3. Different complexes between Fe(III) and biological active ligands

### 3.1. Studies of binary and ternary complexes of sulfamethoxazole (SMZ) and glycine with metal ions

Sulfamethoxazole (4-amino-N-(5-methyl-3-isoxazolyl)-benzenesulfonamide (SMZ)) is the most predominant sulfonamide in human medicine. Sulfonamides are synthetic antimicrobial agents derived from sulfanilamide, whose antibacterial activity was discovered in the early 1930s by Domagk and Tréfouel [31–33].

#### 3.1.1. Stability constants of ternary complexes (metal-SMZ-Gly)

The total stability constant, βlpqr, can be calculated from the equation by:

Table 1. The programs commonly used for calculating formation equilibrium constants.

V, potentiometric experiments; A, spectrophotometric experiments; E, ESR; N, NMR.

Additional data used in calculations are taken from different sources.

30 Descriptive Inorganic Chemistry Researches of Metal Compounds

1

physiological pH range. In this chapter, we reviewed several iron complexes.

3. Different complexes between Fe(III) and biological active ligands

3.1. Studies of binary and ternary complexes of sulfamethoxazole (SMZ) and glycine

Sulfamethoxazole (4-amino-N-(5-methyl-3-isoxazolyl)-benzenesulfonamide (SMZ)) is the most predominant sulfonamide in human medicine. Sulfonamides are synthetic antimicrobial agents

½L1�p½L2�q½H�

where M, L1, L2, and H stand for metal ion, ligand (1), ligand (2), and proton, respectively. For

System Data type<sup>a</sup> Reference

MINIQUAD V [19] MINIQUAD75 V [20] TITAN V [25] SCOGS2a V [21] D SCOGS2ba V [26] MINIQUAD V, A [22] PSEQUAD V, A [23] SPECFIT A(E) [18] PKAS V [15] HYDROUAD V [24] STAR A [29] HYPNMR N [30]

Pettit program computes speciation based on the concentrations of metal ions and the complexing species. This program specifies a certain pH range. Then the former calculates the species distribution of a certain series of complexes and plots it. We enter some data such as the total concentrations of metal and ligand ions and pH range. After that the best-fit set of β values will be used later to compute the equilibrium concentrations of those complex species over the pH range which we have specified before. We can use this program for all types of complexes: mixed complexes, protonated, hydroxo, and polynuclear species. The program produces a graphical recording of the most predominant complex species at any pH and the

<sup>r</sup>ðfor simplicity charges are omitted<sup>Þ</sup> <sup>ð</sup>3<sup>Þ</sup>

β1pqr ¼ ½ MÞlðL1ÞpðL2ÞqðHÞr�=½M�

OH�, the coefficient (r) for H ¼ �1.

2.3. Calculation of speciation

a

with metal ions

Different metal ions Ti(II), Zr(IV), Sr(II), Al(III), Cr(III), Fe(III), Th(IV), Pb(II), La(III), and Co(II) were selected to make further investigation to elucidate the interaction of these metal ions with solution of SMZ and Gly (mixed ligand complexes). The potentiometric equilibrium measurements were made, at constant ionic strength I ¼ 0.1 M NaClO4 at 25 � 0.1�C, for the interaction of SMZ and the selected 10 metal ions, with biologically important secondary ligand glycine (Gly) in a (1:1:1) molar ratio (1 � <sup>10</sup>–<sup>3</sup> M for each). The solutions were titrated pH-metrically against standard carbonate-free NaOH solution, as illustrated in Figure 2.

