Solubility, activity coefficients and acid–base properties of three naphthol derivatives in NaCl(aq) at different ionic strengths and at T = 298.15 K

Journal of Molecular Liquids 158 (2011) 50–56

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Journal of Molecular Liquids j o u r n a l h o m e p a g e : w w w. e l s ev i e r. c o m / l o c a t e / m o l l i q

Solubility, activity coefficients and acid–base properties of three naphthol derivatives in NaCl(aq) at different ionic strengths and at T = 298.15 K Clemente Bretti, Concetta De Stefano, Giuseppe Manfredi, Silvio Sammartano ⁎ Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Università di Messina, viale Ferdinando Stagno d'Alcontres, 31, I-98166 Messina (Vill. S. Agata), Italy

a r t i c l e

i n f o

Article history: Received 8 September 2010 Received in revised form 11 October 2010 Accepted 15 October 2010 Available online 21 October 2010 Keywords: Naphthol derivatives Protonation constants Solubility Activity coefficient Dependence on the ionic strength

a b s t r a c t Total and intrinsic solubilities of three naphthol derivatives (1-Naphthol, 1,5-Dihydroxynaphthalene and 1Amino-2-Naphthol-4-Sulfonic acid) were determined in NaCl(aq) at different salt concentrations (0–3 mol L−1) and in pure water at T = 298.15 K. To characterize the acid–base properties of these naphthols, the protonation constants were determined potentiometrically in the same conditions. The solubility of neutral species in pure water is −1.978, −2.886, −2.594 (expressed in logarithmic scale) for 1-Naphthol, 1,5-Dihydroxynaphthalene and 1-Amino-2-Naphthol-4-Sulfonic acid, respectively. The values of Setschenow coefficients (km) for these three ligands, are 0.209, 0.158 and 0.369. The dependence on the ionic strength of protonation constants KH i from potentiometric measurements was studied using three models, namely Debye–Hückel, SIT (Specific ion Ionic Interaction) and Pitzer equations. Therefore infinite dilution KH i and interaction coefficients (SIT and Pitzer) were calculated. The activity coefficients were obtained from solubility and protonation data, in the molal concentration scale. Activity coefficients in NaCl(aq) shown the same behaviour found for other O-donors ligands, previously studied. © 2010 Elsevier B.V. All rights reserved.

1. Introduction Substituted polycyclic aromatic hydrocarbons (PAHs) are widely used in several industries, and it is well-known that these compounds are pollutants due to their slow degradation and high toxicity (they can cause cancer and other health effects [1–4]). Naphthol and some of its derivatives are used as dye precursors in oxidative hair colorants, but they are also present in different applications in industrial field. These compounds contain protonable or deprotonable groups and generally show low solubility: therefore a study of the acid–base behaviour is very important to model industrial and environmental processes. In this work we considered three naphthalene derivatives (see Fig. 1): 1-Naphthol (L1), 1,5 Dihydroxynaphthalene (L2) and 1-Amino-2-Naphthol-4-Sulfonic acid (L3). 1-Naphthol is a toxic metabolite of the PAH naphthalene; it is a major component of the pesticide napropamide and the main degradation product of the pesticide carbaryl. Recently it has been reported associations between urinary 1-Naphthol levels and several intermediate measures of male reproductive health, namely sperm motility, serum testosterone levels, and sperm DNA damage [5]. 1,5 Dihydroxynaphthalene is used as corrosion inhibitor [6]; it acts as base for epoxy resin to produce carbon films [7]. 1-Amino-2-Naphthol-4-Sulfonic acid was adopted as a reagent for the determination of arsenic and phosphorus

⁎ Corresponding author. Fax: + 39 090 392827. E-mail address: [email protected] (S. Sammartano). 0167-7322/$ – see front matter © 2010 Elsevier B.V. All rights reserved. doi:10.1016/j.molliq.2010.10.008

[8], and also of molybdate-reactive silica [9]; it reacts in presence of aniline to form nanotubes of polyaniline [10]. In the last years we were involved in the study of aqueous solubility (one of the most important physico-chemical factors) of different classes of compounds [11–17]. In this work we have determined total and intrinsic solubilities, and the protonation constants from potentiometric measurements in sodium chloride at different ionic strengths at T =298.15 K, for the three ligands. In Table 1 and in Table 2 literature data of protonation constants [18–20], solubility and Setschenow coefficients [21–26] of some naphthol derivatives are reported.

