**4. Crystal structures, PXRD, and some interesting properties of the metal complexes**

#### **4.1 Crystal structure and catalytic properties of nickel(II) complex with the mono-imine Schiff base ligand (L<sup>1</sup> )**

The neutral monomeric complex [Ni(L<sup>1</sup> )(NCS)2] (**1**) has been found to possess octahedral geometry where central Ni(II) is coordinated by the neutral ligand **L<sup>1</sup>** with tetradentate N4 binding mode and two N-bonded thiocyanate ions occupying the *cis-*position. **Figure 12** depicts the molecular structure of the nickel (II) complex **1** [29].

The coordination environment around the nickel(II) ion is surrounded by N6 fashion (four N from ligand and two N from thiocyanate ions) tending towards distorted octahedral geometry. The Ni2+ center is not lying exactly within the equatorial plane of N4 moiety, and unequal axial and equatorial bond distances (2.112 Å and 2.072 Å, respectively) confirm the distortion. The non-coordinated O–H groups on the ligand L<sup>1</sup> are engaged in H-bonding interactions with thiocyanate S atoms (**Figure 6**) which lead to 1D supramolecular sheet-like arrangement (**Figure 13**). These H-bonding interactions lead to OS separations of 3.132 Å and play prominent role in crystal packing.

#### *4.1.1 Catalytic activity of complex 1*

Analytical grade reagents and freshly distilled solvents, viz., water, acetonitrile, methanol, and dichloromethane, were used to check the catalytic activity. The oxidation reaction was carried out in liquid phase under vigorous stirring in twonecked round bottom flask fitted with a water condenser and placed in an oil bath at 60°C. Substrate (5 mmol) was taken in 10 ml solvent(s) for different sets of reactions along with 5 mg catalyst, to which 10 mmol of *tert*-Butyl hydrogen peroxide

**Figure 12.**

An organometallic complex, [PhHg(HL<sup>3</sup>

*Stability and Applications of Coordination Compounds*

were obtained from their chloroform solutions [34, 35].

)2]Cl.3H2O (**complex 6**), and [Fe(HL<sup>3</sup>

recrystallized from chloroform [37].

plexes **9–11** is [MII(HL3

**Figure 10.**

**Figure 11.**

**6**

The Cd(II), Cr(III), and Fe(III) complexes, i.e., [Cd(HL<sup>3</sup>

the complexes were obtained from their chloroform solutions [35, 36].

*Scheme of formation of metal complexes using H2L<sup>3</sup> ligand indicating different binding pattern.*

*Scheme of preparation of other metal complexes using H2L3 ligand.*

plex [Zn(HL<sup>3</sup>

[Cr(HL<sup>3</sup>

)] (**complex 3**), and the zinc(II) com-

)2]Cl.3H2O (**complex 7**), were

)2)] (**complex 5**),

)(OAc)(H2O)] (**complex 4**) were prepared by the gentle reflux of the

equimolar quantities of diacetylmonoxime (0.5 g, 5 mmol) and morpholine Nthiohydrazide (0.8 g, 5 mmol) (**Figure 10**) in the presence of the respective metal salts, [PhHg(OAc)] (1.68 g, 5 mmol) and [Zn(OAc)2].2H2O (1.1 g, 5 mmol). Yields of the complexes are 65% (**3**) and 75% (**4**). The pure single crystals of the complexes

prepared by the gentle reflux of the mixture of diacetylmonoxime (0.5 g, 5 mmol) and morpholine N-thiohydrazide (0.8 g, 5 mmol) and the respective metal salts ([Cd(OAc)2].2H2O (0.67 g, 2.5 mmol), CrCl3.6H2O (0.67 g, 2.5 mmol), and FeCl3.6H2O (0.68 g, 2.5 mmol) in the ratio of 2:2:1 (**Figure 10**). Yields of the complexes had been recorded as 75% for each complex. The pure single crystals of

Another zinc(II) complex (**complex 8**) was synthesized by the reflux of the mixture of diacetylmonoxime (0.5 g, 5 mmol) and morpholine N-thiohydrazide (0.8 g, 5 mmol) in the presence of Zn(OAc)2.2H2O (0.55 g, 2.5 mmol) in the molar ratio of 2:2:1 in water–methanol (1:1, v/v) mixture (**Figure 11**). The yellow colored complex of composition [Zn(HL3)2]. 2H2O had been separated with 60% yield and

The complexes of Ni(II) (**complex 9**), Co(II) (**complex 10**), and Cu(II) (**complex 11**) were synthesized by gentle reflux of three sets of equimolar quantities of diacetylmonoxime (0.5 g, 5 mmol) and morpholine N-thiohydrazide (0.8 g, 5 mmol) in the presence of the metal salts (Ni(OAc)2.4H2O (1.24 g, 5 mmol), Co (OAc)2.4H2O (1.25 g, 5 mmol), and Cu(OAc)2.4H2O (1.28 g, 5 mmol), respectively) in ethanol (50 ml) (**Figure 11**) [29]. General molecular formula of the metal com-

)(OAc)] where M stands for Ni, Co, and Cu.

*Perspective view of complex 1 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

**Figure 13.** *SO interaction in the crystal of complex 1 (hydrogen atoms are omitted for clarity).*

(TBHP, 70% aq.) was added as oxidant immediately before the start of the reaction. Aliquots of the reaction mixture were withdrawn at various time intervals, and the products were analyzed by using Agilent 6890A gas chromatograph equipped with HP-1 capillary column and FID. All reaction products were identified and estimated by using an Agilent GC–MS (QP-5050 model).