Figure 2 represents the titration curves for the metal-SMZ-Gly system studied. It is observed that the metal ion-SMZ titration curve (c) diverges from the SMZ curve (b) at variable pH values (pH ≈ 2.8 for Fe(III), pH ≈ 3.5 for La(III), pH ≈ 4.2 for Th(IV), pH ≈ 5.5 for Zr(IV), pH ≈ 4.5 for Al(III), and pH ≈ 6.06 for Co(II)) denoting the formation of metal ions-SMZ binary complexes. For the titration curves of the ternary systems studied, it can be observed that the curves (c) and (f), however, overlap with each other at lower pH values in the case of Fe(III) and La (III), whereas that for Sr(II), Pb(II), Cr(III), and Ti(II) are well separated. This indicates the formation of metal ions-SMZ-Gly ternary complexes at lower pH values, which can be considered as an evidence for the formation of protonated SMZ mixed ligand complex. The stability constants of the ternary metal ion complexes containing SMZ and Gly were calculated from Eqs. (4) and (5),

$$\text{M} + \text{SMZ} \leftrightarrow \text{M} \ (\text{SMZ}) \tag{4}$$

$$\mathbf{K}\_{\mathbf{M}(\mathbf{SMZ})(\text{Gly})}^{\mathbf{M}(\mathbf{SMZ})} = \frac{\left[\text{M}(\text{SMZ})(\text{Gly})\right]}{\left[\text{M}(\text{SMZ})(\text{Gly})\right]} \tag{5}$$

using the data obtained from potentiometric titrations (I ¼ 0.1M NaClO4 at 25 � 0.1�C).

Figure 2. Potentiometric curves of SMZ in 0.1 M NaClO4 at 25 � 0.1�C: (a) 0.01 M HClO4, (b) a þ 0.001 M SMZ, (c) b þ 0.001 M Sr (II), (d) b þ 0.001 M Pb(II), (e) bþ 0.001 M Co(II), (f) b þ 0.001 M Fe(III), and (g) b þ 0.001 M Al (III).

#### 3.2. Ternary complexes of iron(III)-glycine( Gly)-nitrilotriacetate (NTA) system

Electrochemical measurements of the dissolved iron(III)-Gly-NTA mixed ligand system in the 0.1 mol�dm–<sup>3</sup> NaClO4 aqueous solution were performed at pH <sup>¼</sup> 8.0 � 0.1and 25 � <sup>1</sup>�C, using differential pulse cathodic voltammetry (DPCV), cyclic voltammetry (CV), and direct current (d. c.) polarography. Iron(III) concentrations were varied from 2.5 � <sup>10</sup>–<sup>5</sup> to 6 � <sup>10</sup>–<sup>4</sup> mol�dm–<sup>3</sup> , NTA total concentrations varied from 2 � <sup>10</sup>–<sup>5</sup> to 1 � <sup>10</sup>–<sup>3</sup> mol�dm–<sup>3</sup> and glycine total concentrations were 0.2, 0.02, and 0.002 mol�dm–<sup>3</sup> . Figure 3 shows the differential pulse voltammograms of iron (III) in a mixture of glycine (0.2 mol�dm–<sup>3</sup> ) and NTA (5 � <sup>10</sup>–<sup>4</sup> mol�dm–<sup>3</sup> ). Reduction peak currents of mixed ligand complex depend on iron(III) concentrations, as shown in Figure 3. Basic line (voltammogram) represents the solution with both ligands present, without iron(III). It does not contain any reduction peak. This later implies electrochemical inactivity of these two ligands under the applied experimental conditions. When iron(III) is added, the reduction peak potentials remain constant at –0.112 V, indicating stability of the formed species. These peaks are the response to iron(III) reduction in mixed ligand complexes.

#### 3.3. Determination of formation equilibria of seven-coordinate Fe(EDTA) complexes with DNA and related biorelevant ligands

Fe(EDTA)-L is a seven-coordinate complex as the coordination number of Fe is seven. L can be a DNA constituent like uracil, uridine, thymine, thymidine, and inosine. To understand the chemistry of this seven-coordinate complex, we did some investigations using methylamine, ammonium chloride, or imidazole. The complexes produced are 1:1 with DNA constituents and other ligands. This complex indicates that the total coordination number of Fe(III) ion is seven. Potentiometric titration is carried out at 25�C and ionic strength 0.1 mol�L�<sup>1</sup> using NaNO3 to measure the stability constant. Besides, the nonlinear least-squares program MINIQUAD-75 is used to deduce the hydrolysis constants of [Fe(EDTA)(H2O)]� and its formation constant in solution. The concentration distributions of the different species formed in solution were evaluated as a pH dependent.

