**1. Introduction**

The stability of metal complex generally means that it exists under favorable conditions without undergoing decomposition and has a considerable shelf life period [1]. The term stability of metal complex cannot be generalized since the complex may be stable to one reagent/condition and may decompose in presence of another reagent/ condition. The stability of metal complexes can be explained with the help of two different aspects, namely, thermodynamic stability and kinetic stability [2]. Nevertheless, a metal complex is said to be stable if it does not react with water, which would lead to a decrease in the free energy of the system, i.e., thermodynamic stability. On the other hand, the complex is said to possess kinetic stability if it reacts with water to form a stable product and there is a known mechanism through which the reaction can proceed. For example, the system may not have sufficient energy available to break a strong bond, although once the existing bond is broken it could be replaced by new bond which is stronger than the older one [1]. Stability of complex compound is assigned to be its existence in aqueous solution with respect to its bond dissociation energy, Gibbs free energy, standard electrode potential, pH of the solution, and rate constant or activation energy for substitution reactions.

## **1.1 Thermodynamic stability**

Thermodynamic stability of a complex refers to its tendency to exist under equilibrium conditions. It determines the extent to which the complex will be formed or be converted into another complex at the point of equilibrium. In other words, thermodynamic stability of complexes is the measure of tendency of a metal ion to selectively form a specific metal complex and is directly related to the metal-ligand bond energies. The thermodynamic stability of complexes is

represented by formation constant. The formation constant is also known as stability constant, which is the equilibrium constant obtained for the formation metal complex [1, 2].

In general, the metal complexes are not prepared from their corresponding starting materials in gaseous phase but are prepared in aqueous solution. In aqueous solution, a metal cation gets hydrated to give aqua complex of the type [M(H2O)x] n+. When a ligand replaces water molecule from aqua complex ion, a new metal complex is formed and equilibrium is established as shown:

$$\begin{array}{rcl} \text{[M(H}\_{2}\text{O)}\_{x}\text{]}^{\text{n}} & + & \text{[K}\_{f} \\ \end{array} \\ \begin{array}{rcl} \text{K}\_{f} \\ \text{[M(H}\_{2}\text{O)}\_{x}\text{]}\_{x}\text{]}\_{x} & + & \text{H}\_{2}\text{O} \\ \end{array} \tag{1}$$

where x is the number of water molecules, n is the oxidation number of the metal cation, and L is the neutral and monodentate ligand. For simplicity, the above reaction can be written in generalized form as given:

$$\mathbf{^M}\mathbf{^L}\mathbf{<^L}\mathbf{<^L}\mathbf{^M}\mathbf{^L}\tag{2}$$

ii. Electrostatic factor

*DOI: http://dx.doi.org/10.5772/intechopen.90894*

*Stability of Metal Complexes*

the formation of [CdBr4]

octahedral [Cd(H2O)3Br3]

**25**

complex [Cd(H2O)6]

iii. Steric hindrance with increase in number of ligands

iv. Statistical factors (number of replaceable positions)

dral. The reaction sequence for the formation of [CdBr4]

However, in some cases, it is found that Kn+1 > Kn because of unusual structural changes and changes in electronic configuration of the metal ion. The change in electronic structure of the metal ion causes the variation in the crystal field stabilization energy (CFSE). The complex with higher CFSE value will be stable, and the equilibrium constant for that complex formation will be high. One such example is

wise formation constants K1, K2, K3, and K4. The order of stepwise formation constants is observed as follows, K1 > K2 > K3 < K4, which is not in agreement with the common trend of K1 > K2 > K3 > K4. Aqua complex of most of the M2+ ions including Cd2+ are octahedral, whereas the halo complexes of Cd2+ ion are tetrahe-

In the final step, there is an unusual structural change from six coordinated

complex in addition to change in the electronic configuration which lead to K4 > K3. The formation constant (K*f*) is related to the standard Gibbs free energy change

Since ΔG° is a thermodynamic property, the formation constant is the measure of thermodynamic stability. From Eqs. (4) to (6), it can be interpreted that the thermodynamic stability of a complex can be measured in terms of formation constant, Gibbs free energy change, and standard electrode potential. A high negative value of ΔG° indicates that the position of equilibrium favors the product

The formation constant describes the formation of a complex from metal cation

(complex); hence the complex formed will be more stable.