The Hg(II) atom remains 0.027(1) Å above the plane. Due to the contribution of electron flow from mercury to the π\* orbitals of the phenyl group, the Hg–C bond distance is found shorter than that in the analogous methylmercury(II) compound, where no such electron drifting is observed. The C–S bond gets partial double bond character in the complex, similar to related thiosemicarbazonates of methylmercury

(II) and dimethylthallium(III). It is interesting to note that there is no

*4.3.2 Crystal structure of the zinc(II) complex (4)*

deprotonated thiol sulfur of ligand [(HL<sup>3</sup>

*DOI: http://dx.doi.org/10.5772/intechopen.90171*

[Zn(HL<sup>3</sup>

**Figure 16.**

**Figure 15.**

**Figure 17.**

**9**

*supramolecular dimer.*

intermolecular π–π interaction between the phenyl rings. But a weak interaction between C(8)–H(8A) and a π group (phenyl ring) links the two phenylmercury molecules into a supramolecular dimer having a C–H π synthon (**Figure 16**) having characteristic HCg distance 2.84 Å, where Cg is the midpoint of the phenyl ring.

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands…*

The X-ray crystal structure shows that due to constrained ligand structure, the

)

torial plane, while the axial position is occupied by oxygen (O4) of coordinated

*Weak interaction (C–Hπ) between the π-electrons of the phenyl ring with the H atom of ligand in a*

*Perspective view of complex 3 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

*Perspective view of complex 4 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

)(OAc)(H2O)].H2O complex **4** (**Figure 17**) has adopted distorted square pyramidal geometry (τ = 0.05) in which the oxime nitrogen, azomethine nitrogen,

], and acetato oxygen defines the equa-

The [Ni(L1 )(NCS)2] (**1**) exhibits good catalytic activity towards the oxidation of styrene in the presence of TBHP which leads to the formation of benzaldehyde as a major product with small amount of styrene oxide as shown in **Figure 14**. By varying solvents, reaction time, temperature, and substrate-to-oxidant ratio, a successful conversion of 70% of the styrene was reached at the optimum condition [29].

#### **4.2 PXRD structure and solid-state properties of oxovanadium(IV) complex with the di-imine Schiff base ligand (H2L<sup>2</sup> )**

Despite our repeated attempts and best effort, the single crystal of the oxovanadium(IV) complex with the ligand H2L<sup>2</sup> could not be grown, and it led us to carry out the powder X-ray diffraction (PXRD) study to characterize the oxovanadium(IV) complex **2**. The composition of the complex is [VO(L<sup>2</sup> )], where the ligand H2L<sup>2</sup> acts as a dibasic [(L<sup>2</sup> ) <sup>2</sup>] tetradentate arrangement. The result of the PXRD shows that the unit cell is monoclinic in nature and the cell parameters of the complex **2** are a = 16.0619 Å, b = 11.5517 Å, c = 12.7262 Å, and β = 122.427°; and the cell volume (V) is 1993.05 Å3 .

Other solid-state properties, viz., electrical, optical, and thermal properties of the complex [VO(L<sup>2</sup> )], have also been studied [24]. The complex is electrically an insulator at room temperature; however, the conductivity is increased as the temperature increases from 330 K, indicating the semiconducting nature of the complex. It behaves as an n-type semiconductor, and the semiconducting behavior of the oxovanadium(IV) complex with the dibasic Schiff base ligand was substantiated by the extended conjugated chemical structure. The said properties are discussed in detail in the following sections.

#### **4.3 Crystal structures and properties of metal complexes with thio-hydrazone Schiff base ligand (H2L<sup>3</sup> )**

#### *4.3.1 Crystal structure of organometallic phenylmercury(II) complex (3)*

The organometallic phenylmercury(II) compound **3** crystallized in triclinic form to give a tricoordinated T-shaped geometry (**Figure 15**). Here the monobasic ligand [(HL<sup>3</sup> ) ] coordinated the central mercury in a bidentate manner through S and N (N of hydrazone function) and the third coordination of the mercury was satisfied by the C atom of phenyl group of the metal precursor. Thus, by choosing the phenylmercury system, the ligand H2L3 has been compelled to act in a bidentate mode leaving another N donor (the oxime N) uncoordinated [35].

**Figure 14.** *Oxidation of styrene using [Ni(L1 )(NCS)2] as catalyst.*

#### *Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands… DOI: http://dx.doi.org/10.5772/intechopen.90171*

The Hg(II) atom remains 0.027(1) Å above the plane. Due to the contribution of electron flow from mercury to the π\* orbitals of the phenyl group, the Hg–C bond distance is found shorter than that in the analogous methylmercury(II) compound, where no such electron drifting is observed. The C–S bond gets partial double bond character in the complex, similar to related thiosemicarbazonates of methylmercury (II) and dimethylthallium(III). It is interesting to note that there is no intermolecular π–π interaction between the phenyl rings. But a weak interaction between C(8)–H(8A) and a π group (phenyl ring) links the two phenylmercury molecules into a supramolecular dimer having a C–H π synthon (**Figure 16**) having characteristic HCg distance 2.84 Å, where Cg is the midpoint of the phenyl ring.

