Interaction of oxovanadium(IV) with carboxylic ligands in aqueous solution: A thermodynamic and visible spectrophotometric study

Journal of Molecular Liquids 142 (2008) 57–63

Contents lists available at ScienceDirect

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

Interaction of oxovanadium(IV) with carboxylic ligands in aqueous solution: A thermodynamic and visible spectrophotometric study Silvia Berto a, Pier G. Daniele a, Claudia Foti b, Enrico Prenesti a,⁎, Silvio Sammartano b,⁎ a b

Dipartimento di Chimica Analitica, Università di Torino, Via Pietro Giuria 5, 10125 Torino, Italy Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Università di Messina, Salita Sperone 31, 98166 Messina (Vill. S. Agata), Italy

A R T I C L E

I N F O

Article history: Received 17 January 2008 Received in revised form 24 April 2008 Accepted 28 April 2008 Available online 6 May 2008 Keywords: Vanadyl hydrolysis Vanadyl complexes Carboxylate ligands Vanadyl absorption spectra Dependence on ionic strength

A B S T R A C T The hydrolysis of VO2+ and the complex with sulfate were studied potentiometrically, spectrophotometrically and calorimetrically, in NaCl aqueous solution (0 b I ≤ 1 mol L− 1) and at t = 25 °C. The formation of two hydrolytic species VO(OH)+ and VO2(OH)2+ 2 and one complex with sulfate was found, with log β = −5.65 for + the reaction VO2+ + H2O = VO(OH)+ + H+, log β = −7.02 for the reaction 2VO2+ + 2H2O = (VO)2(OH)2+ 2 + 2H and log K = 1.73 for VOSO04 species (at I = 0.1 mol L− 1 and t = 25 °C). For these species, using calorimetric data, ΔH and TΔS values were also obtained. By using the above values, interactions of VO2+ with acetate (ac), malonate (mal), succinate (suc), 1,2,3-propanetricarboxylate (tca) and 1,2,3,4-butanetetracarboxylate (btc) ligands were studied potentiometrically and spectrophotometrically. The formation of ML+, ML02 and MLOH0 for ac; ML0, − 0 + − − 0 2− − MLH+, ML2− 2 and ML2H for mal; ML , MLH and MLOH for suc; ML and MLH for tca and ML , MLH and −1 0 MLH2 for btc were found. Formation constants are reported at I = 0.1 mol L , together with SIT parameters for the dependence on ionic strength. By visible spectrophotometric measurements, λmax and εmax values for the relevant species in solution were determined. © 2008 Elsevier B.V. All rights reserved.

1. Introduction The main scientific interest for both vanadium and vanadium compounds is related to their involvement in metallurgical, fuels and coal (energetic), catalytic, electrochemical (vanadium fuel cells) and pharmaceutical fields. Industrial applications of vanadium compounds are very wide, hence, environmental accidents related to vanadium diffusion must be considered, since the toxicity [1] of some species is remarkable. Vanadium stays in the air, water and soil for a long time although the uptake to soil and sediments is relevant, since it does not easily dissolve in water. Vanadium mainly enters the environment from natural sources and from the burning of fuel oils [2,3]. Corrosion phenomena due to vanadium compounds in fuels cause an increase of costs related to the management of energetic facilities. Vanadium-based oxides are appreciated industrial oxidation catalysts [4], mainly used in the SO2/SO3 conversion. Low levels of vanadium have been currently found in plants, but it is not likely to build up in the tissues of animals. The biological effects, bio-distribution, toxicology and pharmacological activity of vanadium compounds are areas of increasing research interest. Despite the amount of the knowledge so far

⁎ Corresponding authors. E-mail addresses: [email protected] (E. Prenesti), [email protected] (S. Sammartano). 0167-7322/$ – see front matter © 2008 Elsevier B.V. All rights reserved. doi:10.1016/j.molliq.2008.04.014