2. Experimental section 2.1. Chemicals 1-Naphthol, 1,5 Dihydroxynaphthalene and 1-Amino-2-Naphthol4-Sulfonic acid (Aldrich products) were used without further purification and their purity, checked alkalimetrically, was found to be N99%. Sodium chloride solutions were prepared by weighing pure salt (Fluka, p.a.) previously dried in an oven at T = 383.15 K. Sodium hydroxide solutions were prepared from concentrated NaOH (Fluka puriss. Electrochemical grade) and standardized against potassium biphthalate. Hydrochloric acid solutions were prepared from concentrated ampoules (Fluka) and standardized against sodium carbonate. All solutions were preserved from atmospheric CO2 by means of soda lime traps. Grade A glassware and twice-distilled water were employed in the preparation of all the solutions.

C. Bretti et al. / Journal of Molecular Liquids 158 (2011) 50–56

OH

OH

NH2 OH

L2

L1 OH

L3 SO3H

Fig. 1. Molecule structures of 1-Naphthol (L1); 1,5-Dihydroxynaphthalene (L2); 1-Amino2-Naphthol-4-Sulfonic acid (L3).

2.2. Apparatus The free hydrogen ion concentration was measured with a Metrohm model 713 potentiometer (resolution ± 0.1 mV, reproducibility ± 0.15 mV) connected to a Metrohm 665 automatic burette and to a model 8101 Ross type Orion electrode, coupled with a standard calomel electrode. The potentiometer and the burette were connected to a personal computer which, using suitable software, allows automatic data acquisition. The measurement cells were thermostatted at (298.15 ± 0.1) K. Purified N2 was bubbled into the solutions in order to exclude the presence of CO2 and O2. To avoid systematic errors, some measurements were carried out using a different apparatus (Metrohm model 809 titrando) and software (Metrohm TiAMO 1.0) for the automatic data acquisition. 2.3. Procedure Saturated solutions of 1-Naphthol (L1), 1,5 Dihydroxynaphthalene (L2) and 1-Amino-2-Naphthol-4-Sulfonic acid (L3) were prepared in the following way: an amount of ligand in a small excess was added to a solution of sodium chloride at a pre-established concentration of salt (0 to 3 mol L−1). These solutions were stirred at T = 298.15 K for at least 24 h. From preliminary conductivity measurements on saturated solutions we established that a so long time of stirring was unnecessary, Table 1 Some of literature data for protonation constants of naphthol derivatives. Ligand

logK1

Ligand

1-Naphthol

9.20a 9.416c 9.64b

6-Methyl-2-Naphthol

9.70a

7-Methyl-2-Naphthol

9.64a

1-Naphthol-4(CH3) 1-Naphthol-(H) 1-Naphthol-4(Cl) 1-Naphthol-4(Br) 1-Naphthol-3(NO2) 1-Naphthol-4(C6H5CO) 1-Naphthol-4(CN) 1-Naphthol-4(CHO) 1-Naphthol-4(NO2) 1-Naphthol-4(NO) 2-Naphthol a b c d e f g

logK1

logK2

b

9.39 8.80b

1-Cl-2-Naphthol 6-Br-2-Naphthol

7.97 9.23a

8.72b

1-Br-2-Naphthol

7.89a

and after 4–6 h of stirring equilibrium conditions were reached. Saturated solutions were centrifuged and filtered through a cellulose membrane filter (Ø= 0.45 μm). A volume of 25 ml of the filtered solution was titrated with standard NaOH. Separate titrations of HCl at the same ionic strength as the sample under study were carried out to determine the standard electrode potential E0, and the junction potential coefficient ja (Ej = ja [H+]). To avoid systematic errors, independent experiments were performed at least three times. 2.4. Calculation All calculations relative to the refinement of parameters (protonation constants, analytical concentration of reagents, formal electrode potential) of ligands investigated in this paper were performed by the computer program BSTAC; this program was also used to check the ligands purity. The general least squares computer program LIANA was used for the refinement of both the parameters for the dependence of protonation constants on ionic strength and on ligand concentrations and of the solubility. Computer programs used in this laboratory are described in reference [27]. 3. Results 3.1. Solubility measurements The treatment of solubility data was already described in previous works [11–17], and here we remind just the essential aspects. Owing to proton dissociation/association equilibria, the total solubility of the ligands is due to the sum of all the species in solution, namely, the neutral and deprotonated or partially protonated species, in fact, we have: h i T 0 − S = HL + ½L 

