Ligands Toward Palladium(II) - American Chemical Society

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Sequestering Ability of Aminopolycarboxylic (APCs) and Aminopolyphosphonic (APPs) Ligands Toward Palladium(II) in Aqueous Solution Concetta De Stefano,† Gabriele Lando,*,† Alberto Pettignano,‡ and Silvio Sammartano† †

Dipartimento di Scienze Chimiche, Università degli Studi di Messina, Viale Ferdinando Stagno d’Alcontres, 31, I-98166 Messina (Vill. S. Agata), Italy ‡ Dipartimento di Fisica e Chimica, Università di Palermo, Viale delle Scienze, I-90128 Palermo, Italy ABSTRACT: The binding capacity of three aminopolycarboxylates [nitrilotriacetic acid (NTA), ethylene-glycol-bis(2-aminoethyl ether)-N,N,N′,N′-tetraacetic acid (EGTA), and diethylenetriamine-N,N,N′,N″,N″-pentaacetic acid (DTPA)] and two aminopolyphosphonates {(1-hydroxyethane-1,1-diyl)bis(phosphonic acid) (HEDP) and [[(phosphonomethyl)imino]bis[2,1ethanediylnitrilobis(methylene)]] tetrakis-phosphonic acid (DTPP)} toward palladium(II) ion was studied by potentiometric and spectrophotometric titrations at different temperatures (283.15 ≤ T/K ≤ 318.15) and ionic strengths (0.1 ≤ I/mol·dm−3 ≤ 1.0) in NaClO4. The hydrolysis of Pd2+ and the protonation of ligands were always taken into account in the speciation models of Pd2+/L systems investigated. Equilibrium reaching experiments were performed to check and confirm the reaching of the equilibrium state. Owing to the high stability of the PdL species (K > 1020), for EGTA, HEDP, and DTPP it was determined using exchange measurements with auxiliary ligands, such as iodide (I−) and ammonia (NH3). For the other ligands the stability of the PdL species was reported in the literature. The general speciation scheme consisted of mononuclear differently protonated species with general formula PdHiL and only in the case of the NTA ligand the formation of the PdL2 species was found. The stability of the PdL species is high: as an example we have log KML = 17.82, 22.60, 36.31, 23.49, and 27.27 for NTA, EGTA, DTPA, HEDP, and DTPP, respectively. Among the ligands, DTPA shows the highest formation constants and sequestering ability, evaluated using the pL0.5 parameter, as well. The complex formation reaction is always exothermic and in general the entropic contribution to the stability is dominant. Some empirical relationships were found to model the dependence of the formation constants on the number of protons and of the sequestering ability on pH.

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INTRODUCTION Among the six platinum group elements (PGE) palladium is the most abundant and constitutes about the 5·10−7 % of the Earth’s crust. It is usually present in the environment in the elemental form or in alloys with iron and nickel and more rarely as mineral like potarite and stibiopalladinite. More than 80 % of the world production of palladium is concentrated in South Africa and in Russia.1,2 Among platinum group elements palladium is the most reactive, above all when it is present as fine powder. The worldwide request of palladium has been increased in the last decades because of its use in different fields. Palladium finds application as catalyst in several redox reactions, together with platinum and rhodium in catalytic converters, in electronics industry (alone or in alloys with silver, gold, platinum, iridium, nickel, or copper) for the production of switches for low voltage contact or relays, in dental implants as copper, gold or silver alloys, in jewelry, etc.1−6 The progressive increase of the amount of palladium used in these fields produces a corresponding increase of its concentration in the environment, especially in urban areas. © 2014 American Chemical Society

With respect to the other PGE, palladium is quickly oxidized in soils2 and this causes an increase of its mobility and also its availability for humans, animals and plants. The remediation of polluted areas as well as the recovery of the metal ion from the environment are of great concern and the use of ligands with high binding ability toward the metal ion could be an adequate response to the problem. In the last years our research group carried out a systematic study on the binding ability of several ligand classes having different kind of binding groups toward Pd2+ ion and some stability data on palladium(II)−ligand complexes have already been published.7,8 Among these ligands, five aminopolycarboxylates (APCs) having different number of amino and carboxylic groups were taken into account. The interest for this ligand class arises from their strong binding ability toward metal ions and for this reason these compounds find interesting applications in different cases in which the sequestration/ Received: January 30, 2014 Accepted: April 24, 2014 Published: May 9, 2014 1970

dx.doi.org/10.1021/je500109m | J. Chem. Eng. Data 2014, 59, 1970−1983

Journal of Chemical & Engineering Data

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Figure 1. Structure of ligands studied in this work.

removal of some metal ions is requested.9−11 Moreover APCs have low or no toxicity toward humans, animals, and plants and are environmental friendly compounds.12,13 APCs are used in environmental field as sequestering agents of metal ions for the remediation of soils, sediments, or natural waters, in medicine as chelating agents for heavy metals and radioactive substances in chelation therapy, in industry for the treatment of wastewaters, and in some cases in industrial processes.11,14−17 Here an improvement of the previous investigation is proposed, taking into account other APCs and also two chelating agents with amino and phosphonic groups. As APCs, also the aminopolyphosphonates (APPs) ligands have a zwitterionic structure like that of betaine, but the amino groups of APPs show an higher basicity. Moreover, phosphonic groups have an high inductive effect and a great nucleophilicity; the stereochemistry of the molecules suggests an higher stability of their complexes with metal ions.11,17 In particular our attention was focused on three APCs [nitrilotriacetic acid (NTA), ethylene-glycol-bis(2-aminoethyl ether)-N,N,N′,N′-tetraacetic acid (EGTA), and diethylenetriamine-N,N,N′,N″,N″-pentaacetic acid (DTPA)] and two APPs {(1-hydroxyethane-1,1-diyl)bis(phosphonic acid) (HEDP) and [[(phosphonomethyl)imino]bis[2,1-ethanediylnitrilobis(methylene)]] tetrakis-phosphonic acid (DTPP); structures are shown in Figure 1}. Due to the great complexing ability of APC and APP ligands, in the last decades some reviews have been published.18−21 As stated by Anderegg et al.,19 the determination of the complex formation constants between Pd2+ and complexones is always difficult because the very high stability of the ML species (K > 1020) hampers the possibility of the use of direct potentiometric or spectrophotometric titrations.19 For this

reason, some potentiometric exchange measurements with auxiliary ligands were carried out. For this purpose, iodide (I−) and ammonia (NH3) were used and the stability of the various PdI k or Pd(NH 3 ) k species has been taken from the literature.7,22 The speciation study of the palladium/ligand systems was made also by spectrophotometric technique in NaClO4 medium, at different ionic strengths, in the temperature range 283.15 ≤T/K ≤ 318.15 and was preceded by the potentiometric determination of the ligand protonation constants in NaNO3 medium. The ionic strength and the temperature dependence of the formation constants was studied and some parameters for the modeling have been proposed, although it was not possible to obtain reliable thermodynamic parameters such as ΔH or TΔS. Finally the data have been compared with literature findings, in particular with the data reported by De Stefano et al.7 for the interaction of Pd2+ cation with ethylenediamine-N,N,N′,N′-tetraacetate (EDTA), (S,S)-ethylenediamine-N,N′-sicuccinate (S,S-EDDS), and triethylenetetramine-N,N,N′,N″,N‴,N‴-hexaaacetate (TTHA). Some empirical relationships were found to model the stability of the metal/ ligand complexes as a function of the number of the binding sites.