**1.2 Stepwise formation of complex and stepwise formation constants**

and ligands. Bjerrum (1941) defined that the formation of a metal complex in aqueous solution takes place by replacing the water molecule by another ligand (L)

(ΔG°) and standard electrode potential (E°) according to following equations:

<sup>2</sup>� complex in aqueous solution. The reaction of aqua

2+ with Br� ligand exhibits four stepwise equilibrium or step-

� complex to four coordinated tetrahedral [CdBr4]

ΔG° ¼ �RTlnK*<sup>f</sup>* (4)

ΔG° ¼ �n*F*E° (5)

hence RTlnK*<sup>f</sup>* ¼ n*F*E° (6)

<sup>2</sup>� is given as follows:

2�

The equilibrium constant K*<sup>f</sup>* of the reaction is given by:

$$\mathbf{K}\_{\circ} - \frac{[\mathbf{ML}]}{[\mathbf{M}][\mathbf{L}]} \tag{3}$$

In the above equation, the concentration of water is not included. Since the solution is dilute, the water molecules which enter the bulk solution do not have much influence on the equilibrium constant. It is observed from Eq. (3) that the higher the value of K*f*, the greater will be the stability of the complex formed. A high value of the equilibrium constant (K*<sup>f</sup>* > 1.0) also indicates that at equilibrium the activity of complex ML is larger than the product of activities of M and L. Thus, large value of K*<sup>f</sup>* indicates that the ligand L binds to the metal ion more strongly than H2O and hence L is a stronger ligand than H2O. If K*<sup>f</sup>* is less than 1.0, then ligand L is weaker than H2O. Thus stability constant is used as a measure of thermodynamic stability of the complex. With a few exceptions, the value of successive stability constants decreases regularly from K1 to Kn, that is, K1 > K2 > K3 > … > Kn�<sup>1</sup> > Kn. This trend is illustrated by taking formation of [Cd(NH3)4] 2+ as an example [3, 4]:

$$\begin{aligned} \text{Cd}^{2+} + \text{NH}\_3 &\xleftarrow{\text{Cd(NH}\_3\text{)}^{2+}} ; \text{ K}\_1 = 10^{2.65} \\\\ [\text{Cd(NH}\_3\text{)}]^{2+} + \text{NH}\_3 &\xleftarrow{\text{Cd(NH}\_3\text{)}\_2\text{)}^{2+}} ; \text{ K}\_2 = 10^{2.10} \\\\ [\text{Cd(NH}\_3\text{)}\_2\text{)}^{2+} + \text{NH}\_3 &\xleftarrow{\text{Cd(NH}\_3\text{)}\_3\text{)}^{2+}} [\text{Cd(NH}\_3\text{)}\_3]^{2+} ; \text{ K}\_3 = 10^{1.44} \\\\ [\text{Cd(NH}\_3\text{)}\_3]^{2-} + \text{NH}\_3 &\xleftarrow{\text{Cd(NH}\_3\text{)}\_4\text{)}^{2-}} [\text{Cd(NH}\_3\text{)}\_4]^{2-} ; \text{ K}\_4 = 10^{0.93} \end{aligned}$$

The steady decrease in the value of stepwise formation constants from K1 to Kn is due to:

i. Increase in the number of ligands in coordination sphere that causes to decrease the number of H2O molecules to be replaced and thus the probability of replacement of water molecules decreased

ii. Electrostatic factor

represented by formation constant. The formation constant is also known as stability constant, which is the equilibrium constant obtained for the formation metal

In general, the metal complexes are not prepared from their corresponding starting materials in gaseous phase but are prepared in aqueous solution. In aqueous solution, a metal cation gets hydrated to give aqua complex of the type [M(H2O)x]

When a ligand replaces water molecule from aqua complex ion, a new metal complex

where x is the number of water molecules, n is the oxidation number of the metal cation, and L is the neutral and monodentate ligand. For simplicity, the above

In the above equation, the concentration of water is not included. Since the solution is dilute, the water molecules which enter the bulk solution do not have much influence on the equilibrium constant. It is observed from Eq. (3) that the higher the value of K*f*, the greater will be the stability of the complex formed. A high value of the equilibrium constant (K*<sup>f</sup>* > 1.0) also indicates that at equilibrium the activity of complex ML is larger than the product of activities of M and L. Thus, large value of K*<sup>f</sup>* indicates that the ligand L binds to the metal ion more strongly than H2O and hence L is a stronger ligand than H2O. If K*<sup>f</sup>* is less than 1.0, then ligand L is weaker than H2O. Thus stability constant is used as a measure of thermodynamic stability of the complex. With a few exceptions, the value of successive stability constants decreases regularly from K1 to Kn, that is, K1 > K2 > K3 >

… > Kn�<sup>1</sup> > Kn. This trend is illustrated by taking formation of [Cd(NH3)4]

The steady decrease in the value of stepwise formation constants from K1 to Kn is

decrease the number of H2O molecules to be replaced and thus the probability

i. Increase in the number of ligands in coordination sphere that causes to

of replacement of water molecules decreased

is formed and equilibrium is established as shown:

*Stability and Applications of Coordination Compounds*

reaction can be written in generalized form as given:

The equilibrium constant K*<sup>f</sup>* of the reaction is given by:

n+.