ð1aÞ

h i h i h i T þ 0  2 S ¼ H3 L þ H2 L þ½HL þ L

ð1bÞ

for L1, L2 and L3, respectively, where ST is the total solubility, [HnL0] = S0 (solubility of neutral species, or specific solubility). By rearranging the Eqs. (1)–(1b) and considering the protonation constants (KH i ) (see hereafter), we have:

2-Cl-1-Naphthol

7.76a

7.33b

4-Cl-1-Naphthol

8.75a

7.08b 6.53b 5.73b 8.18b

1-Amino-2-Naphthol-4-Sulfonic acidc) 2,3-Dihydroxynaphthalene c) 1,2-Dihydroxynaphthalene-4-sulfonic acidc) 2,3-Dihydroxynaphthalene-6-Sulfonic acidc

8.80d

3.10d

10.90e

8.34e

12.64f

8.12f

12.16 g

8.19g

9.45a

Rosenberg et al. [18] , values at I = 0 mol L− 1 at T = 298.15 K. Creamer et al. [19], values at I = 0 mol L− 1 at T = 298.15 K. Pettit and Powell [20]. Values of logK2 and logK3 at I = 0.1 mol L− 1, KNO3 at T = 293.15 K. Values at I = 0.2 mol L− 1, NaCLO4 at T = 303.15 K. Values at I = 0.1 mol L− 1, NaCLO4 at T = 298.15 K. Values at I = 0.1 mol L− 1, KCl at T = 293.15 K.

!

T

0

1+

1 K1H ½Hþ 

T

0

1+

1 1 + H H þ2 K2H ½Hþ  K1 K2 ½H 

T

0

1+

h i 1 1 H þ + H H þ 2 + K3 H H þ K2 ½H  K1 K2 ½H 

S =S 7.86b

ð1Þ

h i h i T 0 − 2− S = H2 L + ½HL  + L

S =S

a

51

S =S

ð2Þ ! ð2aÞ ! ð2bÞ

and then the values of S0 can be calculated. By using Eqs. (2) and (2a) for L1 and L2, we found coincidence between the total solubilities (ST) and the solubilities of the neutral species (S0), whose values are reported in Table 3. For L3 different values of the total solubility and the solubility of the neutral species were obtained, as reported in Table 4. Solubility data were fitted to the function: T

T

logS = logS0 + a⋅cMX

ð3Þ

where cMX is the salt concentration expressed in the molar concentration scale [Eq. (3) can be applied to the molal concentration scale by

52

C. Bretti et al. / Journal of Molecular Liquids 158 (2011) 50–56

Table 2 Literature data for the solubility of naphthol derivatives. Ligand

log ST0

km

1-Naphthol

−2.027b;−2.22f

0.197(NaCl),0.472(Na2SO4), 0.081(NaClO4),−0.032(KSCN)

2-Naphthol 1,5-Dihydroxynaphthalene

−2.28f −2.836a; –2.987e

2,3-Dihydroxynaphthalene

−2.63d

0.144(NaCl),0.393(Na2SO4),0.080(NaClO4),−0.046(KSCN) 0.211(NaCl),0.155(NaBr) −0.498(Me4NBr) b No data

a b c d e f

a

a

0.209(NaCl) b;0.237,0.209(NaCl) −0.360 Me4NClc; 0.207(NaCl) f 0.186(NaCl) c; 0.220(NaCl) f

Bhattacharyya et al. [21]. Perez-Tejeda et al. [22]. Almeida et al. [23]. Yalkowsky, et al. [24]. Korenman et al. [25]. Korenman et al. [26].

substituting cMX with mMX], and ST and ST0 are the total solubility at different salt concentrations, and the total solubility in pure water, respectively; a represent the linear parameter for the dependence on ionic strength. The solubility follows the trend: L2b L3b L1, as shown in Fig. 2. The activity coefficients of neutral species, according to Long and McDevit [28], can be calculated by using the equations: log y = log

