Protonation Constants, Activity Coefficients, and Chloride Ion Pair

Article pubs.acs.org/jced

Protonation Constants, Activity Coefficients, and Chloride Ion Pair Formation of Some Aromatic Amino-Compounds in NaClaq (0 mol·kg−1 ≤ I ≤ 3 mol·kg−1) at T = 298.15 K Clemente Bretti, Francesco Crea, Concetta De Stefano,* Silvio Sammartano, and Giuseppina Vianelli Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Università di Messina, Viale F. Stagno d'Alcontres, 31, I-98166 Messina, Italy S Supporting Information *

ABSTRACT: The protonation constants of five structurally different aromatic amino-compounds [1,2-phenylenediamine (L1), 4-N-(2-hydroxyethyl)2,4-diaminoanisole (L2), 2-(2,4-diaminophenoxyethanol) (L3), 4-5-diamino-1-N-(2hydroxyethyl)pyrazole (L4), 1,3-bis(2,4-diaminophenoxy)propane (L5)] obtained by potentiometric measurements in NaCl(aq) (0 mol·kg−1 ≤ I ≤ 3 mol·kg−1) at T = 298.15 K, are reported. The Debye−Hückel, SIT (Specific ion Interaction Theory), and Pitzer type equations were used to determine the parameters for the dependence of the protonation constants on ionic strength and to calculate the specific interaction coefficients of the ionic species, as well as the protonation constants at infinite dilution. Distribution measurements were also carried out at different ionic strengths in 2-methyl-1-propanol/aqueous solution mixtures to obtain the distribution coefficients of the aromatic-amino compounds. The Setschenow coefficient and the activity coefficient of the neutral species were calculated by using a simple approach from distribution data. The dependence on salt medium was also explained in terms of weak polyammonium cation−chloride complexes.

■

INTRODUCTION This work reports a study on the protonation of some aromatic amino-compounds (see Figure 1) widely used in different

dermatitis due to the 1,4-benzenediamine and nefrotoxic effects due to aminophenols. Some dye intermediates are even carcinogenic. These studies1−9 have revealed an association between occupational exposure to these dyes and incidence of cancers. For example, some studies demonstrated that 1,2phenylenediamine is not only toxic, but also a potential carcinogenic and a possible sensitizer. For these reasons, interest in the development of new analytical methods to separate and analyze dye intermediates and to better define the thermodynamic properties of such ligands continues unabated. The dependence on ionic strength for the protonation constants of the five ligands here investigated, by ISE-[H+] potentiometry, in NaCl aqueous solutions at different ionic strengths, was modeled using the Debye−Hückel,10 SIT11,12 (Specific ion Interaction Theory), and Pitzer type equations.13,14 Distribution measurements in 2-methyl-1-propanol/ aqueous solutions at different ionic strengths were performed to study the distribution of these ligands between the two phases and to determine the Setschenow coefficient and the activity coefficient of the neutral species. Literature reports protonation constants for the 1,2-phenylenediamine and some derivatives (Table 1),15−18 but no information about their dependence on ionic strength and ionic medium. Regarding the other aromatic amino-compounds here

Figure 1. Structure of the ligands.

industrial fields as dye precursors in the manufacturing of hair dyes or as components of engineering polymers and composites. For example, 1,2-phenylenediamine is used in the synthesis of fungicides, corrosion inhibitors, pigments, and pharmaceuticals; it is also suitable for removing elemental sulfur in mining ores and aldehyde color formers in polymeric products. Many epidemiologic studies1−9 are reported in the literature regarding these compounds, since they are dye intermediates of great health concern because of allergic © 2012 American Chemical Society

Received: March 22, 2012 Accepted: May 2, 2012 Published: May 7, 2012 1851

dx.doi.org/10.1021/je300359s | J. Chem. Eng. Data 2012, 57, 1851−1859

Journal of Chemical & Engineering Data

Article

809 titrando) and software (Metrohm TiAMO 1.0) for the automatic data acquisition. Procedure. The protonation constants of each aromaticamine were determined by titrating with carbonate free standard NaOH, 25 mL of solutions, containing the ligand (CL = (5 to 15) mmol·L−1) and NaCl(aq) at different concentrations. For each experiment, the standard electrode potential E° and the junction potential coefficient ja (Ej = ja[H+]) were determined titrating a strong acid solution with a standard base, under the same medium and ionic strength conditions as the systems to be investigated. In this way, the free proton concentration scale was used, pH ≡ −log10[H+]. To check the reliability of the calibration in the alkaline range, the ionic product of water (pKw) was calculated. The distribution measurements of the aromatic-amines were performed by mixing and shaking an aliquot of 25 mL of pure water (or aqueous solution containing NaCl to adjust the ionic strength) containing the ligand at different concentrations (from (5 to 20) mmol·L−1) with 2-methyl-1-propanol at 298.15 K for at least four hours, in a separatory funnel. After separation of the two immiscible phases, the organic phase was opportunely diluted, and the concentration of the ligands was measured by potentiometric titrations. The aqueous phase was also checked for ligand concentration. Calculation. The ESAB2M19 computer program was used to refine all of the parameters (protonation constants, analytical concentration of reagents, formal electrode potential, junction potential coefficient, and ionic product of water) for the titrations of the ligands investigated in this paper. The LIANA computer program20 was used to calculate the parameters for the dependence of the protonation constants and the distribution parameters on ionic strength. The ES2WC computer program was used to calculate the formation constants of the weak protonated amine-Cl− complexes.21 The protonation constants, the concentrations of the ligands, and the ionic strength are expressed in the molar (c) or molal (m) concentration scales. The conversion from the molar to the molal (m, mol·kg−1) concentration scale was obtained using the equation, valid at T = 298.15 K: c/m = d0 + a1c + a2c2, with d0 = density of pure water, a1 = −0.017765, and a2 = −6.525·10−4.22 The molar concentration scale is generally used in analytical and inorganic chemistry, as it can be seen analyzing the thermodynamic data reported in the most common databases.15−18 The molal scale (independent of temperature) is widely used in physical chemistry, and in part, in engineering chemistry; for example, almost all the data on activity coefficients are reported in this scale.14,23,24