ð1Þ

ð2Þ

ð3Þ

2+ as an

complex [1, 2].

example [3, 4]:

due to:

**24**


However, in some cases, it is found that Kn+1 > Kn because of unusual structural changes and changes in electronic configuration of the metal ion. The change in electronic structure of the metal ion causes the variation in the crystal field stabilization energy (CFSE). The complex with higher CFSE value will be stable, and the equilibrium constant for that complex formation will be high. One such example is the formation of [CdBr4] <sup>2</sup>� complex in aqueous solution. The reaction of aqua complex [Cd(H2O)6] 2+ with Br� ligand exhibits four stepwise equilibrium or stepwise formation constants K1, K2, K3, and K4. The order of stepwise formation constants is observed as follows, K1 > K2 > K3 < K4, which is not in agreement with the common trend of K1 > K2 > K3 > K4. Aqua complex of most of the M2+ ions including Cd2+ are octahedral, whereas the halo complexes of Cd2+ ion are tetrahedral. The reaction sequence for the formation of [CdBr4] <sup>2</sup>� is given as follows:

$$\begin{aligned} \begin{aligned} \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{5}]\$^{2+}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{1}} \\ \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{5}\$Br]\$^{+}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{2}} \\ \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{4}\$Br\$]^{2}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{3}} \\ \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{4}\$Br\$]^{2}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{4}} \\ \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{3}\$Br\$]^{+}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{4}} \end{aligned} \\ \begin{aligned} \text{[\$\mathsf{Cd}(\mathsf{H}\_{2}\mathsf{O})\_{3}\$Br\$]^{2}\$} + \text{Br}^{\cdot} &\xrightarrow{\mathsf{K}\_{1}} \end{aligned} \end{aligned}$$

In the final step, there is an unusual structural change from six coordinated octahedral [Cd(H2O)3Br3] � complex to four coordinated tetrahedral [CdBr4] 2� complex in addition to change in the electronic configuration which lead to K4 > K3.

The formation constant (K*f*) is related to the standard Gibbs free energy change (ΔG°) and standard electrode potential (E°) according to following equations:

$$
\Delta \mathbf{G}^{\circ} = -\mathbf{R} \mathbf{T} \ln \mathbf{K}\_{\circ} \tag{4}
$$

$$
\Delta \mathbf{G}^{\circ} = -\mathbf{n} F \mathbf{E}^{\circ} \tag{5}
$$

$$\text{hence RTInK}\_f = \text{nFE}^\circ \tag{6}$$

Since ΔG° is a thermodynamic property, the formation constant is the measure of thermodynamic stability. From Eqs. (4) to (6), it can be interpreted that the thermodynamic stability of a complex can be measured in terms of formation constant, Gibbs free energy change, and standard electrode potential. A high negative value of ΔG° indicates that the position of equilibrium favors the product (complex); hence the complex formed will be more stable.

#### **1.2 Stepwise formation of complex and stepwise formation constants**

The formation constant describes the formation of a complex from metal cation and ligands. Bjerrum (1941) defined that the formation of a metal complex in aqueous solution takes place by replacing the water molecule by another ligand (L) [5, 6]. It is assumed that this reaction does not occur in a single step but occurs in several steps, and each step is characterized by its individual equilibrium constant called as stepwise formation constant (K). For example, consider the formation of a complex [MLn] formed by the following reactions:

$$\begin{array}{c} \text{ML} + \text{L} \xrightarrow{\text{K}^{\text{I}}} \text{ML} \end{array}$$

Above equation indicates that the overall formation constant (β) is equal to the

Kinetic stability is related to the reactivity of the metal complexes in solution and deals with the rate of the reaction, its activation energy, etc. Kinetic stability is also related to how fast a compound reacts rather than how stable it is. It aids in determining the rate at which the reaction occurs to establish the equilibrium [7]. The term kinetic stability of complexes is classified into labile and inert by Taube on the basis of rate of the reactions. When the rate of substitution of ligands is high, the complex is said to be labile. For example, the copper complex of the formula

when concentrated hydrochloric acid is added to this solution, the solution turns

reaction takes place at room temperature when conc. HCl was added to the aqueous solution. However, only one NH3 ligand was found to be substituted by Cl ligand, when the aqueous solution of the complex was heated with 6M hydrochloric acid.

For metal complexes, the stability and reactivity are described in thermodynamic and kinetic terms, respectively. In particular, the terms stable and unstable are related to thermodynamic aspects, whereas labile and inert terms are related to kinetic aspects. As a rule of thumb, a metal complex is said to be labile if it reacts within 1 min at 25°C, and if it takes longer time, it is considered to be inert.