Figure 3. DPC voltammograms; iron(III)-Gly-NTA peak currents on added iron(III). 0.2 molċdm–<sup>3</sup> glycine, 5 � <sup>10</sup>–<sup>4</sup> molċdm–<sup>3</sup> NTA, 0.1 molċdm–<sup>3</sup> NaClO4; pH <sup>¼</sup> 8.0 � 0.1, <sup>E</sup>inc <sup>¼</sup> 2 mV, <sup>a</sup> <sup>¼</sup> 25 mV, <sup>t</sup><sup>p</sup> <sup>¼</sup> 0.05 s, <sup>t</sup>int <sup>¼</sup> 0.2 s.

#### 3.3.1. Calculated equilibria of the [Fe(EDTA)(H2O)]� ion

3.2. Ternary complexes of iron(III)-glycine( Gly)-nitrilotriacetate (NTA) system

peaks are the response to iron(III) reduction in mixed ligand complexes.

were 0.2, 0.02, and 0.002 mol�dm–<sup>3</sup>

(III) in a mixture of glycine (0.2 mol�dm–<sup>3</sup>

32 Descriptive Inorganic Chemistry Researches of Metal Compounds

DNA and related biorelevant ligands

solution were evaluated as a pH dependent.

Electrochemical measurements of the dissolved iron(III)-Gly-NTA mixed ligand system in the 0.1 mol�dm–<sup>3</sup> NaClO4 aqueous solution were performed at pH <sup>¼</sup> 8.0 � 0.1and 25 � <sup>1</sup>�C, using differential pulse cathodic voltammetry (DPCV), cyclic voltammetry (CV), and direct current (d. c.) polarography. Iron(III) concentrations were varied from 2.5 � <sup>10</sup>–<sup>5</sup> to 6 � <sup>10</sup>–<sup>4</sup> mol�dm–<sup>3</sup>

total concentrations varied from 2 � <sup>10</sup>–<sup>5</sup> to 1 � <sup>10</sup>–<sup>3</sup> mol�dm–<sup>3</sup> and glycine total concentrations

currents of mixed ligand complex depend on iron(III) concentrations, as shown in Figure 3. Basic line (voltammogram) represents the solution with both ligands present, without iron(III). It does not contain any reduction peak. This later implies electrochemical inactivity of these two ligands under the applied experimental conditions. When iron(III) is added, the reduction peak potentials remain constant at –0.112 V, indicating stability of the formed species. These

3.3. Determination of formation equilibria of seven-coordinate Fe(EDTA) complexes with

Fe(EDTA)-L is a seven-coordinate complex as the coordination number of Fe is seven. L can be a DNA constituent like uracil, uridine, thymine, thymidine, and inosine. To understand the chemistry of this seven-coordinate complex, we did some investigations using methylamine, ammonium chloride, or imidazole. The complexes produced are 1:1 with DNA constituents and other ligands. This complex indicates that the total coordination number of Fe(III) ion is seven. Potentiometric titration is carried out at 25�C and ionic strength 0.1 mol�L�<sup>1</sup> using NaNO3 to measure the stability constant. Besides, the nonlinear least-squares program MINIQUAD-75 is used to deduce the hydrolysis constants of [Fe(EDTA)(H2O)]� and its formation constant in solution. The concentration distributions of the different species formed in

Figure 3. DPC voltammograms; iron(III)-Gly-NTA peak currents on added iron(III). 0.2 molċdm–<sup>3</sup> glycine, 5 � <sup>10</sup>–<sup>4</sup> molċdm–<sup>3</sup>

NTA, 0.1 molċdm–<sup>3</sup> NaClO4; pH <sup>¼</sup> 8.0 � 0.1, <sup>E</sup>inc <sup>¼</sup> 2 mV, <sup>a</sup> <sup>¼</sup> 25 mV, <sup>t</sup><sup>p</sup> <sup>¼</sup> 0.05 s, <sup>t</sup>int <sup>¼</sup> 0.2 s.

