Speciation of 2-Hydroxybenzoic Acid with Calcium(II), Magnesium(II

Article pubs.acs.org/jced

Speciation of 2‑Hydroxybenzoic Acid with Calcium(II), Magnesium(II), and Nickel(II) Cations in Self-Medium Emilia Furia,* Anna Napoli, Antonio Tagarelli, and Giovanni Sindona Dipartimento di Chimica e Tecnologie Chimiche, Università della Calabria, via P. Bucci 87036 Arcavacata di Rende (CS), Italy ABSTRACT: Equilibria occurring between 2-hydroxybenzoic acid (salicylic acid, H2L) and calcium(II), magnesium(II), and nickel(II) have been studied by potentiometric titrations with glass electrodes at 298.15 K in self-medium constituted of 1.05 mol·kg−1 NaHL. The self-medium composition can be considered constant since only a small amount of ligand HL− is transformed in the reaction products. The hydrogen ion concentration was varied from 10−2 mol·kg−1 to incipient precipitation of a basic salt of each metal, whose formation depends on the particular metal. The total metal concentration, CM, ranged from 3.05·10−3 mol·kg−1 to 10.01·10−3 mol·kg−1. The higher CM value was imposed in order to avoid changes in the ionic medium and in the activity coefficients. The experimental data have been explained by assuming the formation of the complexes: CaL; MgL, MgL22−, and MgL34−; NiL and NiL22−.

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INTRODUCTION

MATERIALS AND METHODS Instrumentation. The cell arrangement, located in a thermostat kept at (298.1 ± 0.1) K, was as previously described.5 The test solutions, stirred during titrations, were purified with a slow stream of nitrogen gas which was passed before through four bottles (a−d) containing: (a) 0.96 mol·kg−1 NaOH, (b) 0.94 mol·kg−1 H2SO4, (c) twice-distilled water, and (d) 1.05 mol·kg−1 NaHL. Reagents and Analysis. Perchloric acid stock solutions were prepared and standardized as previously described.1 A sodium perchlorate stock solution was prepared and standardized according to Biedermann.16 Sodium hydroxide titrant solutions were prepared and standardized as described in a former paper.4 Sodium hydrogen salicylate stock solutions were prepared from the commercial product Aldrich p.a. dried at 378 K. The purity of the salt was assumed to be 99.5 % as stated on the label. Calcium(II) and magnesium(II) perchlorate stock solutions were prepared and standardized according to Bottari and Porto.17 Nickel(II) perchlorate stock solution was prepared and standardized as reported by Biedermann and Ferri.18 All solutions were prepared with twice-distilled water. Potentiometric Measurements. Equilibria occurring in solutions containing salicylate and alternatively calcium(II), magnesium(II), and nickel(II) perchlorate solutions were studied at 298.15 K as reported in a previous paper.1 The silver reference electrode, RE, and test solution, TS, had these general compositions, respectively: RE = Ag/AgCl/1.05 mol·kg−1 NaHL, equilibrated with AgCl/1.05 mol·kg−1 NaHL;

2-Hydroxybenzoic acid (salicylic acid, H2L) is an outstanding antirheumatic and antifungal substance. Moreover, it represents the simplest model for the humic acids present in soil. Thus, the complexation equilibria of H2L with several cations has been studied by potentiometric as well as spectrophotometric methods.1−13 However, despite the relevance of the ligand, only few quantitative data regarding equilibria occurring between H2L and Ca(II), Mg(II), and Ni(II) ions in aqueous media have been reported, and the results are often contradictory (Table 1). Taking into account the discrepancy of the results of previous works and in consideration of the numerous biological and medical applications of salicylic acid, it seems important to reconsider these systems in order to determine the composition of the reaction products. Additional information about complexes formed between ligand and metals may be obtained by working with an ionic medium in which one of the reagents represents the anion or the cation.14 This approach is known as the self-medium method, as distinguished from the inert medium method.15 The purpose of this work was therefore to determine equilibrium constants between salicylic acid and calcium(II), magnesium(II), and nickel(II) ions at 298.15 K by potentiometric titrations with glass electrodes in self-medium constituted of 1.05 mol·kg−1 NaHL, in which Na+ was the inert counterion. The equilibrium constants in this self-medium may be expected to differ from those in an inert medium of the same concentration of inert counterion. In order not to change the ionic medium and the activity factors it was necessary to keep the added amount of metals fairly low. © 2013 American Chemical Society