Thermodynamic stability refers to the energy change that occurs while starting materials are converted to products, that is, ΔG, for the reaction. The change in free energy is given by the equation ΔG = ΔHTΔS = RTlnK, where ΔS is the entropy, ΔH is the enthalpy, and K is the equilibrium constant for the reaction. Kinetic stability refers to reactivity or the ability of the metal complex to undergo ligand substitution reactions. Complexes which undergo extremely rapid ligand substitution reaction are referred to as labile complexes, and complexes that undergo extremely slow ligand substitution reaction are referred to as inert complexes. Sometimes the thermodynamic and kinetic stabilities of complexes are parallel to one another, but often they do not. One of the suitable examples for thermody-

3+ is thermodynamically unstable but kinetically inert. The complex

3+ is thermodynamically unstable since the complex was observed to

decompose very rapidly with rate in the order of 10<sup>25</sup> in acidic solution. However, no ligand substitution reaction is found when the complex is kept in acidic solution for several days; hence the complex is kinetically inert. From the above two examples, it can be interpreted that the stability of a complex mainly depends upon the

rate of ligand exchange is very slow, and the ligands are very exchanged with

**2. Relation between thermodynamic and kinetic stabilities**

namically stable and kinetically inert complex is [Ni(CN)4]

substitution reaction very rapidly. On the other hand, the cobalt complex

2+ is labile. In aqueous solution the complex is blue in color, and

2+. On the other hand, in inert complexes the

3+ reacts slowly, and no

<sup>2</sup> as it undergoes ligand

product of the stepwise formation constant K1, K2, K3, … , Kn.

**1.3 Kinetic stability**

*Stability of Metal Complexes*

*DOI: http://dx.doi.org/10.5772/intechopen.90894*

[Cu(NH3)4(H2O)2]

[Co(NH3)6]

[Co(NH3)6]

**27**

green giving rise to complex [CuCl4]

difficulty. For example, the cobalt complex [Co(NH3)6]

By assuming the value of activity coefficients as unity, the equilibrium constant K1 for the complex (ML) having one ligand (L) will be given as

$$\mathbf{K}\_{\parallel} = \frac{[\mathbf{M}\mathbf{L}]}{[\mathbf{M}][\mathbf{L}]}$$

When the metal complex ML reacts with one more ligand L,

$$\begin{array}{cccc} \text{ML} & + \text{L} & \xleftarrow{\text{K}\_{2}} \text{ML}\_{2} \end{array}$$

and the equilibrium constant K2 will be

$$\mathbf{K\_2} = \frac{[\mathbf{ML}\_2]}{[\mathbf{ML}][\mathbf{L}]},$$

Similarly, for the formation of the complex MLn from MLn<sup>1</sup> and L, the equilibrium constant is represented as follows,

$$\mathbf{K}\_{\mathfrak{n}} = \frac{[\mathbf{M}\mathbf{L}\_{\mathfrak{n}}]}{[\mathbf{M}\mathbf{L}\_{\mathfrak{n}-1}][\mathbf{L}]}$$

The equilibrium constants K1, K2, … , Kn are known as stepwise formation constants. On the other hand, the equilibrium constant for the overall reaction may be considered as

$$\begin{aligned} \text{M} &+ \text{ L} \xrightarrow{\overline{\text{B}\_{1}}} \text{ML}, & \quad \text{B}\_{1} - \frac{\text{[ML]}}{[\text{M}][\text{L}\_{1}]} \\ \text{M} &+ 2\text{L} \xrightarrow{\overline{\text{B}\_{2}}} \text{ML}\_{2}, & \quad \text{B}\_{2} = \frac{\text{[ML}\_{2}]}{[\text{M}][\text{L}\_{2}]^{2}} \\ \text{M} + 3\text{L} \xrightarrow{\overline{\text{B}\_{3}}} \text{ML}\_{3}, & \quad \text{B}\_{3} = \frac{[\text{ML}\_{3}]}{[\text{M}][\text{L}\_{1}]^{3}} \\ \text{M} + \text{nL} &\xrightarrow{\overline{\text{B}\_{n}}} \text{ML}\_{n}, & \quad \text{B}\_{n} = \frac{[\text{ML}\_{n}]}{[\text{M}][\text{L}\_{1}]^{n}} \end{aligned}$$

where β1, β2, β3, … , β<sup>n</sup> are the equilibrium constants called as overall formation constants and K1, K2, K3, … , Kn are stepwise stability or formation constants. The products of stepwise constants are Ks and βs are related one another. For example, consider the product of stepwise formation constants K1, K2, K3, … , Kn.

