Thermodynamics for Proton Binding of Pyridine in Different Ionic

Article pubs.acs.org/jced

Thermodynamics for Proton Binding of Pyridine in Different Ionic Media at Different Temperatures Clemente Bretti, Rosalia Maria Cigala, Concetta De Stefano,* Gabriele Lando, and Silvio Sammartano Dipartimento di Scienze Chimiche, Università degli Studi di Messina, Viale Ferdinando Stagno d’Alcontres, 31, I-98166 Messina (Vill. S. Agata), Italy ABSTRACT: The acid−base behavior of pyridine (py) was studied at different temperatures and ionic strengths in LiCl, NaCl, KCl, RbCl, CsCl, NaNO3, (CH3)4NCl, (C2H5)4NI, MgCl2, and CaCl2 aqueous solutions. The study was developed by means of potentiometric titrations with a ISE-H+ electrode and by the reanalysis of literature data. It was found that the protonation process of pyridine is enthalpic-driven and the values of the protonation enthalpy are in agreement with literature findings. Different models were used to fit the experimental data, namely, the Debye−Hückel equation, the specific ion interaction theory (SIT), and the Pitzer equations. The variation of the apparent protonation constant in the different ionic media was also interpreted in terms of formation of weak complexes between pyridine and the ions of the supporting electrolyte. The formation of (CH3)4Npy+, (C2H5)4Npy+, Mgpy2+, Capy2+, and HpyCl species was found. The modeling ability were comparable for all considered models, although the Pitzer equations can be regarded as the most reliable. SIT and Pitzer coefficients were provided for the interaction between the Hpy+ and the Cl−, and the activity coefficient of the neutral species has been reported. A good agreement was found between experimental data obtained in this work and literature findings.

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INTRODUCTION Pyridine (azine, C5H5N)1 is the most simple azine, and it is a colorless liquid with a disagreeable odor, highly flammable and water-soluble. In natural occurrence pyridine is found among the volatile components of black tea, in the leaves and roots of Atropa belladonna, in the volatile components of various foods,2 but also in the tobacco and marijuana smoke.3 Pyridine production is around 200.000 tonnes per year (only China 120.000 tons/year), and the demand is continuously increasing; it is widely used as a solvent in organic chemistry and industrial practice. Pyridine is used for paint, rubber, and textile water repellents, as raw material for antimicrobial agents, and also for medical products.4−6 Large amounts of pyridine are used as an intermediate in the manufacture of agrochemicals (diquat, chloropyrifos, pyrithione), pharmaceuticals, and other products. Pyridine and its homologous (azines) are water-soluble and can be transported through aquifer materials. Many studies are present in literature about the carcinogenicity and the toxicity effects of pyridine.7 In humans, acute pyridine intoxication affects nervous system, leading to headaches, nausea, and anorexia.8−10 Due to the importance of pyridine in the industrial field, its physicochemical and thermodynamic properties have been studied.11−14 In particular, Chirico et al.11,12 collected in two papers many literature data and calculated new recommended values for some thermodynamic properties, such as heat capacity, enthalpy combustion, vapor pressure, and densities. On the contrary, relatively few data are reported in literature about protonation constants and their dependence on ionic medium and temperature. Marsicano et al.15 studied the free energy of complex formation between Ndonor ligands (including pyridine) and metal cations, such as Ag+, Hg2+, and Cu2+. The authors found that the complex © 2013 American Chemical Society

formation process is generally enthalpy-driven and a linear relationship was found between the free energy of complex formation of Ag+ and Hg2+. Ashton and co-workers16 considered the effect of temperature and ionic strength on the protonation of aromatic heterocyclic amines mostly reanalyzing literature data. In refs 17−19, the acid base and thermodynamic properties of pyridine and some derivatives have been investigated, finding reliable data for log KH and protonation enthalpy. In this work, potentiometric measurements were performed to determine the protonation constant of pyridine (py) in various ionic media (LiCl, NaCl, NaNO3, KCl, RbCl, CsCl, MgCl2, CaCl2, (CH3)4NCl, and (C2H5)4NI) at different temperatures in the range (283.15 ≤ T/K ≤ 318.15) and ionic strengths (0.05 ≤ I/mol·dm−3 ≤ 5.00). In addition, some literature data from Capone et al.18 were re-evaluated on the basis of new computer programs that were not available at that time, 1985, and new data were determined. The whole set was then analyzed with different models, such as Debye−Hückel, specific ion interaction theory (SIT), Pitzer equations and ion pairing formation to provide data in a standard state and to study the medium effect on the acid− base properties of the most simple azine.

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EXPERIMENTAL SECTION Chemicals. Pyridine hydrochloride was used without further purification; its purity was checked alkalimetrically and was always > 99 %. LiCl, NaCl, NaNO3, KCl, RbCl, and CsCl solutions were prepared weighing the pure salts, previously dried Received: November 4, 2013 Accepted: December 13, 2013 Published: December 23, 2013 143

dx.doi.org/10.1021/je4009602 | J. Chem. Eng. Data 2014, 59, 143−156

Journal of Chemical & Engineering Data

Article

The parameters of the acid−base titrations (E0, the junction potential coefficient ja, log Kw, analytical concentration of components) were refined using the ESAB2M computer program. BSTAC and STACO were also used to determine protonation constant of pyridine, and ES2WC was adopted for the determination of the weak complexes between pyridine and the ions of the supporting electrolyte. The general least-squares computer program LIANA was also used to fit different equations. All of the computer programs were reviewed elsewhere.24−26 The experimental data were corrected for temperature-induced volume changes. The protonation data were determined in the molar concentration scale. The corresponding values in the molal concentration scale were obtained using appropriate conversion equations.27

in an oven at T = 383.15 K for two hours. MgCl2 and CaCl2 solutions were standardized by ethylenediaminetetraacetic acid (EDTA) standard solutions.20 (CH3)4NCl and (C2H5)4NI solutions were prepared weighing the salts, previously purified as described in Perrin.21 For measurements in (C2H5)4NI medium, the (C2H5)4NOH was used as the titrant. NaOH, (C2H5)4NOH, and HCl solutions were prepared diluting concentrated solutions and standardized against sodium carbonate (for HCl) and potassium hydrogen phtalate (for bases), previously dried in an oven at T = 383.15 K for two hours. Soda lime traps were used to prevent absorption of atmospheric CO2 by NaOH and (C2H5)4NOH. All solutions were prepared with grade A glassware and using twice distilled water (R ≥ 18 MΩ). All chemicals were purchased from Sigma and its associates at the highest purity available. Apparatus and Procedure for Potentiometric Titrations. Potentiometric measurements were performed to determine the protonation constant of pyridine (py) in various ionic media (LiCl, NaCl, NaNO3, KCl, RbCl, CsCl, MgCl2, CaCl2, (CH3)4NCl, and (C2H5)4NI) at different temperatures in the range (283.15 ≤ T/K ≤ 318.15) and ionic strengths (0.05 ≤ I/mol·dm−3 ≤ 5.00); the pyridine concentration ranged between 0.002 ≤ cpy/mol·dm−3 ≤ 0.020. All of the potentiometric measurements were performed in the range 2 ≤ pH ≤ 10, and the experimental details are summarized in Table 1. A 25 cm3

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MODELS FOR THE DEPENDENCE OF EQUILIBRIUM CONSTANTS ON IONIC STRENGTH AND TEMPERATURE The calculation on potentiometric data, collected in different conditions of medium, temperature, and ionic strength, were performed by BSTAC and STACO computer programs to obtain the pyridine protonation constant reported in Tables 2 to 11. In the same tables, the results of a reanalysis of literature pyridine protonation data are given. In the past, Capone et al.18 reported a systematic study on the protonation of pyridine in all of the ionic media studied in this work (except NaNO3). In this work, these data were reanalyzed using some more recent computer programs (e.g., BSTAC, STACO, ESAB2M) which were not available at that time, 1985. These least-squares programs refine the formation parameters, such as protonation constants, minimizing the error squares sum in the titrant volume (STACO and ESAB2M) and in the experimental potential E/ mV (BSTAC). Moreover, STACO and BSTAC are able to take into account the ionic strength variation between different titrations and within different points of the same titration. The experimental data were then modeled with different equations, each of them demonstrating a good modeling ability with comparable statistical parameters. The four considered models can be divided in two main categories, one dealing with the variation of the activity coefficients with ionic strength and another one which takes into account the formation of weak complexes between the molecule under investigation (pyridine in our case) and the ions of the supporting electrolytes. In the first case, the equilibrium in eq 1 was studied considering the variation of the activity coefficients (γ) as follows:

Table 1. Experimental Conditions for the Potentiometric Measurements Performed in This Work in the Range 283.15 ≤ T/K ≤ 318.15 salt

csalt/mol·dm−3

no. measurements

LiCl NaCl NaNO3 KCl RbCl CsCl (CH3)4NCl (C2H5)4NI MgCl2 CaCl2

0.19−4.93 0.19−4.85 0.10−4.85 0.06−3.06 0.06−3.56 0.06−3.71 0.06−3.05 0.06−0.89 0.02−1.49 0.02−1.06

12 15 18 13 14 15 20 8 12 11

solution, containing pyridine and the ionic medium, was titrated with standard NaOH (or (C2H5)4NOH for measurements in (C2H5)4NI), and for each measurement 50 to 70 data points were collected. To check the repeatability of the system two apparatuses were used, described elsewhere.22,23 Before each experiment, independent titrations of HCl solutions with standard sodium hydroxide (or (C2H5)4NOH) were performed to determine the formal electrode potential in the same experimental conditions (temperature and ionic strength) of the systems under investigation. The free hydrogen ion concentration scale was used (pH = −log[H+]). Calculations. The protonation constant of pyridine refers to the following equilibrium: H+ + py = Hpy +

KH

log K H = log K H0 + log γH+ + log γpy − log γHpy +

The variation of log γ on ionic strength can be interpreted in different ways, according to the Debye−Hückel, specific ion interaction theory (SIT), and Pitzer approaches (see later text). The dependence of the protonation constant on ionic strength was taken into account using the following general equation: log K H = log K H0 + L(I , T )

(4)

where log KH is the protonation constant, log KH0 refers to the protonation at infinite dilution, and L(I, T) is a function of ionic strength and temperature that can be expressed in different ways. Models for the Variation of the Activity Coefficients: Debye−Hückel Equation. Data analysis performed with the Debye−Hückel equation was mainly used as a smoothing function, and the fitting equation was implemented with three

(1)

Other equilibria refer to the general equation: p M z + + q py + r H+ + sCl− = M p(py)q H r Cls(pz + r − s) K pqrs

(3)

(2) 144



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Table 2. Protonation Constant (eq 1) of Pyridine in LiCl Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15

0.294

283.15

0.580

5.486 ± 0.002 5.490 ± 0.002a 5.488 ± 0.002a 5.572 ± 0.002a 5.583 ± 0.002a 5.577 ± 0.002a 5.678 ± 0.003a 5.705 ± 0.003a 5.691 ± 0.003a 5.479 ± 0.001 5.497 ± 0.001 5.984 ± 0.004 5.980 ± 0.004 6.962 ± 0.014 6.926 ± 0.014 5.286 ± 0.001a 5.278 ± 0.001a 5.282 ± 0.001a 5.302 ± 0.001a 5.306 ± 0.001a 5.304 ± 0.001a 5.374 ± 0.001a 5.377 ± 0.001a 5.376 ± 0.001a 5.379 ± 0.001a 5.392 ± 0.001a

T/K

283.15

291.15 291.15 291.15 291.15 291.15 291.15 298.15

298.15

298.15

298.15 a

0.950

0.502 0.540 1.999 2.002 4.833 4.834 0.194

0.293

0.480

0.580

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15

0.580 0.940

298.15

0.950

298.15 298.15 310.15 310.15 310.15 310.15 318.15

4.882 4.926 2.280 2.286 4.783 4.793 0.380

318.15

0.750

318.15

1.240

318.15 318.15 318.15 318.15

1.517 1.518 3.640 3.642

5.386 ± 0.001a 5.519 ± 0.001a 5.523 ± 0.001a 5.521 ± 0.001a 5.481 ± 0.001a 5.502 ± 0.001a 5.491 ± 0.001a 6.813 ± 0.012 6.828 ± 0.012 5.747 ± 0.004 5.779 ± 0.004 6.576 ± 0.010 6.606 ± 0.010 5.080 ± 0.001a 5.084 ± 0.001a 5.082 ± 0.001a 5.174 ± 0.002a 5.186 ± 0.002a 5.180 ± 0.002a 5.284 ± 0.004a 5.314 ± 0.004a 5.299 ± 0.004a 5.443 ± 0.004a 5.442 ± 0.004a 6.047 ± 0.009a 6.049 ± 0.009a

Recalculated from Capone et al.18

Table 3. Protonation Constant (eq 1) of Pyridine in NaCl Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15

0.294

283.15

0.580

283.15

0.950

291.15 291.15 298.15

4.757 4.758 0.197

298.15

0.293

298.15

0.480

298.15

0.580

298.15

0.950

5.502 ± 0.002 5.496 ± 0.002a 5.499 ± 0.002a 5.599 ± 0.002a 5.596 ± 0.002a 5.597 ± 0.002a 5.720 ± 0.003a 5.714 ± 0.003a 5.717 ± 0.003a 6.869 ± 0.016 6.871 ± 0.016 5.284 ± 0.001a 5.285 ± 0.001a 5.285 ± 0.001a 5.312 ± 0.001a 5.312 ± 0.001a 5.312 ± 0.001a 5.373 ± 0.001a 5.373 ± 0.001a 5.373 ± 0.001a 5.402 ± 0.001a 5.405 ± 0.001a 5.403 ± 0.001a 5.512 ± 0.001a 5.515 ± 0.001a 5.514 ± 0.001a

T/K

a

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15

0.960

298.15 298.15 298.15 298.15 298.15 298.15 310.15 310.15 310.15 310.15 318.15

1.921 1.945 2.884 2.909 4.398 4.403 2.018 2.019 4.823 4.852 0.291

318.15

0.570

318.15

0.940

318.15 318.15 318.15

3.778 3.779 4.791

5.521 ± 0.001a 5.521 ± 0.001a 5.521 ± 0.001a 5.809 ± 0.002 5.814 ± 0.002 6.132 ± 0.004 6.086 ± 0.004 6.519 ± 0.009 6.516 ± 0.009 5.660 ± 0.004 5.654 ± 0.004 6.434 ± 0.006 6.453 ± 0.006 5.069 ± 0.001a 5.066 ± 0.001a 5.067 ± 0.001a 5.145 ± 0.002a 5.144 ± 0.002a 5.144 ± 0.002a 5.244 ± 0.004a 5.241 ± 0.004a 5.242 ± 0.004a 5.966 ± 0.009 5.964 ± 0.009 6.256 ± 0.009


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Table 4. Protonation Constant (eq 1) of Pyridine in NaNO3 Aqueous Medium, at Different Temperatures and Ionic Strengths T/K

I/mol·dm−3

log KH ± 2σ

T/K

I/mol·dm−3

log KH ± 2σ

291.15 291.15 291.15 291.15 298.15 298.15 298.15 298.15 298.15

2.015 2.016 4.842 4.847 0.100 0.500 1.000 4.514 4.516

5.923 ± 0.004 5.927 ± 0.004 6.557 ± 0.009 6.543 ± 0.009 5.246 ± 0.001 5.375 ± 0.001 5.515 ± 0.001 6.276 ± 0.006 6.263 ± 0.006

310.15 310.15 310.15 310.15 310.15 310.15 310.15 318.15 318.15

1.983 1.985 2.725 2.766 4.702 4.727 4.752 2.760 2.763

5.516 ± 0.004 5.523 ± 0.004 5.610 ± 0.005 5.605 ± 0.005 5.918 ± 0.009 5.898 ± 0.009 5.898 ± 0.009 5.431 ± 0.007 5.432 ± 0.008

Table 5. Protonation Constant (eq 1) of Pyridine in KCl Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15

0.065

283.15

0.150

5.414 ± 0.002 5.410 ± 0.002a 5.412 ± 0.002a 5.447 ± 0.002a 5.440 ± 0.002a 5.443 ± 0.002a 5.495 ± 0.002a 5.492 ± 0.002a 5.494 ± 0.002a 5.540 ± 0.002a 5.540 ± 0.002a 5.540 ± 0.002a 5.584 ± 0.002a 5.578 ± 0.002a 5.581 ± 0.002a 5.693 ± 0.004a 5.691 ± 0.004a 5.692 ± 0.004a 5.967 ± 0.007 5.946 ± 0.007 6.096 ± 0.012 5.239 ± 0.001a 5.232 ± 0.001a 5.235 ± 0.001a 5.266 ± 0.001a 5.262 ± 0.001a 5.264 ± 0.001a 5.275 ± 0.001a 5.278 ± 0.001a 5.276 ± 0.001a 5.311 ± 0.001a 5.305 ± 0.001a 5.308 ± 0.001a 5.350 ± 0.001a 5.348 ± 0.001a 5.349 ± 0.001a 5.358 ± 0.001a 5.362 ± 0.001a

