Investigations of coordinating properties of oxydiacetate and thiodiacetate anions towards Zn2+ ions in solutions

Inorganica Chimica Acta 405 (2013) 163–168

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Investigations of coordinating properties of oxydiacetate and thiodiacetate anions towards Zn2+ ions in solutions D. Wyrzykowski a,⇑, I. Anusiewicz a, B. Pilarski b, D. Jacewicz a, L. Chmurzyn´ski a a b

´ sk, Sobieskiego 18, 80-952 Gdan ´ sk, Poland Faculty of Chemistry, University of Gdan Cerko Sp. z o.o. Sp. K, Al. Zwycie˛stwa 96/98, 81-451 Gdynia, Poland

a r t i c l e

i n f o

Article history: Received 15 March 2013 Received in revised form 20 May 2013 Accepted 24 May 2013 Available online 3 June 2013 Keywords: Isothermal titration calorimetry Potentiometric titration Zn(II) complexes Oxydiacetate Thiodiacetate Thermodynamic parameters

a b s t r a c t The isothermal titration calorimetry (ITC) method was used to study the coordinating properties in solutions of oxydiacetate (ODA) and thiodiacetate (TDA) anions towards the Zn2+ ion. The calorimetric measurements (ITC) were run in a Mes buffer with a pH of 6.0 at 298.15 K. Furthermore, the potentiometric titration (PT) technique as well as ITC displacement titrations were applied to obtain conditional-independent thermodynamic parameters of the interaction of Zn2+ with ODA (DH = +3.49 [kcal/mol], DS = 30.30 [cal/mol K]) and TDA (DH = +2.93 [kcal/mol], DS = 25.99 [cal/mol K]). The MP2 method with the 6–311++G(d,p) basis set and the polarized continuum (PCM) solvation model were used to compare the relative stabilities of mer and fac conformations of the ODA and TDA ligands in the octahedral coordination sphere of the zinc ion. The relationship between the proposed coordination modes of the ligands and the thermodynamic parameters has been discussed. Ó 2013 Elsevier B.V. All rights reserved.

1. Introduction In biological systems, the zinc(II) ions play important roles in both enzyme catalysis and maintaining structures [1]. Depending on the kind of ligands, the coordination sphere of the zinc(II) ion can adopt different geometries. In an aqueous solution, the Zn2+ ion is hexacoordinated by water molecules, but in zinc-finger proteins and enzymes, the coordination number equal 4 is preferable. On the other hand, in some catalytic binding sites the zinc(II) ion is penta- or, rarely, hexacoordinated [2]. The Zn(II) complexes with polycarboxylate ions such as oxydiacetate (ODA) and thiodiacetate (TDA) (Fig. 1) can serve as simple models for studying the zinc(II) complexation by some biologically important ligands. In the solid state, two polymeric crystal structures of the zinc(II) ion with oxydiacetate are known, namely a 2D polymeric structure of {[Zn(ODA)0.3H2O]}n (1) [3] and a 1D polymeric structure of {[Zn(ODA)(H2O)2]H2O}n (2) [4]. Compound (1) is the first example of structurally characterized zinc oxydiacetate complex. X-ray crystallographic results have shown that in (1) the zinc(II) cation has a trigonal–bipyramidal geometry comprising oxygen atoms of two carboxylic groups and an ether oxygen atom of the tridentate ODA ligand as well as two outer carboxylate oxygen atoms from two adjacent zinc polyhedra. Water molecules occupy isolated positions in the crystal structure and are not bound to the ⇑ Corresponding author. Tel.: +48 697409642. E-mail address: [email protected] (D. Wyrzykowski). 0020-1693/$ - see front matter Ó 2013 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.ica.2013.05.033

