Thermodynamic parameters for the interaction between etidronic acid and inorganic and organic mercury(II)

Accepted Manuscript Thermodynamic parameters for the interaction between etidronic acid and inorganic and organic mercury(II) Donatella Chillè, Claudia Foti, Ottavia Giuffrè PII: DOI: Reference:

S0021-9614(18)30047-8 https://doi.org/10.1016/j.jct.2018.02.009 YJCHT 5329

To appear in:

J. Chem. Thermodynamics

Received Date: Revised Date: Accepted Date:

2 November 2017 10 January 2018 5 February 2018

Please cite this article as: D. Chillè, C. Foti, O. Giuffrè, Thermodynamic parameters for the interaction between etidronic acid and inorganic and organic mercury(II), J. Chem. Thermodynamics (2018), doi: https://doi.org/ 10.1016/j.jct.2018.02.009

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Thermodynamic parameters for the interaction between etidronic acid and inorganic and organic mercury(II)

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Donatella Chillè, Claudia Foti,* Ottavia Giuffrè

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Dipartimento di Scienze Chimiche, Biologiche, Farmaceutiche ed Ambientali, Università di Messina, Viale F. Stagno d’Alcontres, 31 – 98166 Messina, Italy.

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Abstract

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Thermodynamic parameters of etidronic acid [(1-Hydroxy-1,1-ethanediyl)bis(phosphonic acid), HEDPA] interaction with Hg2+ and CH3 Hg+ were determined in NaCl aqueous solution

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at different ionic strength values (0.1 ≤ I ≤ 1 mol·L-1) and at T = 298.15 K. For Hg2+ system, ML2-, MLH20, MLH3+ and ML(OH)3- species were obtained. In presence of CH3 Hg+, ML3-, MLH2-, MLH2-, MLH30 and the dinuclear M2LH- and M2LH20 species were determined with formation constant values lower respect to the Hg2+ species. As an example, logβ = 7.98 and 15.61 were obtained for (CH3 Hg)L3- and HgL2- species, respectively (at I ≈ 0.1 mol·L-1 and T = 298.15 K). The dependence of the stability constants on ionic strength, studied through a Debye-Hückel type equation, is also reported. To enrich the thermodynamic study, the enthalpy values of the complex species were obtained by calorimetric titrations at I = 0.1 mol·L-1 in NaCl and T = 298.15 K. Finally, the sequestering ability of HEDPA towards Hg2+ and CH3 Hg+ was calculated by means of an empirical parameter, known as pL0.5, under blood (I = 0.1 mol·L-1, pH = 7.4, T = 310.15 K) and sea water (I = 0.7 mol·L-1, pH = 8.1, T = 298.15 K) conditions. pL0.5 represents the total ligand concentration required to sequester the 50% of a metal cation present as traces.

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Keywords

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Etidron acid; Hg(II); speciation; thermodynamic parameters; sequestering ability.

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1. Introduction

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Etidronic acid, (1-Hydroxy-1,1-ethanediyl)bis(phosphonic acid), belongs to the class of

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phosphonate compounds, characterized by the presence of a stable C–P bond. The presence of the carbon atom between the phosphate unit and an organic backbone imparts to the molecule

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a high stability often associated to a low toxicity and to a strong complexing ability. First of all, phosphonates are very useful chelating agents but also very potent inhibitors of mineral

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precipitation and growth. They are used for many applications in industrial, technological,

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* Corresponding author; E-mail: [email protected]

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biological and biomedical, and nanotechnological fields. Just some examples of employments are in cooling waters, desalination systems and in oil fields to inhibit scale formation; in detergents, phosphonates are used as a combination of chelating agent, scale inhibitor and bleach stabilizer; in pulp, paper and textile industry to complex heavy metals in chlorine-free bleaching solutions that could inactivate the peroxide; in medicine phosphonates are used to chelate radionuclides for bone cancer treatments; in the drug delivery and treatment; as antimetabolites; to treat various bone and calcium metabolism diseases [1-7]. The broad spectrum of use makes these compounds subject to numerous studies. Quite recently, the properties and the environmental chemistry of phosphonates was reported by Nowack [8] while a IUPAC Technical Report [9] critically evaluates the available experimental data on stability constants of proton and metal complexes for phosphonic acids. The knowledge of the thermodynamic parameters concerning the metal ligand interactions is essential to simulate the speciation of metal-ligand system in any real condition and, therefore, to predict the best performance conditions of the ligand [10; 11]. Here we report the thermodynamic parameters for the interaction of Etidronic acid (1Hydroxy-1,1-ethanediyl)bis(phosphonic acid) (HEDPA, chart 1) and inorganic and organic mercury (Hg2+ and CH3 Hg+), in NaCl at different ionic strengths. Our attention on mercury(II) is due to the environmental importance of its cations, considered highly toxics to living organisms and ecosystems. It is known that mercury toxicity targets the central nervous system and it has a role in many neurodegenerative disorders such as Alzheimer's and Parkinson's

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diseases. The quantitative study of the interactions between mercury(II) and HEDPA is of utmost importance to evaluate the possible use of this ligand in the metal decontamination

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process of waste effluents, groundwater, and seawater.