. Figure 3 shows the differential pulse voltammograms of iron

) and NTA (5 � <sup>10</sup>–<sup>4</sup> mol�dm–<sup>3</sup>

, NTA

). Reduction peak

Different equilibrium models were tested [34] to fit the experimental potentiometric data for the hydrolysis of [Fe(EDTA)(H2O)]� ion. The best-selected model consists of the formation of the 10–<sup>1</sup> species, as given in Eq. (6). This supports the presence of one water molecule coordinated in the [Fe(EDTA)(H2O)]� ion:

½FeðEDTAÞðH2OÞ�� ð100Þ ⇌ Kal ½FeðEDTAÞðOHÞ�<sup>2</sup>� ð10�1Þ þH<sup>þ</sup> ð6Þ

The pH-meter readings (B) recorded in dioxane-water mixtures were converted to hydrogen ion concentrations [Hþ] with the widely used relation given by the Van Uitert and Haas equation [35]:

$$-\log\_{10}[\mathbf{H}^+] = B + \log\_{10} \mathcal{U}\_\mathbf{H} \tag{7}$$

where log10 U<sup>H</sup> is the correction factor for the solvent composition and ionic strength at which β was determined. Values of pK<sup>w</sup> in dioxane-water mixtures were determined as described previously [36, 37]. Different amounts from NaOH of the known concentration were added to a solution of ionic strength 0.1. The amount of base added determines the [OH�], unlike [Hþ], which is calculated from the pH value. The product of ([OH�]. [Hþ]) is used to calculate the mean values (pKw) which is �log10 [Hþ][OH�]. The mean values at 25�C are 14.17, 14.37, 14.50, and 15.44 for 12.5, 25, 37.5, and 50% dioxane, respectively. These percentages are the mass percentage of dioxane in water solution. The equilibrium constants obtained from the titration data (summarized in Table 2) are defined by Eqs. (8) and (9), where M, L, and H stand for [Fe(EDTA)(H2O)]�, ligand, and proton, respectively

$$r\text{ }p\text{M} + q\text{L } + r\text{H }\text{M}\_p\text{L}\_q\text{H}\_r\text{H}\_r\text{H}\_q\text{H}\_r\text{H}\_q\text{H}\_r\text{H}\_q\text{H}\_r\text{H}\_r\text{H}\_q\text{H}\_r\text{H}\_r\text{H}\_q\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\text{H}\_r\tag{8.2}$$

$$\mathcal{B}\_{pqr} = \left[\mathbf{M}\_p \mathbf{L}\_q \mathbf{H}\_r\right] / [\mathbf{M}]^p [\mathbf{L}]^q [\mathbf{H}]^r \tag{9}$$

The speciation distribution diagram for the hydrolysis of [Fe(EDTA)(H2O)]�is given in Figure 4. The fraction of the monohydroxo species increases with increasing pH, attaining a maximum of 99.9% at a pH ¼ 10.6.

#### 3.3.2. Complex formation equilibria of the [Fe(EDTA)(H2O)]�ion

The potentiometric titration curve, given in Figure 5, illustrates the result where imidazole is taken as an example. This curve has two plots one for the [Fe(EDTA)(H2O)]� -imidazole system and the other for imidazole. The complex formation curve for the [Fe(EDTA)(H2O)]�- -imidazole system is lower than the imidazole's one. This is because of the hydrogen ion evolved during the formation of a complex species. This potentiometric data are products for an experiment composed of the species 110.

There are many examples for the pyrimidinic species like uridine, uracil, thymine, and thymidine. The dissociable proton of the pyrimidinic species lies in the N3–C4O group. The acid dissociation constants for pyrimidinic species and the N1 proton of inosine are compared. The


a <sup>p</sup>, q, and <sup>r</sup> are the stoichiometric coefficients corresponding to [Fe(EDTA)(H2O)]�, L, and Hþ, respectively. <sup>b</sup> Standard deviations are given in parentheses.

c Sum of the squares of residuals.

