CHLOROPHENOL WITH ALIPHATIC AMINEB
587
Hydrogen Bonding Interactions of p-Chlorophenol with Aliphatic Amines by Mei-Lan Lin and Ronald M. Scott* Bepartment of Chemistry, Eastern Michigan University, Ypsilanti, Michigan 48157
(Received June SO, 1571)
Publication costs borne completely by The Journal of Physical Chemistry
The formation of hydrogen bonds between p-chlorophenol and a series of aliphatic amines in cyclohexane revealed that tertiary amines react less readily than would be expected from their aqueous pK. This was interpreted as reflecting steric hindrance by the alkyl groups of the tertiary amines.
Introduction The method of spectrophotometric analysis for the study of hydrogen bonding by phenols may first have been utilized by Morton and Stubbs in an investigation of hydroxybenzaldehydes and their methyl ethers. Burawoy, et al.,2-4 made a systematic study of the method to investigate the tautomeric equilibria of phenolic compounds. Other workers5-12 employed spectrophotometric methods to study hydrogen bonding by phenols. The thermodynamics of the phenol hydrogen-bond-formation reaction have been measured for sevcral case~.5--8,13-~6 p-Chlorophenol, the phenol selected for this study, is a weak acid (aqueous pK = 9.39),16with absorbance bands in the ultraviolet region at 283 mp and 289 mp, The previously reported thermodynamic parameters for its hydrogen bond formation are summarized in Table I. Cyclohexane, the solvent in these studies, is a hydrocarbon with a dielectric constant of 2.02 which does not interact either with acidic solutes such as phenols or with basic solutes such as amines. l7 The data reported here were obtained as part of a larger study concerning the types of interaction between phenols and amines. A number of phenols and of amines are being investigated with consideration given to the effects on interaction of the acidity of the phenol, the basicity of the amine, and the dielectric constant of the solvent. The results are interpreted in terms either of formation between phenols and amines of hydrogen bonded complexcs or hydrogen bonded ion pairs, or of the complete transfer of the proton to form phenolate and ammonium ions. Some results are published,s~18~10 and the significance of these results for understanding interactions occurring in biological systems has been considered.20f21
Experimental Section p-Chlorophenol and morpholine were Dow Chemical Company reagont grade. n-Butylamine, N-ethylmorpholine, tri-n-butylamine, and triethylamine were Aldrich Chemical Company reagent grade. Di-nbutylamine, isopropylnmine, and tri-n-propylamine were Eastman Organic Chemical reagent grade. 1,4-
Dioxane was Matheson Coleman and Bell Company Spectrograde. Cyclohexane was Baker Analyzed reagent grade. All reagents except dioxane were redistilled before use. Triethylenediamine was obtained from Houdry Process and Chemical Company, and was purified by crystallization in dry ether twice and sublimation in vacuum, and was thereafter stored in a desiccator. pChloropheno1 and amine stock solutions were prepared by weighing both solvent and solute. Aliquots were then pipetted into volumetric flasks and diluted with cyclohexane, the flask being reweighed after each addition. In each study the amine concentration