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Chart 1. (1-Hydroxy-1,1-ethanediyl)bis(phosphonic acid) - etidronic acid, HEDPA

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2. Experimental

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2.1. Materials

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Etidronic acid (HEDPA) was supplied by Fluka. Its purity was checked by potentiometric titrations and was always >99.5%. Mercury(II) cations were used in the form of chloride salt and were supplied by Riedel-de-Haen (HgCl2) and Stream Chemicals (CH3 HgCl). Sodium chloride aqueous solutions were prepared by weighing the pure salt (Aldrich) previously dried in an oven at T = 383.15 K. Hydrochloric acid and sodium hydroxide solutions were prepared by diluting concentrated ampoules (Fluka). Solutions of HCl were standardized against sodium carbonate and solutions of NaOH against potassium hydrogen phthalate, both previously dried in an oven at T = 383.15 K for 2 hours. Description of chemicals used in this work is reported

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in Table 1. All products were of the highest available purity. Hydroxide solutions were preserved from atmospheric CO2 by means of soda lime traps. All solutions were prepared

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with analytical grade water (ρ = 18 MΩ cm-1) using grade A glassware.

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2.2. Potentiometric equipment and procedure

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Potentiometric measurements were carried out in thermostated cells by means of water circulation in the outer chamber of the titration cell from a thermocryostat (model D1-G Haake) and at T = 298.15 K in NaCl aqueous solutions at different ionic strengths (0.1 ≤ I ≤ 1 mol·L-1) under magnetic stirring and bubbling purified presaturated N2 through the solution to exclude O2 and CO2 inside. In order to reduce systematic errors and to check the repeatability of the systems, two different potentiometric systems were used: i) a Metrohm model 809 Titrando apparatus controlled by Metrohm TiAMO 1.2 software equipped with combined glass electrode (Ross type 8102, from Thermo-Orion); ii) a Metrohm Model 713 potentiometer, equipped with a half-cell glass electrode (Ross type 8101, from Thermo-Orion) and a double-junction reference electrode (type 900200, from Thermo-Orion), and a Model 765 Metrohm motorized burette. The apparatus was connected to a PC, and automatic titrations were performed using a suitable homemade computer program to control titrant delivery, data acquisition and to check for emf stability. Estimated precision was±0.15mV and±0.003 cm3 for the emf and titrant volume readings, respectively, and was the same for both setups. The solutions under study were prepared by mixing different amounts of HEDPA (2 ≤ CL ≤ 6 mmol·L-1), metal cation (2 ≤ CM ≤ 4 mmol·L-1), HCl (5 ≤ CHCl ≤ 10 mmol·L-1) and NaCl in order to establish a prefixed ionic strength value (0.1 ≤ I ≤ 1 mol·L-1). Different metal/ligand ratios were employed (0.5 ≤ CM/CL ≤ 2). Potentiometric measurements were carried out by titrating 25 mL of the solution with standard NaOH solutions, up to pH ~ 10.5. To reach this pH, a volume between 2.5 mL e 4 mL of titrant was added to the initial solution, depending on the concentration of the reagents. For

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each experiment, independent titrations of strong acid solution with standard base were carried out under the same medium and ionic strength conditions as the systems to be investigated,

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with the aim of determining the electrode potential (E0) and the acidic junction potential (Ej =

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ja). In this way, the pH scale used was the free scale, pH ≡ −log [H+], where [H+] is the free proton concentration (not activity). The reliability of the calibration in the alkaline range was

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checked by calculating pKw values. For each titration, 80–100 data points were collected, and the equilibrium state during titrations was checked by adopting the usual precautions, such as

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checking the time required to reach equilibrium and performing back titrations.

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2.3. Calorimetric equipment and procedure

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Calorimetric measurements were carried out at T = 298.150±0.001 K using a Calorimetry Science Corporation (CSC, model 4300) calorimeter. The instrument is an isoperibolic titration calorimeter. It was connected to a PC and automatic titrations were performed using suitable computer programs to control calorimetric data acquisition. The titrations have been carried out with a Hamilton syringe for the titrant delivery, model 1002TLL, having a capacity of 2.5 mL and a glass dewar containing 25 mL of the sample solutions. The solution from syringe was delivered continuously with an injection speed of 0.1 mL/min. Before each experiment the heat of dilution has been measured. The precision of calorimetric system was Q ± 0.015 J and v ±0.001 mL. Its accuracy has been verified titrating a tris(hydroxymethyl)amino-methane (TRIS) buffer solution with standard hydrochloric acid. The enthalpy change obtained for protonation of TRIS (∆H = -46.5±0.5 kJ·mol-1 at T = 298.15 K and I = 0.005 mol·L-1) is close to one reported by Goldberg et al (∆H = -47.55 kJ·mol-1 at T = 298.15 K and I = 0 mol·L-1) [12]. Measurements, for the system Hg2+-HEDPA, were performed using as a titrant, delivered from the syringe, HEDPA salt (CL = 0.2 mol·L-1), as a titrand solution, 25 mL containing Hg2+ (CM = 3 mmol·L-1), HCl (CH = 3 mmol·L-1) and the background salt (NaCl 0.1 mol·L-1). The pH of titrand solution, during the titration, varies from 2.8 to 9.3 (maximum volume of added HEDPA salt, V = 0.6 mL). Titrations, for the system CH3 Hg+-HEDPA, were performed using as a titrant, delivered from the syringe, HEDPA salt (CL = 0.05 mol·L-1) as a titrand solution, 25 mL containing CH3Hg+ (CM = 3 mmol·L-1), HCl (CH = 3 mmol·L-1) and the background salt (NaCl 0.1 mol·L-1). The pH of titrand solution, during the titration, varies from 2.6 to 8.3 (maximum volume of added HEDPA salt, V = 1.5 mL).