Table 1. Literature Data for the Protonation Constants of Some Phenylenediamine Derivatives15−18 at T = 298.15 K I/mol·L−1

ligand o-phenylenediamine

m-phenylenediamine p-phenylenediamine 4-methyl-o-phenylenediamine 4-methoxy-ophenylenediamine 4-methoxy-mphenylenediamine 2-methyl-p-phenylenediamine 1,3-di-(2′-aminophenoxy) propane 4-(5-chloro-2-pyridylazo)-1,3diaminobenzene 4-(2-pyridylazo)-1,3diaminobenzene a

log KH1 log βH2

0.1 KNO3 0.1 KNO3 0.2 NaClO4 0.1 NaClO4 0 0.1 KNO3 0.1 KNO3 0.05 KCla 0.1 KNO3 0.1 KNO3

4.63 4.50 4.70 4.63 4.69 4.89 6.06 6.0 4.79 5.10

0.1 KNO3

5.39

8.04

0.1 KNO3 95 % methanol/H2O, 0.1 Et4NClO4 0

6.02 4.98

8.70 9.05

5.40

6.70

0

6.10

9.20

6.32 5.23

7.29 8.93 8.73

T = 303.15 K.

investigated, in our knowledge the data reported in this paper are the first, both for the protonation constants and for the dependence on ionic strength, as well as the activity coefficients of the neutral and ionic species.

■

EXPERIMENTAL SECTION Chemicals. 1,2-Phenylenediamine (L1), 4-N-(2hydroxyethyl)2,4-diaminoanisole (L2), 2-(2,4-diaminophenoxyethanol)·2HCl (L3), 4-5-diamino-1N-(2-hydroxyethyl)pyrazole (L4), and 1,3-bis(2,4-diaminophenoxy)propane·4HCl (L5) (Fluka or Jos H. Lowenstein, NY, United States) were used without further purification; their purity checked alkalimetrically was always > 99 %. Since L3 and L5 were purchased in their hydrochloride form, they were neutralized with NaOH before distribution measurements. Sodium chloride solutions were prepared by weighing pure salt (Fluka, p.a.) previously dried in an oven at T = 383.15 K. Sodium hydroxide and hydrochloric acid solutions (Fluka p.a.) prepared from concentrated ampules were standardized against potassium hydrogen phthalate and sodium carbonate, respectively. 2Methyl-1-propanol was purchased from Fluka. Soda lime traps were used to preserve all of the solutions from atmospheric CO2. Grade A glassware and twice-distilled water were employed in the preparation of all of the solutions. Apparatus. A Metrohm model 713 potentiometer (resolution ± 0.1 mV, reproducibility ± 0.15 mV) connected to a Metrohm 665 automatic buret and to a Ross type Orion electrode (model 8101), coupled with a standard calomel electrode, was used to measure the free hydrogen ion concentration. The potentiometric equipment was also connected to a personal computer for automatic data acquisition, by means of a suitable homemade software. The measurement cells were thermostatted at (T = 298.15 ± 0.1) K. Purified N2(g) was bubbled into the solutions to exclude any CO2(g) and O2(g) trace. To check the repeatability of measurements and avoid systematic errors, some measurements were carried out using a different apparatus (Metrohm model

■

RESULTS Distribution Measurements. Theoretical aspects of distribution measurements were already described in other papers,25−27 and some considerations can be made; as an example, if the ligand concentration in the organic phase is very low, we can approximate the value of the activity coefficient in this phase at: γorg ≈ 1, so that: γaq =

KD KD0

(1)

Along the manuscript, the lowerscript (aq) and (org) refer to the aqueous and organic phase, respectively. 1852



Article

Table 2. Experimental Values of log KD in NaCl(aq) at Different Salt Concentrations in the Molal Concentration Scale at T = 298.15 K I

I

mol·kg−1 0.099 0.504 0.504 1.024 1.669 1.939 2.749 3.147 a

mol·kg−1

log KD L1 0.334 0.352 0.360 0.364 0.417 0.426 0.443 0.440

± ± ± ± ± ± ± ±

0.008a 0.007 0.007 0.006 0.005 0.006 0.008 0.010

0.389 0.506 0.987 1.442 2.007 3.041

I L2 0.194 0.208 0.270 0.286 0.316 0.412

± ± ± ± ± ±

I

mol·kg−1

log KD 0.008 0.008 0.006 0.005 0.010 0.005

L3 −0.255 −0.121 −0.016 0.106 0.203 0.253

0.101 0.506 1.024 1.551 2.635 3.203

mol·kg−1

log KD ± ± ± ± ± ±

0.006 0.006 0.008 0.008 0.005 0.006

0.116 0.188 0.465 1.159 1.775 2.684 3.198

I mol·kg−1

log KD L4 −0.448 −0.446 −0.387 −0.376 −0.315 −0.281 −0.243

± ± ± ± ± ± ±

0.010 0.010 0.009 0.009 0.007 0.009 0.011

log KD L5 0.224 0.254 0.344 0.417 0.593

0.100 0.506 1.023 1.996 3.050

± ± ± ± ±

0.008 0.007 0.006 0.010 0.015

95 % C.I.

The distribution ratio of the components between the organic phase and the saline solution (KD) or the pure water (KD0) allows to calculate the activity coefficient of i-th component (γ). The activity coefficient of a neutral species in pure water or saline solution can be calculated by the Setschenow equation:28 log γ = k mmsalt

(2)

where km and msalt are the Sestschenow coefficient and the sodium chloride concentration (mol·kg−1), respectively. Rearranging eqs 1 and 2 we obtain: log KD = log KD0 + k mmsalt

(3)

K0D)

Once the distribution ratios (KD and were experimentally determined, it is possible to calculate, using the eq 3, the Setschenow coefficient that highlight the salt effect on the distribution of the ligand, and by eq 2, the activity coefficients of the neutral species in the ionic strength range investigated. Table 2 reports the experimental log KD values, in the molal concentration scale, while the corresponding values at infinite dilution and the Setschenow coefficients (km) are reported in Table 3. From the data reported in Table 3, it is possible to

Figure 2. Dependence of log KD on salt concentration (mol·kg−1) for the ligands □, L1; ○, L3; △, L4.