accumulated, a clearly defined role in the higher forms of life has not yet been assigned to vanadium. In its anionic forms (vanadates, vanadium(V)) it resembles phosphates, but in its cationic forms— mainly as VO2+—it behaves like a typical transition metal ion, which is in competition with other metal cations in coordination with biogenic compounds. This duality, together with the facility with which it changes oxidation states and coordination environments, may be responsible for the very peculiar and somewhat unparalleled behaviour of this metal [5]. Literature fonts report that vanadium can improve the glucose tolerance (the normoglycemic effects is attributed to V(III), V(IV) and V (V) compounds) and reduce the cholesterol biosynthesis [6–14]. There are on sale, for example, pharmaceutical products, mainly classified as nutritional supplements, containing vanadyl salts. The investigation of vanadyl ion chemistry supplies essential information for the knowledge of its biochemical behaviour and the study of chemical equilibrium and structural characteristics of vanadyl organic complexes might provide information essential for synthesising new pharmaceutical aids. Sakurai et al. [11] studied the normoglycemic effect of VO2+ complexes containing four oxygen donor groups on diabetic rats and they found that the coordination compounds show different effectiveness: for example, the bismalonato vanadyl complex is more active than the bis-tartrate one. The study of vanadyl speciation in water solution, and the structural characterisation of its coordination compounds, could lead to the determination of parameters able to explain an effective

58

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63

Fig. 1. Visible absorption spectra of a 0.1 mol L− 1 VOSO4∙5H2O solution. a) Solution freshly prepared; b) solution after a year from preparation.

differentiation and to improve the comprehension of vanadium in the human body. In the literature there are many reports about vanadyl coordination compounds, but most of the scientific works concern vanadium applications in pharmaceutical or catalytic processes. There are not many papers about chemical speciation of vanadyl with simple organic molecules, that are commonly present in environment and in living organisms, and above all, there are not many studies about carboxylic ligands such as mono-, bi-, tri-and tetra-carboxylic acids [15–18]. In this paper we have studied the stability of coordination compounds of vanadyl ion with acetic, malonic, succinic, 1,2,3propanetricarboxylic or 1,2,3,4-butanetetracarboxylic acids in aqueous solution. Since we used vanadyl sulfate in our experiments, the formation of the VO2+–sulfate ion pair has also been investigated. pH-metric, calorimetric and visible spectrophotometric techniques were joined to reach our goals. First the different binary systems have been studied by potentiometric technique at t = 25 °C and, successively, the calculation of the visible absorption spectra were performed for those vanadyl complexes reaching significant percentage of formation in solution, with the aim of giving a better characterization of single species. Since the applicability of our results to real chemical systems must be as wider as possible, we extended the potentiometric investigation to different ionic strengths, in the range 0 b I/mol L− 1 ≤ 1.

2.2. Standardization of vanadyl stock solution Vanadyl sulfate stock solution was standardized by a redox titration followed by photometric detection. This procedure is an improvement of a literature method [19], according to which the titrant, a standard solution of KMnO4, oxidizes vanadium(IV) to vanadium(V). When the vanadyl is completely oxidized, the first permanganate excess (end-point) in the original procedure is visually detected, while in this work we have identified it by photometric measurements at 525 nm, by transferring the sample solution to an optical cell using a peristaltic pump. The accuracy of the method is very good, while the precision is within ±1%. The vanadyl ion is stable with respect to oxidation only at pH below 2 [20]. Since the natural pH of our vanadyl stock solution is about 2.3, the solution stability during the storage is not ensured. Owing to oxidation there is a decrease of the vanadyl concentration (and of pH value) and the consequent increase of absorbance values below 400 nm due to the presence of vanadium(V). Fig. 1 shows, as an example, the absorption spectra of the same solution of vanadyl sulfate, freshly prepared and a year later (the variation of titre in our vanadyl stock solution is about 5%). By recording the spectrophotometric behaviour of vanadyl stock solution in the time, we observed that the degradation process can be considered negligible during a week. Therefore, in order to ensure the absence of vanadium(V), the metal solution was weekly prepared and standardized.

2. Experimental 2.3. Electromotive force measurements 2.1. Chemicals Vanadyl sulfate (vanadium(IV) oxide sulfate pentahydrate, purity ~ 96%, Riedel-de Haën or Aldrich), stock solution (∼0.1 mol L− 1) was weekly prepared without previous purification of the salt; it was standardized by redox titration with permanganate solution (Carlo Erba) [19], followed by photometric detection (see below). Acetic (ac), malonic (mal), succinic (suc), 1,2,3-propanetricarboxylic (tricarballylic, tca) and 1,2,3,4-butanetetracarboxylic (btc) acids (Merck or Fluka) were used without further purification and their purity, checked alkalimetrically, was found to be N99%. Sodium chloride solutions were prepared by weighing pure salt (Fluka, p.a.) previously dried in an oven at 110 °C. Standard NaOH and HCl solutions were prepared by diluting Merck or Fluka concentrate products and standardized against potassium hydrogen phthalate (Fluka, puriss.) and sodium carbonate (Fluka, puriss.), respectively. All solutions were prepared using grade A glassware and ultrapure water (conductivity b 0.1 μS).

