� －15－Kinetics and Mechanism of the Slow Protolysis of Pyridine and 1,10-Phenanthroline Derivatives（K. Itoh et al.）福岡大学理学集報　45�⑴　15～20（2015）

Introduction

A proton transfer is an elementary step and affects many reactions in aqueous solution, e.g. a complex forma-tion, an enzymatic reaction, and an acid and base catalyzed reaction. A proton transfer is of an importance in under-standing those reactions. Thus, the protolysis has been in-vestigated on a wide variety of compounds.[1, 2] It is well known that the protolysis of normal nitrogen and oxygen acids is fast and their protonation reaction is diffusion con-trolled. However, it was reported that the carbon acids un-derwent relatively slow protolysis and the rate constants of protonation for their conjugate bases were 102-104 mol-1dm3s-1.[3, 4]

We measured the rates for the protolysis of pyridine (py), 2,2’-bipyridine (bpy) and 1,10-phenanthroline (phen) in non-buffered aqueous solution.[5] We found that their protolysis proceed through the two concurrent paths; one was the proton transfer step with water molecules of solvent and the other was the proton transfer step with hydroxide ion. Their rate constants were comparable to those of cy-clopentadienes belonged to carbon-acids category. So far it is not cleared out why the protolysis of the aromatic nitro-gen acids like pyridine and 1,10-phenanthroline is very slow.

The protolysis of carbon acids has been investigated for a long time. It is well known that its main retarding factor in their slow protolysis are (i) the small hydrogen bonding abil-ity with solvent, (ii) the electronic and structural reorganiza-

tion on going from the acid to its conjugate base and (iii) the accompanying reorientation of solvation.[2] However, it is difficult for the reason of (i) and (ii) to be the retarding fac-tor in the protolysis of py, bpy, and phen. On the other hand, it was reported that the formation of an intramolecu-lar hydrogen bonding might retard the proton transfer in the protolysis of a nitrogen and an oxygen acid.[6, 7] The forma-tion of an intramolecular hydrogen bonding may take place in the conjugate acid of bpy and phen but not in that of py. Thus it is hard to interpret the slow protolytic behavior of bpy and phen similar to that of py in term of the formation of an intramolecular hydrogen bonding.

In the present study, in order to clarify the mechanism for the slow protolysis of the aromatic compounds involv-ing nitrogen atom, the rates of protolysis were measured on several pyridine and 1,10-phenanthroline derivatives in non-buffered aqueous solution by the pH-jump method using the stopped-flow apparatus and also measured in 1,4-dioxane-water mixture.

Experimental

Materials and Sample Solutions. Samples of pyridine, 2,2’-bipyridine and 1,10-phenanthroline derivatives were purchased commercially available and were used without further purification, unless, other wise noted. 3-Aminopyri-dine, 4-aminopyridine, 4-dimethylaminopyridine, 4-methyl-pyridine, 3,4-dimethylpyridine, 4,7-dimethyl-1,10- phenan-throline, 5,6-dimethyl-1,10-phenanthroline, and 3,4,7,8- tetramethyl-1,10- phenanthroline were purified as

Kinetics and Mechanism of the Slow Protolysis of Pyridine and 1,10-Phenanthroline Derivatives

Kazuhisa Itoh1), Isao Ando1)*, Kikujiro UjImoto1), and Hirondo KUrIhArA1)

(Received December 1, 2014)

Abstract The rate constant for the protolysis of ten pyridine derivatives, two 2,2’-bipyridine derivatives, and seven 1,10-phenanthroline derivatives were determined in aqueous solution by a stopped-flow pH-jump method. The linear free energy relationship (LFER) was examined on their protolysis. It is clarified that their LFER is dependent on the reaction path. Also the rate of the protolysis of pyridine and 1,10-phenanthroline were measured at 0.05 mole fraction of 1,4-dioxane in 1,4-dioxane-water mixture. The acceleration of their protolysis was observed by addition of 1,4-diox-ane. This acceleration was explained in terms of the change in hydration sphere.

1) Department of Chemistry, Faculty of Science, Fukuoka University, 8-19-1, Nanakuma, Jonan-ku, Fukuoka 814-0180, Japan

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mentioned below before use. 3-Aminopyridine, 4-amino-pyridine, and 4-dimethylaminopyridine were reduced by sodium dithionite using a Norite and recrystallized from a benzene-ligroin 4:1 mixture.[8] 4-Methyl- and 3,4-dimethyl-pyridine were converted to their hydrochloride by passing hydrogen chloride gas into them and isolated. These hydro-chloride were used after recrystallization from ethanol. 5,6-Dimethyl- and 3,4,7,8-tetramethyl-1,10-phenanthroline were recrystallized from a water-ethanol 4:1 mixture using activated charcoal.