S00 = kc cMX S0

ð4Þ

log γ = log

S00 = km mMX S0

ð5Þ

in the molar (y) and molal (γ) concentration scale, respectively. S00 is the solubility of the neutral species in pure water, S0 is the solubility of the neutral species at different concentrations of ionic medium and km is the Setschenow coefficient [29]. Eqs. (4) and (5) are valid under the

assumption that in pure water the activity coefficients are equal to the unity (i.e., y0 = γ0 = 1): this assumption is valid when the solubility is low (S b 0.05 mol L−1) and in the absence of interaction of ligand with itself, (self-interactions phenomena) as in our case. By fitting log S0 to the Eqs. (4) and (5) for each ligands we obtained log S00 values and the relative Setschenow coefficients km (molal concentration scale or kc in molar concentration scale) and therefore we calculated the activity coefficients for the neutral species, reported in Table 5. In Table 6 we reported the values of the Setschenow coefficients together with the log S00 values, in the molar and molal concentration scale; while the activity coefficients dependence on salt concentration for L1, L2 and L3 is shown in Fig. 3. The obtained trend is the same for all the naphthol derivatives and, moreover, it is very similar to that found for other O-donor ligands previously studied by our research group [11–17]. 3.2. Protonation constants The experimental values of the protonation constants of L1, L2 and L3, calculated in the same conditions of solubility measurements, are reported in Table 7, (while in the supplementary data the smoothed values are reported).

Table 3 Total solubilities (or solubilities of neutral species) of L1 and L2 at different ionic strength in NaCl(aq) at T = 298.15 K. Ligand

cNaClb

log STc = log S0c

mNaClc

log STm = log S0m

L1

0.000 0.005 0.244 0.932 0.932 1.805 1.805 2.167 2.933 2.933 0.000 0.461 0.484 0.484 0.925 0.925 0.925 1.396 1.489 1.925 1.925 1.999 1.999 3.013 3.013

−1.98 −1.98 −2.04 −2.19 −2.19 −2.39 −2.39 −2.49 −2.66 −2.65 −2.90 −2.96 −2.97 −2.97 −3.06 −3.02 −3.02 −3.15 −3.10 −3.28 −3.29 −3.25 −3.23 −3.42 −3.40

0.000 0.005 0.246 0.951 0.951 1.875 1.875 2.269 3.122 3.122 0.000 0.466 0.490 0.490 0.944 0.944 0.944 1.438 1.536 2.004 2.004 2.085 2.085 3.213 3.213

−1.97 ± 0.01a −1.98 ± 0.01 −2.03 ± 0.01 −2.18 ± 0.01 −2.18 ± 0.01 −2.38 ± 0.01 −2.37 ± 0.01 −2.47 ± 0.01 −2.63 ± 0.02 −2.63 ± 0.02 −2.90 ± 0.01 a −2.96 ± 0.01 −2.97 ± 0.01 −2.96 ± 0.01 −3.06 ± 0.01 −3.01 ± 0.01 −3.01 ± 0.01 −3.14 ± 0.02 −3.09 ± 0.03 −3.26 ± 0.03 −3.28 ± 0.03 −3.24 ± 0.03 −3.21 ± 0.03 −3.39 ± 0.04 −3.37 ± 0.04

L2

STc and STm are the total solubilities in molar and molal concentration scale, respectively. a 95% C.I. b Molar concentration scale. c Molal concentration scale.

Table 4 Total solubilities and solubilities of neutral species of L3 at different ionic strengths in NaCl(aq) in molar and molal concentration scales at T = 298.15 K. cNaClb 0.004 0.004 0.004 0.004 0.164 0.164 0.514 0.514 0.514 0.514 0.928 0.928 0.928 0.928 1.87 1.87 1.87 1.87 2.558 2.558 2.558 3.432 3.432 a b c

mNaClc 0.004 0.004 0.004 0.004 0.165 0.165 0.520 0.520 0.520 0.520 0.947 0.947 0.947 0.947 1.945 1.945 1.945 1.945 2.700 2.700 2.700 3.700 3.700

log STb c −2.53 −2.60 −2.55 −2.53 −2.81 −2.82 −2.65 −2.64 −2.69 −2.68 −2.75 −2.76 −2.78 −2.65 −2.80 −2.80 −2.80 −2.80 −2.90 −2.83 −2.83 −2.89 −3.00