Table 2. Stability constant of mixed complexes in water at 25 � 0.1 �C and 0.1 ionic strength.

Figure 4. Speciation distribution of different species as pH dependence in the Fe(EDTA)-OH system at 1.25 mmol�L�<sup>1</sup> [Fe(EDTA)(H2O)]�, in aqueous solution at 25�C and ionic strength I ¼ 0.1.

latter is slightly more acidic. The anionic form of purinic derivatives is the reason for that as it occurred in a large number of resonance forms. These resonance forms are created by the two condensed rings in the inosine ligand, as shown in Scheme 1. We can conclude that uracil,

Figure 5. Titration curves of the Fe(EDTA)-imidazole system in aqueous medium.

Scheme 1. Structural formulae of the DNA used.

latter is slightly more acidic. The anionic form of purinic derivatives is the reason for that as it occurred in a large number of resonance forms. These resonance forms are created by the two condensed rings in the inosine ligand, as shown in Scheme 1. We can conclude that uracil,

Figure 4. Speciation distribution of different species as pH dependence in the Fe(EDTA)-OH system at 1.25 mmol�L�<sup>1</sup>

System pQr<sup>a</sup> log10 β<sup>b</sup> S<sup>c</sup> [Fe(EDTA)(H2O)]� 1 0 �1 �7.60(0.008) 4.7E–8 Uracil 0 1 1 9.35(0,002) 4.5E–7

Thymine 0 1 1 9.50(0.01) 8.1E–8

Thymidine 0 1 1 9.06(0.01) 8.7E–8

Uridine 0 1 1 9.01(0.02) 1.1E–7

Methylamine�HCL 0 1 1 10.03 (0.04) 4.4E–7

Ammonium chloride 0 1 1 9.32(0.01) 7.2E–5

Imidazole 0 1 1 7.04(0.01) 2.6E–9

Inosine 0 1 1 8.43(0.01) 5.0E–9

<sup>p</sup>, q, and <sup>r</sup> are the stoichiometric coefficients corresponding to [Fe(EDTA)(H2O)]�, L, and Hþ, respectively. <sup>b</sup>

Table 2. Stability constant of mixed complexes in water at 25 � 0.1 �C and 0.1 ionic strength.

[Fe(EDTA)(H2O)]�, in aqueous solution at 25�C and ionic strength I ¼ 0.1.

a

c

Standard deviations are given in parentheses.

34 Descriptive Inorganic Chemistry Researches of Metal Compounds

Sum of the squares of residuals.

1 1 1 13.14(0.06)

1 1 0 5.12(0.1) 8.8E–6

1 1 0 5.98(0.1) 2.5E–6

1 1 0 5.89(0.1) 3.0E–5

1 1 0 4.93(0.05) 2.0E–5

1 1 0 6.13(0.1) 2.2E–5

1 1 0 3.91(0.03) 1.2E–6

1 1 0 2.23(0.04) 5.6E–7

1 1 0 5.96(0.02) 2.9E–5

uridine, thymine, and thymidine do not form protonated complexes. They coordinates through N3 in their deprotonated form; the monoanion one.

Thymine and thymidine have a methyl group that donates an extra electron of an inductive effect. This increases the basicity of the N3 site of thymine and thymidine complexes and stabilizes them more than uracil and uridine ones. Pyrimidines are monodentate ligands of pKa ≈ 9, so their complexes are absent below pH ¼ 6. This indicates that in the neutral or nearly basic pH media, the negatively charged nitrogen donors of pyrimidines bases are vital binding sites.

