2. Analysis of aqueous solution complexes by different methods


There are other experimental techniques that we are going to give some examples for them in the following lines. The specific method for diamagnetic metal ion is the nuclear magnetic resonance (NMR). It gives one separated signal for each unique chemical surrounding. In other words, it can inform us about concentration of the ligand, the free metal ion, the number of species, and their concentration for a given chemical composition. The important feature of this technique is that the positions of these selective signals are responsible for the protonation and deprotonation ones. In the case of fast reactions, we can use a stopped-flow technique.

#### 2.1. The addition of water to iron(III)

It is observed that Fe3<sup>þ</sup> hydrolyses in water goes as follows:

$$\text{Fe}^{3+} + \text{H}\_2\text{O} \leftrightharpoons \text{FeOH}^{2+} + \text{H}^+ \tag{1}$$

The addition of water to Fe3<sup>þ</sup> is carried out through a series of deprotonation reactions, resulting in formation of ferric hydroxides and oxyhydroxides [12, 13] as in Eq. (1), the equilibrium constant for this reaction was calculated to be 6.78 � <sup>10</sup>�<sup>3</sup> at 298 K (total iron(III) concentration of 0.5 mol�dm�<sup>3</sup> and an ionic strength of 0.1 mol�dm�<sup>3</sup> ). Figure 1 shows the speciation of the two ironcontaining species of Eq. (1) as a function of pH, calculated at <sup>T</sup> <sup>¼</sup> 298 K. It seen that FeOH2<sup>þ</sup> will become the common species above pH ¼ 2.17. At lower pH, there are small amount of it; at pH value 1.2, more than 9% of all iron (III) is present as the hydroxide. As a result of that, the calculations for the spectroscopic measurements were carried out at pH values up to 1. The equilibrium constant for the reaction shown in Eq. (1) has been studied at temperature range from

iron intake and bioavailability. On the other side, they are responsible for only about half of the anemia in developing countries [10]. There are other important causes [11] like infectious and inflammatory diseases (especially malaria), blood loss from parasitic infections, and other

1. The first and most common method is ion-selective electrode. It defines the position of dynamic equilibrium. The most important electrode is the glass one. There is also hydro-

2. Metal amalgam electrodes are a second choice. It can be used for some metal ions, but they are not as precise as the hydrogen ion electrode. We prefer the ion-selective electrode in-calculation because the results are collected from series of data taken through a titration procedure. A good method to check for this prerequisite is to make repeated high-resolution electrode readings at predetermined time intervals, since this will make sluggish

3. Spectrophotometry can be used if the metal ion or the ligand is colored, so that the color

There are other experimental techniques that we are going to give some examples for them in the following lines. The specific method for diamagnetic metal ion is the nuclear magnetic resonance (NMR). It gives one separated signal for each unique chemical surrounding. In other words, it can inform us about concentration of the ligand, the free metal ion, the number of species, and their concentration for a given chemical composition. The important feature of this technique is that the positions of these selective signals are responsible for the protonation and deprotonation ones. In the case of fast reactions, we can use a stopped-flow technique.

The addition of water to Fe3<sup>þ</sup> is carried out through a series of deprotonation reactions, resulting in formation of ferric hydroxides and oxyhydroxides [12, 13] as in Eq. (1), the equilibrium constant for this reaction was calculated to be 6.78 � <sup>10</sup>�<sup>3</sup> at 298 K (total iron(III) concentration of 0.5

containing species of Eq. (1) as a function of pH, calculated at <sup>T</sup> <sup>¼</sup> 298 K. It seen that FeOH2<sup>þ</sup> will become the common species above pH ¼ 2.17. At lower pH, there are small amount of it; at pH value 1.2, more than 9% of all iron (III) is present as the hydroxide. As a result of that, the calculations for the spectroscopic measurements were carried out at pH values up to 1. The equilibrium constant for the reaction shown in Eq. (1) has been studied at temperature range from

Fe<sup>3</sup><sup>þ</sup> <sup>þ</sup> H2O ⇋ FeOH<sup>2</sup><sup>þ</sup> <sup>þ</sup> <sup>H</sup><sup>þ</sup> <sup>ð</sup>1<sup>Þ</sup>

). Figure 1 shows the speciation of the two iron-

nutrient deficiencies (vitamin A, riboflavin, folic acid, and vitamin B12).

gen gas electrode that can be used in hydrogen ion calculations.

will change (in intensity and/or frequency) upon complexation.

attainments of equilibrium clearly visible.