was varied while the p-chlorophenol concentration was kept constant. The absorption spectra were obtained with R. A. Morton and A. L. Stubbs, J.Chem. Soc., 1347 (1940). A. Burawoy and I. Markowitsch-Burawoy, {bid., 36 (1936). A. Burawoy and J. T. Chamberlain, ibid., 2310 (1952). A. Burawoy and J. T. Chamberlain, ibid.,3734 (1952). (5) 8. Nagakura and H. Bada, J . Amer. Chem. Soc., 74,5693 (1952). (6) 8. Nagakura, ibid., 76,3070 (1954). (7) L. Bellon, C. R. Acad. Sci., 254,3346 (1962). (8) R. A. Hudson, R. M. Scott, and S. N. Vinogradov, Spectrochim. Acta, Part A , 26,337 (1970). (9) L. Bellon, Trav. Inst. Sci. Cherijien Ser. Sci. Phys., 6, (1960). (10) M . Bonnet and A. Julg, J . Chim. Phys., 59,723 (1962). (11) N. S. Coggeshall and E. M. Lang, J . Amer. Chem. Soc., 70, 3283 (1948). (12) 8. Nagakura a n d M . Gouterman, J . Chem. Phys., 26,881 (1957). (13) R. L. Denyer, A. Gilchrist, J. A. Pegg, J. Smith, T. E. Tomlinson, and L. E. Sutton, J . Chem. Soc., 3889 (1955). (14) D. Gurka, R. W. Taft, L. Joris, and P. von R. Schleyer, J . Amer. Chem. Soc., 89,5957 (1967). (15) E. M. Arnett, T. 8. S. R. Murty, P. von R. Schleyer, and L. Joris, ibid., 89, 5955 (1967). (16) I. Avigal, J. Feitelson, and M. Ottolenghi, J . Chem. Phys., 50, 2615 (1969). (17) A. I. Shatenshtein, “Isotopic Exchange and the Replacement of Hydrogen in Organic Compounds,” (Translation of 1960 Russian edition by C. N. Turton and T . I. Turton) Consultants Bureau, New York, N. Y., 1962. (18) R. Scott, D. De Palma, and 9. Vinogradov, J . Phys. Chem., 72, 3192 (1968). (19) R. Scott andS. Vinogradov, J . Phvs. Chem., 73,1890 (1969). (20) R. A. Hudson, R. M. Scott, and S. N. Vinogradov, Bwchim. Biophys. Acta, 181,353 (1969). (21) S. N. Vinogradov, R. A. Hudson, and R. M . Scott, ibid., 214, 6 (1970). (1) (2) (3) (4)
The Journal of Physical Chemistry, VoZ. 78, N o . 4, 1578
588
MEI-LANLIN AND RONALD M. SCOTT
Table I : Previously Reported Thermodynamic Parameters for Hydrogen Bond Formation with p-Chlorophenol AH, koal/mol
K
Proton acoeptor
Methyl acetate Piperidine Triethylamine Tetrahydropyrane Trimethylamine Dioxane
16.9, 18.3 410 (20') 89 (20') 44 (200) 76
AG, kcal/mol
-6.7 -9.0 -9.55 -6.4 -7.4, - 8 . 6
--I-' -'
-'
-'
-:
Ref
A& eu
-1.7 -4.86 -4.0 -3.8 -3.1 - 1 . 6 (23')
-16.9 -14.2 -18.9 -9.7 -13.0
6
7 7 7 13 15
jr
.. -'
.'
1
> iOP(61
Figure 2. Plots of the log of the ratio of the concentralions of hydrogen bonded to norihydrogen bonded p-chlorophenol (log BHA/HA) us. log base concentration (log B ) calculated from the spectrophotometric data. The bases and the slopes of the lines are: A, dibutylamine 1.1; B, n-butylamine 1.0; C, dioxane 1.0; D, morpholine 1.2; E, triethylamine 1.1; F, tributylamine 1 .O; G, AI-ethylmorpholine 1.O; H, n-propylamine 1.1; I, isopropylamine 1.2; J, tripropylamine 1.2.
?so
m
wu"oIerql/II( "I
Figure 1. Reproductions of the spectra obtained by scanning cyclohexane solutions of p-chlorophenol t o which varying concentrations of amine have been added at 25'. Upper represents 2.10 X 10-3 LVp-chlorophenol with additions of morpholine from 0 (lowest absorbance at 285 mp) to 4.6 X 10-2 M (highest absorbance at. 285 mp). Lower represents 2.12 10-8 M p-chlorophenol with additions of N-ethylmorpholine from 0 to 4.0 X 10-3 M (same order as upper).