Table 4. Activity Coefficients of the Neutral Species in NaCl(aq) Aqueous Solution at T = 298.15 K

Table 3. log K0D Values at Infinite Dilution and Setschenow Coefficients (eq 2) at T = 298.15 K

a

ligand

log K0D

km

L1 L2 L3 L4 L5

0.341 ± 0.008a 0.170 ± 0.010 −0.197 ± 0.018 −0.443 ± 0.006 0.199 ± 0.008

0.032 ± 0.004a 0.080 ± 0.006 0.151 ± 0.012 0.062 ± 0.003 0.121 ± 0.006 kmb = 0.09 ± 0.05c

log γa

I

a

mol·kg−1

L1

L2

L3

L4

L5

0.100 0.500 1.00 2.00 3.00 5.00

0.003 0.016 0.032 0.064 0.096 0.160

0.008 0.040 0.080 0.160 0.240 0.400

0.015 0.076 0.151 0.302 0.453 0.755

0.006 0.031 0.062 0.124 0.186 0.310

0.012 0.060 0.121 0.242 0.363 0.605

Calculated by eq 2

95 % C.I. bMean value. cStandard deviation of the mean values.

observe that the different ligand structures do not affect the km values, so that a mean common value km = 0.09 ± 0.05 was proposed. As an example, Figure 2 reports the dependence of the log KD values of L1, L3, and L4 on the salt concentration (I/ mol·kg−1 [H2O]). Table 4 reports the activity coefficients of the neutral species calculated at different salt concentrations, while Figure 3 shows the trend of the log γi versus I/mol·kg−1 [H2O]. Protonation Constants. The protonation constants of the ligands here studied can be expressed by the general equilibrium,

Figure 3. Activity coefficients of the neutral species vs I (mol·kg−1) calculated by means of distribution measurements □, L1; ○, L2; △, L3; △, L4; ◇, L5. 1853



Article

Table 5. Overall Protonation Constants at Different Ionic Strengths (Molal Concentration Scale) in NaCl(aq) at T = 298.15 K L1 I/mol·kg−1 0.100 0.491 0.906 1.885 2.408 3.266 3.214 3.846

a

log βH2

log KH1 4.638 4.695 4.781 4.953 5.049 5.221 5.316 5.316

± ± ± ± ± ± ± ±

0.008a 0.013 0.008 0.013 0.012 0.017 0.017 0.025 L2

I/mol·kg−1

log KH1

0.029 0.675 1.326 1.548 2.036 2.915

± ± ± ± ± ±

5.218 5.423 5.545 5.621 5.760 5.874

log KH1

0.031 0.094 0.513 1.331 1.657 3.050 3.651

± ± ± ± ± ± ±

5.472 5.467 5.585 5.721 5.810 6.119 6.280

log KH1

0.171 0.506 1.011 1.671 2.580 3.062

± ± ± ± ± ±

5.610 5.677 5.798 5.909 6.080 6.167

I/mol·kg−1

log KH1

0.023 0.657 1.282 1.907 2.204 2.994

± ± ± ± ± ±

5.658 5.960 6.290 6.691 6.871 7.267

0.020 0.032 0.032 0.029 0.031 0.053

a

10.685 11.212 11.647 12.144 12.346 12.880

± ± ± ± ± ±

7.952 8.387 8.610 8.751 9.026 9.252

a

8.226 8.265 8.486 8.906 9.058 9.649 9.943

± ± ± ± ± ±

0.015 0.011 0.010 0.012 0.020 0.043

± ± ± ± ± ± ±

0.023 0.020 0.022 0.027 0.025 0.015 0.018

log βH2

0.004 0.004 0.004 0.003 0.001 0.002 L5

a

log βH2 a

0.017 0.045 0.060 0.053 0.045 0.052 0.051 0.074

log βH2

0.007 0.007 0.011 0.012 0.011 0.017 0.025 L4

I/mol·kg−1

± ± ± ± ± ± ± ±

log βH2

0.009 0.012 0.005 0.004 0.011 0.034 L3

I/mol·kg−1

5.208 5.315 5.311 5.823 6.269 6.711 6.692 6.926

7.238 7.280 7.440 7.672 8.046 8.176 log βH3

0.020 0.033 0.032 0.027 0.028 0.048

13.465 14.416 15.105 15.604 15.866 16.538

± ± ± ± ± ±

0.030 0.025 0.018 0.007 0.009 0.034

± ± ± ± ± ±

0.019 0.010 0.006 0.010 0.026 0.037 log βH4 15.480 16.959 17.603 18.536 18.803 19.796

± ± ± ± ± ±

0.017 0.045 0.045 0.034 0.030 0.043

95 % C.I.

L0 + i H+ = LHi i +

(i = 1 to 4)

βi H

diaminophenol moieties behave as two independent ligands) had similar protonation constants, so that a mean value of log KH = 5.2 ± 0.4 can be proposed. The second step is strongly influenced by other factors, as the relative position of the two aminogroups. Dependence on Ionic Strength. The Debye−Hückel,10 the SIT,11,12 and the Pitzer type equations13,14 were used to model the dependence of the protonation constants on ionic strength. The Debye−Hückel type equation used in our calculation is:

(4)

The experimental protonation constants (in the molal concentration scale) determined in NaCl solutions at different ionic strengths and at T = 298.15 K, are reported in Table 5; corresponding protonation constants in the molar concentration scale are reported as Supporting Information (Table 1S). As expected, the basicity of the aromatic amines is lower than the aliphatic ones by some orders of magnitude (aliphatic amines: log KH1 > 9); this is due to the lower availability of the lone pair on the nitrogen of the amino group. Analyzing the protonation constants of all of the ligands here investigated, it was observed that the first protonation step (including the second one for L5, where it is predictable that the two

log βi H = log βi H0 − 0.51z* 1854

I + Ci·I 1 + 1.5 I

(5)



Article

Table 6. Protonation Constants at Infinite Diluition in the Molal Concentration Scale at T = 298.15 K, Calculated Using Different Equations log log log log log log log log log log log log a

a KH0 1 H0 β2 b KH0 1 H0 β2 c KH0 1 βH0 2 a βH0 3 H0 β4 b βH0 3 βH0 4 c βH0 3 H0 β4

4.623 5.049 4.607 4.729 4.621 4.979

L1

L2

± ± ± ± ± ±

5.206 ± 0.010 7.805 ± 0.016 5.245 ± 0.032 7.787 ± 0.021 5.209 ± 0.006 7.803 ± 0.022 12.933 ± 0.090 14.310 ± 0.160 13.083 ± 0.032 14.749 ± 0.018 13.068 ± 0.052 14.717 ± 0.054

0.011d 0.027 0.018 0.053 0.018 0.066

L3 5.467 8.050 5.450 7.995 5.467 8.039

± ± ± ± ± ±

L4 0.007 0.024 0.012 0.028 0.007 0.022

5.569 7.021 5.581 6.785 5.574 6.959

± ± ± ± ± ±

L5 5.615 ± 0.054 10.453 ± 0.078 5.649 ± 0.021 10.554 ± 0.021 5.641 ± 0.036 10.542 ± 0.040

0.006 0.028 0.010 0.042 0.004 0.046

Equation 9. bEquation 10a. cEquation 10b. d95 % C.I.