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2.4. Calculations

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All the parameters related to the electrode system calibration (formal potential E0 and

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coefficient of the liquid junction potential ja, Ej = ja [H+]), the purity of the reagents and the complex species formation constants were determined using BSTAC and STACO computer

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programs. The enthalpy formation values from calorimetric titration data were obtained by the ES5CM99 program. The dependence of stability constants on the ionic strength was studied

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using the LIANA program. Details on computer programs are reported on ref. [13]. HYSS

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program [14] was employed to calculate the species formation percentages and to draw the distribution diagrams.

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3. Results

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3.1 Etidronic acid - Hg 2+ interactions

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The Hg2+-HEDPA interactions were firstly studied by potentiometric ISE-H+ investigation, in order to define the speciation model and to obtain the complex formation constant values. Experimental data were elaborated by BSTAC program taking into account in all the calculations literature thermodynamic parameters for metal hydrolysis and ligand protonation [15-17], in the same experimental conditions of temperature and ionic strength used in this study (data are resumed in Table 1S of the Supplementary Materials). Several trials were performed on experimental data considering the formation of different species and the best statistical fit [18;19] was selected as the best speciation model, without giving systematic drifts in the residuals. Unknown formation constants were calculated simultaneously, considering at the same time titrations with different metal/ligand ratio. The selected speciation model consists of a simple metal-ligand species (ML2-), two differently protonated species (MLH20 and MLH3+) and a mixed hydrolytic ML(OH)3- one. The agreement between experimental and calculated titration curves is reported as supplementary materials (Fig. 1S). Experimental formation constant values at T = 298.15 K and in NaCl at different ionic strengths are reported in Table 2. The stability of the complexes is high; as an example, for ML2- species log β = 15.61 (at I = 0.18 mol·L-1 and T = 298.15 K). Distribution of different species vs. pH is shown in Figure 1 at two different ionic strengths: I = 0.1 and 1 mol·L-1. As can be observed, at lower ionic strength values (Figure 1a), the diprotonated MLH20 species prevails at pH < 8, while at pH > 8.5 the formation percentage of the hydrolytic M(OH)20 species is high, even in the presence of an excess of ligand. Different is the distribution at I = 1 mol·L-1 (Figure 1b), with

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an increase in the formation percentage of MLOH3- species and, at the same time, the decrease of M(OH)20 percentage. In general, over pH 7, the ionic strength strongly influences the

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distribution.

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3.2 Etidronic acid - CH3Hg + interactions

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Following the same procedure described for Hg2+, potentiometric measurements on CH3 Hg+ (M) - HEDPA (L) system allowed to obtain a speciation model consisting of six species: ML3-,

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MLH2-, MLH2-, MLH30 and two dinuclear M2LH- and M2LH20. Experimental formation constant values are reported in Table 3. The stability of organic mercury(II) complexes is

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lower respect to the inorganic ones. As an example, for MLz- species, logβ = 7.98 and 15.61

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were obtained for CH3 Hg+ and Hg2+ respectively (at I ≈ 0.1 mol·L-1 and T = 298.15 K). Despite this, the distribution of species reported in Figure 2, evidences that the metal cation is present under complex species over all the pH range. Particular relevance assumes the formation of dinuclear species, with ∼40% of metal cation present as M2LH20 between 4 ≤ pH ≤ 5.5 and ∼70% as M2LH- at pH∼8. With the increasing of the ionic strength (Figure 2b), the mononuclear MLHn species (with 0 ≤ n ≤ 3) always prevail.

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3.3 Influence of the ionic strength

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The stability of Hg2+- and CH3Hg+- HEDPA species is strongly influenced by ionic strength

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with formation constant values that may differ up to two orders of magnitude varying from I = 0.1 to I = 1 mol·L-1. Generally, an increase in the stability can be observed by increasing the

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ionic strength, with as an example for HgL2-, log β = 15.61 and 17.21 at I = 0.18 and I = 0.97 mol·L-1, respectively (see Table 2). This is reflected on the distribution of species which may differ significantly with the variation of ionic strength, as outlined in previous paragraphs. This highlights the need to take into account the effect of the ionic strength for a correct evaluation of the species distribution. As reported for other systems [11; 20-22], to have the possibility to simulate the distribution under different conditions from those investigated, the dependence of formation constant values on ionic strength was modeled using the following Debye-Hückel type equation:

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log β = log β 0 - 0.51 z*

I +CI 1 + 1.5 I

(1)

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where z* = Σ (charges)2reactants – Σ (charges)2products, β is the formation constant, β 0 is the

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formation constant at infinite dilution. C is an empirical parameter that depends on the charges involved in the formation reaction. Experimental formation constants of Tables 2 and 3, were

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fitted by equation (1), using the LIANA program to obtain extrapolated logβ0 values, together with the empirical C parameter. Results are reported in Table 4. Known the formation constants at infinite dilution and the C parameter, it is possible to calculate the constants at other ionic strength values.