metal. The structure of compound (2) is isomorphous with the homologous cobalt(II) compound {[Co(ODA)(H2O)2]H2O}n [5]. The environment of the Zn(II) ion consists of a distorted octahedron comprising the tridentate ODA ligand, two molecules of coordination water and oxygen atom from an outer carboxylate group of neighboring ODA ligand which serves as the link in the resulting one-dimensional polymer. In both complexes (1) and (2) ODA ligands adopt a planar (mer) conformation. As far as Zn–TDA complexes are concerned, two structures in the solid state are known as well. They can be achieved by controlling the crystallization conditions. The polymeric structure of {[Zn(TDA)(H2O)]}n (3) was proposed by an inspection of IR spectra that is identical to that of the structurally characterized complex {[Cd(TDA)(H2O)]}n [6,7]. On the other hand, the structure of [Zn(TDA)(H2O)3]H2O (4) consists of discrete units in which the Zn(II) ion has a distorted octahedral environment formed by the sulfur atom of TDA, two oxygen atoms from each of two carboxy-groups and three oxygen atoms of coordination water molecules [8]. It is also worth noting that, opposed to the oxydiacetate complexes, the TDA ligand in (3) and (4) adopts a fac-coordinating conformation. A similar coordination geometry, mer for ODA and fac for TDA ligand, is also observed for other oxy- and thiodiacetate metal complexes [9]. In solution, however, the hexacoordinated complexes of Zn–ODA and Zn–TDA are expected to exist. In this contribution, thermodynamic parameters for complexation reactions of the zinc(II) ions with oxy- and thiodiacetate ligands as determined by the isothermal titration calorimetry (ITC) and potentiometric titration (PT)

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was carried out electrically by using electrically-generated heat pulses. The CaCl2–EDTA titration was performed to check the apparatus, and the results were compared with those obtained for the same samples (test kit) at MicroCal. 2.3. Potentiometric titration (PT)

Fig. 1. Oxydiacetate (ODA) and thiodiacetate (TDA) ligands.

methods are presented. In addition, the relative stabilities of the two possible isomers (mer and fac) of the compounds studied are compared with the predictions provided by theoretical calculations. The knowledge of thermodynamic stability of such species may be helpful for an in-depth interpretation of the Zn–carboxylate bonds’ role in the zinc complexes of biological interest. 2. Experimental 2.1. Materials All reagents, namely 2-(N-morpholino)ethanesulfonic acid hydrate (P99%) (Mes), 2,20 -oxydiacetic acid (P98%) (H2ODA), 2,20 -thiodiacetic acid (P98%) (H2TDA), Zn(NO3)26H2O (P99%), nitrilotriacetic acid (P99%) (H3NTA) and HClO4 were purchased from Sigma–Aldrich Chemical Corp. and used as received. Double-distilled water with conductivity not exceeding 0.18 lS cm1 was used for preparations of aqueous solutions of both titrant and titrand. 2.2. Isothermal titration calorimetry (ITC) All ITC experiments were performed at 298.15 K using an AutoITC isothermal titration calorimeter (MicroCal Inc., GE Healthcare, Northampton, USA) with a 1.4491-mL sample and the reference cells. The reference cell contained distilled water. All details for the measuring devices and the experimental setup were described in [10]. The reagents were dissolved directly into 10 mM Mes buffer of a pH of 6. Then, the pH of solutions was adjusted to 6.0 with 0.1 M HClO4. The experiment consisted of injecting 10.02 lL (29 injections, 2 lL for the first injection only) of ca 30 mM buffered solution of H2ODA or H2TDA into the reaction cell, which initially contained ca 1–2 mM of buffered solution of the zinc salt. The enthalpy of zinc binding to Mes was determined by the displacement titration in 100 mM Mes of a pH of 6. The buffered solution of zinc salt (0.5 mM) was placed in the reaction cell, and 5 mM of nitrilotriacetic acid was titrated into that solution. A background titration, consisting of an identical titrant solution but with the buffer solution in the reaction cell only, was removed from each experimental titration to account for the heat of dilution. All the solutions were degassed prior to the titration. The titrant was injected at 5-min intervals to ensure that the titration peak returned to the baseline before the next injection. Each injection lasted 20 s. For homogeneous mixing in the cell, the stirrer speed was kept constant at 300 rpm. The calibration of the AutoITC calorimeter