1,4-dioxane was purchased from Wako Pure Chemical Industries. This solvent was reagent grade for nonaqueous titration and used without further purification. The water content of 1,4-dioxane was checked on with use of a Mit-subishi CA-02 moisturemeter and was found less than 0.2 w/w %. All other reagents were of guaranteed reagent grade and used further purification.

The stock solution of a sample was prepared by dis-solving a desired amount of it in water. For water-insoluble samples, e.g. 4-phenylpyridine, 4,4’-dimethyl-2,2’-bipyri-dine, and 5-chloro-, 5-methyl-, 4,7-dimethyl-, 5,6-dimeth-yl-, and 3,4,7,8-tetramethyl-1,10- phenanthroline, the stock solutions of them were prepared by dissolving desired amount of them in a small amount of ethanol and diluting with water. Those stock solutions were dilute to the sample concentration of 2.5×10-5 – 1.0×10-4 mol dm-3 for the equilibrium and kinetic measurements. The resulting solu-tion contained ethanol less than 1 v/v %.

Measurements in Aqueous Solution. Acid dissociation constants and rate constants were measure in aqueous solu-tion at the ionic strength of 0.20 mol dm-3 with sodium chlo-ride and 25°C. Acid dissociation constants were measure by the conventional spectrophotometric titration method at the absorption maximum wavelength of the conjugate acid of a sample. The sample solutions contained the following buf-fers at the concentration of less than 0.02 mol dm-3 in the pH range of 4–10 for the adjustment of pH using acetate, phos-phate, and ammonium buffer systems.

The rate of protolysis was followed at the absorption maximum wavelength by a stopped-flow spectrophotometry applicating the pH-jump method without buffer.[9] A dilute sodium hydroxide solution with the same ionic strength as that of a sample solution was used as the perturbing solution to be mixed with a sample solution. An apparent first-order rate constant was obtained from 10 to 20 kinetic runs by the method previously reported.[9] The detailed procedures of equilibrium and kinetic measurement were already de-scribed in the previous paper.[5]

Measurements in 1,4-Dioxane-Water Mixture. Acid dissociation constants and rates for protolysis were also measured in 1,4-dioxane-water mixture containing 0.05 mole fraction of 1,4-dioxane. These measurements were carried out by the same manner as in aqueous solution ex-cept pH measurement. The concentration of hydrogen ion cannot be directly obtained from the reading of a pH meter because there is no standard buffer solution suited for cor-rection of a pH meter in 1,4-dioxane-water mixture. There-fore, pH of a sample solution was determined according to Gran’s method.[10] The measurements were carried out un-der an atmosphere of nitrogen gas saturated with solvent vapor by using TOA HGS-2005 glass electrode in combina-tion with a TOA HS-305DS silver-silver chloride sleeve type electrode as a reference electrode. In each series of measurements, the constant E0 appearing in the Nernst equa-tion is determined by Gran plot[10] and then a potential of sample solution, Es is measured. Because the liquid junc-tion potential is negligibly small in the pH range of 3 to 10, Es is related to the concentration of free hydrogen ion, [H+] and E0 by

　

The concentration of free hydrogen ion was calculated by use of above equation.

Results and Discussion

The absorption spectra of py, phen, and their derivatives changed in depending on pH of the solution. These spectral changes agreed with those reported in literatures.[11, 15] The absorbance of each sample decreased markedly with in-creasing pH of the solution at the absorption maximum wavelength of the protonated species, HL+, of sample. The pKa values of them were determined by a spectrophotomet-ric titration at the wavelength of absorption maximum of HL+. The obtained pKa values are summarized in Table 1. These values are in good agreement with literature values within the difference between experimental conditions.[16, 17] The number of transferred proton, n, obtained together with pKa value was about unity for all samples in this study. These facts imply that the spectral and absorbance change corresponds to the deprotonation of HL+ to the free base, L. The apparent first-order rate constant of deprotonation, (kdp)0, was determined by following the absorbance change at the same wavelength as that for the spectrophotometric titration which was caused by pH-jump perturbation.