*Stability of Metal Complexes DOI: http://dx.doi.org/10.5772/intechopen.90894*

$$\mathsf{K}\_{1}\mathsf{x}\ \mathsf{K}\_{2}\ \mathsf{x}\ \mathsf{K}\_{3}\ \mathsf{x}\ \mathsf{x}\ \dots \mathsf{K}\_{n} \ = \frac{[\mathsf{M}\mathsf{L}\mathsf{L}]}{[\mathsf{M}\mathsf{L}][\mathsf{L}\mathsf{L}]}\ \mathsf{x}\ \frac{[\mathsf{M}\mathsf{L}\mathsf{L}\_{\mathsf{T}}]}{[\mathsf{M}\mathsf{L}][\mathsf{L}\mathsf{L}]}\ \mathsf{x}\ \frac{[\mathsf{M}\mathsf{L}\_{\mathsf{T}}]}{[\mathsf{M}\mathsf{L}\_{\mathsf{T}}][\mathsf{L}\mathsf{L}]}\mathsf{x}\ \cdots \frac{[\mathsf{M}\mathsf{L}\mathsf{n}]}{[\mathsf{M}\mathsf{L}\_{\mathsf{T}\mathsf{n}}][\mathsf{L}\mathsf{L}]}\ \mathsf{x}\ \mathsf{x}\ \cdots$$

$$=\frac{[\mathsf{M}\mathsf{L}\_{\mathsf{T}}]}{[\mathsf{M}\mathsf{M}][\mathsf{L}]}\mathsf{x}\ \mathsf{x}$$

Above equation indicates that the overall formation constant (β) is equal to the product of the stepwise formation constant K1, K2, K3, … , Kn.

#### **1.3 Kinetic stability**

[5, 6]. It is assumed that this reaction does not occur in a single step but occurs in several steps, and each step is characterized by its individual equilibrium constant called as stepwise formation constant (K). For example, consider the formation of a

By assuming the value of activity coefficients as unity, the equilibrium constant

Similarly, for the formation of the complex MLn from MLn<sup>1</sup> and L, the equilib-

The equilibrium constants K1, K2, … , Kn are known as stepwise formation constants. On the other hand, the equilibrium constant for the overall reaction may

where β1, β2, β3, … , β<sup>n</sup> are the equilibrium constants called as overall formation constants and K1, K2, K3, … , Kn are stepwise stability or formation constants. The products of stepwise constants are Ks and βs are related one another. For example, consider the product of stepwise formation constants

complex [MLn] formed by the following reactions:

*Stability and Applications of Coordination Compounds*

and the equilibrium constant K2 will be

rium constant is represented as follows,

be considered as

K1, K2, K3, … , Kn.

**26**

K1 for the complex (ML) having one ligand (L) will be given as

When the metal complex ML reacts with one more ligand L,

Kinetic stability is related to the reactivity of the metal complexes in solution and deals with the rate of the reaction, its activation energy, etc. Kinetic stability is also related to how fast a compound reacts rather than how stable it is. It aids in determining the rate at which the reaction occurs to establish the equilibrium [7].

The term kinetic stability of complexes is classified into labile and inert by Taube on the basis of rate of the reactions. When the rate of substitution of ligands is high, the complex is said to be labile. For example, the copper complex of the formula [Cu(NH3)4(H2O)2] 2+ is labile. In aqueous solution the complex is blue in color, and when concentrated hydrochloric acid is added to this solution, the solution turns green giving rise to complex [CuCl4] 2+. On the other hand, in inert complexes the rate of ligand exchange is very slow, and the ligands are very exchanged with difficulty. For example, the cobalt complex [Co(NH3)6] 3+ reacts slowly, and no reaction takes place at room temperature when conc. HCl was added to the aqueous solution. However, only one NH3 ligand was found to be substituted by Cl ligand, when the aqueous solution of the complex was heated with 6M hydrochloric acid.

#### **2. Relation between thermodynamic and kinetic stabilities**

For metal complexes, the stability and reactivity are described in thermodynamic and kinetic terms, respectively. In particular, the terms stable and unstable are related to thermodynamic aspects, whereas labile and inert terms are related to kinetic aspects. As a rule of thumb, a metal complex is said to be labile if it reacts within 1 min at 25°C, and if it takes longer time, it is considered to be inert.

Thermodynamic stability refers to the energy change that occurs while starting materials are converted to products, that is, ΔG, for the reaction. The change in free energy is given by the equation ΔG = ΔHTΔS = RTlnK, where ΔS is the entropy, ΔH is the enthalpy, and K is the equilibrium constant for the reaction. Kinetic stability refers to reactivity or the ability of the metal complex to undergo ligand substitution reactions. Complexes which undergo extremely rapid ligand substitution reaction are referred to as labile complexes, and complexes that undergo extremely slow ligand substitution reaction are referred to as inert complexes. Sometimes the thermodynamic and kinetic stabilities of complexes are parallel to one another, but often they do not. One of the suitable examples for thermodynamically stable and kinetically inert complex is [Ni(CN)4] <sup>2</sup> as it undergoes ligand substitution reaction very rapidly. On the other hand, the cobalt complex [Co(NH3)6] 3+ is thermodynamically unstable but kinetically inert. The complex [Co(NH3)6] 3+ is thermodynamically unstable since the complex was observed to decompose very rapidly with rate in the order of 10<sup>25</sup> in acidic solution. However, no ligand substitution reaction is found when the complex is kept in acidic solution for several days; hence the complex is kinetically inert. From the above two examples, it can be interpreted that the stability of a complex mainly depends upon the

conditions, and it is always recommended to specify the conditions such as pH, temperature, etc. while mentioning the stability of the complex. In brief, it is not necessary for a stable complex to be inert and an unstable complex to be labile.