Potentiometric titrations were performed in a 30 mL thermostated (298.15 ± 0.10 K) cell using the Cerko Lab System microtitration unit fitted with a 0.5-mL Hamilton’s syringe, a pH combined electrode (Schott – BlueLine 16 pH type) and a self-made measuring cell equipped with a magnetic stirrer. The temperature was controlled using the Lauda E100 circulation thermostat. The electrode was calibrated according to IUPAC recommendations [11]. The syringe was calibrated by a weight method. All the solutions were prepared immediately before measurements. The compositions of the titrand solutions used in the experiments were as follows: (1) Zn2+ (2.00 mM), H2ODA (5.03 mM) and HClO4 (4.98 mM), (2) Zn2+ (2.05 mM), H2TDA (5.03 mM) and HClO4 (4.98 mM), (3) Zn2+ (2.64 mM), HMes (20.13 mM) and HClO4 (4.98 mM) and (4) H3NTA (2.01 mM). The solutions were potentiometrically titrated with a standardized NaOH solution (0.098 M) with the pH ranging from 2.5 to 12.0. The stability constants of the complexes were determined using the CVEQUID program [12] by minimization of the differences between the theoretical model and the experimental data, according to Gauss–Newton–Marquart for nonlinear equations (see Ref [13] for more details). The original CVEQUID algorithm was combined with the CERKOLAB software. After titration, the acquired data was processed and the equilibria model was symbolically described by set of equations. To simplify the preparation of data required for numerical procedures, the ‘‘pK model’’ included statements describing the composition of titrant, titrand, electrode parameters and solvent. The stoichiometric matrix required for the procedure was automatically generated from the model. 2.4. Computational details The equilibrium geometries of the complexes were determined by the second-order Møller–Plesset perturbational method (MP2) with the 6–311++G(d,p) basis set [14,15]. To describe the effect of the solvent (water) on the structures and energies of the molecules, the polarized continuum (PCM) solvation model [16–18] was employed. All calculations were performed with the use of the GAUSSIAN 09 program package [19].

3. Results and discussion The stability constants (KITC) and binding enthalpies (DHITC) of interactions of the ODA and TDA ligands with the zinc(II) ion were obtained from ITC experiments by fitting isotherms (using nonlinear least-squares procedures) to a model that assumes a single set of identical binding sites. These thermodynamic parameters (KITC, DHITC) depend on the condition under which the ITC experiments were performed. In the Mes buffer solution with the pH value of 6, at 298.15 K, log KITC equals 3.78 (±0.04) and 3.25 (±0.01) whereas DHITC equals 3.35 (±0.07) [kcal/mol] and 2.74 (±0.03) [kcal/mol] for the Zn–ODA and the Zn–TDA complexes, respectively. Fig. 2 shows representative binding isotherms for interactions of the Zn(II) ion with the ODA and TDA ligands. To compare ITC parameters to log K and DH values obtained by other methods the pH values of solutions as well as the type of buffers used in calorimetric measurements should be taken into account. The following expression was used to calculate pH- and buffer-independent Zn2+–L (L = ODA2, TDA2) stability constant, KZnL (Eq. (1)):

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Fig. 2. Calorimetric titration isotherms of the binding interaction between Zn2+–ODA and Zn2+–TDA in 10 mM Mes buffer (pH 6) at 298.15 K.