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Examples of calculated formation constants are listed in Table 2S of the Supplementary Materials, considering some useful ionic strength conditions, like that of a fresh water (I =

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0.001 mol·L-1), blood (I = 0.15 mol·L-1) and sea water (I = 0.7 mol·L-1).

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3.4 Enthalpy changes and dependence on temperature

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With the aim to complete the thermodynamic picture of the systems under study and to fill in the lack of thermodynamic data on HEDPA complex systems, the enthalpies of complex species were determined by calorimetric titrations, through a procedure already used for many other systems [23-26]. The measurements were performed at T = 298.15 K and at I = 0.1 mol·L-1 in NaCl, for both the systems. In the evaluation of enthalpy changes values of complex species from calorimetric measurements, the ∆G of all the species calculated by potentiometric measurements were taken into account and they were kept constant in the calculations. In the experimental conditions of calorimetric measurements, the enthalpy change values of the most abundant complex species were obtained. More in detail, the ∆H of MLH2-, M2LHand M2LH20, for Hg2+-HEDPA system, and of MLH20, ML2- and MLOH3-, for CH3Hg+HEDPA system, were determined. The other species (MLH3+ for Hg2+ and ML3-, MLH2- and MLH30 for CH3Hg+) form in very low percentages in the experimental calorimetric conditions. All the calculated values, together with free energy change and entropy change ones are listed in Table 5. An example of agreement between experimental and calculated values of heat is reported as supplementary materials (Fig. 2S). As can be seen from Table 5, Hg2+ and CH3Hg+ -HEDPA species mainly show exothermic enthalpic values, and the main contribution to the free energy is slightly more entropic than enthalpic. As expected, this trend is intermediate to that which is obtained in the case of softsoft and hard-hard interactions. In fact, if we consider as an example of soft-soft interaction, that between the soft Hg2+ cation and the soft S-donor ligands, namely thiolactate and thiomalate, we have enthalpy change values markedly exothermic and the main contribution to the complexation strongly of enthalpic origin [26]. As an example of hard-hard interaction, that between oligophosphates and Al3+ can be considered [27]. In this case, enthalpy values are endothermic, the contribution to the complexation is mainly due to the entropy and the interaction can be considered of non-covalent type, such as electrostatic one. Correlation analyses of compensatory enthalpy-entropy relationships can be performed [28] and expressed by the following equations: T∆S = α∆H + T∆S0 ∆G = (1-α) ∆∆H The slope of the T∆S vs. ∆H plot shows the extension of the enthalpic gain canceled by the entropic loss that accompanies it. This means that only a fraction (1 - α) of the enthalpy gain contributes to the improvement of complex stability. Moreover, the value of the T∆S0 interception indicates the intrinsic stability at ∆H = 0, i.e. that, if the term T∆S0 is positive, the complex is stabilized even without enthalpy stabilization [28]. By considering altogether the ∆H and T∆S here reported and ones of a previous paper (on interactions of Hg2+ with polyamines, polycarboxylates, and amino acids [29]) and plotting T∆S vs. ∆H, a straight line

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with a fairly good linearity was obtained, with a slope, α = 0.7 and an intercept, T∆S0 = 61 kJ mol-1, and a correlation coefficient r = 0.81. The slope and intercept values can be related to

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the degree of conformational change and to the extent of desolvation upon complexation,

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respectively. More in detail, an α value of 0.7 indicates that 30% of the enthalpic gain or loss caused by system changes affects the complex stability. On the contrary, the value of the intercept of 61 kJ mol-1, shows that the entropic gain which comes from desolvation is very significant. Moreover, plotting T∆S values vs. ∆G ones, for the thermodynamic data reported in this paper, according to Bjerrum [30], ∆G and T∆S are linearly correlated, as can be seen from Fig. 3S of the Supplementary Materials. More in detail, the obtained correlation coefficient is fairly good (R = 0.91). The knowledge of the enthalpy change values, allows the calculation of the constants at different temperatures by using the Van’t Hoff equation. Some calculated values of formation constants referring to Hg2+- and CH3Hg+-HEDPA system at T = 288.15, 310.15, 318.15 K, are reported as Supplementary Materials (Table 3S). Most of the formation constants decrease from 288.15 K to 318.15 K. As an example, for HgL2- species logK = 15.85 and 15.18, at T = 288.15 and 318.15 K, respectively, were obtained, and for CH3HgLH2-, logβ varies from 17.30 (at T = 288.18) to 16.21 (at T = 318.15 K). The only exception was HgLH20 species, for which an increase of stability constant value was observed from 288.15 to 318.15 K.

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4. Discussion

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The knowledge of the thermodynamic parameters is of great importance to evaluate the ability of a ligand “to sequester” a metal cation, in systems which differ in composition, pH, ionic strength, temperature. In particular, recently, a graphical/calculation approach was proposed in ref. [31] and was applied to many systems [10; 11; 21; 32-34] for the calculation of sequestering ability. It consists in the calculation of an empirical parameter named pL0.5 that represents the concentration of a ligand able to sequester half of a given metal cation when it is

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considered as trace (CM = 10-12) in a given system, considering all the possible interactions occurring in the system (metal hydrolysis, protonation of the ligand, interactions with other components). The pL0.5 can be calculated by plotting the mole fraction (χ) of the metal ion M

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complexed by the ligand L as a function of the pL (pL = -log CL, CL = total ligand

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concentration). The obtained function is a sigmoidal curve, with asymptotes equal to 1 for pL → -∞ and equal to 0 for pL → + ∞.