Inosine complex chelates as a monodentate has two sites of chelation N(1) and N(7). In the acidic medium, N1 is protonated and N7 is attached to the metal ion. When pH increases, the metal ion moves from N7 to N1. This motion was recorded by nuclear magnetic resonance (NMR) spectroscopy [38, 39]. Chelation depends on the pH range; in basic medium N(1) is a coordination site in the complex formation [40]. The data show the formation of the ternary complexes with stoichiometric coefficients 110 and 111. To know the main features observed in the species distribution in these systems, the speciation diagram obtained for the Fe(EDTA)-uracil and Fe(EDTA)-inosine complexes, as shown in Figure 6, as examples of DNA constituents. The pK<sup>a</sup> value of the N1H group of the protonated complex (log10 β111 log10 β110) amounts to 7.18. This indicates the acidification of the N1H site by 1.25 log units through coordination with the [Fe(EDTA)(H2O)] complex, which is in agreement with previous results for similar systems [41]. Detection of the concentration distribution of the various species in solution provides a useful picture of metal ion binding. At pH 4.0, the mixed complex of Fe(EDTA) with uracil, species 110, occurs. This occurrence increases with increasing the pH of the medium, up to it reaches 84% at pH 8.5. After 8.5 the concentration of Fe(EDTA)-uracil system

Figure 6. Speciation distribution curve of different species as pH dependence in the Fe(EDTA)-inosine and Fe(EDTA)-uracil systems (at 1.25 mmol<sup>L</sup><sup>1</sup> for Fe(EDTA) and 6.25 mmol<sup>L</sup><sup>1</sup> for inosine and uracil), in water at room temperature and 0.1 ionic strength.

drops and [Fe(EDTA)(OH)]2� develop. This result concludes that the Fe(EDTA) complex can interact with bioligands as DNA units. Not only does the pH affects the appearance of Fe (EDTA)-uracil system, it also affects the Fe(EDTA)-inosine system. At low pH mediums, species 111 is present; N7 coordinates to the complex and the N1 nitrogen is protonated. Whereas at high pH like 8.8, species 110 occurs where the concentration of N1 coordinated is 98%. This is the maximum concentration obtained for N1 coordinated. As the pH plays an important role before in the absence and presence of some systems, it has its role with cytosine and cytidine chelates. At low pH, they have their N3 protonated. It was recorded by the NMR spectroscopy in its solution state and with X-ray crystallography in its solid state.

#### 3.3.3. The influence of thermodynamic parameters

metal ion moves from N7 to N1. This motion was recorded by nuclear magnetic resonance (NMR) spectroscopy [38, 39]. Chelation depends on the pH range; in basic medium N(1) is a coordination site in the complex formation [40]. The data show the formation of the ternary complexes with stoichiometric coefficients 110 and 111. To know the main features observed in the species distribution in these systems, the speciation diagram obtained for the Fe(EDTA)-uracil and Fe(EDTA)-inosine complexes, as shown in Figure 6, as examples of DNA constituents. The pK<sup>a</sup> value of the N1H group of the protonated complex (log10 β111 log10 β110) amounts to 7.18. This indicates the acidification of the N1H site by 1.25 log units through coordination with the [Fe(EDTA)(H2O)] complex, which is in agreement with previous results for similar systems [41]. Detection of the concentration distribution of the various species in solution provides a useful picture of metal ion binding. At pH 4.0, the mixed complex of Fe(EDTA) with uracil, species 110, occurs. This occurrence increases with increasing the pH of the medium, up to it reaches 84% at pH 8.5. After 8.5 the concentration of Fe(EDTA)-uracil system

36 Descriptive Inorganic Chemistry Researches of Metal Compounds

Figure 6. Speciation distribution curve of different species as pH dependence in the Fe(EDTA)-inosine and Fe(EDTA)-uracil systems (at 1.25 mmol<sup>L</sup><sup>1</sup> for Fe(EDTA) and 6.25 mmol<sup>L</sup><sup>1</sup> for inosine and uracil), in water at room temperature and 0.1

ionic strength.