T/K

283.15

283.15

0.437

283.15

0.580

283.15

0.950

291.15 291.15 291.15 298.15

2.103 2.104 2.736 0.064

298.15

298.15

298.15

298.15

298.15 a

0.294

0.150

0.197

0.293

0.435

0.480

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15

0.480 0.580

298.15

0.950

298.15

0.960

298.15 298.15 298.15 298.15 310.15 310.15 318.15

1.136 1.137 2.719 2.719 3.063 3.065 0.064

318.15

0.148

318.15

0.292

318.15

0.433

318.15

0.574

318.15

0.940

318.15 318.15 318.15 318.15

1.914 1.916 3.070 3.077

5.360 ± 0.001a 5.393 ± 0.001a 5.387 ± 0.001a 5.390 ± 0.001a 5.490 ± 0.001a 5.490 ± 0.001a 5.490 ± 0.001a 5.488 ± 0.001a 5.492 ± 0.001a 5.490 ± 0.001a 5.539 ± 0.002 5.559 ± 0.002 5.908 ± 0.008 5.915 ± 0.008 5.785 ± 0.006 5.797 ± 0.006 5.001 ± 0.001a 4.994 ± 0.001a 4.997 ± 0.001a 5.026 ± 0.001a 5.023 ± 0.001a 5.024 ± 0.001a 5.065 ± 0.001a 5.064 ± 0.001a 5.064 ± 0.001a 5.102 ± 0.002a 5.095 ± 0.002a 5.099 ± 0.002a 5.133 ± 0.002a 5.131 ± 0.002a 5.132 ± 0.002a 5.223 ± 0.004a 5.225 ± 0.004a 5.224 ± 0.004a 5.412 ± 0.006 5.431 ± 0.006 5.599 ± 0.008 5.642 ± 0.008


parameters that account for the dependence of the protonation constant on temperature. According to the Debye−Hückel model the equilibrium constant and the ionic strength must be given in the molar concentration scale. In this case, all of the data reported in Tables 2 to 11 at different temperatures and ionic strengths were analyzed with eq 4, where L(I, T) is expressed as follows:

L(I , T ) = I(C + EI ) + (T − 298.15) · ⎞ ⎛ A (T − 298.15) + A 2I ⎟ ⎜A 0 + 1 ⎠ ⎝ 1000

(5)

where C and E are empirical parameters for the ionic strength dependence, whereas A0, A1, and A2 are empirical parameters for the temperature dependence. 146



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Table 6. Protonation Constant (eq 1) of Pyridine in RbCl Aqueous Medium, at Different Temperatures and Ionic Strengths T/K 283.15 283.15 283.15 283.15 283.15 291.15 291.15 291.15 291.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 a

I/mol·dm−3

log KH ± 2σ

0.150 0.294 0.439 0.580 0.950 1.873 1.875 3.562 3.562 0.064 0.150 0.294 0.438 0.578 0.950 1.428

5.427 ± 0.002 5.459 ± 0.002a 5.497 ± 0.002a 5.544 ± 0.003a 5.623 ± 0.005a 5.882 ± 0.006 5.902 ± 0.006 6.091 ± 0.015 6.101 ± 0.015 5.230 ± 0.001a 5.251 ± 0.001a 5.284 ± 0.001a 5.314 ± 0.001a 5.356 ± 0.001a 5.417 ± 0.001a 5.624 ± 0.002 a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15 298.15 310.15 310.15 310.15 310.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15

1.429 2.884 2.886 1.919 1.924 3.702 3.702 0.064 0.149 0.291 0.435 0.576 0.940 2.754 2.755

5.592 ± 0.002 5.867 ± 0.007 5.870 ± 0.007 5.462 ± 0.006 5.449 ± 0.006 5.713 ± 0.012 5.730 ± 0.012 4.991 ± 0.001a 5.010 ± 0.001a 5.037 ± 0.002a 5.063 ± 0.003a 5.090 ± 0.003a 5.136 ± 0.005a 5.492 ± 0.014 5.520 ± 0.014


Table 7. Protonation Constant (eq 1) of Pyridine in CsCl Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15 283.15 283.15 283.15 283.15 283.15 291.15

0.064 0.150 0.293 0.437 0.580 0.950 1.861

291.15 291.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15

3.709 3.710 0.064 0.150 0.293 0.436 0.580 0.950 1.470

5.406 ± 0.002 5.428 ± 0.002a 5.478 ± 0.002a 5.518 ± 0.002a 5.555 ± 0.003a 5.638 ± 0.004a 5.824 ± 0.006 5.824 ± 0.006 5.974 ± 0.018 5.974 ± 0.018 5.233 ± 0.001a 5.258 ± 0.001a 5.298 ± 0.001a 5.332 ± 0.001a 5.367 ± 0.001a 5.450 ± 0.002a 5.555 ± 0.003

T/K

a

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15 310.15 310.15 310.15 310.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15

1.496 2.833 1.863 1.868 3.575 3.578 0.064 0.149 0.291 0.433 0.570 0.950 1.345

318.15 318.15

2.857 2.861

5.563 ± 0.003 5.772 ± 0.009 5.526 ± 0.006 5.519 ± 0.006 5.682 ± 0.015 5.674 ± 0.015 4.992 ± 0.001a 5.021 ± 0.001a 5.047 ± 0.002a 5.087 ± 0.002a 5.123 ± 0.003a 5.187 ± 0.004a 5.275 ± 0.006 5.282 ± 0.006 5.448 ± 0.014 5.450 ± 0.014


The C values were fixed to be equal for all the ionic media, except for (CH3)4NCl and (C2H5)4NI because usually the protonation behavior in these ionic media is significantly different than in media containing only alkaline earth metal cations. Similar considerations can be done for the values of the A0, A1, and log KH0 parameters, which were constrained to be equal for all of the ionic media. For the data in MgCl2 and CaCl2 the term L(I, T) is

SIT equation,28−31 and in NaCl at T = 298.15 K, the term L(I) = Δε is Δε = ε(H+, Cl−) ·mCl + k mI − ε(Hpy +, Cl−) ·mCl

L(I , T ) = CI + EI 2 + B0 (T − 298.15) + B1(T − 298.15)2 + B2 I(T − 298.15) + B3I 2(T − 298.15)

−

(7)

−

where ε(H , Cl ) and ε(Hpy , Cl ) are the specific interaction coefficients of the H+ and Hpy+ species with the counterion (Cl−) of the supporting electrolyte and are multiplied for its molal concentration; km is the Setschenow32 coefficient of the neutral (py) species and was calculated considering literature data on similar azines. In particular, the km values of 2,2′,6,2″terpyridine (terpy), 2,2′-bypiridyl (bypy), and 1,10-phenantroline (phen), published in Bretti et al.,17 were averaged, and km ± 0.02 = 0.184 + 0.025 ln(1 + mNaCl). This value, together with the literature ε(H+, Cl−), reported in Bretti et al.,33 was fixed, and the fit of the experimental log KH data (only at T = 298.15 K) versus the ionic strength in molal scale allowed us to determine the unknown ε(Hpy+, Cl−), which resulted ε (Hpy+, Cl−) ± 0.006 = −0.0366 + 0.0799 ln (1 + mNaCl). The activity coefficient of the neutral species (py0) can be calculated multiplying the km values for the ionic strength in the molal concentration scale. The +

(6)

In this case, the log KH0 and the empirical parameters B0, B1, and B3 were considered equal for the two ionic media. Models for the Variation of the Activity Coefficients: Specific Ion Interaction Theory. When both the concentrations and the equilibrium constants are given in the molal concentration scale, the Debye−Hückel equation became the 147

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Table 8. Protonation Constant (eq 1) of Pyridine in (CH3)4NCl Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15 283.15 283.15 283.15 283.15 283.15 291.15 291.15 291.15

0.064 0.150 0.293 0.437 0.580 0.845 1.469 1.474 2.802

298.15 298.15 298.15 298.15 298.15 298.15 298.15

0.064 0.150 0.293 0.435 0.580 0.936 0.942

298.15

1.881

5.404 ± 0.002 5.407 ± 0.002a 5.435 ± 0.002a 5.432 ± 0.002a 5.450 ± 0.003a 5.452 ± 0.004a 5.342 ± 0.004 5.352 ± 0.004 5.236 ± 0.010 5.233 ± 0.010 5.232 ± 0.001a 5.233 ± 0.001a 5.261 ± 0.001a 5.250 ± 0.002a 5.270 ± 0.002a 5.264 ± 0.002 5.267 ± 0.002 5.265 ± 0.002 5.233 ± 0.004

T/K

a

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15

1.881 3.049

310.15

1.464

310.15 310.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15

2.786 2.789 0.064 0.149 0.291 0.433 0.570 0.880 1.461

318.15

2.791

5.237 ± 0.004 5.137 ± 0.009 5.137 ± 0.009 5.137 ± 0.009 5.102 ± 0.003 5.102 ± 0.003 5.010 ± 0.006 5.000 ± 0.006 5.003 ± 0.001a 4.992 ± 0.001a 5.006 ± 0.002a 5.004 ± 0.002a 5.020 ± 0.003a 5.019 ± 0.003a 4.999 ± 0.005 4.999 ± 0.005 4.897 ± 0.008 4.916 ± 0.008