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1

χ=

1 + 10

(pL - pL 0.5 )

(2)

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In Fig. 4, the mole fractions χ of Hg2+ and CH3Hg+ complexed by HEDPA are reported vs pL, considering two different conditions as examples: I = 0.7 mol·L-1, pH = 8.1 and T = 298.15 K 8

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(typical condition of sea water) and I = 0.1 mol·L-1, pH = 7.4 and T = 310.15 K (typical condition of blood). Sequestering ability of HEDPA towards Hg2+ is always higher respect to CH3Hg+. In particular, if we consider the sea water conditions, pL0.5 = 5.5 and 4.4 were obtained for Hg2+ and CH3Hg+, respectively. At T = 310.15 K, sequestering ability towards Hg2+ slightly increases (from 5.5 to 5.9), while the opposite trend can be observed for CH3Hg+ (from 4.4 to 3.6), with differences of almost an order of magnitude. Values are comparable to those of phosphonic derivatives of nitrilotriacetic acid [10], for which 5.5 ≤ pL0.5 ≤ 6.0 for Hg2+ and 2.9 ≤ pL0.5 ≤ 3.1 for CH3Hg+ were obtained.

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5. Conclusion

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The extensive data reported here and derived from both potentiometry and calorimetry in a wide ionic strength range give a complete thermodynamic picture of the HEDPA-Hg2+ and

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HEDPA-CH3Hg+ interactions, at T = 298.15 K. To our knowledge, no data are reported in the literature on these systems. The results can be summarized as follows.

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-

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species. More complex is the speciation model for CH3Hg+ system, with the formations of four mononuclear (ML3-, MLH2-, MLH2-, MLH30) and two dinuclear (M2LH- and M2LH20)

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Different speciation models for HEDPA-Hg2+ and HEDPA-CH3Hg+ systems were obtained. For Hg2+, it consists of four mononuclear ML2-, MLH20, MLH3+ and ML(OH)3-

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species, whose stability is lower respect to the inorganic ones. The knowledge of the interaction thermodynamic parameters and their dependence on ionic strength and temperature allowed us to evaluate the sequestering ability on the basis of empirical parameter pL0.5 that takes into account all the possible interactions occurring in the system. HEDPA showed higher sequestering ability towards Hg2+ respect to CH3Hg+, in both experimental conditions chosen as examples (i.e. sea water and blood conditions), but the trend in the two different conditions is different for the two metal cations: towards Hg2+, sequestering ability of HEDPA is higher in sea water condition, while for CH3Hg+, is higher in blood condition. This highlights that sequestering ability can be influenced by experimental conditions and, therefore, the knowledge of the complete thermodynamic framework allows to predict the behavior and to evaluate the optimum conditions in which a chelating agent can be used to remove a given metal cation.

Acknowledgements We thank MIUR (Ministero dell'Istruzione, dell'Università e della Ricerca) for financial support (cofunded PRIN project with Prot. 2015MP34H3) and FSE regional funds for PhD support to D.C.

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[14] [15] [16] [17] [18] [19] [20] [21]

[22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32]