95% C.I. Molar concentration scale. Molal concentration scale.

log STmc −2.53 −2.60 −2.55 −2.53 −2.81 −2.82 −2.64 −2.64 −2.68 −2.67 −2.74 −2.75 −2.77 −2.64 −2.79 −2.79 −2.79 −2.79 −2.88 −2.81 −2.81 −2.86 −2.96

log S0c −2.61 ± 0.05 −2.68 ± 0.05 −2.62 ± 0.05 −2.61 ± 0.05 −2.86 ± 0.05 −2.87 ± 0.05 −2.71 ± 0.04 −2.71 ± 0.04 −2.75 ± 0.04 −2.74 ± 0.04 −2.95 ± 0.02 −2.96 ± 0.02 −2.97 ± 0.02 −2.95 ± 0.02 −2.96 ± 0.06 −2.96 ± 0.06 −2.96 ± 0.06 −2.96 ± 0.06 −3.66 ± 0.08 −3.60 ± 0.08 −3.60 ± 0.08 −3.65 ± 0.10 −3.70 ± 0.10

log S0m a

−2.61 ± 0.05 −2.68 ± 0.05 −2.62 ± 0.05 −2.61 ± 0.05 −2.86 ± 0.05 −2.87 ± 0.05 −2.71 ± 0.04 −2.70 ± 0.04 −2.75 ± 0.04 −2.74 ± 0.04 −2.94 ± 0.04 −2.95 ± 0.04 −2.97 ± 0.04 −2.84 ± 0.04 −2.95 ± 0.06 −2.95 ± 0.06 −2.95 ± 0.06 −2.95 ± 0.06 −3.66 ± 0.08 −3.60 ± 0.08 −3.59 ± 0.08 −3.60 ± 0.10 −3.65 ± 0.10

C. Bretti et al. / Journal of Molecular Liquids 158 (2011) 50–56

53

Table 6 Total solubilities and solubilities of neutral species for L1, L2 and L3, at T = 298.15 K.

-2.0

-2.5

Ligand

log ST0 = Log S00

km

kc

L1 L2 L3

−1.978 ± 0.001a −2.886 ± 0.003 log S00 −2.684 ± 0.009 log ST0 −2.613 ± 0.002

0.209 ± 0.001a 0.158 ± 0.003

0.230 ± 0.001a 0.177 ± 0.003

0.257 ± 0.04 am 0.089 ± 0.002

0.272 ± 0.04 ac 0.103 ± 0.002

L3 a

95% C.I.

-3.0

0

1

2

where c∞ and c0 are the parameters for the dependence of protonation constants on ionic strength valid for I → ∞ and I → 0, respectively, this formulation of C is already proposed in previous works [16,35] and it can be used up to I = 6 mol kg−1. In the Specific Ionic Interaction equation [36] the activity coefficients of a cation or an anion can be expressed by:

3

Fig. 2. Dependence of the total solubility on ionic strength in NaCl, for 1-Naphthol (●), 1,5-Dihydroxynaphthalene (▲) and 1-Amino-2-Naphthol-4-Sulfonic acid (■).

log γ = −z The protonation constants of the three ligands can be expressed as a function of the activity coefficients, and just for instance we report the equilibrium for L2 as follows: H K1

H0 K1

2

0:51I1 = 2 + ∑ε mi 1 + 1:5I1 = 2

ð10Þ

and for the neutral species: log γ = km I

ð11Þ

+ log γHþ + log γL2− − log γHL−

ð6Þ

where ε is the interaction coefficient expressed by:

log K2 = log K2 + log γHþ + log γL− − log γH2 L0

ð7Þ

ε = ε∞ +

= log

H0

H

0

where log KH i is the protonation constant at infinite dilution and γi is the activity coefficient of the ith component. We used three models to study the dependence of the protonation constants on ionic strength, namely the Debye Hückel type [30], SIT [31,32] and Pitzer [33,34] models. The Debye Hückel type equation is: H