K ZnL ¼ K ITC  aproton  abuffer

ð1Þ

where: KITC is the Zn–L conditional stability constant obtained directly from the ITC experiment, the coefficient aproton ¼ 1þ ½Hþ  K a2

þ 2

 þ K½H is the function of pH and pKa’s of the ligand (H2ODA, a2 K a1

H2TDA), the coefficient abuffer = 1 + K0 ZnMes[Mes] is a function of the metal–buffer (ZnMes) conditional stability constant (K0 ZnMes) and [Mes] is the concentration of the buffer solution. For calculations of aproton Ka values of the ligands obtained in our laboratory by PT method were used. They are as follows: pKa1 (H2ODA = H+ + HODA) = 2.90, pKa2 (HODA = H+ + ODA2) = 4.10 and pKa1 (H2TDA = H+ + HTDA) = 3.20, pKa2 (HTDA = H+ + TDA2) = 4.30. To find abuffer, it was assumed that in the system under study, 1:1 Zn–Mes complex is formed. This assumption was subsequently verified by the potentiometric titration. The following equilibrium models for interactions of the zinc ions with the Mes buffer were used to calculate the metal-buffer stability constant, KZnMes:

MesH ¼ Mes þ Hþ

The log K0 ZnMes equal 1.95 was used for calculations of the condition-independent stability constants KZnL. The log KZnL values calculated basing on Eq. (1) are equal to 4.06 (±0.01) and 3.54 (±0.02) for Zn–ODA and Zn–TDA complexes, respectively. These results are in good agreement, in the range of experimental error, with those obtained from PT measurements: log K = 4.04 (±0.10) and log K = 3.49 (±0.07) for Zn–ODA and Zn–TDA, respectively. The following equilibria were used to calculate the Zn2+–L (L = ODA2 or TDA2) stability constants, KZnL, from PT measurements:

H2 L ¼ LH þ Hþ

pK a1

LH ¼ L2 þ Hþ

pK a2

Zn2þ þ L2 ¼ ZnL

ZnL þ 2OH ¼ ZnLðOHÞ2 2 H2 O ¼ Hþ þ OH

log K ZnMes

ZnMesþ þ 2OH ¼ ZnMesðOHÞ2 H2 O ¼ Hþ þ OH

pK w

Table 1 Individual equilibria that contribute to the overall equilibrium for the formation of Zn2+–L complex in the Mes buffer solution (the charges of ions are omitted for the sake of brevity).

log K ZnMesðOHÞ2

No.

pK w

1 2 3 4 5

The logarithm of the KZnMes value equals 2.41 (±0.08). Next, the conditional (pH = 6) metal-buffer stability constant, KZnMes, was calculated according to Eq. (2) using pKa (MesH = Mes + H+) = 6.27 value [20].

1 1þ

log K ZnLðOHÞ2

pK a ¼ 6:27 ¼ const:

Zn2þ þ Mes ¼ ZnMesþ

K 0ZnMes ¼ K ZnMes 

log K ZnL

½Hþ  Ka

a

ð2Þ

Reactiona ZnMes = Zn + Mes HL = L + H H2L = L + 2H Zn + L = ZnL Mes + H = MesH

Coefficient

D Ho

aZnMes aHL aH2L

DHoZnMes DHoHL DHoH2L DHoZnL DHoMesH

1

aHL + 2aH2L

b

Equilibria are written in the direction in which the reaction occurs (1–3 are dissociations, 4 and 5 are associations). b DHo/kcal mol1 values are for the association reactions.

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The individual equilibria that contribute to the overall equilibrium, as well as the coefficients that indicate the percentage of the particular chemical species in the solution under experimental conditions are presented in Table 1. The overall reaction for the formation of the Zn2+–L complex is given by general Eq. (3).

Table 2 Thermodynamic parameters of Zn2+ binding to ligand (ODA2, TDA2) corrected for metal-binding interactions.