Table 7. Parameters for the Dependence of the Protonation Constants on Ionic Strength at T = 298.15 K L1 i = 1d C Δε c∞ a c0 Δε∞ Δε0 c(0)b c(1) Δε(0)b Δε(1)

0.208 0.186 0.226 0.131 0.193 0.143 0.163 −0.195 0.165 −0.165

± ± ± ± ± ± ± ± ± ±

0.008c 0.008 0.021 0.065 0.021 0.064 0.047 0.180 0.042 0.160

0.474 0.442 0.759 −0.768 0.674 −0.714 0.028 0.428 0.011 0.368 L4

i=1 C Δε c∞a c0 Δε∞a Δε0 c(0)b c(1) Δε(0)b Δε(1) a

0.194 0.294 0.205 0.249 0.175 0.257 0.228 −0.087 0.223 −0.072

± ± ± ± ± ± ± ± ± ±

L2 i = 2d ± ± ± ± ± ± ± ± ± ±

Mde

0.022 0.018 0.072 0.217 0.067 0.224 0.032 0.120 0.028 0.110

0.031

0.037

0.232 0.251 0.191 0.426 0.165 0.428 0.360 0.300 0.353 0.304

± ± ± ± ± ± ± ± ± ±

L3 i=2

0.024 0.024 0.040 0.068 0.033 0.059 0.048 0.058 0.046 0.062

0.338 0.365 0.399 0.261 0.357 0.273 −0.082 0.049 −0.090 0.023

± ± ± ± ± ± ± ± ± ±

i=1

Md 0.018 0.016 0.053 0.094 0.036 0.068 0.042 0.048 0.038 0.050

0.028

0.020

0.020

0.221 0.244 0.279 0.131 0.239 0.152 0.173 0.238 0.175 0.251

± ± ± ± ± ± ± ± ± ±

0.009 0.014 0.022 0.047 0.019 0.047 0.033 0.058 0.032 0.052

i=2

Mde

± ± ± ± ± ± ± ± ± ±

0.028

0.383 0.417 0.515 0.086 0.449 0.127 0.050 0.125 0.030 0.086

0.010 0.016 0.029 0.099 0.024 0.096 0.024 0.036 0.022 0.032

0.015

0.015

L5 Mde

i=2 0.003 0.020 0.013 0.037 0.006 0.022 0.024 0.090 0.022 0.084

0.053

i=1

0.213 0.321 0.496 −0.454 0.441 −0.426 −0.009 0.266 −0.021 0.229

± ± ± ± ± ± ± ± ± ±

0.003 0.024 0.027 0.079 0.029 0.080 0.018 0.066 0.016 0.066

0.026 0.010

0.015

i=1 0.555 0.594 0.663 0.385 0.594 0.417 0.480 0.377 0.493 0.410

± ± ± ± ± ± ± ± ± ±

0.030 0.026 0.049 0.110 0.045 0.129 0.120 0.120 0.112 0.120

i=2 0.646 0.684 0.827 0.022 0.740 0.264 0.084 0.259 0.042 0.188

± ± ± ± ± ± ± ± ± ±

i=3

0.040 0.042 0.047 0.110 0.054 0.129 0.088 0.089 0.086 0.088

0.695 0.738 0.942 0.081 0.834 0.161 0.406 −0.359 0.429 −0.297

± ± ± ± ± ± ± ± ± ±

Mde

i=4 0.048 0.052 0.051 0.135 0.045 0.110 0.120 0.170 0.120 0.170

0.770 0.826 1.400 −1.037 1.247 −0.929 0.199 0.857 0.142 0.741

± ± ± ± ± ± ± ± ± ±

0.088 0.092 0.067 0.182 0.064 0.165 0.090 0.122 0.090 0.133

0.074 0.022

0.031

Equations 6a or 10a. bEquations 6b or 10b. c95 % C.I. di-th protonation steps. eMean deviation on the fit.

where log βH0 is the overall protonation constant at infinite i dilution for the i-th protonation step, z* = ∑(charge)2reactants − ∑(charge)2products, and Ci represents the empirical parameter for the dependence of the protonation constants on ionic strength of the i-th species. In our previous investigations10,29,30 on other classes of ligands, we reported that the C parameter can be considered both constant and dependent on the ionic strength; in this last case, the C parameter can be expressed by the equations: C = c∞ +

c0 − c∞ 1+I

If the SIT approach is taken into account, the molal activity coefficient γi of an ion, having charge zi at T = 298.15 K, is given by: log γi = − 0.51zi 2

I + 1 + 1.5 I

∑ εi ,kmk k

(7)

where ε is the specific interaction coefficient of the i-th ion and the sum is extended over all of the k ions present in the solution at the molality mk. Taking into account eq 4, the log βiH can be expressed as a function of the activity coefficients of each species involved in the equilibria:

(6a)

or:

log βi H = log βi H0 + i log γ(H+) + log γ(L0) C = c(0) + c(1) ln(1 + I )

− log γ(HiLi +)

(6b) 1855

(8)



Article

Using the SIT formulation (eq 7), the ionic strength expressed in the molal concentration scale and eq 8, we obtain: log βi H = log βi H0 − 0.51z*

I + ΔεiI 1 + 1.5 I

(9)

The Δεi value is the difference between the specific interaction coefficients of the reagents and the products. When a ligand neutral species is present, the activity coefficient can be expressed by eq 2. For example, in the case of amines, we have for the first two overall protonation constants: log K1H: Δε1 = ε(H+,Cl−) + km − ε(LH+,Cl−), while for log βH2: Δε2 = 2ε(H+,Cl−) + km − ε(LH22+,Cl−). Equation 9 can be considered equivalent to eq 5, if the molal concentration scale was used, so that Δε = C. The dependence on ionic strength of Δεi can be studied using a similar approach29,31−34 used in eq 6a: Δε = Δε∞ +