A. Clearfield, K. Demadis, Metal Phosphonate Chemistry: From Synthesis To Applications, Royal Society of Chemistry, 2012. J.M.H. de Klerk, v.D. A., A.D. van het Schip, B.A. Zonnenberg, P.P. van Rijk, J. Nucl. Med. 33 (1992) 646-651. J. Galezowskaa, E. Gumienna-Kontecka, Coord. Chem. Rev. 256 (2012) 105–124. D.E. Lorke, A. Stegmeier-Petroianu, G.A. Petroianu, J. Appl. Toxicol. 37 (2017) 13-22. S.J. Popwell, M.D. Schulz, K.B. Wagener, C.D. Batich, R.J. Milner, J. Lagmay, W.E. Bolch, Cancer Biotherapy Radiopharm. 29 (2014) 273-282. P.J. Thornton, H. Kadri, A. Miccoli, Y. Mehellou, J Med Chem. 59 (2016) 10400-10410. R. Lange, R. Heine, R. Knapp, J.M.H. de Klerk, H.J. Bloemendal, H.N. Hendrikse, Bone 91 (2016) 159-179. B. Nowack, Water Res. 37 (2003) 2533-2546. K. Popov, H. Ronkkomaki, L.H.J. Lajunen, Pure Appl. Chem. 73 (2001) 1641-1677. C. De Stefano, C. Foti, O. Giuffrè, D. Milea, New J. Chem. 40 (2016) 1443-1453. P. Cardiano, C. Foti, O. Giuffrè, J. Mol. Liq. 240 (2017) 128–137. R.N. Goldberg, N. Kishore, R.M. Lennen, Pure Appl. Chem. 73 (2001) 1641-1677. C. De Stefano, S. Sammartano, P. Mineo, C. Rigano, Computer Tools for the Speciation of Natural Fluids. in: A. Gianguzza, E. Pelizzetti, S. Sammartano, (Eds.), Marine Chemistry - An Environmental Analytical Chemistry Approach, Kluwer Academic Publishers, Amsterdam, 1997, pp. 71-83. L. Alderighi , P. Gans, A. Ienco, D. Peters, A. Sabatini, A. Vacca, Coord. Chem. Rev. 184 (1999) 311-318. A. De Robertis, C. Foti, G. Patanè, S. Sammartano, J. Chem. Eng. Data 43 (1998) 957-960. C. Foti, O. Giuffrè, S. Sammartano, J. Chem. Thermodynamics 66 (2013) 151-160. K.J. Powell, P.L. Brown, R.H. Byrne, T. Gajda, G. Hefter, S. Sjoberg, H. Wanner, Pure & Appl. Chem. 77 (2005) 739-800. A. Vacca, A. Sabatini, M.A. Gristina, Coord. Chem. Rev. 8 (1972) 45-53. M. Filella, P.M. May, Talanta 65 (2005) 1221-1225. D. Chillè, C. Foti, O. Giuffrè, Chemosphere 190 (2018) 72-79. F. Crea, C. De Stefano, C. Foti, G. Lando, D. Milea, S. Sammartano, Alkali-Metal Ion Complexes with Phosphates, Nucleotides, Amino Acids, and Related Ligands of Biological Relevance. Their Properties in Solution. in: A. Sigel, H. Sigel, R.K.O. Sigel, (Eds.), The Alkali Metal Ions: Their Role for Life, Springer International Publishing AG, Cham, Switzerland, 2016, pp. 1-628. F. Crea, G. Falcone, C. Foti, O. Giuffrè, S. Materazzi, New J. Chem. 38 (2014) 3973-3983. G. Falcone, C. Foti, A. Gianguzza, O. Giuffrè, A. Napoli, A. Pettignano, D. Piazzese, Anal. Bioanal. Chem. 405 (2013) 881-893. C. De Stefano, A. Gianguzza, O. Giuffrè, A. Pettignano, S. Sammartano, Appl. Organomet. Chem. 22 (2008) 30-38. F. Crea, P. Crea, C. De Stefano, O. Giuffrè, A. Pettignano, S. Sammartano, J. Chem. Eng. Data 49 (2004) 658-663. P. Cardiano, D. Cucinotta, C. Foti, O. Giuffrè, S. Sammartano, J. Chem. Eng. Data 56 (2011) 1995-2004. D. Aiello, P. Cardiano, R.M. Cigala, P. Gans, F. Giacobello, O. Giuffrè, A. Napoli, S. Sammartano, J. Chem. Eng. Data 62 (2017) 3981-3990. M.V. Rekharsky, Y. Inoue, Chem. Rev. 98 (1998) 1875-1917. C. Foti, O. Giuffrè, G. Lando, S. Sammartano, J. Chem. Eng. Data 54 (2009) 893–903. E.J. King, Acid-Base Equilibria, Pergamon, New York, 1965. F. Crea, C. De Stefano, C. Foti, D. Milea, S. Sammartano, Curr. Med. Chem. 21 (2014) 38193836. P. Cardiano, C. Foti, O. Giuffrè, J. Mol. Liq. 223 (2016) 360-367.

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[33] P. Cardiano, C. Foti, P. Mineo, M. Galletta, F. Risitano, S. Lo Schiavo, J. Mol. Liq. 223 (2016) 174–181. [34] G. Falcone, C. Foti, S. Sammartano, J. Chem. Eng. Data 57 (2012) 3636−3643.

11

1 2 3 4

Table 1 Description of chemicals used in this work Chemical

Source

Purification

Purity

Method

(%wt.)

Purity Check Potentiometric

Etidronic acid

Fluka

None

> 99.5%

Mercury(II) Chloride

Riedel-de-Haen

None

> 99.5%

-

Methylmercury(II) Chloride

Stream Chemicals

None

> 99

-

Sodium Chloride

Sigma-Aldrich

None

≥ 99.5%

-

Sodium Hydroxide

Sigma-Aldrich

None

≥ 99%

Hydrochloric Acid

Sigma-Aldrich

None

≥ 99%

Potassium Phatalate Monobasic

Sigma-Aldrich

None

≥ 99.5%

-

Sodium Carbonate

Sigma-Aldrich

None

≥ 99.5%

-

titration

Volumetric titration Potentiometric titration

5 6 7

12

1 2

Table 2

3 4

Potentiometric experimental equilibrium constants for the formation of the Hg2+ (M) - HEDPA (L) species, in NaCl aqueous solution at different ionic strengths and at T = 298.15 K (p = 0.1

5

MPa)

6

logβ β

Reaction I /mol·L-1

7

0.18

0.25

0.48

0.97

M2+ + L4- = ML2-

15.61±0.01 a)

15.80±0.08 a)

16.01±0.02 a)

17.21±0.09a)

M2+ + L4- + 2H+ = MLH20 M2+ + L4- + 3H+ = MLH3+ M2+ + L4- + H2O = ML(OH)3-

31.48±0.02 34.64±0.03 6.47±0.01

31.79±0.07 35.06±0.05 6.65±0.04

32.44±0.06 35.70±0.04 6.73±0.04

33.14±0.08 37.26±0.07 7.72±0.08

a)