ð1  aZnMes ÞZn þ aZnMes ZnMes þ ð1  aHL  aH2L ÞL þ aHL HL ð3Þ

The coefficients in Table 1 are defined as follows (Eqs. (4)–(6)):

aZnMes ¼

K 0ZnMes ½Mes 1 þ K 0ZnMes ½Mes 

aHL ¼

Zn–TDA

4.06 (±0.01) 4.04 (±0.10)a 3.49 (±0.01) 30.30 (±0.01) 5.54 (±0.01)

3.54 (±0.02) 3.49 (±0.07)a 2.93 (±0.02) 25.99 (±0.02) 4.82 (±0.01)

a

The values obtained from PT measurements. The values determined using standard DG = RTlnK. b

thermodynamic

relationship:

ð4Þ 1

DHoZnNTA ðZn2þ þ NTA3 ¼ ZnNTA Þ ¼ 0:84 kcalmol

þ

½HL  ½H3 O  ¼ cH2 L K a2  aproton

ð5Þ

1

DHoHNTA ðNTA3 þ Hþ ¼ HNTA2 Þ ¼ 4:90 kcalmol

1

DHoH2NTA ðHNTA2 þ Hþ ¼ H2 NTA Þ ¼ þ0:80 kcalmol

þ 2

aH2 L ¼

Zn–ODA

logK

DH [kcal/mol] DSb [cal/molK] DGb [kcal/mol]

þ aH2L H2 L þ ðaHL þ 2aH2L ÞMes ¼ ZnL þ ðaHL þ 2aH2L ÞMesH

Parameter

½H2 L ½H3 O  ¼ c H3 L K a2  K a1  aproton

ð6Þ

Taking into consideration the fact that the heat absorbed or released during ITC experiments is equal to the sum of all the energetic effects corresponding to the particular equilibrium taking place in complexation process (Table 1), the pH- and buffer-independent enthalpy of the Zn2+–L complex formation (DHoZnL ) can be calculated from the equation based on Hess’s law (Eq. (7)).

DHITC ¼ aZnMes DHoZnMes  aHL DHoHL  aH2L DHoH2L þ ðaHL þ 2aH2L ÞDHMesH þ DHoZnL

ð7Þ DHoZnL ,

As seen (Eq. (7)), to determine the enthalpies of the proton association to the ligand (DHHL, DHH2L) as well as the enthalpy of metal-buffer complex formation (DHoZnMes ) need to be known. The enthalpies of the proton association to the ligand can be determined directly from calorimetric measurements. To do this, three experiments were carried out to determine: (1) DH1 – the enthalpy of titration of the mixture consisting of L2:HL or HL:H2L (15 mM:15 mM) with HClO4 (1 mM), (2) DH2 – the enthalpy of titration of water with HClO4 (1 mM), and (3) DH3 – the enthalpy of titration of the mixture L2:HL or HL:H2L (15 mM:15 mM) with water. Then, the DHHL and DHH2L were determined according to Eq. (8).

DHHL ðor DHH2L Þ ¼ DH1  DH2  DH3

ð8Þ

The values of proton-ligand association enthalpies obtained in the experiments are listed below:

DHðODA2 þ Hþ ¼ HODA Þ ¼ þ1:00ð0:01Þ½kcal=mol DHðHODA þ Hþ ¼ H2 ODAÞ ¼ þ0:23ð0:01Þ ½kcal=mol DHðTDA2 þ Hþ ¼ HTDA Þ ¼ þ0:86ð0:01Þ½kcal=mol DHðHTDA þ Hþ ¼ H2 TDAÞ ¼ þ0:36ð0:01Þ ½kcal=mol The DHoZnMes was calculated indirectly by the ITC displacement titration. In this experiment a weak ligand (Mes) is replaced from a coordination sphere of the Zn(II) ion by a strong-binding, competitive ligand H3NTA (H3NTA–nitrilotriacetic acid). The conditional stability constant (Zn2+ + NTA3 = ZnNTA), log KITC(ZnNTA), and the conditional enthalpy of Zn2+–NTA complex formation, DHITC(ZnNTA), determined from displacement titration are equal 6.23 (±0.13) and 0.35 (±0.01) [kcal/mol], respectively. The condition-independent enthalpy of Zn2+–NTA interactions, DHoZnNTA, as well as the enthalpies of the proton association to the ligand used in the calculations were taken from literature [21,22]. They are listed below:

1

DHoH3NTA ðH2 NTA þ Hþ ¼ H3 NTAÞ ¼ þ0:75 kcalmol

The pKa1, pKa2 and pKa3 values of H3NTA used for calculations were obtained from PT measurements. They are as follows: pKa1 = 2.28 (±0.17), pKa2 = 2.88 (±0.19) and pKa3 = 9.55 (±0.07). The enthalpy of metal-buffer complex formation DHoZnMes equals +0.19 (±0.01) [kcal/mol] and was determined based on Eq. (7) adopted to a triprotonated ligand (H3NTA), Eq. (9).

DHoZnMes ¼ ½DHITC  aHNTA DHoHNTA  aH2NTA DHoH2NTA  aH3NTA DHoH3NTA þ ðaHNTA þ 2aH2NTA þ 3a3NTA ÞDHMesH þ DHoZnNTA =aZnMes

ð9Þ

The conditional-independent formation constant, KZnL, and the enthalpies of the proton association to the ligand were used for calculating the thermodynamic parameters of the Zn2+–L interaction (Table 2). Since a d10 electron configuration is commonly assumed for Zn(II), no ligand field stabilization is expected, hence the stability of polycarboxylate zinc(II) complexes depends mainly on electrostatic forces and steric effects of the ligands. To compare the relative stabilities of mer and fac conformations of the ODA and TDA ligands in six-coordinated complexes with zinc as a central metal atom, ab initio calculations (employing the MP2 method with the 6–311++G(d,p) basis set) were performed. Fig. 3. depicts the optimized geometries of ZnODA(H2O)3 and ZnTDA(H2O)3 in their mer and fac configurations, while the selected optimized geometric parameters for mer and fac isomers as well as the corresponding geometric parameters obtained from X-ray measurements [3,8] (for comparison) are collected in Table 3. As seen (Table 3), there are some differences in bond distances and angles between experimentally obtained and theoretically predicted parameters. These differences are likely due to the fact that in crystals the molecules are packed tightly together as opposed to solutions where the motional degree of freedom of molecules and conformational flexibility of ligands are much higher. According to our findings the mer conformer of ZnODA is lower in the energy by ca. 1 kcal/mol (and 2 kcal/mol in water) only than its fac analog, which indicates the possible co-existence of both isomers in the gas phase (at room temperature) at equilibrium. However, in the case of the ZnTDA complex, the fac conformation of TDA is preferred as its energy is lower by ca. 4 kcal/mol (3 kcal/mol in water) than that of the mer configuration. For this reason, one should expect the fac isomer of ZnTDA(H2O)3 to predominate both in the gas phase and in water solutions. Moreover,

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167

Fig. 3. Optimized geometries of Zn ODA(H2O)3 and Zn TDA(H2O)3 in mer (left) and fac (right) configurations.

Table 3 Calculated and experimental values of selected geometrical parameters (distances in (Å), angles in (°)) of mer and fac isomers of ZnODA(H2O)3 and ZnTDA(H2O)3. Compound