Δε0 − Δε∞ 1+I

(10a)

or eq 6b Δε = Δε(0) + Δε(1) ln(1 + I )

(10b)

Table 6 reports the protonation constant values at infinite dilution, calculated using eqs 5 and 9, 6a and 10a, and 6b and 10b, while Table 7 reports all of the parameters for the dependence on ionic strength. The corresponding protonation constants in the molar concentration scale are reported as the Supporting Information (Table 2S). Independently of the model used to calculate the protonation constants at infinite dilution, quite similar values were obtained, and this allowed us to propose, for each ligand and protonation step, a mean protonation constant value (Table 8). From the analysis of the data reported in Table 8, it

Figure 4. Calculated values of log KH1 and log βH2 (eqs 9 to 10a) vs experimental ones.

compared with the corresponding values of the different unsubstituted aliphatic and n-alkyl substituted diamines30,36 (see Table 9). As it can be seen from the data reported in Table Table 9. Parameters for Dependence of the Protonation Constants on Ionic Strength for Different Class of Ligands

Table 8. Mean Values of the Protonation Constants (in the Molal Concentration Scale) at Infinite Dilution at T = 298.15 K a log KH0 1

L1 L2 L3 L4 L5 a

4.62 5.22 5.46 5.57 5.64

± ± ± ± ±

0.02b 0.01 0.01 0.01 0.04

a log βH0 2

4.92 7.80 8.03 6.92 10.52

± ± ± ± ±

0.07 0.02 0.05 0.05 0.05

a log βH0 3

a log βH0 4

13.03 ± 0.06

14.60 ± 0.08

this work

Calculated using eqs 9 and 10b. b95 % C.I.

may be noted that the Debye−Hückel and the SIT approaches give similar results, so that it is possible to calculate for the parameters of eqs 5, 10a, and 10b a mean common value that has a good predictive power for this ligand class, as already proposed in previous papers for other classes of ligands (e.g., resorcinols,26 aminophenols35). Using the approaches reported in eqs 9, 10a, and 10b, we have for the log K1H: Δε = 0.24 ± 0.05 (eq 9, Δε∞ = 0.27 ± 0.03 and Δε0 = 0.33 ± 0.05 (eq 10a) and Δε(0) = 0.28 ± 0.03 and Δε(1) = 0.17 ± 0.01 (eq 10b). The validity of these calculated parameters can be observed in Figure 4, where the correlation between the experimental and the calculated protonation constants is reported. In each case, using the Δε parameters calculated from eq 9, the linear fit gives a correlation coefficient of R = 0.993. The Δε(i)i and c(i)i values (eqs 5, 6a, and 6b; 9, 10a, and 10b) of the aromatic-amino compounds here investigated can be

n-alkyl substituted aminesc

unsubstituted aliphatic aminesd

parameters

i = 1d

i = 2d

i=1

i=2

i=1

i=2

C c∞ a c0 c(0)b c(1) Δε Δε∞a Δε0 Δε(0)b Δε(1)

0.22 0..23 0.23 0.23 0.01 0.24 0.27 0.33 0.28 0.17

0.35 0.44 −0.22 0.01 0.22 0.38 0.49 −0.10 0.11 0.18

0.34 0.33 0.38 0.32 0.01 0.29 0.38 0.39 0.29 0.08

0.60 0.64 0.38 0.44 0.01 0.55 0.52 0.77 0.64 −0.05

0.25 0.24 0.33 0.29 −0.02 0.22 0.21 0.33 0.26 −0.04

0.40 0.46 0.55 0.49 −0.01 0.43 0.39 0.61 0.51 −0.05

a Equation 6a or 10a. bEquation 6b or 10b. cn-Alkyl substituted amines (ref 30). dUnsubstituted aliphatic amines (ref 36).

9, independently of the structure of ligand, the Δε(i)i and c(i)i values are quite similar for the unsubstituted aliphatic diamines, while significant differences are observed for the n-alkyl substituted diamines. The specific interaction coefficients (SIT model) of the (LHii+,Cl−; i = 1 to 4) species of each ligand were calculated using also the km values determined from the distribution data (Table 2); this allowed us to calculate accurate specific interaction coefficients (see Table 10). For the H + Cl− interaction, the literature values: (ε∞(H+,Cl−) = 0.136, ε0(H+,Cl−) = 0.0848)32 were used. 1856



Article

Table 10. SIT Parameters for the Different Protonated Amine Species at T = 298.15 K ε∞(LH+,Cl−) L1 L2 L3 L4 L5

ε∞(LH22+,Cl−)

ε0(LH22+,Cl−)

0.094 ± 0.024 0.051 ± 0.033 0.048 ± 0.022 0.181 ± 0.006 −0.337 ± 0.045a ε∞(LH33+,Cl−)

a

0.092 ± 0.071 −0.264 ± 0.036 0.083 ± 0.033 −0.111 ± 0.030 −0.212 ± 0.130a ε0(LH33+,Cl−)

−0.251 ± 0.065 −0.005 ± 0.059 −0.026 ± 0.048 −0.107 ± 0.022 −0.347 ± 0.054a ε∞(LH44+,Cl−)

−0.305 ± 0.045

0.212 ± 0.110

−0.582 ± 0.064

L5 a

ε0(LH+,Cl−) a

1.033 ± 0.225a −0.025 ± 0.076 0.192 ± 0.098 0.656 ± 0.080 0.025 ± 0.045 ε0(LH44+,Cl−)

a

1.385 ± 0.165

95 % C.I.

