THERMOCHEMICAL STUDIES OF AMINES
3759
Thermochemical Studies. XV.' Thermodynamics of Protonation of Triethylenediamine, Triethylamine, Trimethylamine, and Ammonia in Aqueous Solution at 25"
by Piero Paoletti, John H. Stern,2and Alberto Vacca Ietituto d i ChimacCr Cfenerale, Universitb, Florence, Italy
(Received April 19,1966)
Thermodynamic functions of protonation AF, AH)and AS were determined via e.m.f. and calorimetric methods for triethylenediamine, triethylamine, trimethylamine, and ammonia. Correlations for the above series, as well as for other related amines in terms of ion-solvent interactions and amine type, were also made.
Introduction Previous communications in this series3#'have dealt with the thermodynamics of stepwise protonation of several amines (B): B zH+ = BHZ2+(1: = 1, 2, 3, . . .). Direct reaction calorimetry and e.m.f. measurements have yielded AH and AF, respectively. The entropy of protonation Ah' was calculated from the above properties. This contribution describes an analogous determination of thermodynamic functions of protonation for a series of related amine bases in aqueous solution at 25". The series consists of triethylenediamine, N(CH2-CH2))N, triethylamine, trimethylamine, and ammonia. Very few thermodynamic data are available for triethylenediamine. Studies on this globular molecule include determination of the heat capacity, heats of fusion, and transition of the pure solidI5and the basicity constantsa in aqueous solution. Triethylamine and trimethylamine were studied t-o provide a basis for thermodynamic comparison, particularly for the first protonation step of the diamine. The values of AH are believed to be the first obtained calorimetrically. Calculated values of have been reported by others from e.m.f. measurements, but this method has in many cases led to heats of much lower accuracy than those determined directly. Finally, the thermodynamics of protonation of ammonia were included in this study since a direct calorimetric determination of this important reaction seemed advi~able.~The observed differences in these thermo-
+
dynamic properties provide an interesting insight on weak base-strong acid behavior, particularly from the standpoint of solute-solvent interactions, type of amine, molecular structure, and their effects on protonation. Analogous studies have been reported by others on weak acid-strong base behavior. lo
Experimental Section Materials. Pure triethylenediamine was obtained as a gift from the Houdry Co. and repurified by vacuum sublimation. Solutions were prepared by weight, with carbon dioxide free water. The concentration was further checked by titrating the amine with standard hydrochloric acid. Triethylamine, trimethyl(1) Part XIV: P. Paoletti, Trans. Faraday soc., 61, 219 (1965). (2).International Faculty Fellow (1964-1965),American Chemical
Society Petroleum Research Fund. (3) (a) L. Sacconi, P. Paoletti, and M. Ciampolini, J . Am. Chem. Soc., 82, 3831 (1960); (b) M. Ciampolini and P. Paoletti, J . Phye. Chem., 65,1224 (1961); (c) P.Paoletti, M. Ciampolini, and A. Vacca, ibid., 67, 1065 (1963). (4) (a) P. Paoletti and M. Ciampolini, Ric. Sci., 33 (11-A), 405 (1963); (b) P. Paoletti and A. Vacca, J . Chem. SOC.,5051 (1964). (5) J. C. Trowbridge and E. F. Westrum, Jr., J . Phys. Chem., 67, 2381 (1963). (6) G. Schwarzenbach, B. Maissen, and H. Ackermann, Helv. Chim. AI%, 35, 2333 (1952). (7) D.H. Everett and W. F. K. Wynne-Jones, Trans. Faraday SOC., 35, 1380 (1939). (8) W. S. Fyfe, J . Chem. SOC.,1347 (1955). (9) E. J. King, "Acid-Base Equilibria." Pergamon Press, Ltd., London, 1965,p. 197. (10) (a) L. P. Fernandea and L. G . Hepler, J . Am. Chem. SOC., 81, 1783 (1959); (b) L.G.Hepler, ibid., 85, 3089 (1963).