Ab initio (MP2/6–311++G(d,p)) parameters

X-ray data (Refs. [3,8])

mer-isomer

fac-isomer

ZnODA

Zn–O(1) 1.959 Zn–O(5) 1.953 Zn–O(3) 2.121 O(1)–Zn–O(5) 161.4

Zn–O(1) 1.981 Zn–O(5) 1.974 Zn–O(3) 2.160 O(1)–Zn–O(5) 121.7

Zn–O(1) 2.085 Zn–O(5) 2.029 Zn–O(3) 2.116 O(1)–Zn–O(5) 151.5

ZnTDA

Zn–O(1) 1.943 Zn–O(4) 1.949 Zn–S 2.499 O(1)–Zn–O(4) 165.4

Zn–O(1) 1.984 Zn–O(4) 1.990 Zn–S 2.456 O(1)–Zn–O(4) 118.0

Zn–O(1) 2.029 Zn–O(4) 2.102 Zn–S 2.601 O(1)–Zn–O(4) 94.3

it is worth noting that the fac conformation of TDA is also observed in the crystal structure of ZnTDA [8]. Our findings indicate that the interactions of the Zn(II) ion with the ODA and TDA ligand are endothermic processes. The thermodynamic stabilities of the resulting complexes are comparable

and strongly depend on the entropy changes (Table 2). Positive values of DH indicate that the energy released by the metal–ligand binding formation is overcompensated by the energy used up in the dehydration process of [Zn(H2O)6]2+. This finding can be explained by comprising the Zn–O (carboxy) bond lengths in ZnODA and ZnTDA complexes and the Zn–O (water) bond lengths in the [Zn(H2O)6]2+ ion, that are shorter in the latter case [3,4,8,23]. The substitution of the ethereal donor atom in the ODA ligand by the sulfur donor atom (the TDA ligand) results in the decrease of DHZnL value (Table 2). The difference in DHZnL between these two complexes is not significant. This phenomenon can be explained by comparing the donor–acceptor bond lengths in both complexes. As seen in Table 3, the Zn–S(thioether) bond length in ZnTDA is much longer than the Zn–O(ethereal) bond length in ZnODA. The elongation of the Zn–S bond causes the ligand configuration to be not as stiffened as it would be if this bond was shorter. On the other hand, due to the smaller difference in the electronegativity between the atoms, the Zn–S(thioether) bond exhibits a more covalent character than the Zn–O(ethereal) bond. This phenomenon is probably the most important factor responsible for the release of a larger amount of energy during the formation of ZnTDA {DHZnODA > DHZnTDA, (see Table 2)} albeit the Zn–S(thioether) bond is longer than the Zn–O(ethereal) bond.

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4. Summary and conclusions The isothermal titration calorimetry (ITC) technique has successfully been applied to determine thermodynamic parameters (DH, DS) for the complexation reaction of the Zn2+ ions with oxydiacetate (ODA) and thiodiacetate (TDA) ligands. The DH and DS values were corrected for metal–ligand interactions as well as for pH values of solutions under which the ITC experiments were done. Based on ab initio calculations it can be concluded that in solution the ODA ligand may adopt both mer and fac configurations in coordination sphere of the zinc ion, whereas for the TDA ligand the fac conformation is preferred. The thermodynamic stabilities of the complexes are similar and ODA as well as TDA act as relatively weak ligands. The interactions of the Zn(II) ion with the ODA and TDA ligands are endothermic processes, however, the energy release during formation of the metal–ligand bonds is higher for ZnTDA. A stronger covalent character of the Zn–S(thioether) bond in comparison to the Zn–O(ethereal) bond is probably the main reason for this phenomenon. Thus, it might be postulated that in biological systems the less endothermic process will be observed for interactions of the zinc ion with a ligand comprising the larger number of sulfur-donor atoms than for oxygen-donor atoms despite the fact that the interatomic distances are expected to be longer when the sulfur atom is involved.

[2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19]

Acknowledgment

[20] [21]

This research was supported by the National Science Centre (grant 2011/03/D/ST5/05920).

[22]

References [1] B.L. Vallee, D.S. Auld, Biochemistry 29 (1990) 5647.