Table 11. Pitzer Parameters for the Different Protonated Species, Taking into Account the C(ϕ) Parameters

a

Mdb

β(0)(LH+,Cl−)

β(1)(LH+,Cl−)

C(ϕ)(LH+,Cl−)

L1 L2 L3 L4 L5

0.031 0.013 0.014 0.011 0.028 Mdb

0.013 ± 0.062 −0.003 ± 0.138 0.253 ± 0.059 0.079 ± 0.047 −0.349 ± 0.440 β(0) (LH33+,Cl−)

0.458 ± 0.215 0.132 ± 0.418 0.015 ± 0.212 0.135 ± 0.178 0.591 ± 0.757 β(1) (LH33+,Cl−)

0.015 ± 0.021 0.044 ± 0.055 −0.052 ± 0.019 −0.002 ± 0.017 0.066 ± 0.150 C(ϕ)(LH3+,Cl−)

L5

0.028

−0.024 ± 0.700

4.067 ± 2.200

0.052 ± 0.460

a

β(0)(LH22+,Cl−)

β(1)(LH22+,Cl−)

Cϕ(LH22+,Cl−)

−0.671 ± 0.223 −0.001 ± 0.204 0.096 ± 0.126 −0.051 ± 0.084 −0.210 ± 0.490 β(0) (LH44+,Cl−)

5.304 ± 0.822 1.930 ± 0.613 2.108 ± 0.438 3.296 ± 0.289 2.086 ± 1.636 β(1)(LH44+,Cl−)

0.220 ± 0.107 0.094 ± 0.115 0.015 ± 0.059 0.064 ± 0.049 0.070 ± 0.240 C(ϕ)(LH4+,Cl−)

7.739 ± 3.003

−0.250 ± 0.359

0.255 ± 0.900

95 % C.I. bMean deviation on the fit.

Table 12. Pitzer Parameters for the Different Protonated Species, Neglecting the C(ϕ) Parameter

a

Mdb

β(0)(LH+,Cl−)

β(1)(LH+,Cl−)

β(0)(LH22+,Cl−)

β(1)(LH22+,Cl−)

L1 L2 L3 L4 L5

0.033 0.013 0.014 0.011 0.028 Mdb

0.028 ± 0.018 0.102 ± 0.028 0.105 ± 0.018 0.133 ± 0.032 −0.175 ± 0.066 β(0)(LH33+,Cl−)

0.325 ± 0.180 −0.177 ± 0.136 0.183 ± 0.096 1.956 ± 0.244 0.045 ± 0.448 β(1)(LH33+,Cl−)

−0.225 ± 0.223 0.161 ± 0.024 0.075 ± 0.018 0.055 ± 0.040 −0.075 ± 0.070 β(0)(LH44+,Cl−)

3.802 ± 0.616 1.453 ± 0.202 0.152 ± 0.152 2.985 ± 0.258 1.655 ± 0.466 β(1)(LH44+,Cl−)

L5

0.028

0.060 ± 0.104

3.787 ± 0.686

−0.067 ± 0.140

8.712 ± 0.950

a

95 % C.I. bMean deviation on the fit.

two ions of the same sign, Ψ is the triple interaction parameters (+ − +, − + −), and λ is the interaction parameter for the neutral species. Generally, the Θ, Ψ, and C(ϕ) parameters can be neglected if the ionic strength is less than 3 mol·kg−1. For HCl and NaCl interactions, the values (β(0)HCl = 0.1775, C(ϕ)HCl = 0.00080, β(1)HCl = 0.2945, C(ϕ)NaCl = 0.00127, β(1)NaCl = 0.2664)14 were used. In the calculation of the Pitzer parameters, β(0), β(1), and C(ϕ), that account for the interactions of the different protonated species of the ligands with chloride (anion of the supporting electrolyte), two different approaches were used: in the first one the β(0), β(1), and C(ϕ) were calculated (Table 11), while in the second case (Table 12), the C(ϕ) parameter was neglected. Independently of the approach used, the results can be considered similar, as can be seen from the errors associated to the β(0), β(1), and C(ϕ) values. This suggests that, in the present experimental conditions, the term C(ϕ) can be neglected.

As already reported in previous papers,25−27 an approach similar to that used in eq 6a can be also used for the ε parameter; if it is considered as dependent on ionic strength, we have: ε − ε∞ ε = ε∞ + 0 (11) 1+I For all the ligands (but L5), the ε values are very small for both (LH+,Cl−) and (LH22+,Cl−) couples with a mean value of 0.0 ± 0.1. The interaction coefficients of the various protonated species of L5 are significantly negative, indicating fairly strong interaction with the anion of supporting electrolyte. According to the Pitzer equations, 13,14 the activity coefficients of a cation M or an anion X can be expressed by: ln γM,γX = z 2f γ + f (I ; β (0); β (1); C(ϕ); Θ; Ψ)

(12)

and for neutral species:

ln γN = 2λI

■

(13)

PROTONATED AMINE-CHLORIDE COMPLEXES In a long series of investigations37 (and refs therein), we proposed a simple model to study the dependence of the protonation and complex formation constants on ionic strength, valid at relatively low ionic strength values and for low molecular weight ligands; this model is dependent only on the stoichiometry and on the charges involved in the formation reaction:

with ⎤ ⎡ I f γ = −0.391⎢ + 1.667 ln(1 + 1.2 I )⎥ ⎦ ⎣ 1 + 1.2 I

(14)

where I is the ionic strength in the molal concentration scale, β(0), β(1), and C(ϕ) are the interaction parameters between two ions of opposite signs, Θ is the interaction parameter between 1857



Article

Table 13. Formation Constants of Cl−-Protonated Amines Species at Infinite Dilution, at T = 298.15 K log log log log log log log log a

a βCl 11 Cl b K11 βCl 12 KCl 12 βCl 13 KCl 13 βCl 24 KCl 24

L1

L2

L3

L4

3.86 ± 0.22c −0.76 5.17 ± 0.15 0.25

4.88 ± 0.10 −0.34 8.08 ± 0.10 0.28

4.90 ± 0.10 −0.56 8.32 ± 0.10 0.29

4.93 ± 0.15 −0.64 6.77 ± 0.10 −0.15

L5 6.02 0.38 11.28 0.76 14.49 1.46 16.86 2.26

± 0.15 ± 0.15 ± 0.20 ± 0.20

Refers to eq 17. bRefers to eq 16. c95 % C.I.

log β H = log β H0 − z*

I + CI + DI 3/2 (2 + 3 I )

deviation on the fit) are obtained, so that all of the above models are suitable for the present ligand class. This allowed us to calculate a mean common value, with a good predictive power, so that the calculation of the protonation constants of other ligands belonging to this homogeneous series can be performed. The dependence of protonation constants on ionic strength was interpreted also in terms of ion pair formation between the polyammonium cations and the chloride anion. The lower stability of L1−L4/Cl− ion pairs, with respect to that of aliphatic polyammonium cations, can be attributed to the lower availability of the lone pair of the aromatic amine group, while the higher stability of L5 ion pairs may be due to intramolecular hydrogen bonding. By means of the distribution (2-methyl-1-propanol/aqueous solution) measurements carried out at different ionic strength values, the KD (distribution coefficient) values were calculated, as well as the Setschenow coefficients and the activity coefficients of the neutral species.

