Quantitative study on the interaction of Sn2+ and Zn2+ with some phosphate ligands, in aqueous solution at different ionic strengths

Journal of Molecular Liquids 165 (2012) 143–153

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Quantitative study on the interaction of Sn 2+ and Zn 2+ with some phosphate ligands, in aqueous solution at different ionic strengths Rosalia Maria Cigala, Francesco Crea, Concetta De Stefano, Gabriele Lando, 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 28 June 2011 Received in revised form 24 October 2011 Accepted 9 November 2011 Available online 19 November 2011 Keywords: Tin(II) Phosphates Formation constants Ionic strength Speciation Sequestration

a b s t r a c t Tin(II) interaction with different phosphate ligands, namely phosphate (PO4), pyrophosphate (PP), tripolyphosphate (TPP), monofluorophosphate (MFP) and adenosine-5′-triphosphate (ATP), was studied at T=298.15 K by potentiometry and voltammetry at different ionic strengths (0.15≤I/mol L−1 ≤1.00) in NaNO3. We also compared our results with those experimentally determined for the Zn/PO4 and Zn/TPP systems. As concerns the Zn/PP, the Zn/ATP and the Zn/MFP systems, we performed a critical literature analysis. In all cases the stability constants observed for the Sn/L species resulted to be higher with respect to the analogous Zn/L ones. The rough correlation (valid for the ML species) log KML (Sn)=3.01·log KML (Zn)−8.13 was obtained from the stability data of the complexes of these cations. In addition, the stability trend found for a given metal cation was: PP~PO4 >>TPP>>MFP~ATP. The ionic strength dependence of the stability constants was studied by the extended Debye–Hückel and the SIT (Specific ion Interaction Theory) equations. Speciation and sequestration studies were also performed, and pL0.5 values (i.e., the total ligand concentration necessary to bind 50% of cation present in trace) were calculated for all the systems at different pH and ionic strengths. In this case, as an example at pH=7.0 and I=0.15 mol L− 1, the sequestration trend was: PO4 >PP~MFP>>TPP>ATP. The dependence of pL0.5 values on pH and ionic strength was modeled by means of two empirical relationships. © 2011 Elsevier B.V. All rights reserved.

1. Introduction Natural waters and biological fluids are multi-component systems containing a variety of metal or organo-metal cations and ligands both inorganic and organic in nature [1–4]. It is widely accepted that the knowledge and the understanding of the speciation of a chemical element or compound (e.g. metal cation) are of fundamental importance for the determination of the quality of a generic “solution” (i.e. natural and biological fluids). The toxicity and the biogeochemical cycling of a heavy metal in the aquatic environment depend on the physicochemical form in which the metal occurs in the dissolved state [5]. Speciation studies put their attention on this topic. For this reason we considered very important the knowledge of the tin(II) behavior in aqueous solutions, although it is not considered a micronutrient neither a pollutant agent, since its concentration is high in the earth crust. The corrosion of tin plated food cans by acidic foods and beverages has caused several intoxications by the soluble tin compounds. This led, for example, the Food Standard Agency in the UK to ⁎ Corresponding author. Tel.: + 39 090 393659; fax: +39 090 392827. E-mail addresses: [email protected] (R.M. Cigala), [email protected] (F. Crea), [email protected] (C. De Stefano), [email protected] (G. Lando), [email protected] (G. Manfredi), [email protected] (S. Sammartano). 0167-7322/$ – see front matter © 2011 Elsevier B.V. All rights reserved. doi:10.1016/j.molliq.2011.11.002

propose upper limits of 200 mg kg− 1 for solid food and 150 mg kg− 1 for beverages [6]. Among the different radionuclides, tin is of interest because of the presence in radioactive waste of one of its isotopes, 126 Sn, coming from fission; its half-life value is close to 105 years [7]. Tin can also exist as organic compounds (e.g., tributyltin, triphenyltin) that are commonly used in various industrial sectors because of their biocide properties [8]. While the chemistry of organotin compounds has been studied extensively because of their high toxicity toward numerous living organisms, less work has been performed on the inorganic forms that are generally considered to be non hazardous [9]. However, the presence of 126Sn in the radioactive waste requires a better knowledge of the chemistry of inorganic tin to understand and model its behavior after disposal in a deep repository. In addition inorganic tin(II) can easily be organicated by ubiquitous methylating and transmethylating bacteria. Phosphates are very important molecules in nature for different reasons. Phosphate is a macronutrient, for both plants and seawater organisms, its concentration in soils and water is very high and is related to the concentration of phytate (see ref [10] and references therein). As an example, the profile of phosphate concentration in seawater is very different with respect to other components, its concentration in surface is close to zero, as phytoplankton and other organisms decompose, the phosphate is regenerated in the water column, reaching the maximum

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of concentration at 1000 m depth in correspondence of the O2 minimum concentration [11]. Pyrophspate is a biological anion of some importance, since many biosynthetic reactions are made irreversible by the hydrolysis of inorganic pyrophosphate. It is also instrumental in the synthesis of terpenes and steroids. In addition stannous pyrophosphate itself has found application in some therapies for more than 35 years [12–14]. Tripolyphosphate and ATP are very important in the biological field for their presence in human cells and in many biochemical reactions. Finally MFP is widely used in cosmetic field. For these reasons, we undertook a potentiometric and voltammetric study on the interaction of tin(II) with different phosphate ligands, such as phosphate (PO4), pyrophosphate (PP), tripolyphosphate (TPP), monofluorophosphate (MFP) and adenosine-5'-triphosphate (ATP) at T = 298.15 K and different ionic strengths (0.15 ≤ I/mol L − 1 ≤ 1.00) in NaNO3. By considering that an inhibitor interaction of dietary tin on zinc retention has been described from different human metabolic studies [15–17], we performed also an analogous study on the interaction of zinc(II) with phosphate and tripolyphosphate, while for the other ligands we used some critically selected literature data, the values for the Zn/MFP system were guessed by means of a predictive equation. This study follows a previous one, on the inorganic speciation of tin (II), and would clarify, together with future studies on different systems, the chemistry of this metal cation and the different behavior with respect to the other bivalent metal cations. The ionic strength dependence was studied by means of both Debye–Hückel and SIT (Specific ion Interaction Theory) approaches. Speciation and sequestration studies were also carried out. Very important is the use of different experimental techniques, especially when dealing with complex systems that can also test the proposed speciation model with information obtained with different techniques. Collected literature data on tin(II) complexes with phosphates are not very accurate, we only know papers reported by Cilley [18] (for Sn/PO4), Ciavatta and Iuliano [19], Duffield et al. [20], Mesmer and Irani [21] as concerns PP and TPP, while no papers were published, to our knowledge, on the interaction of tin(II) with MFP and ATP. A different situation is present for the zinc(II), being this cation one of the most studied in the past. Rich is the literature especially as concerns the Zn/ATP system. In any case an extensive discussion on literature data is made in a proper section.

2. Experimental section 2.1. Chemicals Tin(II) and zinc(II) solutions were prepared by weighing the pure chloride salts (Fluka); the concentration was checked against EDTA standard solutions [22] and the purity was always ≥ 99.5%. In some cases, i.e. for the voltammetric measurements, the metal solutions were prepared by dilution from a Atomic Absorption Standard Solution (1000 ppm). Particular attention was paid to the preparation of these solutions, in order to prevent oxidation of tin(II) to tin(IV). The solutions were acidified with HCl to have pHb 2 and a piece of metallic tin was added to the solutions after the preparation. The solutions were immediately bubbled with purified N2(g) to exclude any O2(g) trace, and used always immediately after the preparation. Nitric acid, hydrochloric acid and sodium hydroxide solutions were prepared by diluting concentrated ampoules; they were standardized against sodium carbonate and potassium hydrogen phthalate, respectively, and previously dried in an oven at T = 383.15 K for at least 2 h. Moreover, NaOH solutions were preserved from atmospheric CO2 by means of soda lime traps. NaNO3 solutions were prepared by weighing the pure salt dried in an oven at T = 383.15 K for at least 2 h. The ligand solutions were prepared by weighing the sodium salts (Fluka) without further purification. The purity was checked potentiometrically by alkalimetric titrations, resulting always ≥ 99%.