Table 9. Protonation Constant (eq 1) of Pyridine in (C2H5)4NI Aqueous Medium, at Different Temperatures and Ionic Strengths I/mol·dm−3

log KH ± 2σ

283.15

0.150

283.15 283.15 283.15 291.15

0.293 0.580 0.890 0.511

291.15

0.859

298.15 298.15 298.15 298.15

0.064 0.150 0.293 0.435

5.401 ± 0.003 5.400 ± 0.003a 5.409 ± 0.004a 5.408 ± 0.005a 5.392 ± 0.007a 5.312 ± 0.004 5.311 ± 0.004 5.298 ± 0.006 5.298 ± 0.006 5.222 ± 0.002a 5.227 ± 0.002a 5.226 ± 0.003a 5.220 ± 0.004a

T/K

a

a

T/K

I/mol·dm−3

log KH ± 2σ

298.15 298.15 310.15

0.580 0.890 0.351

310.15

0.885

318.15 318.15 318.15 318.15 318.15 318.15

0.064 0.148 0.291 0.433 0.576 0.885

5.213 ± 0.003a 5.185 ± 0.006a 5.065 ± 0.003 5.065 ± 0.003 5.054 ± 0.006 5.061 ± 0.006 4.982 ± 0.001a 4.986 ± 0.002a 4.984 ± 0.002a 4.975 ± 0.002a 4.966 ± 0.002a 4.934 ± 0.008a


Table 10. Protonation Constant (eq 1) of Pyridine in MgCl2 Aqueous Medium, at Different Temperatures and Ionic Strengths cMgCl2/mol·dm−3

log KH ± 2σ

283.15 283.15 283.15 283.15 283.15 283.15 291.15

0.022 0.050 0.098 0.146 0.193 0.317 0.970

298.15 298.15 298.15 298.15 298.15 298.15 298.15

0.022 0.050 0.098 0.145 0.193 0.317 0.962

5.403 ± 0.005 5.420 ± 0.005a 5.454 ± 0.005a 5.491 ± 0.006a 5.525 ± 0.008a 5.596 ± 0.012a 5.851 ± 0.012 5.851 ± 0.012 5.222 ± 0.007a 5.240 ± 0.006a 5.264 ± 0.006a 5.293 ± 0.007a 5.321 ± 0.007a 5.383 ± 0.009a 5.751 ± 0.009

T/K

a

a

T/K

cMgCl2/mol·dm−3

log KH ± 2σ

298.15 310.15

0.962 0.984

310.15

1.492

318.15 318.15 318.15 318.15 318.15 318.15 318.15

0.021 0.049 0.097 0.144 0.191 0.315 0.368

318.15 318.15

0.954 0.955

5.752 ± 0.009 5.567 ± 0.010 5.576 ± 0.010 6.045 ± 0.024 6.045 ± 0.024 5.003 ± 0.008a 5.014 ± 0.007a 5.032 ± 0.007a 5.049 ± 0.007a 5.069 ± 0.008a 5.125 ± 0.011a 5.130 ± 0.012 5.074 ± 0.012 5.465 ± 0.014 5.475 ± 0.014


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Table 11. Protonation Constant (eq 1) of Pyridine in CaCl2 Aqueous Medium, at Different Temperatures and Ionic Strengths cCaCl2/mol·dm−3

log KH ± 2σ

283.15 283.15 283.15 283.15 283.15 283.15 291.15

0.021 0.050 0.097 0.146 0.193 0.317 0.973

298.15 298.15 298.15 298.15 298.15 298.15

0.022 0.050 0.098 0.146 0.193 0.317

5.399 ± 0.006 5.417 ± 0.005a 5.455 ± 0.004a 5.486 ± 0.004a 5.519 ± 0.005a 5.593 ± 0.008a 5.870 ± 0.013 5.850 ± 0.013 5.223 ± 0.006a 5.241 ± 0.006a 5.275 ± 0.005a 5.301 ± 0.004a 5.327 ± 0.003a 5.397 ± 0.003a

T/K

a

a

T/K

cCaCl2/mol·dm−3

log KH ± 2σ

298.15

1.063

310.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15 318.15

0.971 0.022 0.051 0.098 0.190 0.315 0.461 0.464 0.475 0.487 0.868 0.882

5.832 ± 0.013 5.845 ± 0.013 5.661 ± 0.014 5.003 ± 0.008a 5.013 ± 0.007a 5.031 ± 0.007a 5.080 ± 0.008a 5.119 ± 0.009a 5.240 ± 0.011 5.245 ± 0.011 5.240 ± 0.011 5.276 ± 0.011 5.534 ± 0.016 5.555 ± 0.016


molal concentration scale, T in Kelvin degrees (K), and Δε is given in eq 7. Both the specific interaction coefficients, the ε and the summation of them (Δε) can be true constants or vary with ionic strength. In this last case, some equations have been proposed in the past (e.g., in Bretti et al.35): p − p∞ P = p∞ + 0 (9) I+1

function of log γpy in NaCl at T = 298.15 K is shown in Figure 1. According to previous works,17,22,34 both the specific interaction

where P can be ε or Δε, p0 = p at I → 0 and p∞ = p at I → ∞ or: P = p0 + p1 ln(I + 1) 0

where p and p do not vary with ionic strength. According to Hepler,36 thermodynamic parameters should be calculated in the molal concentration scale. Data obtained in different ionic media (in the molal concentration scale) were also fitted considering the normalized crystallographic radii of the cation of the supporting electrolytes. The crystallographic radii of Li+, Na+, K+, Rb+, and Cs+ were taken from Shannon37 and are 90, 116, 152, 166, and 181 pm, respectively; these quantities were normalized and are −1.37, −0.67, 0.29, 0.67, and 1.07, respectively. In other words, the ionic strength dependence parameters where parametrized taking into account the normalized ionic radii according to the following equation:

Figure 1. Activity coefficient values of pyridine at different ionic strengths in NaCl and at T = 298.15 K.

coefficients and the Setschenow coefficients can be true constants or vary with ionic strengths according to different equations. For the ionic media different than NaCl it was not possible to determine the single interaction coefficients, because both the km and the ε(Hpy+, X−), where X is the anion of the supporting electrolyte, are unknown, and two parameters cannot be refined with only one equation. For the data analysis performed with the SIT equations, together with the ionic strength dependence terms in eq 7, the well-known van’t Hoff equation was added, and the term L(I, T) of eq 4 is:

L(I , T ) = I(a0 + a1NCR + (a 2 + a3 NCR) ·I ) + F(T ) (10)

⎛ ⎞ ΔcpΔT ⎛ 1 1⎞ F(T ) = ⎜ΔH 0 + − Δε′I ⎟ ·52.23·⎜ − ⎟ ⎝ 298.15 1000 T⎠ ⎝ ⎠ (10a)

Δε′ = b0 + b1NCR

⎜

(8) −1

where ΔH is the enthalpy change at I = 0 mol·kg and T = 298.15 K, Δcp and Δε′ are the dependence parameters of ΔH0 on temperature and ionic strength, respectively, I is expressed in the 0

(10b)

where NCR is the normalized ionic radius, a0, a1, a2, and a3 are fitting parameters equal for all the ionic media, and F(T) is the temperature dependence function, whose ionic strength dependence is in turn a function of the NCR. Models for the Variation of the Activity Coefficients: Pitzer Equations. The experimental data in the molal concentration scale were also analyzed using the Pitzer approach.38,39 According to this model, all of the interactions occurring in solution are considered, even those between ions of the same charge (indicated as θ) and triple interaction (indicated

⎛ ⎞ ΔcpΔT − Δε′I ⎟ ·52.23· L(I , T ) = ΔεI + ⎜ΔH 0 + 1000 ⎝ ⎠ ⎛ 1 1⎞ − ⎟ ⎝ 298.15 T⎠

(9a)

1

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as ψ) such as (+, −, +) or (−, +, −). In all cases, these parameters are often affected by high correlation, and they are neglected if the system is sufficiently simple or for I < 5 mol·kg−1.22 Further details can be found in Pitzer.39 In this work, the triple interaction parameters (ψ) were neglected, whereas some of the homocharged parameters were considered. In addition, considering the big amount of experimental data, the temperature dependence of the Pitzer parameters was proposed for the first time for pyridine. In fact, to our knowledge no papers were published reporting Pitzer calculations on this molecule. The general equation for Pitzer model used in this work is reported in eq 11: H

log K = log K

H0

+ F(p)/ln(10)

(C2H5) 4NI, we also have significant interaction of the unprotonated amine (py) with the tetraalkylammonium cation.41−44 If the formation of both the weak complexes pyHCl and (CH3)4Npy+ (for simplicity we use (CH3)4N+ as cation in this discussion, but the same considerations are valid also for (C2H5)4N+) is taken into account, the mass balance equations for pyridine and proton can be written as: [py]T = [py 0] + [Hpy +] + [(CH3)4 Npy +] + [pyHCl0] (15)

