hydrogen bond with n-butylamine. The lower enthalpy change when o-cresol is compared to p-cresol bonding to n-butylamine is harder to explain.
Enthalpy values do not parallel the equilibrium constants for these adducts. A decreased entropy change is inferred for the ortho-substituted adduct. The protons of the primary amine provide a possible explanation. The primary amine may be interchanging the phenolic proton with its two amine protons in the bond and, as a result, may assume a variety of conformations with respect to the para-substituted phenol. The presence of an ortho substitution places greater additional restraints on the primary amine than on the already rather rigidly prescribed tertiary amine adduct. Without a necessary weakening of the hydrogen bond itself this results in a lowered enthalpy change. The unusually high enthalpy change observed for the primary amine and either p-cresol or phenol may also be the result of a pair of hydrogen bonds being formed between the phenol and the amine (O-H-:N and N-H..-.:O) in the same fashion as the double H bond reported for alcohol dimers.12-14 If so it is likely that a single o-methyl group could block the formation of the N-H--:O bond. Griffiths and Socratesls found that a single ortho substitution sharply restricts the ability of phenols to self-associate in the pattern of cyclic trimers. The 2,6-dimethylphenol values are much lower as expected because of the increased difficulty of approach of the amine to the phenol. Once again the tributylamine has a lower enthalpy change than the triethylamine with its smaller R groups. Here the value for the enthalpy change for the primary amine is so small we would say in a practical sense the bond is not forming. Why it should drop this low is not clear. All the values obtained with 2,6-di-tert-butylphenol are so small we estimate that the bond is not forming, reflecting the degree of masking of the phenol group that the large ortho substitutions provide. Summary The effect on the enthalpy change for hydrogen bond formation between phenols and aliphatic amines of ortho substitution on the phenol has been studied by microcalorimetry. Effects are relatively small for a single ortho substitutent, but subst,itution of both the 2 and the 6 position, even with methyl groups, is sharply inhibiting. Bonding with primary amines is more inhibited than with
J. Phys. Chem. 1980, 84, 2265-2268
tertiary amines. Possible reasons for this behavior are discussed.
References and Notes (1) Pimentel, G. C.; McClellan, A. L. "The Hydrogen Bond"; Reinhold: New York, 1960.
(2) Joesten, M. D.; Schaad, L. J. "Hydrogen Bonding"; Marcel Dekker: New York, 1974. (3) Farah, L.; Giles, G.; Wilson, D.; Ohno, A,; Scott, R. M. J. Phys. Chem. 1979, 83, 2455. (4) Maron, S. H.; Filisko, F. E. J . Macromol. Sci.-Phys. 1972, B6, 57. (5) Joesten, M. D.; Drago, R. S. J . Am. Chem. SOC.1962, 84, 3817.
2205
(6) Epley, T. D.; Drago, R. S. J . Am. Chem. SOC. 1967, 89, 5770. (7) Zharkov, V. V.; Zhltinkina, A. V.; Zhokhova, F. A. Fir. Khim. 1970, 44, 223. (8) Fritzsche, M. Ber. Bunsenges. Phys. Chem. 1964, 68, 459. (9) Gramstad, T. Acta Chem. Scand. 1961, 16, 807. (10) Singh, S.; Rao, C. N. R. Can. J. Chem. 1966, 4 4 , 2611. (11) Lin, M.; Scott, R. M. J . Phys. Chem. 1972, 76, 587. (12) Van Thiel, M.; Eecker, E. D.; Pimentel, G. C. J. Chem. Phys. 1957, 27, 95. (13) Liddel, U.; Becker, E. D. Spectrochim. Acta 1957, 10, 70. (14) Eecker, E. D.; Liddel, U.; Shoolery, J. N. J . Mol. Spectrosc. 1958, 2, 1. (15) Griffiths, V. S.; Socrates, G. J . Mol. Spectrosc. 1966, 21, 302.