Thermodynamic parameters are useful tools for studying the interactions with DNA constituents and understanding the relative stability of the complexes formed. The thermodynamic parameters ΔG<sup>0</sup> , ΔH<sup>0</sup> , and ΔS<sup>0</sup> were easily determined using Van't Hoff relation (Eq. (10)). If we take the protonation of uracil and its complex formation with [Fe(EDTA)], as representative example. Since, we have a known value for protonation constant, the stability formation constant (K) and gas constant (R) in this reaction. So, we can apply the Van't Hoff equation to obtain the value of those parameters at the required temperature (T) in Kelvin. Then we can use the results to draw a graph of ln K versus 1/T and the intercept will be ΔS<sup>0</sup> /R.

$$
\Delta \text{InK} = -\Delta H^0 / \text{RT} + \Delta S^0 / \text{R} \tag{10}
$$

and a slope parameters ΔH<sup>0</sup> . The formation constants and the thermodynamic parameters values are presented in Tables 3 and 4 and can be interpreted as follows:

1. The protonation reaction of uracil can be represented as:

$$\text{L}^+ + \text{H}^+ \text{LH} \tag{11}$$

The thermodynamic processes accompanying the protonation reactions are as follows:



a <sup>p</sup>, q, and <sup>r</sup> are the stoichiometric coefficient corresponding to [Fe(EDTA)(H2O)]�, uracil, and Hþ, respectively. <sup>b</sup> Standard deviations are given in parentheses. c

Sum of the squares of residuals.

Table 3. Protonation constants of uracil and the formation constants of the Fe(EDTA)-uracil complex in aqueous solution and different temperatures and 0.1 ionic strength.


Table 4. Thermodynamic parameters (ΔH<sup>0</sup> , ΔS<sup>0</sup> and ΔG<sup>0</sup> ) for the interaction n of Fe–EDTA with uracil in aqueous solution.

The values of formation constants of the complexes are plotted in Figure 8 at different temperatures. The plotted line shows that the formation constants of the complexes are inversely proportion to increasing temperature. Therefore, the complexation process requires low temperatures. At the end, we can say that

Figure 7. Effect of temperature on the protonation constant of uracil.

Figure 8. Effect of temperature on the stability constant of [Fe(EDTA)(uracil)]2.

The values of formation constants of the complexes are plotted in Figure 8 at different temperatures. The plotted line shows that the formation constants of the complexes are inversely proportion to increasing temperature. Therefore, the complexation process requires low tem-

(1) [Fe(EDTA)(H2O)] ( [Fe(EDTA)(OH)]� þ H<sup>þ</sup> �54.40 � 0.75 27.20 � 2.5 43.4 � 1.5

Table 3. Protonation constants of uracil and the formation constants of the Fe(EDTA)-uracil complex in aqueous solution

<sup>p</sup>, q, and <sup>r</sup> are the stoichiometric coefficient corresponding to [Fe(EDTA)(H2O)]�, uracil, and Hþ, respectively. <sup>b</sup>

(2) L� þ H<sup>þ</sup> ( LH �34.29 � 0.75 63.91 � 2.5 �53.39 � 1.5

(3) [Fe(EDTA)(H2O)] þ L ( [Fe-EDTA-L] þ H2O �43.84 � 0.68 �52.64 � 2.3 �28.15 � 1.4

, ΔS<sup>0</sup> and ΔG<sup>0</sup>

) <sup>Δ</sup>S<sup>0</sup> (J�k1�<sup>1</sup>

�C) pqr<sup>a</sup> log10β<sup>b</sup> S<sup>c</sup>

Uracil 15 0 1 1 9.55(0.002) 1.6E–8 [Fe(EDTA)] 1 0 �1 �7.78(0.009) 1.6E–7 [Fe(EDTA)-uracil] 1 1 0 5.20(0.08) 1.3E–5 Uracil 20 0 1 1 9.46(0.002) 1.1E–8 [Fe(EDTA)] 1 0 �1 �7.69 (007) 1.0E–7 [Fe(EDTA)-uracil] 1 1 0 5.07(0.06) 9.7E–6 Uracil 25 0 1 1 9.35(0.002) 4.5E–7 [Fe(EDTA)] 1 0 �1 �7.60(0.008) 3.7E–8 [Fe(EDTA)-uracil] 1 1 0 4.93(0.05) 8.8E–6 Uracil 30 0 1 1 9.25(0.003) 8.6E–9 [Fe(EDTA)] 1 0 �1 �7.53(0.008) 1.2E–7 [Fe(EDTA)-uracil] 1 1 0 4.80(0.09) 1.0E–5 Uracil 35 0 1 1 9.15(0.003) 3.1E–7 [Fe(EDTA)] 1 0 �1 �7.46(0.008) 1.4E–7 [Fe(EDTA)-uracil] 1 1 0 4.69(0.14) 1.0E–5