Received: January 31, 2013 Accepted: April 11, 2013 Published: April 24, 2013 1349

dx.doi.org/10.1021/je400105c | J. Chem. Eng. Data 2013, 58, 1349−1353

Journal of Chemical & Engineering Data

Article

Table 1. Results from Previous Works. βn Were the Equilibrium Constants According to General Equilibrium M2+ + nHL− ⇆ MLn(2−2n) + nH+ log βn

metal ion

method

T/K

medium

Ca2+

spectrophotometric measurements glass electrode spectrophotometric measurements glass electrode glass electrode

298.15 310.15 298.15 310.15 298.15

0.5 mol·dm−3 NaCl 0.15 mol·dm−3 NaClO4 0.5 mol·dm−3 NaCl 0.15 mol·dm−3 NaClO4 0.1 mol·dm−3 NaNO3

glass electrode glass electrode

308.15 303.15

0.1 mol·dm−3 NaNO3 0.5 mol·dm−3 KNO3

spectrophotometric measurements glass electrode

298.15 298.15

0.15 mol·dm−3 NaClO4 0.1 mol·dm−3 NaClO4

glass electrode

293.15

0.1 mol·dm−3 KCl

Mg2+ Ni2+

TS = CM mol·kg−1 M(ClO4)2, CA mol·kg−1 H2L, CB mol·kg−1 NaOH, 1.05 mol·kg−1 HL−, (1.05 + CB) mol·kg−1 Na+. The total metal concentration, CM, ranged from 3.05·10−3 mol·kg−1 to 10.01·10−3 mol·kg−1. At lower C M no reliable results were obtained, probably due to the influence of small analytical errors. Higher CM values were not included in order to avoid large changes in the medium. The hydrogen ion concentration was varied from 10−2 mol·kg −1 to incipient precipitation of a basic salt of each metal, whose formation depends on the particular metal. Since the effects of composition changes on activity coefficients can be considered negligible, the EMF of cell (G) can be written, in mV, at the temperature of 298.15 K, as follows: E = E° + 59.16 log[H+] + 59.16 log fh + Ej

K a1 = [H+][HL−]/[H 2L]

[HL−] = 1.05 − H

references 6 7 6 7 8 9 10 11 12 13

(5)

In order to calculate E°′ the acidic constant Ka1 is needed; results from other investigations1 indicate that the values of the Ka1, determined at 298.15 K in different ionic media of various molality range from 10−3.03, at zero ionic strength, to 10−3.19 in 3.05 mol·kg −1 NaClO4. Since the Ka1 term in eq 5 amounts to about 0.1 mV, an average value of log Ka1 = −3.1 ± 0.2 is sufficient to estimate reliable E°′ values. The error introduced by this approximation is below the uncertainty of the EMF measurements, that is, ± 0.2 mV. In the second part of each titrations for the investigation of complexes between ligand and metal ions, alkalification was achieved by adding NaOH. When E°′ was known, it was possible to calculate q in the presence of metal ions; combining the q values with the analytical data we can write the quantities reported in eqs 6 and 7

(1)

[H 2L] = q(1 + H )/(2q + 1)

(6)

y = [H 2L]/[HL−]2 = q(2q + 1)/(1 + H )

(7)

from which the function Z, defined as average number of salicylate split off for metal ions (eq 8), was calculated Z = ([H 2L] − H − Ky−1)/CM

(8)

For the calculation of Z the required value of K (which is equal to Ka2/Ka1 ratio) was determined by acid−base titrations with cell G

(2)

RE/TS/glass electrode

(G)

in the absence of metal ions and in the range 0.01 ≤ H ≤ 0.03 mol·kg−1 as follows:

(3)

and

−10.19 −8.71 −8.48 −7.84 −4.95 −13.90 −6.07 −6.87 −13.47 0.63 −6.04 −14.22 −6.66 −15.75

+ 59.16 log[1 + K a1/(1.05 − H )]

log K = (E − E°′)/59.16 + log[q − H(1 + 2q)

where E° = (E°′ + 59.16 log Ka1). In the absence of metal ions and in sufficiently acid solutions, that is, H > 2·10−3 mol·kg−1 in which H is the analytical excess of H+, it was possible to write the stoichiometric conditions (eq 4) [H 2L] = H − [H+]

= = = = = = = = = = = = = =

E°′ = E − 59.16 log H + 59.16 log(1.05 − H )

Taking into account eq 2 it was possible to replace in eq 1 the hydrogen ion concentration [H+] with q = [H2L]/[HL−]