All the solutions were prepared with analytical grade water (R=18 MΩ cm− 1) using grade A glassware. 2.2. Apparatus and procedure for voltammetric measurements The apparatus used in the voltammetric measurements was accurately described elsewhere [23,24]. The voltammetric system was connected, by a remote box purchased by Metrohm (model 6.2148.010), to a Metrohm model 809 Titrando apparatus controlled by Metrohm TiAMO 1.2 software in order to perform automatic voltammetric titrations. The free hydrogen ion concentration was measured before and after each voltammetric run by using a Metrohm glass electrode (model 6.0224.100), calibrated as in the case of potentiometric measurements. The slope of the electrode was randomly checked and was 59.2± 0.1 mV. The time required to reach the equilibrium was 180 to 300 s. The DP-ASV measurements were made using the following procedure. First the solutions were bubbled with purified N2 for 300 s. Then, a plating potential of −1.1 V (−1.8 V for zinc) was applied for 30 s under stirring. After a rest time of 20 s, the voltammogram was recorded from −0.6 V to −0.1 V (−1.5 V to −0.8 V for zinc) with a step potential of 1.5 mV, a scan rate of 3 mV s− 1 and an interval of 0.5 s. The modulation amplitude was 25 mV, with a modulation time of 0.05 s. The above-cited conditions were chosen after preliminary DP-ASV measurements, with and without ligand in solution, in which various parameters were systematically varied to select those giving the best performances in terms of signal/noise ratio, single measurement speed, and repeatability. The effects of the variation of i) the scan rate (0.001 to 0.01 V s− 1), ii) the modulation amplitude (25 to 50 mV), iii) the scan range (−1.8 to 0.0 V), iv) the plating potential (−1.8 to −0.6 V), v) the plating time (10 to 120 s), vi) the concentration of metal (10− 8 to 10− 5 mol L − 1) were investigated. The error of ±1 mV was experimentally determined for the potential reading. For ten replicate experiments the standard error is ±0.3 mV. Measurements were performed on 25 mL of solutions containing known amounts of SnCl2 (or ZnCl2) (cM = 10 − 6 to 10 − 8 mol L − 1), suitable amounts of HNO3 (~20 mmol L− 1) and NaNO3 in order to obtain a pre-established ionic strength value (0.15 ≤ I/mol L − 1 ≤ 1.00). Two different sets of measurements were carried out. The first one, here called ligand titration, consisted of the addition of different amounts of the ligand (up to 15 mmol L − 1) to the solution containing the metal (tin or zinc) at 1.5 ≤ pH≤ 2.0, in order to investigate the formation of stepwise species like MLq, (the formation of hydrolytic species does not occur in this pH range). The second set, here called acid-base titration, was performed as follows: the solution containing both the metal and the ligand, at a fixed cL:cM ratio (10 3:1 ≤ cL:cM ≤ 107:1), was titrated with standard NaOH and, for each point of the titration, both pH value and voltammogram were recorded. The analysis of the acid-base titrations allowed us to determine the formation of the different protonated species MLqHr (or MLq(OH)r). For all the ligands different cL:cM ratios were considered in order to investigate a wide range of experimental conditions. For each titration, 30–40 voltammograms were recorded in the pH range 1.8 ≤ pH ≤ 10.0 (the end titration pH values varied with the different system). In order to check the reproducibility, measurements were carried out by two operators; for each M/L system three cL:cM ratios were investigated and for each one at least two experiments were performed. 2.3. Apparatus and procedure for potentiometric measurements The apparatus used in the potentiometric measurements was widely described elsewhere [23,25]. Before the study of the M/L systems, we determined the protonation constant values of the ligands at different ionic strengths (0.15≤ I/mol L − 1 ≤ 1.00) in NaNO3. 25 mL of a solution containing each ligand (cL = 10 to 30 mmol L − 1), HNO3 and NaNO3, was titrated with standard NaOH solutions.

R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

145

Table 1 Metal hydrolysis and ligand protonation constants used in the calculations at different ionic strengths in NaNO3 and at T = 298.15 K. log β Equilibrium +

−

H2O = H + OH Sn2 + + OH− = Sn(OH)+ Sn2 + + 2 OH− = Sn(OH)2 Sn2 + + 3 OH− = Sn(OH)3− 2 Sn2 + + 2 OH− = Sn2(OH)22+ 3 Sn2 + + 4 OH− = Sn3(OH)42+ Sn2 + + Cl− = SnCl+ Sn2 + + 2 Cl− = SnCl2 Sn2 + + 3 Cl− = SnCl3− Sn2 + + OH− + Cl− = SnOHCl Zn2 + + OH− = Zn(OH)+ Zn2 + + 2 OH− = Zn(OH)2 Zn2 + + 3 OH− = Zn(OH)3− Zn2 + + 4 OH− = Zn(OH)42 − H+ + PO43 − = HPO42 − 2 H+ + PO43 − = H2PO4− 3 H+ + PO43 − = H3PO4 H+ + P2O74 − = HP2O73 − 2 H+ + P2O74 − = H2P2O72 − 3 H+ + P2O74 − = H3P2O7− 5− 4− H+ + P3O10 = HP3O10 5− 3− 2 H+ + P3O10 = H2P3O10 5− 2− = H3P3O10 3 H+ + P3O10 5− − = H4P3O10 4 H+ + P3O10 5− 5 H+ + P3O10 = H5P3O10 H+ + FPO32 − = HFPO3− 2 H+ + FPO32 − = H2FPO3 H+ + ATP4 − = HATP3 − 2 H+ + ATP− = H2ATP2 − a b c d

I = 0.15 mol L− 1

I = 0.50 mol L− 1

I = 1.00 mol L− 1

− 13.75 9.97a 20.97a 24.28a 22.44a 48.61a 0.76a 1.50a 1.46a 11.68a 4.55b 10.40b 12.85b 14.40b 11.49a 18.14a 20.01a 8.42 ± 0.02c 14.39 ± 0.02 16.22 ± 0.03 7.86 ± 0.02c 13.42 ± 0.02 15.23 ± 0.02 15.61d 15.92 4.61 ± 0.01c 5.95 ± 0.02 6.33 ± 0.01c 10.17 ± 0.01

− 13.73 9.82a 20.78a 24.14a 22.26a 48.29a 0.58a 1.23a 1.19a 11.37a 4.63b 10.36b 12.79b 14.52b 11.23a 17.68a 19.51a 7.91 ± 0.02c 13.57 ± 0.02 15.06 ± 0.02 7.37 ± 0.01c 12.51 ± 0.01 14.05 ± 0.01 14.46d 14.84 4.32 ± 0.01c 5.52 ± 0.01 5.99 ± 0.01c 9.99 ± 0.01

− 13.73 9.73a 20.65a 24.04a 22.14a 48.12a 0.49a 1.09a 1.05a 11.12a 4.83b 10.56b 12.89b 14.72b 11.11a 17.46a 19.30a 7.53 ± 0.01c 12.97 ± 0.01 14.19 ± 0.01 7.12 ± 0.02c 12.04 ± 0.02 13.46 ± 0.02 13.99d 14.50 4.09 ± 0.01c 5.17 ± 0.02 5.75 ± 0.01c 9.66 ± 0.01

[unpublished data from this laboratory]. Ref. [29]. ±95% C.I. Ref. [49].

For the Sn/ATP and the Sn/TPP systems, the titrand solutions consisted of different amounts of ligand (1 to 5 mmol L − 1), tin(II) (0.5 to 1.5 mmol L − 1), a suitable amount of HCl and the background salt, NaNO3, in order to obtain pre-established ionic strength values (0.15 ≤ I/mol L− 1 ≤ 1.00). All the measurements were carried out with an excess of the ligand, respect to the concentration of the tin(II), and in different Sn : L molar ratios (from 1:1 to 1:6). The potentiometric measurements were carried out by titrating 25 mL of the titrand solutions with standard NaOH solutions, under pH~ 7.5; up to this value the formation of sparingly soluble species was observed. For each experiment, independent titrations of strong acid solutions (HNO3) with standard NaOH solutions were carried out under the same medium and ionic strength conditions as the systems to be investigated, with the aim of determining electrode potential (E 0) and the acidic junction potential (Ej = ja [H+]). By using this procedure, the pH scale used was the free proton concentration scale, where pH≡ −log [H+]. The reliability of the calibration in the alkaline range was checked by calculating pKw values. For each titration, 80 to 100 data points were collected, and the equilibrium state during titrations was checked by adopting some usual precautions. These included checking the time required to reach equilibrium, the slope and performing back titrations [26].

The LIANA program was used in the calculation of the complex formation constants and to fit different equations. All these computer programs are described in reference [27]. The protonation constants of the ligands are given according to the equilibrium: þ

rH þ L

n−

¼ Hr L

ðr−nÞ

H

βr

ð1Þ

The formation constants of the complexes are given according to the equilibria: þ

2þ

þ rH þ L

2þ

þ rOH þ L

M

M

n−

−

¼ MHr L

n−

ð2þr−nÞ

¼ MLðOHÞr

βr

ð2−r−nÞ

ð2Þ OH

βr

ð3Þ

or 2þ

M

þ Hr L

ðr−nÞ

¼ MHr L

ð2þr−nÞ

ð4Þ

Kr

2.4. Calculations The non-linear least squares computer program ESAB2M was used for the refinement of all the parameters of the acid-base potentiometric titrations (E 0, Kw, liquid junction potential coefficient ja, analytical concentration of reagents). The ES4ECI program was used to draw the speciation diagram and to calculate species formation percentages.

MðOHÞr

ð2−rÞ

þL

n−

¼ MLðOHÞr

ð2−r−nÞ

OH

Kr

ð5Þ

where n is the charge of the different ligands (2 ≤ n ≤ 5, i.e. PO4: n = 3; PP: n = 4; TPP: n = 5; ATP: n = 4; MFP: n = 2); when r index is negative, Eq. (2) refers to hydrolysis reactions.