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MATERIALS AND METHODS Chemicals. Palladium(II) nitrate dihydrate (by SigmaAldrich) stock solutions were prepared by dissolving the weighed salt in 0.1 mol·dm−3 nitric or perchloric acid. Palladium concentration was checked by ICP-AES technique against Pd(NO3)2 standard solutions (Fluka) or gravimetrically using the dimethylglyoxime method.7,8 NTA, EGTA, DTPA, HEDP, and DTPP were supplied by Fluka with analytical grade 1971



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Table 1. Experimental Conditions of Potentiometric and Spectrophotometric Titrations ligand NTA EGTA DTPA HEDP DTPP

cPda 1.06 0.49 0.91 0.15 0.36

to to to to to

1.36 0.61 1.05 0.41 0.50

NTA EGTA DTPA HEDP DTPP

EGTA HEDP DTPP a

cLa

cH a

1.82 to 2.11 0.84 to 1.0 2.03 to 2.12 0.5 to 1.0 0.56 to 0.80 3.50 0.83 3.56 1.47 1.46

to to to to to

3.78 0.95 4.27 4.99 2.52

cPd:cL

I/mol·dm−3

Spectrophotometric Titrations 87.38to 88.28 1:2 0.10 85.30 to 85.94 1:1.5 to 1:2 0.10 92.10 to 92.56 1:2 0.10 2.0 to 4.0 1:1 to 1:4 0.10 50.60 1:1 to 1:2 0.10 to 1.0 Potentiometric Titrations 10.50 to 11.33 0.10 3.31 to 3.80 0.10 17.78 to 21.33 0.10 5.88 to 19.95 0.10 13.18 to 22.71 0.10 to 1.0 Potentiometric Ligand Exchange Measurements

T/K

pH

no. titrat.

283.15 283.15 283.15 283.15 298.15

to to to to

318.15 318.15 318.15 318.15

0.98to 11.70 0.99 to 11.31 0.94 to 11.50 1.17 to 11.70 1.36 to 11.38

15 15 15 15 12

283.15 283.15 283.15 288.15 298.15

to to to to

318.15 318.15 318.15 318.15

2.0 2.0 2.0 2.0 2.0

12 12 12 16 15

to to to to to

10.0 9.5 10.5 10.5 12.0

cPda

cLa

cNaIa

cNH4NO3

I/mol·dm−3

T/K

pH

no. titrat.

0.5 to 1.0 0.5 to 1.0 0.5 to 1.0

1.5 to 3.0 2.0 to 3.0 1.0 to 3.0

0.05 0.05

0.05 0.025 to 0.075 0.002 to 0.006

0.10 0.10 0.10 to 1.00

298.15 298.15 298.15

2.0 to 10.0b 2.0 to 10.0b 2.0 to 8.5

6 6 12

In mmol·dm−3. bFor NaI medium the pH range is 5.0 to 10.0, for NH4NO3 the pH range is 2.0 to 8.5.

purity. NaI, NH4NO3, NaNO3, and NaClO4 solutions were prepared by weighing the pure salts (Fluka) after drying in an oven at 383.15 K. Nitric acid, perchloric acid and sodium hydroxide solutions were prepared by diluting concentrated ampules (Riedel−deHaën) and were standardized against sodium carbonate and potassium hydrogen phthalate, respectively. NaOH solutions were preserved from atmospheric CO2 by means of soda lime traps. All solutions were prepared with analytical grade water (R = 18 MΩ) using grade A glassware. Potentiometric Apparatus and Procedure. Protonation constants of the ligands were determined by potentiometry in the temperature range 283.15 ≤ T/K ≤ 318.15. Ligand exchange measurements were carried out at T = 298.15 K and different ionic strengths in mixed NaClO4/NaI and NaClO4/ NH4NO3 media. Both kind of measurements were done using an 809 Metrohm Titrando apparatus equipped with a combined Orion glass electrode Ross type 8102. The apparatus was connected to a PC, and automatic titrations were performed using the Metrohm TiAMO 1.2 software to check for emf stability and to control titrant delivery and data acquisition. The estimated accuracy was ± 0.2 mV and ± 0.003 cm3 for emf and titrant volume readings, respectively. For each titration 25 cm3 of titrand solution containing known amounts of the ligand under study and sodium nitrate, in order to reach the prefixed ionic strength values, was titrated with standard NaOH. All titrations were carried out under magnetic stirring and presaturated N2 was bubbled through the solution in order to exclude O2 and CO2 inside. Details of the experimental measurements are reported in Table 1. Eighty to hundred points were collected for each titration, and the equilibrium state during titrations was checked by monitoring the time necessary to reach equilibrium (several titrations were carried out with different equilibration periods ranging between 10 and 60 s per data point) and by performing back-titrations to check and confirm the reversibility of the reactions.23 For each experiment, independent titrations of strong acidic solutions with standard base were carried out under the same temperature and ionic strength conditions as in the systems to be investigated, to determine the standard electrode potential.

The ligand exchange measurements with iodide were performed as follows: a suitable amount of ligand was dissolved in a volumetric flask, then standard NaOH(aq) was added to reach pH ∼ 8.0 and then Pd(NO3)2 was weighed ( 8.5 the formation of free ammonia (NH3 (g)) would lead to big errors in the evaluation of both volume and emf. Before the measurements some experiments were performed to ensure the reaching of the equilibrium in the following way: (i) suitable amounts of Pd(NO3)2 (∼ 1 mmol·dm−3), ligand (∼ 3 mmol·dm−3), and NaClO4 (∼ 0.1 mol·dm−3) were dissolved in deionized water; (ii) a reference solution was prepared with NaClO4 (∼ 0.1 mol·dm−3) and HNO3 (∼ 5 mmol·dm−3); (iii) the solutions were putted in a thermostated room at T = 298.15 K; (iv) the pH of the solutions was measured immediately and after 72 h. Spectrophotometric Apparatus and Procedure. The UV−Vis spectra were recorded in the wavelength range 250 to 550 nm, using a Varian Cary 50 UV − Vis spectrophotometer with an optic fiber probe having a fixed 1 cm path length. The spectrophotometer was connected to a PC for the acquisition of the spectra. A volume of 25 cm3 of the solutions containing the ligand (L = NTA, EGTA, DTPA, HEDP, or DTPP), palladium(II) ion, perchloric acid, and sodium perchlorate, in order to reach the prefixed ionic strength values, was titrated with a standard NaOH solution. The pH of titrand solution was monitored with a combined Orion glass electrode Ross type 1972



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involved in the equilibrium and the ions of the supporting electrolytes. The conversion from molar to molal concentration scale was performed using appropriate density values.36 For the generic reaction reported in eq 1, Δε in NaClO4 is

8102 connected to a model 713 Metrohm potentiometer and titrant solution was added by a model 765 Metrohm motorized buret. Estimated accuracy was ± 0.2 mV and ± 0.003 cm3 for emf and titrant volume readings, respectively. NaClO4 was chosen as ionic medium in order to avoid the sorption of NO3− in the wavelength range investigated and the formation of Pd2+ species with Cl− or NO3− that could complicate the speciation picture of the systems investigated. Forty to fifty points were collected for each titration and the equilibrium state during titrations was checked by adopting the same precautions used in potentiometric measurements. Details of experimental measurements are reported in Table 1. Calculations. The nonlinear least-squares computer program ESAB2M25 was used for the refinement of all the parameters of the acid−base titration, such as the standard electrode potential (E0), the ionic product of water (Kw), the acidic junction potential (Ej = ja·[H+]), and the analytical concentration of reagents. For details on the nature of the above parameters and on the refinement procedure see Braibanti et al.23 The BSTAC, STACO26 and HYPERQUAD 200627 computer programs were used in the calculation of protonation constants of ligands from potentiometric titrations. UV−Vis spectra were analyzed by the HYPSPEC28 program, which allows to calculate stability constants and molar absorbance of each absorbing species, using as input the experimental values of absorbance, analytical concentrations of reagents and the chemical model proposed. The ES4ECI29 program was used to draw speciation diagrams and to calculate species formation percentages. The LIANA30 computer program was used to fit different linear and nonlinear functions. Overall and stepwise protonation constants of ligands, hydrolysis constants of palladium(II) ion and formation constants of Pd2+−L complex species are given according to the eqs 1 and 2: j Pd2 + + i H+ + k Lz − = Pd jHiLk(kz − 2j − i) −

