1854
J. Bron and E. L. Simmons
Equilibrium Studies of the Proton-Transfer Reactions between 2,4-, 2,5-, and 2,6-Dinitrophenols and Some Tertiary Amines in Chlorobenzene Jan Bron and E. Lynn Simmons* Department of Chemistry, University of Natal, Durban, South Africa (Received June 12, 1975: Revised Manuscript Received April 27, 1976) Publication costs assisted by the University of Natal
The equilibrium parameters for the proton-transfer reactions between 2,4-, 2,5-, and 2,6-dinitrophenol and triethyl-, tri-n-propyl-, and tri-n-butylamines in chlorobenzene were determined. The strengths of the acids were found to increase in the order 2,5-dinitrophenol < 2,4-dinitrophenol < 2,6-dinitrophenol while the strengths of the bases were found to increase in the order, tri-n-propylamine < tri-n-butylamine < triethylamine. The results were compared with some previously published experimental results for corresponding reactions in benzene and in aqueous solutions. Equilibrium deuterium isotope effects for the reactions were also measured but the isotope effects for the reactions were, within experimental error, unity.
Introduction
Interest in the equilibrium properties of proton-transfer reactions between oxygen acids and amine bases in aprotic solvents particularly became evident1,2in the 1950's and, since that time, considerable work in that area has been done.3-16 Interest in the subject has especially been revived recently because of the development of the microwave-heating5J7-19 and laser-heatinglO temperature-jump techniques, the use of which allows the measurement of kinetic parameters in aprotic solvents. Hence many of the equilibrium studies of protontransfer reactions in aprotic solvents have been carried out in conjunction with kinetic studies of the reactions. Detailed equilibrium studies of three oxygen acids with a range of alkyl amines in various aprotic solvents have been carried out. These acids are 2,4-dinitrophenol (2,4-DNP), tetrabromophenolphthalein ethyl ester (magneta E), and 2,g-dibromophenol (DBP). Proton-transfer reactions in aprotic solvents differ considerably from the corresponding reactions in aqueous solutions for several reasons. In aprotic solutions hydrogen bonding by the solvent is restricted and hence the role of the solvent in the reaction is limited to a greater extent than in aqueous solutions although solvation effects remain important. Also, the dielectric constants of many common aprotic solvents are considerably smaller than that of water and complete ionization is therefore limited. Usually the product of the proton-transfer reaction between an oxygen acid and an amine in an aprotic solvent is therefore a hydrogen-bonded ion pair. The hydrogen-bonded pair is formed through a hydrogen-bonded intermediate, the concentration of which is negligibly small compared with that of the ion pair.Q Over concentration ranges usually employed, neither the formation of separated ions nor the formation of higher aggregate quadruples appears to be i r n p ~ r t a n t .As ~ , has ~ previously been pointed out,20the effects of such reactant factors as inductive effects, steric effects, and molecular size and structure on the equilibrium properties of proton-transfer reactions in aprotic solvents are as yet not well understood. In this investigation, the equilibrium parameters of the reactions between 2,5-DNP and 2,6-DNP and triethyl-, tri-n-propyl-, and tri-n-butylamines in chlorobenzene have been determined. The corresponding reactions of 2,4-DNP have also been reinvestigated. The Journal of Physical Chemistry, Vol. 80, No. 17, 1976
The equilibrium deuterium isotope effects of the reactions were also determined. Experimental Section