#### *4.3.2 Crystal structure of the zinc(II) complex (4)*

The X-ray crystal structure shows that due to constrained ligand structure, the [Zn(HL<sup>3</sup> )(OAc)(H2O)].H2O complex **4** (**Figure 17**) has adopted distorted square pyramidal geometry (τ = 0.05) in which the oxime nitrogen, azomethine nitrogen, deprotonated thiol sulfur of ligand [(HL<sup>3</sup> ) ], and acetato oxygen defines the equatorial plane, while the axial position is occupied by oxygen (O4) of coordinated

**Figure 15.** *Perspective view of complex 3 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

**Figure 16.**

(TBHP, 70% aq.) was added as oxidant immediately before the start of the reaction. Aliquots of the reaction mixture were withdrawn at various time intervals, and the products were analyzed by using Agilent 6890A gas chromatograph equipped with HP-1 capillary column and FID. All reaction products were identified and estimated

of styrene in the presence of TBHP which leads to the formation of benzaldehyde as a major product with small amount of styrene oxide as shown in **Figure 14**. By varying solvents, reaction time, temperature, and substrate-to-oxidant ratio, a successful conversion of 70% of the styrene was reached at the optimum

**4.2 PXRD structure and solid-state properties of oxovanadium(IV) complex**

)

Other solid-state properties, viz., electrical, optical, and thermal properties of

insulator at room temperature; however, the conductivity is increased as the temperature increases from 330 K, indicating the semiconducting nature of the complex. It behaves as an n-type semiconductor, and the semiconducting behavior of the oxovanadium(IV) complex with the dibasic Schiff base ligand was substantiated by the extended conjugated chemical structure. The said properties are discussed in

**4.3 Crystal structures and properties of metal complexes with thio-hydrazone**

The organometallic phenylmercury(II) compound **3** crystallized in triclinic form to give a tricoordinated T-shaped geometry (**Figure 15**). Here the monobasic ligand

] coordinated the central mercury in a bidentate manner through S and N (N of hydrazone function) and the third coordination of the mercury was satisfied by the C atom of phenyl group of the metal precursor. Thus, by choosing the phenylmercury system, the ligand H2L3 has been compelled to act in a bidentate

Despite our repeated attempts and best effort, the single crystal of the oxovanadium(IV) complex with the ligand H2L<sup>2</sup> could not be grown, and it led us to carry out the powder X-ray diffraction (PXRD) study to characterize the oxovanadium(IV) complex **2**. The composition of the complex is [VO(L<sup>2</sup>

result of the PXRD shows that the unit cell is monoclinic in nature and the cell parameters of the complex **2** are a = 16.0619 Å, b = 11.5517 Å, c = 12.7262 Å, and

)(NCS)2] (**1**) exhibits good catalytic activity towards the oxidation

**)**

.

)], have also been studied [24]. The complex is electrically an

)],

<sup>2</sup>] tetradentate arrangement. The

by using an Agilent GC–MS (QP-5050 model).

*Stability and Applications of Coordination Compounds*

**with the di-imine Schiff base ligand (H2L<sup>2</sup>**

where the ligand H2L<sup>2</sup> acts as a dibasic [(L<sup>2</sup>

β = 122.427°; and the cell volume (V) is 1993.05 Å3

**)**

*4.3.1 Crystal structure of organometallic phenylmercury(II) complex (3)*

mode leaving another N donor (the oxime N) uncoordinated [35].

*)(NCS)2] as catalyst.*

The [Ni(L1

condition [29].

the complex [VO(L<sup>2</sup>

[(HL<sup>3</sup> )

**Figure 14.**

**8**

*Oxidation of styrene using [Ni(L1*

detail in the following sections.

**Schiff base ligand (H2L<sup>3</sup>**

*Weak interaction (C–Hπ) between the π-electrons of the phenyl ring with the H atom of ligand in a supramolecular dimer.*

water molecule. The zinc(II) ion is 0.3766 (8) Å out of the basal plane towards the oxygen (O4) of coordinated water molecule. The deprotonation of the ligand results in extensive delocalization of charge, and as a consequence the C–S bond length (1.7459(18) Å) is much closer in length to a C–S single bond (1.82 Å), which has partial double bond character [34].

Due to intramolecular hydrogen bonding, the Zn–N(3) (azomethine) distance is slightly shorter than Zn–N(4) (oxime) distance. Here the H(1) of coordinated water molecule is hydrogen bonded to morpholinic oxygen O(1), while H(2) is hydrogen bonded to acetate oxygen O(3), and such H-bonding forms the 1D supramolecular framework diagonal to the *ab* plane (**Figure 18**).

#### *4.3.3 Crystal structure of the cadmium(II) complex (5)*

The cadmium(II) complex [Cd(HL<sup>3</sup> )2] (**5**) crystallized in a triclinic space group P1̅(2). The central Cd (II) ion is coordinated by two units of monobasic ligand [(HL<sup>3</sup> ) ] forming trapezoidal bipyramid (**Figure 19**), where the equatorial plane was formed by thiol S1, azomethine N3, oxime N4 of one Schiff base unit, and azomethine N7 atoms of another unit (second) and the axial positions are being coordinated by oxime N8 and thiol S2 atoms of the second Schiff base unit. The C–S bond gets partial double bond character in the complex and is shorter than that in phenylmercury(II) complex (3). The C–N average bond distance confirms double bond character, and deprotonation of the hydrazinic NH proton induces double bond character [35].