(15)

where C = 0.1p* + 0.23z* and D = −0.1z* (p* = number of reactants − number of products). This model considers that the deviations of experimentally obtained constants are due to the formation of weak complexes with the cation or the anion of the supporting electrolyte. In the case of the protonation of polyamine (βH = βiH; z* = i − i2; p* = i), the cation LHii+ binds weakly the anion of supporting electrolyte, namely, Cl− in our case. All of the amines considered in this work show significant deviations log βHi(exp) − log βH eq 15, and performing suitable calculations we obtained the complex formation constants reported in Table 13, according to the equilibria LHi + + jCl− = LHiCl(ji − j) + L + i H+ + jCl− = LHiCl(ji − j) +

K ijCl

(16)

βijCl

■

(17)

For these calculations the molar concentration scale was used, as suggested by Johnson and Pytkowicz.38 The formation constant of Cl− protonated diamines are quite low, with mean Cl values: log KCl 11 = −0.57 ± 0.18 (L1 to L4) and log K12 = 0.27 ± 0.02 (L1 to L3). These values are significantly lower than those observed for aliphatic diamines (mean values for seven 36 Cl diamines log KCl 11 = −0.2, K12 = 0.72 at infinite diluition). − On the contrary, the tetramine L5 forms with Cl more stable complexes than some aliphatic tetramines, such as spermine or Cl triethylentetramine (whose mean values are log KCl 11 = −0.3, K12 Cl Cl = 0.6, log K13 = 1.3, K24 ∼ 1.3, at infinite dilution, and T = 298.15 K).36

ASSOCIATED CONTENT

S Supporting Information *

Experimental protonation constants for the ligands at different ionic strengths (Table S1) and protonation constants at infinite dilution in the molar concentration scale (Table S2). This material is available free of charge via the Internet at http:// pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*Tel.: +39-0906765749; e-mail: [email protected].

■

Notes

The authors declare no competing financial interest.

■

CONCLUSIONS The protonation constants of five aromatic-amino compounds 1,2-phenylenediamine (L1), 4-N-(2-hydroxyethyl)2,4-diaminoanisole (L2), 2-(2,4-diaminophenoxyethanol) (L3), 4-5diamino-1-N-(2-hydroxyethyl)pyrazole (L4), 1,3-bis(2,4diaminophenoxy)propane (L5), were determined by ISE-[H+] potentiometry in NaCl aqueous solutions and in a wide ionic strength range from (0 to 3) mol·kg−1 and T = 298.15 K. The dependence of the protonation constants on ionic strength was studied by using different approaches, namely, the Debye− Hückel, the SIT (specific ion interaction theory), and the Pitzer models. This allowed us to calculate the protonation constants at infinite dilution and the corresponding interaction parameters for the dependence on ionic strength. A comparison between the parameters, obtained applying the different approaches, reveals that similar results in terms of statistical analysis (i.e., errors associated to the parameters and mean

REFERENCES

(1) Baan, R.; Straif, K.; Grosse, Y.; Secretan, B.; El Ghissassi, F.; Bouvard, V.; Benbrahim-Tallaa, L.; Cogliano, V. Carcinogenicity of some aromatic amines, organic dyes, and related exposures. Lancet Oncol. 2008, 9 (4), 322−323. (2) Huncharek, M.; Kupelnick, B. Personal use of hair dyes and the risk of bladder cancer: Results of a meta-analysis. Public Health Rep. 2005, 120 (1), 31−38. (3) Kelsh, M. A.; Alexander, D. D.; Kalmes, R. M.; Buffler, P. A. Personal use of hair dyes and risk of bladder cancer: A meta-analysis of epidemiological data. Cancer Causes Control 2008, 19 (6), 549−558. (4) Lee, H. S.; Lin, Y.-W. Permeation Hair dye ingredients, pphenylendiamine and aminophenol isomers, through protective gloves. Ann. Occup. Hyg. 2009, 53 (3), 289−296. (5) Miligi, L.; Costantini, A. S.; Benvenuti, A.; Veraldi, A.; Tumino, R.; Ramazzotti, V.; Vindigni, C.; Amadori, D.; Fontana, A.; Rodella, S.; Stagnaro, E.; Crosignani, P.; Vineis, P. Personal use of hair dyes and