E = E°′ + 59.16 log q

β1 β1 β1 β1 β1 β2 β1 β1 β2 β1 β1 β2 β1 β2

According to eqs 2 and 4 it is possible to write the expression of E°′ (eq 5) as

where E° is constant in each series of measurements and f h is the activity coefficient of H+. The medium components were substituted by the other reacting species by no more than 6 %; therefore according to Biedermann and Sillén15 f h can be considered unitary. Ej = −j [H+] stands for the liquid junction potential between TS and 1.05 mol·kg−1 NaHL.15 The constant j may be assumed to be the same in all experiments, and its value was taken from literature.1 In the range of hydrogen concentration investigated Ej is at most ± 0.2 mV, which represents the reproducibility of the electromotive force measurements, and for this reason its values can be neglected. Each titration was divided in two parts; in the first part, E° was determined in the absence of metal ions as reported in a previous paper.2 For the calculation of standard potential of the cell it is necessary to consider the first acidic constant of salicylic acid (eq 2) H 2L ⇄ HL− + H+

log log log log log log log log log log log log log log

/(1 + H )]

(9)

It was necessary to measure K since it was not possible to calculate it from the Ka2/Ka1 ratio using literature values for Ka2 and Ka1, that are not obtained in our experimental conditions (i.e., self-medium 1.05 mol·kg−1 NaHL). The measured value

(4) 1350



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for log K is −9.02 ± 0.03, in which the uncertainty represents 3σ.

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RESULTS AND DISCUSSION The general equilibrium for all of the three systems can be schematically written as follows: M2 + + 2nHL− ⇄ MLn(2 − 2n) + nH+*βn

(10)

The experimental data were processed by graphical as well as by numerical procedures. The graphical methods consist essentially of the comparison of experimental plots with model functions.19 For the graphical approach it was necessary to evaluate for each point of the titration the experimental data which gave CA, CM, y, and Z by eq 8. The experimental function Z (log y) for the different systems was reported in Figures 1, 2,

Figure 3. Z (log y) for the system Ni2+-salicylate. The continuous curve was constructed with the equilibrium constants given in Table 7. The symbols refer to CM, mol·kg−1: △ = 10.0·10−3; ○ = 5.0·10−3.

Table 2. Summary of the Relevant Data Taken in Two Titrations for the System Ca2+-Salicylate self-medium concentration 1.05 mol·kg

−1

NaHL

CM/(mol·kg−1)

pH range

5.07·10−3 3.05·10−3

5.8 to 6.8 5.7 to 7.2

The function Z (log y) was graphically represented in Figure 1. In order to apply the graphical method it was been formulated a complexation model to fit the experimental data. The preliminary hypothesis was assumed with the formation of the species CaL according to equilibrium 12:

Figure 1. Z (log y) for the system Ca2+-salicylate. The continuous curve was constructed with the equilibrium constant given in Table 3. The symbols refer to CM, mol·kg−1: △ = 5.07·10−3, ○ = 3.05·10−3.

Ca 2 + + HL− ⇄ CaL + H+

Table 3. Best Set of log *β1 for the System Ca2+-Salicylate Obtained by Graphical as well as by Numerical Procedures

and 3. In the numerical treatment, the most probable n value and the corresponding constants *βn according to general equilibrium 10 were computed by the least-squares program Superquad20 to seek the minimum of the function:

∑ (Ei obs − Ei cal)2

(12)

The validity of this assumption has been verified from the graph of Z as a function of log y (Figure 1). Experimental data obtained by titrations performed at different concentrations of metal (i.e., 5.07·10 −3 and 3.05·10 −3 mol·kg −1 ) were completely overlaid to the normalized curve19 in which Z tends to 1 (Figure 1). This experimental evidence confirms that the data are fully explained with equilibrium 12. In the position of best fit it was possible to calculate the values of the constant *β1 which was reported in Table 3. The uncertainty on the complexation constant was

Figure 2. Z (log y) for the system Mg2+-salicylate. The continuous curve was constructed with the equilibrium constants given in Table 5. The symbols refer to CM, mol·kg−1: ○ = 5.01·10−3; △ = 10.01·10−3.