In order to avoid systematic errors, potentiometric measurements were performed using two different equipments. 1) A Metrohm mod. 713 potentiometer (resolution of ±0.1 mV) coupled with a Metrohm 665 Dosimat burette (minimum volume deliverable of ±0.001 cm3)

Table 1 Thermodynamic parameters for the hydrolysis and sulfate complexes of VO2+, at I = 0.1 mol L− 1 (NaCl) and t = 25 °C, together with Δε parameters for the dependence on ionic strength (Eq. (1))

2+

Δεa

log β

Reaction +

+

c

VO + H2O = VO(OH) + H −5.65 ± 0.15 + 2VO2+ + 2 H2O + (VO)2(OH)2+ −7.02 ± 0.03 2 + 2H 2− 0 2+ VO + SO4 = VOSO4 1.73 ± 0.05 a b c

ΔGa,b c

0.41 ± 0.01 −0.41 ± 0.01 0.99 ± 0.03

Calculated from log β values in molal scale (mol kg− 1). In kJ mol− 1. Standard deviation.

33.1 40.1 −9.9

ΔHb

TΔSb c

12 ± 4 −20 36.9 ± 1.0 −3 2.6 ± 1.0 13

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63

59

0 Fig. 2. Speciation diagram of VO2+ at different temperatures. Experimental condition: CVO2+ = 1 mmol L− 1, I = 0.1 mol L− 1 (NaCl). Species: 1. VO(OH)+, 2. (VO)2(OH)2+ 2 , 3. VOSO4.

and equipped with a Metrohm combined glass electrode (mod. 6.0222.100). The potentiometer and the burette were connected to a personal computer which, using suitable software, allows automatic data acquisition. 2) A Metrohm model 809 titrando coupled with a Metrohm 800 Dosino dispenser and equipped with an Orion (Ross mod. 8101) glass electrode and an Ag/AgCl standard electrode. Data were automatically acquired by the software Metrohm TiAMO 1.0. For all the potentiometric measurements the electrode couple was standardized, in terms of pH = −log[H+], by titrating HCl 10 mmol·L− 1 solution (at the same ionic strength value as the solution under study) with standard NaOH in order to determine the standard potential E° before each experiment. The potentiometric titrations were carried out in a stream of purified nitrogen gently bubbled in the titration cell to avoid O2 and CO2 contamination. The measurement cells were thermostated at (25 ± 0.1) by means of a water circulation from a thermocryostat (mod. D1-G Haake). The experiments were carried out in NaCl aqueous solutions at different ionic strength values (I = 0.1, 0.5 and 1.0 mol·L− 1), and at t = 25 ± 0.1 °C. For investigation of VO2+ hydrolysis and sulfate complexes, 25 mL of solution containing VOSO4 (1–10 mmol L− 1) and NaCl was titrated with NaOH standard. For the investigation of VO2+–carboxylate systems, 25 mL of solution containing VO2+ (1– 2 mmol L− 1, but for VO2+–ac system 1–5 mmol L− 1), the carboxylate ligand (ac: 10–100; mal and suc: 2–10; tca and btc: 1–2 mmol L− 1) and NaCl in order to reach the prefixed ionic strength values (I = 0.1, 0.5 and 1.0 mol·L− 1), was titrated with NaOH standard. For each ionic strength value four measurements were performed and, for VO2+–ac, –mal and –suc systems four measurements were performed in KCl 0.1 mol L− 1, too. Each titration was at least twice repeated, usually, in the range 2 ≤ pH ≤ 5.5.

measured before each experiment in the same ionic strength conditions of the measure. The accuracy of calorimetric apparatus was Q ± 0.008 J and v ± 0.001 cm3. 2.5. Spectrophotometric measurements The visible spectrophotometric determinations (500–900 nm) were carried out for VO2+–carboxylate systems, at I = 0.1 mol L− 1 with a model V-550 Jasco spectrophotometer (optical path 1.000 cm). The solution being examined was transferred from the potentiometric to an optical cell using a peristaltic pump. Due to the low values of molar absorptivity coefficients of vanadyl containing species, the concentration of oxocation was always N5 mmol L− 1, with the suitable metal to ligand ratio. 2.6. Data analysis and calculations To determine all parameters of the acid-base titration, the calculations were performed by using the non-linear least squares computer program ESAB2M [21]. This program allows us to refine the analytical concentration of the reagents, the electrode formal potential E°, the coefficient ja relative to the acidic junction potential (according to the equation: Ej = ja [H+]) and the ionic product of water Kw; it is also useful to evaluate the purity of the ligand examined. The refinement of the formation constants was performed by the BSTAC [22] software. It employs an iterative and convergent numerical method, which is based upon the linear combination of the mass balance equations, minimises the error squares sum on electromotive force values and takes into account eventual variations of ionic strength among and/or during titrations.