Reagent grade acids were recrystallized twice from ethanol and dried before use. The amines were dried over powdered calcium hydride under vacuum for several days and then distilled. Chlorobenzene was dried several days over P205 and then fractionally distilled. The equilibrium constants were measured using the spectrophotometric method which has been1!6previously described. Since it is realized that there are inherent inaccuracies in the method, precautions were taken to eliminate as many sources of error as possible. The main precaution involved preparing all solutions by weighing stock solutions and solvents rather than using the less accurate volumetric technique usually employed.l>j-7Concentrations of the acids and bases were in the range 10-5 to M. For the spectrophotometric measurements, a Unicam SP1800 spectrophotometer with a temperature control attachment was employed. The wavelengths used were: 2,4-DNP, 400 nm; 2,5-DNP,440 nm; and 2,6-DNP, 460 nm. Temperature ranges were between 20 and 50 "C. For the isotope studies, the acids were deuterated by crystallization from deuterium oxide and the purity was checked by NMR. Since even trace quantities of water in the chlorobenzene solvent will exchange hydrogen with the deuterated acids, a different technique was used to purify the chlorobenzene for the isotope studies. The chlorobenzene was saturated with deuterium oxide and azeotropically distilled. When this method was employed using water, exactly the same results were obtained as when dried distilled chlorobenzene was used as the solvent. Results and Discussion
The logarithms of the equilibrium constants at 25 "C for the reactions studied are listed in Table I. For comparison, previous values obtained for magneta E and DBP are included.9 In Table 11,the thermodynamic parameters obtained in this investigation are listed. As can be seen, the AH" values for the three tertiary amines are very similar for each acid but differ going from one acid to the other. Previously obtained values for 2,4-DNP are given in Table I1 for c ~ m p a r i s o n . ~ , ~
Proton-Transfer Reactions between Oxygen Acids and Amine Bases
TABLE I: Equilibrium Constant Values at 25 "Cfor the Proton-Transfer Reactions between Oxygen Acids and Triethyl-, Tri-n-propyl-,and Tri-n-butylamines in Chlorobenzene
1855 A
6 -
\
log K Acid
Triethylamine
2,4-DNP 2,5-DNP 2,6-DNP Magneta Ea DBPa
4.13 3.07 5.20
Q
Tri-n-propylamine
Tri-n-butylamine
3.52 2.44 4.49 4.32 3.69
3.65 2.54 4.76 4.50 3.86
5.08
4.53
Reference 9.
0
2
4
6
log K (chlorobenzene)
B
0-
2 -
TABLE 11: Thermodynamic Parameters for the Proton-Transfer Reactions between 2,4-, 2,s-, and 2,6-Dinitrophenols and Triethyl-, Tri-n-propyl-, and Tri-n-butylamines in Chlorobenzene *
02,6-DNP
0 2,4-DNP
4 -
oz,i-DNp I
I 6
lag K (chlorobenzene)
Base Triethylamine Tri-n-propylamine Tri-n-butylamine Triethylamine Tri-n-propylamine Tri-n-butylamine Triethylamine Tri-n-propylamine Tri-n-butylamine
-AHo, kJ/mol
-ASo, J/mol deg
2,4-Dinitrophenol 52.9 f 0.1(69.02.5)b 98.4 f 0.4(152.3 8.3) 51.5 f 0.1(57.3 1.7) 105.4 f 1.3(124.34.2) 52.4 f 0.3(63.4)
106.0 f O.g(l44.6)
2,5-Dinitrophenol 95.0 f 1.2 45.8 f 0.4 100.2 f 2.2 43.8 f 0.7 43.3 f 0.6
96.6 f 1.9
.2,6-Dinitrophenol 60.4 f 0.5 103.1 f 1.7 61.1 f 0.3 119.1 f 0.9 65.6 f 0.6
129.1 f 2.0
Uncertainties are standard deviations. Number of runs per reaction, 6-11. Numbers i n parentheses are values obtained in ref 7. The reason for the large disagreement between values obtained in this investigation and those obtained by Ivin, McGarvey, Simmons, and Sma116,7is not clear. It should, however, be pointed out that it is reasonable to suppose that AH" should not vary significantly going from triethylamine to tri-n-butylamine since the strengths of the nitrogen-hydrogen bond with the oxygen should not differ greatly for the three amines. In contrast with the values obtained in this investigation, the AHo values obtained by Ivin, McGarvey, Simmons, and Small vary erratically going from triethylamine to tri-n-butylamine. According to the data in Table 11, the variations of the equilibrium constant for a given acid going from triethylamine to tri-n-butylamine are mainly due to variations in ASo rather than to variations in AH".It is also interesting to note that the AH" values for 2,4-DNP as acid in Table I1 are very close to that obtained for quinuclidines (-51.2 kJ/mol) which, because the alkyl groups are "tied together" and hence steric factors are smaller, has a much higher equilibrium constant than those of the amines listed in Table 11. Hence the previously discussed13 steric factors involved in
Figure 1. Plots for the reaction between triethylamine and dinitrophenols at 25 O C : (A) log Kin benzene vs. log K i n chlorobenzene; (6) log Kin chlorobenzene vs. -pKa in water.