[23]

T. Dudev, C. Lim, J. Am. Chem. Soc. 122 (2000) 11146. R. Baggio, M.T. Garland, M. Perec, J. Chem. Soc., Dalton Trans. (1996) 2747. R. Baggio, M.T. Garland, M. Perec, Acta Crystallogr., Sect. C 59 (2003) m30. W.E. Hatfield, J.H. Helms, B.R. Rohrs, P. Singh, J.R. Wasson, R.R. Weller, Proc. Indian Acad. Sci. Chem. Sci. 98 (1987) 23. A. Grirrane, A. Pastor, E. Álvarez, C. Mealli, Inorg. Chem. Commun. 9 (2006) 160. S.H. Whitlow, Acta Crystallogr., Sect. B 31 (1975) 2531. M.G.B. Drew, D.A. Rice, C.W. Timewell, J. Chem. Soc., Dalton Trans. (1975) 144. A. Grirrane, A. Pastor, E. Álvarez, C. Mealli, A. Ienco, P. Rosa, A. Galindo, Eur. J. Inorg. Chem. (2007) 3543. A. Panuszko, P. Bruz´dziak, J. Zielkiewicz, D. Wyrzykowski, J. Stangret, J. Phys. Chem. B. 113 (2009) 14797. I. Brandariz, J. Barriada, T. Vilarino, M.S. de Vicente, Monatsh. Chem. 135 (2004) 1475. J. Kostrowicki, A. Liwo, Comput. Chem. 11 (1987) 195. J. Kostrowicki, A. Liwo, Talanta 37 (1990) 645. A.D. McLean, G.S. Chandler, J. Chem. Phys. 72 (1980) 5639. R. Krishnan, J.S. Binkley, R. Seeger, J.A. Pople, J. Chem. Phys. 72 (1980) 650. S. Miertus, E. Scrocco, J. Tomasi, Chem. Phys. 55 (1981) 117. S. Miertus, J. Tomasi, Chem. Phys. 65 (1982) 239. M. Cossi, V. Barone, J. Cammi, Chem. Phys. Lett. 255 (1996) 327. M.J. Frisch, G.W. Trucks, H.B. Schlegel, G.E. Scuseria, M.A. Robb, J.R. Cheeseman, G. Scalmani, V. Barone, B. Mennucci, G.A. Petersson, H. Nakatsuji, M. Caricato, X. Li, H.P. Hratchian, A.F. Izmaylov, J. Bloino, G. Zheng, J.L. Sonnenberg, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O. Kitao, H. Nakai, T. Vreven, J.A. Montgomery Jr., J.E. Peralta, F. Ogliaro, M. Bearpark, J.J. Heyd, E. Brothers, K.N. Kudin, V.N. Staroverov, R. Kobayashi, J. Normand, K. Raghavachari, A. Rendell, J.C. Burant, S.S. Iyengar, J. Tomasi, M. Cossi, N. Rega, N.J. Millam, M. Klene, J.E. Knox, J.B. Cross, V. Bakken, C. Adamo, J. Jaramillo, R. Gomperts, R.E. Stratmann, O. Yazyev, A.J. Austin, R. Cammi, C. Pomelli, J.W. Ochterski, R.L. Martin, K. Morokuma, V.G. Zakrzewski, G.A. Voth, P. Salvador, J.J. Dannenberg, S. Dapprich, A.D. Daniels, O. Farkas, J.B. Foresman, J.V. Ortiz, J. Cioslowski, D.J. Fox, GAUSSIAN 09, Revision B.01, GAUSSIAN Inc., Wallingford, CT, 2009. R.N. Goldberg, N. Kishore, R.M. Lennen, J. Phys. Chem. Ref. Data 31 (2002) 231. L.G. Sillen, A.E. Martel, Stability Constants of Metal–Ion Complexes, The Chemical Society, London, Great Britain, 1966. J.J. Christensen, M.I. Reed, Handbook of Metal Ligand Heats and Related Thermodynamic Quantities, Marcel Dekker Inc., New York, 1983. M. Pavlov, P.E.M. Siegbahn, M. Sandström, J. Phys. Chem. A 102 (1998) 219.