2.4. Calorimetric measurements Calorimetric measurements were performed at 25.000(± 0.001) °C using a Tronac isoperibol model 450 titration calorimeter coupled with a Keithey 196 system Dmm digital multimeter. The titrant was delivered by a 2.5 ml capacity model 1002TLL Hamilton syringe. A computer program was used for the acquisition of calorimetric data. The accuracy was checked by titrating a Tris [tris-(hydroxymethyl) amino-methane] buffer with HCl. 50 ml of a solution obtained by adding to vanadyl sulfate (CVO2+ = 2–4 mmol L− 1) NaOH until the formation of hydrolytic species, and NaCl (I = 0.1 mmol L− 1) was titrated calorimetrically by HCl standard. The heat of dilution was

Table 2 Literature values of formation constants of VO2+ hydrolytic and –SO2− 4 species Ligand

t (°C)

I (mol L− 1)

log βML

log βM2L2

Reference

OH

25 25 25 25 20 25 25

1 0 3 3 1

−6.07 −5.67 −6.4 −6.0 1.74 2.40 2.48

−6.59 −6.67 −7.45 −6.88

27 28 29 30 31 32 33

SO4

0

60

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63

Table 3 Formation constants for VO2+–carboxylate systems, at I = 0.1 mol L− 1 (NaCl) and t = 25 °C, together with parameters of SIT equation (Eq. (1)) Ligand ac

suc

tca btc

b c

Δεb c

110 120 1 1−1 110 111 120 121 110 111 1 1−1 110 111 110 111 112

mal

a

log βpqra

pqr

1.81 ± 0.05 3.30 ± 0.15 − 2.56 ± 0.15 6.12 ± 0.03 8.21 ± 0.06 9.92 ± 0.10 13.22 ± 0.10 3.25 ± 0.05 7.20 ± 0.05 −1.75 ± 0.15 4.16 ± 0.05 8.21 ± 0.10 5.50 ± 0.05 9.11 ± 0.05 14.03 ± 0.15 2+

z−

+

0.42 ± 0.01c 0.68 ± 0.10 0.32 ± 0.12 −0.17 ± 0.03 0.52 ± 0.06 0.13 ± 0.05 1.31 ± 0.08 0.46 ± 0.05 1.00 ± 0.04 0.61 ± 0.10 0.40 ± 0.07 − 0.30 ± 0.07 0.30 ± 0.01 1.20 ± 0.01 0.72 ± 0.01

+ r −qz = (VO)pLqH2p . r −1

Refer to the general reaction: pVO + qL + rH Calculated from log β values in molal scale (mol kg ). ±standard deviation.

Table 4 Protonation constants of carboxylate ligands at I = 0.1 mol L− 1 (NaCl) and t = 25 °C Ligand

i

log βia

Reference

ac mal

1 1 2 1 2 1 2 3 1 2 3 4

4.58 5.275 7.912 5.232 9.225 5.822 10.324 13.806 6.30 11.46 15.57 18.76

35 36

suc tca

btc

a

36 37

38

Index i refers to the reaction: iH+ + Lz− = HiLi − z.

Spectrophotometric data were analysed by means of the HYPERQUAD software [23] which calculates the values of molar absorptivity coefficients (ελ/mol− 1 L cm− 1) of the different complexes by using experimental spectra (absorbance vs. wavelength λ/nm), analytical