H0

log Ki = log Ki −0:51⋅z*

I1 = 2 + C ðIÞ 1 + 1:5I 1 = 2

ð8Þ

where z* = ∑ z2react − ∑ z2prod and C(I) is a function of ionic strength that can be expressed as C = c∞ +

c0 −c∞ I+1

ð9Þ

Table 5 Activity coefficients for the neutral species for each ligands in NaCl(aq) studied, at T = 298.15 K. Ligand

Ia

log γ

L1

0.1 0.5 0.75 1 2 3 0.1 0.5 0.75 1 2 3 0.1 0.5 0.75 1 2 3

0.021 0.105 0.157 0.209 0.418 0.627 0.016 0.079 0.119 0.158 0.316 0.474 0.037 0.185 0.277 0.369 0.738 1.107

L2

L3

a

Molal concentration scale.

ε0 −ε∞ I+1

ð12Þ

By using the SIT formulation, the molal protonation constants can be expressed as a function of molal ionic strength by Eq. (8), with C(I) =ΔεiI. + − For example, in the case of L1, for log KH 1 we have: Δε1 =ε(H , Cl ) + ε(Na+, L−) −km. Note that, in this simple formulation, SIT approach is identical to the one used in Eq. (8), for the molal concentration scale. For this calculation we used, for HCl, the following values ε∞ = 0.136 and ε0 = 0.0839 reported in reference [37]. According to the Pitzer equations [33,34], the activity coefficients of cation M or anion X can be expressed by:   2 γ ð0Þ ð1Þ ðϕÞ ln γΜ ; ln γX = z f + f I; β ; β ; CH;Cl ; Θ; Ψ

ð13Þ

and for neutral species: ð13aÞ

ln γMX0 = 2λI

1,0

γ

log

0,5

0,0 0

1

2

3

Fig. 3. Dependence of log γ on ionic strength in NaCl(aq) for L1 (■), L2 (●) and L3 (▲).

54

C. Bretti et al. / Journal of Molecular Liquids 158 (2011) 50–56

Table 7 Experimental protonation constants of L1, L2 and L3 in NaCl at a different ionic strength at T = 298.15 K. Ligand

Ib

log KH 1

log KH 2

log KH 3

Ic

log KH 1

log KH 2

log KH 3

L1

0.244 0.932 1.805 2.167 2.933 0.461 0.484 0.925 1.396 1.489 1.925 1.999 3.013 0.164 0.514 0.928 1.870 2.558 3.432

9.17 ± 0.01a 9.15 ± 0.01 9.28 ± 0.01 9.36 ± 0.01 9.50 ± 0.01 10.24 ± 0.01 10.23 ± 0.01 10.07 ± 0.01 10.07 ± 0.01 10.05 ± 0.01 10.08 ± 0.01 10.09 ± 0.01 10.30 ± 0.02 9.86 ± 0.02 9.85 ± 0.02 9.91 ± 0.02 10.37 ± 0.02 10.86 ± 0.03 11.40 ± 0.03

– – – – – 8.76 ± 0.01 8.76 ± 0.01 8.72 ± 0.01 8.76 ± 0.01 8.73 ± 0.01 8.82 ± 0.01 8.82 ± 0.01 8.98 ± 0.02 8.19 ± 0.02 8.21 ± 0.02 8.23 ± 0.03 8.53 ± 0.04 8.87 ± 0.03 9.19 ± 0.04

– – – – – – – – – – – – – 1.34 ± 0.02 1.50 ± 0.04 1.80 ± 0.04 1.98 ± 0.05 2.47 ± 0.08 3.47 ± 0.15

0.246 0.951 1.875 2.268 3.122 0.466 0.490 0.944 1.438 1.536 2.004 2.085 3.213 0.165 0.520 0.947 1.945 2.700 3.697

9.17 ± 0.01 9.14 ± 0.02 9.26 ± 0.02 9.34 ± 0.02 9.48 ± 0.02 10.23 ± 0.01 10.22 ± 0.01 10.07 ± 0.01 10.05 ± 0.01 10.03 ± 0.01 10.06 ± 0.01 10.07 ± 0.01 10.27 ± 0.02 9.86 ± 0.02 9.84 ± 0.02 9.89 ± 0.03 10.35 ± 0.03 10.83 ± 0.03 11.36 ± 0.04