βik

(1)

j Pd2 + + HiLk(kz − i) − = Pd jHiLk(kz − 2j − i) −

K ik

(2)

Δεjik = jε(Pd2 +, ClO4 −) + kε(Lz −, Na +) + iε(H+, ClO4 −) − ε(Pd jHLk(kz − 2j

0.51 I + CI 1 + 1.5 I

)−

, Na +/ClO4 −)

(4)

when a neutral species is formed ε is replaced by km, the Setschenow coefficient,37 which is related to the activity coefficient of the neutral species by log γ = k mmNaClO4

(5)

Usually, the km of a neutral species is experimentally determined by solubility or distribution measurements.38−40 In this work, no data for the determination of km have been collected, therefore only Δε values were provided. The dependence of the complex formation constants on temperature has also been taken into account. The amount of data collected in this work is not sufficient to derive reliable values of enthalpy changes using the van’t Hoff equation. Furthermore, as well known the determination of the enthalpy change of a process should be done using appropriate calorimetric titration techniques.41,42 Usually the order of magnitude of the errors associated with the enthalpy values determined with this technique is ± 1 kJ·mol−1. The calculation of enthalpy values with potentiometric titrations at different temperatures comes from the use of a drivative equation, such as the van’t Hoff is and, as reported by Hepler,43 the derivation process always enhance the errors associated with the computed quantity. For these reasons only a temperature gradient (∂logKik/∂T) has been calculated according to the following equation log K ik = log K ikθ +

∂log K ik · (T − θ ) ∂T

(6)

where θ is the reference temperature (298.15 K in our case). From the evaluation of the temperature gradients it is possible to define if a process is exothermic or endothermic and rough values of enthalpy changes have been proposed.

where Lz‑ = NTA3−, EGTA4−, DTPA5−, HEDP4−, or DTPP9−. If j = 1 eqs 1 and 2 refer to complex formation constants, and if j = 0, eqs 1 and 2 refer to the ligand protonation constants and are indicated as βHik and KHik . When k = 0 and i < 0, eqs 1 and 2 refer to the metal hydrolysis constants. Two different approaches were used to evaluate the dependence on ionic strength of protonation constants of ligands and formation constants of their complexes with Pd2+ ion: (i) a Debye−Hückel type equation and (ii) the specific ion interaction theory (SIT) approach. The following Debye−Hückel type equation was used log K ik = log K ik0 − z*

−

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RESULTS AND DISCUSSION Ligand Protonation Constants and Metal Hydrolysis Constants. The study of the metal−ligand systems has been preceded by the determination of the protonation constants of the ligands. For some ligands reliable data are present in the literature, for example in the case of the NTA Bretti et al.44 reported data at T = 298.15 K at different ionic strengths (0 < I/mol·dm−3 ≤ 5) in NaCl. Though these data are reported in NaCl, they can be used in NaClO4 in a narrow ionic strength range. The data reported by Daniele et al.45 were used for the temperature dependence of the NTA protonation constants. For this latter ligand some measurements were performed in NaCl at different ionic strengths and in all cases the results were not significantly different in the experimental error range at 95 % of the C.I. Also the data reported in the most common databases for equilibrium constants22,46,47 are in accordance with the data reported in the cited papers. For DTPA the data reported by Bretti et al.44 and Martell et al.22 were used. For EGTA, results reported by Bretti et al.48 at T = 298.15 K were considered. In all cases some random check of the protonation

(3)

z* = Σ (charges) reactants − Σ (charges)2products, Kik (or the formation constant, K0ik (or β0ik) is the formation 2

where βik) is constant at infinite dilution, and C is an empirical parameter. In this model formation constants, concentrations and ionic strengths are expressed in the molar (mol·dm−3) concentration scale. If both formation constants and ionic strength are expressed in the molal concentration scale (mol·kg−1), eq 3 became the SIT equation,31−35 where C is replaced by Δε, which is the summation of the specific interaction coefficients of the species 1973



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Table 2. Experimental and Literature Protonation Constantsa of the Ligands at Different Temperatures and Ionic Strengths in NaNO3 ligand NTA

EGTA

DTPA

HEDP

DTPP

a

I/mol·dm−3 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.106 0.500 0.745 1.009

T/K 283.15 298.15 307.15 310.15 318.15 283.15 298.15 307.15 310.15 318.15 283.15 298.15 307.15 310.15 318.15 288.15 298.15 307.15 310.15 318.15 298.15 298.15 298.15 298.15

log βH2

log KH1 b

9.60 9.42b 9.32b 9.29b 9.21b 9.62 ± 0.03c 9.36d 9.21 ± 0.01 9.16 ± 0.01 9.04 ± 0.03 10.43e 10.13e 9.9e 9.90e 9.76e 10.89 ± 0.02c 10.77 ± 0.01 10.67 ± 0.02 10.63 ± 0.02 10.54 ± 0.03 12.11 ± 0.02 11.64 ± 0.02 11.48 ± 0.02 11.22 ± 0.02

log βH3

b

12.11 11.94b 11.85b 11.82b 11.74b 18.53 ± 18.08d 17.83 ± 17.76 ± 17.55 ± 19.10e 18.64e 18.38e 18.30e 18.09e 17.68 ± 17.54 ± 17.42 ± 17.38 ± 17.28 ± 23.44 ± 22.88 ± 22.77 ± 22.50 ±

log βH4

log βH5

log βH6

log βH7

b

0.03c 0.03 0.04 0.02

0.03c 0.02 0.02 0.02 0.04 0.05 0.03 0.03 0.04

13.90 13.73b 13.67b 13.64b 13.6b 21.18 ± 20.78d 20.55 ± 20.48 ± 20.30 ± 23.36e 22.84e 22.55e 22.46e 22.22e 20.20 ± 20.03 ± 19.88 ± 19.83 ± 19.70 ± 32.18 ± 31.40 ± 31.26 ± 30.90 ±

0.04c 0.03 0.04 0.02

0.03c 0.02 0.03 0.03 0.04 0.06 0.03 0.06 0.06

23.08 ± 22.71d 22.50 ± 22.44 ± 22.27 ± 26.07e 25.54e 25.24e 25.15e 24.91e

39.46 38.48 38.30 37.84

± ± ± ±

0.04c 0.03 0.05 0.02

0.04 0.04 0.08 0.08

24.29 ± 23.96d 23.78 ± 23.72 ± 23.57 ± 28.19e 27.68e 27.39e 27.30e 27.07e

45.68 44.48 44.24 43.67

± ± ± ±

0.05c 0.05 0.08 0.03 29.30e 28.79e 28.50e 28.41e 28.18e

0.05 0.05 0.08 0.08

51.01 49.60 49.27 48.58

± ± ± ±

0.06 0.06 0.08 0.09

54.39 52.74 52.31 51.49

± ± ± ±

0.06 0.07 0.10 0.10

Refers to eq 1. bCalculated from refs 44 and 45. c95 % C.I. dCalculated from ref 48. eCalculated from refs 22 and 44.