concentrations of the reagents and the proposed chemical model (stoichiometric coefficients and known stability constant values of all complexes) as input. After the calculation of the species distribution, absorption spectra are estimated for each complex formed in solution, only assuming the additivity of the absorbance in the investigated concentration range. Neither assumptions on the shape of the curves nor on the nature of electronic transitions are taken into account by the program. Calorimetric titration data were analysed by the computer program ES5CM [24]. 3. Results and discussion 3.1. Hydrolytic species and sulfate ion pair of vanadyl ion The literature reports the formation of three hydrolytic species of 0 vanadyl ion in aqueous solution: VO(OH)+, VO2(OH)2+ 2 and VO(OH)2 [25–30]. Alkalimetric titrations of solutions containing only the oxocation and NaCl as supporting electrolyte have been performed in order to determine the formation constant of hydrolytic species in the experimental conditions mentioned above. Since our source of vanadium is the vanadyl sulfate salt, the formation of the ion pair VO2+–SO2− 4 is never negligible in our experimental conditions, despite the low concentration of sulfate ion in the solutions under examination. Therefore, the refinement of potentiometric data related to hydrolysis must take into account at the same ion pair, whose formation is fairly independent of pH. The elaboration of pH-metric data obtained as reported in the Experimental section shows that a chemical model which assumes the formation of two hydrolytic 0 species, namely VO(OH)+ and VO2(OH)2+ 2 , together with VOSO4 species, is able to explain the experimental trend of titration curves. Formation constants were determined in the ionic strength range 0 b I ≤ 1 mol L− 1, and are reported in Table 1 at I = 0.1 mol L− 1 and t = 25 °C, together with ΔG, the parameter Δε for the dependence on ionic strength (see below, Eq. (1)), and ΔH and TΔS values determined by the calorimetric measurements. Distribution diagram of this system is shown in Fig. 2, at three different temperatures: 5, 25 and 37 °C and in the experimental condition CVO2+ = 1 mmol L− 1 and I = 0.1 mol L− 1. As can be seen, until pH ∼ 4.5 the only species is VOSO04 that reaches a percentage formation of 60%. At pH 5–5.5 and at + t = 25 °C, ∼ 50% of VO2+ is present as (VO)2(OH)2+ 2 and ∼15% as VO(OH) . Temperature influences the distribution diagram, in particular

Fig. 3. Distribution diagram of VO2+–mal2− system at I = 0.01 and 1 mol L− 1 (NaCl) and at t = 25 °C. Experimental condition: CVO2+ = 1, Cmal = 2 mmol L− 1. Species: 1. VO(mal)0, 2. VO(mal)H+, 3+ 3. (VO)(mal)2+ 2 , 4. (VO)(mal)2H .

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63 Table 5 Literature valuesa for the formation constants of VO2+–carboxylate complexes

61

Table 6 Values of λmax and εmax for the relevant species in solution

Ligand

t (°C)

I (mol L−1)

log βML

log βML2

Ligand

p q ra

λmax (εmax)b

λmax (εmax)b

ac mal

25 25 25 30 21 25 30

1 0.2 1 0.1 1 0.1 0.1

1.86 5.62 5.59 6.10 5.23 3.20 3.65

2.96 9.20 9.48 10.60 8.85 5.60 –

H2O SO2− 4 ac mal

10 0 110 110 111 110 121 120 111 110 111 110 11 2 111 110

765 (16) 765 (19) 770 (21) 768 (19) 782 (27) 785 (24) 802 (39) 774 (19) 794 (34) 790 (32) 795 (21) 785 (26) 793 (17) 797 (26)

635 (7) shc 635 (8) sh 630 (9) sh 630 (8) sh 610 (8) sh 595 (6) 577 (8) 630 (9) sh 610 (17) sh 615(12) sh 610 (8) sh 615 (11) sh 610 (9) sh 600 (11) sh

suc a

Ref. 25,26.

suc tca

concerning the formation of (VO)2(OH)2+ 2 , that varies (at pH = 5.5) from 25% at 5 °C until 60% at 37 °C. Table 2 reports the literature data on formation constants of hydrolytic species and sulfate complexes [27–33]. Values are in accordance with ours: for example at I =0 mol L− 1 and t =25 °C, from the data of Table 1 and Eq. (1) we obtain log β =−5.61, −6.76 and 2.49 for VO(OH)+, VO2(OH)2+ 2 , and VOSO04, respectively. For ΔH values, a comparison can be made only [34] with the values of ΔH =53.1 kJ mol− 1 for (VO)2(OH)2 species [determined by kinetic method, I =0.3 mol L− 1 (NaClO4), t =25 °C]. 3.2. Vanadyl complexes Formation constants for vanadyl–carboxylate complexes are reported in Table 3 (at I = 0.1 mol L− 1, 25 °C), together with Δε parameters for the dependence on ionic strength, as described later. Calculations were performed by considering vanadyl hydrolysis constants determined in this paper (Table 1) and literature protonation constants of carboxylates reported in Table 4 [35–38]. All the species are formed by the involvement of one or two ligand molecules bound to the oxocation coordination sphere, with carboxylate groups. As regards their stability, some aspects can be pointed out. The chelation effect for the six-membered chelate ring of malonate (log β110 = 6.12) is quite relevant with respect to the seven-membered one formed by succinate (log β110 = 3.25); the same behaviour has been reported in the literature [25,26] for many other bivalent cations, although the difference in stability is not so high. As regards 1,2,3propanetricarboxylic and 1,2,3,4-butanetetracarboxylic acids, their formation constants with VO2+, determined for 110 species (log β110 = 4.16 and 5.50, respectively), are significantly higher than that obtained for suc (log β110 = 3.25) and indicate the participation of a third (tca) or of a fourth (btc) carboxylate group in VO2+ coordination