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R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

Table 2 Complex formation constants for the Sn/L systems, at different ionic strengths in NaNO3 and at T = 298.15 K. log βa Ligand PO4

PP

TPP

ATP

MFP

a b c d e

I/mol L− 1

log Kb

MLH2

MLH

c

19.20

ML c

15.90

MLOH c

c

0.15

22.70

0.50

22.16

18.50

14.95

1.00

22.00

18.10

14.1

0.15

22.9 ± 0.1e

21.0 ± 0.1e

16.2 ± 0.2e

0.50

22.11 ± 0.21

20.03 ± 0.10

15.42 ± 0.20

1.00

21.8 ± 0.1

19.5 ± 0.1

15.0 ± 0.2

0.15 0.50 1.00 0.15

– – – –

15.45 ± 0.15 14.32 ± 0.11 13.8 ± 0.15 11.08 ± 0.06

12.00 ± 0.15 11.16 ± 0.12 10.80 ± 0.15 8.31 ± 0.03

0.50

–

10.97 ± 0.05

8.09 ± 0.02

1.00

–

11.05 ± 0.05

8.41 ± 0.07

0.15

–

–

8.9 ± 0.1

0.50

–

–

8.39 ± 0.1

1.00

–

–

8.0 ± 0.1

9.2 22.95d 8.37c 22.10d 7.60c 21.33d 8.5 ± 0.2e 22.25d 7.12 ± 0.10 20.85d 6.0 ± 0.2 19.73d – – – 3.20 ± 0.05e 16.95d 2.78 ± 0.02 16.51d 3.24 ± 0.08 16.97d 4.6 ± 0.1e 18.35d 4.09 ± 0.20 17.82d 3.7 ± 0.2 17.43d

MLH2 4.56

MLH

c

7.71

c

ML 15.90

MLOH c

12.98c

4.48

7.27

14.95

12.28

4.54

6.99

14.1

11.60

8.51

12.58

16.2

12.28

8.54

12.12

15.42

11.03

8.83

11.97

15.0

10.00

– – – –

7.59 6.95 6.68 4.75

12.0 11.16 10.8 8.31

– – – 6.98

–

4.98

8.09

6.69

–

5.30

8.41

7.27

–

–

8.9

8.38

–

–

8.39

8.00

–

–

8.0

7.70

Refers to Eq. (2). Refers to Eqs. (4) and (5). [unpublished data from this laboratory]. Refers to Eq. (3). ± 95% C.I.

3. Results and discussion

constants, and log Kw values, at different ionic strengths, are reported in Table 1.

3.1. Hydrolysis, chloride complexation and protonation constants An extensive study on tin(II) hydrolysis was performed in a parallel investigation [Unpublished data from this laboratory]. The results were encouraging and in good agreement with values reported in the literature [28]; therefore they were used in this paper and are reported in Table 1, together with zinc(II) hydrolysis constant values, taken from Baes and Mesmer [29]. Even if equations proposed for the dependence of hydrolysis constants on ionic strength refer mainly to perchlorate media [30], the accuracy of the data is sufficient for our purposes. In fact, both nitrate and perchlorate anions can be considered as “weakly interacting” with the most of metal cations [30]. Moreover, for the cations investigated in this paper, the hydrolysis is strongly inhibited by the ligand complexation, as also observed in other cases [23,24,31–34]. Experimentally determined ligand protonation

3.2. Formation and stability of tin(II)–phosphates complexes The formation and the stability of tin(II)–phosphate ligands complexes were evaluated by using voltammetric and potentiometric techniques. Voltammetric data, for a reversible redox process and labile complex species, were used as reported elsewhere [23,24,35,36]. Our results are consistent with a reversible process and labile complex species. Reversibility was checked with Cyclic Voltammetry experiments on test solutions and ΔE (difference between cathodic and anodic peaks) resulted 30± 3 mV. As far as the lability is concerned we are supported by the following evidences: i) the peak of the free metal ion shifted to more negative potential values with addition of ligand and/or base without splitting into two or more peaks; ii) all the

Table 3 Complex formation constants for the Zn/L systems, at different ionic strengths and at T = 298.15 K. log βa) Ligand

I/mol L

PO4

0.15 0.50 1.00 0.15 0.50 1.00

TPP

a)

Refers to Eq. (2);

b)

−1

refers to Eq. (4);

c)

log Kb)

MLH2

MLH

ML

MLH2

MLH

ML

– – – 16.53 ± 0.06 14.94 ± 0.05 14.01 ± 0.08

14.3 ± 0.1c) 13.69 ± 0.2 13.4 ± 0.2 12.87 ± 0.03 11.62 ± 0.04 10.94 ± 0.03

7.30 ± 0.05c) 6.72 ± 0.2 6.43 ± 0.15 7.20 ± 0.03 6.36 ± 0.02 5.96 ± 0.03

– – – 3.11 2.43 1.97

2.81 2.46 2.29 5.01 4.25 3.82

7.30 6.72 6.43 7.20 6.36 5.96

± 95% C.I.

R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

1.0

a b

2

ligand; globally, the stability of the simple ML(2 − n) species formed by tin(II) with all the investigated ligands follows the trend:

3

PPePO4 > TPP >> MFPeATP

0.8

χ[Sn2+]

1 0.6

0.4

0.2

4

0.0 2

4

6

pH

Fig. 1. Speciation diagram of the Sn/ATP system. Fraction of Sn2+ with respect to pH. Experimental conditions: cSn = 1 mmol L− 1; cATP = 3 mmol L− 1; I = 0.15 mol L− 1; T = 298.15 K. Curves: 1: MHL; 2: ML; 3: MLOH, 4: M(OH)2. Part (a) represents the situation before the beginning of the precipitation; part (b) represents the hypothetical situation after the beginning of the precipitation (when the precipitation is ignored).

recorded peaks were symmetric and not distorted; iii) the intensity of the peaks did not decrease dramatically, even at high pH values. The conversion from the molar to molal scale for the different supporting electrolytes was obtained using the appropriate density values [37] and was 103 logaH2O = −13.66 m + 0.346 m2. Calculations performed on voltammetric and potentiometric data of investigated Sn/L systems gave evidence of the formation of different stable complexes; the values of the stability constants determined, at different ionic strengths and T = 298.15 K, are reported in Table 2. In any case, the interaction of tin(II) with chloride ions, present in solution for the use of HCl and SnCl2, was taken into account in the calculation by means of the stability constants reported elsewhere [Unpublished data from this laboratory]. Table 2 shows that phosphate and pyrophosphate ligands form four stable complexes with the metal, namely: [MLH2] (4 − n), [MLH] (3 − n), [ML](2 − n) and [M(OH)L](1 − n) with n = 3 for PO4 and n = 4 for PP. The [ML](2 − n) is the only common species determined for all the Sn/L systems; the monoprotonated complex species and the hydrolytic one, [MOHL] (1 − n), were not determined only for the Sn/MFP and Sn/TPP systems, respectively. From the values reported in Table 2 it can be also noted that the higher stability of the complex species is shown by the pyrophosphate

16

PP

PO4 log KSnL

14

TPP

12

10

ATP

8 5

147

6

7

8

log KZnL Fig. 2. Correlation between the stability of the ML species of the Sn/L and the Zn/L systems at I = 0.15 mol L− 1 in NaNO3 and T = 298.15 K. Dotted lines represent the confidence bands at 90% C.I.