Table 3. Complex Formation Constants of the Pd2+/L Systems in NaClO4 or Mixed NaClO4/NaI and NaClO4/NH4NO3 Media in the Molar Concentration Scale log βa ligand NTA

EGTA

EDTA S,S-EDDS DTPA

TTHA HEDP

DTPP

a

T/K

I/mol·dm−3

283.15 298.15 307.15 310.15 318.15 283.15 298.15 307.15 310.15 318.15 298.15 298.15 283.15 298.15 307.15 310.15 318.15 298.15 288.15 298.15 307.15 310.15 318.15 298.15 298.15 298.15 298.15

0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.106 0.500 0.750 1.009

PdL 17.91 ± 17.82d 17.75 ± 17.73 ± 17.68 ± 22.80 ± 22.49 ± 22.27 ± 22.21 ± 22.04 ± 23.60d 23.07d 36.68 ± 36.31d 36.07 ± 36.02 ± 35.82 ± 37.0d 23.50 ± 23.49 ± 23.34 ± 23.31 ± 23.24 ± 27.27 ± 25.78 ± 25.93 ± 26.17 ±

0.04b 0.02 0.03 0.05 0.09 0.08 0.12 0.08 0.10

0.012 0.12 0.13 0.14 0.06 0.06 0.06 0.08 0.10 0.20 0.10 0.12 0.10

PdHL 20.17 ± 20.10d 19.94 ± 19.94 ± 19.86 ± 25.86 ± 25.70 ± 25.55 ± 25.52 ± 25.42 ± 26.63d 26.73d 41.08 ± 40.40d 40.25 ± 39.85 ± 39.52 ± 43.95d 27.11 ± 26.77 ± 26.71 ± 26.61 ± 26.45 ± 35.72 ± 33.59 ± 33.19 ± 33.14 ±

0.04b 0.03 0.05 0.07 0.08 0.10 0.12 0.12 0.15

0.10 0.10 0.12 0.14 0.10 0.08 0.06 0.10 0.15 0.10 0.10 0.11 0.11

PdH2L 25.89 ± 24.50c,d 24.48 ± 24.14 ± 23.64 ± 28.30 ± 28.26 ± 28.01 ± 28.02 ± 27.93 ± 28.40d 28.62d 44.20 ± 43.20d 43.17 ± 42.98 ± 42.63 ± 47.64d 29.89 ± 29.39 ± 29.16 ± 29.02 ± 28.71 ± 42.79 ± 40.28 ± 40.47 ± 40.84 ±

PdH3L

PdH4L

PdH5L

PdLOH

± ± ± ±

11.99 ± 0.05b 10.20d 9.85 ± 0.07 9.44 ± 0.12 8.69 ± 0.17 13.14 ± 0.10 14.06 ± 0.10 14.10 ± 0.10 14.32 ± 0.15 14.66 ± 0.20 13.50d 12.00d 25.15 ± 0.10 24.7d 24.23 ± 0.10 24.15 ± 0.10 24.06 ± 0.10 27.00d 13.71 ± 0.09 13.73 ± 0.09 13.48 ± 0.12 13.48 ± 0.15 13.38 ± 0.18 10.01 ± 0.10 9.55 ± 0.15 9.58 ± 0.12 9.63 ± 0.13

0.10b,c 0.06c 0.13c 0.17c 0.12 0.07 0.09 0.10 0.13

0.12 0.08 0.16 0.14 0.15 0.09 0.08 0.16 0.15 0.10 0.14 0.14 0.21

46.30 ± 45.42d 44.65 ± 44.62 ± 44.11 ± 50.00d

0.13

± ± ± ±

0.10 0.10 0.19 0.21

48.24 46.21 45.91 46.32

0.12 0.16 0.16

52.25 50.68 50.11 50.62

± ± ± ±

0.10 0.10 0.12 0.15

55.13 53.84 53.82 54.84

0.10 0.09 0.26 0.22

Refers to eq 2. b95 % C. I. cAs Pd(NTA)2 species. dTaken from De Stefano et al.7

1974



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be comparable to that of the DTPP. The stability constants of the PdIk species at I = 0.1 mol·kg−1 were taken from De Stefano et al.,7 which recalculated the data of Elding et al.51 The stability of the Pd(NH3)k species was taken from Martell et al.22 Whereas the use of iodide as exchange ligand is reported in the literature, to demonstrate that ammonia is suitable for this purpose, in Figure 2 the trend of the apparent formation

constant values were performed and results consistent with literature findings were obtained. Some data were also reported in the literature for HEDP by Foti et al.,42 where the authors also determined the protonation enthalpies at different ionic strengths in different ionic media. Since it was shown that the medium effect for HEDP is more important, the protonation constants of HEDP have been experimentally determined. In this context it is important to underline that, in terms of formation of weak complexes, the nature of the cation of the supporting electrolyte (e.g., Na+ or K+) is extremely important for the protonation constant values of APPs whereas it is not the same for those of the APCs.49 For EGTA, HEDP and DTPP, the protonation constants were experimentally determined at the same experimental conditions of Pd2+ − L systems investigated. The protonation constant of ammonia at different ionic strengths was taken from Martell et al.22 and is 9.26 at I = 0.1 mol·dm−3, 9.33 at I = 0.5 mol·dm−3 and 9.43 at I = 1.0 mol· dm−3. All the protonation constants, experimentally determined, used in this work are reported in Table 2. As regards the metal cation, the data reported in Baes and Mesmer50 are considered reliable and for the Pd2+ cation, the most important hydrolytic species are the Pd(OH)+ and the Pd(OH)20. Stability of the Palladium/Ligand Complexes. Once determined the ligand protonation constants in different conditions, the Pd2+/ligand systems were studied at different temperatures (283.15 ≤ T/K ≤ 318.15) and ionic strengths (0.1 ≤ I/mol·dm−3 ≤ 1.0) in NaClO4, mixed NaClO4/NaI and NaClO4/NH4NO3 ionic media. In these conditions, the interaction of Pd2+ with NTA, EGTA, DTPA, HEDP, and DTPP, were studied by means of spectrophotometric and potentiometric titrations in the experimental conditions summarized in Table 1. As can be noted, in some cases the ligand concentration is very low, due to the solubility of the molecule, especially in acidic pH conditions. The solubility of DTPA, EGTA, and NTA has been studied elsewhere.48 The equilibrium reaching experiments confirmed that the equilibrium between Pd2+ and the ligands is reached immediately, in fact the emf/pH of the solutions, containing both Pd2+ and the ligands, does not vary significantly ( 8.0. For this reason, the concentration of ammonia necessary to establish the exchange equilibria has to

Figure 2. Values of the apparent formation constants (according to Swarzenbach) for different ligands toward Pd2+ at I = 0.1 mol·dm−3 and T = 298.15 K. 1, I−; 2, NH3; 3, HEDP; 4, EGTA; 5, DTPP.

constants (according to Schwarzenbach) with pH is reported for the Pd2+/EGTA, Pd2+/HEDP, Pd2+/DTPP, Pd2+/NH3, and Pd2+/I− systems. It can be observed that in the range 6 < pH < 8, there is a crossover point between the curve relative to the ammonia and those relative to the other ligands, indicating that using appropriate concentrations an exchange between ammonia and complexone can be established. In fact, looking also at Figure 3, where the formation percentage of the sum of the PdHi(HEDP) and Pd(NH3)k species is plotted as a function of pH, it can be observed that at pH ∼ 7 there is the exchange between HEDP and ammonia. Similar considerations can be done for iodide, also considering EGTA and DTPP ligands.

