Download Enthalpy Change of Hydrogen Bond Formation between

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