Both pyrophosphate and tripolyphosphate form chelate species, as suggested by Ellison and Martell [38], with several metal ions. They can act as bidentate as well as terdentate ligands, depending on the pH of the solution. In any case the binding ability of tripolyphosphate is lower with respect to the one of pyrophosphate and phosphate. As concern the solubility limit, we calculated that in the voltammetric experimental condition no precipitation of Sn(OH)2(s) occurs in the whole investigated pH range. As concerns the Sn/ATP system, the only one studied by potentiometric titrations, we calculated that the precipitation of sparingly soluble species, in the experimental conditions (cL = 3 mmol L− 1; cM = 1 mmol L− 1), starts at pH ~5.5. Distribution diagrams of the Sn/PO4 and Sn/TPP systems are reported in the supporting information. 3.3. Formation and stability of Zinc(II)–phosphates complexes For the Zn/PO4 and Zn/TPP systems, voltammetric data evidenced (by least squares calculations, as it is described in previous section) the formation of stable complex species. With both ligands we determined the formation of two complexes, namely [MLH](3 − n) and [ML](2 − n) with n = 3 for PO4 and n = 5 for TPP, whilst only for the Zn/ TPP system the formation of the diprotonated complex species, [M (TPP)H2] −, was found. The values of the stability constants determined, at different ionic strengths and T = 298.15 K, are reported in Table 3. From these values we can see that the stability of the [ML](2 − n) species is very similar for both the phosphates ligands considered, whilst the stability of the monoprotonated species is higher for the Zn/TPP system respect to the Zn/PO4 system, with values of log K = 5.01 and 2.81 respectively (at I = 0.15 mol L − 1). As concerns the Zn/PP and the Zn/ ATP, we critically selected some literature data, owing to the high number of papers published on these systems. Distribution diagrams of the Zn/PO4 and Zn/TPP systems are reported in the supporting information. 3.4. Comparisons between the stability of Sn/L and Zn/L complexes Fig. 1 shows the speciation diagram related to the Sn/ATP system, drawn at I = 0.15 mol L − 1 and T = 298.15 K, in the experimental conditions, cM = 1 mmol L− 1; cL = 3 mmol L− 1. We can observe that, in the investigated pH range, tin(II) is fully complexed (the same consideration can be extended to the Sn/PO4, Sn/PP, Sn/TPP and Sn/MFP systems). The formation of the Sn(OH)2 species is observed only for the Sn/ATP system above pH= 5.5, accompanied by the formation of precipitate in the potentiometric measurements. In the case of Zn 2+, the phosphate ligands do not complex all the metal in solution. In the Zn/ PO4 system, the free metal ion is present up to pH= 8, when starts the formation of the Zn(OH)2 species, while the protonated MHL(3 − n) species represents the 0.3 of the metal molar fraction and the ML (2 − n) species the 0.8 at pH= 8.5. In the case of the Zn/TPP system, the formation of the Zn(OH)2 species is hampered, and the concentration of the free metal decreases rapidly and; at pH > 5 Zn2+ is totally complexed by the TPP ligand, the peaks of formation of the complex species are always higher than 0.6 of the metal molar fraction at pH= 2.5, pH= 5 and pH = 7.5 for the MH2L −, MHL 2− and ML3−, respectively. Where not specified, the final pH in the distribution diagrams also represents the Sn(OH)2(s) precipitation pH in our experimental conditions. Distribution diagrams for the Sn/PO4, Sn/TPP, Zn/PO4 and Zn/TPP systems are reported in the supporting information. In Fig. 2 it is reported the comparison between the ML(2 − n) stability constant values of the Sn/L and the Zn/L species (see Tables 2 and 3). In this last case the values of the ZnPP 2− and of the ZnATP2– species were selected from literature values. The selected stability constant values

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R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

were chosen on the basis of the experimental conditions in which they were determined, as close as possible to those used in this work. As an example in the case of the Zn/ATP system, the selected value was calculated by averaging the two most reliable values, determined in experimental conditions similar to the ones of this paper. Such values were reported by Sigel et al. [39] and by Cini et al. [40]. The values of the [ZnPO4] − (log K = 7.3) and [ZnTPP]3− (log K = 7.2) species are experimentally determined in this work. As regards the [ZnPP]2− species (log K = 6.7) we used the values reported by Smith et al. [28] at I = 0 mol L − 1 (log K = 8.7), corrected at I = 0.15 mol L − 1 by using the ionic strength dependence parameters, of the Debye–Hückel equation, hereafter reported for the SnPP2− species, considering that this model takes into account only the stoichiometry of the reactions. In this way, we calculated the value of the log βZnPP species at I = 0.15 mol L − 1, that was log K = 6.7. As we can see in Fig. 2, the plot is fitted by a straight line with a correlation coefficient R = 0.87 and represented by the following equation. logK ML ðSnÞ ¼ 3:01  logK ML ðZnÞ–8:13

ð6Þ

This correlation can be carefully used to guess the stability of unknown ML (2 − n) species, of both Sn2+ or Zn2+ complexes with phosphate ligands. In our case, we tested the goodness of this correlation by calculating the stability of the ML(2 − n) species of the Zn/MFP system. The corresponding value of the SnMFP species is log K = 8.9, and the calculated value for the ZnMFP is log KML = 5.66. The only available data for the Zn/MFP system is reported by Christie [41] and is log KML = 2.34 at I = 0.1 mol L − 1 in NaCl at T = 310.15 K. This value seems too low for a similar species, although determined in different experimental conditions, and for this reason we can assert that there is the needing of further experiments for the determination of the stability of the Zn/PP, Zn/ ATP and Zn/MFP species. 3.5. Ionic strength dependence The ionic strength dependence of the formation constants of the studied systems was modeled with two different approaches: the extended Debye–Hückel and the Specific ion Interaction Theory equations. The Debye–Hückel type model is expressed by the equation ⁎

0

logβ ¼ logβ −0:51⋅z ⋅ ⁎

2

1=2

I þ f ðI Þ 1 þ 1:5⋅I1=2

ð7Þ

2

where z ¼ ∑zreact −∑zprod and log β 0 is the value of the stability constant at infinite dilution; f(I) is a linear function of ionic strength

Table 4 Ligand protonation constant values, at infinite dilution together with the relative parameters of the ionic strength dependence, obtained by fitting values in Table 1 with Eq. (7), at T = 298.15 K. log βH0a,b Ligand

HL

H2L

H3L

PO4 C b) PP C TPP C ATP C MFP C

12.23c 0.094 9.42 ± 0.05d −0.222 ± 0.005 9.10 ± 0.01 0.043 ± 0.005 7.33 ± 0.01 0.057 ± 0.007 5.144 ± 0.001 −0.239 ± 0.001

19.36c 0.148 16.24 ± 0.06d − 0.451 ± 0.005 15.66 ± 0.01 0.075 ± 0.005 11.89 ± 0.06 0.660 ± 0.009 6.752 ± 0.001 − 0.359 ± 0.003

21.53c 0.162 18.56 ± 0.06d − 0.688 ± 0.005 18.22 ± 0.03 0.085 ± 0.005 – – – –

a b c d

Refers to Eq. (1). Parameter of Eq. (7). [unpublished data from this laboratory]. ± 95% C.I.

Table 5 Complex formation constant values for the Sn/L and Zn/L systems, at infinite dilution together with the relative parameters of the ionic strength dependence, obtained by fitting values in Tables 2 and 3 with Eq. (7), at T = 298.15 K. log β0a Sn2+ Ligand

MLH2

MLH

ML

MLOH

PO4

24.37 ± 0.05b

20.95 ± 0.05b

17.55 ± 0.05b

C PP

0.095 ± 0.003 25.56 ± 0.05

0.190 ± 0.003 23.48 ± 0.03

0.286 ± 0.003 18.19 ± 0.04

C TPP C ATP

0.190 ± 0.003 – – –

0.286 ± 0.003 18.63 ± 0.05 0.381 ± 0.003 13.41 ± 0.04

0.381 ± 0.003 14.43 ± 0.03 0.476 ± 0.003 10.08 ± 0.05

C MFP

– –

0.286 ± 0.003 –

0.381 ± 0.003 9.95 ± 0.04

C

– Zn2+ – – 20.12 ± 0.02b 0.088 ± 0.002

–

0.190 ± 0.003

10.59 ± 0.05b 24.59c 0.143 ± 0.003 10.22 ± 0.04 24.22c 0.190 ± 0.003 – – 3.96 ± 0.02 17.96c 0.190 ± 0.003 5.65 ± 0.03 19.65c 0.095 ± 0.003

8.71 ± 0.01b 0.088 ± 0.002 9.79 ± 0.01 0.146 ± 0.002

– – – –

PO4 C TPP C a b c

b

16.02 ± 0.02 0.058 ± 0.002 16.09 ± 0.02 0.117 ± 0.002

Refers to Eq. (2). ± 95% C.I. Refers to Eq. (3).

Table 6 Specific interaction coefficients obtained for the Sn/L systems. Cation 2+

Sn SnOH+ H+ Na+ Na+ Na+ H3PO4e Na+ SnHPO4e SnH2PO4+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ SnH2PPe Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ H2MFPe SnMFPe Na+ Na+ Na+ Na+ Na+ Na+ Na+ a b c d e

Anion

ε(+,−)a

NO3− NO3− NO3− PO43− HPO42− H2PO4−

0.18b 0.26b 0.064c 0.112 ± 0.035d 0.088 ± 0.064 0.080 ± 0.064 0.031 ± 0.064 1.237 ± 0.044 0.364 ± 0.023 − 0.064 ± 0.036 1.376 ± 0.049 − 0.605 ± 0.019d − 0.271 ± 0.045 − 0.159 ± 0.045 0.120 ± 0.045 − 0.531 ± 0.022 − 0.490 ± 0.017 − 1.044 ± 0.042 1.434 ± 0.084 0.038 ± 0.030d 0.051 ± 0.038 0.055 ± 0.038 0.034 ± 0.038 − 0.245 ± 0.033 − 0.223 ± 0.031 − 0.501 ± 0.010d − 0.218 ± 0.010 − 0.047 ± 0.013 − 0.040 ± 0.010 0.149 ± 0.021 0.072 ± 0.158d 0.068 ± 0.100 − 0.443 ± 0.262 − 1.290 ± 0.174 − 1.401 ± 0.205 − 0.718 ± 0.252

SnPO4− NO3− SnOHPO42− PP4− HPP3− H2PP2− H3PP− SnPP2− SnHPP− SnOHPP3− TPP5− HTPP4− H2TPP3− H3TPP2− SnTPP3− SnHTPP2− MFP2− HMFP−

SnOHMFP− ATP4− HATP3− H2ATP2− SnATP2− SnHATP− SnOHATP3−

Parameter of Eqs. (10) and (11). [unpublished data from this laboratory]. Ref. [44]. ± 95% C.I. km, parameter of Eq. (12).