Figure 3. Distribution diagram of the Pd2+/HEDP and Pd2+/NH3 species in NaClO4/NH4NO3 at I = 0.1 mol·dm−3 and T = 298.15 K with cPd = 0.001 mol·dm−3, cHEDP = 0.003 mol·dm−3, and cNH3 = 0.050 mol·dm−3. Curves: 1, sum of the PdHi(HEDP) species; 2, sum of the Pd(NH3)k species. 1975



Article

of the Pd2+ cation; the species should then be indicated with the formalism PdH−1L. In addition, the formation of species with many protons indicates that the interaction between Pd2+ and the ligands is very strong in the entire pH range even at very low pH values. The stability constants of the various PdHiLk species are reported in Table 3 at different ionic strengths (in NaClO4) and temperatures. For all of the systems, the dependence on temperature of the formation constants was studied, whereas for the Pd2+/DTPP system the ionic strength dependence has been investigated. From the analysis of the data in Table 3, it is evident that the formation constant values increase with increasing the number of carboxylic groups (nCOO‑) and amine groups (nN) of the APC ligands. As an example, the values of the PdL species is (at T = 298.15 K) log K = 17.82, 22.49, 23.60, 23.07, 36.31, and 37.0 for NTA (nCOO‑ = 3 and nN = 1), EGTA (nCOO‑ = 4 and nN = 2), EDTA (nCOO‑ = 4 and nN = 2), S,S-EDDS (nCOO‑ = 3 and nN = 1), DTPA (nCOO‑ = 5 and nN = 3), and TTHA (nCOO‑ = 6 and nN = 4), respectively. It should be noted that for S,S−EDDS, the amino groups are secondary and not tertiary as for the other ligands. The data of this work confirm the relationship that was obtained by De Stefano et al.7 for the prediction of the stability of the Pd2+/ APCs complexes, based on the number of amine (nN) and carboxylic (nCOO‑) binding sites and on the distance (number of −CH2− spacers) between the amino groups of the molecule (d).

With these kinds of measurements, the stability of the PdL species was determined at I = 0.1 mol·dm−3 and T = 298.15 K for EGTA and HEDP, and at I = 0.1, 0.5, 0.75, and 1.0 mol· dm−3 and T = 298.15 K for DTPP. The spectrophotometric data analysis performed by HypSpec computer program, allowed us to determine different mononuclear species with general formula PdHiLk. Analyzing both direct titrations and back-titrations, the standard deviations of the fits were comparable in all cases. The measurements were analyzed individually and then all together, in order to have an idea of the reproducibility and in all cases a good agreement was found. For the selection of the speciation model an analysis of variance was performed and the model that gave the best results in terms of statistical parameters and reproducibility was selected. In the case of the Pd2+/NTA system it was not possible to refine the Pd(NTA)24− species (see below) in a single measurement, but only analyzing all the titrations together. The results obtained at I = 0.1 mol·dm−3 with spectrophotometric and potentiometric measurements were comparable in most cases, for example the stability of the Pd(EGTA)2− species resulted not significantly different, log KML = 22.49 ± 0.04 and 22.43 ± 0.02 for spectrophotometric (spec) and potentiometric (pot) measurements, respectively. For HEDP, log KML = 23.39 ± 0.03 (spec) and 23.49 ± 0.03 (pot) was obtained and the value of 23.49 was selected. For DTPP log KML = 26.72 ± 0.07 (spec) and 27.27 ± 0.20 (pot). In this case, although associated with higher errors, the value obtained with potentiometry was adopted and kept constant in the analysis of the spectrophotometric measurements for the determination of the other protonated complex species. At the other ionic strengths the differences between the two techniques are higher. For example, at I = 0.5 mol·dm−3, log KML = 22.83 ± 0.02 (spec) and 26.03 ± 0.10 (pot) and at I = 1.0 mol·dm−3 log KML = 22.09 ± 0.33 (spec) and 26.17 ± 0.14 (pot). In these cases, the stability of the PdL species was taken from the potentiometric measurements and that of the other species was determined using the spectrophotometric data. Considering that for EGTA and HEDP, at T = 298.15 K, the stability of the PdL species resulted substantially identical to that obtained with spectrophotometry, the results obtained with this latter technique at the other temperatures were considered reliable. In some cases, the competitive method with auxiliary ligand is comparable with that of simple competitive proton displacement for the complexation of the ligand. Of course the ligands with high protonation constant values are more suitable and in this context the APPs are better than the APCs ones. The general speciation scheme consisted of three species, determined for all the systems, namely the PdL, PdHL and Pd(OH)L species. In addition to these, for the Pd2+/NTA system the PdL2 species was determined; for the Pd2+/EGTA and Pd2+/HEDP the PdH2L, for the Pd2+/DTPA also the PdH2L and PdH3L species were found, whereas the speciation scheme of the Pd2+/DTPP system consisted of seven species, namely PdL, PdHL, PdH2L, PdH3L, PdH4L, and PdH5L and the mixed hydrolytic Pd(OH)L; whose formation can be attributed to the strong hydrolysis of the Pd2+ cation. In fact, even if the formation of both the simple hydrolytic Pd(OH)+ and Pd(OH)20 species is inhibited, probably the PdL complex species can undergoes hydrolysis. Another way to interpret the formation of the hydrolytic species is the displacement of a proton of the ligand through the binding

log KPd(APC) ± 0.6 = 12.8 + 0.8nN 2 + 1.7nCOO − + 0.7d

Please note that in the paper of De Stefano et al.7 there is a misprint, because the parameter for the variable d is indicated as negative (−0.7), whereas it is positive (+0.7). Although the two considered APP ligands, HEDP and DTPP are characterized by a different number of binding sites, the stability of the PdL species is similar (log K = 23.5 and 27.3 at I ∼ 0.1 mol·dm−3 and T = 298.15 K). On the contrary, the stability of the protonated species is significantly higher for DTPP than for HEDP. It was found that the values of the formation constants of the PdHi(DTPP) species, expressed in eq 2, are a function of the number of the proton of the species. In Figure 4 for example this trend is shown at T = 298.15 K and I = 0.106 mol·dm−3; the fitting equation is as follows (95 % C.I.): (log KPdHiL ± 0.2) = 27.1 − 3.6i

This trend was observed in the past for other ligands.52,53 If all of the stability constants reported in Table 3 are considered, the same kind of analysis can be done with common slope, s = −3.8 ± 0.3 at T = 298.15 K and I = 0.1 mol·dm−3. In this last case, the quality of the fit worsens significantly; for example for the Pd2+/HEDP system we have that (95 % C.I.) log KPdHiL ± 1.0 = 21.2 − 3.8i − 0.4i 2

Although it is not a very precise predictive equation, it can be regarded as a general indication to calculate tentative formation constants for other similar systems. For the other systems the intercept are a = 20.6, 35.5, 16.4, 22.3, 21.4, and 37.6 for EGTA, DTPA, NTA, EDTA, S,S-EDDS, and TTHA, respectively. The addition of the quadratic term was necessary for the ligands different than DTPP to account for the curvature. 1976



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Figure 6. Distribution diagram of the PdHi(EGTA) species in NaClO4 at I = 0.1 mol·dm−3 and T = 283.15 K with cPd = 0.001 mol·dm−3 and cL = 0.003 mol·dm−3. Species: 1, PdH2(EGTA); 2, PdH(EGTA)−; 3, Pd(EGTA)2−; 4, Pd(OH)(EGTA)3−.