28 Descriptive Inorganic Chemistry Researches of Metal Compounds

2.1. The addition of water to iron(III)

mol�dm�<sup>3</sup> and an ionic strength of 0.1 mol�dm�<sup>3</sup>

It is observed that Fe3<sup>þ</sup> hydrolyses in water goes as follows:

2. Analysis of aqueous solution complexes by different methods

Figure 1. The speciation of the two species of iron as in Eq. (1) as a function of pH at T ¼ 298 K for all the iron (III) concentration of 0.5 mol�dm�<sup>3</sup> and an ionic strength of 0.1 mol�dm�<sup>3</sup> .

<sup>T</sup> <sup>¼</sup> 298 to 353 K and ionic strength range from <sup>I</sup> <sup>¼</sup> 0.1 until 2.67 mol�dm�<sup>3</sup> in perchlorate media [14].

For Eq. (1) to be done we must prevent the formation of hydrolyzed iron(III). This can be accomplished under the conditions of <sup>T</sup> <sup>¼</sup> <sup>293</sup>–323 K, and <sup>I</sup> ~ 0.1 mol�dm�<sup>3</sup> .

#### 2.2. How can you determine the stability constants of mixed ligand complexes

For inorganic chemistry, it is very important to determine the stability constant, or the equilibria constant or we can refer to it as the formation constant, for the reaction [15]. It is not easy to get the solution equilibria constants between the ligands and the metal ions. Proton ions and a range of metal ions fight for a range of donor sites. There are many factors that decide who will be the winner whether the proton ions or the metal ions. These factors are the concentration and pH. Potentiometry and spectrophotometry are used to determine the stability constants of metal complexes. Legget [16] and Meloum et al. [17] calculated the equilibrium constants from experimental data for the first time. Nowadays, many programs were published for these calculations using microcomputers. Table 1 presents some of these programs [15, 19–30]. These programs are very helpful as they quickly present the best fit. They use the least-square method to reduce the differences between calculated and experimental data. The sum of square of residuals between calculated and experimental values is very small; it is nearly between 10�<sup>6</sup> and 10�<sup>9</sup> . Potentiometry is used to determine the stability constants of metal complexes. It is based on pH-metric titration of the ligand, and the availability of metal ions. Data obtained from potentiometry are analyzed by the least-square method to derive the formation constant. This later can describe the solution equilibria. For the measurements, there must be two conditions: the first one is a constant ionic strength of the solution and the second condition concerns the ionic strength that have to be higher than the concentration of metal ion. The reaction of all mixed complex:

$$\mathbf{L(M)} + p(\mathbf{L\_1}) + q(\mathbf{L\_2}) + r(\mathbf{H}) \leftrightarrow \left[ (\mathbf{M})\_1 (\mathbf{L\_1})\_p (\mathbf{L\_2})\_q (\mathbf{H})\_r \right] \tag{2}$$


V, potentiometric experiments; A, spectrophotometric experiments; E, ESR; N, NMR. a Additional data used in calculations are taken from different sources.

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

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

β1pqr ¼ ½ MÞlðL1ÞpðL2ÞqðHÞr�=½M� 1 ½L1�p½L2�q½H� <sup>r</sup>ðfor simplicity charges are omitted<sup>Þ</sup> <sup>ð</sup>3<sup>Þ</sup>

where M, L1, L2, and H stand for metal ion, ligand (1), ligand (2), and proton, respectively. For OH�, the coefficient (r) for H ¼ �1.

#### 2.3. Calculation of speciation

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