The protolysis of these py and phen derivatives is also considered to proceed through the following two concurrent

� －17－Kinetics and Mechanism of the Slow Protolysis of Pyridine and 1,10-Phenanthroline Derivatives（K. Itoh et al.）

paths as that of py, bpy, and phen in previous study,[5]

　　　 (1)

　　 (2)

where kdp, kp, kOH, and kH2O are the rate constants for the pro-tolysis of Eqs. 1 and 2, respectively. Since the decreasing rate of HL+ is expressed by

　

(3)

(kdp)0 at the reaction time zero is represented as function of hydroxide ion concentration at the reaction time zero, [OH-]0.

　 　 (4)

Figure 1 shows the plots of (kdp)0 against [OH-]0 for 4-phenylpyridine and 5-chloro-1,10-phenanthroline. These plots gave a good linearity. Although the plots in Fig. 1 are

the representative examples, the plots for all compounds in this study also gave a good linear relation. These protolytic behaviors for the compounds in this study are identical to those reported in previous paper.[5] This fact suggested that the protolysis of py and phen derivatives also proceeded through the concurrent paths of Eqs. 1 and 2. Table 2 sum-marizes the values of kdp and kOH evaluated from these plots by the least squares method. The kdp and kOH values of py, bpy, and phen were determined again in this study although the values were already reported.[10]

Since the rate constants in Eqs. 1 and 2 are related each other by Eq. 5 using the acid dissociation constant of HL+ and ionic product of water, KW,[18] the values of kp and kH2O can be calculated from the determined values of Ka and kdp or kOH, respectively.

　 　 (5)

The calculated values of kp and kH2O are summarized together in Table 3.

As can be seen from Tables 2 and 3, the values of kdp and kOH decrease with the increasing pKa and the values of kp is independent on pKa value and nearly constant. Although

Table 1 The pKa values and the number of the transferred proton for conjugate acid of py, bpy, and phen derivatives in aqueous solution at 25 ºC.

compound pKa n

4-Cl-py 3.94±0.01 0.97±0.02

4-Br-py 3.95±0.01 1.01±0.02

2-pheny-py 4.77±0.01 0.97±0.01

py 5.34±0.01 0.95±0.01

4-phenyl-py 5.48±0.01 1.04±0.01

3-NH2-py 6.16±0.01 1.02±0.01

4-CH3-py 6.18±0.01 0.98±0.01

3,4-(CH3)2-py 6.61±0.01 1.00±0.01

bpy 4.51±0.01 1.02±0.01

4,4�-(CH3)2-bpy 5.37±0.01 1.01±0.01

5-Cl-phen 4.26±0.01 1.01±0.02

phen 5.03±0.01 1.00±0.01

5-CH3-phen 5.32±0.02 0.97±0.06

5,6-(CH3)2-phen 5.66±0.01 0.97±0.01

2,9-(CH3)2-phen 5.84±0.02 1.02±0.02

4,7-(CH3)2-phen 6.12±0.01 0.98±0.01

3,4,7,8-(CH3)4-phen 6.56±0.01 0.98±0.01

These values are shown together with the 0.95 confidence

intervals.

Figure 1 The plots of (kdp)0 against [OH-]0 for (a) 4-phenyl-pyridine and (b) 5-chloro-1,10-phenanthroline in aqueous solution at 25 ºC.

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the fact may imply that the protonation of py and phen de-rivatives is diffusion-controlled, the value of kp is smaller than the diffusion controlled rate constant and comparable to that of carbon acids.[3] On the other hand, the value of kp increased with increasing pKa for cyclopentadienes.[5 19] The linear free energy relationship for the protolysis of py, bpy, and phen derivatives seems to be different from that of a carbon acid. The mechanism for the protolysis of py and phen derivatives might differ from that of a carbon acid. Thus, the linear free energy relationship was examined in detail for the protolysis of py, bpy, and phen derivatives.