– – – – – 8.75 ± 0.02 8.75 ± 0.02 8.70 ± 0.02 8.75 ± 0.02 8.72 ± 0.02 8.80 ± 0.02 8.80 ± 0.02 8.95 ± 0.03 8.19 ± 0.01 8.20 ± 0.01 8.22 ± 0.02 8.51 ± 0.03 8.85 ± 0.03 9.16 ± 0.04

– – – – – – – – – – – – – 1.35 ± 0.12 1.49 ± 0.04 1.79 ± 0.03 1.96 ± 0.05 2.45 ± 0.07 3.47 ± 0.14

L2

L3

a b c

95% C.I. Molar concentration scale. Molal concentration scale.

with:      γ 1=2 1 = 2 −1 1=2 ð13bÞ 1 + 1:2I + 1:667 ln 1 + 1:2I f = −0:392 I where I is the ionic strength in molal scale, β(0), β(1), and C(ϕ) represent the interaction parameters between two ions of opposite charge, the Θ interaction parameter between two ions of the same charge, Ψ triple interaction parameter (+ − +, – + –), λ interaction parameter for neutral species. As reported in other studies, this group of research used a simplified version of Pitzer equation depending on three empirical parameters only: H

T

γ

H

ð1Þ

2

ln K = ln K + 2z*f + 2P1 I + P2 I + P3 f1 + 2zβMX f2

ð14Þ

with:

parameters P1, P2 and P3 as follows: for the equilibrium: L 2− + H+ = HL− ð0Þ

ð0Þ

ðϕÞ

ðϕ Þ

ð1Þ

ð1Þ

ð0Þ

P1 = βHCl + βNaL −βNaHL + θHNa P2 = CHCl + CNaL =

pffiffiffi ðϕÞ ðϕÞ 2−CNaHL + CNaCl + ΨHNaCl ð1Þ

P3 = βNaL + βHCl −βNaHL for the equilibrium: HL− + H+ = H2L0, ð1Þ

ð1Þ

ðϕÞ

ðϕ Þ

P1 = βHCl + βNaHL −λ + θHNa ðϕ Þ

P2 = CHCl + CNaHL + CNaCl + ΨHNaCl ð1Þ

ð1Þ

P3 = βHCl −βNaHL

    1=2 1=2 exp −2I f1 = 1– 1 + 2I

ð14aÞ



    1=2 1=2 f2 = −1 + 1 + 2I + 2I exp −2I

ð14bÞ

where z* is defined in Eq. (7) and z is the charge of the species (L2−, HL−, H2L0 depending on the protonation step considered); P1, P2 and P3 are empirical parameters and f1 and f2 are functions defined in the Pitzer equation. By considering, for example, the protonation equilibria of 1Amino-2-Naphthol-4-Sulfonic acid we can express the empirical

for the equilibrium: H2L0 + H+ = H3L+, ð0Þ

ð0Þ 3 LCl

ðϕÞ

ðϕ Þ

ð1Þ

ð1Þ 3 LCl

P1 = βHCl −βH

+ λ + θHNa ðϕ Þ 3 LCl

P2 = CHCl + CNaCl −CH

+ ΨHNaCl

P3 = βHCl −βH

The same considerations can be made for the other two ligands.

Table 8 Protonation constants at infinite dilution and parameters for the dependence of the protonation constants on ionic strength for L1, L2 and L3 at T = 298.15 K. Ligand L1 L2 L3

a b c

95% C.I. See Eq. (8). See Eq. (11).

i 1 1 2 1 2 3

log KH i

0

9.45 ± 0.02 11.17 ± 0.03 9.15 ± 0.03 10.35 ± 0.04 8.44 ± 0.04 1.29 ± 0.05

a

c0b

c∞b

Δε0c

Δε∞c

0.26 ± 0.01 0.36 ± 0.01 0.23 ± 0.01 0.75 ± 0.02 0.48 ± 0.02 0.45 ± 0.02

−0.02 ± 0.03 −0.93 ± 0.05 −0.26 ± 0.05 0.08 ± 0.08 −0.02 ± 0.08 –

0.22 ± 0.01 0.32 ± 0.01 0.20 ± 0.01 0.66 ± 0.02 0.42 ± 0.02 0.42 ± 0.02

−0.006 ± 0.04 −0.904 ± 0.05 −0.245 ± 0.05 0.135 ± 0.08 0.013 ± 0.07 –

C. Bretti et al. / Journal of Molecular Liquids 158 (2011) 50–56 Table 9 SIT parameters for (Na+, L2−), (Na+, HL−), (H3L+, Cl−) interactions, at T = 298.15 K. Ligand