�mol�<sup>1</sup>

) for the interaction n of Fe–EDTA with uracil in aqueous

) <sup>Δ</sup>G<sup>0</sup> (kJ�mol�<sup>1</sup>

)

peratures. At the end, we can say that

Table 4. Thermodynamic parameters (ΔH<sup>0</sup>

Standard deviations are given in parentheses.

and different temperatures and 0.1 ionic strength.

Sum of the squares of residuals.

Equilibrium <sup>Δ</sup>H<sup>0</sup> (kJ�mol�<sup>1</sup>

Fe-EDTA hydrolysis

System T (

38 Descriptive Inorganic Chemistry Researches of Metal Compounds

Fe-EDTA-uracil

L denotes uracil.

solution.

Uracil

a

c


#### 3.3.4. How can you expect the effect of solvent composition

It is well known that the "effective" or "equivalent solution" dielectric constants in a protein [42, 43], or active site cavities of enzymes [44] are small compared to that in bulk water. The dielectric constants detecting in such locations range from 30 to 70 [43, 44]. Therefore, by using aqueous solutions containing ~10–50% dioxane, one may expect to simulate to some degree the situation in active site cavities [45], and hence to extrapolate the data to physiological conditions. We asked what the relation between the solvent occurs in media and formation constant of complexes on the equilibrium constants (Table 4) reveals the following points:


Figure 9. Effect of dioxane on the protonation constant of uracil.

Figure 10. Effect of dioxane on the stability constant of the [Fe(EDTA)uracil)]<sup>2</sup> species.


a <sup>p</sup>, q, and <sup>r</sup> are the stoichiometric coefficient corresponding to [Fe(EDTA)(H2O)]�, uracil, and Hþ, respectively. <sup>b</sup> Standard deviations are given in parentheses. c

Sum of the squares of residuals.

3.3.4. How can you expect the effect of solvent composition

40 Descriptive Inorganic Chemistry Researches of Metal Compounds

Figure 9. Effect of dioxane on the protonation constant of uracil.

Figure 10. Effect of dioxane on the stability constant of the [Fe(EDTA)uracil)]<sup>2</sup> species.

It is well known that the "effective" or "equivalent solution" dielectric constants in a protein [42, 43], or active site cavities of enzymes [44] are small compared to that in bulk water. The dielectric constants detecting in such locations range from 30 to 70 [43, 44]. Therefore, by using aqueous solutions containing ~10–50% dioxane, one may expect to simulate to some degree the situation in active site cavities [45], and hence to extrapolate the data to physiological conditions. We asked what the relation between the solvent occurs in media and formation constant of complexes on the equilibrium constants (Table 4) reveals the following points:

1. The value of pK<sup>a</sup> of uracil (N3-site) increases directly with increasing dioxane in the medium as shown in Figure 9. Dioxane has a low-dielectric constant, which increased the electrostatic forces between the proton and the ligand. Finally, the pK<sup>a</sup> increases. 2. The stability constant (log10 K1) of the Fe(EDTA)-uracil complex increases with increase of the dioxane concentration (Figure 10). This is due to that lowering the dielectric constant of

Table 5. Effect of solvent (dioxane) on the stability constant of [Fe(EDTA)(uracil)] at 25�C.

the medium (by increasing the dioxane content) favors the interaction between Fe(EDTA) and uracil, and consequently the stability constant of the complex increases. These finding, shown in Table 5, is in agreement with literature data [46].