*3.1.2 Size of central metal cation*

*DOI: http://dx.doi.org/10.5772/intechopen.90894*

*Stability of Metal Complexes*

The order of size of dipositive ions is

**3.2 Nature of ligands**

with NH3, H2O, and F� is:

in the following order:

**3.3 The chelate effect**

chloride, cyanide, and water.

ligand 2,2<sup>0</sup>

**29**

CN�, alkene, phenanthroline, etc.

The stability of metal complex increases with decrease in size of the metal

Ba2<sup>þ</sup> < Sr2<sup>þ</sup> <Ca2<sup>þ</sup> < Mg2<sup>þ</sup> < Mn2<sup>þ</sup> < Fe2<sup>þ</sup> < Co2<sup>þ</sup> < Ni2<sup>þ</sup> <Cu2<sup>þ</sup> >Zn2<sup>þ</sup>

This order of stability is also in good agreement with the charge to radius ratio concept because the radii decrease from Ba2+ to Cu2+ and then increased to Zn2+.

Ba2<sup>þ</sup> > Sr2<sup>þ</sup> >Ca2<sup>þ</sup> > Mg2<sup>þ</sup> > Mn2<sup>þ</sup> > Fe2<sup>þ</sup> > Co2<sup>þ</sup> > Ni2<sup>þ</sup> >Cu2<sup>þ</sup> <Zn2<sup>þ</sup>

Basic character of ligands: The greater is the basic character of ligand, the more easily it can donate its lone pair of electrons to the central metal ion and hence greater is the complex stability. In 3D-series metal ion, order of stability of complex

NH3 > H2O >F�

The nature of metal-ligand bond also affects the stability of metal complexes. The higher the covalent character, the greater will be the complex stability. For example, the stabilities of silver complexes have different halide ligands which are

� > AgCl2

Ligands having vacant p- or d-orbital tend to form π bond and hence form stable complexes with metals. Ligands that are capable of forming such π bond are CO,

The chelate effect is that the complexes resulting from coordination of metal ions with the chelating ligand are thermodynamically much more stable than the complexes with non-chelating ligands [10, 11]. Chelating ligands are molecules which can bind to single metal ion through several bonds and are also called as multidentate ligands. Simple (and common) examples include ethylenediamine and oxalate. Non-chelating ligands are ligands that bond to just one site, such as

The chelate effect can be understood by comparing the reaction of a metal ion,


Another such comparison can be made between coordination behavior of chelating 1,2-diaminoethane (ethylenediamine = en) and monodentate ammonia. Such

respectively, with a chelating ligand and with a monodentate ligand having similar/comparable donating groups. During the comparison study, the number of coordination should be maintained equal in both the cases, for example, the value obtained while adding a bidentate ligand is compared with the value obtained for two monodentate ligands. For example, coordination of metal ion with chelating

� > AgF2

�

cations. For M2+ ions, the general trend in stability for complexes is

This trend in stability is known as Irving-Williams series.

AgI2

� > AgBr2

Consider the three complexes [Ni(CN)4] <sup>2</sup>, [Mn(CN)6] <sup>3</sup>, and [Cr(CN)6] 3. All the complexes are thermodynamically stable, but kinetically they behave in a different manner. The rate of exchange can be measured when carbon-14-labelled cyanide ions are reacted with metal complexes in solution. It indicates that [Ni(CN)4] <sup>2</sup> is labile, [Mn(CN)6] <sup>3</sup> is less labile, and [Cr(CN)6] <sup>3</sup> is inert and proves that not all stable complexes are inert and vice versa.