R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153 Table 7 Specific interaction coefficients obtained for all the Zn/L systems.

Table 8 pL0.5 values of the Sn/L and the Zn/L systems at different ionic strengths in NaNO3 at T = 298.15 K.

Cation

Anion

ε(+,−)a

Zn2+ H+ Na+ Na+ Na+ H3PO4e Na+ ZnHPO4e Na+ Na+ Na+ Na+ Na+ Na+ Na+

NO3− NO3− PO43− HPO42− H2PO4−

0.16b 0.064c 0.112 ± 0.035d 0.088 ± 0.064 0.080 ± 0.064 0.031 ± 0.064 0.159 ± 0.037 0.108 ± 0.021 0.038 ± 0.030d 0.051 ± 0.038 0.055 ± 0.038 0.034 ± 0.038 − 0.217 ± 0.042 0.097 ± 0.024 0.056 ± 0.044

a b c d e

ZnPO4− TPP5− HTPP4− H2TPP3− H3TPP2− ZnTPP3− ZnHTPP2− ZnH2TPP−

Parameter of Eqs. (10) and (11). Ref. [2]. Ref. [44]. ± 95% C.I. km, parameter of Eq. (12).

that can be formulated in different ways. The simplest expression for this term is f(I) = CI, where C is the only adjustable parameter. Usually, this simple choice is sufficient to explain the experimental data trend in a wide ionic strength range, generally b1.0 mol L − 1. From a general point of view, the protonation reactions (Eq. (1)) or the complex formation reactions (Eq. (4)) can be expressed as a function of the activity coefficients as follows (charges omitted for simplicity): H

H0

logβr ¼ logβr þ logγ H þ logγL − logγ Hr L 0

ð9Þ

If the molal concentration scale is used, Eq. (7), becomes the SIT (Specific ion Interaction Theory) equation [42], where f(I) = Δε I; ε are the specific interaction coefficients between the species involved in the equilibria and the ions of the supporting electrolyte. If NaNO3 is used as supporting electrolyte, for the equilibrium in Eq. (1), it is:       þ þ n− þ ðr−nÞ − þ ε Na ; L −ε Na ; Hr L Δε ¼ rε H ; NO3

ð10Þ

20

PP

log KSnL

PO4 TPP 10

MFP

5 0.0

0.5

PP

TPP

MFP

ATP

Sn2+ I = 0.15 mol L− 1 5.0 6.50a 7.0 8.48 8.0 9.47

9.44a 8.30 7.64

5.99a 4.60 3.29

8.22a 8.30 8.31

4.69a 3.65 2.71

I = 0.50 mol L− 1 5.0 6.18 7.0 8.13 8.0 9.11

9.83 8.31 7.10

6.06 4.33 2.76

7.86 7.89 7.89

4.87 3.46 2.48

I = 1.00 mol L− 1 5.0 5.53 7.0 7.41 8.0 8.38

9.80 8.18 6.74

6.23 4.31 2.62

6.56 4.56 3.55

5.62 4.10 3.11

Zn2+ I = 0.15 mol L− 1 5.0 – 7.0 2.95 8.0 3.69

– – –

4.44a 6.28 6.91

– – –

– – –

I = 0.50 mol L− 1 5.0 – 7.0 2.67 8.0 3.46

– – –

4.03 5.83 6.20

– – –

– – –

I = 1.00 mol L− 1 5.0 – 7.0 2.38 8.0 3.14

– – –

3.84 5.59 5.81

– – –

– – –

pH

a

PO4

± 0.1 (95% C.I.).

ð8Þ

logK MHr L ¼ logK MHr L þ logγ M þ logγ Hr L − logγMHr L

15

149

1.0

(I / mol L-1)0.5 Fig. 3. Ionic strength dependence of the complex (ML species) stability constants. log KSnL values vs. (I/mol L− 1)0.5 at T = 298.15 K.

For the equilibrium in Eq. (4) (metal/ligand complexes), it is:   2þ − Δε ¼ ε M ; NO3     þ ðr−nÞ þ ðrþ2−nÞ − −ε Na =NO3 ; MHr L þ ε Na ; Hr L

ð11Þ

When, from the interaction between the metal ion and the ligand, the formation of a neutral species is observed, it is:     2þ þ 2− − þ ε Na ; Hðn−2Þ L −km Δε ¼ ε M ; NO3

ð12Þ

where km is the Setschenow coefficient [43] of the neutral species. As concerns the extended Debye–Hückel approach, the equilibrium constant values at infinite dilution and the relative empirical parameter C for the ligand protonations, Sn/L and Zn/L complex formation constants are reported in Tables 4 and 5, respectively. Furthermore, we used SIT approach to model the ionic strength dependence of the stability constants in Tables 1, 2 and 3. We choose for simplicity to use the log K values for the complex formation constants instead of the log β ones. We used known specific interaction coefficients for the ε(H+, NO3−) = 0.064 [44], ε(Sn2+, NO3−) = 0.18 [Unpublished data from this laboratory], ε(Sn(OH) +, NO3−) = 0.26 [Unpublished data from this laboratory] and ε(Zn2+, NO3−) = 0.16 [2]. To determine the specific interaction coefficients an approximation was needed, and we used the one reported by Ciavatta [45], that consider ε(Na+, ML(2 − n)) = 1/2 (ε(M2+, NO3−) + ε(Na +, L n−)). We applied this approximation to the various protonated species of the ligands, as an example ε( Na+, HPP3−) = 1/2 (ε(H2+, NO3−) + ε(Na +, PP4−)) and ε( Na +, H2PP2−) = 1/3 (2 ε(H+, NO3−) + ε(Na +, PP4−)). The calculated specific interaction coefficients are reported in Table 6 for the protonated and Sn/L complex species (together with the known SIT coefficients) and in Table 7 for the Zn/L complex species.

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R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

1.0

Table 9 Available literature data for the Sn/L systems.

x [Sn2+]

0.8

0.6

2

3

4

1

L

Equilibriuma

log K

T/K

I/mol L− 1

Medium

Ref.

PO4

M + L = ML M + HL = MHL M + HL = MHL M + H2L = MH2L M + 2H2L = M(H2L)2 M + 2HL = M(HL)2 M + HL + H2L = M(HL)(H2L) M + 3HL = M(HL)3 M + L = ML M + L = ML M + L = ML M + 2 L = ML2 M + 3 L = ML3 M + 2 L + H = MHL2 M + 2H + 2 L = M(HL)2 M + 3H + 2 L = MH3L2 M + OH + L = M(OH)L 2MO + H3O+ + 2HL = M2H3L 2MO + H3O+ + 2H2L = M2H5L

18.0 9.5 7.8 2.8 5.9 13.4 10.3 12.9 12.0 13.05 14.30 15.5 18.4 22.7 28.3 32.1 6.0 7.25 5.0

298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15

0 0 0.2 0 0 0 0 0 0.1 1.0 1.0 0.1 0.1 0.1 0.1 0.1 0.1 1.0 1.0

NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 NaCl NaClO4 K2SO4 NaCl NaCl NaCl NaCl NaCl NaCl NaClO4 NaClO4

[19] [19] [18] [19] [19] [19] [19] [19] [20] [51] [52] [20] [20] [20] [20] [20] [20] [21] [21]

5

0.4

PP

0.2

0.0 2

4

6

8

10

pL Fig. 4. Sequestration diagram of the Sn/L systems at I = 0.15 mol L− 1 in NaNO3, T = 298.15 K and pH = 7.0. Fraction of Sn2+ vs. pL. Curves: 1: PO4, 2: PP, 3: TPP, 4: ATP, 5: MFP.

TPP a

Furthermore in Fig. 3, is shown as an example the ionic strength dependence of the SnL(2 − n) species with respect to the square root of the ionic strength. It can be noted that, except for the Sn/ATP species, the stability of all the complexes decreases with increasing ionic strength in the investigated pH range (0.15 ≤ I/mol L − 1 ≤ 1.00). 3.6. Sequestering ability of different phosphate ligands towards tin(II) and zinc(II) Recently [31,32,46] we proposed a Boltzmann type equation able to define the sequestering ability of a ligand toward a metal ions by means of the function x (fraction of a generic metal ion M complexed by a generic ligand L) vs. pL, where pL = −log(L)tot, with (L)tot = total ligand concentration. This function is represented by a sigmoid curve (or a dose–response curve), with asymptotes of 1 for pL → −∞ and 0 for pL → +∞:  x¼

1 1 þ 10ðpL−pL0:5 Þ

 ð13Þ

the parameter pL0.5 represents the ligand concentration necessary to sequester 50% of the metal ion (or another ligand). Therefore this parameter can be used as a measure of the sequestering ability of a ligand towards a metal cation (in our case Sn 2+ or Zn 2+).