a Beckman Model DEI-2A spectrophotometer using matched 1-cm silica cells. The temperature of both sample and reference cells was controlled by circulating water from a Lauda K-2/R constant temperature bath through the thermospacers of a brass cell holder similar to that described by Coggeshall and Lang." The cells were capped to prevent solvent loss. The spectra of the solutions were recorded from.240 to 315 mp. Reference cells cancelled the absorbance of both solvent and amine. Each thermodynamic run consisted of at The Journal of Physical Chemistry, Vol. 76, N o . 4, 1972
least 5 different p-chlorophenol : amine ratios at 5 temperatures ranging from 15 to 55'. All calculations of equilibrium and thermodynamic parameters were performed on the IBM 1130 digital computer of Eastern Michigan University. Programs were prepared in Fortran IV which convert data in terms of weight and density into molarities, then using these molarities and absorbance readings determine equilibrium constants by the calculation of Rose and Drago,22 Corrcctions were made for overlapping absorbance bonds in the spectra. Final equilibrium constants obtained in this fashion were tested for their statistical validity using the criterion of ChauvenetZ3applying this test twice to thc data for each series of amine additions. The remaining constants were averaged. Thermodynamic parameters wcre obtained from average log K values at various temperatures in the usual fashion.24 (22) N. J. Rose and R . S . Drago, J . Amer. Chem. Soc., 81, 6138 (1959). (23) H. D. Young, "Statistical Treatment of Experimental Data," McGraw-Hill, New York, N. Y . , 1962. (24) A copy of this program is available from the author on request.
CHLOROPHENOL WITH ALIPHATICAMINES
589
Table IT: Experimental Conditions, Equilibrium Constants, and Wavelength Shifts for the Formation of 1:1 Hydrogen Bonded Complexes by p-Chlorophenol in Cyclohexane a t 25' Concentration Base
Baae pK,"
Dioxane iV-Ethylmorpholine Tri-n-but ylamine Tri-n-prop y lamine Morpholine Triethylenediamine Triethylamine %-Butylamine Isopropylamine Di-n-butylamine
Concentration range of base x 103 M
Equilibrium constant (25')
2.16
6-18.9
5.42 X 10
2.12 2.14 2.13 2.10 2.15 2.18 2.16 2.17 2.15
0.5-40 0.2-15 2.4-50.6 0.6-46 0.3-15 0.3-20 0.8-25 1.6-16 013-15
p-chlorophenol x 103 M
-4.3 in HOAc 2.92in HzSO~ 7.67 9.93 10.26 8.33 8.60b 10.75 10.68 10.63 11.25
1.77 X 2.45 X 2.64 X 5.24 X 7.04 X 1.55 x 1.89 x 2.42 x 2.76 X
lo2 lo2 102 lo2 lo2 103 103 103 loa
Log K
1.72
2.0
2.15 2.09 2.42 2.65 2.84 3.09 3.24 3.38 3.43
3.5 4.0 4.3 4.3 4.5 4.7 4.8 5.0 5.3
a D. D. Perrin, "Dissociation Constants of Organic Bases in Aqueous Solution,'J Butterworths, London, 1966. Process and Chemical Co., Division of Air Products and Chemicals, Inc., Philadelphia, Pa.