Complex **5** forms a 2D supramolecular sheet in the *bc* plane through hydrogen bonding interactions (O–HO and C–HO) (**Figure 20**). The average DA separations and the average D–HA angle is 145°. These supramolecular sheets stack along (100) the plane. There are no directional weak force interactions (π–π, C–Hπ, hydrogen bonding) operating between the two sheets. Therefore, an intermolecular van der Waals interaction might be responsible for such supramolecular stacking. The methyl terminals of the ligand from adjacent sheets face each other in this arrangement.

interaction has been reported. The three H atoms (H9A, H9B, H9C) of the methyl group of the propylidene moiety form a triangular plane which is supported by the angle ∠C8–C9O4\* of 172.0°. The O4 atom of the oxime is directed towards the center of the plane formed by three H atoms (H9A, H9B, H9C) of the methyl group. The C9–H9AO4\*, C9–H9BO4\*, and C9–H9CO4\* angles are 89.0°, 83.8°, and 97.6°, respectively, which are less than the generally accepted ∠C–HO angles (110°), and such short bond angles might be responsible for the greater CO attraction. Thus, the cooperative effect of the three individual interactions acts on the face of the plane of the three methyl hydrogens of the C9 atom due to their high acidic character. This very unusual and nonconventional interaction was termed as

*)2] (5).*

*)2] (5).*

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands…*

a CH3O interaction and proved to be a good supramolecular synthon. The

forms slightly distorted octahedral geometry with [CrN4S2] core having meridional conformation via the oximino N, imine N, and thiol S atoms (**Figure 22**), i.e., the two monobasic tridentate L3 ligands act as N,N,S donor during formation of the

(**6**) with pseudo-octahedral geometry except that the equatorial plane is formed by two oximino N atoms and two thiol S atoms and the axial positions are occupied by two imine N atoms of the ligand (**Figure 23**). The Fe–N(imine) bond lengths are significantly shorter than the Fe–N(oximino) bond lengths because azomethine

)2]Cl3H2O (**7**) is quite similar to that of complex

)2]Cl3H2O (**6**)

The X-Ray single-crystal analysis revealed that complex [Cr(L<sup>3</sup>

trifurcated H-bonding interactions are shown in (**Figure 21**).

*4.3.4 Crystal structure of the chromium(III) complex (6)*

*Trifurcated H-bonding interactions [35] in[Cd(HL<sup>3</sup>*

*Two-dimensional network through H-bonding in [Cd(HL3*

*DOI: http://dx.doi.org/10.5772/intechopen.90171*

*4.3.5 Crystal structure of the iron(III) complex (7)*

The X-ray structure of [Fe(L<sup>3</sup>

complex [36].

**11**

**Figure 21.**

**Figure 20.**

One of the most interesting parts of the structure is that the crystal structure possesses a remarkably short intermolecular C(sp<sup>3</sup> )O(sp<sup>3</sup> ) contact [C9O4\* 2.958 (3) Å]. An interesting packing force using an uncommon C(sp3 )O(sp<sup>3</sup> )

**Figure 18.** *One-dimensional network through H-bonding in [Zn(HL<sup>3</sup> )(OAc)(H2O)].H2O.*

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands… DOI: http://dx.doi.org/10.5772/intechopen.90171*

**Figure 20.** *Two-dimensional network through H-bonding in [Cd(HL3 )2] (5).*

water molecule. The zinc(II) ion is 0.3766 (8) Å out of the basal plane towards the oxygen (O4) of coordinated water molecule. The deprotonation of the ligand results in extensive delocalization of charge, and as a consequence the C–S bond length (1.7459(18) Å) is much closer in length to a C–S single bond (1.82 Å), which has

Due to intramolecular hydrogen bonding, the Zn–N(3) (azomethine) distance is slightly shorter than Zn–N(4) (oxime) distance. Here the H(1) of coordinated water molecule is hydrogen bonded to morpholinic oxygen O(1), while H(2) is hydrogen bonded to acetate oxygen O(3), and such H-bonding forms the 1D supra-

P1̅(2). The central Cd (II) ion is coordinated by two units of monobasic ligand

was formed by thiol S1, azomethine N3, oxime N4 of one Schiff base unit, and azomethine N7 atoms of another unit (second) and the axial positions are being coordinated by oxime N8 and thiol S2 atoms of the second Schiff base unit. The C–S bond gets partial double bond character in the complex and is shorter than that in phenylmercury(II) complex (3). The C–N average bond distance confirms double bond character, and deprotonation of the hydrazinic NH proton induces double

] forming trapezoidal bipyramid (**Figure 19**), where the equatorial plane

Complex **5** forms a 2D supramolecular sheet in the *bc* plane through hydrogen

One of the most interesting parts of the structure is that the crystal structure

*Perspective view of complex 5 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

)O(sp<sup>3</sup>

*)(OAc)(H2O)].H2O.*

) contact [C9O4\* 2.958

)

)O(sp<sup>3</sup>

bonding interactions (O–HO and C–HO) (**Figure 20**). The average DA separations and the average D–HA angle is 145°. These supramolecular sheets stack along (100) the plane. There are no directional weak force interactions (π–π, C–Hπ, hydrogen bonding) operating between the two sheets. Therefore, an intermolecular van der Waals interaction might be responsible for such supramolecular stacking. The methyl terminals of the ligand from adjacent sheets face each

)2] (**5**) crystallized in a triclinic space group

partial double bond character [34].

*Stability and Applications of Coordination Compounds*

molecular framework diagonal to the *ab* plane (**Figure 18**).

*4.3.3 Crystal structure of the cadmium(II) complex (5)*

possesses a remarkably short intermolecular C(sp<sup>3</sup>

*One-dimensional network through H-bonding in [Zn(HL<sup>3</sup>*

(3) Å]. An interesting packing force using an uncommon C(sp3

The cadmium(II) complex [Cd(HL<sup>3</sup>

[(HL<sup>3</sup> )

bond character [35].

other in this arrangement.