[H]T = [H+] + [Hpy +] + [pyHCl0]

(11)

where the subscript T indicates the analytical total concentration, and equilibrium constants for the formation of weak species are:

where F(p) = (p1 + ∂p1 /∂T ·ΔT ) ·2I + (p2 + ∂p2 /∂T ·ΔT ) ·I 2 + (p3 + ∂p3 /∂T ·ΔT ) ·f

K Cl =

[pyHCl0] [pyH+][Cl−]

(17)

(12)

f = [1 − (1 + 2 I ) ·g ]

(12a)

g = exp( −2 I )

(12b)

K (CH3)4 N =

[(CH3)4 Npy +] [(CH3)4 N+][py 0]

[py]T = [py 0]* + K *[py 0]*[H+] = [py 0]*(1 + K *[H+])

(13)

ϕ ϕ p2 = C HX − C HpyX

(13a)

(1) (1) p3 = βHX − βHpyX

(13b)

(18)

Analogously, when considering apparent protonation constants, eqs 15 to 16 become:

p1, p2, and p3 are the Pitzer fitting parameters and are, in turn, a function of Pitzer parameters: (0) (0) p1 = βHX + λpy − βHpyX + θHM + θHMpy

(16)

(19)

[H]T = [H+] + K *[py 0]*[H+]

(20)

where asterisks (*) indicate conditional quantities. Using the two different approaches, the average number of protons bound to the ligand, p̅, can be obtained combining and rearranging eqs 15 to 18 and 19 to 20

for 1:1 supporting electrolyte salts MX, where β(0), λ, θ, Cϕ, and β(1) are Pitzer parameters specific for the single interaction. Some of them, regarding the proton and the background salt, are known and are reported in Pitzer.39 For example, in the case of the NaCl, the single parameters can determined. In fact, knowing that there is a relationship between km (py) and λpy and using the Setschenow coefficient of the neutral species, as in the case of the SIT model, eq 13 can be solved for β(0) HpyX. For this purpose, the parameter θHMpy has been neglected.

p̅ =

K H[H+] + K HK Cl[H+][Cl−] 1 + K H[H+] + K (CH3)4 N[(CH3)4 N+] + K HK Cl[H+][Cl−] (21)

p̅ * =

k ln(10) λpy = m (14) 2 At the same time, eqs 13a to 13b can be solved knowing the Pitzer parameters relative to the interaction of proton with the anion of the supporting electrolyte. Formation of Weak Complexes. As cited above, together with the models involving the variation of the activity coefficients, the ionic strength dependence of the protonation constant in various ionic media can be explained with the formation of weak complexes between the ligand species and the ions of the supporting electrolyte. In particular, pyridine forms weak complexes between its protonated species and the anion of the supporting electrolyte (e.g., Cl−) and between the deprotonated species and the cation of the supporting electrolyte (e.g., (CH3)4N+). Therefore, the protonation constants of amines, measured in aqueous solutions containing a supporting electrolyte, must be considered as conditional constants. As an example, in NaCl only the interaction between chloride and the protonated amine species must be considered, as the sodium cation does not interact significantly with amines.40 In the case of (CH3)4NCl or

K *[H+] 1 + K *[H+]

(22)

the right-hand side of eqs 21 and 22 must be equal and therefore K H[H+] + K HK Cl[H+][Cl−] 1 + K H[H+] + K (CH3)4 N[(CH3)4 N+] + K HK Cl[H+][Cl−] =

K *[H+] 1 + K *[H+]

(23)

the simultaneous analysis of conditional protonation constants in different ionic media (NaCl, (CH3)4NCl, (C2H5)4NI, MgCl2, and CaCl2) at different ionic strengths and temperatures allows us to obtain the formation constants for the weak species by 25 minimizing the squares sum of the differences p̅ − p*: ̅

∑ (p ̅

− p ̅ *)2

(24)

In this analysis the protonation data in various ionic media were included in the ionic strength range 0 < I/mol·dm−3 ≤ 1.5 and in the temperature range 283.15 < T/K ≤ 318.15. According to the model which takes into account the formation of weak complexes, the ionic strength and temperature dependence of a generic equilibrium constants is calculated using the equation: 150



Article

∂ log K θ0 ∂ 2 log K θ0 ·ΔT + 2· ·ΔT 2 2 ∂T ∂T I 3/2 − z*· + CI + DI (25) 2+3 I

log K θ = log K θ0 +

where z* =

∑ (charges)2reag

−

∑ (charges)2prod

(26)

K0θ

Kθ and can be the protonation constant or the weak complex formation constants at different ionic strengths and temperatures and at infinite dilution and T = 298.15 K, respectively. The partial derivatives take into account the temperature dependence of log K0θ, ΔT = (T − 298.15), and the parameters C and D are empirical parameters for the ionic strength dependence which can be expressed as follows: C = (c0 + ∂c0/∂T ·ΔT ) ·p* + (c1 + ∂c1/∂T ·ΔT ) ·z* (27)

D = d0p* + d1z*

(27a)

with p* =

∑ preact

−

∑ pprod

(28)

where p is the stoichiometric coefficient, c0, c1, d0, and d1 are empirical parameters, and the partial derivatives are the temperature gradients of these parameters.

■

Figure 2. Trend of the pyridine protonation constant in different ionic media at T = 298.15 K. 1, LiCl; 2, NaCl; 3, KCl; 4, RbCl; 5, CsCl; 6, MgCl2; 7, (C2H5)4NI.

In tetraalkylammonium halides, in CaCl2 and in MgCl2, pyridine interacts with both halides (in the protonated form) and the cations (in the deprotonated form), whereas in the other ionic media, the interaction is noticed only between chloride and protonated pyridine species. For alkali metal halides the trend traces the dimension of the cation, except for NaNO3, which can hardly be compared with these ionic media. In Figure 2, the ionic strength dependence of the apparent pyridine protonation constant is reported at T = 298.15 K. The behavior is completely different for the tetraalkylammonium and the alkali metal salts. The analysis of the experimental data was performed using different models, and the results, described in this section, are highly comparable in terms of statistical parameters. In particular, the results obtained using the data in the molar concentration scale, applying the Debye−Hückel equation are listed in Table 12 with log KH0 = 5.218 ± 0.001. According to SIT equations, the fit of the experimental data to eqs 4 and 8 to 9a allowed us to determine the pyridine protonation constant at infinite dilution, constrained for all ionic media, and the Δε values in the different ionic media. For NaCl medium at T = 298.15 K it was possible to calculate the

interaction coefficient between the Hpy+ and the Cl− species. In fact, applying eq 7, considering the average km values of azine (km = 0.184 + 0.025 ln(I + 1)) calculated from ref 17 and knowing the specific interaction coefficient of proton in chloride media33 (ε(H+, Cl−) = (0.136 + (−0.0521)/(mNaCl + 1))) it was possible to determine ε(Hpy+,Cl−) ± 0.006 = −0.0366 + 0.0799 ln(mNaCl + 1). Fitting the data in the other ionic media to eq 8 it was possible to determine the SIT parameters Δε and the temperature dependence parameters ΔH0, Δcp, and Δε′ (see eq 8). The ΔH0 and the Δcp values were fixed to be equal for all ionic media, whereas the Δε′ values were determined separately for each ionic medium. Different tries were then performed, for example using eq 9 or eq 9a to express Δε and considering only the data at I < 2 mol·kg−1. In all cases, Δε0 values were constrained for all the chloride media. The results obtained are very interesting, in fact if only the data at I < 2 mol·kg−1 are considered the statistical parameters (as standard deviation of the whole fit) are practically identical (σfit = 0.015), if we use eq 9 or 9a to express Δε. On the contrary, if we use all of the data, up to I ∼ 5.5 mol·kg−1, the statistical data are very different and are σfit = 0.060 and 0.020 using eqs 9 and 9a, respectively. In practice, when the data are considered in a narrow ionic strength range the fit is substantially independent of the model used, while the fit is strongly dependent on the model when the data are considered in the whole ionic strength range. The best model, in terms of standard deviation of the whole fit, was obtained when Δε = Δε0 + Δε1 ln(1 +I). As a test, the thermodynamic parameters of pyridine obtained from data in MgCl2 and CaCl2 (defined as set A) were refined separately from those of the other ionic media (set B). The values obtained for the two sets are highly comparable and demonstrate the robustness of the data treatment. For set A: log KH0 = 5.223 ± 0.003, ΔH0 = −21.6 ±

RESULTS AND DISCUSSION As a general trend, the protonation constant of pyridine, in all of the considered ionic media, decreases with increasing temperature (indicating negative enthalpy values) and increases with increasing ionic strength, although for the tetraalkylammonium halides this latter behavior is very slight. At T = 298.15 K, the following trend for the apparent protonation constant can be explained considering the interaction of pyridine with the ions of the supporting electrolyte: (C2H5)4 NI < (CH3)4 NCl ≪ CaCl 2 < MgCl2 < CsCl < RbCl < NaNO3 < KCl < NaCl < LiCl