~
Wavelength shift (mfi) peak
DABCO, Houdry
~
Table 111: Experimental Conditions, Equilibrium Constants, and Thermodynamic Parameters for the Formation of 1 : I Hydrogen Bonded Complexes of p-Chlorophenol in Cyclohexane
Baae
Dioxane n-Butylamine
p-Chlorophenol concn X 103 M
Base concn X 108 M
Wavelength, mfi
-
15'
-AH,
K , 1. mo;-l 25'
350
450
55O
52.81 44.21 35.07 2.1-2.2 6.2-190 291.5 82.42 52.72 292.5 75.47 44.95 41.47 32.45 56.46 289 79.11 56.29 48.65 31.87 20.76 1728 1603 786.4 426.3 2.1-2.2 0.81-25 285 2156 2037 2036 1392 293 722.8 548.9 2039 2009 1374 295 784.4 509.4
Results Spectrophotometric studies of hydrogen bonding were performed in cyclohexane with p-chlorophenol and the following amines : n-butylamine, morpholine, din-butylamine, tri-n-butylamine, N-ethylmorpholine, triethylamine, isopropylamine, triethylenediamine, and tri-n-propylamine. Hydrogen bonding with dioxane was also studied for the comparison of an oxygen structure with the nitrogenous bases. Concentrations of M and 5-8 adp-chlorophenol were 2.10-2.18 X ditions of amine were made in each case ranging in concentration from about 0.1 to 25 times the concentration of the p-chlorophenol. All spectra obtained had clear isosbestic points inferring that a simple equilibrium was occurring (Figure 1). The ratio of hydrogen bonded to nonhydrogen bonded form was estimated from the absorbances a t 287 and 295 mp. The information was then plotted using log amine concentration as the x axis and log of ratio of hydrogen bonded to nonhydrogen bonded concentration as the y axis. The slope of such a plot indicates the number of amines associating with each p-chlorophenol, and in each case it provided a clear indication of a 1 : 1 complex (Figure 2). The peaks of the p-chlorophenol absorption spectrum shift toward the visible range upon the formation of the
koal/mol
3.54 zt 0.6 3.76 f 0.3 3.42 f 2.4 7.50 i 1.6 7.23 f 1.2 6.83 f 1.2
- A@, kcal/mol
2.44 f 0.04 2.41 f0.02 2.13 f 0.17 4.39 & 0.11 4.37 =t0.09 4.36 f 0.09
-AS, cal/ deg
3.71 4.54 4.33 10.44 9.60 8.30
hydrogen bond. The magnitude of the shift is not the same for each amine considered. The data obtained for the spectral shift of the 281 mp peak are recorded in Table 11. Calculations of t>heequilibrium constant for the formation of the hydrogen bond were performed on data read at three wavelengths and calculated both as separate and as combined data. The values thus obtained for K and log K at 25' are reported in Table 11. For the interactions of p-chlorophenol with n-butylamine and with dioxane the values for log K were obtained at several temperatures. A H , AG, and AS were calculated by use of the computer program earlier described, and the results are presented in Table 111.
Discussion The study was restricted to alkyl amines for a practical rather than a theoretical reason. Aromatic amines are not transparent in the region of the spectrum of interest in this investigation. An interesting pattern emerges as the data are considered. Simple hydrogen bonding is occurring in each case, the conclusion to be drawn from the indication of a 1 : 1 interaction (slopes of Figure 2) and from the clear isosbestic points (Figure 1). The thermodynamic calThe Journal of Physical Chemistry, VoL 76, No. 4, 1972
590
culations provide values for AH and AG that compare well with those obtained previously for other phenolamine hydrogen-bond formation reactions. The magnitudes of the wavelength shifts (Table 11) are seen to correlate with the equilibrium constants of the hydrogen-bond formation reaction. Such shifts should reflect the perturbation of the donor eIectronic structure given by the formation of the hydrogen bond, hence the bond ~ t r e n g t h . ‘ ~ That ~ ~ ~these shifts do not correlate with the aqueous pK values of tlic amines as recorded in the literature is cause for speculation. When a single amine is reactcd with a variety of phenols the resulting values for log K for the reaction were found to be proportional to the aqueous pK of the phenols by Bellon9 using triethylamine in cyclohexane, by Denyer, et al., l 3 with the same phenols and trimethylamine, and by Bonnet and Julg’O using triethylamine, n-butylamine, and di-n-butylamine in heptane. Gordonz6summarized a linear relationship between infrared stretching frequency shifts and the aqueous pK for bases of similar structure. I n our data comparing now a single phenol with various amincs the physical characteristics measured by the aqueous pK of the amine did not explain completely the behavior of the system. A plot of log K for each interaction us. aqueous pK for the particular