$$\begin{array}{ccccc} \text{ $ [\mathsf{Nil}(\mathsf{CN})\_4\]\_2$ ^{2-} + 4} & \text{ $ 4}^{\mathsf{T4}}\mathsf{CN}^- & \xrightarrow{\mathsf{V8}\mathsf{N}\mathsf{V}} & \text{[$ \mathsf{Nil}(\mathsf{"}\mathsf{CN})\_4\]\_2^{2-} + 4} & \text{[ $\mathsf{Nil}(\mathsf{"}\mathsf{CN})\_4\}\_2^{2-} \\\\ & & & \text{t}\_{1\mathsf{N2}} = \mathsf{30}\mathsf{sec} \\\\ \text{[$ \mathsf{M}(\mathsf{CN})\_6\}\_2^{3-} + 6\ \mathsf{T}\mathsf{N}\mathsf{N} & \xrightarrow{\mathsf{s}\mathsf{I}\mathsf{s}\mathsf{C}\mathsf{N}\mathsf{N} \\\\ & & \text{t}\_{1\mathsf{N2}} = \mathsf{T}\mathsf{h}\mathsf{r} \end{array}$$

$$\begin{array}{ccccc} \text{[Cl(CN)\$^{3}\$]}^{3-} + & \text{[} & \text{44\$^{2}\$CN}\$^{-} & \text{[} & \text{\$\text{\$Cl\$^{(4}\$CN)\$^{(8}\$)}\$} \\\\ & & & \text{t\$\_{12}\$=24\$days} \end{array}$$

## **3. Factors affecting the stability of metal complexes**

There are several factors that can affect the stability of the metal complexes [2, 5, 8, 9], which include:


#### **3.1 Nature of central metal ion**

#### *3.1.1 Charge on metal cation*

In metal cations, higher oxidation state forms more stable complex than lower oxidation states with ligands such as NH3, H2O, etc. Even few exceptions are there like CO, PMe3, o-phenanthroline, bipyridyl, CN, which form more stable complex with lower oxidation state metals.

#### *3.1.2 Size of central metal cation*

conditions, and it is always recommended to specify the conditions such as pH, temperature, etc. while mentioning the stability of the complex. In brief, it is not necessary for a stable complex to be inert and an unstable complex to be labile.

All the complexes are thermodynamically stable, but kinetically they behave in a different manner. The rate of exchange can be measured when carbon-14-labelled

cyanide ions are reacted with metal complexes in solution. It indicates that

proves that not all stable complexes are inert and vice versa.

**3. Factors affecting the stability of metal complexes**

[2, 5, 8, 9], which include:

2.Nature of the ligand

3.Chelating effect

4.Macrocyclic effect

5.Resonance effect

1.Nature of the central metal ion

6. Steric effect or steric hindrance

**3.1 Nature of central metal ion**

with lower oxidation state metals.

*3.1.1 Charge on metal cation*

**28**

There are several factors that can affect the stability of the metal complexes

In metal cations, higher oxidation state forms more stable complex than lower oxidation states with ligands such as NH3, H2O, etc. Even few exceptions are there like CO, PMe3, o-phenanthroline, bipyridyl, CN, which form more stable complex

<sup>2</sup>, [Mn(CN)6]

<sup>3</sup> is less labile, and [Cr(CN)6]

<sup>3</sup>, and [Cr(CN)6]

<sup>3</sup> is inert and

3.

Consider the three complexes [Ni(CN)4]

*Stability and Applications of Coordination Compounds*

<sup>2</sup> is labile, [Mn(CN)6]

[Ni(CN)4]

The stability of metal complex increases with decrease in size of the metal cations. For M2+ ions, the general trend in stability for complexes is

$$\text{Ca}^{2+} < \text{Sr}^{2+} < \text{Ca}^{2+} < \text{Mg}^{2+} < \text{Mn}^{2+} < \text{Fe}^{2+} < \text{Co}^{2+} < \text{Ni}^{2+} < \text{Cu}^{2+} > \text{Zn}^{2+} $$

This trend in stability is known as Irving-Williams series.

This order of stability is also in good agreement with the charge to radius ratio concept because the radii decrease from Ba2+ to Cu2+ and then increased to Zn2+. The order of size of dipositive ions is

$$\text{Ca}^{2+} > \text{Sr}^{2+} > \text{Ca}^{2+} > \text{Mg}^{2+} > \text{Mn}^{2+} > \text{Fe}^{2+} > \text{Co}^{2+} > \text{Ni}^{2+} > \text{Cu}^{2+} < \text{Zn}^{2+} $$

#### **3.2 Nature of ligands**

Basic character of ligands: The greater is the basic character of ligand, the more easily it can donate its lone pair of electrons to the central metal ion and hence greater is the complex stability. In 3D-series metal ion, order of stability of complex with NH3, H2O, and F� is:

$$\text{NH}\_3 > \text{H}\_2\text{O} > \text{F}^-$$

The nature of metal-ligand bond also affects the stability of metal complexes. The higher the covalent character, the greater will be the complex stability. For example, the stabilities of silver complexes have different halide ligands which are in the following order:

$$\text{AgI}\_2^- > \text{AgBr}\_2^- > \text{AgCl}\_2^- > \text{AgF}\_2^-$$

Ligands having vacant p- or d-orbital tend to form π bond and hence form stable complexes with metals. Ligands that are capable of forming such π bond are CO, CN�, alkene, phenanthroline, etc.