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Table 12. Ionic Strength and Temperature Dependence Parameters According to the Debye−Hückel Approach, Equations 5 to 6 salt

C

LiCl NaCl NaNO3 KCl RbCl CsCl (CH3)4NCl (C2H5)4NI

0.305 ± 0.002a 0.305 ± 0.002 0.305 ± 0.002 0.305 ± 0.002 0.305 ± 0.002 0.305 ± 0.002 0.047 ± 0.019 0.084 ± 0.004

MgCl2 CaCl2 a

E 0.005 ± 0.001a 0.001 ± 0.001 −0.015 ± 0.001 −0.027 ± 0.001 −0.036 ± 0.001 −0.016 ± 0.001 −0.073 ± 0.026 −0.037 ± 0.002 B0

C

E

0.153 ± 0.008 0.183 ± 0.007

0.031 ± 0.008 0.012 ± 0.007

A0

A1

A2

−0.012 ± 0.001a −0.012 ± 0.001 −0.012 ± 0.001 −0.012 ± 0.001 −0.012 ± 0.001 −0.012 ± 0.001 −0.012 ± 0.001 −0.012 ± 0.001 B1

−0.011 ± 0.004a −0.011 ± 0.004 −0.011 ± 0.004 −0.011 ± 0.004 −0.011 ± 0.004 −0.011 ± 0.004 −0.011 ± 0.004 −0.011 ± 0.004 B2

−0.001 ± 0.001a −0.002 ± 0.001 −0.002 ± 0.001 −0.002 ± 0.001 −0.002 ± 0.001 −0.005 ± 0.001 −0.002 ± 0.001 −0.001 ± 0.001 B3

−0.011 ± 0.001 −0.011 ± 0.001

0.000015 ± 0.000001 0.000015 ± 0.000001

−0.0049 ± 0.0006 −0.0030 ± 0.0006

0.0040 ± 0.0006 a 0.0040 ± 0.0006

± 95 % C.I.

Table 13. Ionic Strength and Temperature Dependence Parameters According to the SIT Approach, Equations 8 to 9

a

log KH0

Δε0

LiCl NaCl NaNO3 KCl RbCl CsCl (CH3)4NCl (C2H5)4NCl MgCl2 CaCl2

0.310 ± 0.005 0.310 ± 0.005 0.330 ± 0.010 0.310 ± 0.005 0.310 ± 0.005 0.310 ± 0.005 −0.021 ± 0.018 0.044 ± 0.004 0.105 ± 0.017 0.105 ± 0.017 a

Δε1

Δε′

σfit

mdfit

−0.012 ± 0.003a −0.027 ± 0.003 −0.080 ± 0.006 −0.057 ± 0.004 −0.085 ± 0.004 −0.107 ± 0.003 −0.089 ± 0.030 −0.059 ± 0.003 0.053 ± 0.012 0.056 ± 0.012

−2.5 ± 0.3a −3.9 ± 0.3 −7.0 ± 0.3 −4.1 ± 0.3 −4.9 ± 0.6 −1.9 ± 0.4 −2.1 ± 0.6 −0.5 ± 0.3 −1.3 ± 0.6 1.5 ± 0.5

0.018 0.014 0.017 0.012 0.040 0.026 0.008 0.013 0.025 0.019

0.014 0.008 0.014 0.007 0.032 0.018 0.006 0.010 0.020 0.014

± 95 % C.I.

Table 14. Pitzer Parameters of Equations 13 to 13b Used in This Work and Determined for NaCl at T = 298.15 K i

−

H H+ Hpy+ a

Cϕij

β(0) ij

j

+

Cl Na+ Cl−

0.1775

a

0.008

β(1) ij

a

θij

0.2945a 0.036a

0.146 ± 0.015b

0.017 ± 0.005b

0.188 ± 0.038b

Taken from ref 39. b± 95 % C.I.

Table 15. Pitzer Parametersa Determined in This Work ionic medium

p1

LiCl NaCl KCl RbCl CsCl NaNO3 (CH3)4NCl (C2H5)4NI MgCl2 CaCl2

0.392 ± 0.022 0.314 ± 0.016 0.368 ± 0.024 0.540 ± 0.036 0.380 ± 0.036 0.304 ± 0.020 −0.078 ± 0.014 0.028 ± 0.134 0.556 ± 0.064 0.783 ± 0.068

a

∂p2

(∂p1/∂T) b

−0.0038 ± 0.0016 −0.0020 ± 0.0016 −0.0020 ± 0.0020 −0.0129 ± 0.0032 −0.0068 ± 0.0026 −0.0144 ± 0.0038 0.0009 ± 0.0120 −0.0001 ± 0.0102 −0.0240 ± 0.0050 −0.0079 ± 0.0048

b

∂p3

(∂p2/∂T)

−0.017 ± 0.006 −0.007 ± 0.006 −0.055 ± 0.012 −0.132 ± 0.014 −0.094 ± 0.012 −0.038 ± 0.006 −0.007 ± 0.004 −0.141 ± 0.126 0.150 ± 0.102 −0.143 ± 0.110

b

0.0008 ± 0.0006 −0.0002 ± 0.0004 −0.0005 ± 0.0010 0.0039 ± 0.0012 0.0020 ± 0.0010 0.0024 ± 0.0008 −0.0002 ± 0.0004 0.0015 ± 0.0098 0.0377 ± 0.0082 0.0229 ± 0.0082 b

−0.187 ± 0.050 0.043 ± 0.038 −0.126 ± 0.050 −0.729 ± 0.092 −0.258 ± 0.088 0.084 ± 0.070 0.251 ± 0.046 −0.217 ± 0.254 0.005 ± 0.084 −0.242 ± 0.088

(∂p3/∂T) b

0.0016 ± 0.0042b −0.0039 ± 0.0040 −0.0031 ± 0.0042 0.0196 ± 0.0074 0.0102 ± 0.0060 0.0343 ± 0.0158 −0.0083 ± 0.0038 −0.0084 ± 0.0180 0.0160 ± 0.0064 −0.0023 ± 0.0070

Parameters of eqs 12 to 13b. b± 95 % C.I.

0.3 kJ·mol−1, and Δcp = 12 ± 24 J·mol−1·K−1; for set B: log KH0 = 5.215 ± 0.002, ΔH0 = −20.1 ± 0.2 kJ·mol−1, and Δcp = −101 ± 9 J·mol−1·K−1. These results are in a good agreement, except for the Δcp values, which in any case, has not a great influence in the considered temperature range (278.15 ≤ T/K ≤ 318.15). In any case, from a thermodynamic point of view the infinite dilution values are equal for all ionic media, and therefore the data were then fitted altogether. In this case, the results are those of set B; therefore, log KH0 = 5.215 ± 0.002, ΔH0 = −20.1 ± 0.2 kJ·mol−1,

and Δcp = −101 ± 9 J·mol−1·K−1. The other temperature and ionic strength dependence parameters for all ionic media are reported in Table 13 together with the statistical results for standard deviation (indicated as σfit) and mean deviation (indicated as mdfit) for the whole fit. Looking at the results in Table 13, it was found that the values of Δε1 and Δε′, for the chlorides of alkali metals have a dependence on the normalized crystallographic radius (NCR), 152



Article

Table 16. Thermodynamic Parameters for the Protonation of Pyridine in NaCl at Different Temperatures and Ionic Strengths (in the Molal Concentration Scale)

a

I/mol·kg−1

T/K

−ΔGa ± 95 % C.I.

ΔHa ± 95 % C.I.

−TΔSb ± 95 % C.I.