btc

a b c

+ r − qz p q r indexes refer to the general reaction pVO2+ + qLz− + rH+ = (VO)pLqH2p . r λmax in nm; εmax in mol− 1 L cm− 1. sh = shoulder.

and furthermore that the stability resulting from the formation of two or three seven-membered chelate rings is lower than that results from the formation of one six-membered chelate ring (log β110 = 6.12 for VO2+–malonate complex). The above formation constant for succinate is close to the value determined for log β120 of acetate (3.30), suggesting that, in this case, the chelate effect can be considered negligible. As regards protonated complexes, the values of log Ki should give us a better estimate of their stability (log Ki is referred to the reaction VO2+ + HiLi − z = VOHiL2 + i − z, with z = charge of the fully deprotonated ligand; all the below values are calculated at I = 0.1 mol L− 1, from the data of Tables 3 and 4). If compared with the value found for vanadyl 110 complex with acetate (log β110 = 1.81), the value of logK1 = 2.94 and 1.97, for 111 complex of mal and suc, respectively, might suggest an involvement, more relevant for hydrogen malonate, of indissociated carboxylic group in coordination. On the contrary, both the values of logK1 = 2.39 for tca and logK2 = 2.57 for btc, if compared with that of log β110 for suc (3.25), and the value of logK1 = 2.81 for btc, if compared with that of log β110 for tca (4.16), show a significantly lower stability, excluding any participation of carboxylic groups in coordination. It is interesting to note that, if considering the stability of 1:1 species for all the studied carboxylic ligand (but malonate), we have a quite regular trend with a free energy contribution for single bond given by ΔGs/n = 10.5 (±0.8)–0.7 (±0.2) n (ΔGs/kJ mol− 1,

Fig. 4. Experimental visible spectra for VO2+–mal2− system. CVO2+ = 8.0, Cmal = 16.0 mmol L− 1 at I = 0.1 mol L− 1. pH values: a. 1.68, b. 2.10; c. 2.51; d. 3.55; e. 4.25.

62

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63

Fig. 5. Calculated absorption visible spectra of VO2+, VOSO40 and VO2+–mal2− complexes (index pqr refer to reaction pVO2+ + qmal2− + rH+ = (VO)p(mal)qHr2p + r − 2q).

n = number of carboxylic groups involved in the coordination). As a consequence VO(mal)0 shows a chelate free energy given by ΔG−1 2+ Chel = 16 ± 2 kJ mol . Fig. 3 shows the distribution diagram of VO –mal system, at two different ionic strength values (I = 0.01 and 1 mol L− 1). In the range 2.5 ≤ pH≤ 4.5, the most important species is VO(mal)0, that reaches 90% of formation at I = 0.01 mol L− 1, but this percentage decreases at 40% at I = 1 mol L− 1, while the protonated species (VO) (mal)2H3+, that is negligible at I = 0.01 mol L− 1 reaches 30% of formation at I = 1 mol L− 1. In the range 4.5 ≤ pH ≤ 5.5, the formation percentage of (VO)(mal)2+ 2 species varies from 50 to 70%, varying the ionic strength from 0.01 to 1 mol L− 1. Comparison with literature data can be made only for ac, mal and suc [25,26], as reported in Table 5. As can be seen, for ML species our values are in accordance with literature ones. There are some differences for ML2 species, probably due to the different speciation model considered. 3.3. Dependence on ionic strength of formation constants Formation constants of VO2+–OH, –SO2− 4 and -carboxylate species were determined in the ionic strength range 0 b I ≤ 1 mol L− 1 (NaCl). Dependence on ionic strength was taken into account by the SIT (Specific ion Interaction Theory) equation [39]. When using the SIT approach, both ionic strength and formation constants must be given in the molal concentration scale. In the short ionic strength range considered in this work (0 ≤ I ≤ 1 mol L− 1), appropriate conversion can be made by simple equations: −1