**Figure 18.**

**Figure 19.**

**10**

**Figure 21.** *Trifurcated H-bonding interactions [35] in[Cd(HL<sup>3</sup> )2] (5).*

interaction has been reported. The three H atoms (H9A, H9B, H9C) of the methyl group of the propylidene moiety form a triangular plane which is supported by the angle ∠C8–C9O4\* of 172.0°. The O4 atom of the oxime is directed towards the center of the plane formed by three H atoms (H9A, H9B, H9C) of the methyl group. The C9–H9AO4\*, C9–H9BO4\*, and C9–H9CO4\* angles are 89.0°, 83.8°, and 97.6°, respectively, which are less than the generally accepted ∠C–HO angles (110°), and such short bond angles might be responsible for the greater CO attraction. Thus, the cooperative effect of the three individual interactions acts on the face of the plane of the three methyl hydrogens of the C9 atom due to their high acidic character. This very unusual and nonconventional interaction was termed as a CH3O interaction and proved to be a good supramolecular synthon. The trifurcated H-bonding interactions are shown in (**Figure 21**).

#### *4.3.4 Crystal structure of the chromium(III) complex (6)*

The X-Ray single-crystal analysis revealed that complex [Cr(L<sup>3</sup> )2]Cl3H2O (**6**) forms slightly distorted octahedral geometry with [CrN4S2] core having meridional conformation via the oximino N, imine N, and thiol S atoms (**Figure 22**), i.e., the two monobasic tridentate L3 ligands act as N,N,S donor during formation of the complex [36].

### *4.3.5 Crystal structure of the iron(III) complex (7)*

The X-ray structure of [Fe(L<sup>3</sup> )2]Cl3H2O (**7**) is quite similar to that of complex (**6**) with pseudo-octahedral geometry except that the equatorial plane is formed by two oximino N atoms and two thiol S atoms and the axial positions are occupied by two imine N atoms of the ligand (**Figure 23**). The Fe–N(imine) bond lengths are significantly shorter than the Fe–N(oximino) bond lengths because azomethine

nitrogen is a stronger base than the oximino nitrogen. The deprotonation of the ligands results in extensive delocalization of charge, and thus the C–S bond gets partial double bond character [36].

The most interesting part of the crystalline structure of complex [Cr(L<sup>3</sup> )2] Cl3H2O (**6**) and [Fe(L<sup>3</sup> )2]Cl3H2O (**7**) is that both the complexes form a 1D supramolecular chain along the *c*-axis through C–HO interactions between each molecule (**Figure 24**).

The 1D chains are arranged in parallel direction to form a supramolecular host having channels along the *c*-axis (**Figure 25**). Such channels are filled by selfassembled "water-chloride" clusters having chair conformation (**Figure 25**).

It is proven by the crystal structure analysis that there are three crystals of water molecules per formula unit of **complexes 6** and **7** along with chloride ion, and these are involved in forming the infinite "water-chloride" cluster [36] with a chair

#### **Figure 22.** *Perspective view of complex 6 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

conformation. The three water molecules and a chloride molecule form a fourmembered cyclic motif, and such three adjacent four-membered rings form the

*Supramolecular channels and "water-chloride" cluster along the c-axis in complexes (6) and (7).*

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands…*

*DOI: http://dx.doi.org/10.5772/intechopen.90171*

The variable temperature (2.5–300 K) magnetic moment study shows the temperature dependence of the magnetic susceptibility. The χ<sup>m</sup> values at 2.5 and 300 K

respectively. The detailed study shows that the magnetic moment value consists of a superimposition of both the low-spin and high-spin states. At very low temperature, the 1-D supramolecular species which is formed by strong intermolecular C– HO interactions and the cooperative interactions with the "water-chloride" cluster between mononuclear spin crossover (SCO) sites stabilize the low-spin state, and thus the high-spin contribution decreases to 21%, and the low-spin contribution increases to 79%. Thus, such variable temperature magnetic behavior may be due to

The Mössbauer spectroscopic study also supports that a spin crossover phenomenon exists in the iron(III) complex (**7**). Both the spin states, low (S = 1/2) and high (S = 5/2), exist in room temperature (300 K) as well as in very low temperature at 20 K. The population density of the electrons decreases from high-spin state to the

a continuous S = 1/2 to 5/2 spin crossover phenomenon of iron centers [36].

low-spin state with decreasing temperature (**Table 1**).

, while the μeff values are 2.61 and 3.46 B.M.,

*4.3.5.1 Magnetic property and Mössbauer spectroscopy of complex (7)*

*Chair conformation formed by "water-chloride" cluster in complexes (6) and (7).*

chair conformation (**Figure 26**).

**Figure 25.**

**Figure 26.**

**13**

are 0.33 and 0.005 cm <sup>3</sup> mol<sup>1</sup>

**Figure 24.** *One-dimensional network through C–HO interactions in complexes (6) and (7).*

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands… DOI: http://dx.doi.org/10.5772/intechopen.90171*

#### **Figure 25.**

nitrogen is a stronger base than the oximino nitrogen. The deprotonation of the ligands results in extensive delocalization of charge, and thus the C–S bond gets

The most interesting part of the crystalline structure of complex [Cr(L<sup>3</sup>

*Perspective view of complex 7 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

*Perspective view of complex 6 with atom numbering scheme (hydrogen atoms are omitted for clarity).*

*One-dimensional network through C–HO interactions in complexes (6) and (7).*

molecular chain along the *c*-axis through C–HO interactions between each mole-

The 1D chains are arranged in parallel direction to form a supramolecular host having channels along the *c*-axis (**Figure 25**). Such channels are filled by selfassembled "water-chloride" clusters having chair conformation (**Figure 25**).