Figure 4. Dependence of the stability constants of the Pd2+/DTPP (log K (PdHiL)) system vs the number of proton of the complexes (i) at T = 298.15 K and I = 0.106 mol·dm−3.

shows molar fraction of 0.9−1.0 from pH ∼ 4.5 to pH ∼ 8. For pH > 8, Pd2+ is mostly bind to EGTA in the form of the mixed hydrolytic species Pd(OH)(EGTA)3−. For the Pd2+/DTPA system, the distribution diagram of the species is reported in Figure 7 under conditions similar to the

To better appreciate the speciation of the studied systems, some distribution diagrams have been reported for some important conditions, where the molar fraction of the various species are plotted vs pH. In Figures 5 to 10 the different distribution diagrams for the PdHiLk species are reported for NTA, EGTA, DTPA, HEDP

Figure 7. Distribution diagram of the PdHi(DTPA) species in NaClO4 at I = 0.1 mol·dm−3 and T = 310.15 K with cPd = 0.001 mol·dm−3 and cL = 0.003 mol·dm−3. Species: 1, PdH3(DTPA); 2, PdH2(DTPA)−; 3, PdH(DTPA)2−; 4, Pd(DTPA)3−; 5, Pd(OH)(DTPA)4−.

Figure 5. Distribution diagram of the PdHi(NTA) species in NaClO4 at I = 0.1 mol·dm−3 and T = 298.15 K with cPd = 0.001 mol·dm−3 and cL = 0.003 mol·dm−3. Species: 1, PdH(NTA); 2, Pd(NTA)−; 3, Pd(NTA)24−; 4, Pd(OH)(NTA)2−.

physiological ones, at T = 310.15 K and I = 0.1 mol·dm−3 in Na+ medium. The most important species is the Pd(DTPA)3−, which covers the entire pH range of the natural fluids conditions in the range 4 < pH < 9. In this case, the hydrolytic species Pd(OH)(DTPA)4− is present in very small amount (∼0.1 molar fraction) at pH > 10.0. The three protonated species are present at pH < 4.0 and each of them reaches 0.5 molar fraction. The distribution diagram of the PdHi(HEDP) species is shown in Figure 8 at I = 0.1 mol·dm−3 and T = 298.15 K and it is similar to those of the APCs. At pH ∼ 2 the predominant species is the PdH2(HEDP) and at pH ∼ 3 the PdH(HEDP)− species is the most important with 0.5 molar fraction. From pH ∼ 4 to pH ∼ 9, the natural fluids pH window, the dominating species is the Pd(HEDP)2− and the hydrolytic Pd(OH)HEDP3− is formed only at pH > 8.5.

and DTPP, respectively in the same conditions of concentrations, cPd = 0.001 mol·dm−3 and cL = 0.003 mol·dm−3 and pH (2 to 10). The distribution diagram of the PdHi(NTA)k species at T = 298.15 K and I = 0.1 mol·dm−3, reported in Figure 5, shows that at pH ∼ 2 the only species present is the PdH(NTA). At pH ∼ 4, the Pd(NTA)− is the dominating species, whereas the Pd(NTA)24− is the most important at pH ∼ 6.5. The Pd(OH)(NTA)2− is formed at pH ∼ 6, but the formation percentage becomes significant (0.2−0.3 molar fraction) for pH > 8.5. In Figure 6 the distribution diagram of the PdHi(EGTA) species is reported at T = 283.15 K and I = 0.1 mol·dm−3. Between pH 2 and pH 3 the two protonated species are the most important and the sum of them reaches the totality of the Pd2+ present in solution. At pH > 3, the Pd(EGTA)2− species 1977



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Figure 10. Distribution diagram of the PdHi(DTPP) species in NaClO4 at I = 1.009 mol·dm−3 and T = 298.15 K with cPd = 0.001 mol· dm−3 and cL = 0.003 mol·dm−3. Species: 1, PdH5(DTPP)2−; 2, PdH 4 (DTPP) 3− ; 3, PdH 3 (DTPP) 4− ; 4, PdH 2 (DTPP) 5− ; 5, PdH(DTPP)6−; 6, Pd(DTPP)7−.

Figure 8. Distribution diagram of the PdHi(HEDP) species in NaClO4 at I = 0.1 mol·dm−3 and T = 298.15 K with cPd = 0.001 mol·dm−3 and cL = 0.003 mol·dm−3. Species: 1, PdH2(HEDP); 2, PdH(HEDP)−; 3, Pd(HEDP)2−; 4, Pd(OH)(HEDP)3−.

Regarding the species distribution, for DTPP the effect of the variation of the ionic strength was also studied. In Figures 9 and

Figure 11. Ionic strength dependence of the Pd(DTPP)7− formation constant at T = 298.15 K in NaClO4. log KML vs I/mol·kg−1. Figure 9. Distribution diagram of the PdHi(DTPP) species in NaClO4 at I = 0.106 mol·dm−3 and T = 298.15 K with cPd = 0.001 mol·dm−3 and c L = 0.003 mol·dm −3 . Species: 1, PdH 5 (DTPP) 2− ; 2, PdH 4 (DTPP) 3− ; 3, PdH 3 (DTPP) 4− ; 4, PdH 2 (DTPP) 5− ; 5, PdH(DTPP)6−; 6, Pd(DTPP)7−.

In this paper, some measurements were performed at different ionic strengths and temperatures, to study the influence of the variation of these parameters on the protonation and on the complex formation constants using eqs 3 and 6. The ionic strength dependence parameters of the complex formation constants are reported in Table 4. As pointed out earlier, the amount of data in this work is not sufficient to derive reliable thermodynamic quantities, therefore in Table 5 some temperature gradients of the complex formation constants have been proposed, together with some rough values of enthalpy changes (ΔH), and entropic contributions (TΔS), calculated using well known thermodynamic relationships. Although these thermodynamic data cannot be considered reliable, their analysis allow us to affirm that all the reactions are exothermic (except for the Pd(OH)(EGTA)3− species). Moreover, looking at the data in Table 5 it can be observed that in general the complex formation process is mainly driven by the entropic contribution. The complex formation constants reported in Table 5 are calculated from eq 6 and for this reason are slightly different than the experimental ones, reported in Table 3. Often it is hard to make comparisons between the binding ability of different ligands, or classes of them, toward a given

10, the distribution diagram of the PdHi(DTPP) species is reported at I = 0.106 mol·dm−3 and I = 1.009 mol·dm−3. It is interesting to note that, at I = 1.009 mol·dm−3, the formation pH of the species is shifted to more alkaline conditions. For example the maximum formation pH of the PdH2(DTPP)5− species is pH ∼ 6 at I = 0.106 mol·dm−3 and pH ∼ 6.5 at I = 1.009 mol·dm−3. The formation of the mixed Pd(OH)(DTPP)8− and the simple Pd(OH)20 hydrolytic species is observed only at pH > 11. From a fast analysis of the data in Table 3, it is evident that the formation constant values decrease with increasing ionic strength (see Figure 11 for the trend of the Pd(DTPP)7− species) and temperature. Some exceptions to this trend are the Pd(OH)(EGTA)3− and the PdH2(EGTA) (expressed in eq 2) species. The temperature effect is observed to different extents for different systems. As an example, for the Pd(DTPA)3− species there is a lowering of ∼0.8 log K units in a range of 35 K degrees, whereas for the Pd(NTA)− species a lowering of ∼0.2 log K units in the same temperature range. 1978