ε∞

ε0

L1 L2

(L−, Na+) 0.12 ± 0.08a (L2−, Na+) −1.16 ± 0.09 (HL−, Na+) −0.17 ± 0.09 (L2−, Na+) 0.35 ± 0.10 (HL−, Na+) 0.30 ± 0.07 (H3L+, Cl−) 0.03 ± 0.04

(L−, Na+) 0.30 ± 0.03a (L2−, Na+) 0.41 ± 0.03 (HL−, Na+) 0.23 ± 0.03 (L2−, Na+) 1.18 ± 0.06 (HL−, Na+) 0.65 ± 0.04 (H3L+, Cl−) 0.09 ± 0.04

L3

a

11

10

95% C.I.

SIT calculations were performed by using Eq. (9) and, for the neutral species, we used the values of Setschenow coefficients calculated with Eqs. (4) and (5); Debye Hückel and SIT parameters are reported in Tables 8 and 9 respectively. About the simplified Pitzer parameters in our case we fitted protonation constants (in molal concentration scale) vs. ionic strength just considering two parameters β(0) and β(1) (Cϕ is generally negligible at I ≤ 3 mol L−1): the relatives values are reported in Table 10 together with the protonation constants at infinite dilution for each ligands. For the calculation carried out with the Pitzer equations, we used for NaCl and HCl the literature interaction parameters NaCl: β(0) =0.00127, β(1) =0.2664; HCl: β(0) =0.1775, C(ϕ) =0.00080, β(1) =0.2945; ΘH,Na = 0.036; ΨH,Na,Cl =−0.004) [34]; while the λ parameter was calculated from Setschenow coefficients by using: λ=

ln 10km 2

ð15Þ

In Fig. 4, as an example, we reported log KH 1 vs. ionic strength. The dependence on (I/mol kg−1)1/2 is, for L1 and L2, very similar to that shown by carboxylates and resorcinols [17,30], whilst for L3 the presence of sulfonic and amino groups leads to a quite different behaviour. In fact we must remember that the function log KH vs. I is affected by the interaction of the considered ligands with the supporting electrolyte, and this interaction has opposite effect on −COO− (or −O−) and NH+ 3 groups. 4. Conclusions Toxicity data on the naphthol derivatives have been reported in literature, but very few data are present about their solution thermodynamic properties, that can help to understand their chemical behaviour. For this reason in this paper we studied three naphthol derivatives and, in particular, we determined the protonation constants, the solubility and activity coefficients in NaCl aqueous solution, at different salt concentrations. These parameters, altogether, allow to model rigorously the acid– base properties of the substituted naphthols here studied, and to predict some ligand class parameters. For example, by considering km values here reported, we have a mean km value of 0.25 ± 0.10 that can be used for the naphthol derivatives class (in accordance with km = 0.20 obtained as the average of the literature data reported in

Table 10 Protonation constants at infinite dilution calculated by Pitzer model at T = 298.15 K. 0

Ligand

i

log KH i

L1 L2

1 1 2 1 2 3

9.45 ± 0.02a 11.18 ± 0.03 9.15 ± 0.04 10.35 ± 0.09 8.45 ± 0.05 1.38 ± 0.40

L3

a

55

95% C.I.

β(0)

β(1)

0.35 ± 0.01a 0.52 ± 0.01 0.27 ± 0.01 1.41 ± 0.06 0.76 ± 0.04 0.11 ± 0.25

−0.06 ± 0.10a −1.73 ± 0.15 −0.47 ± 0.12 −0.25 ± 0.24 −0.41 ± 0.18 0.84 ± 0.30

9 0,5

1,0

1,5

Fig. 4. Dependence of 1-Naphthol (■),1,5-Dihydroxynaphthalene (●) and 1-Amino-2Naphthol-4-Sulfonic acid (▲) on ionic strength, in NaCl and T = 298.15 K.

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