1.0

x [M2+]

0.8

0.6

3

2

1

4

0.4

0.2

0.0 2

4

6

8

10

pL Fig. 5. Sequestration diagram of the M/L systems at I = 0.15 mol L− 1 in NaNO3, T = 298.15 K and pH = 7.0. Fraction of M2+ vs. pL. Curves: 1: Sn/PO4, 2: Zn/PO4, 3: Sn/TPP, 4: Zn/TPP.

Charges omitted for simplicity.

In this paper the pL0.5 values were calculated at three different pH values (i.e. pH= 5.0, 7.0 and 8.0) and at different ionic strengths for all the investigated systems. The values obtained in the case of Sn/TPP and Sn/ATP systems at pH= 8.0 can suffer from the fact that the end titration pH was always pH ≤ 7.5. In Table 8 the pL0.5 values are summarized for the Sn/L and Zn/L systems. As we can see, in all cases the sequestering ability of a ligand decreases with increasing ionic strength, with exception of the Sn/ATP system, for which the pL0.5 values trace the behavior of the stability constant values. As concerns the pH dependence, in the case of the Sn/L systems, we observe an increase of the pL0.5 with respect to pH only for the Sn/PO4 system, while for the Zn/L systems this trend was observed in all cases. This difference can be easily explained with the different hydrolysis behavior of the two metals. In the case of tin(II), the very strong hydrolysis causes a decrease of the pL0.5 values (except for the Sn/PO4 system), while in the case of zinc(II) this problem is not encountered because the hydrolysis of zinc (II) occurs at higher pH values, pH > 7. Some details should also be given, as an example, the strongest decrease of the pL0.5 values with pH is shown by the TPP ligand, due to the absence of an hydrolytic species in the speciation diagram (see Table 2). At pH = 5, the PP and TPP ligands show an increasing of the pL0.5 values with ionic strength, differently to MFP, due to the high numerical value of the SnHL species, that is missing in the Sn/MFP system. Finally, for the Sn/MFP system, the ionic strength dependence is stronger than the other ligands, probably due to the decrease of the stability constant values that in this last case plays a more important role. In Fig. 4 we reported the sequestration diagram of the different Sn/L systems at pH=7.0 and I=0.15 mol L− 1. We can see that phosphate shows the highest sequestering ability (pL0.5 =8.48), slightly higher than PP and MFP, whose sequestering ability is comparable (for both pL0.5 =8.3); much lower is the pL0.5 values of TPP (4.60) and ATP (3.65). Quite surprisingly is the very high sequestering ability of MFP, but we think that the presence of fluoride confers a higher acidity to the molecule, and lower protonation constant values. This, at first sight, strange behavior of MFP demonstrates that from the evaluation of pL0.5 more objective comparisons can be obtained with respect to the simple evaluation of the stability constant values, taking into account also the contribute of the “side-reactions”, like protonation of the ligand, hydrolysis of the metal cation or competition of another ligand or metal. Furthermore a comparison between the sequestering ability of the phosphate ligands toward tin(II) and zinc(II) have been made; in Fig. 5 the sequestration diagram, at I = 0.15 mol L − 1 in NaNO3 and pH= 7, for Sn/PO4, Sn/TPP, Zn/PO4 and Zn/TPP systems are shown. It can be noted that the two ligands considered, phosphate and

R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153 Table 10 Available literature data for the Zn/L systems. Equilibriuma

L PO4

M + HL = MHL M + HL = MHL M + HL = MHL M + HL = MHL M + HL = MHL M + H2L = MH2L M + H2L = MH2L MH2L + H = M(HL)2 M(HL)2 + H = MH3L2 MH3L2 + H = M(H2L)2 M + 2H2L = M(H2L)2 PP M + L = ML M + L = ML M + 2 L = ML2 M + 2 L = ML2 M + 2 L = M2L TPP M + L = ML M + L = ML M + L = ML M + L = ML M + HL = MHL M + HL = MHL M + HL = MHL MFP M + L = ML M + OH + L = M(OH)L ATP M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + L = ML M + H + L = MHL M + H + L = MHL M + H + L = MHL M + HL = MHL M + HL = MHL M + HL = MHL M + 2H + L = MH2L M + 2 L = ML2 2 M + L = M2L 2 M + H + L = M2HL M + OH + L = M(OH)L M + OH + L = M(OH)L 2 M + OH + L = M2(OH)L M(OH)L + H = ML a

log K

T/K

I/mol L− 1 medium

Ref.

2.40 2.4 2.4 4.60 4.16 1.2 0.37 5.76 3.3 4.9 1.10 5.1 8.7 7.19 11.0 7.50 9.7 7.3 6.83 8.43 3.75 3.92 5.13 2.34 − 4.22 5.16 5.23 5.52 5.16 5.81 4.8 4.92 4.85 4.9 4.08 4.85 10.25 9.22 9.72 2.86 2.67 2.94 13.85 8.27 8.26 11.98 − 2.56 − 3.2 1.56 8.76

298.15 298.15 310.15 298.15 298.15 310.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 308.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 310.15 310.15 298.15 298.15 303.15 298.15 298.15 – 298.15 298.15 298.15 298.15 298.15 298.15 298.15 303.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15

0.1 0.1 0.15 0 0.01 0.15 3.0 3.0 3.0 3.0 3.0 1.0 0 1.0 0.5 0 0 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 – 0.1 0.1 0.1 0.12 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1

[53] [39] [63] [54] [55] [63] [56] [56] [56] [56] [56] [57] [28] [57] [58] [64] [28] [38] [65] [28] [65] [38] [28] [41] [41] [59] [61] [66] [67] [40] [68] [69] [70] [60] [71] [72] [40] [61] [66] [59] [72] [69] [40] [40] [40] [40] [40] [61] [40] [28]

NaClO4 NaClO4 KNO3 Self-medium NaClO4 KNO3 NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 NaNO3 – NaNO3 NaClO4 Self-medium – KCl KNO3 R4N+ KNO3 KCl R4N+ NaCl NaCl NaClO4 KNO3 Me4NBr Me4NBr NaClO4 – KCl KNO3 NaNO3 NaCl KNO3 NaClO4 KNO3 Me4NBr NaClO4 KNO3 KCl NaClO4 NaClO4 NaClO4 NaClO4 NaClO4 KNO3 NaClO4 Na+

Charges omitted for simplicity.

tripolyphosphate, have different sequestering ability toward the two metal ions. We can see that PO4 is a much stronger complexing agent than TPP (pL0.5 = 8.48 and 4.60 respectively) towards Sn2+, on the contrary of Zn/L ones (pL0.5 = 6.28 for Zn/TPP and 2.94 for Zn/PO4). The sequestering ability of PO4 and TPP toward the two metal ions follows the general trend: Sn=PO4 > Zn=TPP > Sn=TPP > Zn=PO4

4. Literature comparison The ligand protonation constants here reported are in a good agreement with those reported in literature, although the values have been reported in sodium chloride and in potassium chloride media. Slight difference can be found, that can be ascribed to the different background salt and in turn different activity coefficients. As an example the values of the phosphate protonation constants are very close to those reported