Linear Free Energy Relationship (LFER). The log(k) values were plotted against pKa on py and phen derivatives in order to examine LFER for the protolysis of the reaction paths of Eqs. 1 and 2. Figure 2 shows the log(k) vs pKa plots on kdp, kp, kOH, and kH2O of the py, bpy, and phen deriva-tives. The plot on kdp gave a straight line for both py and phen derivatives and the plot on kp exhibited the same be-havior as that on kdp. Their slopes are -0.95 for the log(kdp) vs pKa plot and 0.05 for log(kp) vs pKa plot. On the other hand, the plots on kOH and kH2O gave the separate lines for both derivatives on each plot although the point on 2,9-di-methyl-1,10-phenanthroline deviates from the straight line on phen derivatives due to its steric hindrance.[20] The slopes of the log(kOH) vs pKa plots are -0.65 and -0.09 for py and phen derivatives, respectively, and those of log(kH2O) vs

pKa plots are 0.35 and 0.90 for py and phen derivatives, re-spectively. It is obscure at present which class bpy deriva-tives should belong to. Thus, these LFER suggests that the protolysis of py and phen derivatives proceeds with identi-cal mechanism in the reaction path of Eq. 1 and with differ-ent mechanism each other in the reaction path of Eq. 2.

It has been reported that the reaction site of 2,9-di-methyl-1,10-phenanthroline is less hydrated than that of other phen derivatives because of the steric hindrance by methyl groups of 2- and 9-positions.[20, 21] On the plots of log(kOH) vs pKa and log(kH2O) vs pKa, the point of 2,9-dimeth-yl-1,10-phenanthroline should fall under the category of py derivatives rather than phen derivatives. Therefore, it is supported that the difference in degree of hydration around the reaction site between py and phen derivatives should af-fect significantly on the LFER of them.

Protolysis in 1,4-Dioxane-Water Mixture. In the pre-liminary experiments, we found that the protolysis of py was accelerated about 2 to 5 times that in aqueous solution by addition of organic solvents such as methanol, dimethyl-sulfoxide and 1,4-dioxane. Ishiguro et al. investigated the solvent effect on the protonation of 1,10-phenanthroline in 1,4-dioxane-water mixture.[22] In their study, its effect was marked at the 0.05 mole fraction of 1,4-dioxane in 1,4-diox-ane-water mixture. Thus, we investigated the protolysis of py and phen at 0.05 mole fraction in 1,4-dioxane-water

Table 2 The rate constants for deprotonation of conjugate acid of py, bpy, and phen derivatives in aqueous solution at 25 ºC.

Table 3 The rate constants for protonation of py, bpy, and phen derivatives in aqueous solution at 25 ºC.

compound kdp / s-1 kOH / mol-1dm3s-1

4-Cl-py (5.7±4.5) 10-1 (6.8±2.9) 104

4-Br-py 1.0±0.5 (4.4±7.1) 104

2-pheny-py (2.6±0.6) 10-1 (1.8±0.5) 104

py (4.9±1.0) 10-2 (3.4±1.3) 103

4-phenyl-py (7.0±1.0) 10-2 (7.9±2.3) 103

3-NH2-py (1.0±0.2) 10-2 (1.1±0.4) 103

4-CH3-py (2.3±3.0) 10-3 (2.3±0.4) 103

3,4-(CH3)2-py (3.9±2.6) 10-3 (1.1±0.5) 103

bpy (3.2±0.5) 10-1 (4.5±1.3) 104

4,4�-(CH3)2-bpy (8.2±1.0) 10-2 (1.3±0.2) 104

5-Cl-phen (7.6±1.0) 10-1 (2.5±0.5) 104

phen (2.9±0.3) 10-1 (1.9±0.5) 104

5-CH3-phen (6.3±2.0) 10-2 (9.8±1.7) 103

5,6-(CH3)2-phen (5.5±1.0) 10-2 (1.7±0.3) 104

2,9-(CH3)2-phen (2.3±0.4) 10-2 (2.3±0.4) 103

4,7-(CH3)2-phen (2.6±0.4) 10-2 (1.2±0.2) 104

3,4,7,8-(CH3)4-phen (5.5±6.9) 10-3 (1.4±0.3) 104

These values are shown together with the 0.95 confidence intervals.