0 0.1 0.5 1.0 3.0 5.0 0 0.1 0.5 1.0 3.0 5.0 0 0.1 0.5 1.0 3.0 5.0 0 0.1 0.5 1.0 3.0 5.0

283.15 283.15 283.15 283.15 283.15 283.15 298.15 298.15 298.15 298.15 298.15 298.15 310.15 310.15 310.15 310.15 310.15 310.15 323.15 323.15 323.15 323.15 323.15 323.15

29.19 ± 0.01 29.38 ± 0.01 30.11 ± 0.01 30.98 ± 0.02 34.23 ± 0.06 37.27 ± 0.12 29.76 ± 0.01 29.94 ± 0.01 30.62 ± 0.01 31.44 ± 0.01 34.44 ± 0.03 37.23 ± 0.07 30.10 ± 0.01 30.27 ± 0.01 30.92 ± 0.01 31.69 ± 0.01 34.50 ± 0.03 37.08 ± 0.05 30.35 ± 0.02 30.51 ± 0.02 31.13 ± 0.02 31.85 ± 0.03 34.45 ± 0.06 36.81 ± 0.09

−18.6 ± 0.5 −19.0 ± 0.5 −20.5 ± 0.7 −22.5 ± 1.2 −30.3 ± 3.3 −38.2 ± 5.5 −20.1 ± 0.4 −20.5 ± 0.4 −22.1 ± 0.7 −24.0 ± 1.2 −31.9 ± 3.3 −39.7 ± 5.5 −21.3 ± 0.5 −21.7 ± 0.5 −23.3 ± 0.7 −25.2 ± 1.2 −33.1 ± 3.3 −40.9 ± 5.5 −22.6 ± 0.6 −23.0 ± 0.6 −24.6 ± 0.8 −26.6 ± 1.3 −34.4 ± 3.3 −42.2 ± 5.5

−10.6 ± 0.5 −10.4 ± 0.5 −9.6 ± 0.7 −8.5 ± 1.2 −3.9 ± 3.3 0.9 ± 5.5 −9.7 ± 0.4 −9.4 ± 0.4 −8.6 ± 0.7 −7.4 ± 1.2 −2.6 ± 3.3 2.5 ± 5.5 −8.8 ± 0.5 −8.6 ± 0.5 −7.6 ± 0.7 −6.4 ± 1.2 −1.4 ± 3.3 3.8 ± 5.5 −7.7 ± 0.6 −7.5 ± 0.6 −6.5 ± 0.8 −5.3 ± 1.3 0.0 ± 3.3 5.4 ± 5.5

In kJ·mol−1. bIn kJ·mol−1·K−1.

therefore, together with the analysis performed with eqs 4 and 8 to 9a, the data were fitted to eqs 4 and 10 to 10b. In this fit all of the data were considered, although those in CsCl are outliers according to results in Table 13. The refined parameters are: a0 = 0.284 ± 0.003, a1 = −0.014 ± 0.002, a2 = −0.017 ± 0.001, a3 = −0.012 ± 0.001, b0 = −3.43 ± 0.09, and b1 = −0.32 ± 0.10. The other fitting parameters are in very good agreement with the results reported above, in fact log KH0 = 5.220 ± 0.001, ΔH0 = −20.2 ± 0.2 kJ·mol−1, and Δcp = −100 ± 10 J·mol−1·K−1. In Capone et al.18 the infinite dilution parameters were log KH0 = 5.207, ΔH0 = −20.06 kJ·mol−1, and Δcp = −36 J·mol−1·K−1. In Martell et al.45 log KH0 = 5.20 and ΔH0 = −20 kJ·mol−1. The analysis of the experimental data in Tables 2 to 11 has been also performed with the Pitzer approach.38,39 Also in this case, for the NaCl medium (at T = 298.15 K) it was possible to determine the Pitzer coefficient of the single interaction of the pyridine species with the counterion of the ionic medium (Cl−). In the case of the other ionic media, it was possible to determine only the cumulative Pitzer coefficients p1, p2, and p3 of eqs 13 to 13b. In addition, by fitting the data at different temperatures to eqs 11 to 12 it was possible to determine the temperature gradient of p1, p2, and p3 in all the ionic media. The Pitzer coefficients used for calculations and determined for the data in NaCl at T = 298.15 K are listed in Table 14, whereas the data for the other ionic media and temperature gradients of the Pitzer coefficients are summarized in Table 15. The whole set of thermodynamic parameters was calculated, and for simplicity these data are shown in Table 16 at different temperatures and ionic strengths only for NaCl. In Figure 3, the three thermodynamic functions are reported against ionic strength in the molal concentration scale. To draw the three functions in the same graph, we reported difference between the

Figure 3. Thermodynamic functions for the protonation of pyridine in NaCl at different ionic strengths and T = 298.15 K. △, TΔS − TΔS0 (kJ· mol−1·K−1); □, ΔG − ΔG0 (kJ·mol−1); ○, ΔH − ΔH0 (kJ·mol−1).

value at a given ionic strength and the value at infinite dilution (ΔG0, ΔH0, TΔS0). It can be observed that the increase of the enthalpic contribution with increasing of the ionic strength is higher than that of ΔG and TΔS. In Capone et al.,18 ΔS0 = 32.4 J· mol−1·K−1 was reported; this value is not significantly different from ΔS0 = 32.5 ± 0.5 J·mol−1·K−1 found in this work. The results obtained with the model that takes into account the formation of weak complexes were satisfactory. The analysis was performed considering the data in NaCl, (CH3)4NCl, (C2H5)4NI, MgCl2, and CaCl2 in the ionic strength range 0 < I/ mol·dm−3 ≤ 1.5 and in the temperature range 283.15 < T/K ≤ 318.15 and evidenced the formation of five weak species, namely, 153



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Table 17. Parameters of the Weak Complex Formation Model equilibrium +

+

H + py = Hpy H+ + py + Cl− = HpyCl (CH3)4N+ + py = (CH3)4Npy+ (C2H5)4N+ + py = (C2H5)4Npy+ Mg2+ + py = Mgpy2+ Ca2+ + py = Capy2+ a

z* 0 2 0 0 0 0

p*

log K0θa

1 2 1 1 1 1

5.220 ± 0.002 4.94 ± 0.03 −0.42 ± 0.02 −0.36 ± 0.02 −0.10 ± 0.01 −0.15 ± 0.01

(∂ log K0θ/∂T)a b

−0.0114 ± 0.0001 −0.020 ± 0.002 −0.014 ± 0.002 −0.0081 ± 0.0012 −0.0022 ± 0.0006 −0.0022 ± 0.0007

2(∂2 log K0θ/∂T2)a b

0.000012 ± 0.000006b

Parameters of eq 25. b95 % C.I.

(log KCl = 5.393), and (CH3)4Npy+ (log K(CH3)4N = −0.247). The weak species in these conditions achieve ∼0.3 molar fraction at pH < 5 for HpyCl and pH > 5 for (CH3)4Npy+.

HpyCl, (CH3)4Npy+, (C2H5)4Npy+, Mgpy2+, and Capy2+. The statistical parameters obtained for the whole fit were mdfit = 0.0012 and σfit = 0.0017. The analysis performed with the SIT equations in the same experimental conditions (with data in the molal concentration scale) resulted comparable as regards the statistical parameters and md = 0.001. The values of c0 and c1 were in agreement with previous findings (d0 was set to zero and d1 was fixed at −0.1),18 c0 = 0.173 ± 0.009 and c1 = 0.148 ± 0.010. The temperature dependence parameters are ∂c0/∂T = −0.0013 ± 0.0004 and ∂c1/∂T = 0.00087 ± 0.0003. The equilibrium constants of the weak species and their dependence on temperature are reported in Table 17. To show the importance and the distribution of the weak species, in Figure 4 the

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CONCLUSIONS In conclusion, the acid−base properties of pyridine were determined in 10 ionic media at different temperatures and ionic strengths, and the new data reported in this work confirmed the findings of Capone et al.18 As expected for an azine the protonation constant increases with increasing ionic strength (in tetraalkylammonium halide this behavior is very slight) and decreases with increasing temperature. This means that the proton binding is exothermic and the enthalpy value at infinite dilution and T = 298.15 K is −20.1 kJ·mol−1, in a very good agreement with literature data in refs 16 and 45 to 47 (e.g., −20 kJ·mol−1 in the same conditions45). The literature data are in good agreement with the data here reported. For example, in the NIST critical database, at 298.15 K and I = (0, 0.1, 0.5, and 1.0) mol·L−1, log KH = 5.200, 5.240, 5.340, and 5.480, respectively. In the same conditions, we have log KH = 5.217, 5.250, 5.359, and 5.494; therefore the highest difference is of 0.019 log K units at I = 0.5 mol·L−1. It is important to stress that pyridine protonation is strongly dependent on the ionic medium. In Table 18 the calculated protonation constant values are reported in different ionic media and at different ionic strengths at T = 298.15 K, and it can be noted that at I = 0.1 mol·kg−1 the values are substantially equal with only small differences; at I = 1.0 mol·kg−1 the discrepancies are more important, and the difference between (C2H5)4NI (lowest value) and LiCl (highest value) is ∼0.4 log K units. For higher ionic strengths the differences are major and are ∼1.5 log K units (I = 3.0 mol·kg−1) and > 1.7 log K units for I = 5.0 mol·kg−1. The data analysis was performed considering different models to interpret the ionic strength dependence, namely, the Debye− Hückel, the SIT, and the Pitzer approaches. The modeling performed with the three models resulted comparable, although

Figure 4. Speciation diagram of pyridine in (CH3)4NCl, molar fraction of pyridine vs pH. cpy = 0.005 mol·dm−1; I = 1 mol·dm−3; T = 298.15 K. Species: 1, Hpy+; 2, pyHCl; 3, py; 4, (C2H5)4Npy+.