Standard uncertainties u(logβ); u(T) = 0.1 K; u(I) = 0.01

8 9

13

1 2

Table 3

3

Potentiometric experimental equilibrium constants for the formation of the CH3Hg+ (M) HEDPA (L) species, in NaCl at different ionic strengths and at T = 298.15 K (p = 0.1 MPa)

4 5

logβ β

Reaction I /mol·L-1

6

0.11

0.23

0.40

0.89

M+ + L4- = ML3M+ + L4- + H+ = MLH2-

7.98±0.04 a) 16.91±0.04

7.95±0.03 a) 17.15±0.06

8.30±0.02 a) 17.26±0.07

8.90±0.02 a) 17.95±0.10

M+ + L4- + 2H+ = MLH2M+ + L4- + 3H+ = MLH30

23.98±0.04 26.74±0.04

23.96±0.04 26.53±0.05

24.07±0.05 26.65±0.09

24.47±.0.10 27.02±0.06

2M+ + L4- + H+ = M2LH2M+ + L4- +2H+ = M2LH20

24.33±0.03 30.70±0.03

24.36±0.03 30.37±0.04

24.47±0.02 30.53±0.02

25.74±0.02 31.99±0.04

a)

Standard uncertainties u(logβ); u(T) = 0.1 K; u(I) = 0.01

7 8

14

1 2 3 4

Table 4 Calculated formation constant values at infinite dilution for Hg2+ and CH3Hg+ -HEDPA (L) species, together with C parameter for the dependence on ionic strength, in NaCl at T = 298.15 K (p = 0.1 MPa)

5

logβ β0

C

ML2-

17.17±0.09a)

3.4±0.1b)

MLH20 MLH3+

33.66±0.12 36.94±0.15

5.1±0.2 5.0±0.2

7.36±0.07

2.4±0.1

ML3MLH2-

8.65±0.04 18.38±0.12

2.1±0.1 2.7±0.2

MLH2MLH30

25.88±0.11 28.79±0.07

2.5±0.2 2.5±0.2

M2LHM2LH20

25.96±0.08 32.34±0.11

3.8±0.2 4.0±0.2

M

Species

Hg2+

ML(OH)3CH3Hg+

6 7

a)

Standard uncertainties u(logβ); b) standard uncertainties u(C).

8

15

1 2

Table 5

3

Thermodynamic formation parameters for the Hg2+ and CH3Hg+ -HEDPA (L) species, in NaCl at I = 0.10 mol·L-1 and T = 298.15 K (p = 0.1 MPa)

4 5

∆H a)

-∆ ∆G a)

T∆ ∆S a)

-39±4b)

89.1

50

M + L + 2H M + L4- + H2O

16.5±1 -41±7

179.6 36.9

196.1 -4

M + L4- + H+

-64±5

96.5

32

-35±2 -30±1

138.8 175.4

104 145

M

Reaction

Hg2+

M + L44-

CH3Hg+

4-

+

+

2M + L + H 2M + L4- +2H+ 6

a)

In kJ· mol-1; b) standard uncertainties u(∆H); u(T) = 0.001 K; u(I) = 0.01.

7 8

16

100

MLH2

80

MH-2

60 ML 40 MLH3

20

MLH-1

% formation relative to M

% formation relative to M

100

80 MLH3

MLH2 ML

60

40 MLH-1 20 MH-2

0

0 3

1 2

5

7 pH

(a)

9

3

5

7

9

pH

(b)

3 4 5 6

Fig. 1. Calculated distribution diagrams of Hg2+-HEDPA species vs pH (a) CM = 1 mmol·L-1, CL = 2 mmol·L-1, I = 0.1 mol·L-1 (NaCl), T = 298.15 K (b) CM = 1 mmol·L-1, CL = 2 mmol·L-1, I = 1 mol·L-1 (NaCl), T = 298.15 K

7 8

17

1

80

80

M2LH

ML

60 MLH2 M2LH2

40

20

MLH

MLH3

MLH2

ML MLH

60

40 MLH3 M2LH

20 M2LH2

0

0

3

2 3

% formation relative to M

100

% formation relative to M

100

5

7 pH

(a)

9

3

5

7

9

pH

(b)

4 5 6 7 8

Fig. 2. Calculated distribution diagrams of CH3Hg+-HEDPA species vs pH (a) CM = 1 mmol· L-1, CL = 2 mmol· L-1, I = 0.1 mol· L-1 (NaCl), T = 298.15 K (b) CM = 1 mmol· L-1, CL = 2 mmol· L-1, I = 1 mol· L-1 (NaCl), T = 298.15 K

18

150

Τ∆S /kJ mol

-1

100 50 0 -50 -100 -200

-150

-100

-50

0

50

-1

1

∆H /kJ mol

2

Fig. 3. T∆S (kJ·mol-1) vs. ∆H (kJ·mol-1) for Hg2+ and CH3Hg+-HEDPA and Hg2+-

3

carboxylates, amines and amino acids species.