compound kp / mol-1dm3s-1 kH 2O / mol-1dm3s-1

4-Cl-py (5.0±3.9) 103 (5.9±2.5) 10-6

4-Br-py (8.9±4.5) 103 (3.9±6.3) 10-6

2-pheny-py (1.5±0.3) 104 (1.1±0.2) 10-5

py (1.1±0.2) 104 (7.5±2.9) 10-6

4-phenyl-py (2.1±0.3) 104 (2.4±0.7) 10-5

3-NH2-py (1.4±0.3) 104 (1.6±0.6) 10-5

4-CH3-py (3.4±4.6) 103 (3.4±0.6) 10-5

3,4-(CH3)2-py (1.6±1.1) 104 (4.4±2.1) 10-5

bpy (1.0±0.2) 104 (1.5±0.4) 10-5

4,4�-(CH3)2-bpy (1.9±0.2) 104 (3.0±0.5) 10-5

5-Cl-phen (1.4±0.2) 104 (4.5±0.9) 10-6

phen (3.1±0.3) 104 (2.1±0.5) 10-5

5-CH3-phen (1.3±0.4) 104 (2.1±0.4) 10-5

5,6-(CH3)2-phen (2.5±0.5) 104 (7.5±1.3) 10-5

2,9-(CH3)2-phen (1.6±0.3) 104 (1.6±0.3) 10-5

4,7-(CH3)2-phen (3.4±0.6) 104 (1.6±0.3) 10-4

3,4,7,8-(CH3)4-phen (2.0±2.5) 104 (4.9±1.2) 10-4

These values are shown together with the 0.95 confidence intervals.

� －19－Kinetics and Mechanism of the Slow Protolysis of Pyridine and 1,10-Phenanthroline Derivatives（K. Itoh et al.）

mixture at 25.0°C in order to examine the participation of solvent in their protolysis.

The acid dissociation constants of the conjugate acids of py and phen were measured in 1,4-dioxane-water mix-ture. The similar pH-titration curves to that in aqueous so-lution were obtained in that mixture solvent. The pKa val-ues were evaluated from these curves. The obtained pKa values are listed in Table 4. The pKa values for the conju-gate acid of phen is in good agreement with that of Ishiguro et al.[22] Furthermore, the number of the transferred proton was about unity for each conjugate acid.

The rates of protolysis for py and phen were also mea-sured at 25.0°C in the same mixture solvent. Figure 3 shows the plots of (kdp)0 vs [OH-]0 for py and phen. As can be seen from Fig. 3, each plot gave a good linear relation. This implies that the protolysis of py and phen proceeds through the two concurrent reaction paths of Eqs. 1 and 2 in 1,4-dioxane-water mixture as that does in aqueous solution. Table 4 summarizes the values of kdp and kOH evaluated from these plots. Table 4 also lists the values of kp and kH2O which were calculated from the obtained kdp or kOH value and Ka value, and Kw value in the literatures[23, 24] by use of Eq. 5. These rate constants were compared to those in aqueous so-lution. The kdp and kOH values are 3 to 5 times larger than those in aqueous solution although the kp and kH2O values are close to those in aqueous solution. Ishiguro et al reported that the addition of 1,4-dioxane affected on hydration of phen and its conjugate acid.[22] Furthermore, our results in 1,4-dioxane-water mixture indicate that the addition of 1,4-dioxane affect only on the hydration of py, phen and their conjugate acids but also on the degree of hydration of the activated state of their protolysis.

For the protolysis of py and phen derivatives, the de-pendence of LFER on the reaction path may be attributable to the difference in degree of their hydration. Furthermore, it is considered that the acceleration of the protolysis of py

Figure 2 Test of linear free energy relationship for the pyridine(○), 2,2’-bipyridine(△), and 1,10-phenanthroline derivatives(□) in aqueous solution at 25.0 ºC. (a) Plots on kdp and kOH, (b) plots on kp and kH2O.

Table 4 The pKa values and rate constants for the protolysis in 0.05 mole fraction 1,4-dioxane-water mixture at 25 ºC.

pyridine 1,10-phenanthroline

pKa 4.97±0.01 4.62±0.01

n 0.95±0.01 0.96±0.01

kdp / s-1 (2.7±0.6) 10-1 1.0±0.1

kOH / mol-1dm3s-1 (1.4±0.3) 104 (8.2±2.2) 104

kp / mol-1dm3s-1 (2.5±0.5) 104 (4.3±0.5) 104

kH2O / mol-1dm3s-1 (3.3±0.6) 10-6 (8.6±2.3) 10-6

These values are shown together with the 0.95 confidence intervals.

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and phen in 1,4-dioxane-water mixture may be owing to any modification of their hydration by the addition of 1,4-dioxane. Therefore, it is conclude that any modification in hydration sphere should play an important role in the protolysis of py and phen derivatives and may be main cause of the slow protolysis.

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24 H. S. Harmed and L. D. Fallon, J. Am. Chem. Soc., 61, 2374(1939).

Figure 3 The plots of (kdp)0 against [OH-]0 for (a) pyridine and (b) 1,10-phenanthroline in 0.05 mole fraction 1,4-dioxane-water mixture at 25.0 ºC.