the proton trahsfer reactions appear to affect A S o and not AHovalues. (ASo for the quinuclidine reaction with 2,4-DNP is -73.93/deg mol, much larger than those in Table 11.) Since it is customary to discuss acid-base behavior in terms of log K the remainder of the discussion is in terms of this parameter. The values for log K going from one acid to the others illustrate that the stability of the ion pair increases in the order 2,5-DNP < 2,4-DNP < 2,6-DNP. This is the same order obtained previously for the corresponding reactions (with triethylamine as base) in benzene solutions.l6 The strengths of the acids in aqueous solutions also increase in the same order.16 However, as can be seen in Figure 1,there appears to be a quantitive relationship between the log K values for the reactions in chlorobenzene and benzene, while no similar quantitive relationship between log K for chlorobenzene as solvent and pK, for aqueous solutions is observed. Similarly no such quantitive relationships between the reactions in benzene and in aqueous solutions were previously observed.16 This was explained on the basis of the differences between the acids in water and in aprotic solvents. In water, solvent molecules may hydrogen bond to the unassociated acid while there is ample e v i d e n ~ e l ~ ,that ~ l -in ~ ~aprotic solvents, the phenolic hydrogen is hydrogen bonded to the o-nitro group. In view of these differences no quantitive correlations between the acid strengths in aqueous solutions and in aprotic solvents are expected. The significance of the order of the strengths of the acids is discussed below. The formation of the hydrogen bonded ion pair, BH+AHA
K
+B
e BH+
-A-
(1)
where HA is the acid and B the amine, can be described in terms of the following three hypothetical reactions: KA
HA+ H+
H+
+ A-
KB +B$ HB+
(2) (3)
The Journal of Physical Chemistry, Vol. 80, No. 17, 1976
1856
J. Bron and E. L. Simmons
+
and A-
+ HB+
Ki
BHf
-
*
A-
(4)
Here it is understood that solvation of the ions occurs and may play an important role. Reactions 2 and 3 have equilibrium constants which describe the acid and base strengths in the traditional ~ e n s e , 2while ~ , ~ ~reaction 4 is an ion association reaction. It is clear from eq 1-4 that
K = KAKBK~
(5)
or log K = log K A
+ log K B + log Ki
(6)
The theory of ion association reactions has been treated in some detail by several a ~ t h o r s . ~ ~The - ~ Oexpression for the equilibrium constant of such a reaction for spherical ions in a medium with a dielectric constant, E , is
Ki = 4nrAB26r exp( WIRT)
(7)
where rAB is the sum of the radii of the two reacting ions, 6r the distance the centers of the two ions can move apart and the ions still remain associated, and
w = Noe 2/€rAB
(8)
for singly charged ions, where N Ois Avogadro's number. It is clear that eq 7 cannot apply to reaction 4 but must be modified in two ways.20 First, because of the formation of a hydrogen bond, W must be expressed as
w = Noe2/crAB+ AGH
(9)
where AGH is the free energy change per mole for the hydrogen bond formation. Second, for the large molecules involved whose surface reactivity is limited, 4mAB26r,must be replaced by a statistical volume,3° u. Theoretically, r j can probably best be described in terms of the sizes of the actual reacting atoms or groups of atoms along with the introduction of a geometric factor, P , of the type suggested by Grunwald3I and which has been successfully employed p r e v i ~ u s l to y ~account ~ ~ ~ for the hindrance of movement caused by groups attached to the reacting atoms. Hence rj
= P4T(?'H
+ ro)2a
(10)
where the subscripts refer to hydrogen and oxygen and 6r has been replaced by a which, since A and B are held together by a hydrogen bond, is a type of vibrational amplitude. Further, it is reasonable to assume that P is the product of the fractions of the surface areas of the reactant groups which are reactive, ^ I
VI
P = PAPB where PA and PBrepresent the respective fractions. The parameter, log K , may now be expressed as log K = log K A + log K B + log PA+ log PB log 4 T ( r H + r d 2 a AGH/RT Noe2/crABRT (12). The term, AGH,probably varies only very little for the different acids and bases in Table I. In particular, it is expected to remain the same for the three bases and a single acid. Since for all three bases the positive charge in the ion pair is centered on nitrogen atom, r A B should remain the same going from triethylamine to tri-n-butylamine for a given acid. Since the negative charge in the ion pair is, however, delocalized over the ring of the acid, rAB should vary going from one acid to
+
+
+
The Journal of Physical Chemistry, Vol. 80. No. 17, 1976