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the analysis of this Table, it is evident that generally the pL0.5 values decrease with increasing pH, except for DTPP at all ionic strengths and EGTA, only at T = 318.15 K. The DTPA ligand shows the highest sequestering ability in all the considered experimental conditions, a very slight and not significant decrease of the pL0.5 has been observed increasing the pH. An interesting comparison can be done between EGTA and HEDP. In fact, the pL0.5 values for the Pd2+/EGTA system are higher than those of the Pd2+/HEDP at high pH values (e.g., pH 9), whereas the contrary is observed at pH = 4. This aspect underlines that the protonated species (formed at low pH) of phosphonates are usually stronger than those of carboxylates, as also seen for the protonation constants of the ligands. Although with some slight discrepancies, due to the different temperature dependence of the protonation and complex formation constants, the effect of the variation of the temperature and the ionic strength on the pL0.5 values traces that of the formation constants. In Figure 12 the sequestration curves of the various Pd2+/L systems are reported at pH 9, I = 0.1 mol·dm−3 and T = 298.15 K. It is evident that NTA has the lowest sequestering ability among the four ligands; the curves of the HEDP and DTPP are almost overlapped (pL0.5 of DTPP is slightly higher) and therefore show similar values of pL0.5, whereas EGTA is characterized by the highest value of pL0.5 among these four ligands. The curve relative to the Pd2+/DTPA system is not reported in the figure because the values are too much higher than those of the other ligands. In Figure 13, the dependence of pL0.5 on pH is shown for the Pd2+/HEDP system at T = 298.15 K and the decrease of the pL0.5 is evident, although the curves at pH 4, 5, and 6 are very close to each other. As stated above, the pL0.5 values are strongly dependent on the experimental conditions in which they are calculated and looking at Table 6 many empirical relationship can be found between pL0.5 and pH, ionic strength or temperature or the number of binding sites of the ligand.

Table 4. Ionic Strength Dependence Parameters for the Pd2+/DTPP System equilibrium Pd+(DTPP) =Pd(DTPP) Pd+H+(DTPP) =PdH(DTPP) Pd+2H+(DTPP) =PdH2(DTPP) Pd+3H+(DTPP) =PdH3(DTPP) Pd+4H+(DTPP) =PdH4(DTPP) Pd+5H+(DTPP) =PdH5(DTPP) Pd+(DTPP)+H2O = Pd(OH)(DTPP)+H a

log K0ika 30.99 ± 0.08

Δεa

Ca b

2.49 ± 0.12

2.44 ± 0.11b

b

41.11 ± 0.08

2.25 ± 0.11

2.22 ± 0.11

49.22 ± 0.09

4.22 ± 0.14

4.11 ± 0.12

55.86 ± 0.14

5.18 ± 0.20

5.04 ± 0.19

60.79 ± 0.20

6.25 ± 0.29

6.07 ± 0.28

64.08 ± 0.19

8.36 ± 0.26

8.10 ± 0.25

12.13 ± 0.09

1.66 ± 0.15

1.63 ± 0.15

Parameter of eqs 3 and 4. b95 % C.I.

metal cation. In fact, the values of the complex formation constants is not sufficient for this purpose, due to different side reactions that occur in solution, as the protonation of the ligand. In this work the pL0.5 (or pL50) parameter has been used. This is an empirical parameter that relates the concentration of the ligand with the percentage of the complexed metal cation. The calculation of this parameter and its symbolism is similar to the dose−response curves (e.g., LD50), but for a wider description of the pL0.5 many papers have been published.54−57 Briefly, the molar fraction (χ) of the sum of the complex species is plotted with respect to the pL or −log[L], where L is the ligand. Calculating χ for different concentrations of the ligand different data points are obtained and then fitted with a Boltzmann type equation, and the pL0.5 is defined as the concentration of the ligand necessary to sequester the 50 % of the concentration of the metal, present in trace concentration. In this work, the pL0.5 values were calculated for the different Pd2+/L systems at different pH, temperatures, and ionic strengths and all the values are summarized in Table 6. From

Table 5. Temperature Dependence Parameters of the Complex Formation Constants of the Pd2+/L Species, at T = 298.15 K and I = 0.1 mol·kg−1 in NaClO4 or Mixed NaClO4/NaI and NaClO4/NH4NO3 Media L NTA

EGTA

DTPA

HEDP

a

species

log Kika

PdL PdHL Pd2L Pd(OH)L PdL PdHL PdH2L Pd(OH)L PdL PdHL PdH2L PdH3L Pd(OH)L PdL PdHL PdH2L Pd(OH)L

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

17.81 20.05 24.89 10.55 22.47 25.67 28.15 13.81 36.31 40.42 43.50 45.36 24.64 23.41 26.86 29.48 13.62

0.01 0.02 0.05 0.05 0.01 0.01 0.02 0.03 0.01 0.03 0.05 0.02 0.03 0.01 0.02 0.02 0.02

∂ log K/∂Ta c

−0.0066 ± 0.0005 −0.0091 ± 0.0014 −0.0627 ± 0.0031 −0.0931 ± 0.0027 −0.0218 ± 0.0005 −0.0127 ± 0.0005 −0.0110 ± 0.0012 0.0425 ± 0.0017 −0.0246 ± 0.0006 −0.0442 ± 0.0020 −0.0435 ± 0.0029 −0.0627 ± 0.0014 −0.0342 ± 0.0020 −0.0086 ± 0.0006 −0.0209 ± 0.0013 −0.0387 ± 0.0011 −0.0120 ± 0.0013

c

−ΔGb

ΔHb

TΔSb

101.7 114.4 142.1 60.2 128.3 146.5 160.7 78.8 207.3 230.7 248.3 258.9 140.6 133.6 153.3 168.3 77.8

−11.2 −15.5 −106.7 −158.4 −37.1 −21.5 −18.8 72.3 −41.9 −75.2 −74.0 −106.8 −58.1 −14.6 −35.6 −65.8 −20.4

90.4 98.9 35.4 −98.2 91.2 125.0 141.9 151.1 165.4 155.5 174.3 152.1 82.5 119.1 117.8 102.5 57.3

Parameter of eq 6. bIn kJ·mol−1. c95 % C.I. 1979



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Table 6. pL0.5 Values for the Pd2+/L Systems at Different pH, Ionic Strengths, and Temperatures in NaClO4 or Mixed NaClO4/ NaI and NaClO4/NH4NO3 Media pH system

T/K

I/mol·dm−3

4.0

5.0

6.0

7.4

8.1

9.0

pL0.5 ± 0.1a Pd2+/NTA

Pd2+/EGTA

Pd2+/DTPA

Pd2+/HEDP

Pd2+/DTPP

a

283.15 298.15 307.15 310.15 318.15 283.15 298.15 307.15 310.15 318.15 283.15 298.15 307.15 310.15 318.15 288.15 298.15 307.15 310.15 318.15 298.15 298.15 298.15 298.15

0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.100 0.1060 0.5004 0.7450 0.1009

9.09 9.18 9.20 9.21 9.24 9.09 9.24 9.30 9.30 9.35 22.49 22.39 22.49 22.35 22.33 10.75 10.72 10.78 10.78 10.80 6.24 5.95 5.88 7.25

8.16 8.20 8.23 8.24 8.27 9.07 9.21 9.26 9.26 9.30 22.41 22.45 22.49 22.49 22.49 10.63 10.65 10.73 10.73 10.76 7.01 6.29 6.31 7.35

7.45 7.21 7.24 7.24 7.27 9.07 9.21 9.25 9.25 9.31 22.38 22.47 22.49 22.51 22.52 10.56 10.58 10.66 10.66 10.69 7.78 6.57 6.78 7.67

7.20 6.02 5.98 5.91 5.89 9.06 9.23 9.34 9.34 9.57 22.36 22.44 22.45 22.48 22.49 9.91 9.93 9.98 9.99 10.02 8.27 6.99 7.04 7.64

7.18 5.75 5.65 5.48 5.32 9.02 9.28 9.53 9.53 9.92 22.28 22.33 22.32 22.35 22.33 9.32 9.33 9.38 9.39 9.41 8.38 7.29 7.45 8.00

7.09 5.54 5.40 5.16 4.89 8.75 9.31 9.67 9.67 10.08 21.84 21.80 21.74 21.74 21.68 8.46 8.49 8.51 8.52 8.54 8.54 7.60 7.85 8.37

95 % C.I.