4 5

19

1 2

χ

1.0

2+

Hg

+

CH3Hg

0.5

0.0 0

2

4

6

8

pL 3 4 5 6

Fig. 4. Mole fraction of Hg2+ and CH3Hg+ complexed by HEDPA vs. pL (−) I = 0.7 mol·L-1, pH = 8.1 and T = 298.15 K (--) I = 0.1 mol·L-1, pH = 7.4 and T = 310.15 K

7 8 9 10

20

1

Supplementary Materials

2 4

Table 1S Literature values of Hg2+ and CH3Hg+ hydrolysis and of HEDPA (L) protonation, in NaCl at

5

different ionic strengths and at T = 298.15 K

3

6

logβ β

Reaction I /mol·L-1

Ref.

0.1

0.25

0.5

1

-3.60

-3.63

-3.62

-3.56

Hg + 2 H2O = Hg(OH)2 + 2 H Hg2+ + 3 H2O = Hg(OH)3- + 3 H+

-6.34 -21.10

-6.36 -21.19

-6.33 -21.29

-6.20 -21.48

2 Hg2+ + H2O = Hg2(OH)3+ + H+

-3.58

-3.10

-3.09

-3.145

Hg2+ + Cl- = HgCl+ Hg2+ + 2 Cl- = HgCl20

6.82 13.36

6.77 13.30

6.78 13.31

6.92 13.50

Hg2+ + 3 Cl- = HgCl3Hg2+ + 4 Cl- = HgCl42-

14.44 15.06

14.36 15.02

14.35 15.08

14.50 15.32

Hg2+ + Cl- + H2O = HgCl(OH)0 + H+

3.68

3.64

3.68

3.88

CH3Hg+ + H2O = CH3(OH)0 + H+

-4.538

-4.556

-4.582

-4.639

[15]

H+ + L4- = HL32H+ + L4- = H2L2-

10.50 17.22

10.07 16.62

9.81 16.25

9.58 15.92

[16]

3H+ + L4- = H3L14H+ + L4- = H4L0

19.77 21.07

19.02 20.29

18.53 19.78

18.05 19.30

Hg2+ + H2O = Hg(OH)+ + H+ 2+

0

+

[17]

[17]

7 8 9

21

2

Table 2S Calculated equilibrium constants for the formation of the Hg2+ and CH3Hg+ (M) - HEDPA (L)

3

species, in NaCl at different ionic strengths and at T = 298.15 K

1

4

M

log β a)

Species I /mol·L-1

Hg2+

CH3Hg+

5 6

a)

0.001

0.15

0.7

ML2-

16.93±0.08 b)

15.68±0.07 b)

16.51±0.02b)

MLH20 MLH3+

33.33±0.12 36.61±0.15

31.68±0.10 34.94±0.12

33.08±0.07 36.25±0.08

ML(OH)3-

7.22±0.09

6.49±0.07

7.19±0.05

ML3MLH2-

8.53±0.04 18.17±0.12

7.97±0.03 17.04±0.09

8.60±0.06 17.61±0.08

MLH2MLH30

25.61±0.11 28.48±0.07

24.01±0.09 26.67±0.05

24.23±0.08 26.78±0.08

M2LHM2LH20

25.68±0.08 32.04±0.11

24.27±0.07 30.44±0.09

25.19±0.07 31.36±0.09

Calculated by eq. (1) and values of Table 4;

b)

standard uncertainties u(logβ).

7

22

2

Table 3S. Calculated values of formation constants of Hg2+- and CH3Hg+-HEDPA species at different

3

temperatures

1

4

M

Hg2+

CH3Hg+

Species

logβ T = 288.18 K

T = 310.15 K

T = 318.15 K

ML2MLH20

15.85 31.38

15.35 31.59

15.18 31.66

MLOH3-

6.72

6.19

6.02

MLH2M2LH-

17.30 24.54

16.48 24.09

16.21 23.94

M2LH20

30.88

30.50

30.37

5 6 7 8

23

400

E /mV

300

200

100

0 0.0

0.5

1.0

1.5

2.0

2.5

mL (NaOH)

1 2 3 4 5 6 7

Fig. 1S. Experimental (symbols) and calculated (continuous line) potentiometric titration curves of Hg2+-HEDPA solution (CHg = 2 mmol·L-1, CHEDPA = 2 mmol·L-1, I = 0.12 mol·L-1, T = 298.15 K)

24

0.4

Q /J

0.2

0.0

-0.2

-0.4 2

4

6

8

10

pH 1 2 3 4 5

Fig. 2S. Experimental (□) and calculated (•) values of heat vs. pH for the determination of enthalpy values of Hg2+-HEDPA species.

25

-1

T∆S /kJ mol

160

80

0 50

100

150

200

-1

-∆G /kJ mol 1 2 3

Fig. 3S. T∆S vs. ∆G for Hg2+- and CH3Hg+- HEDPA interactions at T = 298.15 K and I = 0.1

4

mol·L-1.

5

26

1 2 3 4 5 6 7 8

• For the first time, thermodynamic parameters on etidronic acid-Hg2+ and -CH3Hg + interactions are determined • Dependence on ionic strength and on temperature is defined by potentiometric and calorimetric tecniques • Sequestering ability of etidronic ligand towards Hg2+ and CH3Hg+ is evaluated on the basis of speciation models •

Effect of the experimental conditions on sequestering ability is considered

27