It is proven by the crystal structure analysis that there are three crystals of water molecules per formula unit of **complexes 6** and **7** along with chloride ion, and these are involved in forming the infinite "water-chloride" cluster [36] with a chair

)2]Cl3H2O (**7**) is that both the complexes form a 1D supra-

)2]

partial double bond character [36].

*Stability and Applications of Coordination Compounds*

Cl3H2O (**6**) and [Fe(L<sup>3</sup>

cule (**Figure 24**).

**Figure 23.**

**Figure 22.**

**Figure 24.**

**12**

*Supramolecular channels and "water-chloride" cluster along the c-axis in complexes (6) and (7).*

#### **Figure 26.** *Chair conformation formed by "water-chloride" cluster in complexes (6) and (7).*

conformation. The three water molecules and a chloride molecule form a fourmembered cyclic motif, and such three adjacent four-membered rings form the chair conformation (**Figure 26**).

#### *4.3.5.1 Magnetic property and Mössbauer spectroscopy of complex (7)*

The variable temperature (2.5–300 K) magnetic moment study shows the temperature dependence of the magnetic susceptibility. The χ<sup>m</sup> values at 2.5 and 300 K are 0.33 and 0.005 cm <sup>3</sup> mol<sup>1</sup> , while the μeff values are 2.61 and 3.46 B.M., respectively. The detailed study shows that the magnetic moment value consists of a superimposition of both the low-spin and high-spin states. At very low temperature, the 1-D supramolecular species which is formed by strong intermolecular C– HO interactions and the cooperative interactions with the "water-chloride" cluster between mononuclear spin crossover (SCO) sites stabilize the low-spin state, and thus the high-spin contribution decreases to 21%, and the low-spin contribution increases to 79%. Thus, such variable temperature magnetic behavior may be due to a continuous S = 1/2 to 5/2 spin crossover phenomenon of iron centers [36].

The Mössbauer spectroscopic study also supports that a spin crossover phenomenon exists in the iron(III) complex (**7**). Both the spin states, low (S = 1/2) and high (S = 5/2), exist in room temperature (300 K) as well as in very low temperature at 20 K. The population density of the electrons decreases from high-spin state to the low-spin state with decreasing temperature (**Table 1**).


the current in ampere, Vc the potential drop across the sample of cross-sectional

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands…*

with temperature can be obtained by the Arrhenius equation:

Variation of electrical conductivity of a compound behaving like semiconductor

where σ is the electrical conductivity, σ<sup>o</sup> denotes the pre-exponential factor, *Ea* is the activation energy for this thermally activated process, and *k* is the Boltzmann constant. Generally, "log σ" is plotted against "1000/T" which is expected to have a linearly fitted graph. The thermal energy of activation (*Ea*) is determined from the

If the graph obtained is linear (i.e., fitted with one straight line), then it may be concluded that no molecular rearrangement occurred during heating and the compound will have only one *Ea*. In this case the electrical conduction is mainly due to the intrinsic conducting property, whereas the nonlinear plot clearly indicates molecular rearrangement during heating of the compound. In case of nonlinear plot, more than one straight line are fitted against the linear parts of the plot. One straight line corresponding to lower temperature range is known as region I, and the other straight line corresponding to the higher temperature range is known as region II. In this case

Conduction corresponding to the region I is attributed to the intermolecular conduction via weak force interactions between the molecules. The charge carriers hop near Fermi level within the localized state. Delocalized π-electrons are mainly responsible for this conduction, whereas conduction corresponding to the region II is attributed to intramolecular conduction between the metal center and the ligand center within a metal complex. This conduction occurs due to tunneling of electrons between equivalent HOMO and LUMO of the ligand and metal ion, respectively. Such tunneling of electrons through the intermolecular potential barrier is

reinforced through π–π stacking and extensive H-bonding [24]. Depending on the availability of π-electrons, the compound behaves like n-type semiconductor.

It is also very clear from the Arrhenius plots that the conductivity of metal complexes generally increases with increase in temperature. At room temperature

*Temperature dependence. Electrical conductivity curves of the complexes [24, 28, 29].*

From the Arrhenius plots (**Figure 27**), the electronic parameters, i.e., activation energy of electrical conduction (Ea) and the energy gap for directly allowed transitions of metal complexes (**2**) and (**8–11**), are calculated, and the results obtained are

the compound will have two distinct energy of activations (*Ea1* and *Ea2*).

σ ¼ σ<sup>o</sup> exp ð Þ �*Ea=k*T (1)

area (a), and is the thickness (d).

*DOI: http://dx.doi.org/10.5772/intechopen.90171*

slope of the graph.

summarized in **Table 3**.

**Figure 27.**

**15**

**Table 1.**

*Population density at variable temperature in [Fe(L<sup>3</sup> )2]Cl3H2O (7).*


#### **Table 2.** *Powder X-ray diffraction data.*

#### *4.3.6 PXRD, SEM, and EDX studies of complexes (8)–(11)*

Suitable single crystals of complexes (**8**)–(**11**) could not be grown even after repeated efforts; hence they were characterized by the powder X-ray diffraction study of the compounds [28, 29]. Some of the important lattice parameters of the PXRD study are summarized in **Table 2**.

The SEM investigation of all the above complexes, the ground powders, and the fracture surfaces indicates that the grain size distribution is not uniform, and submicron grains (finely ground powder) as well as grains (fracture surfaces) even above 20 *μ*m (for complex 8) and above 10 *μ*m (for complexes 9–11) have been observed [37, 38].

The formation of metal–ligand complexes and the presence of metal along with C and S within the metal complexes have been substantiated by the EDX analysis.