U=

*β1

species

(log *β1 ± σ)graph

(log *β1 ± 3σ)num

CaL

−7.6 ± 0.1

−7.40 ± 0.02

evaluated taking into account the shift along X-axes that still gave an acceptable fit between the theoretical curve and the experimental points. The main hydrolysis product of Ca2+ has been established as Ca(OH)+ by literature data;21 in the numerical evaluation of complexation constants between Ca2+ and salicylate ions, the equilibrium constant of Ca(OH)+ was invariable because its value was well-known from the literature.21 Therefore the numerical treatment was started assuming the presence of Ca(OH)+ and of the complexes CaL, agreeing with the graphical evaluation. This model with a standard deviation, σ, 0.284 mV, which was comparable with the experimental uncertainty, was assumed as the best describing the data, also in consideration that no other species containing the ligand had produced a significant lowering of the

(11) +

where Eobs = E°′ + 59.16 log [H ] at 298.15 K, while Ecal is a value calculated for a given set of parameters. Calculations of the chi-squared statistic have been considered to test the fit between a theoretical frequency distribution and a frequency distribution of observed data. In the numerical treatment the value of the first acidic constant of salicylic acid (eq 2) has been maintained invariant. The single systems were treated separately. System with Ca2+-Salicylate. The data comprise two titrations with 75 data points. A summary of the relevant data taken in all titrations is reported in Table 2. 1351



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function U. At this model a χ2 value of 12.13 and U of 8.875 mV2 were associated. Results of numerical evaluation for the constant *β1 are reported in Table 3. The agreement between graphical and numerical values of log *β1 is satisfactory. System with Mg2+-Salicylate. The data comprise two titrations with 60 data points. A summary of the relevant data taken in all titrations is reported in Table 4.

System with Ni2+-Salicylate. The data comprise two titrations with 53 data points. A summary of the relevant data taken in all titrations is reported in Table 6. Table 6. Summary of the Relevant Data Taken in Two Titrations for the System Ni2+-Salicylate

Table 4. Summary of the Relevant Data Taken in Two Titrations for the System Mg2+-Salicylate self-medium concentration

CM/(mol·kg−1)

pH range

1.05 mol·kg−1 NaHL

10.01·10−3 5.01·10−3

3.8 to 5.6 5.3 to 6.8

*β1 *β2

(14)

Mg 2 + + 3HL− ⇄ MgL3 4 − + 3H+

*β3

(15)

(log *βn ± σ)graph

(log *βn ± 3 σ)num

−5.79 ± 0.01 −12.65 ± 0.05 −19.85 ± 0.05

−5.803 ± 0.003 −12.60 ± 0.02 −19.79 ± 0.03

4.0 to 5.2 3.2 to 5.9

*β1

(16)

*β2

(17)

The validity of this assumption has been confirmed from the graph of Z as a function of log y, which manifests that experimental data, obtained by titrations performed at different concentrations of Ni2+ (i.e., 5.0·10−3 mol·kg−1 and 10.0·10−3 mol·kg−1), were completely overlaid to the normalized curve19 in which Z tends to 2 (Figure 3). In the position of best fit it was possible to evaluate the values of the constants *β1 and *β2, which are given in Table 7 Table 7. Best Set of log *β1 and log *β2 for the System Ni2+Salicylate Obtained by Graphical as well as by Numerical Procedures

Table 5. Best Set of log *β1, log *β2, and log *β3 for the System Mg2+-Salicylate Obtained by Graphical as well as by Numerical Procedures species

10.0·10−3 5.0·10−3

Ni 2 + + 2HL− ⇄ NiL 2 2 − + 2H+

This hypothesis has been confirmed from the graph of Z as a function of log y (Figure 2). Experimental data acquired by potentiometric titrations achieved at different concentrations of Mg2+ (i.e., 5.01·10−3 mol·kg−1 and 10.01·10−3 mol·kg−1) were completely overlaid to the normalized function19 in which Z tends to 3 (Figure 2); therefore, experimental data are fully explained with the equilibria reported above. In position of best fit it was possible to extrapolate the values of the constant *β1, *β2, and *β3 with the corresponding uncertainties (Table 5). As concerns the numerical treatment,

MgL MgL22− MgL34−

pH range

1.05 mol·kg−1 NaHL

Ni 2 + + HL− ⇄ NiL + H+

(13)

Mg 2 + + 2HL− ⇄ MgL2 2 − + 2H+

CM/(mol·kg−1)

The function Z (log y) for the system Ni2+-salicylate was graphically represented in Figure 3. In order to apply the graphical method it was been formulated a complexation model to fit the experimental data. The preliminary hypothesis was assumed with the formation of the two complexes NiL and NiL22− according to equilibria 16 and 17:

The function Z (log y) for the system Mg2+-salicylate is graphically represented in Figure 2. In order to apply the graphical method it was been chosen a complexation model to fit the experimental data. The formation of the three complexes MgL, MgL22−, and MgL34− according to eqs 13 to 15 was assumed: Mg 2 + + HL− ⇄ MgL + H+

self-medium concentration

species

(log *βn ± σ)graph

(log *βn ± 3 σ)num

NiL NiL22−

−4.21 ± 0.02 −10.06 ± 0.03

−4.12 ± 0.05 −9.90 ± 0.05

with the corresponding uncertainties. The principal hydrolysis product of Ni2+ has been established as Ni(OH)+, Ni2(OH)3+ and Ni4(OH)44+ by literature data.23 In the numerical processing the equilibrium constant for the only hydrolysis product Ni(OH)+, which was considered to prevail in our experimental condition, that is, low concentration of nickel, has been taken from the literature.23 The numerical treatment was started assuming the presence Ni(OH)+ and of the two mononuclear species NiL and NiL22− according to the graphical results. This model with a standard deviation 0.3028 mV, which slightly exceeds the experimental error, was assumed as the best to describe the data. A χ2 value of 23.20 and U of 5.4333 mV2 were associated with this theoretical model. Results of numerical evaluation of the constants *β1 and *β2 are given in Table 7. The agreement between graphical and numerical values is just acceptable. This dissimilarity may be due to low percentage of these species in our experimental conditions, as evident in the distribution diagram of Ni2+ in the different species, which was constructed with the refined equilibrium constants reported in Table 7 (Figure 4). This paper reports a thermodynamic study on the interactions between salicylic acid, H2L, and calcium(II), magnesium(II), and nickel(II) ions at 298.15 K in self-medium constituted of 1.05 mol·kg−1 NaHL. The self-medium method

the main cationic hydrolysis product of Mg2+ has been established as Mg(OH)+ and Mg4(OH)44+ by literature data.22 In the numerical calculation of complexation constants the equilibrium constant for the hydrolysis product Mg(OH)+, which was prevailing in our experimental condition, that is, the low concentration of metal ion, was invariable because it was well-known from the literature.22 The numerical treatment was started assuming the presence Mg(OH)+ and of the three mononuclear complexes MgL, MgL22−, and MgL34− according to the graphical results. This model with a standard deviation 0.3172 mV, which marginally exceeds the experimental error, was assumed as the best describing the data. A χ2 value of 23.20 and U of 5.4333 mV2 were associated with this theoretical model. Results of numerical evaluation of the constants *β1, *β2, and *β3 are given in Table 5. The accordance between graphical and numerical values is noteworthy. 1352