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P. PAOLETTI, J. H. STERN, AND A. VACCA
amine, and ammonia were all prepared from A.R. grade stock solutions and their concentrations determined via titration with standard acid. The purity of all solutions of the three organic amines wm examined by vapor-liquid chromatography. Hydrochloric acid was standardized gravimetrically as silver chloride. Carbon dioxide free solutions of potassium hydroxide, used in e.m.f. measurements, were kept in Jena-glass flasks. Calorimeiric Measurements. The calorimeter“ and details of experimental measurementsa at 25” have been described elsewhere. The estimated over-all experimental accuracy was *0.6%. The ionic medium was 0.1 N KC1 and the concentration of protonated amines ranged from approximately 1.5 to 3 X M. E.m.f. Measurements. The potentiometric experiments were performed with a Radiometer PHM 4 potentiometer using a glass electrode (Radiometer G 2025/B), a 0.1 Jf calomel electrode, and a bridge filled with 0.1 M KC1 solution. All titrations were carried out in a 150-ml. seven-necked flask while the solution was stirred continuously by a magnetic stirrer. Argon was bubbled through the solution which contained 0.1 N KC1, and the cell assembly was thermostated at 25.0 f 0.1’. In each determination 100 ml. of a solution of 0.002 M amine hydrochloride was titrated with varying volumes of 0.1 M KOH added by a piston buret graduated to 0.01 ml. The titer of the KOH solution was checked daily against standard HCI. The hydrogen ion concentrations were calculated from the experimental e.m.f. values, using the formula
[H+] = antilog [ ( E - Eo)/59.154]
-333.1, -3340.4 mv.; e.m.f. (obsd.) = -311.9, -324.2, -332.9, -340.2 mv. ; Eo = 287.5 mv.
Resuits The two-step heats of protonation of triethylenediamine were determined by measuring the heats evolved for different amine-HC1 ratios and by calculating the exact amounts of mono- and diprotonated forms of the polyamines before and after the reaction. The calorimetric measurements for triethylamine, trimethylamine, and ammonia were carried out with a slight excess of amine in the dewar. All heats were OH- = HzO, corrected for the side reaction H + where OH- results from the hydrolysis of the amine in aqueous solution. The heat of formation of water chosen for the correction was 13.34 kcal./mole, which is the confiensus of the most reliable recent determinations.12 In all cases the net correction applied to the calorimetric results was very small. The concentrations of all species present in the calorimeter before and after the HC1 addition were calculated by solving the following system of equations via successive approximation
+
[acidltota~ = [H+I - [OH-] [amineltotal= [Ll
+
+
m j=l
jk,[H+l’[Ll
m
kj[H+l’[Ll
3=1
[OH-] = k,[H+]-l [HJL”]
=
k,[H+]’[L]
[L] is the free amine concentration, j is the number of protons with a maximum of m in the protonated amine HjLj+ with a formation constant k j . k , is the dissociation constant of water in 0.1 M KC1 (1.62 X 10-14).13 This procedure follows the method of Sch~arzenbach’~ and 0the~s.l~ Appropriate programs were written for the IBM 1620 computer to perform the following calculations : (a) least-square determination of basicity constants from the e.m.f. data, (b) equilibrium distribution of all species before and after the reaction in the calorim-