Figure 12. Sequestration diagram of the Pd2+/L systems at pH 9, I = 0.1 mol·kg−1 and T = 298.15 K in NaClO4. Molar fraction of complexed Pd2+ vs pL, −log (cL/mol·kg−1). Curves: 1, NTA; 2, HEDP; 3, DTPP; 4, EGTA.

Figure 13. Sequestration diagram of the Pd2+/HEDP systems at T = 298.15 K and I = 0.1 mol·kg−1 in NaClO4. Molar fraction of complexed Pd2+ vs pL, −log (cL/mol·kg−1). Curves: 1, pH 9; 2, pH 8.1; 3, pH 7.4; 4, pH 6; 5, pH 4.

For example, the pL0.5 values of HEDP were modeled as a function of both temperature and pH by means of the following equation:

already discussed that this work represent an improvement of a previous one,7 reporting data only at T = 298.15 K, on the sequestering ability of the APCs toward Pd2+. As stated by De Stefano et al.,7 the very high stability of the Pd2+ complexes hampers the use of the classical potentiometric technique in the determination of the equilibrium constants. For APCs, discrepancies were evidenced with some literature findings reported in refs.21,58−62 These differences were ascribed to the different experimental techniques and experimental conditions used by the authors for the determination of the equilibrium constants, which in addition do not meet criteria for the critical

pL0.5 ± 0.1 = 7.98 + (1.18 − 0.13pH)pH + 0.003(T − 298.15)

This equation does not have physical meaning and has to be regarded by its empirical nature. Some Pd2+/L systems studied in this work have been studied in the literature in different experimental conditions. It was 1980



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selection in the most common stability constants database. Nevertheless, a satisfactory accordance was found with some literature data, as in the case of log KPdNTA = 17.1 (NaClO4, I = 1.0 mol·dm−3, T = 293.15 K),21 comparable with data in this work (log KPdNTA = 17.82 NaClO4, I = 0.1 mol·dm−3, T = 298.15 K). As regards the stability of Pd2+ complexes with APP ligands, to our knowledge only Rizkalla et al.63 reported some stability constants for the Pd2+/HEDP system (in KNO3). Very high discrepancies are present between these data and the data reported in this work. For example it is reported a value of log KPdL = 5.74 (I = 0.1 mol·dm−3 in KNO3), impossible to be compared with the value of log KPdL = 23.4 (I = 0.1 mol·dm−3 in NaClO4). Also the data of Rizkalla63 were not selected as critical by the most important stability constants database.22,46,47

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complexes involving all the six binding sites but this property is also valid for all the other ligands studied in this work. The contribution of the different binding sites to the stability of the Pd2+/APC complexes have been studied in the paper of De Stefano et al.7 and an empirical relationship, which has been successfully tested with the data of this work, have been proposed on the basis of the number of amine (nN) and carboxylic (nCOO‑) binding sites and on the distance (number of −CH2− spacers) between the amino groups of the molecule (d), namely. log KPd(APC) ± 0.6 = 12.8 + 0.8nN 2 + 1.7nCOO − + 0.7d

Unfortunately, it was not possible to produce a similar equation for the APP ligands, because only few data are available. Although the APC ligands showed better binding ability than the APP, the latter class of ligands can be considered a good replacement for the sequestration of metals, due to their biodegradability and therefore lower permanence time in the environment. In this context it is important to underline that (i) the formation constants of the Pd2+/HEDP complexes are very similar to those of the Pd2+/S,S-EDDS system (see Table 3), which is often regarded as the most biodegradable of the APC ligands, (ii) in contrast to the behavior of all of the others APC and APP ligands, the sequestering ability of DTPP increases with pH. In addition the polyphosphonic ligands are more soluble with respect to the polycarboxylic ones, especially in acidic pH conditions. As a general trend, the values of the stability constants decrease with increasing both ionic strength and temperature, indicating that the complex formation reaction is exothermic. Although it was very difficult to derive quantitatively significant thermodynamic quantities, such as ΔH, it was possible to conclude that the entropic contribution (TΔS) to the stability of the complexes is dominant, because the variation of the formation constants with temperature is small (indicating small ΔH values) compared with the value of the Gibbs free energy (ΔG). Some empirical relationships were also found to model: (a) the dependence of the complex formation constants on the number of proton of the species, founding a common slope of s = −3.8 ± 0.3; (b) the dependence of the pL0.5 of the HEDP toward Pd2+ as a function of pH and temperature; (c) the stability of the Pd2+/APC on the basis of the number of the different binding sites.7 A general good agreement was found with literature findings and the discrepancies between different results can generally be explained with the different experimental conditions, which for the determination of such high complex formation constants are fundamental.

CONCLUSIONS

This work represents an improvement of a previous study reported by De Stefano et al.7 In this work a speciation study of different APC and APP ligands has been carried out to determine their binding ability toward a precious metal cation, such as Pd2+, as a function of temperature, pH and ionic strength. The stability of the PdL species was determined using potentiometric titration with an auxiliary ligand, iodide or ammonia, to ensure the equilibrium between these two ligands and allow a correct determination of the stability of this species, with a such high stability constant (K > 1020). It is noticeable that the PdL species determined for EGTA and HEDP by spectrophotometric and potentiometric technique are very similar. It is important to underline that the uncertainties associated with the formation constants determined with this method are affected by the possible systematic errors in the values of the formation constants between Pd2+ and the auxiliary ligand. For example, if the values of the Pd2+/NH3 complexes are lowered by 0.5 log K units, the value of the refined Pd(HEDP)2− species decreases from log K = 23.49 to 23.24. On the contrary, if the values of the Pd2+/NH3 complexes are increased by 0.5 log K units, the value of the refined Pd(HEDP)2− species goes from log K = 23.49 to 23.70. Therefore, in the estimation of the uncertainties for the stability constants of the Pd2+/APP and Pd2+/APC systems the above considerations must be taken into account. All of the ligands investigated showed a good binding ability toward Pd2+, although that of the DTPA is stronger than the others and for this ligand the highest values of formation constants and pL0.5 were found. These last properties can be useful in the applicative field, because DTPA can be used for the selective sequestration of Pd2+. The binding capacities of the ligands may be attributed mostly to the presence of the amino groups, in fact from the analysis of literature data the values of the formation constants between Pd2+ and amines are higher than those of the Pd2+ complexes with carboxylates. As an example, the formation constant of the simple PdL species is log K = 3.46 and 23.6 for citrate and ethylenediamine, respectively. The stability of the DTPA ligand, with three amino groups is similar to the stability of the diethylenetriamine (dien), and log K = 32.6 and 36.3 for dien and DTPA, respectively (considering that DTPA has also five carboxylic groups). In addition, all the ligands studied in this work form chelate rings, which usually enhance the stability of the complexes. For example in many papers it is reported that EDTA forms

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel: +39-090-6765749. Fax: +39090-392827. Funding

We thank University of Messina and University of Palermo for partial financial support. Notes

The authors declare no competing financial interest. 1981



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Article

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ACKNOWLEDGMENTS The authors greatly acknowledge referee #1 for its helpful suggestions

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Ligands Toward Palladium(II) - American Chemical Society

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