1858



Article

hematolymphopoietic malignancies. Arch. Environ. Occup. Health 2005, 60 (5), 249−256. (6) Rauscher, G. H.; Shore, D.; Sandler, D. P. Hair dye use and risk of adult acure leukemia. Am. J. Epidemiol. 2004, 160 (1), 19−25. (7) Takkouce, B.; Etminan, Y.; Montes-Martinez, A. Personal use of hair dyes and the risk of cancer: A meta-analysis. J. Am. Med. Assoc. 2005, 293 (20), 2516−2525. (8) Zhang, Y.; De Sanjose, S.; Bracci, P. M.; Morton, M. L.; Wang, R.; Brennan, P.; Hartge, P.; Boffetta, P.; Becker, N.; Maynadie, M.; Foretova, L.; Cocco, P.; Staines, A.; Holford, T.; Holly, E. A.; Nieters, A.; Benavente, Y.; Bernstein, L.; Zahm, S. H.; Zheng, T. Personal Use of Hair Dye and the Risk of Certain Subtypes of Non-Hodgkin Lymphoma. Am. J. Epidemiol. 2008, 167 (11), 1321−1331. (9) Bolt, H. M.; Golka, K. The debate on carcinogenicity of permanent hair dyes: New insights. Crit. Rev. Toxicol. 2007, 37 (6), 521−536. (10) Crea, F.; De Robertis, A.; Sammartano, S. Medium and Alkyl Chain Effects on the Protonation of Dicarboxylates in NaCl(aq) and Et4NI(aq) at 25 °C. J. Solution Chem. 2004, 33 (5), 499−528. (11) Brönsted, J. N. Studies on solubility IV. Principle of the specific interaction of ions. J. Am. Chem. Soc. 1922, 44, 887−898. (12) Guggenheim, E. A.; Turgeon, J. C. Specific interaction of ions. Trans. Faraday Soc. 1955, 51, 747−761. (13) Pitzer, K. S. Thermodynamics of electrolytes. I. Theoretical basis and general equations. J. Phys. Chem. 1973, 77, 268−277. (14) Pitzer, K. S. Ion interaction approach: Theory and data correlation. In Activity coefficients in electrolyte solutions; CRC Press: Boca Raton, FL, 1991; pp 75−154. (15) Pettit, L.; Powell, K. J. The IUPAC Stability Constants Database; International Union of Pure and Applied Chemistry: Research Triangle Park, NC, 2004. (16) Martell, A. E.; Motokaitis, R. J.; Smith, R. M. NIST-Database 46; National Institute of Standards of Technology: Gaithersburg, MD, 1997. (17) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum Press: New York, 1997. (18) May, P. M.; Murray, K. Joint Expert Speciation System JESS Primer; Murdoch University: Perth, Western Australia, 2000. (19) De Stefano, C.; Princi, P.; Rigano, C.; Sammartano, S. Computer Analysis of Equilibrium Data in Solution. ESAB2M: An Improved Version of the ESAB Program. Ann. Chim. (Rome) 1987, 77, 643−675. (20) De Stefano, C.; Sammartano, S.; Mineo, P.; Rigano, C. Computer Tools for the Speciation of Natural Fluids. In Marine Chemistry - An Environmental Analytical Chemistry Approach; Gianguzza, A.; Pelizzetti, E.; Sammartano, S., Eds.; Kluwer Academic Publishers: Amsterdam, 1997; pp 71−83. (21) De Robertis, A.; De Stefano, C.; Sammartano, S.; Rigano, C. The Determination of Formation Constants of Weak Complexes by Potentiometric Measurements: Experimental Procedures and Calculation Methods. Talanta 1987, 34, 933−938. (22) Crea, F.; De Stefano, C.; Gianguzza, A.; Piazzese, D.; Sammartano, S. Protonation of carbonate in aqueous tetraalkylammonium salts at 25 °C. Talanta 2006, 68 (4), 1102−1112. (23) Harned, H. S.; Owen, B. B. The Physical Chemistry of Electrolytic Solutions, 3rd ed.; Reinhold Publishing Corp.: New York, 1964. (24) Robinson, R. A.; Stokes, R. H. Electrolytes Solutions, 2nd ed.; Butterworths: London, 1959. (25) Bretti, C.; Cigala, R. M.; Crea, F.; Foti, C.; Sammartano, S. Solubility and activity coefficients of acidic and basic nonelectrolytes in aqueous salt solutions. 3. Solubility and acitivity coefficients of adipic and pimelic acids in NaCl(aq), (CH4)4NCl(aq) and (C2H5)4NI(aq) at different ionic strengths and at t = 25 °C. Fluid Phase Equilib. 2008, 263, 43−54. (26) Bretti, C.; De Stefano, C.; Lando, G.; Sammartano, S. Activity coefficients, acid-base properties and weak Na+ ion pair formation of some resorcinol derivatives. Fluid Phase Equilib. 2010, 292, 71−79. (27) Bretti, C.; De Stefano, C.; Manfredi, G.; Sammartano, S. Solubility, activity coefficients and acid-base properties of three

naphthol derivatives in NaCl(aq) at different ionic strengths and at T = 298.15 K. J. Mol. Liq. 2011, 158 (1), 50−56. (28) Setschenow, J. Z. Ü ber Die Konstitution Der Salzlösunger Anf Grund Irhes Verhaltens Zu Kohlensäure. Z. Phys. Chem. 1889, 4, 117− 125. (29) Bretti, C.; Crea, F.; Giuffrè, O.; Sammartano, S. The effect of different aqueous ionic media on the acid-base properties of some open chain polyamines. J. Solution Chem. 2008, 37, 183−201. (30) Crea, F.; De Stefano, C.; Giuffrè, O.; Sammartano, S. Ionic strength dependence of protonation constants of N-Alkylsubstituted open chain diamines in NaClaq. J. Chem. Eng. Data 2004, 49, 109− 115. (31) Bretti, C.; Foti, C.; Porcino, N.; Sammartano, S. SIT parameters for 1:1 electrolytes and correlation with Pitzer coefficients. J. Solution Chem. 2006, 35 (10), 1401−1415. (32) Bretti, C.; Foti, C.; Sammartano, S. A new approach in the use of SIT in determining the dependence on ionic strength of activity coefficients. Application to some chloride salts of interest in the speciation of natural fluids. Chem. Spec. Bioavail. 2004, 16 (3), 105− 110. (33) Crea, F.; De Stefano, C.; Foti, C.; Sammartano, S. SIT parameter for the dependence of (poly)carboxylate activity coefficients on ionic strength in (C2H5)4NIaq (0 ≤ I ≤ 1.2 mol kg−1) and (CH3)4NClaq (0 ≤ I ≤ 3.9 mol kg−1) in the temperature range 278 ≤ T ≤ 328 K, and correlation with Pitzer parameters. J. Chem. Eng. Data 2007, 52 (6), 2195−2203. (34) Crea, F.; Crea, P.; De Robertis, A.; Sammartano, S. Thermodynamic Study for the Protonation of Branched Poly(EthylenImine) in NaClaq and its Dependence on Ionic Strength. J. Chem. Eng. Data 2007, 52 (1), 279−285. (35) Bretti, C.; De Stefano, C.; Foti, C.; Sammartano, S.; Vianelli, G. Protonation thermodynamics of some aminophenol derivatives in NaCl(aq) (0 < I < 3 mol kg−1) at T = 298.15 K. J. Chem. Thermodyn. 2012, 44 (1), 154−162. (36) De Robertis, A.; Foti, C.; Giuffrè, O.; Sammartano, S. Dependence on ionic strength of polyamines protonation in NaCl aqueous solution. J. Chem. Eng. Data 2001, 46, 1425−1435. (37) Casale, A.; Daniele, P. G.; De Stefano, C.; Sammartano, S. IonicStrength Dependence of Formation Constants - XII. A Model for the Effect of Background on the Protonation Constants of Amines and Amino-acids. Talanta 1989, 36, 903−907. (38) Johnson, K. S.; Pytkowicz, R. M. Ion Association and Activity coefficients in multicomponent solutions. Activity coefficients in electrolyte solutions, Vol. 2; CRC Press: Boca Raton, FL, 1979; pp 1−62.

1859


Protonation Constants, Activity Coefficients, and Chloride Ion Pair

Recommend Documents