#### **4.4 Electrical conductivity**

To explore the utility of the metal complexes as functional materials, the electrical conductivity study was performed, and it shows the semiconducting nature of the complexes [33, 37, 38].

The samples for the measurement of electrical conductivity were prepared from the complexes in the form of tablets of approximately thickness 0.1 cm at a pressure of ca. 1 <sup>10</sup><sup>8</sup> Pascal. These tablets were placed between two copper electrodes covered with silver paste, and contacts of the prepared tablets were to be Ohmic or not. A two-probe method was used to investigate the electrical conductivities of the complex tablets by measuring the current through the probes with a high impedance electrometer (Keithley 6514) upon application of a DC voltage current supplied by a programmable source of voltage (Keithley 230). The conductivities were calculated by using the general equation of σ = (I/Vc)(d/a), where (I) is *Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands… DOI: http://dx.doi.org/10.5772/intechopen.90171*

the current in ampere, Vc the potential drop across the sample of cross-sectional area (a), and is the thickness (d).

Variation of electrical conductivity of a compound behaving like semiconductor with temperature can be obtained by the Arrhenius equation:

$$
\sigma = \sigma\_\circ \, \exp\left(-E\_a/kT\right) \tag{1}
$$

where σ is the electrical conductivity, σ<sup>o</sup> denotes the pre-exponential factor, *Ea* is the activation energy for this thermally activated process, and *k* is the Boltzmann constant. Generally, "log σ" is plotted against "1000/T" which is expected to have a linearly fitted graph. The thermal energy of activation (*Ea*) is determined from the slope of the graph.

If the graph obtained is linear (i.e., fitted with one straight line), then it may be concluded that no molecular rearrangement occurred during heating and the compound will have only one *Ea*. In this case the electrical conduction is mainly due to the intrinsic conducting property, whereas the nonlinear plot clearly indicates molecular rearrangement during heating of the compound. In case of nonlinear plot, more than one straight line are fitted against the linear parts of the plot. One straight line corresponding to lower temperature range is known as region I, and the other straight line corresponding to the higher temperature range is known as region II. In this case the compound will have two distinct energy of activations (*Ea1* and *Ea2*).

Conduction corresponding to the region I is attributed to the intermolecular conduction via weak force interactions between the molecules. The charge carriers hop near Fermi level within the localized state. Delocalized π-electrons are mainly responsible for this conduction, whereas conduction corresponding to the region II is attributed to intramolecular conduction between the metal center and the ligand center within a metal complex. This conduction occurs due to tunneling of electrons between equivalent HOMO and LUMO of the ligand and metal ion, respectively. Such tunneling of electrons through the intermolecular potential barrier is reinforced through π–π stacking and extensive H-bonding [24]. Depending on the availability of π-electrons, the compound behaves like n-type semiconductor.

From the Arrhenius plots (**Figure 27**), the electronic parameters, i.e., activation energy of electrical conduction (Ea) and the energy gap for directly allowed transitions of metal complexes (**2**) and (**8–11**), are calculated, and the results obtained are summarized in **Table 3**.

It is also very clear from the Arrhenius plots that the conductivity of metal complexes generally increases with increase in temperature. At room temperature

**Figure 27.** *Temperature dependence. Electrical conductivity curves of the complexes [24, 28, 29].*

*4.3.6 PXRD, SEM, and EDX studies of complexes (8)–(11)*

PXRD study are summarized in **Table 2**.

*Population density at variable temperature in [Fe(L<sup>3</sup>*

**)2]. 2H2O (8)**

*Stability and Applications of Coordination Compounds*

**[Ni(L<sup>3</sup>**

**[Zn(L<sup>3</sup>**

observed [37, 38].

**Table 1.**

**Cell parameters**

V (Å<sup>3</sup>

**Table 2.**

*Powder X-ray diffraction data.*

**4.4 Electrical conductivity**

the complexes [33, 37, 38].

**14**

Suitable single crystals of complexes (**8**)–(**11**) could not be grown even after repeated efforts; hence they were characterized by the powder X-ray diffraction study of the compounds [28, 29]. Some of the important lattice parameters of the

**Temperature Spin state Occupancy** 300 K Low spin (1/2) 49%

20 K Low spin (1/2) 77%

**)(OAc)] (9)**

System Triclinic Monoclinic Monoclinic Monoclinic

) 1294.88 810.3 1127.7 1044.17 a (Å) 10.297368 6.364172 19.600876 18.953438 b (Å) 11.32531 27.497931 5.53422 6.365518 c (Å) 12.345947 4.686936 12.32786 8.729238 α 111.516869 90 90 90 β 103.288712 98.92 122.51 97.5 γ 91.155464 90 90 90

High spin (5/2) 51%

High spin (5/2) 23%

**)(OAc)]. H2O (10)**

**[Cu(L<sup>3</sup>**

**)(OAc)]. H2O (11)**

*)2]Cl3H2O (7).*

**[Co(L<sup>3</sup>**

The SEM investigation of all the above complexes, the ground powders, and the

The formation of metal–ligand complexes and the presence of metal along with C and S within the metal complexes have been substantiated by the EDX analysis.