151

in Daniele et al. [47] at I = 0.15 mol L − 1, but they slightly differ at higher ionic strengths. Similar consideration can be done as regards pyrophosphate, tripolyphosphate, adenosine-5′-triphosphate and mofluorophosphate ligands. In any case, the literature data (e.g. in Refs. [41,48–50]) are mostly reported in NaCl medium, and our values, determined in NaNO3 aqueous solutions, are slightly higher. Very poor is the literature on the interaction of tin(II) and the ligands studied in this paper. As concerns the Sn/PO4 system, only two papers were published, by Cilley [18] and Ciavatta et al. [19]. Cilley reported the formation of the neutral SnHPO4 species, with a value of log K = 7.8 (at I = 0.2 mol L − 1 in NaClO4), very close to our value of log K = 7.71 (at I = 0.15 mol L − 1 in NaNO3), by solubility measurements at constant pH (pH = 2.8) value and thus without information on the number of protons bounded to the complex species and on the stoichiometry of the species. Ciavatta and Iuliano reported a potentiometric study, by means of a tin amalgam electrode, on the tin(II) phosphate system at T = 298.15 K and I = 3 mol L − 1 in NaClO4. The authors determined seven complex species, three mononuclear species, namely: SnPO4−, SnHPO4 and SnH2PO4+ and four polynuclear, namely: Sn (H2PO4)2, Sn(HPO4)(H2PO4) −, Sn(HPO4)22− and Sn(HPO4)34− owing to concentration of phosphate (0.01 to 0.3 mol L− 1) much higher than the concentration of the metal (0.25 to 2.5 mmol L − 1). The authors report the stability constant values corrected at infinite dilution, these values are in a good agreement with our values at infinite dilution (see Table 3), with the exception of the SnH2PO4+ species, whose value is 2.2 log units higher. This difference can hardly be explained, but the authors consider many polynuclear species, that can presumably lead to underestimation of the stability of this species. In our case we could not determine the stability of polynuclear species, because the concentration of phosphate used in our measurements was too low for this purpose. Mesmer and Irani [21] reported the interaction of tin(II) with both pyrophosphate and tripolyphosphate, basing their study on solubility measurements. Unfortunately, they only determined the formation of polynuclear species and comparisons with our study cannot be made. Only one paper, published by Duffield et al. [20], reports the interaction of tin(II) with pyrophosphate. The stability constants values obtained by Duffield are summarized in Table 9, and are quite different from our data. They also report the solubility product of the neutral Sn2(P2O7) species as log Ks0 = 19.65 (at T = 298.15 K and I = 0.15 mol L − 1 NaCl). Other authors reported values of log KML = 13.05 [51] and log KML = 14.3 [52] in NaClO4 (1 mol L − 1) and K2SO4 (1 mol L − 1), respectively. No other comparison can be made with other ligands, because to our knowledge no papers have been published. In the case of the zinc(II) more investigations are reported in literature (see Table 10). As concerns the Zn/PO4 system, some comparison can be made with Ramamorthy and Manning [53] and Sigel et al. [39]. Both authors performed potentiometric measurements and determined a value of log KMHL = 2.4 (I = 0.1 mol L− 1 in NaClO4). This value is slightly different than our value of log K = 2.9 at the same ionic strength. Ziemniak et al. [54] reported a solubility study at different temperatures, and it was possible to calculate a value of log KM + HL = 4.6 (I = 0 mol L − 1). Vega et al. [55] reported a pseudopolarographic study on the Zn/PO4 system at (I = 0.01 mol L − 1 in NaClO4), and observed the formation of the ZnPO4− and Zn(PO4)37− species, the value of the ZnPO4− species was log KM + HL = 4.16, quite surprisingly they did not determine the Zn(PO4)24− species. Both the values reported by Ziemniak et al. [54] and Vega et al. [55] are slightly higher than our calculated values at the same ionic strength, log KMHL = 3.8 and log KMHL = 3.44 at I = 0 mol L − 1 and I = 0.01 mol L − 1, respectively. Iuliano [56], performing a potentiometric study with glass and zinc amalgam electrodes, found the existence of five complex species, MH2L +, MHL23−, MH2L22−, MH3L2− and MH4L2. No comparisons can be made with our values, because no common species are present. As concerns the Zn/PP system, Costley and Farr in 1968 [57] found the existence of the ZnPP2− and Zn(PP)2 species in NaNO3

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(T= 298.15 K and I = 1.00 mol L − 1) by conductimetric titrations. Mushkina et al. [58] published a paper on the Zn/PP system, and determined the stability of the ML26− species at different ionic strengths in NaClO4, at I = 0.5 mol L − 1 as an example the log β = 11.0. Other values are reported in the common stability constants database [28], as an example a value of log KML = 8.7 at I = 0 mol L − 1. However, a higher stability for the ZnPP 2− species, with respect to both ZnPO4− and ZnTPP 3−, is suggested by the possible formation of a six term chelate species. As regards the Zn/TPP system, all the values (see Table 10) are reported in ionic medium not containing sodium (i.e. K+ or R4N+), and owing to the different interaction between the tripolyphosphate anion and the cation of the ionic medium no significative comparisons can be made. In any case, our value log KML =7.20 is close to those reported in literature. Moreover, our data confirm that the proton affinity of the metal chelates is inversely proportional to the stability of the latter, as reported in Ellison and Martell [38]. For the Zn/MFP system, only one paper was published, by Christie et al. [41], at T = 310.15 K and I = 0.1 mol L − 1 in NaCl. Although no comparisons can be made, we can see that the stability constant values of the two species determined are significantly lower than the analogous values of the other ligands, as we observed in the case of the tin (II) systems. The Zn/ATP system was widely studied by many authors, the stability constant values, determined at various ionic strengths and in different ionic media are reported in Table 10. The values are scattered in a wide range, the ones reported in sodium media are log KML = 5.16 (in NaClO4 at I = 0.1 mol L − 1) by Sigel et al. [59], log KML = 5.81 (in NaClO4 at I = 0.1 mol L − 1) by Cini et al. [40], log KML = 4.9 (in NaNO3 at I = 0.1 mol L − 1) by Khalil [60], and log KML = 4.08 (in NaClO4 at I = 0.12 mol L − 1). Different protonated species were also reported in literature and also in this case there is a spread in the stability constant values. Cini et al. [40] also reported the formation of the ZnH2ATP2+ and the Zn(OH)ATP− species, as well as Arena et al. [61]. The speciation model reported by these last two authors is for this reason similar to the one reported in this paper for the Sn/ATP system. Some dinuclear species were reported by different authors in the case of the Zn/ATP system, as an example by Arena et al. [61], that reported the formation of both ML22− and M2L 2+ species, Sigel et al. [59] the M2L 2+ and the M2HL3+ ones. Handschin and Brintzinger [62] reported the formation of the ML22− species. Cini et al. found the M2LOH+. In this study we could not determine any polynuclear species, because the concentration of reagents used in the experiments was too low for this kind of species, furthermore we adopted in all cases a large excess of ligand, and for this reason the M2L 2+ species has not significance in our experimental conditions. In general, as concerns tin(II) speciation in the presence of phosphate ligands, the major species are the SnL and the SnHL (charges omitted). The SnOHL and the SnH2L specie can be considered as minor species as well as other species reported in the literature. If the ligand concentration is cL ≥ 10 mmol L − 1 also the polynuclear species, reported only by Ciavatta, should be added to the speciation model. Similar considerations can be extended to the zinc(II) speciation. Some of these species are formed in particular conditions of concentration and pH and their determination, with sufficient accuracy, is difficult. 5. Conclusions Potentiometric and voltammetric measurements were performed in this work to study the speciation of the tin(II) cation with different phosphate ligands, such as: PO4, PP, TPP, ATP and MFP at T = 298.15 K in NaNO3 at different ionic strengths (0.15≤ I/mol L − 1 ≤ 1.00). The protonation constants of all the ligands were also determined by potentiometric titrations in NaNO3 at different ionic strengths (0.15 ≤ I/ mol L− 1 ≤ 1.00). In all cases we provided different stability constant

values, with similar speciation schemes for the different ligands. These values were then compared with other Zn/L systems, some experimentally determined (Zn/PO4 and Zn/TPP) and some taken from literature (Zn/PP and Zn/ATP). In all cases the stability of the tin(II) species was higher with respect the analogous zinc species. As concerns the stability of the ML(2 − n) species with respect to the ligand, the pyrophosphate shown the higher binding ability, probably due to the higher tendency of this anion to form chelates (see Ellison and Martell [38]). We also found a correlation between the stability of the Sn/L species and the Zn/L species. By using this relationship we guessed the stability constant value of the ZnMFP species, that is reported in literature only at T = 310.15 K, and was log KML = 5.66. The ionic strength dependence of the complex species was studied by both the extended Debye– Hückel and SIT approaches and the specific interaction coefficient were provided, together with infinite dilution log β values, for all the studied interactions. Speciation and sequestration studies were also carried out, and we found that all the phosphate ligands strongly bind the Sn2+ cation, also the pH region of the metal hydrolysis. Phosphate and pyrophosphate were found to be the best sequestering agents towards tin, mostly at pH = 8.0 for PO4 and pH= 5.0 for PP, whilst TPP showed an higher sequestering ability toward zinc(II) with respect to tin(II) at pH = 7 and I = 0.15 mol L − 1 (see Fig. 5). Acknowledgement We thank the University of Messina for partial financial support (PRA). Appendix A. Supplementary data Supplementary data to this article can be found online at doi:10. 1016/j.molliq.2011.11.002. References [1] A. Gianguzza, E. Pelizzetti, S. Sammartano (Eds.), Chemistry of Marine Waters and Sediments, Springer-Verlag, Berlin, 2002. [2] I. Grenthe, I. Puigdomenech (Eds.), Modelling in Aquatic Chemistry, OECD, Paris, 1997. [3] F.J. Millero, in: E.D. Goldberg (Ed.), The Sea, Wiley, New York, 1974, pp. 3–80. [4] F.J. Millero (Ed.), Chemical Oceanography, 2nd, CRC Press, Inc, Boca Raton, FL, 1996. [5] G.G. Leppard, M. Smies (Eds.), Trace Element Speciation in Surface Waters and its Ecological Implications, Plenum, New York, 1983, p. 177. [6] Tin in canned fruit and vegetables, Food Standard Agency, August 2002. [7] J.P. Schapira, Le dossier des déschets nucléaires, Société Francaise de Physique, 1997, pp. 3–24. [8] S.J. Blunden, A. Chapman, in: P.J. Craig (Ed.), Organometallic Compounds in the Environment, Longman, London, 1986, pp. 111–159. [9] C. Alzieu, L'étain et les Organoétains en Milieu Marin-Biogéochimie et Ecotoxicologie, IFREMER, 1989. [10] F. Crea, C. De Stefano, D. Milea, S. Sammartano, Coordination Chemistry Reviews 252 (2008) 1108–1120. [11] F.J. Millero, in: F.J. Millero (Ed.), Chemical Oceanography, 2nd Edition, CRC Press, Boca Raton, FL, 1996. [12] J.E. Ellingsen, J. EkSreand, Journal of Dental Research 64 (1985) 1250–1252. [13] R.A. Hiles, Toxicology and Applied Pharmacology 27 (1974) 366–379. [14] Merck, in: M. Windholz (Ed.), The Merck Index, Rahway, New Jersey, 1976, pp. 8564–8568. [15] J.L. Greger, M.A. Johnson, Food and Cosmetics Toxicology 19 (1981) 163–166. [16] M.A. Johnson, M.J. Baier, J.L. Greger, American Journal of Clinical Nutrition 35 (1981) 1332–1338. [17] N.W. Solomons, J.S. Marchini, R.M. Duarte-Favaro, H. Vannuchi, J.E. Dutra de Oliveira, American Journal of Clinical Nutrition 37 (1983) 566–571. [18] A.W. Cilley, Inorganic Chemistry 7 (1968) 612–614. [19] L. Ciavatta, M. Iuliano, Polyhedron 19 (2000) 2403–2407. [20] J.R. Duffield, D.R. Williams, I. Kron, Polyhedron 10 (1991) 377–387. [21] R.E. Mesmer, R.R. Irani, Journal of Inorganic and Nuclear Chemistry 28 (1966) 493–502. [22] H.A. Flaschka (Ed.), EDTA Titration, Pergamon, London, 1959. [23] R.M. Cigala, F. Crea, C. De Stefano, G. Lando, D. Milea, S. Sammartano, Journal of Chemical and Engineering Data 55 (2010) 4757–4767. [24] C. De Stefano, G. Lando, D. Milea, A. Pettignano, S. Sammartano, Journal of Solution Chemistry 39 (2010) 179–195. [25] R.M. Cigala, F. Crea, C. De Stefano, G. Lando, D. Milea, S. Sammartano, accepted for publication. Modeling the acid-base properties of glutathione in different ionic