m=mol kg

¼ 1:02 c=mol L−1

Δ ¼ −ð1:3 þ 8:1 cÞ 10−3 where

 −1 −1 Δ ¼ log Km =kg mol − log Kc =L mol According to SIT model [39], formation constants in molal scale Km follow the equation: ! pffiffi pffiffiffi I I0 pffiffiffi þ Δε ðI−I0 Þ pffiffi − log Km ¼ log KmV −z ð1Þ 1 þ 1:5 I 1 þ 1:5 I0 where K' is the formation constant at the reference ionic strength I′ (here I' = 0.1 mol kg− 1), I is ionic strength in the molal concentration scale, z⁎ = Σz2reactants − Σz2products, Δε = Σεreactants − Σεproducts (ε = specific

interaction coefficient). When neutral species is formed Δε parameter includes the salting coefficient of neutral species. Δε values for each equilibrium constant are reported in Tables 1 and 3. 3.4. Visible spectrophotometric spectra The visible spectrophotometric measurements were performed at different pH values, in the same experimental conditions as potentiometric ones, both taking into account the metal cation alone and mixtures of VO2+ with acetate, malonate, succinate, 1,2,3propanetricarboxylate or 1,2,3,4-butanetetracarboxylate. An example of the spectra recorded for the VO2+–malonate system is drawn in Fig. 4. It can be pointed out the presence of two bands which can be assigned, on the basis of Ballhausen–Gray scheme [40], as follows: the stronger band with a bathochromic effect from ~ 770 to ~ 800 nm corresponds to b2 ⇒ e transition; the latter, which often appears as a shoulder, depending both on the pH and on the ligand coordinated, with an ipsochromic effect from ~ 625 to ~ 590 nm, corresponds to b2 ⇒ b1 transition. For both the bands the εmax values rise with increasing pH. A similar behaviour is shown by mixtures of vanadyl ion with ac, suc, tca and btc. Also solutions containing only vanadyl sulfate show two bands in the visible region with similar spectral characteristics, as previously reported [40,41]. For each system investigated, from the whole of the spectra recorded on solutions at different concentrations and pH values, we have also estimated the individual visible spectrum of the relevant species in solution. The principal spectral features of these complexes are listed in Table 6, while an example of spectra calculated for single complexes is drawn in Fig. 5 for vanadyl– malonate complexes. In agreement with the not many literature reports [40–42], all the species are characterized by two bands with higher intensity with respect to the vanadyl cation. Owing to the different coordination modes, the light bathochromic (for the stronger band) and ipsochromic (for the latter band) effects, above outlined for the experimental spectra at different pH, keep in the spectra of the single complex species, as one would expect. As previously shown for copper(II) complexes [43], further investigation will be necessary to verify the eventual existence of a spectrum-structure correlation for vanadyl complexes as well. Acknowledgements We thank ARPA SICILIA (Agenzia Regionale Protezione Ambiente) for financial support (project D.D.G n. 229 del 27.12.02 -Area Tematica

S. Berto et al. / Journal of Molecular Liquids 142 (2008) 57–63

n.6 “Modellistica e nuove tecnologie applicate alla valutazione dello stato dell’ambiente ed alla protezione ambientale”). References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12]

[13] [14] [15] [16] [17]