To explore the utility of the metal complexes as functional materials, the electrical conductivity study was performed, and it shows the semiconducting nature of

The samples for the measurement of electrical conductivity were prepared from

the complexes in the form of tablets of approximately thickness 0.1 cm at a pressure of ca. 1 <sup>10</sup><sup>8</sup> Pascal. These tablets were placed between two copper electrodes covered with silver paste, and contacts of the prepared tablets were to be Ohmic or not. A two-probe method was used to investigate the electrical conductivities of the complex tablets by measuring the current through the probes with a high impedance electrometer (Keithley 6514) upon application of a DC voltage current supplied by a programmable source of voltage (Keithley 230). The conductivities were calculated by using the general equation of σ = (I/Vc)(d/a), where (I) is

fracture surfaces indicates that the grain size distribution is not uniform, and submicron grains (finely ground powder) as well as grains (fracture surfaces) even above 20 *μ*m (for complex 8) and above 10 *μ*m (for complexes 9–11) have been


to the direct energy gap (Egd). The Egd values of the metal complexes are collected in **Table 3**. The comparison of the band gap cannot be directly related to the atomic

*Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands…*

In this review, the synthesis, crystal structure, and solid-state properties of three Schiff base ligands derived from diacetylmonoxime with diethylenetriamine, 1,3 diaminopropane-2-ol, and morpholine N-thiohydrazide and their metal complexes have been vividly discussed. A zwitterionic nitrogen–sulfur heterocyclic compound with nonbonded SS interaction has also been reported to be formed by the reaction of diacetylmonoxime with morpholine N-thiohydrazide under long refluxing (16 h) condition in ethanol. The single X-ray crystal structures have shown many beautiful weak force interactions including a CH3O trifurcated interface communication. Wherever the single-crystal structures could not be grown, the PXRD study has enlightened their structural features. The electrical and optical properties also explored the semiconducting nature of some of the metal complexes. It is also observed that the electron transport process gets influenced by the supramolecular

One of the authors (S.S.) is thankful to the UGC (ERO), Kolkata, for financial grants (MRP) to carry out a part of this work and also to Prof. Y. Aydogdu, Department of Physics, Gazi University, and Dr. S. Biswas, our lab-mate for some useful

1 Department of Chemistry, University of Kalyani, Kalyani, West Bengal, India

3 Department of Chemistry, Kalyani Government Engineering College, Kalyani,

4 Department of Chemistry, Chakdaha College, Chakdaha, West Bengal, India

© 2020 The Author(s). Licensee IntechOpen. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/ by/3.0), which permits unrestricted use, distribution, and reproduction in any medium,

2 Lalgopal High (H.S.) School, Ranaghat, West Bengal, India

\*Address all correspondence to: saikat\_s@rediffmail.com

† Prof. Kamalendu Dey expired on September 01, 2019.

provided the original work is properly cited.

, Kamalendu Dey1† and Saikat Sarkar<sup>4</sup>

\*

number of the metals in the complexes.

*DOI: http://dx.doi.org/10.5772/intechopen.90171*

frameworks of the metal complexes.

**Acknowledgements**

discussion.

**Author details**

West Bengal, India

**17**

Palash Mandal1,2, Uttam Das3

**5. Conclusion**

#### **Table 3.**

*Activation energies and direct band gap values.*

they behave as an insulator, while at higher temperature the semiconducting nature of complexes is observed.

#### **4.5 Optical properties**

Optical absorption spectra was taken by using a UV–VIS spectrophotometer (Perkin Elmer Lambda 2S/45 Double Beam) and measured as function of wavelength in the wavelength range 190–1100 nm.

The energy band gaps and the nature of the optical transitions involved in the metal complex framework systems have been practically determined by the fundamental absorption edge analysis of the recorded optical transitions using the theory of Mott and Davis [39]. It is also observed that the semiconducting behavior of a material increases with rise in temperature which may also damage the actual molecular structure of the material. Hence, Tauc method is used to calculate the energy band gap through optical absorption properties [40].

Utilizing the relation between the optical linear absorption coefficient (α) with photon energy (hν), the energy band gap (Eg) between the top of the valence band and bottom of the conduction band can be determined using equation (Eq. (2)):

$$\mathbf{a}\mathbf{h}\nu = \mathbf{A}\left(\mathbf{h}\nu - \mathbf{E}\_{\mathfrak{g}}\right)^{\mathfrak{n}} \tag{2}$$

where A is a constant characteristic parameter of the respective transition independent of ν.

The values of n depend on the kind of optical transitions. For directly allowed, directly forbidden, indirectly allowed, and indirectly forbidden transitions, the values of n are ½, 3/2, 2, and 3, respectively. Thus the energy band gap for directly allowed (Egd) and indirectly allowed (Egi) transitions can be determined by relating Eq. (2) as follows:

$$\mathbf{a}\mathbf{h}\nu = \mathbf{A}\_{\mathrm{d}} \left(\mathbf{h}\nu - \mathbf{E}\_{\mathrm{gd}}\right)^{1/2} \tag{3}$$

and

$$\alpha \mathbf{h} \nu = \mathbf{A}\_{\mathrm{i}} \left( \mathbf{h} \nu - \mathbf{E}\_{\mathrm{gi}} \right)^{2} \tag{4}$$

where Egd and Egi are direct and indirect energy gaps, respectively.

To calculate the direct and indirect energy band gap, we need to plot a curve of (αhν) <sup>2</sup> against f(hν) and (αhν) 1/2 against f(hν) and then by the extrapolation of the most linear part of the curve to zero.

The satisfactory graphs were obtained for the metal complexes (**2**) and (**8**)–(**11**) by plotting (αhν) <sup>2</sup> against f(hν). Therefore, the energy gaps determined correspond *Crystal Structure and Solid-State Properties of Metal Complexes of the Schiff Base Ligands… DOI: http://dx.doi.org/10.5772/intechopen.90171*

to the direct energy gap (Egd). The Egd values of the metal complexes are collected in **Table 3**. The comparison of the band gap cannot be directly related to the atomic number of the metals in the complexes.