The standard potential of the electrode system, Eo, was determined from the experimental * e.m.f. values. These were obtained during the titration of a solution of 0.001 Jf HC1 (in 0.1 M KCI), with 0.1 M KOH, either in the acidic or basic region. This determination was carried out before and after each basicity constant experiment, and the observed variation in Eo was found to be negligible (less than 0.3 mv.) in all measurements. Each neutralization curve was drawn from at least (11) P. Paoletti, R. Usenza, and A. Vacca, Ric. Sci., 35 (11-A), 201 20 experimental points. A minimum of 12 points and (1965). three curves was used to calculate the basicity constant. (12) (a) J. D. Hale, R. M. Iaatt, and J. J. Christensen, J . Phya. The agreement between the observed e.m.f. values and Chem., 67, 2605 (1963); (b) C. E. Vanderaee and J. A. Swanson, ibid., 67, 2608 (1963); ( c ) L. Sacconi, P. Paoletti, and M. Ciamthose calculated on the basis of the computed equipolini. R k . Sci., 29, 2412 (1959). librium constants was excellent. A representative (13) H. S. Harned and B. B. Owen, “The Physical Chemistry of sample of the data for triethylamine follows: [Et%- Electrolytic Solutions,” 3rd Ed., Reinhold Publishing Corp., New N. Y., 1958,p. 762. N.HCl]in,tia~ = 1.868 X M ; V i n i t i a l = 100.3 York, (14) G.Schwaraenbach, Helv. Chim. Acta, 33, 947 (1950). ml.; volume of 0.1456 M KOH added: 0.40, 0.60, (15) M. J. L. Tillotson and L. A. K. Staveley, J. Chem. Soc., 3613 0.80, 1.00 ml.; e.m.f. (calcd.) = -311.4, -324.2, (1958). The Journal of Physical Chemistry
THERMOCHEMICAL STUDIESOF AMINES
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eter, (3) least-square determination of the heats of protonation together with appropriate standard deviations. Table I shows a summary of all calorimetric measurements. The third column shows the heat evolved in each run, Q, corrected for the heat of dilution of HC1. The last column reports the calculated hydrolysis corrections, Qo. Table I1 shows the values of the thermodynamic functions AF, AH, and AS for all protonations. Values of basicity constants for triethylenediamine reported in ref. 6, measured at 20' and in a different ionic medium, are in poor agreement with present results. The values of ref. 8 for the indirectly determined AH and A S of triethylamine are also in disagreement with this work. All three thermodynamic variables previously determined for trimethylamine' and ammonia,9 however, are in excellent agreement with the data of Table 11.
Table I : Calorimetric Data for the Systems Amine HC1
+
"21,
Base, mmoles
mmoles
Cal.
Triethylenediamine (165 ml.)
4.974 4.980 4.995 4.985 2.412 2.412 2.422
4.822 4.825 4.833 4.824 4.827 4.824 4.855
35.75 35.75 35.83 35.95 23.55 23.55 23.70
1.33 1.34 1.34 1.34 0.93 0.93 0.93
Triethylamine (165 ml.)
5.22 5.22 5.22 5.22
4.697 4.687 4.605 4.645
51.23 50.95 50.00 50.75
10.65 10.68 10.64 10.66
Trimethylamine (165 ml.)
4.35 4.35 4.35
3.969 3.968 3.918
36.32 36.21 35.95
3.45 3.45 3.45
Ammonia (165 ml.)
4.22 4.22 4.22 3.99 3.99 3.99
3.554 3.557 3.551 3.583 3.623 3.564
44.01 44.63 43.80 44.74 45.34 44.90
1.99 1.99 1.99 1.89 1.89 1.89
Base
QV
Q0.O
oal.
a This term is subtracted from the corresponding Q value to yield the net heat of protonation.
Discussion We shall discuss the heat of protonation of triethylenediamine in relation to other similar amines whose thermodynamic properties have been reported previously.