another. It is therefore clear that if the term, log 4 ~ ( rro)2a, ~ is considered constant, then all other terms in eq 1 2 depend either upon the acid or upon the base involved in the reaction. In other words, if the above analysis is valid, then log K = f(A)
+ g(B) + h
(13)
where f(A) is a function only of the acid, g(B) a function only of the base, and h a constant for a series of reactions. The form of eq 13 may be tested using the data in Table I by choosing an arbitrary scale for f(A) such that f(A) = 1for the weakest acid, 2,5-DNP, and selecting values of f(A) for the other acids so that A log K/Af(A) = 1. Naturally, this procedure results in a straight line if log K vs. f(A) is plotted for a single base but it does not guarantee that such curves will be linear on the same scale for all bases unless the theoretical reasoning used to develop eq 13 is correct. In Figure 2, log K is plotted against f(A) for the three bases and five acids using the data in Table I. As can be seen, the data for all three bases form very good straight lines for the same selected values of f(A) for the five acids. As illustrated in Figure 2, trimethylamine9 and three of the acids also form a straight line on the same scale but the slope is different, indicating that this base cannot be included in the same series. Trimethylamine is a weaker base because of smaller inductive effects of the methyl groups. If the values of f(A) may be taken as measures of the acidities of the acids in chlorobenzene and the intercepts (= g(B) h ) taken as measures of basicities of the bases, then the acidities increase in the order 2,5-DNP < 2,4-DNP < DBP < magneta E < 2,6-DNP, and the basicities increase in the order tri-n-propylamine < tri-n-butylamine < triethylamine, as is, of course, also indicated by the relative magnitudes of the various equilibrium constants. The order of the acidities of the acids may be explained mainly on the basis of resonance stabilization and inductive effects of the ring substituents. Hence, 2,4-DNP is a stronger acid than 2,5-DNP because of the resonance stabilization of the nitro group in the para position. The exceptionally strong acidity of 2,6-DNP, cannot, however, be explained on this basis since resonance stabilization should be the same as that for 2,4-DNP. Although steric arguments could be invoked to explain the large acidity of 2,6-DNP in that the bulky nitro groups crowd the hydrogen atom thus facilitating its removal, the major cause of the large acidity is probably due to electrostatic factors. With the electron attracting nitro groups on the end of the ring facing the positive charge in the ion pair, the negative charge is undoubtedly localized nearer to the positive charge for 2,6-DNP than for 2,4-DNP. Hence, the ion pair is more stable for 2,6-DNP than for 2,4-DNP. The larger negative AHo value for the former acid reflects the larger electrostatic energy. The same arguments may be applied to explain the relatively large acidity of 2,6-DNP, this acid being weaker than 2,6-DNP because the bromine groups do not stabilize the negative charge so well as the nitro groups. Magneta E is a stronger acid than DBP because it has the additional group attached to the para position which may stabilize the negative charge through resonance. The order of the basicities of the amines cannot be explained so readily. As previously indicated, it has been assumed that the inductive effects of the alkyl groups do not play a major role in going from triethylamine to tri-n-butylamine. Instead, steric factors are probably more important and, hence, the major differences in the values of the intercepts in Figure 1are probably due to the term log PB.On this basis the fact that log K for triethylamine is greater than that
+
Proton-Transfer Reactions between Oxygen Acids and Amine Bases
1857
since the presence of a hydrogen bond stabilizes a species by 20-30 kJ/mol. In summary, the major conclusions drawn from this study are that the logarithm of the equilibrium constant of the proton transfer reaction of an oxygen acid and a tertiary amine in chlorobenzene can be expressed as the sum of a term dependent only on the base, a term dependent onlyon the acid, and a constant term (for a given series) and that acid strengths of oxygen acids in aprotic solvents can generally be explained on the basis of resonance stabilization by ring substituents and inductive effects.
r M i
Acknowledgments. Financial support by the Atomic Energy Board (J.B.) and the Council for Scientific and Industrial Research (E.L.S.) is gratefully acknowledged. We thank S. Webster for help in this investigation.