speciation diagram of pyridine (0.005 mol·dm−3) is shown in (CH3)4NCl (I = 1 mol·dm−3) at T = 298.15 K, the calculated species in these conditions are Hpy+ (log KH = 5.393), HpyCl

Table 18. Calculated Pyridine Protonationa Constant in Different Ionic Media at Different Ionic Strengths and at T = 298.15 K

a

salt

I = 0.1 mol·kg−1

I = 1.0 mol·kg−1

I = 3.0 mol·kg−1

I = 5.0 mol·kg−1

LiCl NaCl KCl RbCl CsCl NaNO3 (CH3)4NCl (C2H5)4NI MgCl2 CaCl2

5.245 5.252 5.246 5.226 5.240 5.253 5.230 5.211 5.238 5.234

5.507 5.503 5.485 5.446 5.445 5.489 5.216 5.129 5.391 5.411

6.107 6.029 5.919 5.841 5.748 5.889 5.085 4.663 5.771 5.806

6.664 6.527 6.172 5.837 5.747 6.151 4.909 6.209 6.148

± 0.002 (95 % C.I.). 154



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(3) Eatough, D. J.; Benner, C. L.; Bayona, J. M.; Richards, G.; Lamb, J. D.; Lee, M. L.; Lewis, E. A.; Hansen, L. D. Chemical composition of environmental tobacco smoke. Environ. Sci. Technol. 1989, 23, 679−687. (4) Santodonato, J.; Bosch, S.; Meylan, W.; Becker, J.; Neal, M. Monograph on Human Exposure to Chemicals in the Workplace: Pyridine (Report No. SRC-TR-84-1119); Syracuse Research Corporation, Center for Chemical Hazard Assessment: Syracuse, NY, 1985. (5) Scriven, E. F. V.; Toomey, J. E. J.; Murugan, R. Pyridine and pyridine derivatives. In Kirk-Othmer Encyclopedia of Chemical Technology, 4th ed.; Kroschwitz, J. I., Howe-Grant, M., Eds.; John Wiley: New York, 1996; Vol. 20, pp 641−679. (6) Shimizu, S.; Watanabe, N.; Kataoka, T.; Shoji, T.; Abe, N.; Morishita, S.; Ichimura, H. Pyridine and pyridine derivatives. In Ullmann’s Encyclopedia of Chemical Technology, 5th rev. ed.; Elvers, B., Hawkins, S., Russey, W., Schulz, G., Eds.; VCH Publishers: New York, 1993; Vol. A22, pp 399−430. (7) National Toxicology Program. Toxicology and Carcinogenesis Studies of Pyridine (CAS No. 110-86-1) in F344/N Rats, Wistar Rats, and B6C3F1Mice (Drinking Water Studies). In National Toxicology Program Tech. Rep. Series 2000, 470, 1−330. (8) Damani, L. A.; Crooks, P. A.; Shaker, M. S.; Caldwell, J.; D’Souza, J.; Smith, R. L. Species differences in the metabolic C- and N-oxidation, and N-methylation of [14C]pyridine in vivo. Xenobiotica 1982, 12, 527− 534. (9) Harper, B. L.; Legator, M. S. Pyridine prevents the clastogenicity of benzene but not of benzo[a]pyrene or cyclophosphamide. Mutat. Res. 1987, 179, 23−31. (10) Jori, A.; Calamari, D.; Cattabeni, F.; Di Domenico, A.; Galli, C. L.; Galli, E.; Silano, V. Ecotoxicological profile of pyridine. Ecotoxicol. Environ. Saf. 1983, 7, 251−275. (11) Chirico, R. D.; Steele, W. V. 2. Comparison of new recommended values with the literature. J. Chem. Thermodyn. 1996, 28, 819−841. (12) Chirico, R. D.; Steele, W. V.; Nguyen, A.; Klots, T. D.; Knipmeyer, S. E. 1. Vapor pressures, high-temperature heat capacities, critical properties, derived thermodynamic functions, vibrational assignment, and derivation of recommended values. J. Chem. Thermodyn. 1996, 28, 797−818. (13) Dragelj, J. L.; Janjić, G. V.; Veljković, D. Z.; Zarić, S. D. Crystallographic and ab initio Study of Pyridine CH/O Interactions. Linearity of the interactions and influence of pyridine classical hydrogen bonds. CrystEngComm 2013, 15, 10481−10489. (14) Ghatee, M. H.; Fotouhabadi, Z.; Zolghadr, A. R.; Borousan, F.; Ghanavati, F. Structural and phase behavior studies of pyridine and alkyl pyridine at the interface of oil/water by molecular dynamic simulation. Ind. Eng. Chem. Res. 2013, 52, 13384−13392. (15) Marsicano, F.; Hancock, R. D. The linear free-energy relation in the thermodynamics of complex formation. Part 2. Analysis of the formation constants of complexes of the large metal ions silver(I), mercury(II), and cadmium(II) with ligands having “soft” and nitrogendonor atoms. J. Chem. Soc., Dalton Trans. 1978, 228−234. (16) Ashton, L. A.; Bullock, J. I.; Simpson, P. W. Effect of Temperature on the Protonation Constants of Some Arimatic, Heterocyclic Nitrogen Bases and the Anion of 8-Hydroxyquinoline. J. Chem. Soc., Faraday Trans. 1 1982, 78, 1961−1970. (17) Bretti, C.; Crea, F.; De Stefano, C.; Sammartano, S. Solubility and Activity Coefficients of 2,2′-Bipyridyl, 1,10-Phenanthroline and 2,2′,6′,2″-Terpyridine in NaCl(aq) at Different Ionic Strengths and T = 298.15 K. Fluid Phase Equilib. 2008, 272, 47−52. (18) Capone, S.; Casale, A.; Currò, A.; De Robertis, A.; De Stefano, C.; Sammartano, S.; Scarcella, R. The Effect of Background on the Protonation of Pyridine in Aqueous LiCl, NaCl, KCl, RbCl, CsCl, CaCl2, MgCl2, (CH3)4NCl and (C2H5)3NI Solutions at Different Temperatures and Ionic Strengths. Ann. Chim. (Rome) 1986, 76, 441− 472. (19) Casale, A.; De Robertis, A.; Licastro, F. The Effect of Background on the Protonation of Pyridine: A Complex Formation Model. Thermochim. Acta 1989, 143, 289−298. (20) Flaschka, H. A. EDTA Titration; Pergamon: London, 1959.

the results obtained with the Pitzer model can be considered of higher quality, in terms of the statistical parameters standard deviation and mean deviation in the whole fit. The analysis of the thermodynamic parameters in Table 16 demonstrates that the proton binding is enthalpic driven, and the values of the TΔS contribution decrease with increasing both temperature and ionic strength, becoming positive at I > 3 mol· kg−1. The results reported in this work are in a very good agreement with literature data, for example the log KH0 (but also data at higher ionic strengths) and the ΔH0 values are substantially equal to the data in the most common stability constant databases.45−47 The weak complexes between pyridine and the divalent metal cation are slightly higher (∼ 0.2 log K units) than the literature findings. As seen for other amines17,48 at T > 298.15 K, the (C2H5)4N+ cation (log K = −0.27 at T = 298.15 K and I = 0.5 mol·dm−3) forms a more stable species with deprotonated pyridine, although slightly, than the (CH3)4N+ (log K = −0.34 at T = 298.15 K and I = 0.5 mol·dm−3). In the same path, for other azines (terpy, bypy, and phen), the best way to express the Δε as a function of ionic strength is with eq 10a. Some discrepancies are present between the values of the weak (CH3)4Npy+ and (C2H5)4Npy+ species, which in the case of pyridine resulted lower than for terpyridine, but in those papers,17 the HpyCl complexes was not quantified; therefore a rigorous comparison cannot be made. Similar values were also found between the Pitzer parameters (p1, p2, and p3) obtained in this work for pyridine (see Table 15) and reported by Bretti et al.17 for terpy, bypy, and phen. The protonation constant of imidazole is higher than that of the pyridine as well as the values of the weak complexes with Mg2+, Ca2+, (CH3)4N+, and (C2H5)4N+. On the contrary, the stability of the protonated amino complexes with chloride is very similar. In the literature only few papers deal with the variation of the apparent protonation constants with the alkali metal series18,48−52 (Li+, Na+, K+, Rb+, and Cs+). Many ionic strength dependence parameters were found to be dependent on the crystallographic radius of the cation of the supporting electrolyte, with a linear decreasing dependence down along the alkali metal group (from Li+ to Cs+). In conclusion it is reasonable to consider this work as an improvement in the knowledge and in the understanding of the thermodynamic functions and of the solution chemistry of pyridine.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +39-090-6765749. Fax: +39-090-392827. Funding

We thank University of Messina for partial financial support. Notes

The authors declare no competing financial interest.

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REFERENCES

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Thermodynamics for Proton Binding of Pyridine in Different Ionic

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