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(7) Arena, G.; Kavu, G.; Williams, D. R. Metal-ligand complexes involved in rheumatoid arthritis-V: Formation constants for calcium(II)-, magnesium(II)- and copper(II)-salicylate and acetylsalicylate interactions. J. Inorg. Nucl. Chem. 1978, 40, 1221−1226. (8) Khalil, M. M.; Radalla, A. M. Binary and ternary complexes of inosine. Talanta 1998, 46, 53−61. (9) Krishnamoorthy, C. R.; Sunil, S.; Ramalingam, K. The effect of ligand donor atoms on ternary complex stability. Polyhedron 1985, 4, 1451−1456. (10) El-Ezaby, M. S.; El-Khalafawy, T. E. Complexes of vitamin B6IX Ternary complexes of Ni(II), Cu(II) and Zn(II) with pyridoxamine and salicylic acid. J. Inorg. Nucl. Chem. 1981, 43, 831−837. (11) Yousif, Y. Z.; Al-Imarah, F. J. M. Spectrophotometric determination of the stability constants of the complexes between the Ni2+ ion and some substituted salicylic acids in aqueous solution. J. Inorg. Nucl. Chem. 1980, 42, 779−783. (12) Abbas Abbasi, S.; Bhat, B. G.; Singh, R. S. Mixed ligand complexes involving hydroxamic acids I. Complexes of benzohydroxamic acid in aqueous solutions. Inorg. Nucl. Chem. Lett. 1976, 12, 391−397. (13) Perrin, D. D. Stability of Metal Complexes with Salicylic Acid and Related Substances. Nature 1958, 182, 741−742. (14) Hietanen, S.; Sillén, L. G. Studies on the hydrolysis of metal ions 22. Equilibrium studies in Self-Medium; application to the hydrolysis of Th4+. Acta Chem. Scand. 1959, 13, 533−550. (15) Biedermann, G.; Sillén, L. G. Studies on the hydrolysis of metal ions. IV. Liquid junction potentials and constancy of activity factors in NaClO4−HClO4 ionic medium. Ark. Kemi 1953, 5, 425−440. (16) Biedermann, G. Study on the hydrolysis equilibria of cations by emf methods. Svensk. Kem. Tidsk. 1964, 76, 362. (17) Bottari, E.; Porto, R. Complex formation between glycine and magnesium(II), calcium(II), and iron(II) at 25 °C in 3.00 M NaClO4. Monatsh. Chem. 1982, 113, 1245−1252. (18) Biedermann, G.; Ferri, D. On the preparation of metal perchlorate solutions. Chem. Scripta (Stockholm) 1972, 2, 57−61. (19) Sillén, L. G. Some Graphical Methods for Determining Equilibrium Constants. II. On “Curve-fitting” Methods for Twovariable Data. Acta Chem. Scand. 1956, 10, 186−202. (20) Gans, P.; Sabatini, A.; Vacca, A. SUPERQUAD: an improved general program for computation of formation constants from potentiometric data. J. Chem. Soc., Dalton Trans. 1985, 6, 1195−1200. (21) Bates, R. G.; Bower, V. E.; Canham, R. G.; Prue, J. E. The dissociation constant of CaOH+ from 0° to 40 °C. Trans. Faraday Soc. 1959, 55, 2062−2068. (22) Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; A WileyInterscience Publication, New York, 1976. (23) Bolzan, J. A.; Jauregui, E. A.; Arvia, A. J. Hydrolytic equilibria of metallic ions, III. The hydrolysis of Ni(II) in NaClO4 solution. Electrochim. Acta 1963, 8, 841−845. (24) Bottari, E.; Jasionowska, R. Equilibria between cations and glycine in self medium. Ann. Chim. 1980, 70, 569−584. (25) Bottari, E.; Porto, R. Equilibria between cations and serine in self medium. Ann. Chim. 1986, 76, 283−291. (26) Pettit, G. IUPAC: Stability Constants Data Base; Academic Software: Otley, U.K., 1995.

Figure 4. Distribution of Ni(II) species for CM = 10.0·10−3 mol·kg−1. 1: Ni2+; 2: NiL; 3: NiL22−; 4: NiOH+.

promotes formation of complexes MpLq with low values for p or q, and information about such species may be attained with a greater certainty than from work with inert media. It might be held as a disadvantage that the information on the formulas of the species which is obtained with the self-medium method is more restricted than that with the inert medium method, which is more restricted than that obtained from dilute solutions. On the other hand restricted information on equilibrium constants might have been unavailable by the other methods. From equilibrium studies in an inert ionic medium such as NaClO4 it is possible to deduce the formulas and related equilibrium constants of polynuclear species, information which could not have been obtained from studies in dilute solutions as it is not possible to work with high concentration levels in the reagents. The equilibrium constants between salicylic acid and, alternatively, calcium(II), magnesium(II), and nickel(II) ions obtained in this study in self-medium are in each case higher than those reported in Table 1. It has been already reported that the values of the formation constants for the complexation of various cations obtained in self-medium are systematically higher than those obtained in an inert medium.24−26

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +39-0984-492831. Fax: +390984-492044; +39−0984−493307. Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Furia, E.; Porto, R. The effect of ionic strength on the complexation of copper (II) with salicylate ion. Ann. Chim. 2002, 92, 521−530. (2) Furia, E.; Porto, R. Equilibria occurring between beryllium (II) and salicylate ions. Ann. Chim. 2003, 93, 1037−1043. (3) Furia, E.; Porto, R. The hydrogen salicylate ion as ligand. Complex formation equilibria with dioxouranium (VI), neodymium (III) and lead (II). Ann. Chim. 2004, 94, 795−804. (4) Porto, R.; De Tommaso, G.; Furia, E. The second acidic constant of salicylic acid. Ann. Chim. 2005, 95, 551−558. (5) Furia, E.; Sindona, G. Interaction of iron (III) with 2hydroxybenzoic acid in aqueous solutions. J. Chem. Eng. Data 2012, 57, 195−199. (6) Dahlund, M.; Olin, Å. A spectrophotometric study of the complexation of olsalazine and salicylic acid with calcium and magnesium. Acta Chem. Scand. 1990, 44, 321−327. 1353


Speciation of 2-Hydroxybenzoic Acid with Calcium(II), Magnesium(II

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