Anion Radical Solvation Enthalpies as a Function of Cation Size Gerald R. Stevenson" and Yoh-Tz Chang Department of Chemjstty, Illinois State University, Normal, Illinois 6 176 I (Received:March IO, 1980)
Calorimetric methods have been utilized to measure the enthalpies of reaction of solvated anthracene anion radical ion pairs with water. These enthalpies were then used in a thermochemical cycle to obtain the heats of solvation of the separated ions in the gas phase (AH" for AN-., + M+, (AN-.,M+),,~v).These enthalpies were found to be more exothermic when dimethoxyethane (DME) rather than tetrarhydrofuran (THF) served as the solvent. However, for both of these solvents the heats of solvation are more exothermic for the smaller cations (increasingexothermicityCs+ < K+ < Na+ < Li'). In THF this enthalpy varies by more than 50 kcal/mol with the size of the cation. The results are explained in terms of ion solvation and ion association, and these results, are compared to a previous study where the anion size was varied.
-
Neutral hydrocarbons can capture electrons that come from a variety of sourlces including electron beams, alkali metals, and negatively charged electrodes to yield anion radicals in the gase phase,l in solution,2 or in the solid states3 The anion radical systems are known to be much more thermodynamically stable in the two condensed states than in the gas phasea3This is due to the very strong intermolecular stabilizing effects of crystal lattice energy in the solid state3 and of ion solvation and ion association in the solution state: In the gas phase the only important consideration in the stability of the anion radical is the electron affinity of the neutral specie^,^ which is 12.7 kcal/mol for anthracenee6 If this gas-phase anion radical is generated from the transfer of an electron from sodium metal, which has an ionization potential of 118.4 kcal/mol,' the gas-phase ions lie 105.7 kcal/mol higher in energy than However, the neutral gas-phase odium and anthra~ene.~ the very large crystal lattice energy (-166.6 kcal/mol) and solvation enthalpy (including ion association) in THF (-178.7 kcal/mol)8 bring the energies of the solid material and of the THF solvakd ion pairs to 50.9 and 63 kcal/mol lower than that of the gas-phase neutral species, respectively. The immense effect that solvation has upon the thermodynamic stability and thus chemistry of anions prompted us to carry out a systematic study of the enthalpy of solvation and thermodynamic stability of the anthracene anion radical in several solvents and with a variety of alkali metal cations. Despite the obvious importance of solvation enthalpies in controlling the chemistry of organic anion radicals, only one previous report of experimental solvation enthalpies has appeared.8 Szwarc and co-workers have collected a vast amount of qualitative information concerning the solvation of anion radical and dianion ion pairs through studies of anion 0022-3654/%0/2084-2265$0 1.OO/O
radical disproporti~nation.~ For ion pairs involving the anthracene anion radical and K+, Na+, or Li+ serving as the cation in THF (tetrahydrofuran) or in DME (dimethoxyethane), cation solvation appears to decrease as the size of the cation increase^.^ Further, DME appears to have a greater capacity to solvate cations than does THF. For instance, K+ is $oorly solvated when it is associated with the anthracene anion radical in THF, but it is strongly solvated in the same ion pair in DMESga The three most commonly used solvents for alkali generation of anion radicals are THF, DME, and HMPA (hexamethylphosphoramide). In HMPA hydrocarbon anion radicals exist free of ion association.l0J1 Since it was our intention to investigate the effect of cation upon the heat of formation of solvated ion pairs, we have chosen to include DME and THF in this first study of the effect of cation upon the heats of formation of ion pairs from the separated gas-phase ions. In a previous stud9 the enthalpies of solvation of a series of polyacene anion radicals generated via sodium reduction in THF were measured, and we were surprised to find that the enthalpies of solvation of the separated gas-phase ions to form the solvated ion pairs (AH"for the reaction depicted in eq l) are within experimental error for the entire A-e,
+ Na+,
-
(A-.,Na+)THF AH" = -177 kcal/mol
(1)
series of polyacene anion radicals that were included in the study. A-. represents the anion radical of naphthalene, anthracene, tetracene, phenanthrene, pyrelene, or pyrene. The smaller anions evidently form tighter ion pairs with the sodium cation because of their more localized charge densities. These tighter ion pairs then interact more weakly with the solvent (THF). The larger anion radicals 0 1980 American Chemical Society