References and Notes II 1 2,5-DNP
2 2,4-DNP
I
3
DBP
Magneta E 2,6-DNP
f (A), Arbitrary Scale
Figure 2. Plots of log Kvs. f(A) for the proton-transfer reactions between some oxygen acids and tertiary amines in chlorobenzene at 25 O C : (0) triethylamine; ( 0 )tri-n-propylamine; ( 0 )tri-n-butylamine; ( 0 )trimethylamine.
for the other two amines is readily explained. However, neither this nor any other previous argument can explain the fact that tri-n-butylamine is a stronger base than tri-n-propylamine. This anomaly for tri-n-propylamine, which is also apparent in the kinetics of the reactions, has been discussed but no satisfactory explanation has yet been produced.9 Finally, it should be pointed out that no equilibrium deuterium isotope effects within the experimental error of the methods used were observed. This indicates that the environment of the proton in the reactant acid does not differ greatly chemically from that in the ion pair. Likewise, no deuterium equilibrium isotope effects were observed for reactions between triethylamine and 2,4-DNP in chlorobenzene and toluene solutions.33Bell15 points out that the stretching frequency of the N-H bond in the ion pair in the absence of hydrogen bonding is very close to the 0-Hfrequency in the phenols and hence, if there is no hydrogen bonding, no isotope effect would be expected. An alternate explanation, in view of the e ~ i d e n c e l ~ that J - ~ there ~ is an internal hydrogen bond in the o-nitrophenols, is that there is a hydrogen bond in the ion pair and that the stretching frequedes for N-H- - -0and 0- - -H-0 are similar. This explanation appears reasonable
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(16)M.M.'Davis, J. Am. Chem. SOC.,84, 3623 (1962). (17)0. Ertl and H. Gerischer, 2.Elektrochem., 65, 629 (1961). (18)G. Ertl and H. Gerischer, 2.Elektrochem.,66, 560 (1962). (19)E. F. Caldin and J. E. Crooks, J. Sci. Instrum., 44, 449 (1967). (20)E. L. Simmons, Prog. React. Klnet,, in press. (21)P. M. Boll, Acta Chem. Scand., 12, 1777 (1958). (22)J. H. Richards and S . Walker, Trans. Faraday SOC., 57, 406 (1961). (23)J. A. Davidson, J. Am. Chem. SOC.,67, 228 (1945). (24)G. E. Hibert, 0. R. Wulf, S. R. Hendricks, and U.Liddei, J. Am. Chem. Soc., 58,548 (1936). (25) H. N. Alyea, J. Chem. Educ., 18, 206 (1941). (26)J. R. Jones, Prog. React. Kinet., 7, l(1973). (27)J. E. Prue, J. Chem. Educ., 46, 12 (1969). (28)N. Bjerrum, K. Dan. Videskab. Selskab, Mat.-Fys. Medd., 7, No. 9 ( 1926). (29)R. M. Fuoss, J. Am. Chem. SOC.,80, 5059 (1958). (30)M. Elgen, 2.Phys. Chem. (FrankfurtamMain), 1, 176 (1954). (31)E. Grunwald in "Progress in Physical Organic Chemistry", Vol. 111, Cohen, Streltweisser, and Taft, Ed., Interscience, New York, N.Y., 1965. (32)E. L. Simmons, 2.Phys. Chem., 96, 47 (1975). (33)R. P. Bell and J. E. Crooks, J. Chem. SOC.,3513 (1962).
The Journal of Physical Chemistry, Vol. 80, No. 17, 1976