amine (Figure 3) emphasizes that a linear relationship does in fact exist between these parameters for primary and secondary amines. Conventional tertiary amines (tri-n-butylamine, tri-n-propylamine, triethylamine) show less tendency to form the hydrogen bond than their aqueous pK would indicate. We propose that the three alkyl groups of the tertiary amine sterically hinder the interaction of amine and phenol. I n support of this proposal we point out that the deviation from the “unhindered” behavior of the primary or secondary amines becomes greater as the length of the alkyl groups becomes greater (butyl > propyl > ethyl). Furthermore in the case of N-ethylmorpholine the deviation from “unhindered” behavior is less due to the folding back of two of three alkyl substitutions by ring formation. Finally when we consider the case of a tertiary amine all of whose alkyl substitutions are folded back into rings, triethyIenediamine, we find that it falls on the line formed by the primary and secondary amines in the plot. Coggeshall27 studied the effect steric hindrance by large ortho alkyl groups on the phenol has on self association by hydrogen bonding. The pcak at 3.0 1.1 characteristic of the phenol hydroxyl groups shifts toward shorter wavelengths with the formation of the hydrogen bond, and the magnitude of the shift was used as a measure of the strength of the hydrogen bond formed.z8 That those phenols having large ortho substituents were hindered in forming hydrogen bonds by self association was indicated by the occurrence of shifts of less than 0.04 p as contrasted with expected shifts of over 0.15 p . This evidence of hindrance correlated with the The Journal of Physical Chemistry, Vol. 76, N o . 4, lQ72
MEI-LANLIN AND RONALD M. SCOTT
4.0-
/
la5 1
.o
t’,
8
9
10
T r i propy Iamlne
11
12
p K a o f Amine
Figure 3. Plot of the log of the equilibrium constant for the formation of the hydrogen bond between p-chlorophenol and the indicated amines at 25” vs. the aqueous p K , of the amine.
general loss of phenolic character observed for these compounds exemplified by inability to react with sodium metal in ether solution, insolubility in aqueous alkali, and nonreactivity with alcoholic ferric Parallel work was done later’’ with ultraviolet spectroscopy using the same series of hindered and unhindered phenols. A shift of absorbance peaks toward longer wavelengths was observed whcn comparing the spectra of solutions in the hydrogen bonding solvent ethanol to those taken in isooctane. The spectral shift for hindered phenols was markedly less than for unhindered phenols. In our studies the degree of spectral shift correlated roughly with the log of the equilibrium constant calculated for the formation of the hydrogen bond (Table 11). Studies relating to the steric effects of amine structures in the formation of hydrogen bonds between phenols and amines are not available. B r o ~ n ~ O has. ~ ~ studied the steric characteristics of tertiary amines, however. He noted that in accepting a proton in aqueous solution trimethylamine is a weaker base than tri(25) A. Julg and M. Bonnet, Theoret. Chim. Acta, 1 , 6 (1962). (26) J. E. Gordon, J.Org. Chem., 26,738 (1961). (27) N. D. Coggeshall, J . Amer. Chem. SOC.,69,1620 (1947). (28) L. Pauling, {bid., 58, 94 (1936). (29) G. H. Stillson, D. W. Sawyer, and C. K. Hunt, ibid., 67, 303 (1945). (30) H. C. Brown and M. D. Taylor, ibid., 69,1332 (1947). (31) H. C. Brown and S. Sujiski, ibid., 70,2878 (1048).
591
LEGHEMOGLOBIN WITH NITROGEN AND WITH XENON ethylamine, yet in the vapor state trimethylamine can form a complex with trimethylboron while triethylamine cannot. Support for the interpretation that failure to bond reflects blocking of the binding site on nitrogen was provided when he showed that quinuclidine does form a stable complex with trimethylboron. Earlie? Perrin and Williams demonstrated that at 80" trimethylamine reacts 61-fold faster than triethylamine with isopropyl iodide. These results all suggest that all three of the alkyl groups of triethylamine are not accommodated to thc rear of the molecule and are supportive of our findings. Dioxane was studied as a nonnitrogenous proton acceptor in the hydrogen bonding interaction for contrast. The results are in general agreement with other such studies performed with phenol and p-nitrophenoL5z8 It was assumed because of the slope of approximately 1 obtained for the dioxane interaction (Figure 2) and because of the excess of dioxane over p-chloro-
phenol in that study that the reaction of two phenols per dioxane reported by Baba and NagakuraSawas not occurring.