It may be observed that -AH1 of this compound is lower than that of any other similar diamine. For example, one may consider the related series ethylenediamine, l6 H,N-CH2-CH2-NH~, piperazine,ac HN(CH2-CH&NH, and triethylenediamine, where each of the molecules contains the basic chain N-CH2CH2-N. The values of AH1 are -11.91, -10.17, and -7.30 kcal./mole, respectively. Triethylenediamine is a tertiary amine, and heats of protonation of these are generally lower than those of primary and secondary amines; thus, this result is not unexpected. One may compare in this context the series methylamine,' dimethylamine,' and trimethylamine, whose heats of protonation are -13.09, -11.88, and -8.86 kcal./mole, respectively. In the tertiary amine series triethylamine, trimethylamine, and triethylenediamine, the inductive effect and -AH1 decrease in the given order. The value of A& is approximately equal in all three cases, and it is likely to be due mainly to solvent-ion and solventmolecule interactions, Le., external effects,lob which do not change appreciably during the prothnations of the above amines. Thus, the diminution of enthalpy corresponds to a decrease in the strength of the N-H bond formed. The differences in the enthalpy of protonation between methylamine' and trimethylamine and between ethylenediamine18 and triethylenediamine are 4.3 and 4.9 kcal./mole, respectively; ie., the differences are approximately equal in both pairs of amines. If internal strain changed significantly during the protonation of triethylenediamine, the differences should certainly be larger than those observed. We may thus conclude that changes in internal strain are not likely to be among the major contributing factors in the observed enthalpy lowering of triethylenediamine. The decrease in -AH2 with respect to -AH1 for triethylenediamine is in part a consequence of the electrostatic repulsion between the two charged ammonium groups. Repulsion also explains the diminution of - AH1-2 for the series ethylenediamine.H22+, piperazine. HZ2+, and triethylenediamine .HZ2+as the distance between th! two nitrogens decreases from 3.9 through 2.9 to 2.5 A. in the above order of amines." It is clear that the two charges of the diprotonated diamines are not sufficiently separated from each other to behave as single charges. These charge fields overlap in accord with the observed lowering of
-
AH1-2.
(16) T. Davies, S. S. Singer, and L. A. K. Staveley, J. Chem. Soc., 2304 (1954). (17) Calculated from standard crystallographic distances and angles.
Volume 69,Number 11
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Table 11: Thermodynamic Functions of Protonation in 0.1 M KC1 at 25’: Triethylenediamine (Dabco), Triethylamine, Trimethylamine, and Ammonia”
+ + + + +
Dabco H+ DabcoH+ Dabco 2 H + F? Dabco&*+ EtaN H+G Eta”+ MeaN H+Ft Me*“+ NHa H+
PKb
- AF, kcal./mole
-AH, kaal./mole
8.82 i 0.01 11.79 f 0.02 10.75 f 0.01 9.79 f 0.01 9.29 f 0.02
12.03 16.09 14.66 13.36 12.67
7.30f 0.01 10.31 f 0.03 10.38 f 0.03 8.86 0.02 12.43i 0.10
*
A& B.U.
15.9i 0.05 19.4i 0.2 14.4 i 0.2 15.1 i 0.2 0.8 i 0.4
Uncertainty intervals associated with all values are standard deviations.
The total entropies of protonation are positive for all amines under discussion. This may be attributed to a net release of water molecules during the protonation. Certain trends as a function of amine order may also be observed. For example, the series ammonia, methylamine,’ dimethylamine,’ and trimethylamine has entropies of protonation which increase regularly by approximately 5 e.u. in the above order. The values of AS1 as well as their differences are similar in the diamine series ethylenediamine, piperazine, and triethylenediamine. The second-step protonation entropy ASZin the series ethylenediamine, piperazine, and triethylenediamine is lower than AS1. This is in accord with the interpretation of Everett and Pinsentla that the extent of hydration of a short-chain diammonium ion is more than twice that of a singly charged monoammonium ion. The observed trends of ASz in the above series may be explained by the following argument: the diprotonated forms of the amines become more rigid as a
The Journd of P h u h l Chem&yi
result of the repulsion between the two like charges. The structure stiffening’* relative to the unprotonated forms decreases from the chainlike ethylenediamine, through the cyclic piperazine, to the globular triethylenediamine. The last compound is already constrained to a rigid form prior to protonation. Lowlying vibrational states which normally contribute significantly to the free amine energy are raised as a consequence of the stiffening, to such an extent that their contribution to the entropy diminishes. This relative effect also decreases in the above order of amines.
Acknowledgments. The authors wish to thank Professor L. Sacconi for encouragement of this work and the Italian Consiglio Nazionale delle Ricerche and the American Chemical Society Petroleum Research Fund for financial assistance. (18) D. H. Everett and B.R. W. Pinsent, Proc. Roll. Sac. (London), A215,416 (1952).