#### **3.3 The chelate effect**

The chelate effect is that the complexes resulting from coordination of metal ions with the chelating ligand are thermodynamically much more stable than the complexes with non-chelating ligands [10, 11]. Chelating ligands are molecules which can bind to single metal ion through several bonds and are also called as multidentate ligands. Simple (and common) examples include ethylenediamine and oxalate. Non-chelating ligands are ligands that bond to just one site, such as chloride, cyanide, and water.

The chelate effect can be understood by comparing the reaction of a metal ion, respectively, with a chelating ligand and with a monodentate ligand having similar/comparable donating groups. During the comparison study, the number of coordination should be maintained equal in both the cases, for example, the value obtained while adding a bidentate ligand is compared with the value obtained for two monodentate ligands. For example, coordination of metal ion with chelating ligand 2,2<sup>0</sup> -bipyridine can be compared with that of monodentate pyridine ligand. Another such comparison can be made between coordination behavior of chelating 1,2-diaminoethane (ethylenediamine = en) and monodentate ammonia. Such

comparison studies revealed that the metal complex formed from chelating ligands are thermodynamically more stable than the complex formed from monodentate ligand. For example, formation of complexes from hydrated cadmium ion, [Cd(H2O)4] 2+ with methylamine (CH3NH2), ethylenediamine (en) and triethylenetetramine (trien), and their stability is in the following order:

**3.6 Steric effect**

*Stability of Metal Complexes*

*DOI: http://dx.doi.org/10.5772/intechopen.90894*

shown in **Figure 3**.

stability of the complex.

**Figure 3.**

**3.7 Crystal field stabilization energy (CFSE)**

*Chelating complexes of Ni(II) ion showing steric effect.*

**4. Determination of stability constants**

**4.1 Spectroscopic methods**

**31**

The presence of bulky substituents in the ligands can affect the stability of the metal complex, and this type of destabilization of metal complex due to bulkiness of the substituent is called as steric effect [13]. For example, consider the ligand 8-hydroxy quinoline and its methyl substituted derivative 2-methyl-8-hydroxy quinolone. Both are bidentate ligands and form chelated complexes with Ni2+ ion as

The complex (II) is less stable than complex (I) because of bulky group attached to an atom adjacent to donor atom which cause a steric hindrance and lower the

The crystal field stabilization energy (CFSE) is one of the most important factors that decides the stability of the metal complexes. CFSE is the stability that arises when a metal ion coordinates to a set of ligands, which is due to the generation of a crystal field by the ligands. Thus, a higher value of CFSE means that the complex is thermodynamically stable and kinetically inert. Some of the notable examples of complexes that have high CFSE are low spin 5d6 complexes of Pt4+ and Ir3+ and square planar 5d8 complexes of Pt2+. All these complexes are thermodynamically stable and kinetically

The determination of metal complexes involves several methods including spectroscopic and potentiometric methods. The determination of stability constant is very significant to understand the role and behavior of ligand(s) in stabilizing the metal complexes and found applications in the fields of biology, environmental study, metallurgy, food chemistry, and many other industrial processes. Some of the methods that are used for the determination of stability constants are given as follows.

UV-Vis spectroscopic technique has been used to determine the stability constant and composition of a complex [14]. The formation of metal complex is indicated by the change in absorbance in the UV-Vis spectroscopy. The relationship

A ¼ ԑ*:*c*:*l*:*

between absorbance (A) and concentration is given by Beer's law as shown.

inert, which undergo ligand substitution reactions extremely slowly [3, 4].

$$\left[\text{Cd}(\text{CH}\_3\text{NH}\_2)\_4\right]^{2+} < \left[\text{Cd}(\text{en})\right]^{2+} < \left[\text{Cd}(\text{trien})\right]^{2+} $$

### **3.4 Macrocyclic effect**

A macrocyclic ligand is a cyclic molecule that contains nine or more atoms in the cyclic structure and has three or more potential donor atoms which can coordinate to the metal ion. It has been observed that the stability of metal complexes in the presence of macrocyclic ligand of appropriate size is higher than the stability of complexes coordinated to open-ended multidentate chelating ligands. Some notable examples of macrocyclic ligands include cyclic crown polyether, heme, etc. [12].

#### **3.5 Resonance effect**

Resonance increases the stability of the complexes. For example, acetylacetonate anion ligand shows resonance, and as a result it forms stable complexes upon reacting with metal ion (**Figure 1**). The ligand-metal π bonding increases the delocalization of electrons compared to free enolate as shown below and leads to increased stability (**Figure 2**).

**Figure 1.** *Resonance structure of acetonylacetonate ligand.*

**Figure 2.** *Acetonylacetonate-metal complex.*