R.M. Cigala et al. / Journal of Molecular Liquids 165 (2012) 143–153

[26] [27]

[28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38] [39] [40] [41] [42] [43] [44] [45] [46] [47]

media, with particular reference to natural waters and biological fluids, Amino Acids (in press), doi:10.1007/s00726-011-1110-0. A. Braibanti, G. Ostacoli, P. Paoletti, L.D. Pettit, S. Sammartano, Pure and Applied Chemistry 59 (1987) 1721–1728. C. De Stefano, S. Sammartano, P. Mineo, C. Rigano, in: A. Gianguzza, E. Pelizzetti, S. Sammartano (Eds.), Marine Chemistry—An Environmental Analytical Chemistry Approach, Kluwer Academic Publishers, Amsterdam, 1997, pp. 71–83. A.E. Martell, R.M. Smith, R.J. Motekaitis, National Institute of Standard and Technology, NIST, Vers. 8 2004 (PC-based Database: Gaithersburg, MD), 2004. C.F. Baes, R.E. Mesmer (Eds.), The Hydrolysis of Cations, Wiley, New York, 1976. P.G. Daniele, C. Foti, A. Gianguzza, E. Prenesti, S. Sammartano, Coordination Chemistry Reviews 252 (2008) 1093–1107. F. Crea, C. De Stefano, D. Milea, S. Sammartano, Journal of Solution Chemistry 38 (2009) 115–134. P. Crea, C. De Stefano, D. Milea, N. Porcino, S. Sammartano, Biophysical Chemistry 128 (2007) 176–184. C. De Stefano, D. Milea, N. Porcino, S. Sammartano, Analytical and Bioanalytical Chemistry 386 (2006) 346–356. C. De Stefano, D. Milea, N. Porcino, S. Sammartano, Journal of Agricultural and Food Chemistry 54 (2006) 1459–1466. I. Cukrowski, M. Adsetts, Journal of Electroanalytical Chemistry 429 (1997) 129–137. I. Cukrowski, R.D. Hancock, R.C. Luckay, Analytica Chimica Acta 319 (1996) 39–48. M.J. Astle, W.H. Beyer, R.C. Weast (Eds.), Handbook of Chemistry and Physics, 66th, CRC Press, Inc, Boca Raton, FL, 1985–1986. H. Ellison, A.E. Martell, Journal of Inorganic and Nuclear Chemistry 26 (1964) 1555–1560. S.H. Sigel, K. Becker, D.B. Mc Cormick, Biochimica et Biophysica Acta 148 (1967) 655–664. R. Cini, A. Cinquantini, R. Seeber, Inorganica Chimica Acta 123 (1986) 69–73. G.L. Christie, D. Cummins, J.R. Duffield, S.R. Hurford, C.R. Morris, P.I. Riley, J.A. Vesey, D.R. Williams, Journal of Inorganic Biochemistry 42 (1991) 273–287. G. Biederman, in: E.A. Jenne, E. Rizzarelli, V. Romano, S. Sammartano (Eds.), Metal Complexes in Solution, Piccin, Padua, Italy, 1986, pp. 303–314. J.Z. Setschenow, Zeitschrift für Physikalische Chemie 4 (1889) 117–125. C. Bretti, C. Foti, N. Porcino, S. Sammartano, Journal of Solution Chemistry 35 (2006) 1401–1415. L. Ciavatta, Annales de Chimie (Rome) 80 (1990) 255–263. F. Crea, C. De Stefano, D. Milea, S. Sammartano, Monatshefte fur Chemie 141 (2010) 511–520. P.G. Daniele, A. De Robertis, C. De Stefano, A. Gianguzza, S. Sammartano, Journal of Solution Chemistry 20 (1991) 495–515.

153

[48] A. De Robertis, C. De Stefano, S. Sammartano, R. Calì, R. Purrello, C. Rigano, Journal of Chemical Research 164 (1986) 1301–1347. [49] C. De Stefano, C. Foti, A. Gianguzza, Journal of Chemical Research 464 (1994) 2639–2661. [50] J.W. Larson, B. Su, Journal of Chemical and Engineering Data 39 (1994) 33–35. [51] I.Y. Tur'yan, V.I. Kravtsov, V.V. Kondrat'ev, Elektrokhimiya 22 (1986) 1618–1624. [52] V.V. Orekhova, F.K. Andryushenko, N.D. Sakhnenko, Elektrokhimiya 16 (1980) 1304–1310. [53] S. Ramamoorthy, P.G. Manning, Journal of Inorganic and Nuclear Chemistry 36 (1974) 695–698. [54] S.E. Ziemniak, M.E. Jones, K.E.S. Combs, Journal of Solution Chemistry 21 (1992) 1153–1176. [55] M. Vega, R. Pardo, M.M. Herguedas, E. Barrado, Y. Castrillejo, Analytica Chimica Acta 310 (1995) 131–138. [56] M. Iuliano, Annales de Chimie (Rome) 84 (1994) 187–209. [57] B.D. Costley, J.P.G. Farr, Chemistry and Industry (1968) 1435–1437. [58] M.G. Mushkina, M.S. Novakovskii, N.C. Uen, Russian Journal of Inorganic Chemistry 12 (1967) 1239–1242. [59] H. Sigel, R. Tribolet, R. Balkrishnan-Malini, B. Martin, Inorganic Chemistry 26 (1987) 2149–2157. [60] M. Khalil, Journal of Chemical and Engineering Data 45 (2000) 70–74. [61] G. Arena, R. Calì, V. Cucinotta, S. Musumeci, E. Rizzarelli, S. Sammartano, Journal of the Chemical Society, Dalton Transactions (1983) 1271–1278. [62] Y. Handschin, H. Brinzinger, Helvetica Chimica Acta 45 (1962) 1037–1040. [63] C.W. Childs, Inorganic Chemistry 9 (1970) 2465–2469. [64] B.C. Haldar, Current Science (India) 19 (1950) 283. [65] M.M. Taqui Khan, P. Rabindra Reddy, Journal of Inorganic and Nuclear Chemistry 35 (1973) 179–186. [66] D.D. Perrin, V.S. Sharma, Biochimica et Biophysica Acta 127 (1966) 35–41. [67] R.M. Smith, A.E. Martell, Y. Chen, Pure and Applied Chemistry 63 (1991) 1015–1080. [68] H. Brinzinger, Helvetica Chimica Acta 44 (1961) 935–939. [69] N.K. Davidenko, P.A. Manorik, K.B. Yatsimirskii, Russian Journal of Inorganic Chemistry 25 (1980) 883–890. [70] H.A. Azab, F.S. Deghaidy, A.S. Orabi, N.Y. Farid, Journal of Chemical and Engineering Data 45 (2000) 709–715. [71] K.S. Rajan, S. Mainer, J.M. Davis, Bioinorganic Chemistry 9 (1978) 187–203. [72] M.M. Taqui Khan, A.E. Martell, Journal of the American Chemical Society 88 (1966) 668–671.