Summary A detailed study of the hydrogen bonding interactions of p-chlorophenol with various nitrogenous proton acceptors and with dioxane was performed. The thermodynamics of the interaction with n-butylamine and with dioxane were studied, The relationship between the aqueous pK of the amines and the log of the equilibrium constant for the formation of the hydrogen bond was studied. It is proposed that this relationship is linear unless the alkyl groups of the amine hinder the interaction. This occurs with tertiary amines having normal alkyl groups. (32) M.W. Perrin and E. G . Williams, Proc. Roy. SOC.,Ser. A , 159, 162 (1937). (33) H.Baba and S. Nagakura, J . Chem. SOC.,Jap., Pure Chem. Sect., 72,3 (1951).
The Interaction of Leghemoglobin with Nitrogen and with Xenon' by Gordon J. Ewing" and Lavinel G. Ionescu Department of Chemistry, N e w Mexico State University, L a s Cruces, N e w Mexico 88001
(Received J u l y 8, 1071)
Publication costs assisted by the National Science Foudation
The absorption of nitrogen in solutions containing the main components of leguminous hemoglobin has been determined a t 5, 15, and 25" in the 0-5 atni pressure range. Kitrogen interacts with both ferri- and ferroleghemoglobin in an approximate 1:1 molar ratio, The thermodynamic quantities for these interactions were determined. For the interaction of nitrogen and ferrilegoglobin a t 15", AGO = -4.06 k 0.03 kcal/mol, AHo = +15.7 f 0.1 kcal/niol, and AS" = 69 1 eu. For the nitrogen ferrolegoglobin interaction a t 15", AGO = -3.8 i 0.21 kcal/mol, AH" = +13 =t1 kcal/mol, and ASo = 57 =k 2 eu. For both interactions ACp was very large, about -1350 cal/(mol deg). The large positive entropy changes, particularly a t lower temperatures, suggest a considerable disordering of the leghemoglobin molecule during the absorption of nitrogen. A similar study of the interaction of xenon with leghemoglobin was attempted. While evidence for a weak interaction was obtained, the experimental errors were so large in comparison to the interaction that the stoichiometry could not be determined.
*
Introduction Leguminous hemoglobin, legoglobin or leghemoglobin, is the respiratory pigment found in the root of most leguminous plants. The hemoglobin nature of the red pigment was first discovered-by Kubo2 in 1939 and was confirmed later by the work of Virtanen3f4 and of Keilin and Wang.5 Leguminous hemoglobin is always foulld in the interior of nodules growing the roots and in association with bacteria, usually of the genus Rhizobium legunzinosarum. Purification of leghemoglobin usually yields two main components Of
slightly different molecular weight with values ranging ".Ind ~ 6 ~ o ~ ~ - ~ 7 ~ o o o ~ 6 - s (1) This work was partially supported by grant GB 7829 from the National Science Foundation. (2) H. Kubo, Acta Phytochim., 11, 195 (1939). (3) A. I. Virtanen, Sitzungsber. F i n n . A k a d . Wias., 12, 1 (1945). (4) A. I. Virtanen, Nature (London), 155, 747 (1945). (5) D. Keilin and Y . L. Wang, ibid., 155, 227 (1946). (6) N.Ellfolk, Acta Chem. Scand., 14,609 (1960). (7) N.Ellfolk, ibid., 14, 1819 (1960). (8) N.8. Subba Rao and C. L. Chopra, J. Sci. I n d . Res., 26 (8), 329 (1967);Chem. Abatr., 67, 113616f (1967). T h e Journal
of
Physical Chemistry, Vol. 76, N o .
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