J.L. Domingo, Reprod. Toxicol. 10 (1996) 175. T. Soldi, C. Riolo, G. Alberti, M. Gallorini, G.F. Peloso, Sci. Total Environ.181 (1996) 45. B.K. Hope, Biogeochemistry 37 (1997) 1. B.M. Weckhuysen, D.E. Keller, Catal. Today 78 (2003) 25. E.J. Baran, J. Braz. Chem. Soc. 14 (2003) 878. A.M. Fiabane, D.R. Williams, Principles of Bio-inorganic Chemistry, The Chemical Society, London, 1977. I. Goldwaser, D. Gefel, E. Gershonov, M. Fridkin, Y. Shechter, J. Inorg. Biochem. 80 (2000) 21. S. Bhattacharyya, A.S. Tracey, J. Inorg. Biochem. 85 (2001) 9. G.R. Willsky, A.B. Goldfine, P.J. Kostyniak, J.H. McNeill, L.Q. Yang, H.R. Khan, D.C. Crans, J. Inorg. Biochem. 85 (2001) 33. S. Takeshita, I. Kawamura, Y. Yasuno, C. Kimura, T. Yamamoto, J. Seki, A. Tamura, H. Sakurai, T. Goto, J. Inorg. Biochem. 85 (2001) 179. H. Sakurai, Y. Kijima, Y. Yoshikawa, K. Kawabe, H. Yasui, Coord. Chem. Rev. 226 (2002) 187. S. Almada, M.T. Amorim, L. Afonso, H.P. Caldeira, E. Silva, J.A. Silva, J.J.R. Fraùsto da Silva, M.C. Vaz, Abstract of the 7th International Conference on Bioinorganic Chemistry, Lübeck, Germany, 1995, p. 589. C. Orving, K.H. Thompson, B.D. Liboiron, S.A. Dikanov, Abstract of the 10th International Conference on Bioinorganic Chemistry, Florence, Italy, 2001, p. 41. H. Sakurai, K. Kawabe, H. Yasui, Y. Kojima, Y. Yoshikawa, Abstract of the 10th International Conference on Bioinorganic Chemistry, Florence, Italy, 2001, p. 94. P. Di Bernardo, G. Tomat, P. Zanonato, P. Portanova, M. Tolazzi, Inorg. Chim. Acta 145 (1988) 285. L.D. Pettit, J.L.M. Swash, J. Chem. Soc., Dalton Trans. (1978) 286. R.E. Tapscott, R.L. Belford, Inorg. Chem. 6 (1967) 735.

[18] [19] [20] [21] [22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38] [39] [40] [41] [42] [43]

63

J.H. Dunlop, D.F. Evans, R.D. Gillard, G. Wilkinson, J. Chem Soc. A (1966) 1260. F.P. Treadwell, in: F. Vallardi (Ed.), Chimica Analitica, 1996. K. Pyrzynska, T. Wierzbicki, Talanta 64 (2004) 823. C. De Stefano, P. Princi, C. Rigano, S. Sammartano, Ann. Chim. (Rome) 77 (1987) 643. C. De Stefano, P. Mineo, C. Rigano, S. Sammartano, Ann. Chim. (Rome) 83 (1993) 243. P. Gans, A. Sabatini, A. Vacca, Talanta 43 (1996) 1739. A. De Robertis, C. De Stefano, C. Rigano, Thermochim. Acta 138 (1986) 141. R.M. Smith, A.E. Martell, R.J. Motekaitis, NIST Critical Selected Stability Constant of Metal Complexes Databases, version 6.0, 2002. L. Pettit, K.J. Powell, The IUPAC Stability Constants Database, Academic Software, 2001. I. Nagypal, I. Fabian, Inorg. Chim. Acta 61 (1982) 109. R. Henry, P. Mitchell, J. Prue, J. Chem. Soc. Dalton Trans. (1973) 1156. S. Mateo, F. Brito, An. Quim. 68 (1972) 37. F. Rossotti, H. Rossotti, Acta Chem. Scand. 9 (1955) 1166. A. Golub, A. Tishchenko, A.M. Kalinichenko, I.S. Kononenko, Isvest. VUZ. Khim. 16 (1973) 342. G. Kohler, H. Wendt, Ber. Buns. Phys. Chem. 70 (1966) 674. H. Strehlow, H. Wendt, Inorg. Chem. 2 (1963) 6. B. Lutz, H. Wendt, Ber. Buns. Phys. Chem. 74 (1970) 372. A. De Robertis, C. De Stefano, C. Rigano, S. Sammartano, R. Scarcella, J. Chem. Res. (S) 42 (1985) 629 (M). F. Crea, A. De Robertis, S. Sammartano, J. Solution Chem. 33 (2004) 497. C. De Stefano, C. Foti, A. Gianguzza, Talanta 41 (1994) 1715. A. De Robertis, C. Foti, A. Gianguzza, Ann. Chim. (Rome) 83 (1993) 485. L. Ciavatta, Ann. Chim. 70 (1980) 551. C.J. Ballhausen, H.B. Gray, Inorg. Chem. 1 (1962) 111. C. Klixbüll Jørgensen, Acta Chem. Scand. 11 (1957) 73. J.C. Costa Pessoa, T. Gajda, R.D. Gillard, T. Kiss, S.M. Luz, J.J.G. Moura, I. Tomaz, J.P. Telo, I. Török, J. Chem. Soc. Dalton Trans. (1998) 3587. E. Prenesti, P.G. Daniele, M. Principe, G. Ostacoli, Polyhedron 1 (1999) 8 3233.